Texas STAAR Review & Practice Chemistry - Pearson

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Feb 14, 2012 ... Practice Test A Answer Sheet . ... Practice Test B Answer Sheet . ..... Which of the tests would involve a chemical change in the mineral? A Tests ...

Texas STAAR Review & Practice Chemistry

TEKS

Boston, Massachusetts • Chandler, Arizona • Glenview, Illinois • Upper Saddle River, New Jersey

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Content Reviewer C. Alton Hassell, Ph.D. Director of Undergraduate Programs, Chemistry and Biochemistry Baylor University Waco, Texas

Copyright © 2012 Pearson Education, Inc., or its affiliates. All Rights Reserved. Printed in the United States of America. This publication is protected by copyright, and permission should be obtained from the publisher prior to any prohibited reproduction, storage in a retrieval system, or transmission in any form or by any means, electronic, mechanical, photocopying, recording, or likewise. For information regarding permissions, write to Rights Management & Contracts, Pearson Education, Inc., One Lake Street, Upper Saddle River, New Jersey 07458.

ISBN-13: 978-0-13-319135-6 ISBN-10: 0-13-319135-4

1 2 3 4 5 6 7 8 9 10 V088 15 14 13 12

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TEKS REVIEW

Contents About this Book . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . v Texas Essential Knowledge and Skills (TEKS) for Chemistry . . . . . . . . . . . . . . . . . vii Lesson Reviews TEKS 1A Demonstrating Safe Practices . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1 TEKS 1B

Hazards of Chemical Substances . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4

TEKS 1C

Conservation of Resources and Proper Disposal . . . . . . . . . . . . . . . . . . . . 7

TEKS 2A Definition of Science . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10 TEKS 2B

Scientific Hypotheses . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13

TEKS 2C

Scientific Theories . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16

TEKS 2D Scientific Hypotheses and Scientific Theories . . . . . . . . . . . . . . . . . . . . . 19 TEKS 2E

Planning and Implementing Investigative Procedures . . . . . . . . . . . . . . 22

TEKS 2F

Accuracy and Precision . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 25

TEKS 2G Dimensional Analysis, Scientific Notation, and Significant Figures . . . . 28 TEKS 2H Data Analysis, Inferences, and Predictions . . . . . . . . . . . . . . . . . . . . . . . 31 TEKS 2I

Communicating Conclusions Based on Scientific Data . . . . . . . . . . . . . . 34

TEKS 3A Analyzing, Evaluating, and Critiquing Scientific Explanations . . . . . . . 37 TEKS 3B

Communicating and Applying Scientific Information . . . . . . . . . . . . . . 40

TEKS 3C

Drawing Inferences from Promotional Materials . . . . . . . . . . . . . . . . . . 43

TEKS 3D The Impact of Scientific Research . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 46 TEKS 3E

Chemistry and Careers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 49

TEKS 3F

The History of Chemistry and Contributions of Scientists . . . . . . . . . . . 52

TEKS 4A Physical and Chemical Changes and Properties . . . . . . . . . . . . . . . . . . . 55 TEKS 4B

Intensive and Extensive Properties . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 58

TEKS 4C

States of Matter . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 61

TEKS 4D Classifying Matter . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 64 TEKS 5A Development of the Periodic Table . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 67 TEKS 5B

Chemical Families in the Periodic Table . . . . . . . . . . . . . . . . . . . . . . . . . 70

TEKS 5C

Trends in the Periodic Table . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 73

TEKS 6A Atomic Theory . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 76 TEKS 6B

Electromagnetic Waves . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 79

TEKS 6C

Calculating Wavelength, Frequency, and Energy of Light . . . . . . . . . . . 82

TEKS 6D Calculating Average Atomic Mass . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 85 TEKS 6E

Electron Configurations and Lewis Valence Electron Dot Structures . . 88

TEKS 7A Naming Compounds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 91 TEKS 7B

Writing Chemical Formulas . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 94

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TEKS 7C

Constructing Electron Dot Formulas . . . . . . . . . . . . . . . . . . . . . . . . . . . . 97

TEKS 7D The Nature of Metallic Bonding . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 100 TEKS 7E

Molecular Structures . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 103

TEKS 8A Defining and Using the Mole . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 106 TEKS 8B

Calculating Atoms, Ions, or Molecules Using Moles . . . . . . . . . . . . . . . 109

TEKS 8C

Calculating Percent Composition, Empirical Formulas, and Molecular Formulas . . . . . . . . . . . . . . . . . . . . 112

TEKS 8D Balancing Chemical Equations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 115 TEKS 8E

Stoichiometric Calculations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 118

TEKS 9A Gas Laws . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 121 TEKS 9B

Gas Stoichiometry . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 124

TEKS 9C

Kinetic Molecular Theory . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 127

TEKS 10A The Importance of Water . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 130 TEKS 10B Solubility Rules . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 133 TEKS 10C Calculations Involving Molarity . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 136 TEKS 10D Calculating Dilutions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 139 TEKS 10E Types of Solutions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 142 TEKS 10F Factors Influencing Solubility . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 145 TEKS 10G Acids and Bases . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 148 TEKS 10H Types of Reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 151 TEKS 10I pH of a Solution . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 154 TEKS 10J Degrees of Dissociation for Acids and Bases . . . . . . . . . . . . . . . . . . . . . 157 TEKS 11A Energy and Its Forms . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 160 TEKS 11B Conservation of Energy and Heat Transfer . . . . . . . . . . . . . . . . . . . . . . 163 TEKS 11C Energy Changes in Chemical Reactions . . . . . . . . . . . . . . . . . . . . . . . . . 166 TEKS 11D Heat, Mass, Temperature Change, and Specific Heat . . . . . . . . . . . . . 169 TEKS 11E Calorimetry . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 172 TEKS 12A Alpha, Beta, and Gamma Radiation . . . . . . . . . . . . . . . . . . . . . . . . . . . 175 TEKS 12B Describing Radioactive Decay Using Nuclear Equations . . . . . . . . . . . . 178 TEKS 12C Fission and Fusion Reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 181 Answers to End-of-Course-Assessment Reviews . . . . . . . . . . . . . . . . . . . . . . . . . . 185 Glossary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 197 Test-Taking Tips . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 201 Chemistry Reference Materials . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 209 Practice Test A . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 215 Practice Test A Answer Sheet . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 237 Practice Test B . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 239 Practice Test B Answer Sheet . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 261

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About this Book This review and practice book focuses on the basic content that may be tested on the end-of-course State of Texas Assessments of Academic Readiness (STAAR) in Chemistry. Each Texas Essential Knowledge and Skills (TEKS) is reviewed in sequence. You can use the book in any order, as each TEKS review is independent. A

TEKS REVIEW

5C

Each three-page lesson begins with the TEKS. Readiness standards, which have been identified as those most important for in-depth understanding of a particular topic, are clearly indicated. The content summaries specifically address the concepts called out in the TEKS. The lessons include numerous illustrations to help you visualize and understand the concepts and vocabulary of chemistry. You should carefully review the illustrations as well as the explanations within the text.

A

Vocabulary atomic radius ion cation anion electronegativity ionization energy

What periodic trend s in atomic radii can identified in the perio be dic table

? The size of an atom is expressed as an atomic radius (plural, radii). The atomic radius is one half the distance between the nuclei of two atoms of the same element when the atoms are joined. In general, atomic radii decrease from left to right across a period on the periodic table and increase from top to bottom within a group, or family. a period in the periodic table, the number of protons As you move across increases, but the electron in s remain in the same energy the atoms outer-level electrons are level. Therefore, pulled more strongly toward the nucleus from to right across a period. left This increasingly stronge r pull results in a smaller radius from left to right across a period. The principal quantum number, n, of the outer-le vel electrons increases one from period to period. by For example, for element For elements in period s in period 1, n = 1. 2, n = 2, and so on. As n increases down a family, the outer-level electron s have an average position that is farther from the nucleus. As a result, the atoms are larger.

B

What periodic trend s in ionic radii can identified in the perio be dic table

? An ion is an atom or group of atoms that has a positive There are two types of or negative charge. ions—cations and anions. Cations are atoms that have lost one or more electrons and thus have a positive charge. Atoms that lose electrons become smaller. For example, the calcium ion, Ca2+, smaller than a calcium is atom, Ca, because Ca2+ has two fewer electron s. Anions are atoms that have gained one or more electrons and thus have negative charge. Atoms a that gain electrons are bigger. For example, a bromide ion, Br–, is larger than a bromine atom, Br, because Br– has one more electron. Cations and anions exhibit trends that are similar to those of their parent atoms across periods and down families. From period, the radii of cations left and anions decrease because to right across a protons in the nucleus the number of increases. From top to bottom within a family, radii of cations and anions the increase because the principa number, n, increases. l quantum

B

You will need to know the definitions of the vocabulary words listed at the beginning of each lesson in order to answer many of the end-of-course assessment questions. These words are shown in bold type within the topic where they are first defined. Each bold word is accompanied by a simple definition in the text. Vocabulary words are also defined in the glossary at the end of the book. Words that are in italicized type are words that you need to know to understand basic chemistry concepts. Although you are not likely to be tested on the specific definitions of non-vocabulary words, these words may be used in end-of-course assessment questions.

Study Tip

Use the Periodic Table to identify and explain TEKS_TX T periodic trends, includin radii, electron egativity, and ionizatio g atomic and ionic n energy.

TEKS 5C • Copyright ©

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or its affiliates. All Rights

Reserved.

What periodic trends in electronegativity can be identified in the periodic table?

Study Tip Draw a rough outline of the periodic table leaving room for labels above and to the left of the table. For each trend that you study, draw a line above the table and on the left side, and use arrowheads to indicate the direction in which the trend increases. Label each arrow.

Electronegativity is the ability of an atom to attract electrons when the 7/1/11 1:12 atom is in a compound. The greater an atom’s electronegativity, the PM greater its ability to attract electrons. The number of protons in the nucleus and the principal quantum number influence the periodic trends for electronegativity. Generally, the electronegativity increases from left to right across a period of the periodic table because the number of protons in the nucleus increases. Electronegativity generally decreases from top to bottom within a family because outer energy level electrons are farther from the nucleus.

C

Ionization energy is the minimum energy required to remove an electron from an atom or ion. The energy required to remove the first electron from an atom is referred to as the first ionization energy. The greater an element’s ionization energy, the more difficult it is to remove an electron. Generally, ionization energy increases from left to right across a period and decreases from top to bottom within a family.

What periodic trends in ionization energy can be identified in the periodic table?

Ionization energy depends on the force of attraction the nucleus exerts on the electron. As with the other periodic trends, this attraction depends upon the number of protons in the nucleus and the distance of the electron from the nucleus. More protons exert more force, making electrons harder to remove. Therefore, from left to right across a period, ionization energy increases because the number of protons in the nucleus increases. Electrons that are closer to the nucleus are pulled more strongly toward the nucleus, making them harder to remove. Therefore, as atomic radii increase from top to bottom within a family, electrons that are farther from the nucleus are easier to remove. These trends can be seen in Figure 1 below.

C

Figure 1 Trends in Ionization Energy

Each lesson contains a study tip that either refreshes your memory of relevant information that may have been covered in previous science courses or helps reinforce understanding of a particular concept.

First ionization energy (kJ/mol)

Vocabulary

Trends in the Per iodic Table TEKS 5C REadinE SS

TEKS

First Ionization Energy vs. Atomic Number

2500

He

Ne

2000 1500 1000 500 0

Ar

N H

Kr Xe

P

Be

Zn As

Mg Li

Na

10

Rb

K

20

Cd

30

40

Atomic number

Cs

50

60

TEKS 5C • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

74

End-of-Course Assessment Review Questions D

TEKS

End-of-Course Assessment Review STR12_ANC_CHEM_05C.indd 74

7/1/11 1:12 PM

1. Identify Which of the following trends can be identified on the periodic table? A Atomic radii increase from left to right across a period.

Following each content summary is a set of questions that will help you clarify and reinforce your understanding of the content. Answering these questions will help gauge your understanding of a particular TEKS. Answers and explanations can be found at the end of the book.

D

B Ionization energy increases from top to bottom within a family. C Electronegativity decreases from left to right across a period. D Ionic radii of cations decreases from left to right across a period.

2. Explain Which of the following correctly explains why the sizes of atoms decrease from left to right across a period? A The principal quantum number increases. B The number of electrons increases. C The distance from the nucleus increases. D The number of protons increases. 3. Explain An increase in principal quantum number explains which of the following trends? A The ionization energy decreases from top to bottom within a family. B The ionization energy increases from left to right across a period. C The electronegativity increases from left to right across a period. D The atomic radius increases from left to right across a period.

4. Apply Concepts Use the table below to determine which of the Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

v

following relationships is correct.

Selected Trends in the Periodic Table Element Li C Rb Be

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Period 2 2 5 2

Family 1A 4A 1A 2A

A Li has a smaller atomic radii than C. B Li+ has a larger atomic radii than Li.

2/14/12 9:54 AM

6C

Frequency, and Energy of Light TEKS 6C Calculate the wavelength, frequency, and energy of light using Planck’s constant and TEKS_TXT the speed of light.

How can you calculate the wavelength of light using its frequency and the speed of light?

Study Tip Keep track of units as you do your calculations. Knowing the units of the value that you are calculating can help you determine which equation to use.

Sample Problems

E

Sample Problem 1

Multiplying the numerator and denominator by s leaves only the unit of m. Divide 3.00 × 108 m by 4.0 × 1014.

Visible light has frequencies between about 4.0 × 1014 hertz (Hz) and about 7.9 × 1014 Hz. What are the wavelengths of the lowest frequencies of visible light?

λ=

First, substitute the lowest frequency, 4.0 × 1014 Hz, into the equation: c λ= 4.0 × 1014 Hz

λ=

4.0 × 1014

λ = 0.75 × 10–6 m or 7.5 × 10 –7 m The longest wavelengths of visible light are about 7.5 × 10 –7 m.

3.00 × 108 m/s 4.0 × 1014 Hz

F

Notice that, in this form, the units do not cancel. Recall that 1 Hz equals 1 s−1. Substitute 1 s−1 for the unit Hz:

Interpreting Diagrams

3.00 × 10 m/s 8

The correct answer is A. To identify the correct answer, you need to apply some basic facts about atoms. As choice A and the diagram suggest, protons and neutrons are located in the atomic nucleus. Electrons move in the space around the nucleus. Choices B, C, and D are all features that the diagram does not represent accurately about atoms. Remember that all models are inaccurate in at least some ways.

Scientists use diagrams to show the parts of objects or the steps of a process. Begin by looking carefully at the diagram to get an overall sense of what it shows. Look for any labels with lines connected to parts of the diagram. If you don’t immediately recognize what the diagram shows, the labels may you a clue. have TEKS 6C • Copyright © Pearson Education, Inc., or give its affiliates. All After Rightsyou Reserved. examined the diagram, read the question. After you 82 an answer, recheck the diagram to make choose sure that your answer is correct.

4.0 × 1014 s−1

Sample Question 2

Sometimes you will be asked to interpret a set of related diagrams. Each part of the diagram could represent a test group of a science experiment or demonstration. Or the diagrams could show the steps of a process, such as a chemical reaction.

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Four identical bottles of carbonated water are each placed in a glass bowl. The conditions of the bowl and the bottle are shown in the set of diagrams below.

B

Bottle 2 No cap

No cap

Ice

Warm water Bottle 4

Bottle 3

Proton

REVIEW

Warm water

Bottle 1

In the diagram below, which feature of protons, neutrons, and electrons is represented most accurately?

A

Cap

Ice

Sample Question 1

TEKS

7/25/11 8:03 AM

Cap

To interpret a set of diagrams, make sure that you understand the meaning of arrows and other symbols. You should also compare the different diagrams. Determine the features that are similar and different. Then read the question and answer choices carefully.

Test-taking tips provide strategies for answering multiple-choice questions. Accompanying sample questions allow you to practice your test-taking skills.

Neutron

Which bottle will contain the least amount of dissolved carbon dioxide after 15 minutes?

Electron

A Bottle 1

C Bottle 3

B Bottle 2

D Bottle 4

The correct answer is D. Bottle 4 is uncapped, so the carbon dioxide in Bottle 4 is under less pressure than position inside or outside of the atomic nucleus the carbon dioxide in Bottles 1 and 2. Because the distance between them solubility of a gas decreases as the pressure decreases, Bottle 4 will contain less carbon dioxide markings on their surfaces than Bottles 1 and 2. The warm water surrounding relative sizes State of Texas Bottle 4 also decreases the solubility of carbon Assessments of dioxide. Bottle 3 is uncapped like Bottle 4, but is Academic Readiness sitting in ice.

C STAAR CHEMISTRY Chemistry Reference D REFERENCE MATERIALS

Reference Materials

3.00 × 108 m

Recall that the exponent of the denominator, 14, is subtracted from the exponent of the numerator, 8.

Next, substitute the speed of light for c:

λ=

F

c = λν Because the product of wavelength and frequency is equal to a constant, you can always calculate one of these variables if you know the value of the other. For example, if you know the frequency of a wave, you can calculate its wavelength by dividing both sides of the equation by frequency. The result is: c λ= ν

E

Numerous solved sample problems appear throughout this book to provide you with examples of typical problems found on the end-of-course assessment. The step-bystep detailed solutions guide you in the problem-solving process and help reinforce content knowledge.

Test-Taking Tips

In a vacuum, all electromagnetic waves, including light waves, travel at the speed of light. The speed of light, c, is a constant value of 300,000,000 meters per second (3.00 × 108 m/s). The product of the wavelength, λ, and frequency, ν, of a wave equals the speed of light, c.

GAR Materials STA

TM

ATOMIC STRUCTURE STRUCTURE ATOMIC

G

Speed of light = (frequency)(wavelength)

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205

Energy = (Planck’s constant)(frequency)

Chemistry reference materials include formulas, constants, conversions, the periodic table, and other information that may be needed to answer questions on the exam. You will get a copy of these reference materials when you take your end-of-course test.

Energy =

E photon = hf

(Planck’s constant)(speed of light) (wavelength)

STR12_ANC_CHEM_EM.indd 205

BEHAVIOROF OFGASES GASES BEHAVIOR Total pressure of a gas =

(

E photon =

sum of the partial pressures of the component gases

)

(Initial pressure)(initial volume)

(Initial pressure)(initial volume) (Initial volume) (Initial temperature) (Initial volume) (Initial moles)

=

=

PV = nRT PV 1 1

(final pressure)(final volume)

n1T1

(final moles)(final temperature)

= (final pressure)(final volume)

=

P2V2

n2T2

PV = P V 1 1

V1

(final vo l lume)

=

1/18/12 2:53 PM

PT = P1 + P2 + P3 + . . .

(Pressure)(volume) = (moles)(ideal gas constant)(temperature)

s )(initial temperature) (Initial moles

hc λ

T1

(final temperature) (final volume)

V1

n1

(final moles)

2 2

=

=

V2 T2

V2

n2

SOLUTIONS SOLUTIONS Molarity =

moles of solute liter of solution

M =

Ionization constant of water =

(

STAAR Chemistry

Practice Tests

TEKS

H

Volume of solution 1

)(

) (

molarity of solution 1 =

(

hydrogen ion concentration

volume of solution 2

)(

)(

)

hydroxide ion concentration Practice Test A

molarity of solution 2

ent e Asse pH = −logarithm ionssm concentration) ours(hydrogen End-of-C A THERMOCHEMISTRY istry Practice Test ChemTHERMOCHEMISTRY

( )( ) ( TEKS )

Two sample full-length exams mimic the STAAR test in both format and layout. Following each test is an answer sheet for you to fill in the correct answers.

K w = [H+][OH−]

)

V1M1 = V2M2 pH = −log[H+]

STAAR specific Chemistrychange in

1

mol L

)

Practice Test B

Q = mcp∆T Heat gained or lost = (mass) its heat temperature mineral sample from ine the identity of a following tests to determ A student performs the Enthalpy enthalpy enthalpy ies. of − of reactants propert ∆H = ∆H of (products) − ∆H of (reactants) reaction = of products physical and chemical its crystals l to see the shape of examination of the minera Test 1 Microscopic 209 ine its hardness of the mineral to determ Test 2 Scratch test of its to determine the color mineral on a streak plate the g Rubbin 3 Test 1 A student performs the following tests to determine the identity of a mineral sample. powdered form s will form mineral to see if bubble STR12_ANC_CHEM_EM_RM.indd 209 loric acid on the 2/13/12 1:28 PM hydroch dilute 1. The mineral is heated in chlorine gas to see if it will produce a different substance. Test 4 Dropping

(

Which of the tests would A Tests 2 and 4 B Tests 3 and 4

End-of-Course Assessment Chemistry Practice Test B

H

mineral? themineral 2. inThe is immersed in an ammonia solution to see if any new color is produced. involve a chemical change 3. The mineral is rubbed on a streak plate to determine the color of its powdered form.

4. Dilute hydrochloric acid is placed on the mineral to see if bubbles will form.

C Tests 1, 2, and 4 D Test 4 only

Which of the above tests were for physical properties of the mineral? A Test 1 B Test 2 C Test 3

2

D Test 4 of a small sample tions and measurements makes careful observa A chemistry student ines the following: of matter, and determ Appearance

silver solid

Density

11.85 g 2 A sample of sulfur obtained from the crater of a volcano is carefully measured. Which of the 3 5.9 g/cm following is an intensive property of the sample?

Melting point

30ºC A Density of 2.07 g/cm3

Mass

B Mass of 3.85 g of the following ce is gallium (Ga). Which that the unknown substan 3 The student determines ?C Volume of 1.86 cm y of the gallium sample is an extensive propert D Temperature of 20°C A Silver solid B Mass of 11.85 g 3 C Density of 5.9 g/cm 30°C D Melting point of Copyright © Pearson Education,

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Texas Essential Skills and Knowledge (TEKS) for Chemistry (1) Scientific processes. The student, for at least 40% of instructional time, conducts laboratory and field investigations using safe, environmentally appropriate, and ethical practices. The student is expected to: (A) Demonstrate safe practices during laboratory and field investigations, including the appropriate use of safety showers, eyewash fountains, safety goggles, and fire extinguishers. (B) Know specific hazards of chemical substances such as flammability, corrosiveness, and radioactivity as summarized on the Material Safety Data Sheets (MSDS). (C) Demonstrate an understanding of the use and conservation of resources and the proper disposal or recycling of materials.

(2) Scientific processes. The student uses scientific methods to solve investigative questions. The student is expected to: (A) Know the definition of science and understand that it has limitations. (B) Know that scientific hypotheses are tentative and testable statements that must be capable of being supported or not supported by observational evidence. Hypotheses of durable explanatory power which have been tested over a wide variety of conditions are incorporated into theories. (C) Know that scientific theories are based on natural and physical phenomena and are capable of being tested by multiple independent researchers. Unlike hypotheses, scientific theories are well-established and highly-reliable explanations, but may be subject to change as new areas of science and new technologies are developed. (D) Distinguish between scientific hypotheses and scientific theories.

es (E) Plan and implement investigative procedures, including asking questions, formulating testable hypotheses, and selecting equipment and technology, including graphing calculators, computers and probes, sufficient scientific glassware such as beakers, Erlenmeyer flasks, pipettes, graduated cylinders, volumetric flasks, safety goggles, and burettes, electronic balances, and an adequate supply of consumable chemicals. (F) Collect data and make measurements with accuracy and precision. (G) Express and manipulate chemical quantities using scientific conventions and mathematical procedures, including dimensional analysis, scientific notation, and significant figures. (H) Organize, analyze, evaluate, make inferences, and predict trends from data. (I) Communicate valid conclusions supported by the data through methods such as lab reports, labeled drawings, graphs, journals, summaries, oral reports, and technology-based reports.

(3) Scientific processes. The student uses critical thinking, scientific reasoning, and problem solving to make informed decisions within and outside the classroom. The student is expected to: (A) In all fields of science, analyze, evaluate, and critique scientific explanations by using empirical evidence, logical reasoning, and experimental and observational testing, including examining all sides of scientific evidence of those scientific explanations, so as to encourage critical thinking by the student. (B) Communicate and apply scientific information extracted from various sources such as current events, news reports, published journal articles, and marketing materials. (C) Draw inferences based on data related to promotional materials for products and services. (D) Evaluate the impact of research on scientific thought, society, and the environment. (E) Describe the connection between chemistry and future careers. (F) Research and describe the history of chemistry and contributions of scientists.

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Chemistry TEKS continued (4) Science concepts. The student knows the characteristics of matter and can analyze the relationships between chemical and physical changes and properties. The student is expected to: (A) Differentiate between physical and chemical changes and properties.

Readiness Standard

(B) Identify extensive and intensive properties.

Supporting Standard

(C) Compare solids, liquids, and gases in terms of compressibility, structure, shape, and volume.

Supporting Standard

(D) Classify matter as pure substances or mixtures through investigation of their properties.

Readiness Standard

(5) Science concepts. The student understands the historical development of the Periodic Table and can apply its predictive power. The student is expected to: (A) Explain the use of chemical and physical properties in the historical development of the Periodic Table.

Supporting Standard

(B) Use the Periodic Table to identify and explain the properties of chemical families, including alkali metals, alkaline earth metals, halogens, noble gases, and transition metals.

Readiness Standard

(C) Use the Periodic Table to identify and explain periodic trends, including atomic and ionic radii, electronegativity, and ionization energy.

Readiness Standard

(6) Science concepts. The student knows and understands the historical development of atomic theory. The student is expected to: (A) Understand the experimental design and conclusions used in the development of modern atomic theory, including Dalton’s Postulates, Thomson’s discovery of electron properties, Rutherford’s nuclear atom, and Bohr’s nuclear atom.

Supporting Standard

(B) Understand the electromagnetic spectrum and the mathematical relationships between energy, frequency, and wavelength of light.

Supporting Standard

(C) Calculate the wavelength, frequency, and energy of light using Planck’s constant and the speed of light.

Supporting Standard

(D) Use isotopic composition to calculate average atomic mass of an element.

Supporting Standard

(E) Express the arrangement of electrons in atoms through electron configurations and Lewis valence electron dot structures.

Readiness Standard

(7) Science concepts. The student knows how atoms form ionic, metallic, and covalent bonds. The student is expected to: (A) Name ionic compounds containing main group or transition metals, covalent compounds, acids, and bases, using International Union of Pure and Applied Chemistry (IUPAC) nomenclature rules.

Readiness Standard

(B) Write the chemical formulas of common polyatomic ions, ionic compounds containing main group or transition metals, covalent compounds, acids, and bases.

Readiness Standard

(C) Construct electron dot formulas to illustrate ionic and covalent bonds.

Readiness Standard

(D) Describe the nature of metallic bonding and apply the theory to explain metallic properties such as thermal and electrical conductivity, malleability, and ductility.

Supporting Standard

(E) Predict molecular structure for molecules with linear, trigonal planar, or tetrahedral electron pair geometries using Valence Shell Electron Pair Repulsion (VSEPR) theory.

Supporting Standard

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(8) Science concepts. The student can quantify the changes that occur during chemical reactions. The student is expected to: (A) Define and use the concept of a mole.

Supporting Standard

(B) Use the mole concept to calculate the number of atoms, ions, or molecules in a sample of material.

Readiness Standard

(C) Calculate percent composition and empirical and molecular formulas.

Supporting Standard

(D) Use the law of conservation of mass to write and balance chemical equations.

Readiness Standard

(E) Perform stoichiometric calculations, including determination of mass relationships between reactants and products, calculation of limiting reagents, and percent yield.

Supporting Standard

(9) Science concepts. The student understands the principles of ideal gas behavior, kinetic molecular theory, and the conditions that influence the behavior of gases. The student is expected to: (A) Describe and calculate the relations between volume, pressure, number of moles, and temperature for an ideal gas as described by Boyle’s law, Charles’ law, Avogadro’s law, Dalton’s law of partial pressure, and the ideal gas law.

Readiness Standard

(B) Perform stoichiometric calculations, including determination of mass and volume relationships between reactants and products for reactions involving gases.

Supporting Standard

(C) Describe the postulates of kinetic molecular theory.

Supporting Standard

(10) Science concepts. The student understands and can apply the factors that influence the behavior of solutions. The student is expected to: (A) Describe the unique role of water in chemical and biological systems.

Supporting Standard

(B) Develop and use general rules regarding solubility through investigations with aqueous solutions.

Readiness Standard

(C) Calculate the concentration of solutions in units of molarity.

Supporting Standard

(D) Use molarity to calculate the dilutions of solutions.

Supporting Standard

(E) Distinguish between types of solutions such as electrolytes and nonelectrolytes and unsaturated, saturated, and supersaturated solutions.

Readiness Standard

(F) Investigate factors that influence solubilities and rates of dissolution such as temperature, agitation, and surface area.

Readiness Standard

(G) Define acids and bases and distinguish between Arrhenius and Bronsted-Lowry definitions and predict products in acid-base reactions that form water.

Supporting Standard

(H) Understand and differentiate among acid-base reactions, precipitation reactions, and oxidation-reduction reactions.

Readiness Standard

(I) Define pH and use the hydrogen or hydroxide ion concentrations to calculate the pH of a solution.

Supporting Standard

(J) Distinguish between degrees of dissociation for strong and weak acids and bases.

Supporting Standard

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Chemistry TEKS continued (11) Science concepts. The student understands the energy changes that occur in chemical reactions. The student is expected to: (A) Understand energy and its forms, including kinetic, potential, chemical, and thermal energies.

Supporting Standard

(B) Understand the law of conservation of energy and the processes of heat transfer.

Supporting Standard

(C) Use thermochemical equations to calculate energy changes that occur in chemical reactions and classify reactions as exothermic or endothermic.

Readiness Standard

(D) Perform calculations involving heat, mass, temperature change, and specific heat.

Supporting Standard

(E) Use calorimetry to calculate the heat of a chemical process.

Supporting Standard

(12) Science concepts. The student understands the basic processes of nuclear chemistry. The student is expected to: (A) Describe the characteristics of alpha, beta, and gamma radiation.

Supporting Standard

(B) Describe radioactive decay process in terms of balanced nuclear equations.

Readiness Standard

(C) Compare fission and fusion reactions.

Supporting Standard

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TEKS REVIEW

1A

Demonstrating Safe Practices TEKS 1A Demonstrate safe practices during laboratory and field investigations, including the appropriate use of safety showers, eyewash fountains, safety goggles, and fire extinguishers.

What safety procedures should be followed during investigations? Because laboratory and field investigations often involve the use of hazardous or potentially hazardous materials and equipment, the risk of accidents or injury is always present. However, by following some general safety practices and procedures during investigations, accidents and injuries can be prevented. Basic safety requirements for working in the laboratory include knowing emergency procedures and the locations of all safety equipment, following directions, and never working in the laboratory alone. It is also important to be aware of the hazards and handling procedures for all the materials and equipment you use. Other important safety policies are listed in Figure 1. Figure 1

Laboratory Safety Do’s and Don’ts • Do not touch any chemical with your hands.

• Notify the teacher of any sensitivities or allergies to chemicals or other substances.

• Never inhale chemical vapors by placing the container directly under your nose.

• Do not leave an experiment unattended.

• Never pour chemicals down the sink drain unless you’ve been specifically instructed to do so.

• Never chew gum, eat, or drink in the laboratory. • Always wear shoes and avoid wearing loose-fitting clothing and dangling jewelry.

• When diluting an acid, always pour the acid slowly into the water, stirring to dissipate the heat. CAUTION: Never pour water into a concentrated acid.

• Secure long hair and loose clothing; roll up loose sleeves when working with burners or flames. • Inspect all equipment for damage prior to use; do not use damaged equipment.

• Know the location of all emergency exits in the laboratory and the building.

• Keep the floor clear of all objects such as personal items, spilled liquids, and any other item that may cause someone to trip or fall.

•Wash your hands with soap and water at the end of each investigation.

• Never point the open end of a test tube containing a chemical at other people.

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Study Tip Look for all symbols that represent safety equipment and protective devices in the laboratory. Learn the meaning of the symbols and purpose of each device so that you can act quickly in an emergency.

What is the appropriate use of important safety equipment? Safe practices should always be the highest priority during laboratory and field investigations. Each person has a responsibility to learn the dangers that may be present while working in the laboratory and the purpose and operation of all emergency safety equipment. Although everyone should know how to use safety equipment, a teacher’s instructions should always be followed in an emergency. One of the most important safety practices is to be familiar with the proper use of protective safety devices. These devices include safety showers, fire blankets, eyewash fountains, safety goggles, aprons, gloves, and fire extinguishers. Such knowledge is vital to responding to accidents and to eliminating the risk of serious personal injury.

When is it appropriate to use a safety shower or fire blanket? Whenever the skin or clothing is exposed to a significant amount of corrosive or toxic chemicals, the contaminants must be immediately washed away with large quantities of water. An emergency safety shower is the most effective way to quickly eliminate contaminants on your skin or clothing and avoid injury. In any circumstance in which the use of a safety shower is necessary, it is important to act quickly and remove any affected articles of clothing to avoid further exposure. Once the hazardous materials have been washed away, obtain medical attention immediately. If clothing or hair catches fire, do not run because running fans flames. A fire blanket can be used to smother flames. Or, a safety shower can be used if there is one nearby.

When is it appropriate to use safety goggles? As a rule, safety goggles should be worn at all times in the laboratory and during field investigations. During an experiment, it is likely you will use wet or dry chemicals. There is a constant danger of splashes or particles entering the eye. Safety goggles can also help protect eyes from damage due to explosions and flying debris.

When is it necessary to use an eyewash fountain? Even though the use of safety goggles is required, eyewash fountains are located in all laboratory environments where the eyes may be exposed to hazardous chemicals. In the event that one or both eyes are exposed to a hazardous substance, an eyewash fountain should be used without delay. Hold the eyelids open, and flush the affected eyes thoroughly for several minutes to ensure the substance has been purged. Then seek immediate medical attention.

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What types of fire extinguishers are appropriate for use in laboratory environments? Not all fires are the same, and no single fire extinguisher works on every type of fire. Fires are designated by type and listed by class. The most common fire classifications, combustibles, and appropriate extinguishing agents are shown Figure 2. Figure 2

Fire Classifications, Combustibles, and Extinguishers Class A B C D

Combustibles Ordinary materials such as paper, wood, cardboard, and plastics Flammable or combustible liquids such as gasoline, kerosene, and most organic solvents Electrical equipment Combustible metals such as magnesium, potassium, and sodium

Extinguishing Agent Dry chemical CO2 or dry chemical CO2 or dry chemical Dry powder agents

Water extinguishers are never used in the laboratory as they are rated only for Class A fires. The use of a water extinguisher on a Class B, C, or D fire in the laboratory or field is extremely hazardous. Pouring water on fires involving combustible liquids, electrical equipment, or combustible metals may spread the fire or make it worse.

TEKS

End-of-Course Assessment Review

1. Classify  If isopropyl alcohol were to catch fire in the laboratory, what class fire would this be? A. Class A B. Class B C. Class C D. Class D 2. Infer  If an electric hot plate caught fire in the lab, which class of fire extinguisher would you use and why? 3. Differentiate  What is the appropriate personal safety response to a significant amount of corrosive liquid chemical splashed on clothing? 4. Evaluate  A student says that he does not need to wear safety goggles for an experiment because no liquid chemicals are being used. What would your response be to this student?

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TEKS REVIEW

1B

Hazards of Chemical Substances TEKS 1B Know specific hazards of chemical substances such as flammability, corrosiveness, and radioactivity as summarized on the Material Safety Data Sheets (MSDS).

Vocabulary Material Safety Data Sheets (MSDS) flammable substance corrosive substance radioactive substance

Why is it important to know specific hazards of chemical substances? Before working with any chemical substance, it is important to be thoroughly familiar with its properties, specific hazards, safety precautions, and handling procedures. Hazardous chemical substances are generally classified according to their hazard types as listed in Figure 1 below.

What are Material Safety Data Sheets (MSDS)? Material Safety Data Sheets (MSDS) are data forms that contain detailed information on the properties, hazards, and health effects of chemical substances. MSDS also provide guidelines for the safe handling, storage, and disposal of hazardous substances. The sheets are prepared and made available by chemical suppliers.

Figure 1

The U.S. Occupational Safety and Health Administration (OSHA) requires the presence of MSDS wherever hazardous materials are produced, shipped, or used. These locations include all chemistry laboratories. An MSDS includes the information listed in Figure 2.

Hazardous Chemical Substances Type of Hazard

Description

Examples

Flammable/Combustible Liquids

Generate vapors that will burn when ignited

Acetone Methanol

Corrosives

Corrode metal and damage living tissues (acids, bases, and others)

Sulfuric acid Sodium hydroxide

Oxidizers

Cause other materials to combust

Bromine Hydrogen peroxide

Water Reactives

React with water to form heat and flammable gases

Alkali metals

Pyrophorics

Ignite spontaneously in air

Diethylzinc Diphsophine

Peroxide-forming

Explodes if subject to shock or sparks

Isoprophyl ether Potassium amide

Compressed Gases

Disperse forcefully and quickly if released

Oxygen Acetylene

Cryogens

Freeze human tissue quickly (super-cooled fluids)

Liquid nitrogen

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Figure 2

Material Safety Data Sheets Section Chemical Identity

Name, weight, and chemical formula

Manufacturer

Name, address, and phone number

Hazardous Ingredients

Hazardous components by chemical identity

Physical/Chemical Characteristics

Boiling point, vapor pressure, melting point, and other physical and chemical properties

Fire & Explosion Hazard

Flash point, flammability limits, extinguishing method, and firefighting procedures

Reactivity

Stability, a list of materials and conditions to avoid, and hazardous by-products

Health Hazard

Routes of physical entry such as inhalation, ingestion, or skin, symptoms of exposure, and emergency/first aid procedures

Safe Handling & Use

Precautions for handling and storage, and steps to be taken if a spill or release occurs

Control Measures

Protective measures for handling

In addition to Materials Safety Data sheets, many hazardous materials are also labeled with a “hazard diamond”, published by the National Fire Protection Association (NFPA). The NFPA warning label rates materials for health (blue), flammability (red), and instability (yellow). The three colorcoded sections range from 0 (the least severe hazard) to 4 (the most severe hazard.) The bottom section is usually blank. It may be used to present specific hazards or special fire-fighting measures.

Study Tip Review the MSDS of chemicals that you use in the laboratory. Become familiar with the format and the information presented in the forms.

What are the specific hazards of flammable, corrosive, and radioactive substances? Flammables  A flammable substance gives off combustible vapors that can easily ignite. Flammables include solids, liquids, and gases. These substances have a flash point of below 100º F. Their vapors can ignite at temperatures near room temperature.

Figure 3 “Hazard Diamond”

As their MSDS specifies, flammable substances should only be used only with proper ventilation and away from heat, electric sparks, and flames. Flammable substances should also be stored in approved fire-retardant storage cabinets and should never be placed near corrosives. Flammable substances include diethyl ether, acetone, gasoline, toluene, and methyl alcohol.

3

Flammability

1

3

Instability Specific Hazards

2

Instability Specific Hazards

2

Health

Flammability

h

Description

0 1 2 3 4

Low Slight Moderate High Extreme

0 1 2 3 4

Low Corrosives  A corrosive substance is a compound that is highly reactive Slight and will cause serious damage to living tissue. The MSDS for corrosives Moderate include warnings against contact with skin, eyes, or any part of the body. High They also recommended first aid measures that include flushing a Extreme damaged area with water for up to 15 minutes, followed by immediate

medical attention. Corrosives should be stored in well-ventilated areas and away from flammables and heat. Typical corrosives found in laboratories include sulfuric acid, hydrochloric acid, and sodium hydroxide.

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Radioactivity  A radioactive substance spontaneously emits ionizing radiation. These substances include solids, liquids, and gases. Depending on the level of radioactivity of the substance and the level of exposure, ionizing radiation can damage cells, injure tissues and organs, and can lead to cancer. Use of radioactive materials is tightly controlled. Anyone who uses radioactive substances must be trained and certified. As with all hazardous substances, the MSDS provides safe handling procedures and safety precautions, as well as first aid and containment measures.

TEKS

End-of-Course Assessment Review

1. Infer  The reason that radioactive substances are dangerous is A they will react violently with other chemicals. B the radioactivity will spread to any chemical placed nearby. C they can burn in air. D the radioactivity can damage living cells. 2. Identify  Which is a characteristic of a corrosive substance? A The substance forcefully and quickly disperses if released. B The substance quickly freezes human tissue. C The substance seriously damages skin cells on contact. D The substance causes other materials to combust. 3. Evaluate  While some students are working in the lab, one student spills a corrosive substance and asks the others to help him mop it up with paper towels and carry them to the trash can. Another student objects and says they should warn everyone nearby and inform the instructor. Which approach is the appropriate safe response, and why? 4. Infer  Give at least one reason why diethyl ether, a flammable liquid, should not be stored in a commercial refrigerator. 5. Evaluate  Why is storing chemicals in alphabetical order not a safe approach?

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TEKS REVIEW

1C

Conservation of Resources and Proper Disposal TEKS 1C Demonstrate an understanding of the use and conservation of resources and the TEKS_TXT proper disposal or recycling of materials.

Vocabulary decontamination

Figure 1 Recycling Symbol

Why is the conservation of resources and proper disposal or recycling of materials important? Most natural resources are limited. Conservation of these resources will ensure that they are available for future generations. Similarly, proper disposal or recycling of waste materials is essential to maintaining human and environmental health. Figure 1 is the universal sign for recycling. But even if you see this symbol, always follow your teacher’s instructions regarding proper disposal or recycling of materials in the laboratory. Waste reduction is especially important in a chemistry laboratory. Chemistry experiments involve many chemicals and they may produce hazardous wastes. One way to reduce waste is to use the smallest amounts of chemicals required whenever possible. Another way to reduce waste is to use materials that can be recovered, or recycled, instead of being discarded. When chemicals cannot be recycled, disposal must follow strict guidelines and comply with local, state, and federal regulations. These regulations were implemented to avoid the health problems and expense caused by pollution and contamination. By effectively managing the use of materials through conservation and recycling, and by properly disposing chemical waste, both human health and the environment can be protected.

What are the proper ways to reuse and recycle materials in the laboratory? Many materials used in the laboratory can be reused after thorough cleaning and safe removal of chemical residues. These materials include containers and instruments made of sturdy glass, plastic, and metal. Materials that cannot be reused, but can be recycled, should be collected and sent to recycling plants. Recyclable items should always be discarded in containers designated for each type of material. Before placing items in recycling containers, care should be taken to make sure they are decontaminated. Decontamination is the removal of hazardous compounds. Decontamination should only be performed under the supervision of your teacher or qualified laboratory personnel.

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Study Tip Think about the lab investigations you performed throughout the year. Recall the disposal methods used for various chemicals and other materials.

Some chemicals used in the laboratory can also be reused or recycled. Solvents such as acetone, methanol, and toluene are routinely reused because they can easily be purified. Chemicals that cannot be purified in the laboratory may still be good candidates for recycling because of the value of the chemical. For example, solutions containing silver ions are frequently recycled because silver is very expensive. Designated containers should always be used for each chemical to be recycled because mixing certain materials can be extremely hazardous.

What are the proper ways to dispose of materials in the laboratory? Whenever chemical substances and equipment are used in the laboratory, some waste will be generated. Because of the potential risks, very little laboratory waste can be disposed of in public waste containers. However, some laboratory consumables—materials that cannot be recycled and were not exposed to chemicals—may be discarded as regular trash. Your teacher will tell you which, if any, of the lab materials you use can be disposed of in the regular trash. Hazardous or toxic chemical wastes must be disposed of separately and according to proper guidelines for safe disposal. They also must be properly labeled. The manner in which hazardous materials are disposed of is determined by the reactive properties of each substance. The recommended disposal method for a substance is provided on its Material Safety Data Sheet (MSDS).

Figure 2 No Disposal Sign

NOTICE DO NOT

DUMP CHEMICALS DOWN THIS DRAIN

Chemicals such as dilute acids, bases, and certain organic compounds can be discarded by pouring them down the drain with large quantities of water. Materials that are acceptable for this disposal method are water soluble, have very low toxicity, and, if organic, are readily biodegradable. However, this type of disposal should only be performed with the approval of your teacher and after a thorough review of disposal guidelines. Figure 2 shows one example of a warning that may be found posted on some laboratory sinks. In some situations, no chemicals can be poured down a drain. Some hazardous chemical wastes can be neutralized, or made non-hazardous. Wastes that cannot be neutralized must be shipped to a hazardous waste landfill by a licensed company approved by the Department of Transportation. You will most likely not be working with this type of chemical wastes in your chemistry labs. Disposal of all hazardous chemical wastes must comply with local, state, and federal regulations. The federal agencies responsible for regulating waste disposal are the U.S. Environmental Protection Agency (EPA) and the Occupational Safety and Health Administration (OSHA). Failure to properly dispose of chemical wastes is dangerous and unhealthy, and can result in fines and lawsuits.

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TEKS

End-of-Course Assessment Review

1. Identify  A student uses a glass beaker to mix hydrochloric acid and water in an experiment. After the experiment is complete, should she place the beaker in a recycling bin suitable for glass? A Yes. Glass beakers should always be recycled after any use. B No. The beaker should be discarded in the trash. C Yes. Hydrochloric acid is dangerous and any item that contains it should be recycled immediately. D No. Glass beakers are reusable. They can easily be cleaned after an experiment and stored. 2. Analyze  A researcher finishes an experiment and is unsure of how to dispose of a particular chemical. Which of the following describes the safest approach? A Contain the chemical tightly and put it in the trash. B Combine it with other chemical waste. C Label the chemical as “Waste” and leave it out for someone else to dispose of. D Refer to the chemical’s MSDS for the recommended disposal method. 3. Evaluate  At the end of a laboratory experiment, a student disposes of all liquid chemicals by flushing them down the sink drain with water. Explain what is wrong with this action. 4. Demonstrate Understanding  During an investigation, Evan spills some compound on the lab table. His lab partner tells him that the compound should be recycled. Evan collects the substance and disposes of it in a container marked for general recycling. Explain two mistakes Evan made. 5. Infer  Why would a laboratory never have a single container for all waste chemicals?

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TEKS REVIEW

2A

Definition of Science TEKS 2A Know the definition of science and understand that it has limitations, as specified in subsection (b)(2) of this section. (b)(2) Nature of science. Science, as defined by the National Academy of Sciences, is the “use of evidence to construct testable explanations and predictions of natural phenomena, as well as the knowledge generated through this process.” This vast body of changing and increasing knowledge is described by physical, mathematical, and conceptual models. Students should know that some questions are outside the realm of science because they deal with phenomena that are not scientifically testable.

What is the definition of science?

Study Tip Remember that scientific concepts must be part of the natural and physical world and must be testable and falsifiable. Concepts that do not fit into these categories are not scientific.

Science is the use of evidence to construct testable explanations and predictions of natural phenomena, as well as the knowledge generated through this process. In other words, science is the study of the natural and physical world using physical, mathematical, and conceptual models. Scientific explanations must be both testable and falsifiable—able to be proven incorrect. Observation, experimentation, research, and the use of models produce evidence that allow scientists to understand natural phenomena. Scientists study patterns and make predictions about natural phenomena and processes to understand how the world works. In many cases, because of the use of observation, experimentation, research, and models, scientists can predict the results of a natural process even if they do not have all the information about that process. Many explanations of natural processes are accepted as valid because there is so much evidence supporting them, and because they have been observed and/or tested under a wide variety of conditions. When scientific explanations have been tested and widely accepted, predictions about future events usually end up to be accurate. Science is not the same as technology. Technology is the application of science, often for industrial or commercial uses. Science identifies how or why a natural or physical phenomenon occurs. Technology identifies how to apply that phenomenon for a practical use.

Why study science? There are many different reasons why people study science. Chemists might study compounds for potential use in medicines. Meteorologists might study weather patterns to predict hurricanes and tornadoes. Geologists might study natural processes to recognize how events in the past might influence events, such as earthquakes, in the future. Doctors, dentists, veterinarians, nurses, and pharmacists study science to provide health care to you and your pets. Physicists study the physical world from the smallest of particles to the vastness of the universe. What would you be most interested in studying through science? TEKS 2A • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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What are the limitations of science? Because science is the study of natural and physical phenomena, science has limitations. Science is not emotion, art, or feeling. Science cannot determine which painting is more appealing or who is the best choice for president. Science cannot answer questions regarding faith or personal feelings. Such phenomena are outside the realm of science because they are not scientifically testable. Science can provide information, but nonscientific factors decide how we use science. Current scientific knowledge is limited to the information presently known about the natural and physical world. This is why all scientific hypotheses and models are subject to change. As we learn new information, current scientific understandings sometimes become outdated. As new information becomes available, new technologies may also arise, making old technologies obsolete. For example, Figure 1 shows how scientists’ model of the atom has changed over time. As scientists conducted new experiments and gathered new evidence and data, they discarded older models of the atom in favor of revised or new models. Each newer model supported the new information that scientists had gathered. Figure 1 A Changing Model of The Atom +

J. J. Thomson’s model, about 1904

Ernest Rutherford’s model, 1911

+

Niels Bohr’s model, 1913

Modern Model

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Many ideas and explanations that are currently known about the natural and physical world are a result of the use of physical, mathematical, and conceptual models. Sometimes a process or idea is too large (such as the universe) or too small (such as atoms) to be studied directly, or because it is too dangerous or too expensive to be studied directly. It is important to note that models never represent a process or idea perfectly, and will change as more knowledge is gained through additional scientific research. As models become more sophisticated, they can more accurately predict the system they are concerned with.

TEKS

End-of-Course Assessment Review

1. Define  Which of the following questions is not a scientific question? A What caused dinosaurs to become extinct? B How is hydrochloric acid produced and contained in the stomach? C How are atoms of nitrogen different from atoms of carbon? D Was Isaac Newton the greatest scientist that ever lived? 2. Define  The circuitry for computers was invented after scientists learned how electrons flow through certain materials, such as silicon. Computer circuitry is an example of A a prediction. B a limitation of science. C technology. D a model. 3. Evaluate  Your friend tells you that commercial space travel to other planets will never be possible. The technology to get people into space is too expensive, too dangerous, and too complicated. Knowing what you know about science, what would you tell your friend? 4. Compare and Contrast  Describe an area of study that is not science. Write three to five sentences describing why the area you chose is not science and how it could be changed to qualify as science.

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TEKS REVIEW

2B

Scientific Hypotheses TEKS 2B Know that scientific hypotheses are tentative and testable statements that must be capable of being supported or not supported by observational evidence. Hypotheses of TEKS_TXT durable explanatory power which have been tested over a wide variety of conditions are incorporated into theories.

What is a scientific hypothesis?

Vocabulary Hypothesis

Suppose you slice an apple in half. You place half the apple in the refrigerator, and set the other half on the counter and leave the room. When you return an hour later, you notice that the apple half on the counter has turned brown. You look in the refrigerator and observe that the half of the apple you placed inside is only slightly brown. You recall that a change in color of a substance can be an indicator that a chemical reaction has taken place. After observing both apple halves, you hypothesize that the chemical reactions that cause an apple to brown occur more quickly at higher temperatures. In the example above, you developed a tentative explanation, or scientific hypothesis, for why the apple half on the counter undergoes a chemical reaction more quickly than the half in the refrigerator—chemical reactions that cause an apple to brown occur more quickly at higher temperatures. A scientific hypothesis is a tentative statement or explanation for an observation in nature. Scientific hypotheses are capable of being tested and supported, or not supported, through further observation and experimentation.

Why must a scientific hypothesis be testable?

Study Tip The root words contained in hypothesis tell you the meaning of the word. Hypo- means “under,” and -thesis means “proposition.” In a way, a hypothesis is an underlying proposition for an experiment or observation.

Typically, once a scientific hypothesis is stated, the next step is to develop an experiment or conduct observational research to identify evidence that either supports or does not support the hypothesis. This is because a hypothesis has no meaning unless there is observational evidence or data that supports it. For example, the scientific hypothesis, “Chemical reactions that cause an apple to brown occur more quickly at higher temperatures” has no meaning without an experiment or data to support it. Therefore, a scientific hypothesis must be testable. Then, information and evidence gathered can be analyzed to draw conclusions about the hypothesis. In some cases, a hypothesis will be supported by the evidence that accumulates. In other cases it will not be supported.

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What if a scientific hypothesis is supported? Suppose you design a scientific experiment to test the effect of temperature on the rate of the chemical reactions that cause an apple to brown. If you find that your experiment supports your hypothesis, are you finished? Not quite. The experiment should be repeated several times to confirm the results and to ensure that no errors have been made. Additionally, a hypothesis should be tested over a variety of conditions to confirm that all variables have been considered that might alter the results. In the best scientific tradition, it is also important to have others repeat the experiment separately to confirm the results. It is always possible that a single investigator may accidentally introduce some form of bias into the results. The more investigators who have been able to replicate the results, the less likely it will be that any bias is involved. Hypotheses that have undergone significant testing by multiple independent scientists over a variety of conditions can be said to have durable explanatory power—that is, they have stood up to multiple tests by many scientists. If hypotheses have been consistently supported through multiple tests, they are incorporated into a theory (if one exists) related to its given topic. At that point, additional research and tests are often devised in attempt to ensure that the revised theory is supported under all known conditions.

What if a scientific hypothesis is not supported? On the other hand, suppose the experiment did not support the hypothesis. Was the experiment then a failure? Not necessarily. A scientific investigation is never a failure, as long as it leads to information and knowledge you did not previously have. But if it is not supported, the original hypothesis itself is not very useful for making predictions or understanding observations, so it is typically modified, or in some cases, discarded. If it is modified, the cycle begins again with additional observational research and data or a new experiment to test the modified hypothesis. This cycle of development is illustrated in Figure 1. Figure 1 Development of a Hypothesis

Make an observation

Yes: Repeat experiment or gather additional data to confirm.

Form a scientific hypothesis based on the observation

Gather observational evidence and data or develop an experiment to test the hypothesis

Is the hypothesis supported?

No: Modify or reject the hypothesis

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Why is a scientific hypothesis only a tentative explanation? It is important to note that hypotheses are only tentative explanations and are not proven facts, regardless of how many experiments and observations support the hypothesis. This is because as new data and evidence become available, a hypothesis may need to be revised or even rejected altogether. As such, a hypothesis can never be proven true or accepted as absolute truth. It can only be supported through further observation, evidence, and experimentation.

TEKS

End-of-Course Assessment Review

1. Infer  Which of the following is a logical next step if a scientist’s repeated experiments do not support his hypothesis? A Alter the experiment so that the hypothesis would be supported. B Incorporate the scientific hypothesis into a theory. C Modify the hypothesis and conduct a new experiment. D End the failed investigation. 2. Identify  Why can a scientific hypothesis never be proven true? A A scientific hypothesis is unreliable. B Supporting evidence is difficult to identify. C New information might become available that contradicts the scientific hypothesis. D It is impossible to design an experiment that can directly test a scientific hypothesis. 3. Explain  What does it mean when a hypothesis is said to have durable explanatory power? Explain. 4. Hypothesize  Suppose you find that a battery-operated flashlight is not working. Write a hypothesis explaining why your flashlight might not work. Then, explain why your statement qualifies as a hypothesis. 5. Analyze  Suppose you designed several experiments to test your hypothesis in Question 4 above. It turns out that none of your experiments supported your hypothesis. Was your experimentation and hypothesis a failure? Explain. Then, describe what steps you should take next.

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TEKS REVIEW

2C

Scientific Theories TEKS 2C Know that scientific theories are based on natural and physical phenomena and are capable of being tested by multiple independent researchers. Unlike hypotheses, scientific theories are well-established and highly-reliable explanations, but they may be subject to change as new areas of science and new technologies are developed.

Vocabulary scientific theory

What does it mean that scientific theories are considered well-established and highly-reliable explanations? A scientific theory is different from the common use of the word theory. If you say that “I have a theory as to why Texas A&M lost that game,” you mean that you suspect that you know the reason; you have a guess. But that is very different from a scientific theory. (In this review, when we use the term theory we mean scientific theory.) A scientific theory is a well established, highly-reliable explanation of a natural or physical phenomenon. Natural phenomena include every part of our physical environment. Natural phenomena also include the forces and energies that operate on and within our environment, such as gravity. Physical phenomena include anything that can be observed with one or more of our senses. A theory cannot be based on a nonnatural or a nonphysical cause and still be considered to be a scientific theory. Being established and highly reliable means that a theory has been repeatedly and consistently upheld by numerous, extensive scientific investigations conducted by many independent researchers. Theories are capable of unifying a broad range of observations and hypotheses. A theory is powerful because it can be used to predict a wide variety of future events. A theory also explains how or why an event or process occurs. For example, the kinetic molecular theory explains how gas particles move. This explanation can be applied to predict the behavior of any gas under a wide variety of circumstances.

Why must a theory be capable of being tested by multiple independent researchers? Anyone could propose an explanation for events in nature. A scientific theory, however, has been tested by multiple independent researchers. This means that many scientists working separately from one another have verified the results of experiments related to the theory. This is important because individual researchers can make errors, introduce investigator bias, or use faulty methods. When a large number of independent researchers conduct investigations, the results are much more likely to be accurate.

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How are theories subject to change as new areas of science and new technologies are developed? Although theories are thoroughly tested and evaluated, they can be changed if further scientific study supports a better explanation for the phenomenon being studied. That is, a theory is the most useful and powerful explanation of the data available at the current time. Strictly speaking, a theory is neither accurate nor inaccurate. All theories are subject to change as new areas of science and new technologies are developed. If new evidence is identified that is not consistent with an existing theory, the theory might be revised or rejected. For example, part of Dalton’s atomic theory explained that atoms are indestructible and never change into other atoms. Years later, with the aid of new technology, other scientists observed changes to atoms resulting from radioactivity. They also observed nuclear fission, the process in which large atoms break apart into smaller atoms. As a result, parts of Dalton’s atomic theory were revised.

Study Tip To remember the power of a theory, think of the letters WEHR, which stand for Well Established, Highly Reliable.

Figure 1

What is the relationship between a scientific theory and a scientific law? A common misconception is that when enough evidence is gathered, a scientific theory can become a law. In fact, scientific laws and theories are very different. Both laws and theories are supported by large bodies of evidence gathered by multiple independent researchers. However, a theory explains a phenomenon, while a law does not offer an explanation. A scientific law is a concise statement that summarizes the results of many observations and experiments. For example, the law of conservation of energy states that energy can be changed from one form to another, but it is neither created nor destroyed. This law is well-supported by the results of many experiments. But because it does not explain how energy is conserved, it is a law instead of a theory. Some common misconceptions about theories are listed in Figure 1 below.

Common Misconceptions About Theories Misconceptions

Facts

A theory is a fact.

A theory can never be proven true. New technology, discoveries, and information can lead to its modification or rejection.

A theory is a guess.

In everyday language, theory refers to a guess or a suspicion. However, a scientific theory is an explanation that is both reliable and well-supported.

Over time, a theory can become a law.

A scientific theory cannot become a scientific law. A theory explains events, while a law does not.

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TEKS

End-of-Course Assessment Review

1. Identify  Which of the following best defines a scientific theory? A a well-established, highly-reliable explanation of a natural event B a preliminary guess or idea about an event in nature C a true statement about an event in nature D a well-established principle that does not include an explanation 2. Explain  When would an established scientific theory most likely be revised or replaced? A when one scientist argues against the theory B when public opinion amasses against the theory C when new evidence is gathered that does not support the theory D when the theory is promoted to a law 3. Evaluate  Suppose that you were to hear that a talented chemistry student in another class had just discovered a new theory of chemistry. Evaluate that claim based on three characteristics of a scientific theory outlined in this TEKS. 4. Describe  If scientific theories cannot be proven true, why are they so powerful and useful? 5. Evaluate  Your friend is describing the concept of gravity. She states that if she drops an object, it will fall to the ground every time. Most likely, does the friend’s description of gravity involve a hypothesis, a theory, or a law? Explain. 6. Explain  How does the development of new technologies affect scientific theories? Explain.

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TEKS REVIEW

2D

Scientific Hypotheses and Scientific Theories TEKS 2D

Distinguish between scientific hypotheses and scientific theories. TEKS_TXT

What is a scientific hypothesis? A scientific hypothesis is a proposed explanation for observations or an answer to a scientific question for which you can gather objective data that supports or refutes it. To create a hypothesis, a scientist asks a question about how or why a specific event occurs (or does not occur) and constructs a statement that explains the phenomenon. This statement, if testable, can be the hypothesis for an investigation. Hypotheses are proposals. They are starting points for specific, controlled research projects. Scientists must test hypotheses. When possible, scientists test hypotheses by setting up experiments that involve independent variables, factors that change during an experiment, to determine if they affect the outcome. When a scientist changes an independent variable, he or she records any changes in the outcome, the dependent variable. Once the experiment is complete, the resulting data can be examined to determine if they support the hypothesis. If the hypothesis is not supported, the scientist must construct a new hypothesis, and the process begins again. For example, suppose that a researcher observed that bath towels seem to lose their absorbency when they are dried using fabric softener sheets. He constructs the following hypothesis: Using fabric softener causes towels to become less absorbent. The researcher can test this hypothesis by devising and performing an experiment. One experiment could involve two test groups of identical towels. Each test group is dried under almost identical conditions. The only difference is that a fabric softener sheet is added to the dryer for one test group only. After the towels are dried, the absorbency of each towel is measured. If the towels dried with the fabric softener sheets are less absorbent than the other towels, then the hypothesis is supported.

Study Tip The prefix hypo- means “beneath” or “less than.” You can think of a hypothesis as a less powerful statement than a theory.

How can you distinguish between scientific hypotheses and scientific theories? In everyday conversation, the word theory generally means “suspicion.” You might say, “I have a theory as to why you got a C on the chemistry test.” But in science the word theory has a different meaning. A scientific theory is a well-supported explanation for observations made in many situations. So theories are much broader than hypotheses.

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Most hypotheses refer to a specific situation or case, but a theory provides an explanation for a broad range of observations. In chemistry, scientists might test hypotheses concerning the physical properties of oxygen gas, the electrical conductivity of a solution of sodium chloride, or the chemical reactivity of compounds such as acetic acid or ammonia. In contrast, the theories of chemistry explain the physical or chemical properties of a large number of elements or compounds. A theory might explain the chemical reactivities of a large class of compounds, such as acids or bases. In the flowchart in Figure 1, you can compare the roles that hypotheses and theories play in scientific methodology. Figure 1 Role of Hypotheses and Theories

Observations

Hypothesis

Experiments

Theory

A hypothesis may be revised based on experimental data.

An experiment can lead to observations that support or disprove a hypothesis.

A theory is tested by more experiments and modified if necessary.

Scientific Law A scientific law summarizes the results of many observations and experiments.

How do scientific theories develop? One of the most important theories in chemistry is the atomic theory of matter, which states that all matter is made of very tiny particles called atoms. Today, this theory is universally accepted. Scientists have explained the structure of atoms and identified the forces that hold atoms together. With the aid of very powerful microscopes, scientists now can photograph and manipulate individual atoms. Yet similar to other theories, the atomic theory of matter developed over time. The origin of the theory can be traced to ancient Greece. Democritus, a Greek philosopher, proposed that matter was made of tiny indivisible particles. The word atom comes from the Greek word meaning “indivisible.” While Democritus’ arguments for atoms had merit, Aristotle and other leading Greek philosophers rejected them. Hundreds of years passed before scientists seriously considered an atomic theory of matter. Technology is often essential in the development of a new theory or the revision of an existing theory. For example, in the mid-1800s, scientists began experimenting with a device called a cathode ray tube. This led to the discovery of the electron, one of the particles that compose atoms. Further discoveries of atomic structure also depended on new technology, well-constructed experiments, and logical reasoning. TEKS 2D • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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TEKS

End-of-Course Assessment Review

1. Identify  Which of the following statements describes a hypothesis that might be useful for a scientific experiment? A All atoms are made of protons, neutrons, and electrons. B Increasing the surface area of a sample of iron will increase the rate at which it rusts. C The ratio of hydrogen to oxygen atoms is 2:1 for every water molecule. D For ideal gases, the pressure and volume are inversely proportional. 2. Evaluate  Which of the following statements is the best definition of a scientific hypothesis? A a suspicion or hunch about an event in nature B a proposal that can be tested in an experiment C an idea or explanation that most people agree with D a well-supported explanation for a broad set of observations 3. Distinguish  Suppose you were to read about a scientific statement based on data from hundreds of years of research and observation that applies to a broad set of naturally occurring events. Would you consider it a hypothesis or a theory? A a hypothesis B a theory C both D neither 4. Distinguish  Suppose that your lab partner states that he has just come up with a theory that explains the results of your last investigation. What are three reasons you could give him that his explanation is a hypothesis and not a theory?

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TEKS REVIEW

2E

Planning and Implementing Investigative Procedures TEKS 2E Plan and implement investigative procedures, including asking questions, formulating testable hypotheses, and selecting equipment and technology, including graphing calculators, computers and probes, sufficient scientific glassware such as beakers, Erlenmeyer flasks, pipettes, graduated cylinders, volumetric flasks, safety goggles, and burettes, electronic balances, and an adequate supply of consumable chemicals.

How do you plan investigative procedures? When planning investigative procedures, members of the scientific community use the same scientific methodology that you use in a lab setting. Scientific methodology usually includes the following steps: 1. Making observations 2. Asking questions 3. Formulating a hypothesis 4. Testing a hypothesis These steps represent a logical approach to explaining phenomena in the natural world and are a useful tool for planning scientific investigations. When planning investigations, it is not always necessary to follow the steps in order or even to use all the steps. Scientific methodology provides guidelines for all kinds of scientific inquiry.

Why is it important to ask questions about observations? Observations are always a part of scientific investigative procedures. An observation often leads to another step in scientific methodology: asking questions. For example, a scientist might observe that the same hen can make different vocalizations. When following investigative procedures, scientists ask questions about their observations. If a scientist notices a phenomenon that is unexplained, he or she may ask a question about it. The observation that hens have different vocalizations can lead to the question of “What do the different vocalizations of a hen mean?”

Why is it important to formulate a testable hypothesis? Scientists formulate testable hypotheses to test the answers to the questions that have arisen from their observations. These hypotheses form the basis of investigative procedures. A hypothesis is a reasonable explanation as to why something occurs. In order to be used in investigative procedures, a hypothesis must be testable. It must be able to be supported or refuted by evidence. TEKS 2E • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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Testable hypotheses contain clear wording and clearly state what happens to the dependent variable (the phenomenon being observed) when the independent variable (the factor that changes in the experiment) is changed. A testable hypothesis may follow this pattern: “If [independent variable] happens, then [dependent variable] happens.” In the question of hen vocalizations, a testable hypothesis could be “If a hen makes a cluck-cluck-cluck vocalization, her chicks come to feed.” This hypothesis is testable because a scientist can listen for the sound and then observe the chicks’ behavior. If they come and feed only when they hear the cluck-cluck-cluck vocalization, the hypothesis is supported. The independent variable is the cluck-cluck-cluck sound; the dependent variable is the chicks’ behavior. If the hen makes a different sound that produces the same outcome (the chicks running to feed), the hypothesis is not supported.

How do scientists implement procedures to test hypotheses? Scientists implement procedures to test their hypotheses by setting up experimental conditions that enable the dependent variable to be observed while the independent variable is changed. In the case of the hen and chicks, a scientist could implement a procedure that would involve listening for the hen’s cluck-cluck-cluck and observing the chicks’ response to it. Over time, and by carefully eliminating the possibility of other explanations, the hypothesis may be accepted and be used by others to predict when chicks will come to eat.

How can you select equipment and technology for an investigation? An investigative procedure needs the appropriate equipment and technology to be conducted properly and accurately. Laboratory equipment must meet standards for both safety and accuracy. Some equipment and technology helps in taking measurements. An electronic balance, for example, can measure mass quickly and accurately. A graduated cylinder provides quick and accurate measurements for the volume of liquids. This equipment is also reliable, meaning repeated measurements of the same subject will yield very similar results.

Study Tip Smaller quantities will require smaller units; larger quantities require larger units. As you are planning an experiment, consider the quantities you are measuring and choose equipment that is appropriately scaled with the units you need.

As you plan a scientific investigation, make a list of all the equipment and technology you will need. You may need to revise your list to use only the equipment available to you, and to use this equipment efficiently. For example, instead of using three volumetric flasks to measure the same volume of a liquid, you could use the same flask three times. Perhaps you can reduce your use of expensive chemicals, or replace them with less expensive ones. Computers and graphing calculators can help you store and analyze data. If your investigation will generate a large amount of numerical data, consider entering the data directly into a computer or graphing calculator. Your teacher can instruct you on how to use these pieces of technology.

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In the chemistry lab, remember to use only the scientific equipment that your teacher provides. Always follow your teacher’s instructions on the care and use of laboratory equipment.

TEKS

End-of-Course Assessment Review

1. Identify  Which of these would you use to measure 0.2 mL? A a pipette B a beaker C a measuring spoon D a volumetric flask 2. Explain  Why is it important to use clear wording and clearly stated outcomes in a hypothesis? A Hypotheses are not facts or scientific theories, so they need to be explained clearly. B Hypotheses have to be clear to explain natural phenomena to other scientists. C Hypotheses have to be clearly expressed so one can determine if an experiment confirms or refutes them. D Hypotheses may become part of a scientific theory, so it is important to write them well. 3. Evaluate  Which of these hypotheses is most easily tested? A If plants are overwatered, they die. B If the weather is cold, people get sick. C If 2 g ammonium nitrate is added to 2 L of water, the temperature of the water will decrease by 1°C. D If 2 g baking soda is added to a solution of water and vinegar, a chemical reaction happens. 4. Plan  A procedure requires you to test the pH of 500 mL of an unknown solution your instructor gives you by dropping phenolphthalein in 0.5 mL increments. List the materials and equipment you need for this procedure. Explain how you would use each item. 5. Plan  Formulate a hypothesis and design an experiment to determine how the amount of available sunlight affects the size of peppers growing on a plant. Determine your independent and dependent variables and what kind of equipment you will need.

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TEKS REVIEW

2F

Accuracy and Precision TEKS 2F

Collect data and make measurements with accuracy and precision. TEKS_TXT

What is accuracy?

Vocabulary accuracy precision

Suppose two students are performing the same experiment, but are working separately. One step in the procedure requires them to make a 10% salt solution. To make the solution, they each need to mix 10.0 g of salt in 90.0 mL of water. Using a balance, Student A measures 10.4348 g of salt for her solution. Student B measures 10.1927 g of salt for his solution. Which measurement is more accurate? The accuracy of a measurement identifies how close that measurement is to an accepted or true value. If the true value of salt needed for the experiment is 10.0 g, then Student B’s measurement is more accurate. Student B’s measurement of 10.1927 g is closer to the value of 10.0 g than Student A’s measurement of 10.4348 g. Collecting accurate data is important, because if the data are accurate, they reflect what was expected or what truly happened. With accurate data, one can be more confident in the reliability of the results. Because Student B prepared a more accurate 10% salt solution than did Student A, the results of Student B’s experiment might be more reliable than Student A’s results. To achieve reliable results, scientists need to make accurate measurements. For example, a hospital pharmacist preparing a solution must accurately measure the amounts of compounds needed to prevent over- or undermedicating a patient. A marine biologist who dives deep into the ocean needs to accurately measure the amount of oxygen she needs in her tank to remain safe. Scientists launching a satellite into space have to launch the satellite at an accurate velocity to achieve the preferred orbit of the satellite.

Study Tip When trying to remember the difference between accuracy and precision, keep in mind that an accurate measurement is close to the accepted or actual measurement.

What is precision? Precision refers to how much a series of measurements varies. In other words, the more precise a series of measurements, the closer they are to one another. Now, suppose Student A and Student B decide to repeat the experiment two more times to confirm their results. They obtain the measurements shown in Figure 1 on the next page. Whose measurements are more precise?

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Figure 1

Student A

Student B

Trial 1

10.4348 g of salt

10.1927 g of salt

Trial 2

10.4127 g of salt

10.3281 g of salt

Trial 3

10.4821 g of salt

9.8233 g of salt

Actual value of salt needed: 10.0 g

Because Student A’s three measurements are closer in value to one another than Student B’s measurements, Student A’s measurements are more precise. However, because Student B’s measurements overall are closer to the actual value of 10.0 g, they are more accurate than Student A’s.

Why are accuracy and precision important in data collection? The goal of any scientific investigation is to obtain measurements that are both accurate and precise. Knowing the degree of accuracy and precision for a given measurement or group of measurements is important in science. The accuracy and precision of a measurement depend on the units used in a measurement and the number of significant digits stated. For example, if you want to know the average distance from Earth to the sun, you might search the Internet to find the data. One source tells you that the average distance is 150,000,000 km. A second source tells you that the average distance is 149,597,870 km. The second measurement is more accurate because it is closer to the accepted average distance of 149,597,870.691 km. The second measurement is also more precise because it includes a greater number of significant digits. The accuracy of a measurement is related to the instrument used. Typically, an instrument that measures in smaller units or smaller increments can provide more accurate results than a tool that measures in larger units or larger increments. For example, suppose you are trying to measure 10 mL of hydrochloric acid. A graduated cylinder that measures to the nearest milliliter would allow you to measure 10 mL more accurately than would a graduated cylinder that measures to the nearest 10 milliliters. The degree of precision of a measurement is also related to the instrument used. For example, suppose a chemist extracts a salt from a solution in the lab. When determining the mass of the extracted salt, he uses a balance that measures to the nearest gram. He determines that the mass of the salt is 2 g. Another chemist, who is working on the same experiment, uses a balance that measures to the nearest ten thousandth. She reports that she extracted 2.3484 g of salt from her solution. Which chemist extracted more salt from his or her solution? In this situation, it is impossible to determine. Because the first chemist’s balance measures only to the nearest gram, any mass from 1.50 g to 2.49 g will result in a reading of “2 g.” If the first chemist uses an instrument that measures mass to the same degree of precision as the second chemist, then the two masses can be compared. When comparing measurements, the degree of precision of the tool used can greatly influence the results. TEKS 2F • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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TEKS

End-of-Course Assessment Review

1. Identify  The accepted value for the boiling point of water is 100˚C. During an experiment, students recorded the temperature observations listed below. Which one is the most accurate temperature for the boiling point of water? A 102˚C C 95˚C B 97˚C

D 103˚C

2. Infer  Two students need to determine the density of iron for part of an experiment. Using the information in the table below, calculate the average of each student’s trials. Which student has the most precise set of measurements? Which student’s average is the most accurate? Sarah

Felipe

Trial 1

7.96 g/cm

3

7.24 g/cm3

Trial 2

8.05 g/cm3

7.75 g/cm3

Trial 3

7.92 g/cm3

8.57 g/cm3

Accepted value of the density of iron = 7.86 g/cm3

A Sarah’s measurements are more precise and her average is more accurate. B Felipe’s measurements are more precise and his average is more accurate. C Sarah’s measurements are more precise but Felipe’s average is more accurate. D Felipe’s measurements are more precise but Sarah’s average is more accurate. 3. Explain  Suppose you need to measure 2.25 mL of water. There are two different graduated cylinders, shown at right, that you can use to measure the water. Using the terms precise and accurate, explain why you would choose graduated cylinder a or b. All measurement markings shown are in milliliters.

15

3

10

5

a

2

b

4. Describe  Suppose you work at a theme park. Your supervisor wants you to make a sign displaying the maximum weight that a roller coaster train can carry. Your supervisor knows that the maximum weight is 1686.5 kg. However, he wants the sign to be quickly understood and tells you to make a sign that says: Maximum Weight 1700 kg. How could the lack of precision in this example cause problems? TEKS 2F • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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TEKS REVIEW

2G

Dimensional Analysis, Scientific Notation, and Significant Figures TEKS 2G Express and manipulate chemical quantities using scientific conventions and mathematical procedures, including dimensional analysis, scientific notation, and significant figures.

How does dimensional analysis help you express and manipulate chemical quantities?

Vocabulary dimensional analysis scientific notation significant figure

Dimensional analysis involves using ratios called conversion factors. Multiplying by the appropriate conversion factor lets you convert a quantity from one unit to another. Every conversion factor is numerically equal to 1.

Sample Problem For example, to convert between kilograms and grams, begin by writing the equation:

Use the first conversion factor to convert from grams to kilograms. For example: 600 grams ×

1 kilogram = 1000 grams This equation leads to two conversion factors: 1 kilogram 1000 grams and 1000 grams  1 kilogram 

1 kilogram = 0.6 kilograms 1000 grams

Notice that the unit gram cancels out of the expression. To convert from kilograms to grams, use the conversion factor that has grams in the numerator and kilograms in the denominator. 1000 grams 7.5 kilograms × = 7500 grams 1 kilogram

How does scientific notation help you express and manipulate quantities? Scientific notation is used to express very large numbers. A quantity expressed in scientific notation is in the format of a number multiplied by 10 raised to a power. For example, the average distance between Jupiter and the sun is 778,300,000 km. In scientific notation, this measurement is 7.783 × 108 km. In scientific notation, the coefficient (the number being multiplied) must be equal to or greater than 1 and less than 10. For numbers greater than 10, the exponent is positive and is equal to the number of places the original decimal point was moved to the left. For example, light travels through empty space at a speed of 1,079,000,000 kilometers per hour. To convert this number to scientific notation, begin by moving the decimal point to the right of the leftmost digit: 1.079,000,000. TEKS 2G • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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Study Tip As you perform calculations, keep track of an excess of significant figures. Then round off the answer to the correct number of significant figures.

Then count how many places the decimal moved. Since the decimal moved 9 places to the left, you raise 10 to that power (109). The speed of light expressed in scientific notation is 1.079 × 109 kilometers per hour. For numbers less than 1, the exponent is negative and equals the number of places the original decimal point was moved to the right. The radius of a helium atom is 0.000000000031 meters. To express this measurement in scientific notation, move the decimal point 11 places so that it is just to the right of the 3. This gives the result of 3.1 × 10–11 meters.

How do significant figures help you express and manipulate quantities? When you measure something, there is always a degree of uncertainty in the measurement. For this reason, the last digit is considered an estimate. This digit represents the degree of precision of the measuring tool. For example, depending on the ruler that is used, the width of the door shown in Figure 1 might be reported as 0.8 m, 0.77 m, or 0.772 m. Figure 1

0.8 m 1m

0.77 m 10

20

30

40

50

60

70

80

90

1m

0.772 m 10

20

30

40

50

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1m

Significant figures are the digits that represent the precision of a measurement. For example, the population of Tyler County, Texas, might be reported to be 21,000 people, 20,600 people, or 20,556 people. The first measurement has two significant figures (2 and 1), the second measurement has three significant figures (2, 0, and 6), and the third measurement has five significant figures. The following are rules for identifying significant figures in measurements. 1. All nonzero digits are significant. The numbers 543, 54.3, and 0.543 each have three significant figures. 2. All zeros that appear between nonzeros are significant. The numbers 2034, 2.034, and 20.34 each have four significant figures.

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3. Z  eros at the end of a number and to the right of a decimal point are significant. The numbers 23.00, 2.040, and 4.000 each have four significant figures. 4. C  ounted numbers are significant. If you count 24 blossoms on a tomato plant, your measurement has two significant figures. Figures that are not significant include the zeros that precede or follow the significant digits up to the decimal point. The numbers 0.0000013, 0.20, and 0.0034 each have two significant figures. The numbers 340, 520, and 72,000 also have two significant figures. To clearly show significant figures, you can use scientific notation. The numbers 3.4 × 102, 5.2 × 10–3, and 7.0 × 104 each have two significant figures. Notice that the powers of 10 are not counted as significant.

TEKS

End-of-Course Assessment Review

1. Evaluate  Multiplying by which conversion factor would allow you to convert 35 liters to milliliters? A

1000 milliters 1 liter

B

1 liter 1000 milliliters

C

1000 liter 1 milliliter

D

1 liter 1 milliliter

2. Apply  Which of these measurements is expressed with the most significant figures? A 0.000342 grams B 1.5 × 104 light years C 150 liters D 0.20008 kilometers 3. Calculate  Determine the number of seconds in 4.3 years. Express the answer in scientific notation and with the proper number of significant digits.

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TEKS REVIEW

2H

Data Analysis, Inferences, and Predictions TEKS 2H

Organize, analyze, evaluate, make inferences, and predict trends from data.

How do scientists organize, analyze, and evaluate data?

Vocabulary inference

Scientists often collect a large amount of data during an investigation or experiment. Organizing the data in a useful way allows scientists to analyze and evaluate the data. Scientists often organize numerical data in graphs, charts, or tables. These methods help show patterns or trends in the data that might otherwise be difficult to observe. Scientists also analyze numerical data with concepts and formulas from statistics, the branch of mathematics that deals with numerical data.

How do scientists use data to make inferences and predict trends? An inference is a logical interpretation based on prior knowledge and experiences. For example, look at Figure 1. Of the eleven plotted data points, three fall outside the normal pattern. These points often indicate an error in the experimental procedure or design, or an error in measurements or calculations. Yet the data show a trend that can be used to make inferences. For example, although a pH of 7 was not measured, it can be inferred that the yield would be a little less than 20 percent. And, although a pH of 12 was not measured, but it can be inferred that the percent yield would be about 30 percent. Figure 1

Percent Yield of a Reaction vs. pH 30

Percent Yield

25 20 15 10 5 0

0

2

4

6

8

10

12

pH TEKS 2H • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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Graphing is a helpful way to predict trends in data. Trends in data graphs may be linear, meaning the data points fall on a straight line; or they may fall on a curved line. Sometimes a graph shows no clear trend in data. Figure 2 shows several types of graphs that can be used to make predictions. Figure 2 Various Types of Graphs

Population Growth 80

5

Bacterial Cells

Number of Baskets Made

Baskets and Distance 4 3 2 1 0

1

2

3

4

5

Distance From Hoop (m)

20

6

0 20 40

60 80 100 120

Time (min)

This graph shows a linear trend. Fewer baskets are made as the player moves farther from the hoop.

The rising curve in this graph shows exponential growth. Every twenty minutes, the population of bacteria becomes twice as large.

Seasonal Rainfall

Hours of TV per Day 5

Hours of TV

Rainfall (cm)

8 6 4 2 0 Jan.

4 3 2 1 0

June

Jan.

June

Jan.

Month This graph shows a cyclical trend in data. Rainfall amount changes in the same way from year to year.

When reading graphs, ask yourself the following questions: What information is contained in the graph? What are the variables? What happens to one variable as the other variable changes?

40

0 0

10

Study Tip

60

0

2

4

6

8

10 12 14

Age (years)

These data points show no trend. The data are not useful for making predictions.

Consider the following data that were obtained from an investigation on the time it takes different volumes of water to boil. A scientist placed several different samples of water on a hot plate, then measured the time each sample needed to boil. Graphing the data can reveal trends for making inferences. Figure 3 shows the four data points as small circles. It also shows a line that connects the data points. Joining the data points by a line shows the “best fit.” The graph shows a linear trend related to the volume of water and the boiling time. As the volume of water increases, the time it takes the water to boil also increases. You can use the line to predict that 1200 mL of water will boil in about 20 minutes, or that 300 mL of water will boil in about 5 minutes.

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Volume of Water (mL)

Time to Boil (min)

500 1000 1500 2000

7.8 16.6 26.0 33.7

Boiling Time of Water by Volume 40

Boiling Time (min)

Figure 3 Graphing Data

30 20 10 0

0

500

1000

1500

2000

Volume of Water (mL)

TEKS

End-of-Course Assessment Review

1. Evaluate  A chemist is studying the boiling points of 12 related compounds. She reports that the average boiling point is 65.5˚C. What can be concluded from this information? A If any compound boils at a temperature greater than 65.5˚C, then at least one compound must boil at a temperature less than 65.5˚C. B All 12 compounds boil at a temperature of 65.5˚C. C None of the 12 compounds boil at a temperature of 65.5˚C. D A boiling point of 65.5˚C is the most common boiling point among the 12 compounds. 2. Analyze  A chemist sets up the same chemical reaction eight times. In seven of the eight trials, 30 to 32 g of product is obtained. In one of the eight trials, 78 g of product is obtained. How should the chemist analyze these data? A calculate the average amount of product using the data from all eight trials B exclude the outlier from the report C study the details of the trial with the outlier result and look for possible errors in the procedure, measurements, and calculations D conclude that on rare occasions, more than double the typical amount of product can be obtained 3. Make Graphs  Plot points and draw lines of best fit to represent the following set of data: the days per month that high school students spend in school over a period of three years.

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TEKS REVIEW

2I

Communicating Conclusions Based on Scientific Data TEKS 2I Communicate valid conclusions supported by the data through methods such as lab reports, labeled drawings, graphs, journals, summaries, oral reports, and technologybased reports.

How do scientists communicate conclusions from their investigations? Scientists use many different methods to communicate valid conclusions of their research. These methods include lab reports, labeled drawings, graphs, journals, summaries, oral reports, and technology-based reports. Each method has benefits and drawbacks, and scientists may use a combination of several methods to present their conclusions.

Study Tip Review the lab reports you have recently produced. Note the different ways you communicated results based on the type of experiment or investigation you performed.

Lab Reports  Scientists use lab reports to communicate the procedure and results of an experiment. A lab report includes all the information that another scientist would need to repeat the experiment. It includes the hypothesis, list of materials and equipment, important background information, procedure, observations, and conclusions. It may also include graphs or other representations of data. Labeled Drawings  Labeled drawings may help communicate the procedure, observations, or conclusions of an experiment. These drawings help explain findings that may be difficult to communicate with only words. A drawing might show information such as the arrangement of laboratory equipment, a proposed structure of a chemical compound, or the pathway of a series of chemical reactions. Graphs  Scientists often use graphs to communicate conclusions that involve a trend in numerical data. Trends are typically easier to see in a graph than in a chart or in a written description of data. Common types of graphs include bar graphs, line graphs, and circle graphs. As shown in Figure 1, line graphs show how the dependent variable changes in response to the independent variable. Bar graphs are used to display data in a number of separate, or distinct, categories. Circle graphs (sometimes called pie charts) are especially useful when the set of data represents a whole, or 100 percent. Each section of the circle represents a particular component of the whole. Journals  Science journals allow scientists to communicate their conclusions to the scientific community. When scientists have completed an investigation, they may present their conclusions as part of an article or report. Then they submit the article or report for publication in a journal that is appropriate to their field. TEKS 2I • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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Plastics Packaging in U.S. Waste

Mass of Elements in Samples

15,000

20

15

10

5

0

13,630

12,500

Waste produced (thousands of tons)

Mass of Iron (g)

Figure 1 Types of Graphs

0

2

4

6

8

11,190 10,000 7500 5000 3400 2500 0

Mass of Oxygen (g)

6900

2090 120 1960

1970

1980

1990

Year

2000

2007

Before an article or report is published in a science journal, it undergoes peer review by independent scientists. Peer review is a process in which scientists carefully review the work of other scientists. The article is published only if the reviewers are convinced that the findings are genuine. When scientific reports are published, they become available for other scientists to read, evaluate, and test. These reports enable the scientific community to keep informed about the work of other scientists and research groups. Summaries  A summary is a brief restatement of the purpose, procedure, and findings of an experiment. Scientists use summaries to communicate with the general public, their employers or the people providing their funding, and other scientists. In science journals, summaries are called abstracts. They are typically located at the beginning of an article and are designed to give the reader a general idea of what the article is about. Oral Reports  In an oral report, the scientist speaks directly to an audience about the results of an investigation. Oral reports may be formal and follow the format of a journal article. Or an oral report could be more conversational. Many scientists use visual aids to accompany their presentations, such as projections from computers. An oral report allows the presenter to interact with the audience. The audience may ask questions about the conclusions drawn from the experiment. Oral reports are especially useful for venues that showcase multiple experiments, such as science conferences or school science fairs. Technology-Based Reports  With a technology-based report, scientists can use computers and other types of technology to communicate conclusions. By using appropriate software, scientists can illustrate their conclusions with diagrams, animations, photographs, and narration or other audio components.

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TEKS

End-of-Course Assessment Review

1. Identify  Which section of a journal article introduces the purpose of the article? A procedures C abstract B conclusion

D data tables and graphs

2. Analyze  What is an advantage of an oral report compared to other methods of communicating scientific conclusions? A The scientist can communicate directly with an audience and respond to their questions and ideas. B The scientist can precisely describe the hypothesis, procedure, and data obtained from an investigation. C The scientist can present graphs and labeled diagrams to show trends in the data. D The scientist can have conclusions accepted without rigorous peer review. 3. Explain  Scientists may submit reports of their investigations to science journals for publication. How do science journals ensure that the reports are accurate? A Science journals publish only reports of experiments conducted in laboratories. B Science journals accept only reports from leading, famous scientists. C Independent scientists review and evaluate the reports, and only approved reports are published. D Science journals publish reports after they are published in magazines or newspapers. 4. Evaluate  Which set of data would be best represented by a graph? A the pathway of chemical reactions that the body uses to break apart sugars B the atmospheric pressure when an experiment was conducted C a scientist’s observations of ice cubes at their melting point D the percentage by mass of each element in the human body

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TEKS REVIEW

3A

Analyzing, Evaluating, and ­Critiquing Scientific Explanations TEKS 3A In all fields of science, analyze, evaluate, and critique scientific explanations by using empirical evidence, logical reasoning, and experimental and observational testing, including examining all sides of scientific evidence of those scientific explanations, so as to encourage critical thinking by the student.

How can empirical evidence be used to analyze, evaluate, and critique scientific explanations? When you analyze a scientific explanation, you determine what claim is being made and the evidence that is being used to support the claim. Often this evidence is in the form of empirical evidence, which is data based on observation or experience—that is, scientific experiments. Evaluating an explanation involves determining whether it is valid or useful. Part of the evaluation process is to review the evidence that supports the explanation. Empirical evidence needs to be analyzed and evaluated carefully, and it may or may not agree with other types of evidence. For example, consider a laboratory experiment in which scientists test the effect of soil pH on the growth of soybeans. The empirical evidence from the experiment suggests that soybeans grow best at a pH of 6.5. However, a group of farmers report that their soybeans grow poorly at this pH. How can this difference be explained? Does this mean that the scientists are incorrect, or that the experiment was badly designed? No, this is not necessarily the case. The conditions that were used in the laboratory experiment may have differed from those in the farmers’ fields. The empirical evidence from the experiment is not incorrect. Rather, more data are needed to construct a logical explanation. In their next experiment, the scientists compared how soybeans grow in different types of soil, while keeping pH constant. To evaluate how well the soybeans grow, they measure the concentration of manganese (Mn) in the leaves. Manganese is a mineral that plants need in small amounts and that scientists can use to measure overall plant health. The table in Figure 1 below shows how manganese concentration in leaves varies according to the type of soil the plants are growing in. Figure 1

Soil Type

pH

Mn in Leaves (ppm)

Clay-rich soil Sandy soil

6.5 6.5

210 25

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With this added empirical evidence, the scientists can construct a more useful explanation for why the soybeans in the farmers’ fields did not grow well: The soybeans growing in the farmers’ fields, which contained sandy soil, grew poorly because there was not enough manganese available in the soil for plant uptake; the manganese concentration in the plants grown in sandy soil was too deficient to support proper growth. As with all scientific studies, this study is subject to critique. When you critique a scientific explanation, you identify specific problems with the scientific investigation or with the claims that the scientist is making based on the data. In this example, a useful critique is that the first study was not incorrect, but it was incomplete—more evidence was necessary to construct an explanation.

Study Tip Remember that a scientific explanation should explain more than merely a single set of observations from an experiment. A useful explanation has the power to predict future events.

How can logical reasoning be used to evaluate scientific explanations? In science, logical reasoning involves drawing valid conclusions or formulating ideas based on empirical evidence. As part of the experiment on soybean growth, the researchers measured the amount of manganese dissolved in the soil at different pH values. They found that the amount of soluble manganese in soil decreases as pH increases. From this evidence, is it logical to conclude that soil pH can change the amount of soluble manganese available in the soil? Yes. Is it logical to conclude that lower soil pH will increase soybean growth? No, not from this evidence alone.

How can experimental and observational testing be used to evaluate scientific explanations? Experimental testing involves testing the effect of a variable while keeping other factors the same between the groups you are testing.. In observational testing, variables are not directly tested. Instead, situations or models are directly observed and the observations are recorded as evidence. Both types of testing can provide useful evidence for different scientific explanations. In the example of soybean growth, the experimental testing in the laboratory suggested a relationship among soil pH, soil manganese concentration, and the health of soybeans. The observational testing in the farmers’ fields helped show that soil pH alone did not explain the health of soybeans.

How can all sides of scientific evidence be analyzed? Sometimes, two sets of evidence may seem to disagree or contradict with one another. In the example of soybean growth, the ideal soil pH measured in the laboratory experiment disagreed with the results observed in the farmer’s fields. Part of the scientist’s job is to make sense of all sides of scientific evidence. In this example, evidence from additional experiments helped formulate a useful explanation.

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In some cases, different scientists explain the same set of evidence in different ways. This is especially true when many variables could affect an outcome. For example, the health of soybeans could be affected by different plowing or planting techniques, quality of the seeds, trace minerals in the soil, and the weather. A scientist might cite any or all of these factors to explain poor soybean growth. When scientists propose many explanations for the same observation, how can we decide which explanation is best? It is always important to evaluate how evidence was obtained and whether the evidence is reliable. Evidence gathered from properly run scientific investigations—those applying scientific methodology to obtain precise and accurate data—is more likely to be reliable than reports from untrained observers. Yet even scientists are subject to biases and personal opinions that may affect the results, which must be taken into consideration. Ultimately, however, the power of a scientific explanation depends on its ability to predict future events. If a scientist claims that low manganese levels in the soil are the cause of poorly growing soybeans, then increasing these levels should improve the soybean crop.

TEKS

End-of-Course Assessment Review

1. Analyze  Which of the following describes empirical evidence? A data from surveys and questionnaires B opinions from leading scientists C measurements and observations from scientific experiments D predictions from accepted scientific theories 2. Evaluate  Which statement best explains why scientists often propose different explanations for the same event in nature? A Many variables may affect an event, and not all variables are easy to test in laboratory experiments. B Scientists often obtain different results from the same laboratory experiment. C Scientists often ignore data that disagree with the data they obtain. D Events in nature occur randomly and cannot be explained by the methods of science. 3. Critique  A brand of laundry detergent is marketed with the slogan “Keeps your red clothes red and your white clothes white—even in the same wash!” What does this slogan suggest? How could you ­scientifically evaluate and critique the accuracy of the slogan?

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TEKS REVIEW

3B

Communicating and Applying ­Scientific Information TEKS 3B Communicate and apply scientific information extracted from various sources such as TEKS_TXT current events, news reports, published journal articles, and marketing materials.

Why is scientific communication important? Whenever you talk on the phone, text someone, or listen to your teacher at school, you are communicating. Communicating is the process of sharing ideas and information with other people. The way that scientists communicate with each other and with the public has changed over the centuries. In earlier centuries, scientists exchanged ideas through letters. They also formed societies to discuss the latest work of their members. When societies began to publish journals, scientists used the journals to publish the results of their research and to keep up with new discoveries. Today, the Internet has become both a means of communication and a major source of information. One advantage of the Internet is that anyone with a computer can access the information. One disadvantage is that anyone can post information on the Internet without first having that information reviewed. To judge the reliability of information you find there, you have to consider the source. This same advice applies to articles in journals, magazines, newspapers, or the news you encounter on television or radio. If a media outlet has a reporter who specializes in science, the chances are better that a report will be more scientifically accurate.

How can you evaluate scientific information from various sources? In science, you often need to do research to learn more about a particular topic and communicate your findings. Therefore, you need relevant, reliable information. Information qualifies as reliable if it comes from a person or organization that is reputable in a particular field and not biased. Generally, universities, museums, and government agencies are good sources of reliable information. Personal blogs and politically motivated news reports are not reliable sources of information. Current Events and News Sources  News and entertainment media often report scientific discoveries. Sometimes reports are based on articles in scientific journals or interviews with scientists. In other cases, the reports are based on statements from companies, the government, or universities.

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Study Tip News reports, science journals, and marketing publications present scientific information in different ways and for different purposes. Remember these differences when you evaluate the scientific information you read or hear.

People use science-based stories in the news to help them make decisions. By applying the information they learn, people might decide to try a new medicine, to purchase a new product, to revise safety practices at work or home, or to change their diet or exercise habits. Not all reports of science news present unbiased scientific information. Sometimes, information is deliberately misrepresented, or mixed up to present a certain point of view. When you hear science news, consider the following questions: • Does the news report directly quote a scientific publication or does it summarize the publication? • Could the news medium or reporter have biases that could affect the way the information is being presented? • Does the report include conclusions or comments from more than one scientist? (Remember that scientists may draw different conclusions from the same data.) Basically, you need to ask yourself “Am I hearing a scientist’s words, or the version of the scientist’s words someone else wants me to hear?” Marketing Materials  Marketing materials include commercials, advertisements, brochures, posters, and product packaging. A wide variety of products are marketed with scientific information or claims based on scientific information. Laws dictate that companies cannot make false claims or invent scientific data to support a claim. Yet even when marketing claims contain accurate scientific information, they can still be misleading. Consider this marketing logo on a floor cleaner box:

Floorbrite will make your floors sparkle! Our scientists have mixed borax with fresh lemon scent and TLC to make the best floor cleaner you can buy. In test after test, users agree that Floorbrite beat competing brands in cleaning away stains, dirt, and grime.

Like other marketing materials, these claims present both scientific information, such as the chemicals in the cleaning product, and nonscientific information, such as the claim about “TLC.” The information about the tests is presented as scientific information, but it could be misleading. Were the tests completely unbiased? How many tests were conducted? How many users agreed that Floorbrite was the best product? How was “competing” defined? Before communicating or applying scientific information from marketing materials, you always need to ask yourself such questions.

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Scientific Journals  Scientists publish the results of their investigations in scientific journals. Typically, the reports are very detailed and technical, and they are intended for other scientists. Journal articles communicate much more information than stories in media that are meant for the general public. Journal articles are also subject to peer reviews, during which other scientists in the same field evaluate them. Peer reviews help ensure that journal articles present accurate, reliable information. To approach a journal article, begin by reading the abstract, which is a summary of the article. It briefly describes the purpose of the investigation, the methods, and the conclusion. You can use the abstract to decide if the journal article will be useful to you. Other sections of the journal article describe the investigation in more detail. You can apply scientific information from journal articles in many ways. Journal articles might help you develop a hypothesis or procedure for a new investigation. They might help you better understand a topic in your scientific studies. Or they might raise questions or help you understand issues in your daily life.

TEKS

End-of-Course Assessment Review

1. Evaluate  On an infomercial, the presenter is selling a new ­fertilizer. According to the presenter, scientific data show that the fertilizer doubles the speed of plant growth. Which activity would best help you evaluate the presenter’s claim? A Reading marketing materials for the new fertilizer. B Reading the journal article in which scientists report their plant growth findings. C Reading a newspaper or magazine article about the fertilizer. D Discussing the fertilizer with friends who have gardens. 2. Apply Concepts  Which of the following statements best describes the scientific information in marketing materials? A The information is usually inaccurate because the company is trying to sell its products. B The information may be accurate or inaccurate because no laws insist on accuracy. C The information is accurate but is often invented. D The information is accurate but sometimes misleading. 3. Apply Concepts  How do people apply the scientific information that they learn from newspapers, magazines, and other sources? Why should people carefully evaluate this information?

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TEKS REVIEW

3C

Drawing Inferences From ­Promotional Materials TEKS 3C Draw inferences based on data related to promotional materials for products and services.

What is an inference? You may recall that an inference is a logical interpretation based on prior knowledge and experience. For example, by observing the trees in your neighborhood, you might draw inferences about the shapes of trees and the maximum height that trees can grow to.

How can you draw inferences based on data related to promotional materials? Companies publish a variety of materials to help promote the products and services they offer. They hope that you will purchase these goods or services based on the data in these promotional materials. Promotional materials take many forms, including advertisements, Web sites, and packaging. You encounter promotional materials everywhere: at home, on television, at shopping malls, even at school.

Study Tip Remember that a promotional claim can be accurate yet still be misleading. When evaluating a claim, decide whether or not data can support it. Then evaluate the reliability of the data.

In many cases, promotional materials make scientific claims which are based on evidence from scientific investigations. Many companies are required to publish this evidence on the packaging of their products. Food companies, for example, must publish specific information for a food product on nutrition labels. Drug companies must publish the chemicals contained in the drugs, drug-safety data, and other information. For many reasons, consumers should be wary of advertisers’ claims. For example, consider these advertising claims for Milan Ice Water, a fictitious brand of bottled water: • Milan Ice Water is the most environmentally responsible consumer product in the world. • Our bottles are made of 75 percent recycled plastic and are more ecofriendly than our competitors. • Milan Ice Water is the Healthy, Eco-Friendly Choice. How can you evaluate these claims? One way is to decide whether they can be supported by data, and then to analyze that data. The first and third claims, even if they have merit, are not scientifically testable because they are too vague. When a company makes claims that are impossible to evaluate, you might choose to discount the claims altogether.

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The second claim contains data, so you could evaluate it. You would need to gather more information, however. First, can you verify that their bottles contain 75 percent recycled plastic—will Milan provide these data to consumers? Second, what is the typical recycled plastic content in plastic water bottles—are Milan’s bottles really more “eco-friendly” than others? Even if their bottles are made of 75% recycled plastic, additional information might reveal it as unremarkable. For example, other companies’ bottles may be 95% reycled plastic.

How can you draw inferences based on data related to promotional materials for services? Suppose that one morning you wake up and find your sugar bowl is full of ants. You notice that a flyer has come in your mail from Pest-Be-Gone, a local pest control service. In that flyer is the table in Figure 1, which Pest-Be-Gone is using to show the frequency with which they get repeat business from the same customers, which for most businesses, is a positive statistic. Figure 1

Pest-Be-Gone’s Faithful Customers Company

Customers Returned in One Year

Customers Returned in Two Years

Customers Returned in Five Years

Pest-Be-Gone

75%

50%

5%

No Mo’ Bugs

10%

20%

50%

Bugging Out

15%

20%

40%

But remember this is a pest-control business. While some of Pest-Be-Gone’s repeat business might be from pleased customers who wanted them to remove other pests, if that is the case, why does their business drop off drastically after two years? Could it be they never removed the pests entirely in the first place, and after repeated calls over two years, most customers gave up on them? The data certainly support that inference. And why does the business for the other two pest-control businesses show the opposite pattern? The data support the inferences that those companies removed the original pests, made a favorable impression, and were hired back for increasing amounts of additional work. Despite Pest-Be-Gone’s favorable impression of themselves, you may want to think twice before hiring them.

Why is drawing inferences from promotional materials a useful skill? Companies publish promotional materials to encourage you to buy their products and services. Laws regulate the claims that a company may make, but not the way that people interpret those claims. A company may or may not provide you with all the data you need to properly evaluate their products or services. And as with the Pest-Be-Gone example above, you must always ask yourself if the data actually shows what the company is telling you it shows. TEKS 3C • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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By drawing the proper inferences from promotional materials, you can make wise purchasing decisions. Your choices can help you save money, stay fit and healthy, and buy the quality of products and services that you demand and expect.

TEKS

End-of-Course Assessment Review

1. Infer  Which of the following types of data would best help you evaluate a DVD delivery service? A a table showing average delivery times B a table showing movies available for streaming C a list of upcoming releases D reviews on a competitor’s Web site 2. Identify  Which of these claims or slogans could best be evaluated by analyzing scientific data? A ”Super Bran Cereal has 25 percent more nutritional value than the leading brand.” B “More than 60 percent of students improve their grades by one letter after two months in our tutoring program.” C “Our bottled water contains more hydrogen than the competitor’s brand.” D “The best brand of milk for your body.” 3. Discuss  Why do you think many promotional claims are accurate but can be misleading? 4. Infer  VitaPro Cereal is marketed with the slogan: “The Healthiest Way to Start Your Day.” Included in the promotional materials for the cereal is the data table shown below. The company states that the data was compiled by researchers who work for a consumer testing agency. The five nutrients included in the table are all important parts of a ­balanced diet.

Percent of Minimum Daily Requirements for 4 Leading Cereals, per serving Brand VitaPro Brand A Brand B Brand C

Iron 95 35 2 20

Vitamin B6 100 10 5 25

Vitamin B12 100 10 8 20

Calcium 100 5 12 18

Magnesium 100 15 20 19

What inference can you draw from these data about the nutritional value of VitaPro compared to the other cereals? What two questions would you ask that could help you better evaluate the data and the company’s claims about VitaPro? TEKS 3C • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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TEKS REVIEW

3D

The Impact of Scientific Research TEKS 3D

Evaluate the impact of research on scientific thought, society, and the environment. TEKS_TXT

What is scientific research? Scientific research involves investigation and experimentation. Scientists may have many goals for their research. They may want to discover a pattern or trend in nature, provide evidence in support of an idea or theory, or develop a new technology or product. Chemists conduct research on the composition of matter and the changes that matter undergoes. Sometimes a chemist’s goal is to explain observations made in another branch of science. For example, a chemist may help a botanist investigate how some plants resist drought better than other plants do. Or chemists might conduct research to develop a new chemical, such as a drug, a food additive, or a fiber for a new fabric. Scientific research and its results have had major effects on scientific thought, society, and the environment. Only 150 years ago, the most effective forms of transportation were animal-drawn wagons and wind-powered boats. Most people lived on small farms and grew all their own food. Diseases such as smallpox and tuberculosis killed thousands every year. Scientific research has changed life dramatically in the past 150 years, and it continues to do so every day.

How does research affect scientific thought? The goal of science is to explain observations and processes in the natural world. These explanations are useful because they are based on logical reasoning and evidence. Scientific research provides this evidence. Scientists apply the results of their research to evaluate explanations and to revise or replace explanations when necessary. For example, through the years, scientists developed many models of the atom—the smallest particle of an element that retains its identity in a chemical reaction. Each model improved the explanatory power of the previous model. Today, scientists use the model of the atom in a wide variety of ways. The model helps them predict the structures of compounds and the properties of those compounds. The model also helps them develop new technologies, such as smoke detectors and powerful magnets (like those based on neodymium, an element with strong magnetic properties).

TEKS 3D • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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What is the impact of research on society?

Study Tip Remember that scientists study a wide variety of topics in the natural and physical world. Their research affects everyone’s life.

Scientific research can lead to many changes for a society. Sometimes these changes develop in ways that no one could predict. In the 1860s, for example, billiards players were looking for a new material for billiard balls. Balls made of wood or clay worked poorly. Ivory worked well but was very expensive. In response to this challenge, researchers began experimenting with different types of materials. American inventor John Wesley Hyatt developed celluloid, one of the first materials that we now call plastic. Celluloid proved very useful for making billiard balls and other products, including early motion picture films. (Motion pictures themselves have had a powerful effect on human society.). Today, plastics are used to make a wide variety of products. Thin, flexible plastics are used in food wrap; tough, sturdy plastics are used to make containers and furniture. Perhaps most importantly, plastics are used to make sterile, disposable medical equipment that greatly reduces the transmission of infections between patients. In the 1800s, not many people would have predicted the wide use of plastics today. New research and the technology that results from it also change the worklife of a society. New jobs arise and existing jobs become unnecessary. For example, 100 years ago, workers in the northern United States cut ice from frozen lakes and stored it in warehouses to keep food cool in the summer. The invention of the refrigerator made these jobs unnecessary.

What is the impact of research on the environment? Scientific research has helped identify environmental problems. It also has helped solve or improve many of these problems. But research that solves some problems can cause other problems. One example of the impact of research on the environment is DDT, a commonly used pesticide from the 1940s through the 1960s. In the 1950s and 1960s, scientist Rachel Carson researched the effects of DDT on the environment. Carson’s results showed that DDT was collecting in the bodies of eagles, hawks, and other birds of prey. Because of the accumulation of DDT, the birds were producing eggs with thin, fragile shells, reducing hatching rates and causing a drop in bird populations. Carson’s research led to the ban of DDT in the United States and many other nations. Since then, the populations of eagles and other birds of prey have recovered. The research of others has led to regulations on other toxic substances, including lead, mercury, and asbestos. The invention of plastics has greatly improved human life, but it has also affected the environment. Most plastics are not biodegradable, meaning that natural processes break them apart very slowly, if at all. Because plastic is so widely used, the amount of plastic waste keeps increasing, as you see in Figure 1 (on the next page). More and more space is needed for landfills to store discarded plastic. Researchers continue to study effective ways to recover plastic from consumer waste and to recycle, and responsible manufacturers continue to reduce the amount of plastic in their products and packaging. TEKS 3D • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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Figure 1 15,000

Plastics Packaging in U.S. Waste 13,630

Waste produced (thousands of tons)

12,500 11,190 10,000 7500

6900

5000 3400 2500 0

120 1960

2090 1970

1980

1990

2000

Year

TEKS

2007

End-of-Course Assessment Review

1. Evaluate  A chemist could best help achieve which of the following research goals? A identifying the insects that are ruining a corn crop B processing corn husks into diesel fuel C explaining the rise of the Andes Mountains D designing the most useful shape for a parachute 2. Infer  What do the history of both plastics and DDT demonstrate about the results of scientific research? A Research results can cause unexpected changes to society and the environment. B With enough time, scientists always achieve the goals of a research project. C Research results may have unintended effects, but those effects can be remedied. D Research results cause changes to society that last 50 years at the longest. 3. Evaluate  Write a short paragraph in response to this statement: ­“Scientific research can result in both benefits and drawbacks.” Include a specific example in your answer (other than DDT or plastics).

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TEKS REVIEW

3E

Chemistry and Careers TEKS 3E

Describe the connection between chemistry and future careers.

How can studying chemistry help you in your future career? Chemistry is a growing and profitable field. Every year, billions of dollars are invested in the United States in the research and development of new products and technology. Often, chemists are employed in this research and development. In 2006, according to the American Chemistry Council, chemists were hired for almost 900,000 new jobs in the United States. This number is even greater today. Why are chemistry jobs important? Reading the news will help you answer this question. Common news topics include alternative energy resources, medical research, global environmental problems, and diet and nutrition. In these types of news stories and many others, you can learn about the work of chemists or scientists whose work involves chemistry.

What are some careers that are related to chemistry? Chemistry applies to careers in all branches of science, including life sciences, physics, earth science, and space. It applies to careers in other fields, too. Doctors apply their knowledge of chemistry when they prescribe drugs to patients. Farmers apply chemistry knowledge when they measure the pH of soil and add soil supplements and fertilizers. Real estate brokers apply chemistry knowledge when they evaluate environmental reports. Here are several careers related to chemistry. Environmental Scientist  Environmental scientists study the interactions among humans and the environment. Some environmental scientists may study how to restore polluted land or water. Others may research how we can increase our food supply for the growing human population. These scientists depend on their knowledge of chemistry to achieve their goals. Materials Scientist  Make a list of some useful products whose raw materials are natural resources. Your list might include plastics, paper products, ceramics, glass, fabric dyes, protective waxes and other coatings, and ink. These materials are made from natural resources, including petroleum, wood, and sand, that undergo many physical and chemical changes in order to be made into products we rely on every day. Knowledge of chemistry helps identify potentially useful substances and mixtures, as well as effective techniques for making them on a large scale. TEKS 3E • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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Study Tip Remember that a strong background in chemistry is important to many careers, not just those that have the word chemist in their title.

Energy Researcher  Today, fossil fuels provide most of the developed world’s energy needs. The three main types of fossil fuels are coal, oil, and natural gas. Knowledge of chemistry helps scientists find, extract, and process fossil fuels. Because the supply of fossil fuels is limited, companies and governments are investing money in research on alternative energy resources. Much of this research involves chemistry. For example, researchers are looking for efficient ways to develop biofuels, fuels made from plant matter. Researchers also continue to work to develop motor vehicles that can economically run on hydrogen without the use of fossil fuels. Biochemist  Biochemists study the chemistry of living things. Biochemistry has thousands of applications for the improvement of life on Earth—or wherever humans are. Spacecraft, vehicles, space suits, and other objects are subject to extremes of temperature, radiation exposure, and other conditions not found on Earth. Chemists and materials scientists help develop, test, and evaluate the materials used in space travel. The work of biochemists has helped keep astronauts safe and healthy during space travel. Biochemists also are studying ways to grow food in space and to help recycle wastes. Their work also helps evaluate chemicals collected by robotic vehicles on Mars. The chemicals may provide clues about whether life on Mars exists or ever existed.

Figure 1 Making Bioplastics Corn is grown, harvested, and ground. A sugar called glucose is extracted.

Biochemists also may work in the production of plant-based plastics, called bioplastics. Unlike traditional plastics, a bioplastic is not petroleum based. Figure 1 shows one way that a bioplastic, polylactic acid (PLA), is made.

Bacteria are added to convert glucose into lactic acid.

Lactic acid molecules are linked into long chains called polymers.

The polylactic acid polymer (PLA) is formed into small pellets. The pellets can be spun into fibers or melted to take almost any form.

Medical or Pharmaceutical Careers  The human body is made of a huge number and variety of chemical compounds. These compounds are constantly undergoing chemical changes. Diseases often produce harmful compounds or disrupt the normal chemical reactions in the body. Medicines and drugs affect these chemical reactions. For these reasons, all doctors and nurses study chemistry as part of their training. Drug companies employ chemists to develop new medicines and drugs. These chemists may experiment with thousands of different chemicals before developing the proper compound for treating a disease. Once compounds have been developed, pharmacists are the professionals responsible for dispensing those compounds. The education of a pharmacist involves a great deal of chemistry, as you might imagine.

TEKS 3E • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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Nanotechnology Engineer  Nanotechnology is the development of technology on an atomic scale. It combines ideas from chemistry, physics, and engineering. One example of research in nanotechnology involves very small cagelike objects. These objects could deliver medicines precisely where they are needed in the body. In theory, the medicine would be protected from the immune system, and side effects of the medicine would be minimized. Nanotechnology is also used to manufacture chemicals. These chemicals are now part of sunscreens, fabrics, and antibacterial coatings for many products.

TEKS

End-of-Course Assessment Review

1. Describe  Which goal would best be accomplished by a biochemist? A finding the best way to manufacture a new type of paper B isolating the toxin in a poisonous plant C removing mercury from a polluted tract of land D determining the age of a fossil 2. Evaluate  Could a professional baker benefit from chemistry knowledge? If so, how? A Yes. Bakers apply knowledge of chemistry when they develop new recipes. B Yes. Bakers apply knowledge of chemistry when they measure ingredients. C Yes. Bakers apply knowledge of chemistry when they choose menu items that are popular and nutritious. D No. Baking does not involve chemistry in any way. 3. Apply Concepts  Why is knowledge of chemistry important in careers in every scientific field, including life science, physics, earth science, and space science? Give examples. 4. Form an Opinion  Which of the careers mentioned in this review do you think you would enjoy the most? What aspect of the chemistry involved would be most interesting to you?

TEKS 3E • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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TEKS TEKS REVIEW

3F

The History of Chemistry and ­Contributions of Scientists TEKS 3F

Research and describe the history of chemistry and contributions of scientists.

What is the history of chemistry? The history of chemistry spans from early prehistory to the modern age. Many individuals from around the world have made significant contributions to the advancements in chemistry throughout time. Alchemy was practiced in parts of the world as early as 400 b.c. Similar to modern chemistry, alchemy involved the study of matter. However, it also involved philosophical or mystical ideas that were not based on evidence. Alchemists had many goals, including prolonging life, curing disease, and turning ordinary metals such as lead or copper into gold. Although their methods were not scientific, they did develop many tools, such as flasks and tongs, and techniques, such as separating mixtures, that are still in use today. In the 1500s, there was a shift from alchemy to science in Europe. A new era began with scientists who based conclusions on experimental evidence.

What contributions have scientists added to chemistry over the ages? As in all sciences, our current body of chemistry knowledge is built upon the work of many contributors. A few of these are mentioned below. Lavoisier  Antoine-Laurent Lavoisier, who worked in the mid- to late-1700s, is known as the father of modern chemistry. Just a few of his accomplishments include validating the law of conservation of mass, determining that oxygen is required for materials to burn, and working with other chemists to develop a method for naming chemical compounds. Democritus, Dalton, Thomson, Rutherford, and Bohr  Scientists build on the ideas and discoveries of other thinkers and scientists. The Greek philosopher Democritus (460 B.C.–370 B.C.) was among the first to suggest the existence of atoms. He argued that atoms are indestructible and indivisible, but could not support his ideas with evidence. More than 2000 years later, John Dalton used experimental methods to gather data in support of the idea of atoms. On the basis of his observations, Dalton developed an atomic theory of matter. According to this theory, an element is composed of only one kind of atom. A compound is composed of particles that are combinations of different kinds of atoms. Figure 1 summarizes the work of other scientists who contributed to our current understanding of the atom. TEKS 3F • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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Figure 1

Contributors to the Atomic Model Scientist J.J. Thomson Eugene Goldstein James Chadwick Ernest Rutherford Niels Bohr Erwin Schrödinger

Contribution Discovered the electron Discovered the proton Discovered the neutron Developed the theory of the atomic nucleus Proposed that an electron moves in circular paths, or orbits Developed an equation to describe the movement of electrons, which move through regions called orbitals

Planck, Einstein, De Broglie  These scientists contributed to the understanding of the quantization of energy. Max Planck discovered that the energy of an entity changes in small units called quanta. He developed an equation to explain this phenomenon. In 1905, Albert Einstein applied Planck’s work to explain the photoelectric effect. Einstein proposed that light exists as quanta called photons. Photons behave like particles even though they have no mass at rest. In 1924, Louis de Broglie proposed that if light could act like both waves and particles, then perhaps particles could act as waves and light. This idea led to the development of quantum mechanics, the description of atoms and subatomic particles as waves. Mendeleev  In 1869, the work of Russian Dmitri Mendeleev led to the development of the periodic table. He organized the known elements based on increasing atomic mass. By applying his ideas, Mendeleev predicted the existence of two elements. These elements were discovered in 1875 and 1886, and each fit his descriptions precisely. The modern periodic table lists elements in order of increasing atomic number, a property unknown in Mendeleev’s time. Figure 2 describes other contributors to the modern periodic table. Figure 2

Contributors to the Development of the Periodic Table Scientist Mendeleev Meyer

Contribution Published a table that organized elements by properties in order of increasing atomic mass Published a very similar table shortly after Mendeleev

Moseley

Determined atomic number for each element

Boyle, Charles, Gay-Lussac, Avogadro  The work of many scientists led to the understanding of the behavior of gases. In 1662, Robert Boyle studied the relationship of pressure and volume. He devised what is now known as Boyle’s law, which states that the volume of a gas is inversely proportional to pressure. In 1787, Jacques Charles found that the temperature and volume of a gas increase at the same rate (Charles’ law). Then in 1802, Joseph Gay-Lussac found that the pressure of a gas is directly proportional to the temperature in Kelvins when volume is a constant (Gay-Lussac’s law). Related to these three gas laws is Amedeo Avogadro’s hypothesis that equal volumes of gas at the same temperature and pressure contain equal numbers of particles. TEKS 3F • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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Study Tip To remember the scientists in chemistry history, group them according to the work they accomplished. For example, Arrhenius, Brønsted and Lowry, and Lewis all described acids and bases. Remember their names as a group.

Le Chatelier  Henri Le Chatelier devised an equation to describe how equilibrium in a chemical reaction shifts as a result of changing conditions. His work led to Le Chatelier’s principle: If a stress is applied to a system in dynamic equilibrium, the system changes in a way that relieves the stress. Arrhenius, Brønsted, Lowry, Lewis  These scientists were pioneers in acid-base theories. In 1887 Svante Arrhenius defined acids as hydrogencontaining compounds that ionize to yield hydrogen ions (H+) in aqueous solution and bases as compounds that ionize to yield hydroxide ions (OH–) in aqueous solution. In 1923, Johannes Brønsted and Thomas Lowry separately devised another definition. According to the Brønsted-Lowry definition, an acid is a hydrogen-ion donor and a base is a hydrogen-ion acceptor. Gilbert Lewis devised a third acid-base definition: An acid accepts a pair of electrons and a base donates a pair of electrons during a reaction. Becquerel, the Curies, Rutherford  Nuclear chemistry started in 1896, when the Henri Becquerel discovered that uranium fogged photographic plates. This led to the idea of radioactivity, the spontaneous emission of energy from an element. Marie and Pierre Curie, who were working with Becquerel, discovered the radiation that caused the film to fog. This research led to Ernest Rutherford’s discovery of the proton and three kinds of radioactivity: alpha, beta, and gamma rays. Current work in nuclear chemistry involves particle accelerators, which are long tubes used to accelerate subatomic particles to very high speeds. Scientists are using particle accelerators to research the particles that make up protons, neutrons, and electrons, as well as other types of very small particles. Their work may lead to some fundamental changes in how scientists view the matter that makes up the universe.

TEKS

End-of-Course Assessment Review

1. Apply  A researcher claims to have measured the exact speed of an electron moving around the atomic nucleus. This claim violates an equation developed by A Robert Boyle. B Svante Arrhenius. C Erwin Schrodinger and Werner Heisenberg. D Dmitri Mendeleev. 2. Apply  Suppose you want to know the characteristics of an ­undiscovered element, ununnovium. The work of which of these ­scientists would best help you predict this element’s characteristics? A Dmitri Mendeleev C Marie Curie B Ernest Rutherford

D Antoine Lavoisier

3. Research  What do you think is the most significant discovery in chemistry during the last 100 years? Conduct research to help you form an answer. TEKS 3F • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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TEKS REVIEW

4A

Physical and Chemical Changes and Properties

Re adi n ess

TEKS 4A

Vocabulary physical property chemical property physical change chemical change

Differentiate between physical and chemical changes and properties. TEKS_TXT

What is a physical property? Properties of substances can be classified as either physical or chemical. A physical property is a quality or a condition of a substance that can be observed or measured without changing the substance’s composition. Examples of physical properties include shape, length, mass, volume, melting point, boiling point, state of matter, color, hardness, density, and solubility. Physical properties can be used to identify a substance. Figure 1 shows some physical properties of several substances.

Figure 1

Physical Properties of Some Substances Substance Neon Oxygen Chlorine Ethanol Mercury Bromine Water Sulfur Aluminum Sodium chloride Gold Copper

State

Color

Gas Gas Gas Liquid Liquid

Colorless Colorless Greenish-yellow Colorless Silvery-white

Liquid Liquid Solid Solid Solid Solid Solid

Reddish-brown Colorless Yellow Silver White Yellow Reddish-yellow

Melting Point (°C) −249 −218 −101 −117 −39 −7 0 115 660 801 1064 1084

Boiling Point (°C) −246 −183 −34 78 357 59 100 445 2519 1413 2856 2562

What is a chemical property? A chemical property describes the ability of one substance to change into a different substance. A chemical property differs from a physical property in that we can observe it only by changing (or attempting to change) the composition of a substance. Examples of chemical properties include reactivity, flammability, heat of combustion, electronegativity, and the ability to oxidize. Chemical properties can also be used to identify a substance. But chemical properties can be observed only when a substance undergoes a chemical change.

TEKS 4A • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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What is a physical change? A physical change is a change that does not alter the chemical identity of a substance. The substance that exists before the change still exists after the change. For example, when an aluminum can is cut, crushed, or broken into smaller pieces, its chemical identity—aluminum—does not change. It is still aluminum. Cutting, crushing, and breaking into smaller pieces are all physical changes. Similarly, when aluminum melts, changing state from a solid to liquid, it is still aluminum. Changes of state, such as melting, freezing, and condensation, are physical changes. Dissolving is another example of a physical change. When salt dissolves in water, it is still salt. We can separate the salt from the water by evaporating the water. Although one or more physical properties of a substance might change during a physical change, its chemical properties stay the same.

What is a chemical change? A chemical change occurs when one or more substances change into one or more different substances. A substance present before the change occurred is called a reactant; a substance produced as a result of the change is called a product. In a chemical change, the products have chemical and physical properties that are different from those of the reactants. For example, when iron is exposed to moist air, the iron reacts with oxygen molecules; the oxygen and iron combine to form new substances, including iron oxide (Fe2O3)—“rust.” This reaction is shown below.



4Fe  +  3O2  ➝  2Fe2O3 Iron

Oxygen

Iron(III) oxide

The physical and chemical properties of iron oxide are different from those of the iron and the water. Oxidation, therefore, is an example of a chemical change. Combustion, or burning, is another example of a chemical change. When wood burns in the presence of oxygen, two new substances—carbon dioxide and water—are formed. All chemical changes involve a transfer of energy. For example, in the burning of wood, energy is given off in the form of heat and light.

How can you recognize a chemical change?

Study Tip You may have heard the term chemical reaction in other science classes. The term chemical reaction has the same meaning as the term chemical change.

When chemical changes occur, they often exhibit telltale signs. Some of these signs are the formation of a gas, a change in color or odor, a change in temperature, the release of light or sound, and the formation of a precipitate. A precipitate is a solid that forms from a chemical change, not from a change of state. For example, if you mix two clear liquids and a yellow solid forms, it is likely that a chemical change has occurred. If a piece of metal changes color over time, it is likely that a chemical change has occurred. Although these signs can indicate that a change might have occurred, the only way to be certain of a chemical change is to test the chemical properties of the reactants and products. If the properties are different, then a chemical change has occurred.

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TEKS

End-of-Course Assessment Review

1. Identify  Which of the following observations likely describes a physical change? A A solid was added to water and the mixture was stirred until the solid was no longer visible. B A piece of metal was added to a liquid and bubbles formed on the surface of the metal. C A white solid was exposed to ultraviolet light and turned brown. D Two clear liquids were mixed together and a yellow solid formed. 2. Draw Conclusions  A student places a beaker of liquid water on a hot plate. After a few minutes, bubbles form in the water. The student tests the bubbles and determines that they contain gaseous water. Is the formation of bubbles by the heating of water a physical change or a chemical change? A The change is chemical because a new substance was formed. B The change is chemical because formation of a gas indicates a chemical change. C The change is physical because no new substances were formed. D The change is physical because the physical properties of the substance changed. 3. Explain  The equations below describe two changes. The first equation describes a change that occurs when sodium chloride is added to water. NaCl + H2O ➝ Na+ + Cl– + H2O The equation below describes the change that occurs when the mixture of sodium chloride and water is left out for several days. Na+ + Cl– + H2O ➝ NaCl + H2O Explain what is happening to the atoms in each equation. Identify each change as either physical or chemical and explain your reasoning. 4. Differentiate  List two ways in which chemical changes differ from physical changes.

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TEKS REVIEW

4B

Intensive and Extensive Properties TEKS 4B

Identify extensive and intensive properties.

What is an intensive property?

Vocabulary intensive property extensive property

Study Tip As you learn about chemical and physical properties of substances, practice classifying these properties as extensive or intensive.

Figure 1

One way to classify physical and chemical properties of matter is as intensive or extensive. An intensive property is a property that is not dependent on the amount of a substance. Different amounts of a substance will have the same intensive properties if all other conditions are equal. For example, density is an intensive property. Density is the ratio of mass to volume of a substance. A large volume of liquid water will have the same density as a small volume of liquid water. Similarly, a large volume of water will boil at the same temperature as a smaller volume of water. Boiling point is an intensive property. Reactivity is another example of an intensive property. Other examples of intensive properties include melting point, hardness, solubility, the ability to oxidize, and specific heat. Knowing how a substance reacts with other substances can help you identify it. An intensive property depends on the identity of a substance. Different substances have different intensive properties. Several intensive properties of seven substances are shown in Figure 1. The more intensive properties that are known for a given substance, the easier it is to identify the substance.

Intensive Properties of Selected Substances Substance

Melting Point (°C)

Boiling Point (°C)

Density (g/cm3)

Iron Water Aluminum Gold Silver Mercury Platinum

1538 0 660.45 1064.43 961.78 −38.83 1768.3

2862 100 2519 2856 2162 356.73 3825

7.86 1 2.702 19.31 10.49 13.534 21.45

TEKS 4B • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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What is an extensive property? An extensive property is one that changes with the amount or size of a substance. For example, mass is an extensive property. A larger amount of water has a greater mass than a smaller amount of water. Length, volume, and energy are other examples of extensive properties. An extensive property is not dependent on the identity of a substance. This is because different substances can have the same extensive property. For example, knowing that a sample of matter has a mass of 5 g does not help identify it. Samples of iron, water, aluminum, and gold all can have a mass of 5 g. Because extensive properties depend on the size of a sample and not on the identity of a sample, a single extensive property, such as mass or volume, is not helpful in determining the identity of a substance. However, sometimes knowing two or more extensive properties can help you identify a substance. For example, knowing both the mass and the volume of a substance can yield information that is dependent on the identity of the substance. Some of the differences between extensive and intensive properties are listed in the table below. Figure 2

Extensive Properties and Intensive Properties Extensive properties • Dependent on the amount of a substance

Intensive properties • Not dependent on the amount of a substance

• Not dependent on the identity of a substance

• Dependent on the identity of a substance

Examples

Examples

Mass

Density

Length

Boiling point

Volume

Reactivity

Energy

Solubility Specific heat Hardness

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TEKS

End-of-Course Assessment Review

1. Identify  Which of the following is an example of an intensive property? A mass B thermal energy C length D density 2. Conclude  A student measures the volume of a substance to be 3.25 mL. Can the student determine the identity of the substance from this measurement? A No, because volume is an intensive property. B No, because volume is an extensive property. C Yes, because volume is an intensive property. D Yes, because volume is an extensive property. 3. Apply Concepts  You have measured the length and density of a pure metal object in your chemistry lab. You will be able to identify the metal the object is made of because A length is an intensive property. B length is an extensive property. C density is an intensive property. D density is an extensive property. 4. Evaluate  In a classroom discussion, one of your classmates argues that mass and volume are intensive properties, not extensive properties, because they give you information that helps determine the identity of a substance. Is your classmate correct? Explain your answer. 5. Infer  Iron oxidizes faster than gold. Is oxidation an example of an extensive property or an intensive property? Explain. 6. Analyze Data  A student is given Substance X to analyze. He measures the volume of the substance to be 3.25 mL, and the mass to be 2.55 g. The student also measures the boiling point to be 82.5 °C. Which of the following substances in the table below is the identity of Substance X? Explain two ways you can confirm the substance’s identity. Substance

Mass (g) 2.55 5.11 3.25 4.60

Methanol Isopropyl alcohol Water Chloroform

Volume (cm3)

Boiling Point (°C)

3.22 6.50 3.25 3.10

64.7 82.5 100.0 61.2

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TEKS REVIEW

4C

States of Matter TEKS 4C Compare solids, liquids, and gases in terms of compressibility, structure, shape, and TEKS_TXT volume.

What are states of matter?

Vocabulary solid liquid gas

The three states of matter found commonly on Earth are solids, liquids, and gases. All of these states have different properties and characteristics. All matter is made of particles—atoms, molecules, or ions. The three states differ in how the particles they consist of are arranged and how these particles move in relation to one another. These differences are responsible for the unique characteristics of each state.

What is a solid? A solid is a state of matter that has a definite shape and volume. The particles in a solid are held closely together, usually in a regular arrangement, or pattern. Figure 1 shows what a solid might look like if you could zoom in to see the particles that make up the solid. You can see that the particles are packed closely together in a rigid arrangement. Figure 1 A Solid

All particles in matter are in constant motion, even the particles in a solid. But because the particles in a solid are held very close together, they are able only to vibrate in place. The particles cannot move around much relative to one another, and therefore solids do not flow. The shape of a solid does not change to fit the shape of a container it is in. For example, the shape of a rock does not change when it is moved from inside a small beaker to a large beaker. Also, because the particles of a solid are held closely together, the volume of solids is nearly fixed. The particles that make up a solid can move only very slightly. As a result, solids are almost incompressible; that is, it is difficult to squeeze a solid into a smaller volume. In addition, solids expand only slightly when heated.

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What is a liquid? A liquid is a state of matter that has an indefinite shape, flows, yet has a fixed volume. The particles in a liquid are held closely together. However, in a liquid, particles have enough energy to slide past one another. Figure 2 shows what a liquid might look like if you could zoom in to see the particles that make up the liquid. Figure 2 A Liquid

The ability of particles in a liquid to move around relative to one another leads to two characteristics that are different from those of a solid. First, liquids are able to flow. You can see this when you pour a liquid from one container to another or when water in a river flows downstream. Second, liquids are able to take the shape of the container they are in. If you pour a liquid from one container to another, the shape of the liquid changes to match the shape of the new container. Because the particles in a liquid are only slightly farther apart than they are in a solid, the volume of liquids is nearly fixed. Liquids are only slightly more compressible than solids. They tend to expand slightly when heated.

Study Tip Make flash cards to study solids, liquids, and gases. On one side of each card, write a characteristic, such as “only slightly compressible.” Then write the appropriate state or states, such as “solids and liquids” on the other side.

What is a gas? Similar to liquids, gases take on the shape of their container. But, unlike liquids or solids, gases can expand to fill the container that they are in. A gas is a state of matter that takes both the shape and volume of its container. The large amount of space between gas particles allows a gas to be compressed. For example, if you add more air to a full bicycle tire, the air in the tire compresses—the particles that make up the air in the tire move closer together. When you let some of the air out of the tire, the particles move farther apart and the gas expands. Figure 3 shows what a gas might look like if you could zoom in to see the particles that make up the gas. Notice that the particles in a gas are far apart and can move more quickly then in a solid or a liquid.

Figure 3 A Gas

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TEKS

End-of-Course Assessment Review

1. Compare  Which of the following correctly describes the particles in a solid compared to the particles in a liquid? A Particles in a solid are farther apart than particles in a liquid. B Particles in a liquid are much farther apart than particles in a solid. C Particles in a solid do not move at all, and particles in a liquid slide past one another. D Particles in a solid vibrate in place and particles in a liquid slide past one another. 2. Identify  Which observations were likely made of a gas? A This substance completely filled only the bottom half of a container. B The volume of this substance did not change when pressure was applied. C When this substance was transferred from a smaller container to a larger container, its volume increased. D When this substance was moved from one container to another, its shape stayed the same. 3. Analyze  Which of the following correctly describes what happens when an inflated balloon is compressed? A When the gas is compressed, its particles get smaller. B When the gas is compressed, its particles get closer together. C When the gas is compressed, its particles get larger. D When the gas is compressed, its particles get farther apart. 4. Apply  Explain how inflatable packing materials and foam packaging peanuts both rely on the properties of gases to keep objects safe during shipping.

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4D Classifying Matter Re adi n ess

TEKS 4D

Classify matter as pure substances or mixtures through investigation of their properties. TEKS_TXT

How is matter classified?

Vocabulary pure substance element compound mixture solution

One way matter can be classified is by the type and arrangement of the particles that it consists of. Using this method, matter can be classified into one of two broad groups—pure substances or mixtures. Each of these groups can be classified into subgroups. Matter

Can be separated physically

Pure substance Definite composition (homogeneous)

Can be separated chemically Element

Compound

Mixture of substances Variable composition

Homogeneous mixture Uniform; also called a solution

Heterogeneous mixture Nonuniform; distinct phases

What is a pure substance? A pure substance is matter that has a definite composition. This means that the composition of a pure substance is always the same wherever it is found. Water is an example of a pure substance. Water is always made of molecules composed of one oxygen atom bonded to two hydrogen atoms. Another example of a pure substance is aluminum. Aluminum is always composed of only aluminum atoms. TEKS 4D • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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What are elements and compounds?

Study Tip As you read about the different types of matter, use a graphic organizer, such as a spider map, to keep track of definitions, examples, and sketches of what each type of matter might look like at the atomic level.

Pure substances can be classified as elements or compounds. An element is matter that is made up of only one type of atom. All elements are pure substances. Aluminum is an element. Other familiar elements are gold, helium, copper, iron, and oxygen. A compound is matter that is made up of two or more different atoms chemically combined in a fixed proportion. All compounds are pure substances. Water is a compound. Other familiar compounds are carbon dioxide, sodium chloride (table salt), and glucose. You cannot classify a sample of matter as an element or compound simply by looking at it. Both will have a uniform texture and color. The only way to identify whether the sample is an element or compound is to test its chemical properties. If the sample can be broken down by chemical changes into two or more different elements, it is a compound.

How do mixtures differ from a pure substance? A mixture is a physical blend of two or more substances that are not chemically combined. Nearly all of the matter that you come into contact with every day is a mixture. Wood, fabric, food, beverages, plastics, plants, and air are all examples of mixtures. Unlike a pure substance, the components of a mixture can be separated using physical processes such as boiling and filtering. The composition of a mixture can vary. For example, air is a mixture of nitrogen, oxygen, hydrogen, carbon dioxide, and other pure substances. The percentage of each of these gases can vary in different mixtures. For example, the percentage of oxygen in room air is about 21 percent, while the percentage of oxygen in air within an underwater dive tank can be below 20 percent. To classify a sample of matter as a pure substance or a mixture, you must study its physical and chemical properties. If the texture and color of the matter vary, then you know that it is a mixture because all pure substances are uniform. However, if the substance has uniform texture and color, then you must determine whether physical means can separate the components. If so, the sample is a mixture because pure substances cannot be separated into different components by physical means.

What are homogeneous and heterogeneous mixtures? Mixtures can be either homogeneous or heterogeneous. A homogeneous mixture is evenly mixed at the atomic level. Another name for a homogeneous mixture is solution. Solutions have a uniform texture and color throughout. Variations in texture and color cannot be seen with the unaided eye or even with traditional microscopes. Solutions can be liquids, solids, or gases. Vinegar and fruit drinks made with a drink mix are examples of liquid solutions. Alloys such as brass or steel are examples of solid solutions. Air is an example of a gaseous solution. TEKS 4D • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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A heterogeneous mixture is one that is unevenly mixed. Different parts of a heterogeneous mixture have different compositions. You can see these differences with the unaided eye or with a traditional microscope. For example, a mixture of cereal, fruit, and milk is heterogeneous.

TEKS

End-of-Course Assessment Review

1. Infer  A student is given a sample of matter to classify. The sample is uniform in color and texture. Based on this information, the student can determine which of the following? A The sample is an element. B The sample can be an element or a compound. C The sample is a homogeneous mixture. D The sample can be a homogeneous mixture or a pure substance. 2. Classify  A new material was produced in a laboratory by bonding atoms of different elements together in a unique arrangement. Other laboratories were able to reproduce these results, producing the same material with the same elements in the same arrangement. What type of matter was produced? A element B compound C homogeneous mixture D heterogeneous mixture 3. Infer  Vinegar is classified as a mixture. What does this tell you about vinegar? A Vinegar is made of atoms of one type of element. B Vinegar is made of atoms of different elements chemically bonded together. C Vinegar has a composition that is the same wherever it is found. D Vinegar contains two or more pure substances mixed together. 4. Classify  When you look at blood with your unaided eyes, it has a uniform texture and color. However, when you look at a blood sample under a microscope, you can see it contains platelets, red blood cells, and other particles suspended in a liquid. Is blood an element, compound, homogeneous mixture, or heterogeneous mixture? Explain your reasoning. 5. Conclude  A student does a series of experiments on a sample of unknown material. The student determines that physical processes are not able to break the sample into simpler substances. However, a chemical reaction produces carbon, hydrogen, and oxygen. What type of matter is the sample? Explain your reasoning. TEKS 4D • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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5A

Development of the Periodic Table TEKS 5A Explain the use of chemical and physical properties in the historical development of the Periodic Table.

How were chemical and physical properties used in the development of the periodic table?

Vocabulary periodic table

As scientists learned how to isolate elements from compounds, the number of known elements increased rapidly in the early 1800s. Scientists looked for patterns and similarities in the properties of elements in order to classify the growing number of discovered elements. In 1869, two chemists—Dmitri Mendeleev from Russia and Lothar Meyer from Germany—independently published very similar classification systems. Mendeleev is often given more credit because he published his table first. Mendeleev knew that many of the known elements shared similar chemical and physical properties. In order to show the relationships among the known elements, Mendeleev organized the elements into a periodic table. The periodic table is an arrangement of elements in which the elements are separated into groups based on a set of repeating properties. Mendeleev arranged the elements in his periodic table in order of increasing atomic mass. By using this organization, Mendeleev noticed that the physical and chemical properties repeated in a predictable pattern. For example, the melting points of lithium (Li), beryllium (Be), boron (B), and carbon (C) increased with increasing atomic mass. However, the melting point of the next element, nitrogen (N), dropped sharply. This same pattern was seen with the elements sodium (Na), magnesium (Mg), aluminum (Al), silicon (S), and phosphorus (P). The following table shows the pattern of rise and sudden drop in melting point. Figure 1 Melting Points of Certain Elements

Element

Li

Atomic Mass Melting Point (°C)

Be 7

180.5

9 1278.0

B

C 11

2300.0

12

N 14

3550

–209.9 P

Element

Na

Mg

Al

Si

Atomic Mass

23

24

27

28

Melting Point (°C)

97.7

650.0

660.3

1410.0

31 41.4

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Study Tip Use the mnemonic “Very Good Hot Pies” to remember that the vertical columns in the Periodic Table are called groups, and the horizontal rows are called periods.

Mendeleev noticed that occasionally, an element placed in a group according to its properties was slightly out of place according to atomic mass. For example, the atomic mass of tellurium (Te) is slightly greater than that of iodine (I), yet the properties of the two elements suggested that tellurium be placed before iodine, not after it. Mendeleev assumed that the measured atomic masses of these elements were not accurate and maintained the order based on properties. Scientists soon learned that the atomic masses were correct. The problem was using atomic mass to organize the periodic table. Though Mendeleev and Meyer’s tables were nearly identical, Mendeleev took his classification system one step further. He left several empty spaces in his table for which no known element existed at the time with properties matching others in the column. Mendeleev boldly predicted that new elements would be discovered with properties that would fit into these empty spaces. For example, he referred to one of these undiscovered elements as eka-aluminum and predicted that its properties would be similar to those of aluminum. In fact, in 1875 an element with the atomic mass and properties predicted by Mendeleev was discovered. This element was named gallium (Ga). Mendeleev also predicted the properties of the element germanium (Ge), which was discovered in 1886.

How does the modern periodic table differ from the earlier table? Years later, in the early 1900s, a scientist named Henry Moseley determined an atomic number for each known element. The atomic number is the number of protons in an element. If the elements were ordered by atomic number, tellurium would come before iodine which made more sense. In the modern periodic table, elements are arranged in order of increasing atomic number. The elements in the modern periodic table start with hydrogen (H) with an atomic number of 1 (one proton in the nucleus). The elements are organized into seven rows called periods. Each period corresponds to a principal energy level. There are more elements in higher-numbered periods because there are more orbitals in higher energy levels. The properties of the elements within a period change as you move across a period from left to right. However, the pattern of properties within a period repeats as you move from one period to the next. This pattern gives rise to the periodic law: When elements are arranged in order of increasing atomic number, there is a periodic repetition of these physical and chemical properties. The arrangement of the elements into periods has an important consequence. Elements that have similar chemical and physical properties end up in the same column (or group) in the periodic table. An example of the modern periodic table can be found in the Chemistry Reference Materials at the end of this book.

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TEKS

End-of-Course Assessment Review

1. Infer  What can you infer about elements that are in the same group? A They will have the same atomic mass. B They will contain the same number of protons. C They will display similar chemical and physical properties. D They will emit X rays of similar frequency. 2. Think Critically  Why did Mendeleev switch the order of some elements in his periodic table? A He determined that the atomic masses were incorrect. B He decided that elements should be placed in order of atomic mass even if this meant placing elements with different properties in the same group. C He decided that elements should be grouped based on properties even if this meant placing elements out of order by atomic mass. D He determined that elements should be placed in order of increasing atomic number. 3. Identify  The periods and groups on the periodic table are numbered. For example, the element neon (Ne) is in period 2, group 18. Neon is an inert gas and therefore does not react readily with any other element. Which of the following elements is likely to have properties that are similar to those of neon? A lithium (Li) in period 2, group 1 B boron (B) in period 2, group 13 C iron (Fe) in period 4, group 8 D krypton (Kr) in period 4, group 18 4. Explain  How did Mendeleev’s organization of elements help him predict the properties of elements that had not yet been discovered?

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5B

Chemical Families in the Periodic Table

Readi ness

TEKS 5B Use the Periodic Table to identify and explain the properties of chemical families, including alkali metals, alkaline earth metals, halogens, noble gases, and transition metals.

What are the main chemical families on the periodic table?

Vocabulary alkali metal alkaline earth metal halogen noble gas transition metal

On the modern periodic table, elements are placed in order of atomic number and then into groups, or families, based on repeating, or periodic, properties. There are 18 families on the periodic table. The families are grouped into representative elements (groups 1, 2, and 13 through 18), transition elements (groups 3 through 12) and inner transition elements (actinides and lanthanides), as shown in Figure 1. Elements within each of these families share many physical and chemical properties.

Figure 1 Chemical Families in the Periodic Table

Noble Gases

Alkali Metals Alkaline Earth Metals

Halogens

Transition Metals

What properties do alkali metals share? Alkali metals make up Group 1 on the Periodic Table. The elements in this family are all metals and share metallic properties, such as a shiny luster, malleability, and high thermal and electric conductivity. Alkali metals are soft solids at room temperature and have low melting points and densities. Notably, all alkali metals are very reactive. For example, alkali metals react vigorously with water to produce hydrogen and a basic solution. Because they are so reactive, alkali metals are found in nature only within compounds and never as free elements. When atoms of alkali metals react, they usually lose one electron, resulting in an atom with a +1 oxidation state. TEKS 5B • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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What properties do alkaline earth metals share? The alkaline earth metals make up Group 2 on the periodic table. The elements in this family are soft metals. Alkaline earth metals are also very reactive, although slightly less reactive than alkali metals. Except for beryllium (Be), they react with water to produce basic solutions. Most occur naturally only in compounds. When atoms of alkaline earth metals react, they usually lose two electrons, resulting in an atom with a +2 oxidation state.

What properties do halogens share? The halogens make up Group 17 on the periodic table. All halogens are nonmetals and are generally poor conductors of heat and electricity. Unlike the alkali metals and alkaline earth metals, the halogens do not share a common state of matter. In fact, this is the only family on the periodic table to contain elements in all three states of matter. At room temperature, fluorine (F) and chlorine (Cl) are gases, bromine (Br) is a liquid, and iodine (I) and astatine (At) are solids. Halogens are very reactive and are rarely found in nature as free elements. In fact, fluorine is one of the most reactive elements and will even react with glass, a relatively inert material. The elemental form of a halogen is a diatomic molecule, such as F2, Cl2, and Br2. When halogen atoms react, they usually gain one electron, resulting in an atom with a −1 oxidation state.

What properties do the noble gases share? Noble gases make up Group 18 on the periodic table. As the name suggests, all noble gases are gases at room temperature. They are also odorless and colorless. Unlike elements in Groups 1, 2, and 17, noble gases have a very low reactivity and exist as unbounded single atoms in nature. For this reason, noble gases were once called inert, meaning nonreactive. However, scientists have since synthesized noble gas compounds in the laboratory, demonstrating that noble gases are not completely inert.

What properties do the transition metals share? Study Tip As you read about families in the Periodic Table, organize the information into a table that contains the family names in columns and row headers such as Elements, State of Matter, Metal or Nonmetal, and Reactivity.

Transition metals include elements in Groups 3 through 12 on the periodic table, although there is disagreement about including the elements zinc (Zn), cadmium (Cd), and mercury (Hg) because these elements do not exhibit some of the characteristics of other transition metals. Transition metals share metallic properties with elements in other families. Unlike elements in Groups 1 and 2 and Groups 13 through 18, transition metals can exist in several common oxidation states and are more likely to form metal complexes in which the charges are not balanced and there is an − excess number of electrons. Tetrachloroferrate(III), FeCl4 , is one such complex. Although compounds containing transition metals are very common, transition metals also exist as free elements and, as a family, are not considered exceptionally reactive.

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TEKS

End-of-Course Assessment Review

1. Identify  Which of the following groups contain elements that are gaseous at room temperature? A alkali metals and alkaline earth metals B alkali metals and transition metals C noble gases and transition metals D noble gases and halogens 2. Classify  When this element was discovered, it exhibited luster and malleability, and it reacted very vigorously with water. This element is never found as a free element in nature and always exists in a compound. To which group does this element most likely belong? A alkali metals B halogens C noble gases D transition metals 3. Identify  Which sequence contains elements listed from most reactive to least reactive? A transition metals, noble gases, halogens B transition metals, alkali metals, alkaline earth metals C alkali metals, alkaline earth metals, noble gases D alkaline earth metals, alkali metals, halogens 4. Predict  Based on each family’s ability to either gain or lose electrons, predict what might happen if an element in Group 1 came into contact with an element in Group 17. 5. Explain  Why was the term inert gas once used to refer to noble gases and why is it no longer in common use?

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5C

Trends in the Periodic Table

Re adi n ess

TEKS 5C

Vocabulary atomic radius ion cation anion electronegativity ionization energy

Use the Periodic Table to identify and explain periodic trends, including atomic and ionic TEKS_TXT radii, electronegativity, and ionization energy.

What periodic trends in atomic radii can be identified in the periodic table? The size of an atom is expressed as an atomic radius (plural, radii). The atomic radius is one half the distance between the nuclei of two atoms of the same element when the atoms are joined. In general, atomic radii decrease from left to right across a period on the periodic table and increase from top to bottom within a group, or family. As you move across a period in the periodic table, the number of protons in the atoms increases, but the electrons remain in the same energy level. Therefore, outer-level electrons are pulled more strongly toward the nucleus from left to right across a period. This increasingly stronger pull results in a smaller radius from left to right across a period. The principal quantum number, n, of the outer-level electrons increases by one from period to period. For example, for elements in period 1, n = 1. For elements in period 2, n = 2, and so on. As n increases down a family, the outer-level electrons have an average position that is farther from the nucleus. As a result, the atoms are larger.

What periodic trends in ionic radii can be identified in the periodic table? An ion is an atom or group of atoms that has a positive or negative charge. There are two types of ions—cations and anions. Cations are atoms that have lost one or more electrons and thus have a positive charge. Atoms that lose electrons become smaller. For example, the calcium ion, Ca2+, is smaller than a calcium atom, Ca, because Ca2+ has two fewer electrons. Anions are atoms that have gained one or more electrons and thus have a negative charge. Atoms that gain electrons are bigger. For example, a bromide ion, Br–, is larger than a bromine atom, Br, because Br– has one more electron. Cations and anions exhibit trends that are similar to those of their parent atoms across periods and down families. From left to right across a period, the radii of cations and anions decrease because the number of protons in the nucleus increases. From top to bottom within a family, the radii of cations and anions increase because the principal quantum number, n, increases. TEKS 5C • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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What periodic trends in electronegativity can be identified in the periodic table?

Study Tip Draw a rough outline of the periodic table leaving room for labels above and to the left of the table. For each trend that you study, draw a line above the table and on the left side, and use arrowheads to indicate the direction in which the trend increases. Label each arrow.

Electronegativity is the ability of an atom to attract electrons when the atom is in a compound. The greater an atom’s electronegativity, the greater its ability to attract electrons. The number of protons in the nucleus and the principal quantum number influence the periodic trends for electronegativity. Generally, the electronegativity increases from left to right across a period of the periodic table because the number of protons in the nucleus increases. Electronegativity generally decreases from top to bottom within a family because outer energy level electrons are farther from the nucleus.

What periodic trends in ionization energy can be identified in the periodic table? Ionization energy is the minimum energy required to remove an electron from an atom or ion. The energy required to remove the first electron from an atom is referred to as the first ionization energy. The greater an element’s ionization energy, the more difficult it is to remove an electron. Generally, ionization energy increases from left to right across a period and decreases from top to bottom within a family.

Figure 1 Trends in Ionization Energy

First ionization energy (kJ/mol)

Ionization energy depends on the force of attraction the nucleus exerts on the electron. As with the other periodic trends, this attraction depends upon the number of protons in the nucleus and the distance of the electron from the nucleus. More protons exert more force, making electrons harder to remove. Therefore, from left to right across a period, ionization energy increases because the number of protons in the nucleus increases. Electrons that are closer to the nucleus are pulled more strongly toward the nucleus, making them harder to remove. Therefore, as atomic radii increase from top to bottom within a family, electrons that are farther from the nucleus are easier to remove. These trends can be seen in Figure 1 below.

First Ionization Energy vs. Atomic Number

2500

He

Ne

2000 1500 1000 500 0

Ar

N H

Kr Xe

P

Be

Zn As

Mg Li

Na

10

Rb

K

20

Cd

30

40

Atomic number

Cs

50

60

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TEKS

End-of-Course Assessment Review

1. Identify  Which of the following trends can be identified on the periodic table? A Atomic radii increase from left to right across a period. B Ionization energy increases from top to bottom within a family. C Electronegativity decreases from left to right across a period. D Ionic radii of cations decreases from left to right across a period. 2. Explain  Which of the following correctly explains why the sizes of atoms decrease from left to right across a period? A The principal quantum number increases. B The number of electrons increases. C The distance from the nucleus increases. D The number of protons increases. 3. Explain  An increase in principal quantum number explains which of the following trends? A The ionization energy decreases from top to bottom within a family. B The ionization energy increases from left to right across a period. C The electronegativity increases from left to right across a period. D The atomic radius increases from left to right across a period. 4. Apply Concepts  Use the table below to determine which of the following relationships is correct.

Selected Trends in the Periodic Table Element Li C Rb Be

Period 2 2 5 2

Family 1A 4A 1A 2A

A Li has a smaller atomic radii than C. B Li+ has a larger atomic radii than Li. C Rb has a smaller atomic radii than Li. D Li has a larger atomic radii than Be. 5. Explain  Why does an increase in the number of protons in the nucleus of an atom increase the ionization energy of atoms within a period?

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6A Atomic Theory TEKS 6A Understand the experimental design and conclusions used in the development of modern atomic theory, including Dalton’s Postulates, Thomson’s discovery of electron properties, Rutherford’s nuclear atom, and Bohr’s nuclear atom.

What are Dalton’s Postulates?

Vocabulary atom electron nucleus proton neutron

In the early 1800s, English chemist John Dalton expanded on the 2000-yearold ideas of Greek philosophers about matter being made of tiny particles. Dalton proposed an atomic theory that include these postulates: 1. Elements are made of very small, indivisible particles called atoms. 2. All atoms of a given element are identical. 3. Atoms of a given element are different from those of any other element and have different atomic masses. 4. In a chemical reaction, atoms of one element can combine with atoms of another element in whole-number ratios to form compounds. 5. Chemical reactions change the arrangement of atoms, but do not change atoms of one element into atoms of another element. Some of these postulates have been confirmed, while others are now known to be inaccurate. For example, we now know that atoms are divisible. An atom is now defined as the smallest particle of an element that retains its identity in a chemical reaction.

What did Thomson discover about electron properties? Study Tip As you study the observations that shaped our current understanding of the atom, make a twocolumn table. Keep track of conclusions made about the atom in one column. In the other column, list the observations used to form those conclusions.

In the late nineteenth century, scientists experimented with glass tubes containing gas at low pressure with electrodes at each end. When the electrodes are connected to an electric current, a cathode ray is produced. English physicist J. J. Thomson observed that cathode rays are attracted to positively charged metal plates and repelled by negatively charged plates. Thomson hypothesized that cathode rays are composed of negatively charged particles. He performed more experiments to calculate the chargeto-mass ratio of a single charged particle, and concluded that the mass of one of these particles was much less than the mass of the smallest known atom. This subatomic, negatively charged particle was eventually called the electron. Thomson proposed a new model of the atom, called the “plum-pudding model.” This model described an atom as electrons submerged in a positively charged substance, similar to raisins stuck in dough.

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How did Rutherford’s experiments lead to the nuclear atom model? In 1911, Ernest Rutherford, a former student of Thomson, set out to test the plum-pudding model. He passed positively charged high-energy alpha particles, which are helium nuclei, through a thin sheet of gold foil. Rutherford predicted that the alpha particles would pass straight through the foil because the atom’s positive charge was spread out evenly and because electrons were too small to deflect an alpha particle. As predicted, when the gold foil was bombarded with alpha particles, most of the particles passed straight through the foil without deflection. However, to Rutherford’s surprise, a very small fraction of the particles were strongly deflected and some were reflected straight back toward the source, as shown in Figure 1.

Figure 1 Rutherford’s Experiment

Rutherford hypothesized that the alpha particles struck a small, dense region of positive charge within the gold atoms called the nucleus. He proposed a new model of the atom that contained a nucleus surrounded by electrons—the nuclear atom. Further research showed that the nucleus contains positively charged particles called protons and neutral particles called neutrons. Fluorescent screen

Gold foil

Nucleus

Lead shield Alpha particles

Source of alpha particles

Atoms of gold foil

Beam of alpha particles

How did Bohr refine the model of the nuclear atom? In 1913, Danish physicist Niels Bohr, a student of Rutherford, developed a new atomic model. He changed Rutherford’s model to incorporate newer discoveries about how the energy of an atom changes when the atom absorbs or emits light. Bohr proposed that an electron is found only in specific circular paths, or orbits, around the nucleus. Each electron orbit in Bohr’s nuclear atom has a fixed energy. The fixed energies an electron can have are called energy levels. An electron can move from one energy level to another by gaining or losing just the right amount of energy. A quantum of energy is the amount of energy required to move an electron from one energy level to another energy level. TEKS 6A • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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The amount of energy that an electron gains or loses in an atom is not always the same. The higher energy levels are closer together. So, the higher the energy level occupied by an electron, the less energy it takes for the electron to move from that energy level to the next higher energy level. While Bohr’s ideas about specific energy levels are part of the modern atomic model, scientists now know that electrons do not travel in orbits. The modern model—the quantum mechanical model—proposes that electrons travel throughout an electron cloud, a region of space surrounding a nucleus that contains protons and neutrons.

TEKS

End-of-Course Assessment Review

1. Conclude  Which of the following observations led Thomson to conclude that cathode ray particles were negatively charged? A He observed that negatively charged plates repelled the rays. B He observed that a metal plate exposed to cathode ray particles became positively charged. C He observed that cathode ray particles traveled in a straight path through electric fields. D He observed that positively charged plates repelled the rays. 2. Infer  Given what Rutherford learned about the structure of atoms from the gold-foil experiment, what would Rutherford likely have observed if he had used platinum foil instead of gold foil? Platinum is an element with properties similar to those of gold, but its atoms have slightly less mass than gold. A All of the alpha particles would travel straight through the platinum without deflection. B A few alpha particles would be slightly deflected but none would be strongly deflected by the platinum foil. C A few alpha particles would be strongly deflected by the platinum foil. D All of the alpha particles would be strongly deflected by the platinum foil. 3. Describe  How did the observations of Dalton, Thomson, Rutherford, and Bohr all contribute to the modern model of the atom?

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TEKS REVIEW

6B

Electromagnetic Waves TEKS 6B Understand the electromagnetic spectrum and the mathematical relationships between TEKS_TXT energy, frequency, and wavelength of light.

How do energy, frequency, and wavelength characterize light waves?

Vocabulary frequency hertz wavelength Planck’s constant

A wave is a disturbance that transfers energy from one point to another. For example, water waves carry energy through the water from the source of the disturbance, such as a splash, outward in all directions. Sound waves carry energy through air from the source of the sound, such as a radio speaker, outward in all directions. Other examples of waves include seismic waves, which are waves produced by an earthquake, and electromagnetic waves, which consist of oscillating electric and magnetic fields. Visible light is one type of electromagnetic wave. A wave can be described by its frequency, wavelength, and energy. Frequency is represented by the Greek letter nu, ν. The frequency of a wave is the number of waves that pass a given point per unit of time. The units of frequency are usually cycles per second. The SI unit of cycles per second is called the hertz (Hz). A hertz can be expressed as s−1. For example, 1000 Hz is 1000 waves per second, or 1000 s−1. The wavelength of a wave is the distance between two equivalent points on a wave. In Figure 1, the wavelength is measured between two troughs. Wavelength is represented by the Greek letter lambda, λ, and is measured in units of length, such as meters (m), centimeters (cm), and nanometers (nm).

Figure 1 Frequency and Wavelength

A. Higher Frequency, Shorter Wavelength Time

Study Tip Note that sometimes the letter “f” is used to represent frequency, rather than the Greek letter nu (ν).

1s

2s Wavelength (λ)

B. Lower Frequency, Longer Wavelength Time

1s

Wavelength (λ)

2s

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What are the mathematical relationships between wavelength and frequency? All electromagnetic waves travel at the speed of light in a vacuum. The speed of light, c, is a constant value of 3.00 × 108 m/s, or 300,000,000 meters per second. The product of the wavelength and frequency of a wave equals the speed of light. c = λν Frequency and wavelength are inversely proportional to each other. This means that as one increases, the other decreases such that the product of the two is always the constant c. As the frequency increases, the wavelength decreases. Conversely, as the wavelength increases, the frequency decreases. Because the product of wavelength and frequency is equal to a constant, you can always calculate one of these values for an electromagnetic wave if you know the value of the other. For example, if you know the frequency of a wave, you can calculate its wavelength. And, if you know the wavelength of a wave, you can calculate its frequency.

What are the mathematical relationships between energy and frequency? For much of history, scientists thought of light as strictly wavelike in its nature. But this wavelike nature of light did not explain some phenomena. For example, when some metals were heated, they emitted light. The color and intensity of the light emitted depended on the temperature of the object. Hotter objects emitted light of greater intensity and higher frequency. In 1900, German physicist Max Planck determined a relationship between energy and frequency. He determined that matter absorbs electromagnetic radiation in small, specific amounts, called quanta. Planck showed that the amount of radiant energy (E) of a single quantum absorbed or emitted by a body is proportional to the frequency of radiation (ν). E = hν The constant (h), which has a value of 6.626 × 10–34 J•s, is called Planck’s constant. Energy and frequency are directly proportional to one another. This means that as one increases, the other also increases such that the ratio of the two is equal to the constant h.

What is the electromagnetic spectrum? Visible light is only one type of electromagnetic radiation. Other types include radio waves, microwaves, infrared waves, ultraviolet waves, X-rays, and gamma rays. The types of waves differ by their wavelengths and frequencies. For example, radio waves have long wavelengths, low frequencies, and low energies. Gamma rays have short wavelengths, high frequencies, and high energies.

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The electromagnetic spectrum, shown in Figure 2, is a visual representation of all the wavelengths and frequencies of electromagnetic radiation. Notice that the visible spectrum of light makes up only a small portion of the entire electromagnetic spectrum.

Figure 2 The Electromagnetic Spectrum

Red light, low energy (7.0 × 10 –7 m = 700 nm)

Violet light, high energy (3.8 × 10 –7 m = 380 nm)

Visible light

Frequency 𝛎 (s −1) 3 × 106

3 × 10 8

Radio waves

10 2

1

3 × 10 10

Radar

3 × 10 12

Microwaves

10 –2

10 –4

3 × 10 14

Infrared

3 × 10 16

Ultraviolet

10 –6

3 × 10 18

X-rays

10 –8

3 × 10 20

3 × 10 22

Gamma rays

10 –10

10 –12

10 –14

Wavelength 𝛌 (m)

TEKS

End-of-Course Assessment Review

1. Identify  FM radio waves have higher frequencies than AM radio waves. If a particular AM radio wave has a wavelength of 375 m, which of the following wavelengths is likely of an FM radio wave? A 3.00 m C 4500 m B 650 m

D

70,000 m

2. Organize  Which of the following lists types of electromagnetic waves in order of increasing wavelength? A gamma, radio, ultraviolet, infrared B infrared, ultraviolet, radio, gamma C radio, infrared, ultraviolet, gamma D gamma, ultraviolet, infrared, radio 3. Calculate  What is the distance that radio waves travel in 10 seconds? 3.00 × 109 m A 3.00 × 108 m C B 6.00 × 108 m

D

6.00 × 109 m

4. Infer  Use what you have learned about energy and electromagnetic waves to explain why ultraviolet rays can cause sunburn but visible light from a light bulb does not damage skin.

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TEKS REVIEW

6C

Calculating Wavelength, ­Frequency, and Energy of Light TEKS 6C Calculate the wavelength, frequency, and energy of light using Planck’s constant and TEKS_TXT the speed of light.

How can you calculate the wavelength of light using its frequency and the speed of light?

Study Tip Keep track of units as you do your calculations. Knowing the units of the value that you are calculating can help you determine which equation to use.

In a vacuum, all electromagnetic waves, including light waves, travel at the speed of light. The speed of light, c, is a constant value of 300,000,000 meters per second (3.00 × 108 m/s). The product of the wavelength, λ, and frequency, ν, of a wave equals the speed of light, c. c = λν Because the product of wavelength and frequency is equal to a constant, you can always calculate one of these variables if you know the value of the other. For example, if you know the frequency of a wave, you can calculate its wavelength by dividing both sides of the equation by frequency. The result is: c λ= ν

Sample Problem 1 Visible light has frequencies between about 4.0 × 1014 hertz (Hz) and about 7.9 × 1014 Hz. What are the wavelengths of the lowest frequencies of visible light? First, substitute the lowest frequency, 4.0 × 10 Hz, into the equation: c λ= 4.0 × 1014 Hz 14

Next, substitute the speed of light for c:

λ=

3.00 × 108 m/s

Multiplying the numerator and denominator by s leaves only the unit of m. Divide 3.00 × 108 m by 4.0 × 1014.

λ=

3.00 × 108 m 4.0 × 1014

Recall that the exponent of the denominator, 14, is subtracted from the exponent of the numerator, 8.

λ = 0.75 × 10–6 m or 7.5 × 10 –7 m The longest wavelengths of visible light are about 7.5 × 10 –7 m.

4.0 × 1014 Hz

Notice that, in this form, the units do not cancel. Recall that 1 Hz equals 1 s−1. Substitute 1 s−1 for the unit Hz:

λ=

3.00 × 108 m/s 4.0 × 1014 s−1

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How can you calculate the frequency of light using its wavelength and the speed of light? If you know the wavelength of a wave, you can calculate its frequency by dividing both sides of the speed-of-light equation by wavelength. The result is c ν= λ

Sample Problem 2 If a particular green light has a wavelength of 4.9 × 10–7 m, what is its frequency? First, substitute the value for the wavelength into the equation. c

ν=

4.9 × 10–7 m

Next, substitute the speed of light for c.

ν=

Next, divide 3.00 × 108 by 4.9 × 10–7. Recall that the exponent of the denominator, –7, is subtracted from the exponent of the numerator, 8. Also notice that the units m cancel leaving 1 , which is written as s–1. s ν = 0.61 × 1015 s–1 or 6.1 × 1014 s–1 Recall that s–1 is 1 Hz. Therefore, green light with a wavelength of 4.90 × 10–7 m has a frequency of 6.1 × 1014 Hz.

3.00 × 108 m/s 4.9 × 10–7 m

How can you calculate the energy of light from its frequency using Planck’s constant? Physicist Max Planck discovered a relationship between the energy of a photon and frequency of the associated electromagnetic wave. He determined that the energy of radiation E was equal to the product of a constant h and the frequency of emitted or absorbed radiation ν. E = hν The constant h is known as Planck’s constant. Planck’s constant is equal to 6.626 × 10–34 J∙s, read as joule-seconds. To determine the units of energy, multiply the unit for h, J∙s, by the unit for ν, s−1, to obtain J, the quantity of joules.

Sample Problem 3 The human eye can see light with a frequency about as high as 7.9 × 1014 Hz, which appears violet. Calculate the energy that one photon of violet light carries.

Next, substitute the value for the constant, h, into the equation and change Hz to s−1. E = (6.626 × 10−34 J∙s)(7.9 × 1014 s–1) Multiply 6.626 × 10−34 J∙s and 7.9 × 1014 s–1. Recall that s and s–1 cancel.

First, write the equation to solve for:

E = hν

Then,substitute the frequency into the equation for energy.

E = 52.3 × 10−20 J or 5.2 × 10−19 J One photon of violet light carries 5.2 × 10−19 J of energy.

E = h(7.9 × 1014 Hz)

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TEKS

End-of-Course Assessment Review

1. Calculate  Laser light is found in many types of technology including DVD players and surgical tools. If a laser emits light with a frequency of 5.0 × 1014 Hz, what is the wavelength of the light? A 1.7 × 106 m B 1.3 × 10–34 m C 1.5 × 1023 m D 6.0 × 10–7 m 2. Calculate  An argon laser emits light at a wavelength of 489 nm (1 nm = 1 × 10–9 m). What is the frequency of this light? A 6.13 × 1014 s–1 B 6.13 × 105 s–1 C 6.13 × 1016 s–1 D 6.13 × 1018 s–1 3. Calculate  Ultraviolet (UV) radiation has enough energy to penetrate the top layer of human skin, causing sunburn. How much energy is carried by one photon of ultraviolet radiation with a frequency of 1.0 × 1014 Hz? A 6.626 × 10−22 J B 6.626 × 10−48 J C 6.626 × 10−20 J D 3.00 × 1022 J 4. Draw Conclusions  Visible light has frequencies between about 4 × 1014 Hz and about 7.9 × 1014 Hz. Can electromagnetic radiation with a wavelength of 3.00 × 10−7 m be seen with human eyes? Show the calculation that supports your answer. 5. Apply Concepts  How would you calculate the energy carried by a photon of light that has a wavelength of 4.9 × 10−7 m? Which equation or equations would you need? Set up and perform the calculations, showing your work.

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TEKS REVIEW

6D

Calculating Average Atomic Mass TEKS 6D

Use isotopic composition to calculate average atomic mass of an element. TEKS_TXT

What is an isotope?

Vocabulary atomic number isotopes mass number atomic mass unit (amu) atomic mass

Recall that atoms contain particles called protons, neutrons, and electrons. Atoms of the same element have the same number of protons. For example, two atoms that each contain 6 protons are both carbon atoms. The number of protons in an atom represents its atomic number. An element’s atomic number is listed in the element’s square on the periodic table. Atoms with the same numbers of protons can differ in their numbers of neutrons, however. For example, atoms of the element carbon can have 5, 6, 7, or 8 neutrons. Atoms of the same element with different numbers of neutrons are called isotopes. The mass number of an atom is the sum of the atom’s protons and neutrons. The mass number is used to identify an isotope and is written after the element name. For example, carbon-14 identifies an isotope of carbon with 6 protons and 8 neutrons. The different isotopes of carbon are shown in Figure 1.

Figure 1

Carbon Isotopes Carbon-11

Carbon-12

Carbon-13

Carbon-14

Number of protons

6

6

6

6

Number of neutrons

5

6

7

8

11

12

13

14

Mass number

How is isotopic composition used to calculate average atomic mass of an element? Chemists describe the mass of an atom using a unit that is based on the mass of a carbon-12 atom. An atomic mass unit (amu) is defined as 1 the 12 mass of a carbon-12 atom and is equal to 1.66054 × 10–24 g. The mass of  a carbon-12 atom is exactly 12 amu. The mass of a carbon-13 atom is 13.003355 amu. Each element has several isotopes. Rather than listing the mass numbers for all isotopes, the periodic table lists the average atomic mass of each element. The atomic mass of an element is an average mass that is weighted based on the abundance of the isotopes, or isotopic composition, of the element. TEKS 6D • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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Some isotopes are more abundant than others. For example, about 98.93 percent of all carbon atoms exist as carbon-12 atoms. About 1.07 percent of carbon atoms exist as carbon-13, and less than 0.01 percent exist as other isotopes. The atomic masses and abundances of the two most abundant isotopes of carbon are shown in Figure 2. Figure 2

Isotopic Composition of Carbon Isotope

Study Tip To convert a percentage to decimal form, divide the percentage by 100 (move the decimal point in the percentage to the left two places).

Atomic mass (amu)

Carbon-12

12

Carbon-13

13.003355

Approximate abundance (percent) 98.93 1.07

To calculate the average atomic mass of carbon, multiply the atomic mass (in amu) of each isotope by its percentage abundance in decimal form. Then add the products together. (Because the abundances of carbon-11 and carbon-14 are less than 0.01 percent, their impact is very small and they are not included in the calculation.) (12 amu) (0.9893) + (13.003355 amu) (0.0107) = 11.87 amu + 0.1391 amu = 12.01 amu The resulting value, 12.01 amu, is the average atomic mass of carbon. In other words, if you take any naturally occurring sample of carbon, it will contain enough atoms that are carbon-13 isotopes and enough atoms that are carbon-12 isotopes so that the average atomic mass for the sample is 12.01 amu. The two most abundant isotopes of chlorine are chlorine-35 and chlorine-37. Their atomic masses and isotopic abundances are provided in Figure 3.

Figure 3

Isotopic Composition of Chlorine Isotope

Atomic mass (amu)

Approximate abundance (percent)

Chlorine-35

34.969

75.78

Chlorine-37

36.966

24.22

How can you calculate the average atomic mass of chlorine? Start by multiplying the atomic mass of each isotope by its abundance. Then add the products. (34.969 amu) (0.7578) + (36.966 amu) (0.2422) = 26.50 amu + 8.953 amu = 35.45 amu The average atomic mass of chlorine is 35.45 amu.

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TEKS

End-of-Course Assessment Review

1. Calculate  One atom contains 29 protons and 34 neutrons. Another atom of the same element has a mass number of 65. How many protons and neutrons does this unknown atom have? A 28 protons, 37 neutrons B 29 protons, 36 neutrons C 29 protons, 35 neutrons D 31 protons, 34 neutrons 2. Calculate  The two naturally occurring isotopes of hydrogen are hydrogen-1 (1.00783 amu; 99.9885 percent) and hydrogen-2 (2.01410 amu; 0.0115 percent). What is the average atomic mass of hydrogen? A 1.0079 amu B 1.0097 amu C 1.0309 amu D 1.5110 amu 3. Calculate  The most abundant isotopes of silicon are silicon-28, silicon-29, and silicon-30. Given the atomic masses and abundances in the table below, what is the average atomic mass of silicon?

Isotopic Composition of Silicon Isotope

Atomic mass (amu)

Approximate abundance (percent)

Silicon-28

27.97693

92.23

Silicon-29

28.97649

4.68

Silicon-30

29.97377

3.09

4. Infer  Fluorine has an average atomic mass of 18.998. What is likely the most abundant isotope of fluorine? Use your inference to identify the following data: the isotope’s mass number, atomic number, and number of neutrons. 5. Evaluate  Why are the atomic masses and mass numbers for each isotope not listed on the periodic table?

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TEKS REVIEW

6E

Electron Configurations and Lewis Valence Electron Dot Structures

Re adi n e ss

TEKS 6E Express the arrangement of electrons in atoms through electron configurations and Lewis valence electron dot structures.

How do quantum numbers describe atomic orbitals?

Vocabulary atomic orbital

As explained by modern atomic theory, the locations of electrons aren’t exact. Instead, mathematical expressions called atomic orbitals describe the probability of finding an electron at various locations around the nucleus. Atomic orbitals differ by size, shape, and energy. The energy levels of electrons are labeled by principal quantum numbers (n). These numbers have positive integer values of 1, 2, 3, and so on. Electrons with the same principal quantum number are in the same principal energy level. Each level contains one or more sublevels. Each sublevel contains one or more atomic orbitals. The second quantum number describes the shape of the atomic orbitals in a sublevel. Each shape is denoted by a letter. The number of sublevels is equal to the principal energy level number. For example, level n = 1 has one sublevel, the s sublevel. Level n = 2 contains two sublevels—the s and p sublevels. Level n = 3 has three sublevels—s, p, and d. The third quantum number describes the orientation of the orbital in space. It also describes the number of orbitals in a particular sublevel. The s sublevels contain one orbital, p sublevels contain three orbitals, d sublevels contain five orbitals, and f sublevels contain seven orbitals. The sublevels and atomic orbitals in each principal energy level are shown in Figure 1. Each atomic orbital may describe a maximum number of two electrons, each with opposite spin direction. This is called the Pauli exclusion principle. Figure 1

Atomic Orbitals and Electrons in Principal Energy Levels Principal energy level

Type of sublevel

Number of orbitals in sublevels

Maximum number of electrons

1

s

1

2

2

s, p

1+3=4

8

3

s, p, d

1 + 3 +5 = 9

18

4

s, p, d, f

1 + 3 + 5 + 7 = 16

32

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How is the arrangement of electrons expressed by electron configurations? An electron configuration describes which atomic orbitals hold the atom’s electrons. Electrons occupy atomic orbitals with lowest energies first. This is called the aufbau principle. You can see an illustration of this principle in Figure 2, which is called an aufbau diagram. Figure 2 Aufbau Diagram

6p 6s 5p 5s

5d

4f

4d

Increasing energy

4p 3d

4s 3p 3s 2p 2s

1s

The electron configuration for hydrogen is 1s1. The first numeral 1 indicates the principal energy level, the letter s indicates the sublevel and type of atomic orbital, and the superscript 1 describes the number of electrons in the s orbital. The electron configuration for an atom of helium (He), which has two electrons, is 1s2. In the second row of the periodic table, the second principal energy level (n = 2) is filled. The electron configuration for lithium (Li) is 1s22s1. An atom of beryllium (Be) has four electrons, and its configuration is 1s22s2. The fifth electron in boron (B) begins to fill the p orbitals. Boron has an electron configuration of 1s22s22p1. Recall that a p sublevel contains three p orbitals and each p atomic orbital can hold two electrons. The electron configurations of atoms of carbon (C), nitrogen (N), oxygen (O), fluorine (F), and neon (Ne) fill the remaining p orbitals. The electron configuration for neon, with 10 electrons, is 1s22s22p6.

Study Tip To remember the order of the types of atomic orbitals, you can use a mnemonic such as “Saint Paul did fall,” for s, p, d, f.

The n = 3 level contains three sublevels—s, p, and d. Similar to the n = 2 level, electrons fill s and p orbitals first. However, electrons fill the 4s orbital before filling the 3d orbital. The electron configuration for an atom of scandium (Sc), with 21 electrons, is 1s22s22p63s23p64s23d1. After the 3d orbitals are filled, electrons fill the 4p orbitals.

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How is the arrangement of valence electrons expressed by Lewis valence electron dot structures? Valence electrons are electrons in the outermost principal energy level of an atom. These are the s and p electrons with the highest n value. For example, neon has an electron configuration of 1s22s22p6. The outermost principal energy level is n = 2, and it contains 8 electrons (2 electrons in the 2s orbital and 6 electrons in 2p orbitals). Therefore, neon has 8 valence electrons. Because the s and p orbitals in a principal energy level can contain a maximum of 8 electrons, the maximum number of valence electrons is 8. Lewis valence electron dot structures represent valence electrons with the symbol of the element surrounded by a dot for each valence electron. The first four dots are arranged individually on four sides of the symbol. Each additional dot is paired with one of the first four dots. Figure 3 Lewis Valence Electron Dot Structures

Carbon

Neon

Scandium

C

Ne

Sc

You can see some examples in Figure 3. Carbon has four valence electrons, so it has four dots. Neon has eight. Note that scandium has an electron configuration of is 1s22s22p63s23p64s23d1. The outermost electrons are in principal energy level n = 4, which contains 2 electrons.

TEKS

End-of-Course Assessment Review

1. Identify  What is the correct Lewis valence electron dot structure for silicon (Si), which has 14 electrons? A Si

B Si C Si D

Si

2. Explain  Is it possible for an atom to have an incomplete principal energy level before starting to fill the next larger level? Explain why or why not. 3. Explain  An atom of gallium (Ga) has 31 electrons. Express the electron arrangement of gallium using an electron configuration and a Lewis valence electron dot structure. 4. Infer  Suppose an atom has a completely filled n = 2 principal energy level. At minimum, which sublevels are filled, and how many electrons are in each sublevel? 5. Apply Concepts  A shorthand notation for electron configurations uses the expressions [He], [Ne], or [Ar] to represent the configurations of an atom’s inner-level electrons. For example, the electron configuration for lithium, which is 1s22s1, can also be written as [He]2s1. Using this notation, write the electron configuration for an atom of magnesium (Mg), which has 12 electrons. TEKS 6E • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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TEKS REVIEW

7A

Naming Compounds

Re adi n ess

TEKS 7A Name ionic compounds containing main group or transition metals, covalent TEKS_TXT acids, and bases, using International Union of Pure and Applied Chemistry compounds, (IUPAC) nomenclature rules.

How do you name compounds?

Vocabulary covalent bond acid base

All compounds must be electrically neutral, that is, the sum of the charges must equal zero. Compounds achieve neutrality by having an equal number of positive and negative charges. When a sodium ion (Na+) and a chloride ion (Cl-) combine, they do so in a 1:1 ratio, resulting in a neutral compound NaCl. Compounds are named according to the types of elements that form them. Ionic compounds are named by one method; covalent compounds by another method.

How are ionic compounds named using IUPAC rules? An ion forms when an atom or a group of atoms loses or gains electrons. If the ion has a positive charge, it is a cation. If it has a negative charge, it is an anion. When ionic bonds form between cations and anions, an ionic compound is formed. The names of ionic compounds convey information about the ions in the compound as well as information about the bonds holding them together. The positively charged particle (often a metallic ion) is placed first. The negatively charged ion will end the formula. For example, a compound containing a potassium ion and an ion of chlorine would begin as potassium. The name of the negative ion is changed to end in -ide, making the negative ion chloride. The name of the compound is potassium chloride. Generally, elements from Group 1 form +1 ions, elements from Group 2 form +2 ions, and elements from Group 3 form +3 ions. Some metals have more than one common oxidation state. For example, iron can have either a +2 or +3 oxidation number. In these cases, you include a roman numeral in parentheses to indicate the charge. Some examples are shown in Figure 1. Figure 1 Common Ion Names

Ion Name Iron(II) Iron(III) Copper(I) Copper(II)

Ion Fe

2+

Fe3+ Cu+ Cu2+

Ion Name

Ion

Silver(I) Silver(II) Mercury(I) Mercury(II)

Ag+ Ag2+ Hg+ Hg2+

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Study Tip To identify the charge on a cation that can have different possible charges, use the formula of the compound. You can usually find the charge of each ion by the amount of the other ion in the compound. For example, in Fe2O3, to balance the three –2 charges from the oxygen, the charge on the iron must be +3.

To name compounds containing polyatomic ions, you can use the same method. When the positive portion is a metal, use the metal name and name of the negative polyatomic ion. Most polyatomic anions end in -ite or -ate. For example, the name of KNO3 is potassium nitrate. For example, how would you name the compound FeO? Since oxygen almost always forms a –2 charge, iron must have an equal and opposite charge. Iron in FeO must have a +2 charge. The compound’s name is iron(II) oxide.

How are covalent compounds named using IUPAC rules? Atoms in a covalent compound are held together by covalent bonds— bonds formed by the sharing of electrons between atoms. A covalent bond forms between two atoms with high electronegativity, usually nonmetals. Because their attractions for a given electron are about equally strong, they share a pair of electrons. This sharing pulls the atoms into close contact with each other. When two nonmetallic elements combine, they often do so in more than one way. For example, carbon and oxygen can combine to form CO and CO2. In order to distinguish between these two compounds, the IUPAC nomenclature rules makes use of prefixes. Some of these prefixes are shown in Figure 2.

Figure 2

Prefixes used for naming covalent compounds Number

Prefix

1

mono- *

2

di-

3

tri-

4

tetra-

5

penta-

6

hexa-

7

hepta-

8

octa-

9

nona-

10

deca-

*The prefix mono- is only used for the more electronegative element.

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How are acids and bases named? Acids are substances that contain one or more hydrogen atoms and donate a proton in a reaction. Acids are named based on the suffix of the anion. Figure 3 summarizes the three rules that can be used to name acids with the general formula HnX dissolved in water. Figure 3

Naming Common Acids Anion ending

Example

-ide

chloride, Cl

-ite

sulfite, SO3

-ate

nitrate, NO3



− −

Acid name

Example

hydro-(stem)-ic acid

hydrochloric acid

(stem)-ous acid

sulfurous acid

(stem)-ic acid

nitric acid

Bases are compounds that produce hydroxide ions when dissolved in water. Bases are named in the same way as other ionic compounds. The name of the cation is given first, followed by hydroxide, for example, barium hydroxide (Ba(OH)2).

TEKS

End-of-Course Assessment Review

1. Name  What is the name of the compound with the formula H2SO4? A hydrosulfuric acid B sulfur hydroxide C dihydrogen sulfate D hydrosulfurous acid 2. Analyze  In a sample of Al(NO3)3, the ratio of aluminum (Al) ions to nitrate ions (NO3) is A 1:1 B 1:2 C 1:3 D 1:6 3. Interpret  A compound has the name iodine pentafluoride. What kind of bond is present between the atoms? A ionic B covalent C metallic D hydrogen 4. Name  Write the formula for the salt calcium bromide.

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TEKS REVIEW

7B

Writing Chemical Formulas

Readi ness

TEKS 7B Write the chemical formulas of common polyatomic ions, ionic compounds containing main group or transition metals, covalent compounds, acids, and bases.

How do you write the chemical formulas of common polyatomic ions?

Vocabulary polyatomic ion

Many ionic compounds are made of polyatomic ions. A polyatomic ion is a tightly-bound group of atoms that has a positive or negative charge and behaves as a unit. There is no simple set of rules for naming most polyatomic ions, so you should memorize the chemical formulas for some of the most common ones, such as those in Figure 1. Figure 1

Chemical Formulas for Some Common Polyatomic Ions Cations

Anions

Name

Formula

Ammonium

NH4

Hydronium

H3O

+ +

Anions Name

Formula

Name

Formula

Hydroxide

OH

Nitrite

NO2−

Nitrate

NO3−

Phosphate

PO43−



Carbonate

CO32−

Sulfite

SO32−

Hydrogen carbonate

HCO3−

Sulfate

SO42−

How do you write the chemical formulas of ionic compounds?

Study Tip It might help you to remember the formula for oxygen-containing ions that also contain sulfur and nitrogen by remembering that the ion with more oxygen atoms is the one whose name has the suffix -ate. The ion with fewer oxygen atoms is the one whose name has the suffix -ite.

When two elements combine, forming a compound, the product is called a binary compound. Binary compounds can be ionic compounds or covalent compounds. If you know the name of a binary ionic compound, you can write the formula. To write the formula of a binary ionic compound, first write the symbol of the cation and then the anion. Then add subscripts as needed to balance the charges. The positive charge of the cation must balance the negative charge of the anion so that the net ionic charge of the formula is zero. For example, how could you determine the formula for the compound beryllium chloride? Beryllium chloride is made up of beryllium cations (Be2+) and chloride anions (Cl–). The two ions do not have equal numerical charges. Thus each beryllium ion with its +2 charge must combine with (or be balanced by) two chloride ions, each with a –1 charge.

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That means that the ions must combine in a 1:2 ratio, so the formula for beryllium chloride is BeCl2. The net ionic charge of the formula unit is zero. How could you determine the formula for the compound mercury fluoride? Mercury has a variable charge, so the chemical name indicates that in this compound mercury has a +1 charge. Fluoride is in Group 17 on the periodic table, so it has a –1 charge. Because they have equal but opposite charges, they will combine in a 1:1 ratio, so the formula is HgF.

How do you write the chemical formulas of covalent compounds? Because the names of covalent compounds usually indicate the amounts of each atom as a prefix, writing the formula is often a matter of taking the prefixes and translating them into subscripts after each element symbol. For example, how could you write the formula for dinitrogen difluoride? You know that the prefix di- means two. The covalent compound is made up of nitrogen and fluoride. So the formula is N2F2. Some common covalent compounds are not based on the rules for naming compounds. For example, water (H2O) and ammonia (NH3) are not named according to the conventional system. You will need to memorize molecular formulas such as these.

How do you write the chemical formulas of acids and bases? If you know the name of an acid, you can write its formula by using the rules for writing the name of an acid, in reverse. Then balance the ionic charges just as you would for any ionic compound. Recall that the IUPAC naming system depends on the suffix in the name of the anion. Each rule deals with an anion with a different suffix: -ide, -ite, and -ate. • When the name of the anion ends in -ide, the acid name begins with the prefix hydro-. The stem of the anion has the suffix -ic and is followed by the word acid. Therefore, HCl (X = chloride) is named hydrochloric acid. • When the anion name ends in -ite, the acid name is the stem of the anion with the suffix -ous, followed by the word acid. Thus, H2SO3 (X = sulfite) is named sulfurous acid. • When the anion name ends in -ate, the acid name is the stem of the anion with the suffix -ic, followed by the word acid. Thus, HNO3 (X = nitrate) is named nitric acid. Bases are indicated by the term hydroxide in their names. The formula of a base includes a cation and at least one hydroxide ion. The subscript following the hydroxide ion is determined by the charge of the cation. So, for example, how could you write the formula for aluminum hydroxide? Aluminum is the cation. Hydroxide (OH–) is the anion. Three hydroxide anions are needed to balance the +3 charge on the aluminum ion, so the formula is Al(OH)3.

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TEKS

End-of-Course Assessment Review

1. Identify  Which is the correct formula for the compound dinitrogen monoxide? A N–O B N2O2 C NO2 D N2O 2. Identify  Hydrosulfuric acid is a colorless, poisonous gas formed from the reaction of hydrogen gas and molten sulfur. What is its chemical formula? A HS− B H2S C HSO4− D H2SO4 3. Identify  Phosphoric acid contains which polyatomic ion? A P3+ B PO43– C PO33– D P2O74–

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TEKS REVIEW

7C

Constructing Electron Dot Formulas

Re adi n ess

TEKS 7C

Construct electron dot formulas to illustrate ionic and covalent bonds. TEKS_TXT

Vocabulary electron dot formula valence electron

How can you construct electron dot formulas? Different kinds of bonds often form between different atoms. A simple way to show how electrons are transferred or shared during bond formation is with an electron dot formula. Learning how to draw and interpret these diagrams is an important skill. An electron dot formula (or diagram) consists of a chemical symbol surrounded by one to eight dots representing valance electrons. Valence electrons are the electrons in the highest occupied energy level of an element’s atoms. All of the elements within a given group (with the exception of helium) have the same number of electron dots in their structures. Hydrogen has 1 valence electron, so you would write the electron dot formula for hydrogen as: H Iodine has 7 valence electrons, so you would write the electron dot formula for iodine as: I Ions can be represented in electron dot formulas by depicting the charge and the corresponding number of valence electrons. For example, iodide (I−), which has gained one electron, has eight valence electrons, shown below. I



Magnesium ion (Mg2+), which has lost both of its valence electrons, is illustrated below without any dots. Mg2+

How can you construct electron dot formulas that illustrate ionic bonds?

The electrons involved in an ionic bond are removed from the cation and found on the anion. Electron dot formulas illustrate this by indicating oppositely charged ions side by side.

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For example, you would write the electron dot formula for sodium chloride (NaCl) as shown. Na+ Cl



The electron dot formula for aluminum oxide (Al2O3) is written as: Al3+ Al

3+

O

2−

O

2−

O

2−

How can you construct electron dot formulas that illustrate covalent bonds? Study Tip When you are constructing electron dot formulas for covalent molecules, the atoms that often form double bonds are C, N, O, and S. You can remember these with the mnemonic Snow CONes form double bonds.

Covalent bonds are formed from two electrons that are shared between two atoms. They can be depicted in electron dot formulas as two dots, or one short straight line, such as in the bond between two hydrogen atoms. HH H H Single bonds are represented by single lines. Double bonds are represented by double lines. Triple bonds are represented by triple lines. H

H C

H

H

C

C

C

H

H

How can you construct electron dot formulas for more complex covalent molecules or ions? Use these steps to help you construct the electron dot formulas for more complex covalent molecules or ions. 1. Add the total number of valence electrons for every atom in the molecule. If it is an ion, add or subtract electrons to produce the correct charge. 2. Write the structure of the skeleton. The more electronegative atom usually belongs in the center. Connect atoms with lines (or pairs of dots). 3. Distribute electrons to each of the outer atoms to satisfy the octet rule. 4. Distribute the remaining electrons to the central atom. Is the octet rule satisfied for the central atom? If not, then there is probably at least one multiple bond in the molecule or ion. Move pairs of electrons from neighboring atoms to form double bonds with the central atom until the octet rule is satisfied for all of the atoms in the molecule.

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For example, you would write the electron dot formula for nitrogen trifluoride (NF3) as follows: F N

F

F The electron dot formula for carbon dioxide (CO2) would be written as: O

TEKS

C

O

End-of-Course Assessment Review

1. Interpret Diagrams  Which of the following is the correct electron dot formula for ethane (C2H6), one of the components of natural gas? A

H

H C

H

H

C

C

H

H H H C C H H H

H C

3+

B

C

3+

H2− H2−

D

H

H2−



H

H

C

C

H

H

H

2. Predict  A chemist heats a sample of beryllium and inserts it into a flask containing chlorine gas. Which of the following electron dot formulas correctly represents the product of the resulting reaction? A

B

Cl

Be

Be

2+

Cl

Cl



Cl



C

Cl−

D

Cl

Be2+ Be

Cl − Cl

3. Construct  Magnesium carbonate (MgCO3) is a component of chalk. Construct the electron dot formula for this ionic compound. 4. Construct  Hydrogen cyanide (HCN) is a poisonous liquid sometimes called “prussic acid.” Construct the electron dot formula for this compound.

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7D

The Nature of Metallic Bonding TEKS 7D Describe the nature of metallic bonding and apply the theory to explain metallic properties such as thermal and electrical conductivity, malleability, and ductility.

How can you describe the nature of metallic bonding?

Vocabulary metallic bond malleable ductile alloy

Metallic atoms have few valence electrons and low ionization energies. The bonds holding metallic atoms together in the solid and liquid phases, however, are apparently strong, as metals have fairly high melting and boiling points. A metallic atom may be considered to have a central portion, or core, made up of its nucleus and its nonvalence electrons. The atom’s valence electrons surround the core. The cores of the metallic atoms making up a metallic solid are arranged in the fixed positions of a crystalline lattice. The valence electrons of metal atoms can be modeled as a “sea of electrons.” That is, the valence electrons are mobile and can drift freely from one part of the metal to another. Metallic bonds consist of the attraction of the free-floating valance electrons for the positively charged metal ions. These bonds are the forces of attraction that hold metals together. A representative of this model in shown in Figure 1.

Figure 1 The “Sea of Electrons” Model

Sea of electrons Metal cation





− − −

+ + + −





− −

+ + +− − −



+ + + − −



How can you apply metallic bonding theory to explain metallic properties? The nature of metallic bonding explains many physical properties of metals. For example, most metals are excellent conductors of thermal energy. When a difference in thermal energy is applied across a metal, it is quickly and evenly transmitted throughout. This is because the crystal lattice of cations making up a metal is held together very tightly. Thermal energy makes cations move about more rapidly. Since the lattice is held together so tightly, the energy of motion is quickly distributed throughout the solid.

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Study Tip If you remember that metals have a tendency to give up electrons, then it will be easy to remember the nature of metallic bonds as positively charged ions in a “sea” of electrons.

Most metals are also excellent conductors of electrical energy. When a difference in electrical energy is applied across a metal, it is quickly and evenly distributed throughout. This happens because the electrons flow freely. Electrons move in response to the difference in electrical energy, moving from high electric potential energy to low electric potential energy. Metals are also malleable, which means that they can be shaped and hammered into thin sheets. A force, such as the strike of a hammer, applied to the solid reshapes the lattice of cations because the cations can move through the “sea” of electrons without breaking the metallic bonds. For this same reason, metals are also ductile, which means that they can be drawn into long wires without breaking. Another property of metals is their very high boiling points. Figure 2 lists the boiling points of some metals from period 4 of the periodic table. It takes a great deal of thermal energy to break the bonds holding the ions of a metal together.

Figure 2

Boiling Points of Selected Period 4 Metals Element

Boiling Point (Kelvins)

Potassium

1032

Calcium

1757

Scandium

3103

Titanium

3560

Vanadium

3680

Chromium

2944

Manganese

2334

Iron

3134

Cobalt

3200

Nickel

3186

Copper

3200

Zinc

1180

What are alloys? Metallic bonds bind elemental metals, such as pure sodium. However, metallic bonds also bind alloys of two or more elements. Alloys are mixtures composed of two or more elements, at least one of which is a metal. Brass, for example, is an alloy of copper and zinc. Alloys are important because their properties are often superior to those of their component elements. The most important alloys today are steels. The table in Figure 3 on the next page lists some common alloys and the different elements that compose them.

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Figure 3

Common Alloys Composition

Common Name of Alloy

Common Uses

Copper and zinc

Brass

Locks, gears, musical instruments

Copper and tin

Bronze

Sculptures, bearings, springs

Iron and chromium

Stainless steel

Cookware, appliances

Tin and copper

Pewter

Statuettes, figurines

Silver and copper or other metal such as zinc

Sterling silver

Fine tableware, jewelry

TEKS

End-of-Course Assessment Review

1. Evaluate  Which of the following statements is true? A Malleability is caused by the movement of electrons. B Ductility is caused by a strong lattice of cations. C Malleability describes the ability of a metal to be drawn into long wires. D Ductility describes the ability of a metal to be drawn into long wires. 2. Apply Concepts  Which of the following statements is best explained by the nature of metallic bonding? A Electricity from a bolt of lightning travels through salt water. B The metal handle of a frying pan becomes too hot to touch soon after it is placed upon a stove. C Iron rusts in the presence of oxygen. D As a solid, calcium phosphate (Ca3(PO4)2) forms a crystal pattern. 3. Compare and Contrast  How are thermal conductivity and electrical conductivity similar and how are they different? 4. Describe  How does the nature of metallic bonding affect the physical properties of metals? 5. Evaluate  A student states that metallic bonds and ionic bonds are very similar. In what ways do you agree and in what ways do you disagree? Justify your answer.

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TEKS REVIEW

7E

Molecular Structures TEKS 7E Predict molecular structure for molecules with linear, trigonal planar, or tetrahedral electron pair geometries using Valence Shell Electron Pair Repulsion (VSEPR) theory.

Vocabulary Valence Shell Electron Pair Repulsion (VSEPR) theory

Figure 1 Nitrogen Trifluoride (left) and Aluminum Trifluoride (right).

Why is molecular structure important? The actual three-dimensional shape of a molecule or ion is its molecular structure. Molecular structure includes the angles of each of the bonds with respect to the others. For example, consider nitrogen trifluoride (NF3) and aluminum trifluoride (AlF3). Their chemical formulas are similar, and their electron dot structures suggest nearly identical two-dimensional shapes, shown in the figures below. The only difference between the two electron dot structures is the presence of a lone pair of electrons on nitrogen, which affects the shape of the entire molecule. Experimental evidence of their chemical properties shows that they have very different molecular structures. F

F N

Al

F

F

F

F

How can you predict molecular structures?

Study Tip The names of the electron pair geometries can help you to remember their molecular shapes. Linear molecules are in the shape of a line. Trigonal planar contains the root tri-, which refers to three. The root tetra- refers to four, so tetrahedron refers to a structure with four plane faces.

Different methods are used to analyze how the valence electrons in molecules and ions affect their overall structures. The simplest way involves the Valence Shell Electron Pair Repulsion theory. The Valence Shell Electron Pair Repulsion (VSEPR) theory states that the repulsion between electron pairs causes molecular shapes to adjust so that the valence-electron pairs stay as far apart as possible. Valence electrons, like all electrons, repel each other. The repulsions of valence electrons affect the angles of the covalent bonds, giving the overall shape to each molecule. The electrons in covalent bonds repel the electrons in other covalent bonds, pushing them away. The electrons in lone pairs, such as in the pair on the nitrogen atom of nitrogen trifluoride, as a result, occupy even more space than the electrons in a covalent bond. The electrons of lone pairs push away the other electrons in the covalent bonds of a molecule, which changes the entire structure of the molecule.

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How can I use VSEPR theory to predict geometric shapes? Once you have drawn the electron dot structure of a molecule and identified any lone pairs of electrons, you can use the VSEPR theory to help determine the geometric shape of the molecule. Since they repel each other, the electron pairs are as far away as possible from each other. The shape of a molecule depends on the arrangement of electron pairs. Linear Arrangement  If there are just two groups of valence electrons on an atom in a molecule, they tend to be on opposite ends of the atom. This gives the molecule a linear electron pair geometry, as shown below. The two groups could either be single electron pairs, or a double or triple bond. Figure 2 Linear Arrangement

Trigonal Planar Arrangement  If there are three electron pairs on an atom in a molecule, they tend to spread out from each other on the same plane. This gives the molecule a trigonal planar electron pair geometry (Figure 2). Figure 3 Trigonal Planar Arrangement

Tetrahedral Arrangement  If there are four electron pairs on an atom in a molecule, they tend to spread out from each other into four points that form a pyramid, as shown in Figure 4. This gives the molecule a tetrahedral electron pair geometry. Figure 4 Tetrahedral Arrangement

The presence of a lone electron pair on an atom pushes the other bonds further away, as you saw in the structure of nitrogen trifluoride. The repulsion of the electrons pushed the fluoride atoms toward each other. The lone pair of electrons can be visualized as a cloud extending away from the molecule, shown below. H 109.5°

H

C

H

H

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TEKS

End-of-Course Assessment Review

1. Compare and Contrast  What is the difference between a molecule with a trigonal planar electron arrangement and one with a tetrahedral arrangement? A Trigonal planar refers to four equal faces. B A tetrahedral arrangement has corners with equal angles. C A tetrahedral arrangement has three equal faces. D A tetrahedral arrangement has four equal faces. 2. Predict  Use the VSEPR theory to determine the electron configuration of methane (CH4), a major component of natural gas. A linear B trigonal planar C tetrahedral D triangular 3. Predict  Which compound has the molecular structure shown below? Explain your answer.

4. Predict  The compound trichlorofluoromethane (Cl3FC) is a gas that was once widely used in refrigerants until it was discovered that it was damaging the ozone layer. Draw the molecular structure and identify the electron arrangement.

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TEKS REVIEW

8A Defining and Using the Mole TEKS 8A

Define and use the concept of a mole.

How do chemists define a mole?

Vocabulary mole Avogadro’s number representative particle

Atoms and molecules are so small that even a small amount of a substance consists of an extremely large number of them. A drop of water, for example, contains trillions of water molecules. In chemistry, the number of particles of a substance is expressed by a unit called a mole, which is the SI unit for the amount of a substance. It is abbreviated as mol. A mole of a substance is 6.02 × 1023 representative particles of that substance. The number of representative particles in a mole, 6.02 × 1023 , is called Avogadro’s number. When you refer to the particles in a mole of a substance, you are referring to representative particles of that substance. A representative particle is the type of particle that makes up the substance. It is the smallest particle into which the substance can be divided and still be that substance. As shown in Figure 1 below, representative particles might be atoms, molecules, or ions.

Figure 1

Representative Particles of Selected Substances Substance Lithium Iron Argon gas Bromine gas Sodium ion Sodium chloride Water

Representative particle Atom Atom Atom Molecule Ion Formula unit Molecule

Chemical formula

Number of representative particles per mole

Li Fe Ar Br2 Na− NaCl H2O

6.02 6.02 6.02 6.02 6.02 6.02 6.02

× 1023 × 1023 × 1023 × 1023 × 1023 × 1023 × 1023

The number of representative particles in a mole is always the same, but the number of atoms or ions varies. For example, a mole of lithium contains 6.02 × 1023 lithium atoms. A mole of iron contains 6.02 × 1023 iron atoms. Each formula unit of sodium chloride is composed of one ion of sodium and one ion of chloride. As a result, one mole of sodium chloride consists of one mole of sodium ions and one mole of chloride ions. A mole of water contains 6.02 × 1023 molecules. Because each molecule consists of three atoms, a mole of water consists of one mole of oxygen atoms and two moles of hydrogen atoms. As a result, a mole of water has three times as many atoms as a mole of lithium. TEKS 8A • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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How can you determine the representative particle of a given substance? For most elements, the representative particle is an atom. The diatomic elements hydrogen, nitrogen, oxygen, fluorine, chlorine, bromine, and iodine, however, exist normally as diatomic molecules. A molecule is the representative particle of these elements and all molecular compounds, such as water (H2O) and sucrose (C12H22O11). Ionic compounds, such as calcium carbonate (CaCO3) and magnesium oxide (MgO), are composed of large arrays of ions that are bonded together. Their representative particle, however, is a formula unit, consisting of the simplest whole-number ratio of ions in the substance.

How can you use moles to calculate the number of representative particles? The relationship 1 mol = 6.02 × 1023 representative particles is the basis for the following conversion factors that you can use to convert number of representative particles to moles: 1 mol 6.02 × 1023 representative particles You can also use the conversion factor below to convert from moles to number of representative particles. 6.02 × 1023 representative particles 1 mol If you know the number of moles of a substance in a sample, you can calculate the number of representative particles. Conversely, you can calculate the number of moles of a substance if you know the number of representative particles. These conversions can be useful in calculations involving the mass of a substance. Suppose a sample consists of 6.85 × 1020 atoms of carbon. How can you determine the number of moles?

Study Tip

First list what you know.

When dividing powers of 10, as long as the bases are the same, you subtract the exponents.

• number of atoms: 6.85 × 1020 atoms C • 1 mol C = 6.02 × 1023 atoms C Next, multiply the number of atoms of C by the conversion factor, and this will give you the answer. 6.85 × 1020 atoms C ×

1 mol C = 1.14 × 10−3 mol C 6.02 × 1023 atoms C

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Another sample consists of 2.58 moles of water, H2O. How can you determine the number of water molecules in the sample? 2.58 mol H2O ×

6.02 × 1023 molecules H2O 1 mol H2O

= 1.55 × 1024 molecules H2O

Notice that the representative particle in the first example is an atom. Because the conversion is from atoms to the number of representative particles, you divide by Avogadro’s number. The representative particle in the second example is a molecule. In this case the conversion is from moles to the number of representative particles, and you multiply by Avogadro’s number.

TEKS

End-of-Course Assessment Review

1. Define  Which number corresponds to a mole? A 6.02 × 10−23 B 6.02 × 1023 C 6.23 × 10−02 D 6.23 × 1002 2. Calculate  How many moles of carbon dioxide (CO2) are equivalent to 4.55 × 1024 molecules? A 4.55 moles B 7.56 moles C 15.1 moles D 22.7 moles 3. Apply Concepts  The representative particle of oxygen gas (O2) is a molecule. If NA is Avogadro’s number, which of the following calculations would you use to determine the number of representative particles in 6 moles of oxygen gas (O2)? A 3 × NA C 6 × NA B 3 ÷ NA

D 6 ÷ NA

4. Apply Concepts  Does a mole of sucrose (C11H22O11) have the same number of representative particles as a mole of magnesium (Mg)? Explain. 5. Evaluate  A chemist adds 9.0 moles of H2O to a mixture. In her journal she records the amount as 6.0 moles of hydrogen (H) and 3.0 moles of oxygen (O). Is she correct? Explain why or why not.

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TEKS REVIEW

8B

Calculating Atoms, Ions, or Molecules Using Moles

Re adi n ess

TEKS 8B Use the mole concept to calculate the number of atoms, ions, or molecules in a sample TEKS_TXT of material.

How can you use moles to calculate the number of atoms in a sample of material?

Vocabulary molar mass

You can use the mole concept to calculate how many particles make up the mass of a sample of material. In order to perform the calculation, you need to know the molar mass. Molar mass is the mass of one mole of the material. If the representative particle of a material is an atom, then the molar mass (in grams per mole) is numerically equal to atomic mass (in atomic mass units). As you see in Figure 1, each box on the periodic table gives the atomic mass of the element. Figure 1

13

Al

Atomic mass

Aluminum 26.982

The atomic mass for a single atom of aluminum is 26.982 atomic mass units. The following calculation shows how to use an element’s atomic mass in the conversion of mass to number of atoms. Suppose you needed to find the number of atoms in 7.85 g of aluminum. As shown below, first you would convert the mass to moles, and then you would convert the moles to atoms. 7.85 g Al ×

Study Tip Remember that to get from moles to mass, you multiply by the molar mass. Think “Ma’am” for “to get mass, multiply.” To get from mass to moles, you divide by the molar mass. Think “mod” for “to get moles, divide.”

1 mol Al 6.02 × 1023 Al atoms × = 1.75 × 1023 Al atoms 26.982 g Al 1 mol Al

You can also perform the inverse calculation and determine the mass of a given number of atoms of a material. Suppose you needed to find the mass of a certain number of atoms of magnesium. As shown below, first convert the number of atoms to moles. Then use the atomic mass of magnesium (atomic mass = 24.305) to find the mass of the moles. 1 mol Mg 24.305 g Mg × = 17.7 g Mg 4.38 × 1023 Mg atoms × 23 6.02 × 10 Mg atoms 1 mol Mg In both of these calculations, the key is to convert the quantity you know into moles. From moles, it is often easier to convert into other units.

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How can you use moles to calculate the number of molecules or formula units in a sample of material? The representative particle for a molecular compound is a molecule, and the representative particle for an ionic compound is a formula unit. In order to calculate the number of molecules or formula units that make up a given mass of a compound, you also have to know the compound’s formula. The formula identifies the number of atoms or ions in each representative unit. The following are examples: • Each formula unit of iron bromide (FeBr2) consists of one iron ion (Fe2+) and two bromide ions (Br–). • Each molecule of ammonia (NH3) consists of one atom of nitrogen (N) and three atoms of hydrogen (H). To calculate molar mass, recognize that a mole of Mg3N2 will contain three moles of magnesium and two moles of nitrogen. You can add these masses to determine the molar mass of a single mole of Mg2N3.

3 mol Mg ×

24.305 g Mg = 72.915 g 1 mol Mg

2 mol N ×



14.007 g N = 28.014 g 1 mol N

molar mass Mg3N2 = 100.929 g/mol



After you know the molar mass of the compound, you can calculate the number of molecules or ions from the mass of the compound, as shown in the example below for 16.2 g of the ionic compound magnesium nitride: 16.2 g Mg3N2 ×

1 mol Mg3N2 100.929 g Mg3N2

×

6.02 × 1023 mol Mg3N2 formula units 1 mol Mg3N2

= 9.66 × 1022 Mg3N2 formula units

First you calculate the number of moles by dividing the mass of the sample by the molar mass of the compound. You then calculate the number of formula units by multiplying by Avogadro’s number. If you wanted to, you could also perform the inverse calculations to determine the mass of a sample if you know the number of molecules or ions. To solve for this, you would divide by Avogadro’s number to solve for moles and then multiply the result by the compound’s molar mass to solve for mass.

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How can you use moles to calculate the number of ions in a sample of material? You can also use a compound’s formula to calculate the number of particles other than representative particles. In this case, you have to multiply by the subscript of the atom or ion in which you are interested. The following example shows how to calculate the number of ions in 16.2 g of magnesium nitride. 16.2 g Mg3N2 ×

1/mol Mg3N2 100.929 g Mg3N2

×

3 mol Mg 1 mol Mg3N2

×

6.02 × 1023 Mg− ions 1 mol Mg

= 2.90 × 1023 Mg− ions



TEKS

End-of-Course Assessment Review

1. Define  The molar mass of Al2Cl6 is the mass of one mole of the compound. It is also which of the following? A the mass of 1 mole of Al and 1 mole of Cl B the mass of 2 moles of Al and 6 moles of Cl C the mass of 1/2 mole of Al and 1/2 mole of Cl D the mass of 1/4 mole of Al and 3/4 mole of Cl 2. Calculate  How many atoms make up 3.29 g of silicon (Si)? A 2.14 × 1022 B 7.05 × 1022 C 6.02 × 1023 D 1.98 × 1024 3. Calculate  How many molecules make up 12.8 g of N2O4? A 8.37 × 1022 B 5.02 × 1023 C 7.71 × 1024 D 7.09 × 1026 4. Construct  Write a step-by-step procedure for converting the mass of an ionic compound to the number of formula units that make up the compound. 5. Evaluate  A student calculates the number of chloride ions (Cl–) in 7.0 g of aluminum chloride (AlCl3). He incorrectly states that the answer is 3.16 × 1022 Cl– ions. What mistake did the student most likely make? What is the correct answer?

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8C

Calculating Percent Composition, Empirical Formulas, and Molecular Formulas TEKS 8C

Calculate percent composition and empirical and molecular formulas.

How can you calculate percent composition?

Vocabulary percent composition empirical formula

The relative amount of mass of each element in a compound is its percent composition, or the percent by mass of each element in the compound. You can calculate a percent composition for each element that makes up a compound. The total percent compositions for a compound equal 100. For example, Figure 1 shows the percent compositions of the elements making up the compound potassium dichromate, K2Cr2O7. A 100-g sample of potassium dichromate has 26.5 g of potassium, 35.4 g of chromium, and 38.1 g of oxygen. That adds up to 100 g. A 200-g sample would have exactly twice as much of each element. The relative masses of the elements in a given compound are always the same, regardless of sample size. Potassium Dichromate (K2Cr2O7) Percent Compositions

Figure 1

K 26.5%

Cr 35.4 %

O 38.1%

To determine the percent composition of an element in a compound if you know the mass of the element in the compound, divide the mass of the element by the mass of the compound and multiply by 100 percent. % mass of element =

mass of element mass of compound

× 100%

Suppose, for example, that you combine 7.10 g of copper (Cu) with 17.9 g of bromine gas (Br2) to make 25.0 g of copper bromide (CuBr2). The percent compositions of the elements in the compound are:

% Cu =

7.10 g × 100% = 28.4% 25.0 g

% Br =

17.9 g × 100% = 71.6% 25.0 g

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You can also determine the percent compositions for elements of a compound if you know the compound’s chemical formula. First, determine the compound’s molar mass. Then, divide each part by the compound’s molar mass. % by mass of element =

mass of element in 1 mol compound molar mass of compound

× 100%

For the compound MgCl2, for example: 1 mol Mg ×

24.305 g Mg = 24.305 g 1 mol Mg

2 mol Cl ×

34.453 g Cl = 70.906 g 1 mol Cl

molar mass MgCl2 = 95.211 g/mol



The percent compositions of the elements, then, are: % mass of Mg = % mass of Cl =

24.305 g/mol × 100% = 26% 95.211 g/mol

2(35.453 g/mol) × 100% = 74% 95.211 g/mol

How can you calculate an empirical formula? An empirical formula is a formula with the lowest whole-number ratio of elements in a compound. It is not always the same as the molecular formula, which tells exactly how many atoms of each element are contained in a molecule or a formula unit of a compound. For example, methane’s molecular formula is C2H6, and its empirical formula is CH3. If you know the percent compositions of the elements in a compound, you can calculate the empirical formula. First, find the number of moles of each element by dividing the percent composition of each element by its atomic mass. Next, divide each result by the lowest number of moles of all the elements contained in the compound. If the results are not integers, multiply by the smallest factor that will make them integers.

Study Tip Because formulas contain whole atoms, the subscripts in an empirical formula should be integers. Sometimes, because of rounding in an earlier calculation, a subscript might be slightly greater or less than an integer.

The following example demonstrates how to find the empirical formula for a compound that is 79.8 percent C and 20.2 percent H. To make the calculations easier, consider a 100 g sample. When you assume a 100 g sample, the percentage values for each element can also be used as the number of g of that element. If 100 percent is equivalent to 100 g, then every percent equals 1 g. 79.8 g C ×

1 mol C = 6.64 mol C 12.011 g C

20.2 g H ×

1 mol H = 20.0 mol H 1.008 g H

Dividing by the smallest number of moles and writing the result as a formula gives C6.64H20.0. Dividing each subscript by the smallest number of moles gives the empirical formula: CH3.

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How can you calculate a molecular formula? If you know the empirical formula and the molar mass of a compound, you can determine the molecular formula. Divide the compound’s molar mass by the empirical formula mass. Then multiply each subscript in the empirical formula by this number. Suppose, for example, a compound has a molar mass of 78.054 g/mole and an empirical formula of CH: molar mass of compound 78.054 g/mol = =6 molar mass of CH 12.001 g/mol + 1.008 g/mol The molecular formula of the compound is C6H6.

TEKS

End-of-Course Assessment Review

1. Calculate  What is the percent composition of hydrogen in beryllium hydride (BeH2) if 69.6 g of beryllium (Be) react with 15.6 g of hydrogen to produce 85.2 g of BeH2? A 18.3 percent B 22.4 percent C 54.0 percent D 81.7 percent 2. Calculate  What is the empirical formula of a compound that is 36.8 percent nitrogen (N) and 63.2 percent oxygen? A NO B NO1.5 C NO2 D N2O3 3. Calculate  What is the molecular formula of a compound that has a molar mass of 180.156 g per mole and an empirical formula of CH2O? A C3H3O3 B C3H6O3 C C6H6O6 D C6H12O6 4. Calculate  What is the formula for a compound with a molar mass of 42.08 g/mol and an empirical formula of CH2?

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TEKS REVIEW

8D

Balancing Chemical Equations

Re adi n ess

TEKS 8D

Use the law of conservation of mass to write and balance chemical equations. TEKS_TXT

Vocabulary law of conservation of mass

What is the law of conservation of mass? The law of conservation of mass states that in any physical change or chemical change, mass is conserved. Mass is neither created nor destroyed. When applied to a chemical reaction, this law means that the atoms present after the reaction (the products) are exactly the same as the atoms present before the reaction (the reactants). Figure 1 illustrates this concept in the making of water (H2O) from the elements hydrogen (H2) and oxygen (O2).

Figure 1 Conservation of Mass

2H2

+

(2 × 2.02g)

+

O2 32.00 g

2H2O =

(2 × 18.02 g)

Four hydrogen atoms and two oxygen atoms are present before and after the reaction. Notice, however, that the atoms have been rearranged. Before the reaction, only particles of hydrogen gas and oxygen gas are present. Within these gases, like atoms are bonded to like atoms. Hydrogen atoms are bonded to other hydrogen atoms to form H2, and oxygen atoms are bonded to other oxygen atoms to form O2. After the reaction, only water molecules are present, in which each oxygen atom is bonded to two hydrogen atoms, which are no longer bonded to each other. Also notice that in the reaction the mass has not changed. The sum of the masses of the two moles of hydrogen gas molecules and of the one mole of oxygen gas molecules equals 36.04 g. This is exactly the mass of the two moles of water molecules that are present after the reaction. Mass is always the same before and after any chemical reaction.

What are balanced equations? A balanced equation tells you what amounts of reactants to mix and what amounts of product to expect. When you know the quantity of one substance in a reaction, you can calculate the quantity of any other substance consumed or created in the reaction. Quantity usually means the amount of a substance expressed in grams or moles. However, quantity could also be in liters, tons, or molecules. TEKS 8D • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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How do you use the law of conservation of mass to write and balance chemical equations? The following steps describe how to write and balance a chemical equation: 1. Write a skeleton equation. Write formulas for reactants on the left and formulas for products on the right. Place an arrow between them. 2. Count the number of atoms of each element in the reactants and of each element in the products. For simplicity, if a polyatomic ion such as NO3–, SO42–, CO32–, or NH4+, is present unchanged on both sides of the equation, treat it as a single unit. 3. Use coefficients to make the number of each element and polyatomic ion the same on both sides of the equation. 4. Make sure the coefficients are in the lowest possible ratio that balances. For example, the following steps describe balancing an equation for the decomposition of iron hydroxide, Fe(OH)3, into iron oxide, Fe2O3, and water, H2O. 1. Write the skeleton equation:

Fe(OH)3



Fe2O3 +

H2O

2. Count the atoms.

Reactants:

1 Fe, 3 O, 3 H



Products:

2 Fe, 4 O, 2 H

3. Write the equation again with a coefficient added to balance Fe.

2Fe(OH)3



Fe2O3 +

H2O

4. Count the atoms again.

Reactants:

2 Fe, 6 O, 6 H



Products:

2 Fe, 4 O, 2 H

5. Write the equation again with coefficients added to balance H and O.

2Fe(OH)3



Fe2O3 +

3H2O

6. Count the atoms again.

Study Tip



Reactants:

2 Fe, 6 O, 6 H

When balancing a chemical equation, start with elements present in only one reactant and one product.



Products:

2 Fe, 6 O, 6 H

Because the quantities of each atom are the same in the reactants and the products, the equation is balanced and obeys the law of conservation of mass.

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TEKS

End-of-Course Assessment Review

1. Define  Because mass is conserved in a chemical reaction, which of the following is the same for the reactants and the products? A amount of energy B number of atoms C moles of molecules D types of substances 2. Apply Concepts  Which of the following chemical equations is balanced? A 4Al + 3O2 ➝ 2Al2O3 B Mg3N2 + H2O ➝ 3MgO + 2NH3 C C3H8 + 5O2 ➝ 3CO2 + H2O D H2SO4 + 2NaCN ➝ HCN + 2Na2SO4 3. Apply Concepts  Which of the following chemical equations is balanced? A 2Al + Fe3N2 ➝ 2AlN + Fe B 4H2O + 7CO2 ➝ C7H8 + 7O2 C C6H12O6 + 3O2 ➝ 6H2O + 6CO2 D 2KNO3 + H2CO3 ➝ K2CO3 + 2HNO3 4. Identify  What coefficient should you place before HCl on the right side of this chemical equation in order to balance it? 2AsCl3 + 3H2S ➝ As2S3 + ■ HCl A 2 B 3 C 4 D 6 5. Construct  Write a balanced chemical equation from this unbalanced equation: Fe + H2O ➝ H2 + Fe3O4 6. Apply Concepts  In your own words, use the law of conservation of mass to explain how to balance a chemical equation.

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TEKS REVIEW

8E

Stoichiometric Calculations TEKS 8E Perform stoichiometric calculations, including determination of mass relationships between reactants and products, calculation of limiting reagents, and percent yield.

How can you determine mass relationships between reactants and products?

Vocabulary stoichiometry limiting reagent excess reagent theoretical yield actual yield percent yield

The calculation of quantities in chemical reactions is a subject of chemistry called stoichiometry. Calculations using balanced equations are called stoichiometric calculations. When preparing chemicals for a reaction, chemists often measure the mass of each reactant. A balanced chemical equation, however, describes the relationships among moles of each substance. To describe mass relationships for the reaction, one must first calculate the molar mass of each reactant and product, and then use the molar masses to convert between moles and grams.

Sample Problem 1 Suppose, for example, 8.75 g of propane (C3H8) react with oxygen gas (O2) to produce carbon dioxide (CO2) and water (H2O). How many grams of water are produced? First, write and balance a chemical equation for the reaction: C3H8 × 5O2  ➝  3CO2 × 4H2O Calculate the molar masses of propane and water. propane: 3(12.011 g/mol) + 8(1.008 g/mol) = 44.097 g/mol

Start with a given quantity and convert from mass to moles. 8.75 g C3H8 ×

1 mol C3H8 44.097 g C3H8

Then convert from moles of reactant to moles of product: 1 mol C3H8 4 mol H2O 8.75 g C3H8 × × 44.097 g C3H8 1 mol C3H8 Finish by converting moles to grams using the molar mass of water. 8.75 g C3H8 ×

water: 2(1.008 g/mol) + 1(15.999 g/mol) = 18.015 g/mol

4 mol H2O 1 mol C3H8

List what you know: mass of propane = 8.75 g

×

1 mol C3H8 44.097 g C3H8

18.015 g H2O 1 mol H2O

×

= 14.3 g H2O

4 mol H2O 1 mol C3H8

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How can you determine the limiting reagent in a reaction? The limiting reagent in a chemical reaction is the reactant that determines the amount of each product that will form. When all of the limiting reagent has been used, the reaction stops. Any other reactants that remain are excess reagents. In Figure 1, the hydrogen gas (H2) is the limiting reagent because it is used up in the reaction. The oxygen gas (O2) is the excess reagent because some of it does not get used in the reaction. Figure 1 Gas as the Limiting Reagent in the Synthesis of Water

Study Tip Remember that the limiting reagent is the reactant that produces the only possible amount of the product given the reactants available.

Before reaction

After reaction

10 H2 and 7 O2

10 H2O and 2 O2

If you know the mass or number of moles of each reactant, you can use the chemical equation to determine the limiting reagent. First, choose one of the reactants. Then use coefficients of the chemical equation and atomic masses to determine how much of each of the other reactants would be needed to completely use up the available mass of the first reactant. The reactant with the least available mass compared to how much is needed is the limiting reagent.

Sample Problem 2

Then convert from moles of reactant to moles of product:

For example what is the limiting reagent when 385 g of sodium (Na) reacts with 125 g of chlorine gas (Cl2)?

385 g Na ×

Write the balanced equation for the reaction:

List what you know and what you need to solve for.

385 g Na ×

mass of sodium = 385 g mass of chlorine gas = 125 g Cl2

1 mol Cl2

1 mol Cl2

2 mol Na

2 mol Na Start with a given quantity and convert from mass to moles. 385 g Na ×

1 mol Na

22.990 g Na

×

1 mol Cl2 2 mol Na

Determine the amount of Cl2 that would react with 385 g of Na.

2Na + Cl2  ➝  2NaCl

1 mol Na

×

1 mol Na 22.990 g Na

70.906 g Cl2 1 mol Cl2

×

= 594 g Cl2

Since only 125 g of chlorine gas (instead of 594 g) is available, it must be the limiting reagent. Sodium is the excess reagent.

22.990 g Na

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How can you calculate percent yield? The amount of a product that you calculate is actually a theoretical yield. Frequently, however, reactions do not run to completion and some reactants remain unused. The amount of a product that actually forms is the actual yield. The ratio of actual yield to theoretical yield, multiplied by 100 percent, is the percent yield. percent yield =

actual yield × 100% theoretical yield

Suppose, for example, that you calculate a theoretical yield of 58.2 g, but the actual yield is 51.9 g. The percent yield is: percent yield =

TEKS

51.9 g × 100% = 89.2% 58.2 g

End-of-Course Assessment Review

1. Calculate  Which of the following reactions had the greatest percent yield? A theoretical yield 52.3 g; actual yield 50.7 g B theoretical yield 17.1 g; actual yield 15.7 g C theoretical yield 38.8 g; actual yield 36.2 g D theoretical yield 24.6 g; actual yield 22.5 g 2. Predict  What is the mass of Br2 that will form if 5.35 g of KBr reacts with excess Cl2? The chemical equation is Cl2 + 2KBr ➝ 2KCl + Br2? A 1.80 g B 2.68 g C 3.59 g D 7.18 g 3. Evaluate  The equation 3Mg × N2 ➝ Mg3N2 describes the synthesis of magnesium nitride (Mg3N2). In a specific example, 2.6 moles of magnesium (Mg) and 4.5 moles of nitrogen gas (N2) are combined to form magnesium nitride (Mg3N2). A student concludes that Mg is the limiting reagent. Do you agree? Defend your answer.

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TEKS REVIEW

9A Gas Laws Re ad i ness

TEKS 9A Describe and calculate the relations between volume, pressure, number of moles, and temperature for an ideal gas as described by Boyle’s law, Charles’ law, Avogadro’s law, Dalton’s law of partial pressure, and the ideal gas law.

What is the relationship between volume, temperature, and pressure?

Vocabulary standard temperature and pressure (STP) Boyle’s law Charles’ law Avogadro’s law Dalton’s law of partial pressure ideal gas law

Four things are generally used to describe a gas—pressure (P), volume (V), temperature (T), and number of moles (n). The amount of gas, the volume, and the temperature are the factors that affect gas pressure. The volume of a gas varies with a change in temperature or a change in pressure. Due to these variations, the volume of a gas is usually measured at a standard temperature and pressure. Standard temperature and pressure (STP) is equal to a temperature of 0°C and a pressure of 101.3 kPa or 1 atmosphere (atm). At STP, one mole (mol) of a gas occupies 22.4 L.

What are the gas laws? There are a number of gas laws that describe the relationship between volume, temperature, and pressure. Boyle’s Law  Robert Boyle was the first person to propose a law that described the relationship between pressure and volume of a gas. Boyle’s law states that for a given mass of a gas at a constant temperature, the volume of the gas varies inversely with the pressure. The mathematical expression of Boyle’s law is: P1 × V1 = P2 × V2

Sample Problem 1 The pressure on a balloon with a volume of 300 mL increases from 1.10 to 2.00 atmospheres (atm). Use Boyle’s law to calculate the new volume of the balloon.

Rearrange the equation algebraically to isolate the unknown. V2 =

First, identify the variables. P1 = 1.10 atm   V1 = 300 mL   P2 = 2.00 atm   V2 = ?

P1 × V1 P2

Insert the known values and solve. 1.10 atm × 300mL V2 = 2.00 atm V2 = 165 mL

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Charles’ Law  The volume and absolute temperature of a gas are also closely related. This relationship, Charles’ law, states that if you increase or decrease the temperature, T, of a gas, but keep the pressure constant, its volume (V) will increase or decrease by the same proportion. V1 V = 2 T1 T2

Sample Problem 2 The air in a balloon with a volume of 25.0 liters (L) is heated from 20°C to 60°C. If the pressure stays the same, what will be the new volume of the balloon?

Rearrange the equation algebraically to isolate the unknown. V1 × T2 V2 = T1

First, identify the variables. V1 = 25.0 L   T1 = 2  0°C (convert °C to K by adding 273) = 293 K

Insert the known values and solve.

T2 = 60°C = 333 K   V2 = ?  

V2 =

25.0 L × 333 K 293 K V2 = 28.4 L

Gay-Lussac’s Law  A third law, called Gay-Lussac’s law, relates the pressure and temperature of a gas kept at constant volume. P1 P = 2 T1 T2

Study Tip When calculating using the gas laws, use dimensional analysis to cancel units out, leaving only the appropriate unit in the answer. This will help you to catch errors and to convert units.

Combined Gas Law  Often, the pressure, volume, and temperature of a gas change at the same time. To explain these cases, the separate gas laws may be combined into one equation, called the combined gas law. The combined gas law summarizes the relationships among pressure, volume and temperature. (P1 × V1) (P2 × V2) = T1 T2 Avogadro’s Law  Another gas law, Avogadro’s law, states that equal volumes of gases at the same temperature and pressure contain equal numbers of particles. Dalton’s Law of Partial Pressure  Each gas in a gas mixture exerts a partial pressure. In a mixture of gases, the total pressure is the sum of the partial pressures of the gases. This is Dalton’s law of partial pressure. Dalton’s law of partial pressures is shown by the equation: Ptotal = P1 + P2 + P3 + …

Sample Problem 3 A sample of gas in a 1.00 L flask at 1.50 atm contains 75.0 percent CO2 and 25.0 percent H2O gas. Calculate the partial pressures of each gas.

PCO = (0.750 × 1.50 atm) = 1.125 atm 2

PH

2

O

= (0.250 × 1.50 atm) = 0.375 atm

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Ideal Gas Law  The gas law that contains all four variables—P, V, T, and n—is called the ideal gas law. The volume, pressure, number of moles, and temperature of an ideal gas are proportionately related. In the ideal gas law equation, a conversion factor called the ideal gas constant, R, is used (R = 0.08206 L atm/mol K). P × V = n × R × T or PV = nRT

Sample Problem 4 Calculate the temperature of 5.85 mol N2 gas in a 12.0 L steel bottle under 10.0 atm of pressure.

Rearrange the equation algebraically to isolate the unknown.

First, identify the variables. T=

P = 10.0 atm     V = 12.0 L     n = 5.85 mol R = 0.08206 L atm/mol K

T =?

P×V n×R

Insert known values and solve. T=

10.0 atm × 12.0 L 5.85 mol × 0.08206 L atm/mol K T = 250 K = –23.0°C

TEKS

End-of-Course Assessment Review

1. Apply Concepts  Which of the following statements about ideal gases is supported by Charles’s law? A An increase in pressure will cause a proportional decrease in volume. B If you decrease the temperature of a gas, its volume will decrease proportionally. C At STP, 1 mole occupies 22.4 L. D The total pressure of a gas mixture is the sum of the partial pressures of each gas. 2. Calculate  How many moles of propane (C3H8) are in a 7.00-L tank at 20.0°C and 5.45 atm of pressure? A 0.629 moles C 23.2 moles B 1.59 moles

D 917 moles

3. Calculate  The volume of air in a diving bell changes as it descends into deep water. If the air in a diving bell occupies 9.50 L at 1.02 atm and 20.1°C, what will be the volume of the air at 1.55 atm and 15.0°C?

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9B

Gas Stoichiometry TEKS 9B Perform stoichiometric calculations, including determination of mass and volume relationships between reactants and products for reactions involving gases.

How can you determine the mass relationships between reactants and products for reactions involving gases? You can obtain mole ratios from a balanced chemical equation. By using the mole ratios, you can calculate any measurement unit that is related to the mole. This is true for chemical reactions involving gases. For example, suppose you want to determine the mass relationships between reactants and products for a reaction involving gases. You first convert the given mass to moles. Then you use the mole ratio from the balanced equation to calculate the number of moles of the wanted substance. Finally, you convert the moles to mass. The relationship between moles and mass (1 mol = molar mass) is useful in solving mass-mass stoichiometric problems. The mole-mass relationship gives you two conversion factors. 1 mol molar mass and molar mass 1 mol

Sample Problem 1

1 mol H2O = 18.0 g H2O (molar mass) 1 mol O2 = 32.0 g O2 (molar mass)

What is the mass of oxygen gas produced when 29.2 g of water is decomposed by electrolysis according to this balanced equation?

mass of oxygen = ? g O2

2H2O → 2H2 + O2 In order to solve this equation, you have to perform the following calculations: g H2O → mol H2O → mol O2 → g O2

Then convert from moles of reactant to moles of product. 1 mol H2O 1 mol O2 29.2 g H2O × × 18.0 g H2O 2 mol H2O

The appropriate mole ratio relating mol O2 to mol H2O from the balanced equation is 1 mol O2 . 2 mol H2O

Finish by converting from moles to mass

First, list the knowns and the unknown.

29.2 g H2O ×

mass of water = 29.2 g H2O 1 mol O2 2 mol H2O

Next, calculate for the unknown. Start with the given quantity, and convert from mass to moles. 1 mol H2O 29.2 g H2O × 18.0 g H O 2

1 mol H2O × 18.0 g H2O

32.0 g O2 1 mol O2

(from balanced equation)

1 mol O2 2 mol H2O

×

= 26.0 g O2

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Study Tip If these stoichiometry problems seem too challenging, review STAAR Review 8E on stoichiometric calculations. Practice dimensional analysis to make sure your calculations have the appropriate conversion factors.

How can you determine volume relationships between reactants and products for reactions involving gases? Recall that the mole can be related to the volume of a gas at STP. The relationship between moles and volume at STP (1 mol of a gas = 22.4 L at STP) is useful in solving volume-volume stoichiometric problems. The mole-volume relationship gives you two conversion factors. 1 mol 22.4 L

Sample Problem 2 Nitrogen monoxide and oxygen gas combine to form the brown gas nitrogen dioxide, which contributes to photochemical smog. What volume (in L) of nitrogen dioxide gas is produced when 34 L of oxygen gas react with an excess of nitrogen monoxide? Assume conditions are at STP. 2NO + O2→ 2NO2 Recall that the following calculations need to be performed: L O2 →mol O2 → mol NO2→ L NO2 First, list the knowns and unknowns: volume of oxygen = 34 L O2 2 mol NO2 1 mol O2

and

22.4 L 1 mol

Then, solve for the unknown. Start with the given quantity, and convert from volume to moles by using the mole-volume ratio. 1 mol O2 34 L O2 × 22.4 L O2 Then, convert from moles of reactant to moles of product by using the correct mole ratio. 2 mol NO2 1 mol O2 34 L O2 × × 1 mol O2 22.4 L O2 Finish by converting from moles to liters. Use the mole-volume ratio. 1 mol O2 2 mol NO2 34 L O2 × × × 22.4 L O2 1 mol O2 22.4 L NO2 1 mol O2 = 68 L NO2

(from balanced equation)

1 mol O2 = 22.4 L O2 (at STP) 1 mol NO2 = 22.4 L NO2 (at STP) volume of nitrogen dioxide = ? L NO2

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TEKS

End-of-Course Assessment Review

1. Calculate  How much oxygen will react with 6.0 liters hydrogen in the reaction below? 2H2 + O2→ 2H2O A 3.0 liters B 6.0 liters C 12 liters D 18 liters 2. Analyze  Use the reaction equation below to determine the volume ratio of acetylene (C2H2) to carbon dioxide (CO2). 2C2H2 + 5O2 → 2H2O + 4CO2 A 1 to 1 B 1 to 2 C 2 to 5 D 5 to 4 3. Apply Concepts  Octane (C8H18) in gasoline reacts with oxygen as shown in the equation below. 2C8H18 + 25O2→ 16CO2 + 18H2O   If one gallon of gasoline contains 995 grams octane, what is the volume of carbon dioxide produced in the combustion of a gallon of gasoline at STP? A 69.7 liters B 179 liters C 1560 liters D 178,000 liters 4. Calculate  If 14 g carbon monoxide (CO) gas react with oxygen (O2) gas, producing carbon dioxide (CO2) in a pressurized flask at 2.0 atmospheres and 290 kelvins, how many liters of carbon dioxide will be produced?

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9C

Kinetic Molecular Theory TEKS 9C

Describe the postulates of kinetic molecular theory.

Vocabulary kinetic molecular theory

What are the postulates of kinetic molecular theory? You may recall that the word kinetic refers to motion. The energy an object has because of its motion is kinetic energy. The kinetic molecular theory states that all matter consists of small particles that are always in motion. Kinetic molecular theory—as it applies to gases—is based on three postulates, also called assumptions. Postulate 1: The particles in a gas are considered to be small, hard spheres with insignificant volumes. Gases are made of particles so small that, compared to the distances between each other, their own sizes are negligible. In other words, the volume of a gas can be imagined to be composed of the empty space between the particles. The particles are assumed to be spheres that bounce off each other in every direction. Many real gases do not perfectly fit this postulate. For example, when gases are compressed or are at low temperature, the volumes of the particles become a significant factor in the volume of the gas. Also, many common gases, such as nitrogen (N2), oxygen (O2), and carbon dioxide (CO2) contain molecules that are not spherical. Even so, the predictions of kinetic molecular theory come very close to describing most gases under most conditions. Postulate 2: The motion of the particles in a gas is rapid, constant, and random. The particles of a gas move randomly, in straight lines, at different speeds, and in every direction. The particles change direction only when they collide with other particles or the walls of their container, as represented in Figure 1 below. In other words, the particles of a gas do not attract each other and change these straight-line paths.

Figure 1 Gases in Their Containers

Container wall

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But real gases have some attractions for each other. Some gases, such as water vapor and water, are even polar molecules. Other gas molecules, such as nitrogen (N2) and oxygen (O2), have only very weak and temporary attractions for each other. In either case, the attractions or repulsions caused by the presence of a nearby molecule can change the path of another molecule. Despite these imperfections, the kinetic molecular theory’s predictions remain very close to the values observed in experiments. Postulate 3: The collisions between particles in a gas are perfectly elastic. When the particles of a gas collide, the total kinetic energy, which is the energy of motion, stays the same. In other words, when the particles collide, they don’t lose energy by reshaping each other in a collision. The kinetic energy in a sample of gas will remain constant unless thermal energy is either added or removed. Additionally, the average kinetic energy of the particles in a gas depends on the absolute temperature of the gas. This means that you can think of the temperature of a gas as a measure of its average kinetic energy.

Study Tip Remember that temperature is a measure of the average of the kinetic energies of a gas’s molecules. Some molecules move slower than most, and others move faster than most.

Figure 2 Gas Pressure

For the most part, this postulate is true of real gases. But a collision of gas molecules can cause changes in them. In fact, if a mixture of gases is involved, the collision could even spark a chemical reaction between the molecules. Usually, the kinetic theory is applied primarily to gas molecules that are not in mixtures or not expected to react suddenly.

How do the postulates of kinetic molecular theory explain the behavior of gases? One consequence of kinetic molecular theory is that, over time, given the randomness of collisions, any sample of a gas will have a range of speeds (and thus kinetic energies) for its particles. The temperature of a gas is actually a measure of the average kinetic energy of that gas’s particles. As the temperature increases, the average value of the kinetic energy increases, but there are still some very slow molecules, and there are still some very fast ones. Pressure is also explained by kinetic molecular theory. Pressure is defined as force per unit area, and any gas presses outward on its surroundings with some pressure. This force can be imagined as the sum of all of the tiny forces caused by the individual molecules of the gas colliding with the sides of its container. The postulates of kinetic molecular theory can even explain each of the gas laws. For example, the postulates explain Boyle’s law, which relates pressure and volume. When the volume of a gas decreases, assuming that the temperature of the gas does not change, the particles that compose it will be forced into a smaller space, so they will collide with each other and the walls of their container more frequently, resulting in an increase in pressure, as you can see in Figure 2.

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TEKS

End-of-Course Assessment Review

1. Differentiate  Which of the following is the postulate of kinetic molecular theory that explains why a liter of gas is essentially a liter of empty space? A The particles in a gas are considered to be small, hard spheres with insignificant volumes. B The motion of the particles in a gas is rapid, constant, and random. C Equal volumes of gases at the same temperature and pressure contain equal numbers of particles. D The collisions between particles in a gas are perfectly elastic. 2. Apply Concepts  Which of the following is explained by the postulates of kinetic molecular theory? A Smaller particles have more surface area than larger particles. B Water boils at 100°C. C Propane releases energy stored in its bonds when it reacts with oxygen. D An aerosol can is cold to the touch while gas is sprayed out. 3. Describe  Summarize, in your own words, the postulates of kinetic molecular theory. 4. Apply Concepts  Use the postulates of kinetic molecular theory to explain why air pressure in car tires tends to be lower on cold days than on hot days. 5. Evaluate  A student predicts that a mole of uranium hexafluoride gas (UF6) will have a much greater volume than a mole of hydrogen gas (H2), because the molecules of UF6 are so much larger and more massive than those of H2. Explain what is wrong with this reasoning.

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10A The Importance of Water TEKS 10A

Describe the unique role of water in chemical and biological systems.

What factors contribute to water’s unique properties?

Vocabulary cohesion surface tension adhesion

The two factors that contribute to water’s unique properties are (1) that water is a polar molecule, and (2) hydrogen bonding. Both factors are a result of the distribution of electrons and the shape of the molecule. An individual water molecule consists of an oxygen atom covalently bonded to two hydrogen atoms. You may recall that oxygen has high electronegativity. Because oxygen has greater electronegativity than hydrogen, the electrons in a water molecule are attracted more to the oxygen than to the hydrogen molecules. Thus, the oxygen end of the molecule has a partial negative charge and the hydrogen end has a partial positive charge, as shown in Figure 1a. The polarity of the two O—H bonds do not cancel out. Water molecules are polar. Polar molecules are attracted to one another— the negative end of one polar molecule is attracted to the positive end of another polar molecule. Hydrogen atoms in one water molecule form attractions to the oxygen atom of other water molecules. The strong intermolecular attractions between water molecules lead to the formation of hydrogen bonds. Each water molecule forms up to two hydrogen bonds with other molecules. These hydrogen bonds contribute to many of the unique properties of water such as high surface tension, high specific heat, adhesion, and cohesion. (a) Water Molecule

Figure 1 Polarity and Hydrogen Bonding

(b) Atomic Pattern of Liquid Water

δ+

Hydrogen bond

2δ–

δ+

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What role does water play in chemical systems? Water as a Solvent  Because water is both strongly polar and forms hydrogen bonds, polar substances and ionic substances dissolve in water. An aqueous solution is water that contains dissolved substances. High Specific Heat  Because it is nonreactive, water is a useful medium for many chemical reactions. Water has high specific heat, which allows it to resist changes in temperature. As a result, it has an unusually low freezing point and high boiling point. As water is heated, much of the energy is absorbed by the bonds within and between each molecule. A gram of water absorbs 2270 joules before it evaporates. Because water absorbs energy with only minimal changes in temperature or state, it is useful for cooling other systems. Neutrality  Water is also unusual in its neutrality. Water dissociates into hydrogen ions (H+) and hydroxide ions (OH–). Water is a neutral substance because it produces hydrogen ions and hydroxide ions at the same rate. Density of Solid Water  The solid phase for most substances is denser than the liquid phase. That is not the case with water. When water freezes, the hydrogen atoms and oxygen atoms align themselves into a crystal pattern that has more space between each molecule than in liquid water (compare the two structures in Figure 1b and Figure 2). As a result, ice is less dense than liquid water. Figure 2 Atomic Patterm of Ice Hydrogen bond

What role does water play in biological systems? Study Tip Use the mnemonic “Oh no!” to remember that O (oxygen) is more “negative” in most molecules.

Water is an important part of every living thing. It is an excellent solvent that can transport nutrients and wastes. Hydrogen bonding between water molecules gives water a property called cohesion—the attraction between molecules of the same substance. Drops of water form on substances, such as table tops, because surface water molecules are drawn inward due to cohesion. Surface water molecules are hydrogen bonded only on the inside of the drop. This inward force is called surface tension.

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When water comes in contact with certain substances that attract water, called hydrophilic substances, water will rise through adhesion—an attraction between molecules of a different substance. Adhesion leads to water traveling from a plant’s roots to its stems and leaves. Water also helps regulate the temperature of living organisms because it has a high heat capacity. It resists rapid change in temperature and absorbs a great deal of heat when it evaporates. Bodies of water also resist changes in temperature, which affects climate. Finally, because ice is less dense than liquid water, ice floats on bodies of water. The ice insulates the liquid water below it from the cold air. This usually prevents bodies of water from freezing completely, which allows fish and other aquatic life to survive winters.

TEKS

End-of-Course Assessment Review

1. Predict  Ammonia (NH3) and water have similar structures, but the forces between the molecules of ammonia are weaker. How would the properties of ammonia differ from those of water? A Ammonia would have a lower heat capacity. B Ammonia would have greater molecular weight. C Ammonia would have a higher boiling point. D Ammonia would be more cohesive. 2. Describe  Which of the following describes cohesion? A the ability of a liquid to dissolve substances B the force needed to break through the surface of a liquid C the density of a liquid in a long tube D the rise of a liquid in a narrow tube 3. Analyze  Which aspect of water makes it an excellent solvent? A It has a high heat capacity. B It has a low freezing point. C Its crystal form is less dense than its liquid form. D It is polar and forms hydrogen bonds. 4. Describe  List four roles of water in chemical and biological systems.

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10B Solubility Rules Rea di ness

TEKS 10B

Vocabulary solubility spectator ion

Develop and use general rules regarding solubility through investigations with aqueous solutions.

What is solubility? The solubility of a substance is a measure of the amount of solute that can dissolve in a given quantity of a solvent at a specified temperature and pressure. Solubility of solids is often expressed in grams of solute per 100 g of solvent (g/100 g H2O). For example, sodium chloride (NaCl) is soluble in water—almost 40 g of sodium chloride can dissolve in 100 g of water. A substance that does not dissolve well in a solvent at a given temperature and pressure is described as insoluble. If a substance is insoluble in water, only a small amount of the substance will dissolve. The rest of the insoluble substance precipitates from the solution. When two solutions containing ionic compounds are combined, the ionic compounds may or may not react. If they do not react, the ions all remain in solution. If they do react, a product may precipitate from the solution. A precipitate forms if a new compound that forms is insoluble.

How can you develop solubility rules through investigations? Most ionic substances and polar substances are soluble in water, but there are exceptions. Knowing the differences in solubility of various substances can be helpful for predicting and identifying the products of reactions. Some trends in solubility can be easy to identify. For example, compounds containing sodium, potassium, and cesium dissolve readily in water. On the basis of this observation, you can conclude that compounds formed from Group 1 elements are typically soluble. You can also identify a solubility rule from one reaction, and use it to develop rules about the products of another reaction. For example, consider the reaction of silver nitrate (AgNO3) and potassium iodide (KI). One of the products, silver iodide (AgI), is an insoluble precipitate. AgNO3(aq)   +   KI(aq)   ➝   AgI(s)   +   KNO3(aq)

silver nitrate

potassium iodide

silver iodide

potassium nitrate

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Study Tip If all of the possible products of a reaction in a solution are soluble, it means that there will be no reaction between the products.

In another reaction, dilute solutions of silver chloride (AgCl) and lithium iodide (LiI) are combined, and a precipitate forms. On the basis of the previous reaction, you know that silver iodide is insoluble. This precipitate can also be identified as silver iodide. AgCl(aq)   +   LiI(aq)   ➝   AgI(s)   +   LiCl(aq)

silver chloride

lithium iodide

silver iodide

lithium chloride

You may have noticed that some symbols have appeared in parentheses after the formulas. The symbol (s) is used to show that a substance is a solid, (l) indicates that it is a liquid, and (g) shows that the substance is a gas. In addition, (aq) is used to indicate that the material is dissolved in water; that is, it is an aqueous solution. In the above formula, the silver iodide (AgI) forms is a solid precipitate. The lithium chloride (LiCl) is in an aqueous solution.

How can solubility rules be used? General rules of solubility are sometimes listed according to which ionic compounds dissolve the most readily and which ones are the least soluble. A table summarizing solubility rules, such as that in Figure 1, can be helpful for predicting the products of a reaction. For example, the rules can be used to determine the products of a reaction involving aluminum chloride and potassium phosphate. AlCl3(aq)   +   K3PO4(aq)   ➝   3KCl(aq)   +   Al PO4(s)

aluminum chloride

potassium phosphate

potassium chloride

aluminum phosphate

In this reaction, you can use solubility rules to determine that potassium chloride would be soluble, because potassium is an alkali metal, and solubility rules state that salts of alkali metals are soluble. Then, you can verify that aluminum phosphate is a precipitate because phosphates tend to be insoluble in water. Aluminum is not listed as one of the exceptions to this rule. Figure 1

Solubility Rules for Ionic Compounds in Aqueous Solutions Compounds

Solubility

Exceptions

Salts of alkali metals and ammonia

Soluble

Some lithium compounds

Nitrate salts and chlorate salts

Soluble

Few exceptions

Sulfate salts

Soluble

Compounds of Pb, Ag, Hg, Ba, Sr, and Ca

Chloride salts

Soluble

Compounds of Ag and some compounds of Hg and Pb

Carbonates, phosphates, chromates, sulfides, and hydroxides

Most are insoluble

Compounds of the alkali metals and of ammonia

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How do you write a balanced net ionic equation? A net ionic equation shows only those particles involved in the reaction and is balanced with respect to both mass and charge. Consider the equation for the reaction of lead with silver nitrate. Pb(s) + AgNO3(aq) ➝ Ag(s) + Pb(NO3)2(aq) The nitrate ion is a spectator ion in this reaction. Spectator ions appear on both sides of an equation but are not directly involved in the reaction. The net ionic equation is as follows: Pb(s) + Ag+(aq) ➝ Ag(s) + Pb2+(aq) Why is this equation unbalanced? Notice that a single unit of positive charge is on the reactant side of the equation. Two units of positive charge are on the product side. Placing the coefficient 2 in front of Ag+(aq) balances the charge. A coefficient of 2 in front of Ag(s) also rebalances the atoms. Pb(s) + 2Ag+(aq) ➝ 2Ag(s) + Pb2+(aq)

TEKS

End-of-Course Assessment Review

1. Define  What is a precipitate? A an insoluble product of a reaction between substances in a solution B a soluble product of a reaction between substances in a solution C any product in a reaction between substances in a solution D a covalent product from a reaction between ionic compounds 2. Predict  Based on the information in the review and Figure 1, which of the following compounds would form a precipitate in water? A KCl C (NH4)2S B RbOC

D

BaSO4

3. Predict  Based on the information in the review and Figure 1, in a reaction between calcium hydroxide (Ca(OH)2) and lithium sulfide (Li2S), which of the following, if any, would be the solid product(s)? A calcium sulfide (CaS) B lithium hydroxide (LiOH) C calcium sulfide (CaS) and lithium hydroxide (LiOH) D no solid products 4. Apply  Use the general solubility rules to write the balanced net ionic equation for a reaction between nickel chloride (NiCl2) and sodium carbonate (Na2CO3).

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10C Calculations Involving Molarity TEKS 10C

Calculate the concentration of solutions in units of molarity.

Vocabulary solvent solute molarity (M)

How can you describe the concentration of a solution? Solutions are homogeneous mixtures that may be solid, liquid or gaseous. In a solution, the dissolving medium is the solvent. The dissolved particles in a solution are the solute. The concentration of a solution is a measure of the amount of solute dissolved in a given quantity of solvent. A solution that contains a relatively small amount of solute is a dilute solution. By contrast, a solution that contains a large amount of solute is a concentrated solution. However, dilute and concentrated are relative terms and not precise regarding the amount of solute involved. One of the most useful ways to describe the concentration of a solution is its molarity, also called its molar concentration. The molarity (M  ) of a solution is the number of moles of solute per liter of solution. For example, a 1M, or 1-molar, solution of copper sulfate (CuSO4) has one mole (mol) of copper sulfate for each liter (L) of the solution.

How can you calculate the molarity of a solution? You can calculate molarity by dividing the amount of solute in moles by the volume of the solution in liters. moles of solute Molarity (M) = liters of solution For example, suppose 0.25 moles of sodium sulfate (Na2SO4) is dissolved in 1.5 L solution. You could calculate the molarity of the solution as follows: 0.25 mol Na2SO4 = 0.17 mol Na2SO4/L = 0.17 M Na2SO4 1.5 L solution You can also use molarity to find the amount of solute in a solution by multiplying molarity by liters of the solution. Moles of solute = molarity × liters of solution For example, if you have 825 mL of 0.25-molar hydrochloric acid (HCl), you can determine the number of moles of HCl in the solution by multiplying as follows: mol HCl = (0.25 mol HCl/L ) × (0.825 L HCl) = 0.21 mol HCl TEKS 10C • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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Molarity can also be used to calculate the volume of a solution.

Study Tip When calculating using molarity, consider it a conversion factor between moles and liters. This will help you to apply molarity using dimensional analysis.

Liters of solution =

moles of solute Molarity (M)

For example, consider a 2.0-molar solution of glucose (C6H12O6). How many liters of the solution would contain 3 moles glucose? liters of solution =

moles C6H12O6 Molarity

3 mol C6H12O6

=

(2 mol C6H12O6    / L)

= 1.5 L

How can you calculate molarity using mass? In many problems, the mass of solute is given instead of the number of moles. To solve this type of problem, convert grams of solute to moles of solute. For example, how could you calculate the molarity of 0.510 L of solution that contains 110 g sodium chloride (NaCl)? To calculate molarity, first divide by the molar mass to find the amount of NaCl in moles. The molar mass of NaCl is 58.44 grams. mol NaCl = (110 g NaCl ×

1 mol NaCl (58.44 g NaCl)

) = 1.9 mol NaCl

Then, calculate the molarity of NaCl. molarity NaCl =

1.9 mol NaCl 0.510 L

= 3.7M NaCl

The molarity of a solution can also be used to calculate the mass of solute, such as the mass of Mg(OH)2 in 275 mL 0.33M Mg(OH)2. The molar mass of Mg(OH)2 is 58.31 grams. First, use the molarity to find the moles of Mg(OH)2 in 275 mL 0.33M Mg(OH)2. mol Mg(OH)2 = 0.275L ×

0.33 mol Mg(OH)2

= 0.091 mol Mg(OH)2 L Then, calculate the mass of that amount of Mg(OH)2: mass Mg(OH)2 = mol Mg(OH)2 × molar mass Mg(OH)2 mass of Mg(OH)2 = 0.091 mol Mg(OH)2 ×

58.31 g (Mg(OH)2 1 mol Mg(OH)2

= 5.3 g Mg(OH)2

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TEKS

End-of-Course Assessment Review

1. Define  Molarity is calculated by dividing the moles of a solute by which quantity? A moles of the solvent B volume of the solvent C moles of the solution D volume of the solution 2. Calculate  A chemist dissolves 49 g manganese (IV) oxide (MnO2) in water, and adds enough water to make 510 mL of solution. Calculate the molarity of the solution. The molar mass of MnO2 is 86.9 g. A 0.096M B 0.91M C 1.1M D 10.4M 3. Calculate  A 0.10M sports drink is prepared using sucrose (C12H22O11) in 390 mL of water. What is the mass of sucrose added to the water? The molar mass of sucrose is 342 g. A 0.11 g B 0.88 g C 2.5 g D 13 g 4. Compare and Contrast  Which is the more concentrated solution: a solution made of 0.70 moles of Na2SO4 dissolved to make 0.35 L of solution, or 75 g of Na2SO4 dissolved to make 300 mL of solution? Explain your conclusion.

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10D Calculating Dilutions TEKS 10D

Use molarity to calculate the dilutions of solutions.

How is molarity used to calculate the dilutions of solutions? Molarity is important for calculating changes in the concentration as a solution is diluted. Calculations of dilutions are based on the fact that the amount of the solute stays the same in the solution, even after more solvent is added. Diluting a solution reduces the number of moles of solute per unit volume, but the total number of moles in solution does not change. This concept can be expressed with the following equation: Moles of solute before dilution = moles of solute after dilution Remember the definition of molarity and how it can be rearranged to solve for moles of solute: Molarity (M) =

moles of solute liters of solution (V)

For instance, you might start with a small amount of a concentrated solution and dilute it with water. The final solution will have a larger volume but a lower concentration, and the moles of solute present will be the same after as before the dilution. This relationship can be expressed by the following formula, which states that the product of the molarity and volume of the first solution (M1 × V1), is equal to the product of the molarity and volume of the second solution (M2 × V2). Moles of solute = M1 × V1 = M2 × V2 Suppose 1.5 moles (mol) of sulfuric acid (H2SO4) are dissolved in 0.10 L of solution. You would calculate the molarity of the solution as follows: 1.5 mol H2SO4 0.10 L solution

= 15 mol H2SO4/L = 15M H2SO4

No matter how much the solution is diluted, the amount of H2SO4 in the solution will still be 1.5 moles. As the sulfuric acid solution is diluted, its volume increases and its concentration decreases proportionately.

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Sample Problem 1 Suppose the solution in the example above were diluted to a volume of 1.0 L. What would the resulting concentration be? The relationship M1 × V1 = M2 × V2 can be used to find the new concentration. First, identify the variables. Also apply the definition of molarity (M), which is moles of solute per liter of solution. In this solution, the solute is H2SO4.

?M 1.0 L

M1 = 15M = 15 mol H2SO4 /L V1 = 0.10 L

15M 100 mL

1.5 mol H2SO4

M2 = ? V2 = 1.0 L

1.5 mol H2SO4

Rearrange the equation to solve for the unknown variable, M2. M1 × V1 = M2 × V2 M2 =

M1 × V1 V2

Substitute the variables into the equation. Then solve for M2. M2 =

15 mol H2SO4/L × 0.10 L 1.0 L

= 1.5 mol H2SO4/L = 1.5M H2SO4

The new concentration is 1.5M.

Sample Problem 2 A chemist has a container of concentrated 3.0M sodium hydroxide (NaOH). If she wants to prepare 450 mL of 0.10M sodium hydroxide, how much of the concentrated solution will she need to use? The relationship M1 × V1 = M2 × V2 can be used to find the new volume. First, identify the variables. Then, express molarity (M) as moles of NaOH per liter of solution. M1 = 3.0M NaOH = 3.0 mol NaOH/L V1 = ? M2 = 0.10M NaOH = 0.10 mol NaOH/L V2 = 450 mL = 0.450 L

Study Tip When calculating dilutions, remember that the volume of an aqueous solution as a whole tends to be just slightly more than that of the solvent alone.

Rearrange the equation to solve for the unknown variable, V1. M × V2 V1 = 2 M1 Finally, substitute the variables into the equation, and solve for V1. V1 =

(0.10 mol NaOH/L × 0.450 L) 3.0 mol NaOH/L

= 0.015 L

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TEKS

End-of-Course Assessment Review

1. Predict  If a solution is diluted by doubling its volume with water, what will happen to its concentration? A It will increase by a factor of 10. B It will also double. C It will decrease by a half. D It will decrease by a factor of 10. 2. Calculate  A bottle of industrial-grade hydrochloric acid (HCl) is 8.0M. The molar concentration of hydrochloric acid for household cleaning is 2.7M. How much industrial-grade hydrochloric acid is required to prepare 10.0 L for household cleaning purposes? A 0.30 L B 2.2 L C 3.0 L D 3.4 L 3. Calculate  A chemistry student dilutes 550 mL of 3.0M barium hydroxide (BaOH) to prepare a 0.033M solution. What is the volume of the new diluted solution? A 0.17 L B 0.60 L C 5.0 L D 17 L 4. Calculate  If you were to add 5.0 g table salt (NaCl) to a beaker of water, what would be the volume of solution required to make a 0.05M solution? 5. Explain  A bottle of solution contains 500 mL of 0.20M potassium iodide (KI). By diluting this solution with water, what is the largest volume of 0.15M KI you could prepare? Explain why you could not prepare a larger volume.

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TEKS REVIEW

10E Types of Solutions Rea di ness

TEKS 10E Distinguish between types of solutions such as electrolytes and nonelectrolytes and unsaturated, saturated, and supersaturated solutions.

How can you distinguish between electrolytes and nonelectrolytes?

Vocabulary electrolyte nonelectrolyte saturated solution unsaturated solution supersaturated solution

Forces between particles determine how a substance behaves in a liquid. Ionic compounds, for example, have relatively weak forces holding the ions together. When ionic compounds are dissolved in an aqueous solution or in a liquid phase, the ions move freely. One way to distinguish solutions is based on the strengths of the forces between their particles. Solutions with weak forces between their particles are called electrolytes. Electrolytes are compounds that conduct an electric current when they are in aqueous or molten states. As shown in Figure 1, a solution can be joined into an electric circuit that is powered by a battery. The particles of the electrolyte (such as dissolved ions) align in the solution to allow electrons to flow.

Figure 1 An Electrolyte Solution Circuit

Flow of electrons

Power source

Current meter Flow of electrons

Inert metal electrodes +



Cl– Na+

Molten salt (801°C—1412°C)

To (+) To (–) electrode electrode

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Some compounds are stronger electrolytes than others. A weak electrolyte conducts only a small amount of electricity in a solution. The strongest electrolytes are ionic compounds that completely dissociate into ions in water. Solutions with strong forces between their particles are called nonelectrolytes. Covalent substances tend to be held together with more force than ionic substances. When covalent substances dissolve in water or are in a liquid phase, the molecules will not align themselves to permit the flow of electrons. Nonelectrolytes are compounds that do not conduct an electrical current in aqueous solutions or in the molten state.

How can you distinguish between unsaturated, saturated, and supersaturated solutions? You can also distinguish solutions based on how saturated they are, in other words, by the amount of solute present in the solvent at a given temperature and pressure. Every solution can be characterized by the relation between the amount of solute and its solubility. Solubility and Saturated Solutions  Solubility is the amount of a solute that can dissolve in a given quantity of a solvent at a specified temperature and pressure to produce a saturated solution. A saturated solution is a solution containing the maximum amount of solute for a given amount of solvent at a constant temperature and pressure.

Study Tip Although supersaturated and saturated solutions exist, they are relatively uncommon. Nearly all of the solutions you are likely to work with are unsaturated.

Unsaturated Solutions  For solid and liquid solutes, higher temperatures allow more solute to dissolve. For gas solutes, higher pressure allows more solute to dissolve. Any solution in which more solute can be dissolved under existing conditions is an unsaturated solution. Each substance has a unique tendency to dissolve in a solvent at a specific temperature. For example, 203.9 g sucrose can be dissolved in 100 mL of water at 20°C. Any more than 203.9 g of sucrose will remain undissolved and settle to the bottom of the solution. Supersaturated Solutions  On the other hand, if you gradually lower the temperature of a saturated solution, crystals may not begin to form. In this case, more of the solute remains dissolved than would otherwise be possible at that temperature or pressure. This type of solution is called a supersaturated solution. A supersaturated solution is unstable. Any disturbance or addition of solute can cause the excess solute to suddenly precipitate.

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TEKS

End-of-Course Assessment Review

1. Distinguish  Which of the following describes a solution containing an electrolyte? A The solute particles are so firmly bonded that they do not break apart. B The solute particles permit the passage of an electric current. C It is unstable, and all the solute will precipitate if the solution is disturbed. D It cannot contain any more solute particles. 2. Apply Concepts  A chemistry student prepares a saturated solution of KNO3 in 100 g water at 50°C. Then she rapidly cools the solution to 40°C. Use the graph to estimate how much solute will likely precipitate. A about 17 g B about 58 g C C

out 75 g

D about 90 g

Solubility Varies With Temperature Solubility (g/100 g H2O)

160

KNO3

140

NaNO3

120

KBr

100

NH4Cl

80 60

NaCl

40 20

Yb2(SO4)3 0

10

20

30

40

50

60

Temperature (ºC)

70

80

90

3. Evaluate  Critique the statement, “Any substance that dissolves in water is an electrolyte.” 4. Apply Concepts  At your lab table, your instructor has placed a beaker containing a cloudy solution with a solid substance resting on the bottom surface of the beaker. On the basis of what you have learned about solubility and saturation, give two plausible explanations for the appearance of the solution.

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TEKS REVIEW

10F Factors Influencing Solubility Re adi n ess

TEKS 10F Investigate factors that influence solubilities and rates of dissolution such as TEKS_TXT temperature, agitation, and surface area.

How are solubilities and rates of dissolution determined?

Vocabulary dissolution

Dissolution, which is also called solvation, is the process by which one substance, the solute, dissolves into another, the solvent. Dissolution is brought about by interactions at the particle level. Solute particles are surrounded by solvent particles until both are spread evenly throughout the entire solution. The rate of dissolution depends both on how quickly the solute comes into contact with the solvent and on the intermolecular forces between them. Generally, when a solute is added to a solvent, dissolution begins quickly. Over time, dissolution slows. Consider a crystal of table salt, sodium chloride (NaCl), as it dissolves in water. Water interacts with each ion of salt. The number of dissolved ions around the crystal increases. Soon the ions interfere with further mixing of the water with the part of the crystal that remains undissolved. You can increase the rate of dissolution by raising the temperature, increasing the surface area of the undissolved solute, and by agitating the solution as the solute dissolves. Solubility is the maximum amount of a substance that can be dissolved in a given amount of solvent. Solubility depends on temperature and pressure, as well as on the type of solvent being used. Figure 1 Solvation of an Ionic Solid in Water

Solvated ions Cation



Water molecule +

+

Anion + – + – + – + – + –

– + – +

+ – + –

– + – + – + – + – +



Surface of ionic solid TEKS 10F • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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Many compounds, such as the ionic solid shown in Figure 1 on the previous page, dissolve readily in water, resulting in aqueous solutions. Solubility is not affected by an increase in the surface area of the solute or by agitation of the solution.

How does temperature influence solubilities and rates of dissolution? For most solutions in which the solvent is a solid, an increase in temperature increases both the solubility and the rate of dissolution. For solutions in which the solute is a gas, temperature has the opposite effect on solubility. As the temperature increases, the increased kinetic energy of the gas particles in solution causes more of them to leave the solution. The solubilities of several substances are shown in Figure 2 below. Figure 2

Solubility (g/100 g H2O) at Standard Pressure Substance

0°C

Sodium chloride (NaCl) Calcium hydroxide (Ca(OH)2) Sucrose (C12H22O11) Carbon dioxide (CO2)

20°C

60°C

35.700

35.900

37.100

0.189

0.173

0.121

179.200

203.900

287.300

1.713

0.878

0.359

How does agitation influence solubilities and rates of dissolution? Agitating a solution by stirring or shaking it speeds up the rate of dissolution. This is because the movement helps carry dissolved particles away from undissolved solute, allowing more solvent to interact with it.

How does surface area influence solubilities and rates of dissolution?

Study Tip To remember how different factors affect solubility, think about the interactions between the solute particles and the solvent particles. Generally, if the solute and solvent are able to interact more with one another, then solubility increases and dissolution speeds up.

If you increase the surface area of a solid solute, it will speed up the rate of dissolution. This is because increased surface area permits more interaction between particles of solute and solvent. The surface area of a solid solute can be increased by breaking it into smaller pieces or crushing it before adding it to the solvent.

How does pressure influence solubilities and rates of dissolution? Pressure does not have a significant effect on the solubilities and rates of dissolution for most solutions, except for solutions of dissolved gases. As pressure increases, the solubility of a gas increases proportionately. At higher pressures, there are more interactions between solute particles and the solvent, and more solute-solvent interactions lead to more dissolution.

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TEKS

End-of-Course Assessment Review

1. Predict  Which of the following actions is most likely to speed up dissolution? A Crystals of sodium carbonate (Na2CO3) are crushed into a powder before being added to the solvent. B Water is cooled to 15°C before the solute is added. C An aqueous solution of sodium chloride (NaCl) is placed in a pressurized container. D Lithium iodide (LiI) is added to water and left to sit without stirring. 2. Interpret Graphs  Which of the following statements is true based on information in the graph? A Oxygen has higher solubility compared to nitrous oxide (NO) and nitrogen gas (N2). B For each of the gases, at 0°C, almost 0 g would dissolve. C For each of the gases, at 100°C, almost 0 g would dissolve. D Gases become more soluble at higher temperatures.

Solubilities of Three Gases in Water

Solubility (g/100 g H2O)

0.0100 0.0080

NO

0.0060 0.0040

O2

0.0020

N2

0

20

40

60

80

100

Temperature (ºC)

3. Evaluate  A new manager at a carbonated soft drink bottling plant is proposing to increase the temperature of the manufacturing process so that the soft drink syrup will dissolve evenly and quickly. Describe a drawback of this plan.

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TEKS REVIEW

10G Acids and Bases TEKS 10G Define acids and bases and distinguish between Arrhenius and Brønsted-Lowry definitions and predict products in acid-base reactions that form water.

What are acids and bases? Acids and bases are two classes of compounds that have opposing properties and that react readily with each other. Acids have a sour taste and turn blue litmus paper red. Bases have a bitter taste and turn red litmus paper blue.

What is the Arrhenius definition of an acid? The Arrhenius definition was the first one proposed for acids and bases. An Arrhenius acid is any substance that produces hydrogen ions (H+) in water. For example, when hydrochloric acid (HCl) is dissolved in water, it ionizes into hydrogen and chloride ions. HCl(aq)    ➝    H+(aq)   +   Cl–(aq)

Hydrochloric acid

Hydrogen ion

Chloride ion

What is the Arrhenius definition of a base? An Arrhenius base is any substance that produces hydroxide ions (OH−) when it dissolves in water. For example, when calcium hydroxide, Ca(OH)2, dissolves in water, it ionizes into calcium and hydroxide ions.

Ca(OH)2(aq)    ➝    Ca2+(aq)   +   2 OH–(aq)

Calcium hydroxide

Calcium ion

Hydroxide ion

What is the Brønsted-Lowry definition of an acid? Not all acid-base reactions take place in aqueous solutions, so hydroxide ions are not present in every reaction between an acid and a base. Brønsted and Lowry resolved this by defining acids and bases according to how protons are exchanged. A hydrogen ion, which is a hydrogen atom that has lost its only electron, is a proton. The Brønsted-Lowry definition of acids and bases regards every reaction between an acid and a base as a transfer of a proton. A Brønsted-Lowry acid, then, is any substance that donates a proton in a reaction, whether or not this takes place in an aqueous solution. For example, in a reversible reaction of ammonium (NH4+) and ammonia (NH3), ammonium reacts with hydroxide and donates a proton to form ammonia.

NH4+(aq)   +   OH–(aq)    ➝    NH3(aq)   +   H2O(l)

Ammonium ion

Hydroxide ion

Ammonia

Water

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Study Tip A chemical formula can give you a clue as to whether it represents an acid or a base. A formula for an acid will often be written with the H for hydrogen first, to emphasize that the compound can react to lose a proton. A formula for a base will often end in OH, with hydrogen at the end.

What is the Brønsted-Lowry definition of a base? A Brønsted-Lowry base is any substance that accepts a proton in a reaction. For example, consider the reversible reaction below. When ammonia (NH3) reacts with water, ammonia accepts a proton to form ammonium (NH4+).

NH3(aq)   +   H2O(aq)    ➝    NH4+(aq)   +   OH–(aq)

Ammonia

Water

Ammonium ion

Hydroxide ion

You’ll notice that while ammonia, the base, accepts a proton, water donates a proton. The water in this reaction, then, serves as an acid. In any reaction in which a proton is exchanged, the substance that accepts the proton is the base, while the substance that donates the proton is the acid.

How can you distinguish between the Arrhenius and Brønsted-Lowry definitions? The Arrhenius definitions require acid-base reactions to occur in aqueous solutions. The Brønsted-Lowry definitions require acid-base reactions to involve proton transfer. Both the Arrhenius and Brønsted-Lowry definitions of acids involve substances that readily give up protons. Unlike the Arrhenius definition of bases, which describes bases by their production of hydroxide ions, the Brønsted-Lowry definition describes bases by their ability to accept a proton. In this way, the Brønsted-Lowry definition includes all Arrhenius bases that produce hydroxide ions that can accept a proton. But the Brønsted-Lowry definition also covers other bases, such as ammonia, NH3, that do not contain hydroxide ions.

Can you predict other products in acid-base reactions that form water? When acids and bases react, the reaction usually forms water and always forms a salt. A salt is any substance that is formed from a positive and a negative ion. Consider the reaction of sodium hydroxide (NaOH) and hydrochloric acid (HCl) shown in Figure 1 below. When the acid loses a proton and the base provides a hydroxide ion that accepts a proton, the product is water. The sodium (Na+) and the chloride (Cl–) ions remain in solution as a dissolved salt, sodium chloride, NaCl(aq). Figure 1 A Neutralization Reaction

+++

NaOH NaOH NaOH Sodium Sodium Sodium hydroxide hydroxide hydroxide

+++

HCl HCl HCl

−−−

+ ++ − −− Na Na Na ClCl Cl

Hydrochloric Hydrochloric Hydrochloric acid acid acid

+++

HH H OO O 2 22

Dissolved Dissolved Dissolved sodium sodium sodium chloride chloride chloride

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TEKS

End-of-Course Assessment Review

1. Define  Which of the following describes a Brønsted-Lowry base? A It is a substance that produces a hydrogen ion in a solution of water. B It is a substance that produces a hydroxide ion in a solution of water. C It is a substance that readily accepts a proton in a reaction. D It is a substance that readily donates a proton in a reaction. 2. Predict  A chemist is analyzing a substance known to be an Arrhenius acid. Which of the following substances could it be? A H2O B HNO3 C KOH D NH3– 3. Predict  Acid indigestion is sometimes neutralized with an antacid such as magnesium hydroxide (Mg(OH)2). What products will be released when the antacid reacts with the hydrochloric acid found in the stomach? A H2O and OH– B MgOH– and H3O2 C Cl(OH)2 and HNa D MgCl2 and H2O 4. Infer  An acid-base reaction produces water and barium iodide, a salt. What were the likely reactants? A hydroiodic acid and barium hydroxide B baric acid and iodine hydroxide C hydrobarium and iodine peroxide D water and barium iodide 5. Predict  Hydrocyanic acid (HCN) is neutralized when it reacts with sodium hydroxide (NaOH). Predict the products of this reaction and write the balanced chemical equation.

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TEKS REVIEW

10H Types of Reactions Re adi n ess

TEKS 10H Understand and differentiate among acid-base reactions, precipitation reactions, and TEKS_TXT oxidation-reduction reactions.

Vocabulary acid-base reaction precipitation reaction oxidation-reduction reaction

How can you identify acid-base reactions? Acid-base reactions are reactions that occur between an acid and a base in which a proton is transferred. The product includes a salt and, often, water. If a chemical equation shows the transfer of a hydrogen ion (H+), and the formation of both a salt and water, then you can identify it as an acid-base reaction. In the equation below, a proton is transferred from the chloride ion to the hydroxide ion. As a result, a salt, calcium chloride (CaCl2), and water are formed. Ca(OH)2 (aq)   +   HCl(aq)   ➝   CaCl2 (aq)   +   H2O

calcium hydroxide

hydrochloric acid

calcium chloride

water

In the reaction below, water was not formed, but one compound contributed a proton to the other compound and a salt, ammonium sulfate, formed. This indicates that the reaction is an acid-base reaction. H2SO4 (aq) + 2NH3 (aq) ➝ (NH4)2SO4 (aq)

How can you identify precipitation reactions? Precipitation reactions are reactions that occur when two aqueous solutions react and produce a solid precipitate. The occurrence of a precipitate, either physically or in the chemical equation, is an indicator of a precipitation reaction. If the phases of the compounds are not indicated in the equation, you might need to recall the general rules of solubility to determine whether a precipitate is formed. For example, in the reaction of potassium chromate and silver nitrate, you would need to recall that chromate is insoluble unless it is combined with a group 1 metal or ammonia. Since silver chromate (Ag2CrO4) would precipitate in an aqueous solution, this reaction is a precipitation reaction. K2CrO4    +   2AgNO3   ➝   2KNO3   +   Ag2CrO4

potassium chromate silver nitrate

potassium nitrate

silver chromate

As you can see from the notation of the phases in the reaction below, nickel phosphate, (Ni3(PO4)2), is a solid. This tells you that it is also a precipitation reaction. 3NiCl2(aq) + 2Na3PO4(aq) ➝ Ni3(PO4)2(s) + 6NaCl(aq)

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How can you identify oxidation-reduction reactions? An oxidation-reduction reaction, also called a redox reaction, is any reaction that involves the exchange of an electron or oxygen. The exchange can take place between ionic substances or in the formation of a covalent bond. The substance that loses electrons or gains oxygen is said to be oxidized. The substance that gains electrons or loses oxygen is said to be reduced. Oxidation and reduction always occur simultaneously. No oxidation occurs without reduction, and no reduction occurs without oxidation. How can you know if an electron is transferred during a reaction? One indicator is light. Sometimes, though, the light produced is not visible, or the chemical equation does not indicate whether light was produced. In those cases, you need to recall the rules in Figure 1 to identify whether an electron was exchanged. In general, when the oxidation number of a certain atom is different in one of the reactants than it is in one of the products, it means that at least one electron was exchanged. Figure 1

Rules for Determining Oxidation Numbers 1. The oxidation number of an atom in its elemental form is zero. Example: For O2 the oxidation number of oxygen is 0. 2. In monoatomic ions, the oxidation number is the same as the charge on the ion. Example: For Al3+ the oxidation number of aluminum is +3. 3. In molecules and complex ions, the oxidation number of the most electronegative atom is the same as the charge on it when the atom is an ion. Example: For HCl the oxidation number of chloride is −1. 4. The oxidation number of hydrogen is usually +1. If hydrogen is bonded to a metal, then the oxidation number is −1. 5. In a neutral compound, the sum of oxidation numbers is always zero. Example: In Fe2O3 the oxidation number of iron is +3 and of oxygen is −2 because 2(+3) + 3(−2) = 0. 6. In a polyatomic ion, the sum of oxidation numbers is equal to the charge of the ion. Example: For NH4+ the oxidation number of nitrogen is −3 and of hydrogen is +1 because 1(−3) + 4(+1) = +1.

Study Tip Many types of reactions are oxidation-reduction reactions. In fact, all combustion reactions and single-replacement reactions, and many combination and decomposition reactions, are oxidation-reduction reactions.

Consider the reaction between methane (CH4) and oxygen (O2). CH4( g)  +  2O2( g)  ➝  CO2( g)  +  2H2O( g)

methane

oxygen

carbon dioxide

water

Based on the rules, you can see that the oxidation number of oxygen gas is zero, while the oxidation number of the oxygen in both products is −2. You can also see that the oxidation number of the carbon in methane is −4, and the oxidation number in the carbon dioxide in the product is +4. Since electrons were exchanged, you can conclude this was an oxidationreduction reaction.

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In the reaction below, both reactants began with an oxidation number of zero. In the products, the oxidation number of magnesium is +2, and the oxidation number of nitrate is −3. This reaction is also an oxidationreduction reaction. 3Mg(s) + N2(g) ➝ Mg3N2(s)

TEKS

End-of-Course Assessment Review

1. Differentiate  What is a difference between an acid-base reaction and an oxidation-reduction reaction? A Electrons, not protons, are transferred in all acid-base reactions. B Electrons, not protons, are transferred in all oxidation-reduction reactions. C Neutrons, not protons, are transferred in all acid-base reactions. D Protons, not electrons, are transferred in all oxidation-reduction reactions. 2. Classify  Which of the following is a precipitation reaction? A 3NH4OH(aq) + AlCl3(aq) ➝ 3NH4Cl(aq) + Al(OH)3(s) B 2H2( g) + O2( g) ➝2H2O(l) C 2HCl(aq) + Na2S(aq) ➝ H2S( g) + 2NaCl(aq) D H2CO3(aq) ➝ H2O(l) + CO2( g) 3. Apply Concepts  A chemistry student added a sample of molten sodium to a flask of chlorine gas. A bright light was produced, and a white solid formed on the inside of the flask. What kind of reaction occurred? A acid-base reaction B precipitation reaction C oxidation-reduction reaction D neutralization reaction 4. Differentiate  Dilute hydrochloric acid (HCl) is added to a test tube containing a grain of magnesium ribbon. How might you determine whether the reaction that occurs is an acid-base reaction, precipitation reaction, or oxidation-reduction reaction? 5. Evaluate  Is it possible for a reaction to be both an acid-base reaction and a precipitation reaction? Explain why or why not.

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TEKS REVIEW

10I pH of a Solution TEKS 10I Define pH and use the hydrogen or hydroxide ion concentrations to calculate the pH of a solution.

What is pH?

Vocabulary pH

The strength of an acid depends on the concentration of hydrogen ions (H+) in a solution, which can be expressed in terms of molarity. For example, the concentration of hydrogen ions in water is 1.0 × 10−7 M. This means that for each liter of water, there is 0.0000001 mole of hydrogen ions. A much easier way to express the strength of an acid is by its pH, which is the negative logarithm of the hydrogen-ion concentration. pH = −log[H+] The pH of water is the negative of the log of 1.0 × 10−7 M, which is 7. Water has an equal amount of hydrogen ions and hydroxide ions, so it is considered neutral. Solutions with a pH lower than 7 are considered acidic, and solutions with a pH higher than 7 are considered basic. Each change of a whole number in pH represents a difference of a factor of 10 in the concentration of hydrogen ions. You can see the relative strengths of some familiar acids and bases in Figure 1. Figure 1

Relative Strengths of Common Acids and Bases HCl

Nitric acid

HNO3

Sulfuric acid

H2SO4

Phosphoric acid

H3PO4

Ethanoic acid

CH3COOH

Carbonic acid

H2CO3

Hypochlorous acid

HClO

Ammonia

NH3

Sodium silicate

Na2SiO3

Calcium hydroxide

CA(OH)2

Sodium hydroxide

NaOH

Potassium hydroxide

KOH

Strong acids

Increasing strength of acid

Hydrochloric acid

Relative strength

Neutral solution Increasing strength of base

Formula

Substance

Strong bases

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Study Tip The log of a number is taken from how many times 1 has to be multiplied by 10 to reach that number. The log of 1000, or (1 × 103) is 3. For numbers that are very small, the log is a negative number because 1 is divided by 10. The 1 log of , or (1 × 10−3), 1000 is −3.

How is the hydrogen ion concentration used to calculate the pH of a solution? Consider an aqueous solution with a hydrogen ion (H+) concentration of 1.0 × 10−5 M. The pH is calculated using the negative of the log of the hydrogen ion concentration. pH = −log (1.0 × 10−5) = −(−5) = 5 The hydrogen ion concentration of a solution can also be determined from its pH through a reverse of the calculation. The molar concentration of any substance is described using brackets, such as [H+]. [H+] = antilog (−pH) (Note: Sometimes the antilog is referred to as the “inverse of the log” or as “log–1.” On some calculators, the button for this function is labeled “10 x.”) For example, to find the hydrogen ion concentration of a solution with a pH of 11, you can calculate the antilog of −11. [H+] = antilog (−11) = 1.0 × 10−11 M Here is how you could calculate the pH of a solution with a hydrogen ion concentration of 6.3 × 10−4 M pH = −log (6.3 × 10−4) = −(−3.2) = 3.2

How is the hydroxide ion concentration used to calculate the pH of a solution? The concentration of hydroxide ions is closely related to the pH of a solution. Recall that as water molecules disassociate, equal amounts of hydroxide ions and hydrogen ions are formed. H O(l)   ➝   H+(aq)   +   OH−(aq) ➝



2

Hydrogen ion

Hydroxide ion

For all aqueous solutions, this self-ionization equilibrium is occurring, with [H+][OH–] = 10–14. The pOH, or relative concentration of hydroxide ions, can be calculated similar to the way in which pH is calculated. pOH = −log [OH−] So, taking the log of the equilibrium expression, one can derive this relationship: pH + pOH = 14 This is why, in water, the pH and the pOH are both 7. If an acid is added to the water, the concentration of hydrogen ions increases, lowering the pH. As the hydrogen ion concentration increases, the hydroxide concentration decreases. But the sum of the pH and the pOH is always 14. On the basis of this relationship, you can calculate the pOH of a solution as long as you know its pH.

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The hydroxide concentration can be used to find pH based on the same relationship. For example, if a solution has a hydroxide ion concentration of 1.3 × 10−10, then you can calculate the pOH and find the pH. pOH = −log (1.3 × 10−10) = 9.9 Therefore, pH = 14 – 9.9 = 4.1 Alternatively, if you know the pOH of a solution, then you can calculate the hydrogen ion concentration. Consider a solution with a pOH of 4. How could you calculate the concentration of hydrogen ions of the solution? The pH is 10. So the concentration of hydrogen ions is 1.0 × 10−10.

TEKS

End-of-Course Assessment Review

1. Define  Which of the following most accurately describes pH? A a negative logarithm of hydroxide ion concentration B a scale indicating the number of hydroxide ions C a scale indicating the number of hydrogen ions D a negative logarithm of hydrogen ion concentration 2. Compare and Contrast  Solution A has a pH of 7, and Solution B has a pOH of 9. Which solution has a higher concentration of hydroxide ions? A Solution A B Solution B C They have the same concentration of hydroxide ions. D The concentrations of hydroxide ions cannot be determined. 3. Predict  What will happen to the pH of a substance if the concentration of hydrogen ions increases from 1.0 × 10−8 to 1.0 × 10−6? A The pH will change from 8 to 6. B The pH will change from −8 to −6. C The pOH will change from 8 to 6. D The pH will change from 6 to 8. 4. Calculate  An average cup of coffee is composed of many organic chemicals, including caffeine. Some of them are slightly acidic, and some are slightly basic. Overall, coffee has a hydroxide concentration of 1.9 × 10−9. What is its pH?

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TEKS

Degrees of Dissociation 10J for Acids and Bases REVIEW

Re adi n ess

TEKS 10J

Distinguish between degrees of dissociation for strong and weak acids and bases. TEKS_TXT

To what degree do strong acids and strong bases dissociate?

Vocabulary strong acid strong base weak acid weak base

In general, strong acids and strong bases dissociate, or ionize, completely in aqueous solution. In other words, almost 100 percent of a strong acid or strong base interacts with water and ionizes. A strong dissociation is described using one arrow pointing to the right, as shown. HA (aq)    ➝      H+ (aq)      +    A– (aq) If a given number of moles of a strong acid, represented as HA, is added to a solution, then that same amount of its dissociated ions will be in the solution. Figure 1 also describes the dissociation of a strong acid. The bars represent the relative amounts of the acid and the ions it forms in solution. Figure 1 Relative number of moles

Dissociation of a Strong Acid H 3O +

HA

A–

Complete dissociation

HA(aq) + H2O(l)

H3O+(aq) + A–(aq)

Strong bases generally dissociate into cations and hydroxide ions in water. For example, consider the dissociation of sodium hydroxide (NaOH). NaOH (aq)    ➝    Na+ (aq)    +    OH– (aq)

To what degree do weak acids and weak bases dissociate? Weak acids and weak bases ionize only slightly in aqueous solution. This weaker dissociation can be expressed using arrows pointing in both directions. The arrows indicate that after a certain amount of the weak acid or base dissociates, the reverse reaction occurs at an equal rate. HA (aq)    ➝      H+ (aq)      +    A– (aq) ➝

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If a given number of moles of a weak acid, represented as HA in Figure 2, are added to water, then nearly that same amount of the weak acid will remain in the solution. Only a very small portion of the molecules will dissociate, representing a small fraction of the entire solution. The same is true for a weak base; only a small fraction of the molecules will dissociate. Consider hydrofluoric acid, which is a weak acid. In an aqueous solution at a given temperature, only a fixed amount dissociates. Most of the solution is composed of non-ionized hydrofluoric acid. HF (aq)    ➝    H+ (aq)    +    F– (aq) ➝

Figure 2 Relative number of moles

Dissociation of a Weak Acid HA HA Little dissociation H3O + HA(aq) + H2O(l)

A–

H3O+(aq) + A–(aq)

How can you distinguish between the degrees of dissociation for weak acids and bases?

Study Tip When comparing values for Ka or Kb, remember to interpret the scientific notation correctly. Negative exponents indicate values that are smaller than 1.0. Stronger acids and bases will have constants that are closer to, or greater than a value of 1 than will weaker acids and bases.

The degree of dissociation of a weak acid in water is represented by the acid dissociation constant (Ka). The degree of dissociation of a weak base in water is represented by the base dissociation constant (Kb). Both are calculated using the equilibrium concentrations of the acid or base, dividing the product of the concentration of the dissociated form of an acid by the concentration of the undissociated form. For example, the dissociation constant of hydrofluoric acid is calculated by the product of the concentrations of the dissociated ions of the hydrogen ion and fluoride ion divided by the concentration of hydrofluoric acid. Ka =

[H+][F–] [HF]

= 6.3 × 10−4

At 25°C, Ka of hydrofluoric acid is 6.3 × 10−4.

The strength of acids and bases depends on the degree to which they dissociate. The greater the degree of dissociation, the greater their strength. Algebraically, this corresponds to a larger numerator, which increases the value of Ka or Kb. This explains why stronger acids and bases have larger dissociation constants than weaker acids and bases. Figure 3 lists the dissociation constants of several weak acids and bases. Note that some acids can dissociate more than one proton. The Ka for the first proton is a smaller value than the Ka for the second proton.

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Figure 3

Dissociation Constants at 25ºC º Weak Acids

Weak Bases

Name (formula)

Ka

Carbonic acid (H2CO3)

4.3 × 10−7

Ammonia (NH3)

1.8 × 10−5

Carbonic acid (HCO3−)

4.8 ×10−11

Dimethylamine ((CH3)2NH)

5.1 × 10−4

Nitrous acid (HNO2)

4.5 × 10−4

Hydrazine (N2H4)

1.7 × 10−6

Phosphoric acid (H3PO4)

6.9 × 10−3

Hydroxylamine (NH2OH)

1.1 × 10−8

Phosphoric acid (H2PO4−)

6.3 × 10−8

Pyridine (C5H5N)

1.4 × 10−9

Sulfurous acid (H2SO3)

1.3 × 10−2

Urea (NH2CONH2)

1.5 × 10−14

TEKS

Name (formula)

Kb

End-of-Course Assessment Review

1. Distinguish  Which option below represents the strongest acid? A Ka = 4.0 × 10−10 B Kb = 4.0 × 10−10 C complete dissociation of hydroxide ions D complete dissociation of hydrogen ions 2. Compare and Contrast  Based upon their dissociation constants at 25°C, which would be the weakest base? A 0.5M ammonia C 0.5M hydroxylamine B 0.5M dimethylamine

D 0.5M urea

3. Compare and Contrast  Based on their dissociation constants at 25°C, which of the following would be the best conductor of electricity? A sulfurous acid C phosphoric acid B urea

D pyridine

4. Summarize  In a few sentences, state generally how strong acids and bases dissociate and how weak acids and bases dissociate. 5. Analyze  A technician in a chemistry lab prepares 100 mL of 1.0M nitrous acid at 25°C. Write the balanced chemical equation showing the dissociation of the acid.

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TEKS REVIEW

11A Energy and Its Forms TEKS 11A Understand energy and its forms, including kinetic, potential, chemical, and thermal energies.

What is energy?

Vocabulary energy kinetic energy potential energy chemical energy thermal energy

Energy is the capacity to do work or produce heat. Energy does not have mass or volume. Units called joules (J) are used to measure energy. A common non-SI unit of energy is the calorie. One calorie (cal) is the quantity of heat that raises the temperature of 1 g of water by 1 degree Celsius. The conversion relationships for joules and calories are:

1 J = 0.2390 cal

1 cal = 4.184 J

Energy values can also be expressed in kilojoules (kJ) and kilocalories (kcal). A kilojoule is 1000 joules; a kilocalorie is 1000 calories.

What are energy’s forms? Energy exists in various forms, such as light, heat, sound, electrical energy, gravitational potential energy, chemical energy, and nuclear energy. However, all of these forms belong to one of two fundamental categories: kinetic energy or potential energy. Moving objects and particles have kinetic energy, and objects that are not moving have potential energy. Figure 1 classifies some forms of energy into the two major categories. Figure 1

Forms of Kinetic and Potential Energy Kinetic energy

Study Tip Remember that energy is neither created nor destroyed. It may transform from one type to another, and less of it may be available to do work, but it never goes away.

Potential energy

Light

Nuclear energy

Heat

Gravitational potential energy

Sound

Chemical energy

Mechanical motion

Mechanical potential energy

Energy is always being transformed from one form to another. Some energy changes involve single transformations, while others involve many transformations. A toaster performs a single transformation as it transforms electrical energy to thermal energy. In a car engine, a series of transformations are required to make the car move. The chemical energy in the gasoline is ignited, producing light, heat, and sound in the engine. The ignition of the gasoline also pushes a piston, which makes the car move, resulting in mechanical motion.

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What is kinetic energy? Kinetic energy is the energy of matter in motion. In the simplest case, the kinetic energy of a particle in motion can be calculated using the equation KE = ½ mv2. If m is the object’s mass in kilograms (kg), and v is the speed in meters per second (m/s), the kinetic energy is in units of joules (J). The kinetic energies of the molecules of a chemical substance differ from each other because their individual velocities differ in both speed and direction. Often, the quantity of interest is not the kinetic energy of a specific molecule, but the average kinetic energy of all of the molecules. The average kinetic energy of a substance is also its temperature. At any given temperature, the particles of all substances have the same average kinetic energy. Absolute zero is the temperature at which the motion of particles theoretically is no longer possible.

What is potential energy? Potential energy is the ability of a body to do work because of its position, composition, or configuration. It is essentially stored energy. Some examples of potential energy include a compressed spring, a book lifted off the ground, and a stretched rubber band. Applying force by compressing the spring, lifting the book, or stretching the rubber band does work that is stored as potential energy. Once those forces are removed, the stored energy in the spring, the book, or the rubber band is transformed into kinetic energy. The spring or rubber band bounces back and the book falls to the floor. The book example describes gravitational potential energy, which is potential energy related to an object’s height. The gravitational potential energy of an object is calculated by multiplying the object’s weight by its height (the distance the object was moved).

What is chemical energy? Chemical energy is a form of potential energy stored in the chemical bonds of a substance. The kinds of atoms and the arrangement of the atoms in a substance determine the amount of energy stored in the substance. When bonds are broken in reactants, there is an input of energy. When new bonds are formed in the products of a reaction, there is an output of energy. During a reaction, energy is absorbed from or released into the surroundings, depending on whether more or less energy is required to break the old bonds than is released in forming the new ones. If energy is absorbed, the chemical potential energy of the substance increases; if it is released, chemical potential energy decreases. Released energy can be in the form of heat, sound, light, or a combination of these. Thermochemistry is the study of energy changes that occur during chemical reactions and changes in state. A major goal of thermochemistry is to examine the flow of heat between a system and its surroundings.

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What is thermal energy? Thermal energy is the total potential and kinetic energy of the molecules of a substance caused by their arrangement and constant motion. The thermal energy of a substance can be changed by exchanging heat with its surroundings. Heat is thermal energy moving from a warmer object to a cooler object. A simple example of thermal energy is eating ice cream on a hot day. Fastmoving particles in the warm air make the particles of ice cream move faster. As the kinetic energy of the particles increases, so does the thermal energy of the ice cream. Eventually, the ice cream melts.

TEKS

End-of-Course Assessment Review

1. Define  Energy is best described as the capacity to A transfer heat. B do work. C cause a chemical reaction. D compress a spring. 2. Evaluate  Which of the following statements is true of chemical energy? A Chemical energy may be increased by applying heat. B In a solid, the atoms or molecules vibrate but do not move relative to each other because of chemical energy. C In a water molecule, the oxygen atom bonds to two hydrogen atoms because this minimizes the chemical energy. D Energy may be released as heat and light during a chemical reaction. 3. Classify  List four examples of kinetic energy you can identify as you sit at a computer. 4. Classify  List and describe the energy transformations that occur as a log is burned in a fireplace. 5. Evaluate  A chemistry student is describing a piece of iron on a lab table and writes “Because the iron is not moving or reacting, it doesn’t have any energy.” Explain three reasons why this statement is not true.

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TEKS

Conservation of Energy 11B and Heat Transfer REVIEW

TEKS 11B

Understand the law of conservation of energy and the processes of heat transfer. TEKS_TXT

What are the roles of systems and surroundings in energy conversions?

Vocabulary system surroundings law of conservation of energy heat endothermic process exothermic process

Thermochemistry is the study of energy changes that occur during chemical reactions and changes of state. Every substance has a certain amount of chemical potential energy. During a chemical reaction, reactants are transformed into products with more or less energy than the reactants. When gasoline is burned in a car’s engine, some of the chemical potential energy is transformed into mechanical energy, which is used to propel the car. At the same time, thermal energy is released in the form of heat. Energy changes occur as work or heat transfer, or a combination of both. A major goal of thermochemistry is to examine the flow of heat between a system and its surroundings. A system is the process being studied. In the example above, the system is the car’s engine. But every system occurs in some context, its surroundings—everything else in the universe. In thermochemical experiments, you can consider the region in the immediate vicinity of the system as the surroundings. In a system, only a portion of input energy is converted into output energy. The rest of the energy is lost to the surroundings as thermal energy. Figure 1 illustrates this concept.

Figure 1 Flow of Energy

Input Energy

Process

Output Energy

Thermal energy to surroundings

What is the law of conservation of energy? The law of conservation of energy states that in any chemical or physical process, energy is neither created nor destroyed. During any chemical or physical process, the amount of energy in the universe remains unchanged. If the energy of the system increases during that process, the energy of the surroundings must decrease by the same amount. Likewise, if the energy of the system decreases during that process, the energy of the surroundings must increase by the same amount. TEKS 11B • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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This law may not seem obvious. After all, a bouncing ball and a swinging pendulum eventually come to rest. Motors and engines are not perfectly efficient. All of the energy available in gasoline is not used to power the engine. The light energy emitted by an electric lamp is always less than the electrical energy consumed by the lamp. What happens to the energy that is not used to power the engine or emit light from the lamp? Well, what do a car engine and light bulb have in common? They both transfer heat to their surroundings.

How is heat transferred between the system and the surroundings? Heat, represented by q, is energy that transfers from one object to another because of a temperature difference between the objects. One of the effects of adding heat to an object is an increase in its temperature. Heat flows spontaneously from a warmer object to a cooler object. If two objects remain in contact, heat will flow from the warmer object to the cooler object until the temperatures of both objects are the same. Like other forms of energy, heat is measured in joules and calories. However, it is incorrect to say that an object has “heat energy”; heat is the transfer of thermal energy (which may increase or decrease an object’s thermal energy).

Study Tip You can find examples of heat transfer all around you. As you observe them, try to classify them as exothermic or endothermic processes.

In thermochemical calculations, the direction of heat flow is given from the point of view of the system. Heat is absorbed from the surroundings in an endothermic process. In an endothermic process, the system gains heat as the surroundings lose heat. In Figure 2, the system (the body) gains heat from its surroundings (the fire). Heat flowing into a system from its surroundings is defined as positive; q has a positive value in this situation. An exothermic process is one that releases heat to the surroundings. In an exothermic process, the system loses heat as the surroundings gain heat. In the figure below, the system (the body) loses heat to the surroundings (the perspiration on the skin and the air). Heat flowing out of a system into its surroundings is defined as negative; q has a negative value in this situation.

Figure 2 Endothermic and Exothermic Processes

Surroundings System

Surroundings HEAT System

HEAT

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TEKS

End-of-Course Assessment Review

1. Infer  The glass of a lit 75-watt incandescent bulb is hotter to the touch than the glass of a 25-watt compact fluorescent (CFL) bulb that emits the same amount of light. What could be a reason for this? A The CFL bulb emits more thermal energy than an incandescent bulb. B More of the electricity used by an incandescent bulb is converted to thermal energy than that of the CFL bulb. C More of the electricity used by an incandescent bulb is converted to light energy. D The CFL bulb contains more chemical potential energy than the incandescent bulb. 2. Analyze  An engine makes a car move because of the force generated by small explosions as gasoline is burned in air and produces carbon dioxide and water. Why isn’t the stored chemical energy converted completely into kinetic energy of the car’s motion? A During the explosion reaction, some energy is lost to the surroundings as heat. B Some of the energy from the gasoline is destroyed during the process. C Because the car is so heavy, it takes extra energy to get started. D Some energy is transferred from the surroundings to the car’s engine. 3. Infer  A forgotten ice pop lies melting on a deck on a hot summer day. What is the direction of heat flow as the ice pop melts? Is the process endothermic or exothermic? Explain.

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TEKS

Energy Changes in 11C Chemical Reactions REVIEW

Readi ness

TEKS 11C Use thermochemical equations to calculate energy changes that occur in chemical reactions and classify reactions as exothermic or endothermic.

Vocabulary bond dissociation energy

How can you classify energy changes in chemical reactions as exothermic or endothermic? All chemical reactions involve a change in chemical potential energy. These energy changes occur whenever bonds between atoms are broken in the reactant molecules and formed in the product molecules. Because bonded atoms represent a configuration of minimum potential energy, breaking a bond between a pair of atoms always absorbs energy from the surroundings; it is an endothermic process. Conversely, forming atomic bonds always releases energy to the surroundings, an exothermic process.

Study Tip Do not confuse an exothermic reaction with a spontaneous reaction. Even though an exothermic reaction may ultimately release energy to its surroundings, it still often needs some external activation energy to start.

The chemical potential energy change for a reaction can be estimated by adding up all of the energy required to break the bonds in the reactant molecules and then subtracting all of the energy released when new bonds are formed in the product molecules. A more refined calculation would account for energy changes as intermolecular bonds are formed and broken in solids and liquids. While it is convenient to think of bond breaking and bond formation as separate steps, they actually occur simultaneously. Energy released by bond formation is available to participate in breaking other bonds. The bonded atoms are considered the system, and everything else is the surroundings. This means that the energy change of an endothermic reaction is a positive value, indicating that energy is added to the system. After an endothermic change, the system has more energy “stored” in its chemical bonds. For exothermic reactions, the energy change is a negative value. This indicates that energy has left the system. After an exothermic change, the system will have less energy “stored” in its chemical bonds. The chemical potential energy change of a reaction (∆H) can be calculated by subtracting the energy of the reactants from the energy of the products. ∆H = Hproducts − Hreactants The difference is positive in an endothermic reaction, indicating that more energy is required to break bonds in the reactants than is released by forming the new bonds of the products. The difference is negative in an exothermic reaction, indicating that more energy is released from the formed bonds than is required to break bonds.

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What is bond dissociation energy? For every pair of bonding atoms, the bond dissociation energy is the energy required to break the bond. It is always a positive number, and it is typically reported in units of kilojoules per mole (kJ/mol). The bond energy is a measure of the strength of a chemical bond. The amount of energy released when the bond is formed is numerically the same as the bond energy, but it has a negative value instead of a positive one, because forming bonds is an exothermic process. Note that for a bond between two atoms, the bond energy will vary somewhat depending upon the molecule that contains the bond. Therefore, an average bond energy is often used. Figure 1 shows some common average bond energies. Figure 1

Average Bond Energies Bond H H Cl

Cl

kJ/mol 435

Bond H Cl

243

H O

kJ/mol 432

Bond C C

464

C C

kJ/mol

kJ/mol

347

Bond O O

657

O O

498

142

How do thermochemical equations show energy changes that occur in chemical reactions? A thermochemical equation is a balanced chemical equation that shows the phase of all reactants and products and that includes the energy change (∆H) associated with the reaction. If the reaction is endothermic, the energy change is written with the reactants because it is absorbed by the reaction. If the reaction is exothermic, the energy change is written as a product because it is released by the reaction. Energy change can be determined experimentally, or it can be estimated from the bond dissociation energies when all reactants and products are in the gaseous phase. When ∆H is determined from reactions carried out under standard conditions (one atmosphere pressure and 298 kelvins (K)), the symbol ∆H° is used. Consider the following thermochemical equation for a reaction between gaseous hydrogen and chlorine to form hydrogen chloride: H2( g)  +  Cl2( g)  ➝  2HCl( g)   ∆H = –184 kJ This reaction involves the breaking of an H—H bond and a Cl—Cl bond, together with the formation of two H—Cl bonds. The energy change of the reaction can be estimated from average bond energies. Notice that the estimated energy change is slightly different from the observed value. ∆H = (energy required to break bonds) – (energy required to form bonds) ∆H = (435 kJ + 243 kJ) – 2 (432 kJ) = –186 kJ This is an exothermic reaction. The energy released by the formation of H—Cl bonds is greater than the energy required to break apart the hydrogen and chlorine molecules. TEKS 11C • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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TEKS

End-of-Course Assessment Review

1. Infer  A reaction takes place in a vessel, causing the temperature inside to increase. Which statement best describes the process? A Energy is being released because chemical bonds are forming. B More energy is being released by forming bonds in the products than is required to break the bonds in the reactants. C The reaction vessel is insulated. D The reaction vessel contains a combustion reaction. 2. Calculate  What is the estimated ∆H for this thermochemical equation?   2H2( g)  +  O2( g)  ➝  2H2O( g)   ∆H = ? A 488 kD

C 442 kJ

B −922 kD

D −488 kJ

3. Calculate  Assume that the average bond energy of the F−F bond is 158 kJ/mol. Use the following thermochemical equation to estimate the average bond energy of the H−F bond:   H2( g)  +  F2( g)  ➝  2HF( g)   ∆H = −546 kJ A 1139 kJ/mol

C 23.5 kJ/mol

B 570 kJ/mol

D −570 kJ/mol

4. Classify  Why do endothermic reactions absorb energy from their environment?

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TEKS

Heat, Mass, Temperature Change, 11D and Specific Heat REVIEW

TEKS 11D

Perform calculations involving heat, mass, temperature change, and specific heat. TEKS_TXT

Vocabulary heat capacity specific heat molar heat of fusion molar heat of vaporization molar heat of solidification molar heat of condensation molar heat of solution

What is specific heat? Heat capacity is the quantity of heat required to raise the temperature of a substance by one degree Celsius. Heat capacity is an extensive property, because it depends on the mass and chemical composition of an object. Specific heat is an intensive property. It does not vary with the amount of the substance. The specific heat of a substance is the amount of heat needed to raise the temperature of one gram by one degree Celsius. The unit of specific heat is joules per gram-degree Celsius (J/g•°C). The equation for specific heat is: Q = mc∆T, Q represents the heat input, m is for mass, c is for specific heat, and ∆T is for the change in temperature.

Study Tip When ice is placed under a heat lamp, three processes can be involved: heating ice to 0°C, the melting of ice at 0°C, and the heating of liquid water.

For example, how much heat is necessary to raise the temperature of 10 g of liquid water from 25°C to 45°C? The specific heat of liquid water is 4.18 J/g•°C. If you substitute the known quantities, you can see that it is 836 J.

Q = mc∆T



= (10 g)(4.18 J/g•°C)(45°C – 25°C)



= 836 J

What are the molar heat of fusion and the molar heat of vaporization? Most of the time, adding heat to a substance causes its temperature to increase. However, during phase changes, such as melting or boiling, the temperature stays the same until the change is complete. When the temperature of a substance is at its melting point or boiling point, the added heat overcomes the forces that hold the substance together as a solid (at the melting point) or as a liquid (at the boiling point). The molar heat of fusion (∆Hfus) is the amount of heat needed to change one mole of a substance from a solid to a liquid at the melting point—to melt it. Similarly, the molar heat of vaporization (∆Hvap) is the amount of heat needed to change one mole of a substance from a liquid to a vapor at the boiling point. TEKS 11D • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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The equations that use molar heats include: For melting: Q = L∆Hfus For boiling: Q = L∆Hvap (L is the number of moles of the substance). Figure 1 shows the values of specific heat and constants for water. Figure 1

Specific Heat and Other Constants for Water Temperature

Phase

c (J/g∙°C)

𝚫Hfus (kJ/mol)

𝚫Hvap (kJ/mol)

6.01

40.7

Water, –10°C

Solid

2.05

Water, 25°C

Liquid

4.18

Water, 100°C

Gas

2.08

For example, how could you figure out how much heat is necessary to melt 36.0 g of ice into liquid water at 0°C? The molecular mass of water is 18.0 g/mol. First, you would convert grams to moles. 36.0 g ×

1 mol = 2.0 mol 18.0 g

Next, find the molar heat of fusion in Figure 1 and substitute it into the equation. Q = (2.0 mol)(6.01 kJ/mol) = 12.0 kJ Notice that for water, both ∆Hfus and ∆Hvap are positive values. This fact indicates that both melting and boiling are endothermic processes (heat is absorbed from surroundings). The two inverse processes—solidification and condensation—are exothermic processes (heat is released). Their energy changes are called the molar heat of solidification (∆Hsolid) and the molar heat of condensation (∆Hcond). The molar heat of solidification is the heat lost when one mole of a liquid substance solidifies at a constant temperature. The molar heat of condensation is the amount of heat released when one mole of a vapor condenses at its normal boiling point. These values are equal in quantity but opposite in sign to the energy changes for melting and boiling: ∆Hcond = –∆Hvap ∆Hsolid = –∆Hfus During the formation of a solution, heat is either released or absorbed. The change in heat caused by the dissolution of one mole of substance is the molar heat of solution (ΔHsoln).

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TEKS

End-of-Course Assessment Review

1. Calculate  How much energy must be added to 50.0 g of water to raise its temperature from 40.0°C to 45.0°C? A 1045 joules B 58 joules C 1500 joules D 41.8 joules 2. Calculate  How much energy must be added to 36 g of ice at 0°C in order to raise its temperature to 25°C? (The molecular mass of water is 18.0 grams/mol). A 220 kilojoules B 15.8 kilojoules C 12.0 kilojoules D 3.8 kilojoules 3. Calculate  When a 250-g block of aluminum is cooled by 8°C, it is found to have given up 1794 J. What is the specific heat of aluminum? A 1.794 J/g•°C B 0.897 J/g•°C C 0.449 J/g•°C D 1.115 J/g•°C 4. Infer  Why is the molar heat of vaporization of a substance usually higher than its molar heat of fusion? 5. Apply Concepts  Provide one example of the fact that water has a relatively high specific heat.

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TEKS REVIEW

11E Calorimetry TEKS 11E

Use calorimetry to calculate the heat of a chemical process.

What is calorimetry?

Vocabulary calorimetry calorimeter

Calorimetry is the measurement of the heat flow that occurs during a chemical process or reaction. Heat flow cannot be directly measured in the same way that volume or other properties can. However, the heat of a chemical process or reaction will either increase or decrease the temperature of its surroundings. One way to measure heat flow is to measure temperature change. A calorimeter is an instrument that measures the absorption or release of heat in chemical or physical processes. One type of calorimeter, a constantpressure calorimeter, is useful for reactions that do not involve gases and that can be carried out in solutions. A simple version of a constant-pressure calorimeter uses two nested foam cups. The inner cup contains water and a thermometer. The temperature changes of the water can be used to calculate the heat of a reaction. Another type of calorimeter is a constant-volume calorimeter, sometimes called a bomb calorimeter. This calorimeter is useful for reactions that produce gases. The bomb is a high-pressure stainless steel reaction chamber that contains the substance being studied and oxygen.

Figure 1 Two Types of Calorimeter

Constant-Volume (“Bomb”) Calorimeter Electrical leads

Constant-Pressure Calorimeter Stirrer

Thermometer

Thermometer

Foam lid Insulated container

Water

Oxygen

Nested foam cups

Stirrer

Sample Water

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Study Tip The word calorimetry is made of the root word calor, meaning “heat,” and the word part -metry, meaning “measurement.”

Upon ignition, the substance combines with oxygen in a combustion reaction. The heat is transferred to a fluid, such as water, and then the temperature increase can be measured. Because a calorimeter is isolated from the external environment, the energy that is released or absorbed by the chemical reaction or process occurring within it changes the temperature of the calorimeter. If the calorimeter’s heat capacity is known, it is possible to calculate that energy. In a perfect calorimeter, no thermal energy would be transferred to or from the external environment. This would mean that the heat flow inside the calorimeter is the heat of the chemical process being studied. Actually, no calorimeter is perfect. But transfers of thermal energy to or from the environment can be minimized by careful insulation of the calorimeter.

How can you calculate energy changes with a calorimeter? A calorimeter usually contains a carefully measured mass of a substance of known specific heat, such as water. As the reaction absorbs or releases energy, the temperature of the water will change. From that temperature change, the energy that the water absorbed or released can be calculated.

Sample Problem 1 A student is using a calorimeter to measure the molar heat of solvation of calcium chloride (CaCl2). The calorimeter contains 100.0 g of water at 25.1°C. After 10.0 g of CaCl2 is fully dissolved, the temperature of the water rises to 42.7°C. (Hint: Remember that the specific heat of water is 4.18 J/g•°C.) First, write the process as a chemical equation: CaCl2(s) ➝ Ca2+(aq) + 2Cl–(aq)   ∆H = ∆Hsolution Then calculate the temperature change (∆T ) of the water. ∆T = 42.7°C – 25.1°C = 17.6°C Next, calculate the energy absorbed by the water using the following equation.

Qwater = mwatercwater∆T

However, Qwater is the energy released when 10.0 g CaCl2 dissolves, not 1 mol CaCl2. To calculate the heat per mole, first divide 10.0 g by the molar mass of CaCl2, which is 110.98 g/mol. (10.0 g CaCl2) (1.00 mol CaCl2)



mol CaCl2 =



mol CaCl2 = 0.0901 mol CaCl2

110.98 g CaCl­2

Then divide the heat change of the water by 0.0901 mol CaCl2. Notice that because the heat change of the water is positive, the heat change of forming the solution is negative.

∆Hsolution = –Qwater/number of mol CaCl2



∆Hsolution = –7360 J/0.0901 mol CaCl2



∆Hsolution = –81,700 J/mol CaCl2



∆Hsolution = –81.7 kJ/mol CaCl2

Qwater = (100.0 g) (4.18 J/g•°C) (17.6°C)

Qwater = 7360 J

Qwater is positive, indicating that the water absorbs energy from the reaction. In a perfect calorimeter, Qwater equals the amount of energy released by the dissolving of calcium chloride.

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Sample Problem 2 Liquid hydrazine (N2H4) and oxygen combine in anexplosive reaction described by the following equation: N2H4(l) + O2(g) ➝ N2(g) + H2O(g) A sample of 1.50 g N2H4 is burned in a bomb calorimeter, raising the temperature of the calorimeter by 5.27˚C. The heat capacity (Cbomb) of the calorimeter is 5.510 kJ/˚C. Calculate the heat released in the combustion of 1 mole of hydrazine. First, calculate the energy absorbed by the bomb calorimeter using the following equation:

Qbomb = (5.510 kJ/˚C) (5.27 °C)



Qbomb = 29.0 kJ

Qbomb is positive, indicating that the bomb absorbs energy from the reaction. If the calorimeter were perfect, the heat released by the combustion reaction would be −Qbomb = −29.0 kJ. Because 1.50 g N2H4 was used, the heat released per mole of hydrazine is calculated as follows:

Qbomb = Cbomb∆T

TEKS



∆H = (−29.0 kJ)

32.0 g N2H4 g/mol 1.50 g N2H4

= −619 kJ/mol N2H4

End-of-Course Assessment Review

1. Calculate  In a constant-pressure calorimeter, the temperature of 60.0 g of water increases by 4.50°C. What amount of heat is transferred to the water? A 3.20 D C 1.13 kJ B 64.6 D

D 7.88 kJ

2. Calculate  A calorimeter contains 240 g of an unidentified fluid at 20.5°C. An immersed heating element transfers heat to the fluid at a rate of 120 J/s for 10.0 seconds, and the temperature rises to 24.2 °C. What is the specific heat capacity of the fluid? A 0.14 J/g•°C B 0.21 J/g•°C C 0.74 J/g•°C D 1.4 J/g•°C 3. Calculate  A calorimetric measurement of the “cold pack” reaction is carried out by dissolving 5.0 g of ammonium nitrate (NH4NO3) in 100.0 g of water. A temperature drop of 4.0 °C is observed. What is the molar heat of solution of ammonium nitrate? A −334.4 kJ/g B −26.8 kJ/mol C 26.8 kJ/mol D 334.4 kJ/g 4. Infer  Why is a simple calorimeter made of two foam coffee cups not appropriate for reactions that release gases? TEKS 11E • Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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TEKS

Alpha, Beta, and Gamma 12A Radiation REVIEW

TEKS 12A

Describe the characteristics of alpha, beta, and gamma radiation. TEKS_TXT

What is radiation?

Vocabulary nucear radiation radioactive decay alpha particle beta particle gamma ray

The particles that make up an atom are held together by a balance of attractive and repulsive forces. On the one hand, the protons within a nucleus all repel each other. On the other hand, the so-called “strong nuclear force” provides an attraction between each particle in the nucleus and all the other particles in the same nucleus. If the attractive and repulsive forces are unbalanced, the nucleus of the atom is unstable. This instability can cause a rearrangement of particles within the atom, and the nucleus can break apart. The particles and energy emitted by an unstable nucleus is called nuclear radiation. The breakdown of a nucleus by the emission of radiation is radioactive decay. An isotope involved in a decay reaction is usually identified by the notation A ZX, where X is the particle’s chemical symbol, A is the mass number (sum of the numbers of protons and neutrons), and Z is the atomic number (number of protons). Three main types of nuclear radiation are alpha radiation, beta radiation, and gamma radiation.

What are the characteristics of alpha radiation?

Figure 1 Alpha Decay Alpha particle 4 He 2

238 92 U

234 90 Th

An alpha particle is a helium nucleus emitted by a radioactive source. Another name for alpha particles is alpha radiation. The radioactive decay that results in alpha radiation is often called alpha decay. Each alpha particle consists of two protons and two neutrons and has a double positive charge. Because an alpha particle consists of 2 protons and 2 neutrons, Z = 2 and A = 4. Thus the notation for an alpha particle is 42 He, where He is the chemical symbol for helium. An alpha particle is also sometimes represented by the Greek letter alpha, α. Figure 1 shows an example of alpha decay in a uranium isotope with Z = 92 and A = 238. Note the following: • When the nucleus emits alpha radiation, its atomic number decreases by 2, so Z becomes 90. Because the atomic number has changed, the remaining nucleus is no longer a uranium nucleus, but a thorium nucleus. • The mass number decreases by 4, so A becomes 234. • Since the number of protons has decreased, the overall charge of the nucleus has decreased by 2.

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What are the characteristics of beta radiation? An electron resulting from the changing of a neutron into a proton is called a beta particle. The neutron breaks apart a proton, which remains in the nucleus, and a fast-moving electron, which is released. 1 0



n   ➝ 11 p   +   –10 e

neutron

proton electron (beta particle)

A beta particle is identified by the Greek letter beta, β, or the notation –10 e. Electrons have mass, but the superscript 0 indicates an insignificant mass compared to the mass of a proton. The subscript indicates a -1 charge. Figure 2 shows an example of beta decay. Note the following points in the figure:

Figure 2 Beta Decay Beta particle 0e –1 14 C 6 14 N 7

• The mass number A does not change. • The number of protons increases by 1. Thus, Z increases by 1, and the atom’s identity changes from carbon to nitrogen. A related radioactive decay process that is sometimes classified as beta plus radiation is positron emission. In this process, a proton in an unstable nucleus changes to a neutron and releases a positively charged electron. A positron is represented by the notation +10 e. 1 1



p   ➝ 10 n   +   +10 e

proton

neutron

positron

What are the characteristics of gamma radiation? A third type of radioactive decay is called gamma radiation. A high-energy photon emitted by a radioisotope is called a gamma ray. The high-energy photons are a form of electromagnetic radiation. Nuclei often emit gamma rays along with alpha or beta particles. In nuclear equations, the symbol for a gamma ray is the Greek letter gamma, γ. A gamma ray has no electrical charge and no mass. Neither the atomic number nor the mass number of an atom changes during the emission of gamma radiation. Gamma radiation differs from other types of electromagnetic radiation, such as visible light and radio waves, because it has much higher energy and because it is emitted by a nucleus.

How do the energy, mass, and charge of alpha, beta, and gamma radiation compare?

Study Tip Remember that the types of radiation from greatest to least mass are alpha, beta, and gamma. Alpha particles are much more massive than beta particles, and photons have no mass at all.

Energy, mass, and charge affect the ability of radiation to penetrate matter. Alpha radiation has high energy, but because of its high mass and electrical charge, alpha particles do not travel very far and are not very penetrating. A sheet of paper will block alpha particles. A beta particle has less electrical charge than an alpha particle and much less mass than an alpha particle. Thus, beta particles are more penetrating than alpha particles. Aluminum foil or a thin piece of wood can block beta particles. Gamma rays have no mass or electrical charge and therefore are the most penetrating, and can be very dangerous. Lead or thick concrete is needed to block gamma rays.

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Figure 3 shows some common forms of radiation, as well as their mass, charge, symbol, and penetrating powers. Figure 3

Common Forms of Radiation Particle

Mass

Charge

Alpha

4 amu

+2

Beta

0 amu

Positron Gamma

TEKS

Symbol

Penetrating Power

He or α

low

–1

0 –1

e or β

moderate

0 amu

+1

0 +1

e

moderate

0 amu

none

4 2

γ

high

End-of-Course Assessment Review

1. Describe  Which of the following characteristics is consistent with alpha radiation? A +1 charge B massless C low energy D easily blocked 2. Describe  What happens to a thorium-230 nucleus when it experiences radioactive decay and emits alpha radiation? A The mass number will increase by 2 B The overall energy will remain unchanged. C The mass number will decrease by 4. D The atomic number will remain unchanged. 3. Infer  A nucleus of 199  Pt experiences radioactive decay and becomes 78  lr. Which type of radiation must have been emitted? a nucleus of 199 77 Explain. 4. Sequence  List alpha, beta, and gamma radiation from lowest to highest penetrating powers. 5. Classify  For alpha radiation, beta radiation, and gamma radiation, tell whether each increases atomic number, decreases atomic number, or has no effect on it.

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TEKS

Describing Radioactive Decay 12B Using Nuclear Equations REVIEW

Rea di ness

TEKS 12B

Describe radioactive decay process in terms of balanced nuclear equations.

What are balanced nuclear equations? Balanced nuclear equations can be used to describe radioactive decay. Each isotope involved in a nuclear reaction is represented in the equation by a symbol similar to A ZX. The symbol A is the mass number. The symbol Z is the atomic number. Because atomic number is the number of protons in the nucleus, Z also indicates the charge of the nucleus. An example of an isotope’s symbol is shown in Figure 1. Figure 1 Mass number 238 92

Particle symbol

U

Atomic number

For particles other than isotopes, such as electrons, positrons, and photons, Z is replaced by the charge of the particle. The symbol for an electron is shown in Figure 2. Figure 2

Mass Charge

0 –1

e

Particle symbol

Because of the law of conservation of mass, the mass before and after a nuclear reaction must be the same. Likewise, charge remains the same before and after a reaction. If there was a net charge before a reaction, there will be the same net charge after the reaction. If there was no net charge before a reaction, there will be no net charge after the reaction.

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How can you describe radioactive decay processes with balanced nuclear equations? Each type of radioactive decay has a different nuclear equation. Alpha Decay  During alpha decay, an unstable nucleus releases an alpha particle, 42 He. Because the alpha particle has a mass number of 4, the isotope produced during the reaction has a mass number that is 4 less than that of the original isotope. The value of A therefore decreases by 4. Because the alpha particle has an atomic number of 2, the isotope produced during the reaction has an atomic number that is 2 less than that of the original isotope. The value of Z therefore decreases by 2. A 4 For example, in the equation for alpha decay: 226 88Ra → 2He + Z X , how could you figure out what element X is?

The equation for the superscripts is: 226 = 4 + A, so A = 222 The equation for the subscripts is: 88 = 2 + Z, so Z = 86 Since Z = 86, the new element must be radon, Rn. 226 4 The balanced equation is : 226 88Ra → 2He + 86Rn

Study Tip Sometimes the notation 0  β is used for a beta –1 particle instead of –10 e.

Beta Decay  During beta decay, an unstable nucleus releases a beta particle, –10 e. Because its mass number is 0, the isotope produced during the reaction has a mass number that is the same as that of the original isotope. Because the beta particle’s charge is –1, the isotope produced during the reaction has a nuclear charge that is 1 more than that of the original isotope. This means its atomic number, Z, is 1 more than the atomic number of the original isotope. A 0 In the equation for beta decay: 89 36Kr → –1e + Z X, how could you solve for X?

The equation for the superscripts is : 89 = 0 + A , so A = 89 The equation for the subscripts is: 36 = –1 + Z, so Z = 37 Since Z = 37, the new element must be rubidium, Rb. 89 0 The balanced equation: 89 36Kr → –1e + 37Rb

Gamma Decay  During gamma decay, an unstable nucleus releases a photon, 00γ. Because both the mass number and the photon charge are 0, photons do not change either the mass number or the atomic number of the original isotope. For example: 99 Te 43

→ 00γ + 99 Te 43

Photons are almost always emitted during both alpha decay and beta decay. They are usually left out of the nuclear equations for these processes, however, because they do not affect the calculation of A or Z. The mass number and the charge must balance in equations describing nuclear reactions. This means the sum of the superscripts (the masses) and the sum of the subscripts (the charges) must be the same on both sides of the equation.

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These steps summarize how to write a balanced nuclear equation for radioactive decay: • Write as much of the nuclear equation as you know. • Represent any unknown particle with the symbol A ZX. • Write an equation relating the superscripts. Solve this equation for A. • Write an equation relating the subscripts. Solve this equation for Z. • Substitute the values of A and Z into the symbol A ZX. • Determine the identity of the new isotope from its atomic number, Z. Check a periodic table to see which element has that atomic number. Substitute the symbol of the new isotope for X in the symbol. • Double-check that the equation is balanced by counting superscripts and subscripts again.

TEKS

End-of-Course Assessment Review

1. Apply  Which equation correctly shows the alpha decay of neptunium-239? 239 4 243 A 93Np → 2He + 95Am 239

0

239

B 93Np → –1 e + 92U 239

4

235

C 93Np → 2He + 91Pa 239

0

239

D 93Np → –1 e + 94Pu 2. Predict  Which isotope should you use to complete the following nuclear equation? 208TI 81

→ 00γ +

A 208 TI 79 B 208 TI 84

C 204 TI 79 D 208 TI 81

3. Construct  Write a balanced nuclear equation for the beta decay of bismuth-210. 4. Evaluate  A student wrote this nuclear equation for the radioactive decay of thorium: 230 Th → 42He + 226 Th. Is the equation correct? If so, 90 88 explain how you know. If not, describe what is wrong.

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TEKS REVIEW

12C Fission and Fusion Reactions TEKS 12C

Compare fission and fusion reactions. TEKS_TXT

Vocabulary fission fusion

Figure 1 Uranium Fission Reaction

What is a fission reaction? When the nuclei of certain isotopes are bombarded with neutrons, they undergo fission, the breaking apart of a nucleus into smaller fragments. Fission can occur if an atom’s nucleus is unstable. Figure 1 shows how uranium-235 breaks into two fragments of roughly the same size when struck by a neutron. Neutron

91 36 Kr

Krypton-91 1 3 0n

Energy

235 92 U

Uranium-235 (fissionable)

236 92 U

Uranium-236 (very unstable)

142 56 Ba

Barium-142

The energy released by the fission of a single nucleus can be millions of times greater than the energy released by the combining of atoms during a chemical reaction. The mass of any atom’s nucleus is always less than the sum of the masses of the protons and neutrons of which it is made. This missing mass, called the mass defect, takes the form of energy stored in the atom. This energy is released during fission. The nuclear equation below describes the reaction in Figure 1: 235U 92

+ 10n → 91 Kr + 142 Ba+ 10n+ 10n+ 10n 36 56

U) and In this reaction, a neutron (10 n) strikes a uranium isotope (235 92

produces a krypton isotope (91 Kr) and a barium isotope (142 Ba). Three 36 56

neutrons are also released during the reaction.

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Other isotopes could have been produced in the reaction as long as the mass numbers (shown by the superscripts) and the atomic numbers (shown by the subscripts) are the same before and after the reaction. The ejection of neutrons as a result of the fission reaction is important. These neutrons can, in turn, strike more nuclei, causing them to undergo fission. The result is a chain reaction. Chain reactions release a tremendous amount of energy in a very short time. Fission must be controlled when it is used for the production of electricity at nuclear power plants. The neutrons can be slowed down so that they are more likely to cause more fission reactions, rather than passing straight through a nucleus without causing a reaction. At the same time, materials must be used to absorb neutrons, because if all of them caused a new fission reaction, the process would get out of control and possibly result in a dangerous situation.

What is a fusion reaction? The combining of nuclei to form a nucleus with a greater mass is called fusion. The fusion of nuclei releases a tremendous amount of energy. The reason for this is that mass can change to energy in nuclear reactions, because the mass of a nucleus is less than the sum of the masses of the protons and neutrons of which it is made. When nuclei combine by fusion to form a larger nucleus, the mass defect of the nuclei is released in the form of energy. Fusion can release much more energy than fission, but fusion is not used for the production of electricity at power plants. An enormous amount of energy is required to start the reaction. Thus the material used for fusion must be at such extreme temperatures that it cannot easily be contained. Such extreme temperatures exist, however, in stars, and fusion is responsible for the energy released by the sun. The equation below describes a fusion reaction. In a series of steps, four hydrogen nuclei (11 H) combine to produce two helium nuclei (42 He) and two positrons ( +10 e): 411H → 42He + 2+10e + energy

Study Tip If you have trouble remembering which process is fission and which is fusion, notice that fission is spelled with two of the letter “s” and produces two nuclei. Fusion has one “s” and produces one nucleus.

Each hydrogen nucleus consists of one proton. The sum of the masses of these protons is greater than the mass of the helium nucleus. The mass defect changes to energy that is released as kinetic energy of the particles produced during the reaction.

How do fission and fusion compare? Fission and fusion are both nuclear reactions. However, as you can see in Figure 2 on the following page, they have significant differences.

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Fission vs. Fusion Heavy nucleus breaks into lighter nuclei.

Light nuclei combine to form a heavier nucleus.

Can emit neutrons that can cause other fission reactions

Products do not result in a chain reaction.

Both release tremendous amounts of energy, but fusion releases more energy than fission does. Much higher energy is needed to initiate fusion than fission. • Radioactive byproducts

• No radioactive byproducts

• Uses uranium, a nonrenewable resource

• Uses easily available hydrogen isotopes • Extreme temperatures required make it unusable for production of electricity.

Figure 2

TEKS

End-of-Course Assessment Review

1. Identify  Which of the following is produced during both fission and fusion reactions? A heavier nuclei B enormous energy C hydrogen isotopes D radioactive particles 2. Compare  Which of the following enable a chain reaction in fission but not in fusion? A free neutrons B beta particles C helium nuclei D hydrogen isotopes 3. Compare  Which characteristic pertains to fusion reactions but not fission reactions? A It involves changes of mass to energy. B It often starts with a uranium isotope. C Extremely high temperatures are required to initiate it. D It is used to produce electricity at nuclear power plants. 4. Infer  How is it that fusion occurs so easily in the sun, when it is so hard to cause it on Earth? 5. Evaluate  Current nuclear power plant designs include the use of control rods, which absorb neutrons and prevent the fission reaction from proceeding in an uncontrolled manner. Are neutron-absorbing control rods likely to be a part of the design for a fusion reactor? Explain.

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TEKS REVIEW

Answers to End-of-Course Assessment Reviews TEKS 1C

Sample answers have been provided to open-ended questions. Some questions have many correct answers. If you are unsure about the accuracy of your answer, review it with your teacher.

1. D  A glass beaker is reusable. Since it contained a mixture of hydrochloric acid and water, it can easily be cleaned by thoroughly rinsing with water before returning it to the stockroom to be used again. 2. D  The recommended disposal method for a substance is provided on its Material Safety Data Sheet (MSDS).

TEKS 1A

1. B  Isopropyl alcohol is a flammable liquid. 2. Fire caused by electrical equipment is an example of a Class C fire. Any extinguisher rated for that kind of fire, either an ABC or BC type extinguisher, can be used. 3. When a significant amount of a corrosive chemical is splashed on clothing, there is risk that skin may be exposed to the chemical. The appropriate response in this circumstance is to immediately use a safety shower to wash the contaminant from the clothing, remove the affected articles of clothing, and seek medical attention. 4. His attitude could endanger his eyes. Safety goggles do much more than protect eyes from chemical splashes. They also protect the eye from exposure to powdered chemicals, explosions, and flying debris.

TEKS 1B

1. D  Radioactivity from a radioactive material can cause damage to cells. 2. C  Corrosive substances damage human tissues on contact. 3. The second student is describing the appropriate safe response to the spill. The first student’s plan will likely injure that student and those who help with the clean-up. In addition, the corrosives in the paper towels might react with something in the trash. The instructor should be notified because he or she has responsibility for the lab and more experience and knowledge about how to handle such a situation. 4. Diethyl ether should not be stored in a commercial refrigerator because flammable liquids produce combustible vapors that can easily ignite; it should be stored only in a flammables cabinet. 5. Sample answer: If chemicals are stored alphabetically there is the risk that corrosives and flammables could be stored next to each other. This could lead to a dangerous, unintended reaction.

3. One should never flush a chemical substance (liquid or solid) down the drain without first knowing the approved disposal methods for that chemical. Many substances must not be flushed down the drain because they can contaminate a community’s water supply. In addition, some chemicals react dangerously with water. 4. Evan’s actions were incorrect. Evan should have first checked with the teacher to confirm that the compound should be recycled. If so, it should have been placed in a designated container for that compound. 5. Chemicals that may not be dangerous alone can be dangerous when combined with certain other substances. If a single waste container is used for all chemicals, there is the possibility of a hazardous reaction.

TEKS 2A 1. D  The answer to this question involves people’s judgment and opinions, and is not testable or falsifiable. 2. C  Technology is an application of science. 3. The methods of science have led to many accomplishments. Although commercial space travel to other planets may not occur in the immediate future, it could eventually be possible. As new information is learned, previous scientific beliefs and explanations might be revised or discarded. The new information and ideas could be applied to develop new technology. 4. Answers will vary, but you should note a way of looking at the world that is not science, such as astrology. Answers should imply that science is the study of physical phenomena and the process by which people attempt to understand those phenomena by constructing testable explanations and predictions about them and gathering data that support or falsify those expectations.

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Answers to End-of-Course Assessment Reviews continued TEKS 2B

5. The friend’s description involves a law. A scientific law describes a relationship or pattern in nature, but it does not provide an explanation for the pattern. The description of gravity is not a useful hypothesis because it applies to a broad set of events. A testable hypothesis on gravity would involve a specific set of circumstances.

1. C  An unsupported hypothesis is not very useful for making predictions or understanding observations, so it should be modified and retested. 2. C  As new data and evidence become available, a hypothesis may need to be revised or even rejected altogether. 3. A hypothesis that has durable explanatory power has been tested by multiple scientists over a variety of conditions, and is useful in making predictions under those conditions.

6. A new technology may help researchers identify evidence that may or may not be consistent with a particular theory. If the evidence does not support the theory, the theory may be revised or rejected.

4. Sample answers: The flashlight does not work because the batteries are dead; the flashlight does not work because the bulb is burned out; the flashlight does not work because the wiring inside has shorted. Each statement qualifies as a hypothesis because each is a tentative explanation that is capable of being tested.

TEKS 2D 1. B  This hypothesis could be tested by observing the rusting of iron pieces of different sizes. 2. B  This is the definition of a hypothesis. 3. B  This is the definition of a theory.

5. A hypothesis that is not supported by experimentation is never a failure. Valuable knowledge is gained. For example, if your hypothesis was that the bulb is burned out, and you tried several new bulbs and the flashlight still does not work, your hypothesis would not be supported. However, this information is valuable because it can lead you in another direction. Your next step might be to formulate another hypothesis about why the flashlight is not working and to test the new hypothesis.

4. Sample answers: His idea likely does not explain observations made in many situations. It also cannot yet explain a broad range of observations made in those situations. Theories are also well supported by evidence from many experiments; his idea has not yet been supported by other experiments.

TEKS 2C

2. C  Hypotheses need to be clearly stated so they can be tested.

1. A  A theory is a well-established, highly-reliable explanation of an event. 2. C  Scientists demand that theories be supported by evidence. When evidence amasses against a theory, the theory is revised or replaced. 3. It cannot be a new theory. A scientific theory requires multiple independent researchers—not a single researcher; it also must be highly reliable—a single investigation cannot demonstrate reliability as an explanatory/predictability tool; it also must be well-established—a single investigation cannot establish a theory.

TEKS 2E 1. A  Pipettes are capable of measuring small quantities such as milliliters with precision.

3. C  This hypothesis has very specific independent and dependent variables and therefore is testable. 4. Answers will vary, but should include: safety goggles–to protect eyes; burette—to drop the 0.5 mL of phenolphthalein; burette stand and clamp—to hold burette; beaker or Erlenmeyer flask in 100 mL increments—to hold 500 mL of unknown solution; unknown solution—500 mL to test for pH; phenolphthalein—to titrate using burette

4. A scientific theory involves an explanation; it is not a factual statement that can be proven true or false. A theory is powerful because it has undergone an immense amount of testing over time, and it has been applied to make valid predictions in many cases.

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3. 4.3 years × 365.25 days/year × 24 hours/day × 60 minutes/hour × 60 seconds/minute = 135697680 seconds = 1.35697680 × 108 seconds (assuming one leap year in 4.3 years)

5. Sample answer: Hypothesis: Peppers on plants that are in direct sunlight are larger than peppers that are not in direct sunlight; Independent variable: amount of sunlight; Dependent variable: length of peppers; Equipment: centimeter ruler; identical garden plots (one in shade, one in sun); pepper plants; Experiment: Choose two plants from the same kind of pepper. Place one plant in the shade and one plant in the sun. Give all equal amounts of water. Record the length of the peppers in centimeters over a period of several weeks. Compare the results to evaluate the validity of the hypothesis.

TEKS 2H 1. A  Values greater than the average are balanced by values less than the average. 2. C  Reviewing these details could help explain the unusual result of the eighth trial. 3. The line of best fit would be high for 9 months and then the line would drop for 3 months. This would continue in a cyclical pattern.

TEKS 2F 1. A  The measurement of 102˚C is closest to the accepted value of the boiling point of water, 100˚C, when compared to the other measurements. 2. C  Because Sarah’s three trials are closer together, her measurements are more precise. Sarah’s average = 7.98 g/cm3; Felipe’s average = 7.85 g/cm3. Because Felipe’s average is closer to the accepted value of 7.86 g/cm3, his average is more accurate.

TEKS 2I

3. Graduated cylinder b would allow for a more precise and accurate measurement. Since the measurements on graduated cylinder b are in smaller increments, a more precise measurement of 2.25 mL could be identified. And, using graduated cylinder b would most likely result in a measurement closer to 2.25 mL than using graduated cylinder a. Therefore, the measurement would also be more accurate because it would be closer to the actual value of 2.25 mL.

4. D  Circle graphs are ideal for showing percentages of a whole.

4. If the maximum weight of the train is precisely 1686.5 kg, but the sign has been rounded up to 1700 kg, it is possible that someone might think that a weight of greater than 1685.5 but less than 1700 kg is safe. In this way, the lack of precision in the sign’s value may contribute to a lack of accuracy in properly identifying the safe limits of the ride and ultimately in a dangerous situation.

3. Sample answer: The slogan suggests that the detergent will not allow red clothing dye to stain white clothes. It can be tested by taking photographs of red and white clothing before and after they are washed with the detergent. The experiment should involve several trials under different conditions, such as wash water temperature, presence of fabric softener or other additives, and type of washing machine used. Under all tested conditions, the amount of detergent used and articles of clothing should be kept constant.

1. C  The abstract is a kind of summary that introduces a journal article. 2. A  Direct communication with an audience is the key advantage of an oral report. 3. C  This process, called peer review, helps ensure that the reports in science journals are accurate.

TEKS 3A 1. C  Empirical evidence comes from the results of science experiments. 2. A  The more variables that affect an event, the more difficult it becomes for scientists to evaluate the role these variables play in the outcome.

TEKS 2G 1. A  Multiplying by this conversion factor will cancel out the unit of liters and leave the unit of milliliters. 2. D  The three zeros before the last digit are significant. This measurement is expressed with five significant figures.

TEKS 3B 1. B  Scientists use journal articles to report the results and conclusions of their investigations, and they include data that support their conclusions.

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Answers to End-of-Course Assessment Reviews continued 2. D  The law states that scientific information must be accurate. But marketing materials can present information in misleading ways. 3. People apply the scientific information they learn to make choices about diet, safety, lifestyle, and purchasing products. Because of the importance of these choices, people should carefully evaluate the source of the information for its reliability, and then apply critical thinking skills to draw proper conclusions.

TEKS 3C 1. A  These data are objective, and will give you a clear idea of an aspect of the service’s performance that will be important to you. 2. B  This statement has data that can be scientifically tested. 3. Companies are in business to sell products and services. While laws prevent them from making inaccurate statements, these laws do not prevent them from making statements that could mislead people who are not thinking critically. 4. The data in the table seem to support the claim that VitaPro is a more nutritious cereal than the three other brands. Sample additional questions: How similar to VitaPro were Brands A, B, and C? Was the consumer testing agency that performed the tests connected to VitaPro in any way? What other nutrients does VitaPro contain?

TEKS 3D 1. B  Chemists study matter and changes to matter. 2. A  Though it saved many human lives, DDT caused unexpected harm; plastics have brought many unexpected benefits to society as well as drawbacks such as pollution. 3. Sample answer: Scientific research results in both benefits and drawbacks. For example, computers and the Internet help some people perform their jobs more efficiently and conveniently but take away other jobs, such as travel agents.

3. Chemistry is the study of matter and the changes to matter, and all careers in science involve matter in some way. Experts in life science know how to classify the compounds in living things and describe the changes they undergo. Experts in physics must understand chemistry in order to understand how changes in matter affect changes in energy and other physical properties. Experts in earth science know the chemical composition of Earth’s rocks, soil, water, and air, and they know how these materials undergo physical and chemical changes. Space scientists apply chemical knowledge to help astronauts survive in the harsh atmosphere of space and to develop materials that are suitable for space travel. 4. Answers as to career type will vary, but you should accurately note the type of chemistry involved in your career choice.

TEKS 3F 1. C  Heisenberg’s uncertainty principle, which Schrödinger helped develop, explains that the momentum and position of an electron can never be measured with exact certainty. 2. A  Mendeleev created the periodic table and predicted properties of elements. 3. Answers will vary but should state a discovery, indicate what research has been done to back up your choice, and provide reasons for the discovery you chose.

TEKS 4A 1. A  This describes dissolving, which is a physical change. 2. C  Only water exists before and after the change. Therefore, no new substances were formed. 3. In the first equation, sodium chloride is dissolved in water. This is a physical change because the sodium and chloride ions do not react with other ions. In the second equation, the water is evaporated and sodium chloride reforms. This is also a physical change because the ions do not react to form new substances.

TEKS 3E 1. B  Biochemists study the chemistry of all living things, including plants, animals, and humans. 2. A  Developing a new dish involves managing the physical and chemical changes to food. Chemistry involves the study of these changes. Copyright ©˙Pearson Education, Inc., or its affiliates. All Rights Reserved.

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4. Sample answer: (1) During chemical changes, new substances are formed. During physical changes, the substance(s) that exist before the change still exist after the change, and no additional substances are formed. (2) A chemical change results in a change of both chemical and physical properties of substances. A physical change results in a change in physical properties only.

TEKS 4B 1. D  Density is an intensive property. 2. B  The student cannot determine the identity of the substance because the measurement of one extensive property is not enough to identify a substance. 3. C  Because density is an intensive property, it will allow you to identify the substance the object is made from. 4. No; your classmate is incorrect. While mass and volume considered together can give you density, an intensive property that can give you the identity of a substance, neither on its own can give you the identity of a substance, so neither on its own is an intensive property. 5. Oxidation is an example of an intensive property because it is dependent upon the identity, and not the amount, of the substance. 6. The intensive properties of boiling point and density can be used to determine the identity of the unknown substance. The boiling point of the unknown substance is 82.5°C. Of the substances in the table, only isopropyl alcohol has this boiling point. The identity can be confirmed by calculating the density of the unknown and comparing it to that of isopropyl alcohol. The density of the unknown substance is determined by dividing the mass by the volume, with a result of 0.785 g/mL. The density of isopropyl alcohol is determined by dividing the mass by the volume as shown in the table, with a result of 0.786. Accounting for rounding, these two densities are the same; Substance X is isopropyl alcohol.

TEKS 4C 1. D  Particles in a solid vibrate in place. Particles in a liquid slide past one another.

3. B  When the gas is compressed, its particles get closer together. 4. Both the inflatable packing material and foam packaging peanuts contain air, a gas. This packaging is placed between the contents of the package and the outside of the box. If the contents receive an impact from an object outside of the box, the air within the packaging material would compress, but the items in the package would be protected from the impact.

TEKS 4D 1. B  Both elements and compounds are uniform in color and texture. 2. B  The new material contained atoms of different elements bonded together, which is the definition of a compound. 3. D  Vinegar contains two or more pure substances mixed together. 4. Blood has a uniform texture and color when viewed with unaided eyes, but because it does not appear uniform under a microscope, it is not evenly mixed at the atomic level. Therefore, it is a heterogeneous mixture. 5. The material is a compound. The substance contains more than one element, so it is not an element. It does not break down using physical processes, which means it is not a mixture. Compounds are made from more than one element but break down only during chemical reactions.

TEKS 5A 1. C  Elements in a group will display similar chemical and physical properties. 2. C  He decided that elements should be grouped based on properties even if that meant placing them out of atomic mass order. 3. D  Krypton is in group 18. Elements in the same group share similar properties. 4. The elements were organized first by placing them in order of increasing atomic mass and then by placing elements with similar properties in the same column. Mendeleev left spaces in the columns (groups) for elements with similar properties that had not yet been discovered.

2. C  When a gas is transferred from a smaller container to a larger container, its volume increases. Copyright ©˙Pearson Education, Inc., or its affiliates. All Rights Reserved.

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Answers to End-of-Course Assessment Reviews continued TEKS 5B

2. C  The observations would be similar to those for gold. A few alpha particles would be strongly deflected by the platinum foil.

1. D  Both noble gases and halogens contain gaseous elements.

3. All of Dalton’s postulates except #1 and #2 have contributed to the modern model of the atom. In addition, the modern model of the atom includes a small, dense nucleus (Rutherford) that contains protons (Rutherford) and neutrons (Rutherford). The electrons (Thomson) are located in an electron cloud surrounding the nucleus (after Bohr) and emit energy when they change energy levels (Bohr).

2. A  Alkali metals exhibit luster and malleability, they react vigorously with water, and they are found only as compounds in nature. 3. C  Alkali metals are more reactive than alkaline earth metals and all groups/families are more reactive than noble gases. 4. Elements in Group 1 tend to lose one electron when they react with other elements. Elements in Group 17 tend to gain one electron when they react with other elements. It is likely that an electron from the Group 1 element would be transferred to the Group 17 element.

TEKS 6B 1. A  If FM radio waves have higher frequencies than AM radio waves, they must have shorter wavelengths. (3.00 m is shorter than 375 m.)

5. Noble gases were once thought to be nonreactive because compounds of noble gases do not exist in nature. Since then, compounds of noble gases have been synthesized in the laboratory and scientists now know that noble gases are not inert.

2. D  These are correctly listed from shortest (gamma) to longest (radio). 3. C  30.0 × 108 m = 10(3.00 × 108 m), the distance waves travel in 10 seconds. 4. Ultraviolet rays have higher frequencies than light rays and thus have more energy than light rays. Waves with higher energy are more likely to be able to damage human tissue than waves with lower energy.

TEKS 5C 1. D  Ionic radii of cations decrease from left to right across a period. 2. D  The number of protons increases from left to right across a period. The more protons in the nucleus, the greater their pull on the electrons, so the smaller the atomic radius is.

TEKS 6C

3. A  An increase in principle quantum number indicates that electrons spend a greater time farther from the nucleus and are therefore easier to remove.

2. A 489 nm = 4.89 × 10−7 m;

1. D  λ =

4. D  Li and Be are in the same period, and Li appears to the left of Be. Atomic radii decrease from left to right, so Li is larger than Be.

ν= 3.

3.00 × 108 m/s 5 × 1014 Hz 3.00 × 108 m/s 4.89 ×

10−7

m

C  E = (6.626 × 10−34 = 6.626 × 10−20 J

= 6 × 10−7 m

= 6.13 × 1014 s–1

J•s)(1.0 × 1014 s−1)

4. No, because this wavelength of light has a frequency of 1.00 × 1015 Hz, which is outside the visible range. c ν=

5. Protons in the nucleus exert a force of attraction on the electrons that surround the nucleus. The more protons there are in the nucleus, the greater the force of attraction and the greater the energy required to remove one of the electrons.

λ

TEKS 6A

c



ν=



= 1.00 × 1015 s–1

3.0 × 10−7 m

=

3.0 × 108 m/s 3.0 × 10−7 m

1. A  The rays were repelled by negatively charged plates and attracted to positively charged plates. He expected the alpha particles to pass straight through the gold foil, and most did.

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5. ν =

ν=

c

4. In order for the n = 2 level to fill, the n = 1 level must also be filled. The n = 1 level contains one s sublevel with two electrons. The n = 2 level contains one s sublevel with two electrons and one p sublevel with six electrons. It is possible that other sublevels are filled also, but these are the minimum that would be filled if an atom has a completely filled n = 2 level.

λ

3.00 × 108 m/s

4.90 × 10−7 m = 6.1 × 1014 s–1 E = hν E = (6.626 × 10–34 J•s)(6.1 × 1014 s–1) E = 4.04 × 10–19 J

5. The symbol [Ne] stands for 1s2 2s2 2p6. The shorthand electron configuration for magnesium is [Ne]3s2.

TEKS 6D 1. B  The unknown atom has 29 protons and 36 neutrons (65 mass number − 29 protons) 2. A  (1.00783 amu) (0.999885) + (2.01410 amu) (0.000115) = 1.0079 amu

TEKS 7A

3. (27.97693 amu) (0.9293) + (28.97649 amu) (0.0468) + (29.97377 amu) (0.0309) = 25.80 amu + 1.36 amu + 0.926 amu = 28.09 amu 4. The most abundant isotope of fluorine is fluorine-19. Its mass number is 19; it has 9 protons and 10 neutrons. It is the most abundant isotope because the average atomic mass is very close to 19. If there were other, more abundant isotopes, the average atomic mass would be farther from a whole number.

1. A  The anion is sulfate, so the name of the acid has the suffix -ic. 2. C  The subscript 3 indicates that there are 3 nitrate ions for every aluminum ion. 3. B  The name gives the first clue that it is covalent, since it uses the prefixes for the amounts of each atom. The second clue is that iodine and fluorine are both nonmetals. 4. CaBr2

5. The atomic mass and mass number describe a single atom with a specific number of protons and neutrons. Each element listed on the periodic table is composed of several isotopes, each with a different number of protons and neutrons.

TEKS 6E

TEKS 7B 1. D  According to the chemical name, there are two nitrogen atoms and one oxygen atom for each molecule of the substance. 2. B  According to the name, the acid is a binary acid. And sulfide has a –2 charge, so the hydrogen ion should have the subscript two. 3. B  Phosphoric acid consists of phosphate (PO43–).

1. D  Si has four valence electrons represented by a dot on each of four sides of the element symbol.

TEKS 7C

2. Yes, this is possible. For example, electrons start to fill the 4s atomic orbital in principal energy level n = 4 before filling the 3d atomic orbital in principal energy level n = 3.

1. D  This is a covalently bonded compound with hydrogen atoms surrounding carbon atoms.

3. 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p1. The outermost electrons in gallium are in principal energy level 4, which contains three electrons—two are s electrons and one is a p electron. The Lewis dot diagram for gallium is; Ga

2. B  The product is beryllium chloride (BeCl2), an ionic compound. 3.

[ ] O

Mg2+ C

2−

O

O

4.

− H+[ C N ]

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Answers to End-of-Course Assessment Reviews continued TEKS 7D 1. D  This describes ductility, as opposed to malleability. 2. B  Metallic bonding explains how thermal energy is transferred through a metal. 3. Sample answer: Thermal conductivity and electrical conductivity are similar in that they are both a result of metallic bonds. Thermal energy is conducted through the vibrations of the lattice of cations, and electrical energy is conducted through the movement of the delocalized electrons. 4. Sample answer: A “sea” of delocalized electrons holding together a lattice of cations causes metals to hold together and yet be reshaped enough to be drawn into wires or hammered flat. The delocalized electrons carry electricity and the lattice of ions easily transfers thermal energy. 5. Sample answer: Ionic and metallic bonds are similar in that they both involve metal cations. In a metallic bond, however, there is no corresponding anion. The cations are attracted to the delocalized “sea” of electrons.

TEKS 7E 1. D  A tetrahedral arrangement has four equal faces. 2. C  Carbon has four electron pairs, giving methane a tetrahedral arrangement. 3. The molecule is carbon dioxide; it has two double bonds, no lone pairs of electrons, and is linear like this molecule. 4.

5. The chemist is incorrect. The representative particle of water is a molecule. Each water molecule has two hydrogen atoms and one oxygen atoms. As a result, the 9 moles of water consists of 18 moles of hydrogen atoms and 9 moles of oxygen atoms.

TEKS 8B 1. B  There are 2 moles of Al and 6 moles of Cl in 1 mole of Al2Cl6. 2. B  To calculate the number of atoms, divide the mass by the molar mass and then multiply by Avogadro’s number. 3. A  The number of molecules is the mass divided by the molar mass times Avogadro’s number. 4. Sample answer: (1) Calculate the molar mass of each ion in the formula unit. (2) Multiply each molar mass by the number of ions that make up the formula unit, and then add them to get the formula unit’s molar mass. (3) Divide the sample’s mass by the molar mass. (4) Multiply by Avogadro’s number. 5. Sample answer: The student neglected to multiply by the conversion factor of 3 Cl– ions per formula unit of AlCl3. The correct answer is 9.49 × 1022 Cl– ions.

TEKS 8C 1. A  This is the percent composition of hydrogen, calculated by dividing the mass of hydrogen by the total mass.

F Cl

4. Yes. The type of representative particle is different (a molecule for sucrose and an atom for magnesium), but a mole describes a number of particles. One mole is 6.02 × 1023 regardless of the type of particle.

C Cl

Cl

The arrangement is tetrahedral.

TEKS 8A 1. B  A mole is equivalent to Avogadro’s number, or 6.02 × 1023. 2. B  Dividing by Avogadro’s number converts number of molecules to moles.

2. D  The moles of each element were used as subscripts then divided by the lowest number of moles. The result was multiplied by 2. 3. D  The subscripts of the empirical formula were multiplied by the quotient of the molar mass and the empirical mass formula. 4.

molar mass of compound

=

molar mass of CH



42.081 g/mol

=3

12.001 g/mol + 2(1.008 g/mol)

The molecular formula is C1×3H2×3 , or C3H6.

3. C  The number of molecules is the number of moles multiplied by Avogadro’s number.

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TEKS 8D

TEKS 9B

1. B  Because mass is conserved, the atoms that make up the products must be the same as the atoms that make up the reactants.

1. A  The amount of oxygen that will react will be one half the amount of the hydrogen.

2. A  There are 4 Al atoms and 6 O atoms before and after the reaction. 3. D  There are 2 K atoms, 2 N atoms, 9 O atoms, 2 H atoms, and 1 C atom before and after the reaction. 4. D  A coefficient of 6 would result in 6 H atoms and 6 Cl atoms on both sides of the equation.

2. B  Two moles of acetylene react to produce 4 moles of carbon dioxide. This is simplified to a ratio of 1 to 2. 3. C  First convert mass of octane to moles; then use the volume ratio and Avogadro’s law to find the volume of carbon dioxide. 4. 2CO + O2→ 2CO2 14 g CO ×

5. 3Fe + 4H2O → 4H2 + Fe3O4 6. Sample answer: According to the law of conservation of mass, mass can be neither created nor destroyed. This means that to balance a chemical equation, you must use coefficients before the formulas so that the atoms of each element that are present on the left side of a chemical equation equal the atoms of each element that are present on the right side.

1. B  Charles’ law describes the relationship between the temperature and volume of an ideal gas. 2. B  The amount of propane is calculated by multiplying pressure and volume and dividing by the product of R and temperature (in K). 3. The new volume is calculated by multiplying P1 × V1 × T2 and dividing it by the product of T1 and P2 = 4.67 liters.

1 mol CO



V=

1. A  The percent yield is 96.9 percent.

TEKS 9A

1 mol CO2

P = 2.0 atm V=? n = 0.50 mol R = 0.08205 L•atm/mol•K T = 290 K

=

3. Yes, magnesium (Mg) is the limiting reagent. Multiplying the given 2.6 moles Mg by 1 mole of N2 per 3 moles of Mg, you find that about 0.87 moles of N2 is needed. Because 4.5 moles of N2 is available, N2 is the excess reagent, and Mg is the limiting reagent.

28 g CO

×



TEKS 8E 2. A  This answer correctly divides the mass by a molar mass of 119.002 g for KBr, divides by 2 moles KBr to 1 mol Br2, and multiplies by 159.808 g per mole for Br2.

1 mol CO

= 0.50 mol CO2

n×R×T P 0.50 mol × 0.08205 (L • atm/mol • K) × 290 K

= 5.9 L

2.0 atm

TEKS 9C 1. A  Because the particles are so small, their volumes can be considered negligible. 2. D  As gas expands, the kinetic energy is distributed over a larger volume, so the temperature decreases. 3. Answers will vary, but should summarize the three postulates of kinetic molecular theory: that gas molecules have insignificant volumes, that they have rapid and random motion, and that their collisions are elastic. 4. On cold days, the particles that make up the gases in the tires move with less average kinetic energy, so the pressure inside the tires is less than on hot days, when the particles move with more average kinetic energy. 5. Nearly all the volume of a gas is made up of empty space, according to the first postulate of kinetic molecular theory. So the fact that the uranium hexafluoride molecules are larger than the hydrogen molecules will have only a very small impact on the overall volume of the gas. In other words, both will contain about the same volume of empty space—nearly all of the volume.

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Answers to End-of-Course Assessment Reviews continued TEKS 10A

2. D  To find the volume of undiluted solution, you multiply the concentration and volume of the new solution and divide by the concentration of the old solution.

1. A  Because of the weakness of the forces, they cannot absorb as much energy. 2. B  Surface tension is related to cohesion.

3. C  To calculate the new diluted volume, you multiply the old concentration and the old volume and divide by the new concentration.

3. D  Water’s polar nature and ability to form hydrogen bonds means it will dissolve both polar and ionic compounds.

4. The 5.0 g of NaCl is equivalent to 0.086 mol of NaCl. In order to make a 0.05M solution, the volume would need to be 1.7 L.

4. It is a strongly polar molecule; it exhibits cohesion (surface tension); it exhibits capillary action in hydrophilic substances; its liquid phase is denser than its solid phase; it has a high heat capacity; it is an excellent solvent.

5. Write the equation 0.20M × 0.500 L = 0.15M × V2; so V2 = 670 mL. The stock solution does not contain enough potassium iodide for you to prepare any volume greater than 670 mL.

TEKS 10B

TEKS 10E

1. A  A precipitate will settle to the bottom because it is not soluble.

1. B  Ability to conduct electricity describes an electrolyte.

2. D  While many sulfates are soluble, barium is an exception.

2. A  At 50°C the solution would be saturated with about 75g KNO3. At 40°C, about 58 g would dissolve. So about 17 g of solute would precipitate.

3. A  Sulfides of Group 2 ions, such as calcium, are only slightly soluble.

3. The statement is incorrect. While ionic compounds dissolve in water and are electrolytes, some covalent substances also dissolve in water, and covalent substances are not electrolytes.

4. NiCl2(aq) + Na2CO3(s) → NiCO3(s) + 2NaCl(aq)

TEKS 10C 1. D  Molarity is calculated by dividing the moles of a solute by liters of the solution.

4. It may be supersaturated, and the solid substance may be precipitate, or the substance may just not be very soluble in the solvent under the given conditions.

2. C  To find the molarity, you divide the moles of solute (0.56 mol) by the volume in liters (0.510 L).

TEKS 10F

3. D  The mass can be calculated by multiplying molarity by volume to calculate moles needed, and then multiplying by molar mass to find number of grams.

1. A  Crushing the solute will increase its surface area, speeding dissolution.

The first solution is more concentrated; it has a concentration of 2.0M, found by dividing number of moles by volume of solution. The second solution has a concentration of 1.8M, found by dividing mass by the molar mass to find the number of moles (0.53 mol), and dividing the result by the volume of the solution in liters (0.30 L).

TEKS 10D 1. C  The diluted solution will be less concentrated, but when the new concentration (M2) and the new volume (V2) are multiplied, they equal the old concentration (M1) and the old volume (V1).

2. C  The solubility for each of the gases is nearly zero at 100°C. 3. Sample answer: While an increase in temperature will speed the rate of dissolution for the syrup, it will also decrease the solubility of the carbon dioxide gas used to carbonate the soft drink, causing the carbon dioxide to leave the solution. The resulting soda may therefore taste “flat.”

TEKS 10G 1. C  This is the definition of a Brønsted-Lowry base. 2. B  Nitrous acid (HNO3) donates a hydrogen ion, so it is an Arrhenius acid.

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3. D  A salt (MgCl2) and water are products of a neutralization reaction.

2. A  Solution A has a pOH of 7 since it has a pH of 7, and Solution B has a pOH of 9. The lower the pOH, the higher the hydroxide concentration.

4. A  The salt would be formed by an acid of the anion, I−, and a base of the cation, Ba2+.

3. A  The pH will decrease by two units, from 8 to 6.

5. The products are sodium cyanide and water. HCN + NaOH → NaCN + H2O

4. 5.3 pOH = log (1.9 × 10−9) = 8.7   pH = 14 − 8.7 = 5.3

TEKS 10H

TEKS 10J

1. B  Electrons are transferred in all oxidation– reduction reactions, and not in all acid-base reactions.

1. D  Complete dissociation of hydrogen ions represents a strong acid. 2. D  Urea has the smallest Kb of all the bases listed; therefore it would be the weakest base.

2. A  According to the equation, the product, Al(OH)3, is a precipitate, because it is noted as a solid (s).

3. A  The best conductor of electricity has the highest dissociation constant (K) and therefore dissociates the most in water.

3. C  Light was produced, which indicates that it was an oxidation-reduction reaction. A chemical equation summarizing the reaction would also show that an electron was transferred from the sodium to the chlorine in the reaction.

4. In general, strong acids and bases dissociate completely or nearly completely. Weak acids and bases dissociate very little.

4. You could determine if it was an acid-base reaction by writing the chemical equation for the reaction and identifying whether a proton was transferred. Mg(s) + 2HCl(aq) → H2(g) + MgCl2(aq). A proton was not transferred from the hydrochloric acid to the magnesium, so the reaction was not an acid-base reaction. You could determine if it was a precipitation reaction if the products make the solution appear cloudy or if a substance appears in the solution. You could also determine whether any of the products in the equation are solids. This reaction is not a precipitation reaction. You could determine if it was an oxidation-reduction reaction by checking if the oxidation numbers of any of the elements changed from the reactant to the product. In this reaction, magnesium begins with an oxidation number of zero and ends with an oxidation number of +2. It is an oxidationreduction reaction.

5. HNO2(aq)  →  H+(aq)  +  NO2–(aq)

TEKS 11A 1. B  Energy is the ability to do work. 2. C  The formation of covalent bonds occurs in order to minimize potential energy of the bonding electrons. 3. Heat: from the monitor; light: from the monitor; sound: from the speakers; mechanical motion: movement of keys 4. Chemical potential energy is stored within the chemical bonds inside the log. As it is ignited, a chemical reaction occurs, which releases heat, light, and sound. The reactants, which have higher chemical potential energy, are converted into products that have lower potential chemical energy.

5. Yes. An acid-base reaction produces water and a salt. A precipitation reaction involves the combining of two aqueous solutions resulting in a solid precipitate. If an acid and base are aqueous solutions, and the salt they form is insoluble in water, the reaction would be both an acid-base reaction and a precipitation reaction.

TEKS 10I 1. D  pH is a negative logarithmic scale of hydrogen ion concentration.

5. Sample answer: Gravitational potential energy: it’s on a lab table, not the floor; chemical potential energy: the atoms could be ready to react; thermal energy: the atoms in the piece of iron have some motion.

TEKS 11B 1. B  As electrical energy is converted to light energy, more energy is lost as thermal energy in the incandescent bulb.

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Answers to End-of-Course Assessment Reviews continued 2. A  Energy is converted to thermal energy when the reaction takes place. The car’s engine temperature will get so high that you will burn your fingers if you touch it.

4. Because the coffee-cup calorimeter is not airtight, the volume of the gases could change, and some of the energy would be involved in this change, instead of just the change in the temperature of the fluid in the calorimeter.

3. Heat is flowing toward the ice pop. The process is endothermic because the ice pop “system” is gaining heat from its surroundings, the warm deck.

TEKS 12A 1. D  Alpha radiation has such high mass and charge that it can be blocked by a piece of paper.

TEKS 11C

2. C  Because the alpha particle has two protons and two neutrons, the total decrease in mass number for the isotope is 4.

1. B  The temperature is rising because this reaction results in the release of energy; it must be exothermic because more energy is released in the bond formation of the products than was required in the breaking of bonds in the reactants.

3. Positron; the decrease in Z from 78 to 77 indicates a decrease of one proton. This type of decrease occurs with positron emission.

2. D  (2 × 435) + (498) – (4 × 464)

4. gamma, beta, alpha

3. B  ½ × [435 + 158 – (–546)]

5. Increases atomic number: beta radiation; decreases atomic number: alpha radiation, positron emission; no effect on atomic number: gamma radiation

4. Endothermic reactions absorb energy from their environment because they require more energy to break bonds in reactants than is released by the formation of new bonds.

TEKS 12B 1. C  This shows alpha particle with the mass number decreasing by 4 and the atomic number decreasing by 2.

TEKS 11D 1. A  Q = (50.0 g)(4.18 J/g•˚C)(5.0˚C) = 1045 J 2. B  First calculate the heat to melt the ice, which is 12.0 kJ. Then add the heat to raise the temperature to 25°C, which is 3.77 kJ. The sum is 15.77 kJ.

2. D  Gamma radiation does not change the mass number or the atomic number of an isotope. 3. 210 Bi → 0 e + 210 Po –1 84 83 4. The equation is incorrect. The atomic number (Z) changes from 90 to 88, but the student did not change the identity of the isotope to radium (Ra).

3. B  Use the equation Q = mcΔT, and solve for c. 4. The energy absorbed during the melting process must disrupt the intermolecular attractive forces only enough to allow the molecules to slide past each other in relative motion. During the boiling process, sufficient energy must be supplied to completely separate the molecules from each other.

TEKS 12C

5. Answers will vary but may include water’s function as temperature regulation of the human body, coolant in automobiles, a friendly habitat for ectothermic animals that do not regulate their own temperature, and hot water bags for treatment of muscle pain.

TEKS 11E 1. C  Remember that heat can be expressed in different units, including joules and kilojoules. 2. D  The specific heat is c = Qsubstance/(msubstanceΔT). 3. C

1. B  Enormous amounts of energy are released by both combining and breaking apart nuclei. 2. A  The free neutrons produced by fission can cause other nuclei to break apart, leading to a chain reaction. 3. C  Extremely high temperatures are required to initiate fusion but not fission. 4. Sample answer: Fusion takes place in the sun because of the high temperatures there. There is no place on Earth with temperatures this hot, and construction of vessels that could contain such temperatures is not currently possible for large-scale use. 5. No, because fusion reactions do not depend upon neutrons to get them started, and they do not take place in a chain reaction. Therefore, the use of control rods is not likely because absorbing neutrons will not slow a fusion reaction.

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TEKS REVIEW

Glossary accuracy: the closeness of a measurement to the true value of what is being measured, 25 acid: a compound that contains one or more hydrogen atoms and donates a proton in a reaction, 93 acid-base reaction: a type of reaction that occurs between an acid and a base in which a proton is transferred, 151 actual yield: the amount of product that actually forms when a reaction is carried out, 120 adhesion: the attraction between molecules of different substances, 132 alkali metal: any metal in Group 1 (or 1A) of the periodic table, 70 alkaline earth metal: any metal in Group 2 (or 2A) of the periodic table, 71 alloy: a mixture composed of two or more elements, at least one of which is a metal, 101 alpha particle: a helium nucleus emitted from certain radioactive nuclei; represented by the Greek letter alpha, α, 175 anion: any atom or group of atoms with a negative charge, 73 atom: the smallest particle of an element that retains its identity in a chemical reaction, 76 atomic mass: the weighted average of the masses of the isotopes of an element, 85 atomic mass unit (amu): a unit of mass equal to onetwelfth the mass of a carbon-12 atom, 85 atomic number: number of protons in an atom, 68, 85 atomic orbital: a region in an atom in which an electron of a particular amount of energy is most likely to be located, 88 atomic radius: one-half the distance between the nuclei of two atoms of the same element when the atoms are joined, 73 Avogadro’s law: law that states that equal volumes of gases at the same temperature and pressure contain equal numbers of particles, 122 Avogadro’s number: the number of representative particles contained in one mole of a substance; equal to 6.02 × 1023 particles, 106 base: a compound that produces hydroxide ions when dissolved in solution, 93 beta particle: an electron resulting from the changing of a neutron into a proton; identified by the Greek letter beta, β, 176

bond dissociation energy: the energy required to break the bond between two covalently bonded atoms; usually expressed in kJ per mol of substance, 167 Boyle’s law: law that states for a given mass of gas at constant temperature, the volume of the gas varies inversely with pressure, 121 calorimeter: an instrument that measures the absorption or release of heat in chemical or physical processes, 172 calorimetry: the measurement of the heat flow that occurs during a chemical process or reaction, 172 cation: any atom or group of atoms with a positive charge, 73 Charles’ law: law that states that the volume of a fixed mass of gas is directly proportional to its Kelvin temperature if the pressure is kept constant, 122 chemical change: a reaction that occurs when one or more substances change into one or more different substances, 56 chemical energy: a form of potential energy stored in the chemical bonds of a substance, 161 chemical property: the ability of a substance to change into a different substance, 55 cohesion: the attraction between molecules of the same substance, 131 compound: matter that is made up of two or more different elements chemically combined in a fixed proportion, 65 corrosive substance: a compound that is highly reactive and will cause serious damage to living tissue, 5 covalent bond: a bond formed by the sharing of electrons between atoms, 92 Dalton’s law of partial pressure: law that states that at constant volume and temperature, the total pressure exerted by a mixture of gases is equal to the sum of the partial pressures of the component gases, 122 decontamination: the removal of hazardous compounds, 7 dimensional analysis: a method of problem solving that uses the units that are part of a measurement to help solve the problem, 28 dissolution: the process by which one substance (the solute) dissolves in another (the solvent); also called solvation, 145 ductile: property of a metal that enables it to be drawn into a wire, 101

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Glossary continued

electrolyte: a compound that conducts an electric current when it is in an aqueous or molten state, 142 electron dot formula: a notation that consists of a chemical symbol surrounded by one to eight dots representing valence electrons, 97 electron: a subatomic, negatively charged particle, 76 electronegativity: the ability of an atom to attract electrons when the atom is in a compound, 74 element: matter that is made up of only one type of atom, 65 empirical formula: a formula with the lowest whole number ratio of elements in a compound, 113 endothermic process: a process that absorbs heat from the surroundings, 164 energy: the capacity for doing work or producing heat, 160 excess reagent: a reagent present in a quantity that is more than sufficient to react with a limiting reagent; any reactant that remains after the limiting reagent is used up in a chemical reaction, 119 exothermic process: a process that releases heat to its surroundings, 164 extensive property: a property that depends on the amount or size of a substance, 59 fission: splitting of a nucleus into smaller fragments accompanied by the release of neutrons and a large amount of energy, 181 flammable substance: a substance that gives off combustible vapors that can easily ignite, 5 frequency: the number of wave cycles that pass a given point per unit of time, 79 fusion: the process of combining nuclei to produce a nucleus of greater mass, 182 gamma ray: a high-energy photon emitted by a radioisotope; represented by the Greek letter gamma, γ, 176 gas: a state of matter without definite shape or volume, 62 halogen: a nonmetal found in Group 17 (or 7A) of the periodic table, 71 heat: energy that transfers from one object to another because of a temperature difference between the objects, 164 heat capacity: the amount of heat needed to increase the temperature of an object 1°C, 169

hertz (Hz): the SI unit of frequency, equal to one cycle per second, 79 heterogeneous: a mixture in which the substances are not uniformly mixed, 66 homogeneous: a substance in which the particles are uniformly mixed, 65 hypothesis: a tentative statement or explanation for an observation in nature, 13 ideal gas law: the relationship PV = nRT, which describes the behavior of an ideal gas, 123 inference: a logical interpretation based on prior knowledge and experiences, 31 intensive property: a property that is not dependent on the amount of a substance, 58 ion: an atom or group of atoms that has a positive or negative charge, 73 ionization energy: the minimum amount of energy needed to remove an electron from an atom or ion, 74 isotopes: atoms of the same element that have the same atomic number but different atomic masses due to a different number of neutrons, 85 joule: the SI unit of energy; 4.184 J = 1 calorie, 160 kinetic energy: the energy an object has because of its motion, 161 kinetic molecular theory: a theory that states that all matter consists of small particles that are in constant motion, 127 law of conservation of energy: law that states that in any chemical or physical process, energy is neither created nor destroyed, 163 law of conservation of mass: law that states that in any physical change or chemical reaction, mass is conserved; it is neither created nor destroyed, 115 limiting reagent: the reactant that determines the amount of each product that can be formed in the reaction, 119 liquid: a state of matter having definite volume but no definite shape (takes the shape of its container), 62 malleable: the property of metals that allows them to be hammered into thin sheets, 101 mass number: the sum of an atom’s protons and neutrons, 85 Material Safety Data Sheets (MSDS): data forms that contain detailed information on the properties, hazards, and health effects of chemical substances, 4

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metallic bond: the force of attraction that holds metals together; consists of the attraction of valence electrons for the positively charged metal ions, 100 mixture: a physical blend of two or more substances that are not chemically combined, 65 molar heat of condensation: the amount of heat released by one mole of a vapor as it condenses to a liquid at a constant temperature, 170 molar heat of fusion: the amount of heat needed to change one mole of a substance from a solid to a liquid at the melting point, 169 molar heat of solidification: the amount of heat lost by one mole of a liquid as it solidifies at a constant temperature, 170 molar heat of solution: the change in heat caused by the dissolution of one mole of a substance, 170 molar heat of vaporization: the amount of heat needed to change one mole of a substance from a liquid to a vapor at the boiling point, 169 molarity (M ): the number of moles of solute dissolved in one liter of solution, 136 molar mass: the mass of one mole of any substance, 109 mole: the amount of a substance that contains 6.02 × 1023 representative particles of that substance, 106 neutron: the uncharged particle in the nucleus of an atom, 77 noble gas: an element that is in Group 18 (or 8A) of the periodic table, 71 nonelectrolyte: a compound that does not conduct an electric current in aqueous solution or in the molten state, 143 nuclear radiation: the particles and energy emitted by an unstable nucleus, 175 nucleus: small, dense, central portion of an atom; composed of protons and neutrons, 77 oxidation-reduction reaction: a type of reaction that involves the transfer of electrons between reactants; also called a redox reaction, 152 percent composition: the relative amount of mass of each element in a compound, 112 percent yield: the ratio of the actual yield to the theoretical yield for a chemical reaction expressed as a percentage, 120 periodic table: an arrangement of elements in which the elements are separated into groups based on a set of repeating properties, 67

pH: a number used to denote the hydrogen-ion concentration, or acidity, of a solution; it is the negative logarithm of the hydrogen-ion concentration of a solution, 154 physical change: a change that does not alter the chemical identity of a substance, 56 physical property: a quality or condition of a substance that can be observed or measured without changing the substance’s composition, 55 Planck’s constant: a number used to calculate the radiant energy (E) absorbed or emitted by a body based on the frequency of radiation, 80 polyatomic ion: a tightly-bound group of atoms that has a positive or negative charge and behaves as a unit, 94 potential energy: the ability of a body to do work because of its position, composition, or configuration, 161 precipitation reaction: type of reaction that occurs when two aqueous solutions react and produce a solid precipitate, 151 precision: the closeness, or reproducibility, of a set of measurements taken under the same conditions, 25 proton: the positively charged particle in the nucleus of an atom, 77 pure substance: matter that has a definite composition; either an element or a compound, 64 radioactive decay: the breakdown of a nucleus by the emission of radiation, 175 radioactive substance: a substance that spontaneously emits ionizing radiation, 6 representative particle: the smallest unit into which a substance can be broken down without a change in composition, usually atoms, molecules, or ions, 106 salt: the product (other than water) of a neutralization reaction, 149 saturated solution: a solution containing the maximum amount of solute for a given amount of solvent at a constant temperature and pressure, 143 scientific notation: an expression of numbers in the form of m × 10n, where m is equal to, or greater than 1 and less than 10, and n is the integer, 28 scientific theory: a well-established, highly reliable explanation of a natural or physical phenomenon, 16 significant figure: all the digits that can be known precisely in a measurement, plus a last estimated digit, 29 solid: a state of matter having a definite shape and volume, 61

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Glossary continued solubility: a measure of how much solute will dissolve in a certain amount of solvent at a specific temperature and pressure, 133, 145 solute: the dissolved particles in a solution, 136, 145 solution: a homogeneous mixture; consists of solutes dissolved in a solvent, 65 solvent: the dissolving medium in a solution, 136 145 specific heat: the amount of heat needed to increase the temperature of 1 g of a substance by 1°C, 169 spectator ion: an ion that is not directly involved in a chemical reaction, 135 standard temperature and pressure (STP): the conditions under which the volume of a gas is usually measured; standard temperature is 0°C and standard pressure is 1 atmosphere (atm) or 101.3 kPa, 121 stoichiometry: the calculation of quantities of substances involved in chemical equations, 118 strong acid: an acid that is completely (or almost completely) ionized in aqueous solution, 157 strong base: a base that completely dissociates into metal ions and hydroxide ions in aqueous solution, 157 supersaturated solution: a solution that contains more solute than it could theoretically hold at a given temperature, 143 surface tension: an inward force that tends to minimize the surface area of a liquid, 131 surroundings: everything in the universe outside of the system, 163 system: a part of the universe on which you focus your attention, 163

theoretical yield: the maximum amount of a product that could be formed from a given amount of reactant, 120 thermal energy: the total potential and kinetic energy of the molecules of a substance caused by their arrangement and constant motion, 162 transition metal: one of the Group B elements in which the highest occupied s sublevel and a nearby d sublevel generally contain electrons, 71 unsaturated solution: a solution that contains less solute than a saturated solution at a given temperature and pressure, 143 valence electron: an electron in the highest occupied energy level of an atom, 90, 97 Valence Shell Electron Pair Repulsion (VSEPR) theory: theory that states that the repulsion between electron pairs causes molecular shapes to adjust so that valence electron pairs are as far apart as possible, 103 wavelength: the distance between two equivalent points of a wave, 79 weak acid: an acid that is only slightly ionized in aqueous solution, 157 weak base: a base that reacts with water to form the hydroxide ion and the conjugate acid of the base, 157

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TEKS REVIEW

Test-Taking Tips Multiple-choice questions make up the entire endof-course assessment tests. Therefore, it is important to be good at deciphering this type of question. We have included a variety of strategies that will help you. For any one question, not all strategies will need to be used. In a multiple-choice question, several possible answers are given to you, and you need to figure out which one of those answers is best. The first part of the question is called the stem. The stem can be a question or an incomplete statement. Read the stem carefully before you look at the answer choices. The answer choices are indicated by letters, A, B, C, and D. One answer choice is correct. The other answer choices, called distractors, are incorrect.

Anticipating the Answer A useful strategy for answering multiple-choice questions is to provide your own response before you look at the answer choices. After you provide your own response, compare it with the answer choices. You will then be able to quickly identify the correct choice. This technique is especially useful for questions that test vocabulary.

Sample Question Gravy is a suspension of flour in water. What type of mixture is gravy? A  heterogeneous mixture B  homogeneous mixture

Using the Process of Elimination Suppose you are not sure of the correct answer. You may be able to determine the correct answer through the process of elimination. Look at each answer choice and eliminate the choices that are least likely to be correct. Reduce you answer choices to two, if possible. Then reread the question and select the better of the two choices.

Sample Question Selenium is a nonmetal that is a solid at room temperature. Which element has chemical properties most like those of selenium (Se), an element in Group 16 of the periodic table? A  krypton, a noble gas B  gallium, located three boxes to the left of selenium in the periodic table C  sulfur, located one box above selenium in the periodic table D sodium, located on the far left side of the periodic table The correct answer is C. You may not remember anything about selenium, but you know that Group 16 is close to the far right of the periodic table. Sodium, at the far left of the table, is a metal, so it is not likely to have similar properties as selenium; you can thus eliminate choice D. The noble gases are at the very right of the periodic table, which is Group 18, so you can eliminate Choice A. If you remember that elements in columns, or groups, have similar properties, then you can choose choice C as the correct answer.

C  solution D  emulsion The correct answer is A, which is easy to choose if you know that a suspension is a type of heterogeneous mixture. You might also remember that a solution is another term for a homogeneous mixture. Because the question has only one correct answer choices, B, C, and D must both be incorrect.

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Test-Taking Tips continued

Identifying What Is Being Asked

Identifying an Event’s Place in a Sequence

In some questions, you must distinguish between background information and what is being asked. In the sample question, the first paragraph presents background information. The question being asked is about the qualities of gases.

Some questions ask you to identify the correct place of an event in a sequence of events. For example, the question may ask you which event comes first or last, which event precedes another event, or which event happens at the same time as another event.

Sample Question Two steel plates can be joined by a process called arc welding. In this process, tiny droplets of molten metal are deposited on the joint between the plates. The droplets then cool and harden, joining the plates together. Many metals will react with oxygen when they are exposed to the high temperatures of arc welding. Therefore, the droplets of molten metal are often shielded from oxygen in the surrounding air by a layer of argon gas. Why would argon gas be a more suitable shield for arc welding than hydrogen gas? A  Unlike hydrogen, argon is a nonmetal. B  Unlike hydrogen, argon is an inert gas. C  Argon has a larger atomic radius than hydrogen. D  Argon has a greater atomic mass than hydrogen. The correct answer is B. Choice A is incorrect because both argon and hydrogen are nonmetals. Hydrogen has only a single electron and has the smallest atomic radius of any element, but this does not explain why argon makes a better shield for arc welding; therefore, choice C is incorrect. Choice D is incorrect because although argon does have a greater atomic mass (39.948) than hydrogen (1.008), this fact does not explain why argon makes a better shield for arc welding.

Before you answer this kind of question, try to recall as much of the entire process as you can. Then look at all the answer choices. Begin by eliminating those that you know to be incorrect. For example, if you are asked to identify the last event in a series, eliminate the answer choices that indicate steps that occur early in the process. Some multiple choice questions that involve sequence will not ask you about where an event falls in a sequence, but rather ask about the correct sequence of events or steps in a process.

Sample Question A student measures the volume of a small steel marble using the water displacement method. Which of the following steps must be performed twice? A  Add a steel marble to the graduated cylinder. B  Measure the level of the water in the graduated cylinder. C  Add water to the graduated cylinder. D  Remove the steel marble from the graduated cylinder. The correct answer is B. As you can conclude from the question stem and the answer choices, the water displacement method involves water, the object to be measured such as a steel marble, and a graduated cylinder. Recall how this method works: A water level is measured in the graduated cylinder, the object is added, and the water level is measured again. The rise in the water level equals the volume of the object. If you understand this process, you can eliminate all three of the distracters. Both the water and the marble are added to the cylinder only once.

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Interpreting Line Graphs A line graph enables you to see patterns in data and to make predictions. You must understand what the x and y axes represent and how to interpret the slope of the line. In the line graph below, the x-axis represents temperature and the y-axis represents solubility. This information tells you that the lines relate the temperature of a solution to the solubility of the solute.

A  30°C B  0°C C  50°C D  70°C

KNO3

140

Solubility (g/100 g H2O)

At which temperature do sodium chloride (NaCl) and potassium nitrate (KNO3) have approximately equal solubilities?

The correct answer is A. Find the lines labeled NaCl and KNO3. They intersect at a temperature of about 30°C and a solubility of 40 g/100 g H2O.

Solubility Varies With Temperature 160

Sample Question 2

120

NaNO3

100

KBr

80

NH4Cl

60

NaCl

40 20

Yb2(SO4)3 0

10

20

30

40

50

60

70

80

90

Temperature (ºC)

Sample Question 1 As the temperature of an aqueous solution increases, the solubility of which compound decreases? A  potassium nitrate (KNO3) B  sodium chloride (NaCl) C  potassium bromide (KBr) D  ytterbium sulfate (Yb2(SO4)3) The correct answer is D. As the lines in the graph show, solubility increases with temperature for most compounds. But for ytterbium sulfate (Yb2(SO4)3), the line slopes downward.

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Test-Taking Tips continued

Interpreting Data Tables

Sample Question 2

Data tables are one way in which scientists present the results of experiments. Before answering a question about a data table, look at the table to get a general sense of what kind of information it contains. If the table has a title, it usually indicates what the experiment was about. If a table does not have a title, you can often infer the purpose of the experiment by examining the headings of the columns and/or rows to see what information they provide.

By analyzing and performing calculations on the data in the table to the left, what can be concluded about the reaction? A  the rate of the reaction B  the mechanism of the reaction C  the equilibrium constant D  the solubility product The correct answer is D. Recall that the equilibrium constant is a ratio between the concentrations of products and reactants at the equilibrium position. It is a property of a reversible reaction that does not depend on the initial concentrations. The data table provides the equilibrium concentrations for three trials. Each trial should yield a similar equilibrium constant.

Equilibrium Positions Sample Question 1 2NO N2O4 for the Reaction 2 Nitrogen dioxide to dinitrogen 2) converts Trial Initial (NOInitial Equilibrium Equilibrium tetroxide (N2O4) in a reversible reaction. A scientist [N2O4] [N O ] [NO2] [NO2] places different amounts of these two gases in2 a4 (M) She then (M) measures (M)the changing (M) reaction vessel. concentrations equilibrium0.014 is 1 0.20of the0gases until0.17 reached. The data obtained is shown in the table 0 0.24 0.028 2 0.30 below. 0.20 0.31 0.045 3 0

Sample Question 1 How was Trial 3 different from Trials 1 and 2? A  Equilibrium was reached more quickly. B  Equilibrium was reached more slowly. C  Initially, nitrogen dioxide converted to dinitrogen tetroxide. D  Initially, dinitrogen tetroxide converted to nitrogen dioxide. The correct answer is D. The most noticeable difference in Trial 3 is that the gas initially in the reaction vessel was dinitrogen tetroxide (N2O4), not nitrogen dioxide (NO2). This means that the initial reaction was the conversion of N2O4 to NO2 , not the opposite reaction. Notice that the table presents no information about the time in which equilibrium was reached, so you can eliminate choices A and B.

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Interpreting Diagrams Scientists use diagrams to show the parts of objects or the steps of a process. Begin by looking carefully at the diagram to get an overall sense of what it shows. Look for any labels with lines connected to parts of the diagram. If you don’t immediately recognize what the diagram shows, the labels may give you a clue. After you have examined the diagram, read the question. After you choose an answer, recheck the diagram to make sure that your answer is correct. Sometimes you will be asked to interpret a set of related diagrams. Each part of the diagram could represent a test group of a science experiment or demonstration. Or the diagrams could show the steps of a process, such as a chemical reaction.

The correct answer is A. To identify the correct answer, you need to apply some basic facts about atoms. As choice A and the diagram suggest, protons and neutrons are located in the atomic nucleus. Electrons move in the space around the nucleus. Choices B, C, and D are all features that the diagram does not represent accurately about atoms. Remember that all models are inaccurate in at least some ways.

Sample Question 2 Four identical bottles of carbonated water are each placed in a glass bowl. The conditions of the bowl and the bottle are shown in the set of diagrams below. Cap

To interpret a set of diagrams, make sure that you understand the meaning of arrows and other symbols. You should also compare the different diagrams. Determine the features that are similar and different. Then read the question and answer choices carefully.

Warm water

Ice Bottle 1

Sample Question 1

Bottle 2 No cap

In the diagram below, which feature of protons, neutrons, and electrons is represented most accurately? Bottle 3

C  markings on their surfaces

Warm water Bottle 4

Neutron

Which bottle will contain the least amount of dissolved carbon dioxide after 15 minutes?

Electron

A  Bottle 1

C  Bottle 3

B  Bottle 2

D  Bottle 4

A  position inside or outside of the atomic nucleus B  distance between them

No cap

Ice

Proton

D  relative sizes

Cap

The correct answer is D. Bottle 4 is uncapped, so the carbon dioxide in Bottle 4 is under less pressure than the carbon dioxide in Bottles 1 and 2. Because the solubility of a gas decreases as the pressure decreases, Bottle 4 will contain less carbon dioxide than Bottles 1 and 2. The warm water surrounding Bottle 4 also decreases the solubility of carbon dioxide. Bottle 3 is uncapped like Bottle 4, but is sitting in ice.

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Test-Taking Tips continued

Interpreting a Text Passage Sometimes you will be asked to read a text passage consisting of one or more paragraphs and then answer questions about the text. Always read the text passage entirely before answering any questions. The order of the questions may be different from the order of information in the text passage. For instance, the first question may ask about something found in the last paragraph of the text, and the last question may relate to information in the first sentence of the text. When you’ve chosen an answer, check the text passage to make sure your answer is correct. If you are unsure of an answer to a question, skim the passage again to see whether you can find the information you need. The answer may not be stated directly; you may need to make inferences about what you have read or apply a concept that is not stated in the text. Sample Questions 1 and 2 refer to the following passage. In the 1800s, several scientists attempted to explain the properties of the chemical elements that were known at the time. Several scientists arranged the elements in order of increasing atomic mass, then noticed that the properties of the elements appeared to repeat in a pattern. But the pattern did not hold true for all the known elements. In 1869, Dmitri Mendeleev published his first periodic table of the elements. This table improved on other arrangements of elements in several ways. Mendeleev had the insight to leave gaps in the table for elements that he believed were undiscovered. He also switched the order of some elements to help them fit in the table.

Sample Question 1 As can be inferred from the passage, which discovery helped scientists accept the value of Mendeleev’s periodic table? A  the discovery that atomic number, not atomic mass, caused periodic properties B  the discovery of noble gases, which have nearly identical chemical properties C  the discovery of the elements that filled the gaps in Mendeleev’s table D  the discovery of protons, neutrons, and electrons The correct answer is C. Notice that the passage does not discuss how scientists came to accept Mendeleev’s table. You need to use information in the passage to infer which discovery would be most useful. Choice A might serve to undermine Mendeleev’s work, not support it. Choices B and D are not discussed in the passage. Choice C, the correct answer, relates to the third sentence in the second paragraph of the passage. Mendeleev deliberately left gaps in his table. The discovery of elements that filled those gaps provided evidence that supported the table’s organization.

Sample Question 2 As the passage suggests, Mendeleev’s periodic table is an example of A  a new discovery that overturned older ideas. B  a successful revision of older ideas. C  an accidental discovery based mostly on luck or chance. D  a failure that led to future success. The correct answer is B. The first paragraph describes the work of Mendeleev’s predecessors, who also had noticed that the elements seem to repeat in a pattern. Mendeleev improved on their work by revising it, not overturning it.

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Griddable Items

Sample Question

Some questions involve writing answers on a bubble grid that has a decimal point. Begin by figuring out the correct number (the answer) and where the decimal point goes. Then, write the number in the spaces at the top of the grid, with the decimal point in the correct place. Finally, fill in the correct bubble beneath each numeral.

What is the volume of liquid shown in the graduated cylinder? 4 . 5

4

5

0

0

0

0

0

0

0

1

1

1

1

1

1

1

2

2

2

2

2

2

2

3

3

3

3

3

3

3

4

4

4

4

4

4

4

5

5

5

5

5

5

5

6

6

6

6

6

6

6

7

7

7

7

7

7

7

0

0

0

0

0

0

0

8

8

8

8

8

8

8

1

1

1

1

1

1

1

9

9

9

9

9

9

9

2

2

2

2

2

2

2

3

3

3

3

3

3

3

4

4

4

4

4

4

4

5

5

5

5

5

5

5

6

6

6

6

6

6

6

7

7

7

7

7

7

7

8

8

8

8

8

8

8

9

9

9

9

9

9

9

To solve, look at the bottom of the curve, called the meniscus, shown in the graduated cylinder. It is halfway between 4 and 5 markings on the graduated cylinder, so the volume is 4.5 mL. Enter the symbols 4, decimal point, and 5 into the boxes. Then fill in the appropriate bubbles below the grid.

Here are important points to remember about the bubble grid used on the end-of-course assessment. • The correct answer can be a positive or a negative number. If the answer is a negative number, you must enter a negative sign. Otherwise, the answer will be recognized as a positive number. • If the answer is a decimal number, enter a decimal point in the proper location. • If you are taking the test online, enter numbers into the boxes. If you are taking the printed version of the test, write the numbers in the boxes and then fill in the bubbles. • You do not need to use all of the boxes. You may place an answer in any set of consecutive boxes. • You may fill in extra zeroes either before or after the answer, so long as the extra zeroes do not change the answer. An entry of “20” is the same as “00020”, but not “2000.”

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Test-Taking Tips continued

Interpreting Experiments

Recognizing Units

A problem may include a description of an experiment, an illustration of a lab setup, or a graph or table showing the results of an experiment. You may be asked to answer one or more related questions about the experiment.

All measurements are expressed in certain units. The mass of an object could be 5.13 grams or 20 kilograms, but not 52 liters (a measurement of volume) or 71 joules (a measurement of energy.) Sometimes the distracters will present numerical answers with a variety of units. If so, look at the units carefully. You can eliminate distracters that include incorrect or unreasonable units.

Before trying to answer any of the questions, carefully read or look at the information you have been given and see what you can learn directly or infer from it. Does it tell you the scientist’s purpose in conducting the investigation, or the variables? If not, can you use other information to figure this out?

Sample Question 1 The density of water is 1 g/cm3. What might be the density of an object that floats in water? A  3.7 g/cm3 B  0.69 g/cm3

Sample Question A piston with a freely-movable top holds 2.5 L of helium. A scientist raises the temperature of the piston from 300 K to 350 K, and then measures the change in volume of the helium. What principle of science is the scientist evaluating? A  the ideal gas law B  the octet rule, which explains why helium is nonreactive C  the change in reactivity of helium at different temperatures

C  0.075 kg/cm3 D  0.32 m/s2 The correct answer is B. Less dense objects float on denser objects. Because the number 0.075 is so small, you might be tempted to choose choice C. But the unit of density in this distracter includes kilograms, not grams. The value is equivalent to 75 g/cm3. Note also that choice D is a measurement of acceleration, not density.

D  the conservation of mass The correct answer is A. The experiment is testing the effect of temperature on the volume of helium, and the result can be predicted by the ideal gas law. The scientist is not measuring the mass or reactivity of helium, so the other answer choices can be eliminated.

Sample Question 2 Compound Y is produced at an initial rate of 3.0 × 10−4 mol/L·s. At this rate, how many moles of Compound Y will be produced in a 3.0-L reaction vessel after 20 seconds? A  6.0 × 10−4 mol B  3.0 × 10−4 mol C  1.8 × 10−2 mol D  3.0 × 10−2 mol The correct answer is C. You might not remember solving problems that involve reaction rates, but look at the units in the question stem. The initial rate is expressed in a unit of mol/L·s. If you multiply by both the liters (L) and seconds (s), the unit becomes moles (mol), which is the unit of the answer choices. Multiplying the three quantities stated in the problem will give you the correct answer.

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TEKS REVIEW

STAAR CHEMISTRY Chemistry Reference REFERENCE MATERIALS

Materials STAAR

TM

State of Texas Assessments of Academic Readiness

ATOMIC STRUCTURE STRUCTURE ATOMIC Speed of light = (frequency)(wavelength)

c = fλ

Energy = (Planck’s constant)(frequency) Energy =

E photon = hf

(Planck’s constant)(speed of light)

E photon =

(wavelength)

BEHAVIOROF OFGASES GASES BEHAVIOR Total pressure of a gas =

(

sum of the partial pressures of the component gases

)

PT = P1 + P2 + P3 + . . .

(Pressure)(volume) = (moles)(ideal gas constant)(temperature) (Initial pressure)(initial volume) s )(initial temperature) (Initial moles (Initial pressure)(initial volume) (Initial volume) (Initial temperature) (Initial volume) (Initial moles)

=

=

(final pressure)(final volume)

=

hc λ

(final moles)(final temperature)

= (final pressure)(final volume)

PV = nRT PV 1 1

n1T1

=

PV = P V 1 1

V1

(final vo l ume)

T1

(final temperature)

(final volume)

V1

n1

(final moles)

P2V2

n2T2

2 2

=

=

V2 T2

V2

n2

SOLUTIONS SOLUTIONS Molarity =

moles of solute liter of solution

M =

Ionization constant of water =

(

Volume of solution 1

)(

) (

molarity of solution 1 =

(

hydrogen ion concentration

volume of solution 2

)(

)(

hydroxide ion concentration

molarity of solution 2

pH = −logarithm (hydrogen ion concentration)

)

)

mol L

K w = [H+][OH−] V1M1 = V2M2 pH = −log[H+]

THERMOCHEMISTRY THERMOCHEMISTRY

(

Heat gained or lost = (mass)

(

specific heat

) (

)(

change in temperature

enthalpy enthalpy Enthalpy of − of reactants reaction = of products

)

)

Q = mcp∆T ∆H = ∆H of (products) − ∆H of (reactants)

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STAAR CHEMISTRY REFERENCE MATERIALS OTHERFORMULAS FORMULAS OTHER Density =

mass volume

( (

Percent error = Percent yield =

D =

accepted value − experimental value accepted value actual yield theoretical yield

)

)

m V

(100)

(100)

CONSTANTSAND ANDCONVERSIONS CONVERSIONS CONSTANTS 23

Avogadro’s number = 6.02 × 10

particles per mole −34

h = Planck’s constant = 6.63 × 10

8

c = speed of light = 3.00 × 10 K

J ⋅s

m s −14

w

= ionization constant of water = 1.00 × 10

alpha particle (α) =

4 He 2

beta particle (β) =

0 e −1

( ) mol

2

L

neutron =

1 n 0

standard temperature and pressure (STP) = 0°C and 1 atm 0°C = 273 K volume of ideal gas at STP = 22.4

L mol

3

1 cm = 1 mL = 1 cc 1 atm = 760 mm Hg = 101.3 kPa R = ideal gas constant = 0.0821

L ⋅ atm mol ⋅ K

= 8.31

L ⋅ kPa mol ⋅ K

= 62.4

L ⋅ mm Hg mol ⋅ K

1 calorie (cal) = 4.18 joules (J) 1000 calories (cal) = 1 Calorie (Cal) = 1 kilocalorie (kcal)

RULES FIGURES RULESFOR FORSIGNIFICANT SIGNIFICANT FIGURES 1. Non-zero digits and zeros between non-zero digits are always significant. 2. Leading zeros are not significant. 3. Zeros to the right of all non-zero digits are only significant if a decimal point is shown. 4. For values written in scientific notation, the digits in the coefficient are significant. 5. In a common logarithm, there are as many digits after the decimal point as there are significant figures in the original number.

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STR12_ANC_CHEM_EM_RM.indd 211

Compounds of NH+ and the alkali metal cations Compounds of NH+ and the alkali metal cations Compounds of NH4+ , the alkali metal cations,

SO2− 4

Insoluble compounds contain

CrO2− 4 Cr2O2− 7 OH−

ClO3− ClO2− CrO2− 4 CN− Cr2O2− 7 HCO3− OH− ClO− NO3− NO2− ClO4− MnO− 4 PO3− 4 SO2− 4 SO2− 3

Chlorate

Chlorite

Chromate

Cyanide

Dichromate

Hydrogen carbonate

Hydroxide

Hypochlorite

Nitrate

Nitrite

Perchlorate

Permanganate

Phosphate

Sulfate

Sulfite

None

CN−

S2−

PO3− 4

CO2− 3

I−

Cl−

Br

None

ClO2− ClO3− ClO4− −

Ca2+, Sr

2+

2+

, and Ba2+

, and Ba2+

Compounds of NH4+ , the alkali metal cations,

Ca2+, Sr

4

4

4

Compounds of NH+ and the alkali metal cations

4

Compounds of NH+ and the alkali metal cations

Common exceptions

Compounds of Sr 2+ , Ba2+ , Pb2+ , and Hg2+ 2

Compounds of Ag+ , Pb2+ , and Hg2+ 2

Compounds of Ag+ , Pb2+ , and Hg2+ 2

Compounds of Ag+ , Pb2+ , and Hg2+ 2

None

None

None



ClO

None

NO3−

None

CO2− 3

C2H3O2− , CH COO− 3 NH4+

None

Common exceptions

Carbonate

Soluble compounds contain

NH4+

C2H3O2− , CH3COO−

SOLubILITY COMMON SOLUBILITY OFOf COMMON IONIC COMPOuNdS WATER IONIC COMPOUNDS IN IN WATER

Ammonium

Acetate

POLYATOMIC POLYATOMIC IONS IONS

STAAR CHEMISTRY REfERENCE MATERIALS

Gold

Platinum

Silver

Mercury

Copper

(Hydrogen)

Lead

Tin

Nickel

Cobalt

Iron

Chromium

Zinc

Manganese

Aluminum

Magnesium

Sodium

Calcium

Barium

Potassium

Lithium

Metal

ACTIvITY ACTIVITY SERIES SERIES

Increasing Activity

211

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Increasing Activity

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7

6

5

4

3

2

1

19

104

103

88 (262)

(226)

Radium

(223)

Francium

(267)

Rf

Actinide Series

Lanthanide Series

60

231.036 Protactinium

232.038 Thorium

(227) Actinium

91

Pa

Ac

Th

89

90

140.908

Cerium

144.242

(145)

61

Pm

Hassium

(270)

108

Hs

190.23

Osmium

Os

76

Ruthenium

101.07

44

Ru

Iron

55.845

Fe

26

8

Silicon

28.086

Si

14

Uranium

238.029

U

92

Neptunium

(237)

93

Np

Praseodymium Neodymium Promethium

140.116

138.905

Nd

59

Lanthanum

La

Pr

58

Bohrium

(271) Seaborgium

(272)

107

Bh

(268)

Sg

106

Rhenium

186.207

Re

75

Dubnium

105

183.84 Tungsten

180.948 Tantalum

74

W

Db

(98)

43

Tc

Manganese

54.938

Mn

25

7 7B

Molybdenum Technetium

95.96

42

Mo

Chromium

51.996

Cr

24

6 6B

Ta

73

Niobium

92.906

41

Nb

Vanadium

50.942

V

23

5 5B

Atomic mass

Symbol

Atomic number

Ce

57

Lawrencium Rutherfordium

Lr

Ra

87

Fr

178.49 Hafnium

Lutetium

Barium

Cesium

174.967

137.328

132.905

Hf

Ba

Lu

72

71

Cs

56

55

Zirconium

91.224

40

Zr

Yttrium

88.906

87.62

Strontium

85.468

Rubidium

Y

Sr

39

Rb

37

38

47.867 Titanium

44.956 Scandium

40.078

Calcium

39.098

Ti

22

4 4B

Sc

21

3 3B

Potassium

Ca

20

Sodium

K

24.305

Magnesium

22.990

Mg

12

11

Na

9.012

Beryllium

6.941

4

Be

2 2A

Lithium

Li

3

Hydrogen

1.008

H

1

1 1A

PERIODIC OFOF THE ELEMENTS PERIODICTABLE TABLE THE ELEMENTS

STAAR CHEMISTRY REFERENCE MATERIALS

(281)

Ds

110

Platinum

195.085

Pt

78

Palladium

106.42

46

Pd

Nickel

58.693

Ni

28

10

(280)

111

Rg

Gold

196.967

Au

79

Silver

107.868

47

Ag

Copper

63.546

Cu

29

11 1B

Plutonium

(244)

94

Pu

Samarium

150.36

62

Sm

(247) Curium

(243)

96

Cm Americium

95

Am

157.25 Gadolinium

151.964

64

Gd Europium

63

Eu

Meitnerium Darmstadtium Roentgenium

(276)

Mt

109

Iridium

192.217

Ir

77

Rhodium

102.906

45

Rh

Cobalt

58.933

Co

27

9 8B

Name

26.982

Thallium

204.383

Tl

81

Indium

114.818

49

In

Gallium

69.723

Ga

31

Aluminum

Lead

207.2

Pb

82

Tin

118.711

50

Sn

Germanium

72.64

Ge

32

Silicon

28.086

Si

14

Carbon

12.011

6

C

14 4A

Bismuth

208.980

Bi

83

Antimony

121.760

51

Sb

Arsenic

74.922

As

33

30.974

Phosphorus

P

15

Nitrogen

14.007

7

N

15 5A

Berkelium

(247)

97

Bk

Terbium

158.925

65

Tb

Californium

(251)

98

Cf

Dysprosium

162.500

66

Dy

Einsteinium

(252)

99

Es

Holmium

164.930

67

Ho

Fermium

(257)

100

Fm

Erbium

167.259

68

Er

Mass numbers in parentheses are those of the most stable or most common isotope.

Mercury

200.59

Hg

80

Cadmium

112.412

48

Cd

Zinc

65.38

Zn

30

12 2B

Al

13

Boron

10.812

5

B

13 3A

Mendelevium

(258)

101

Md

Thulium

168.934

69

Tm

Polonium

(209)

Po

84

Tellurium

127.60

52

Te

Selenium

78.96

Se

34

Sulfur

32.066

S

16

Oxygen

15.999

8

O

16 6A

4.003

Radon

(222)

86

Rn

Xenon

131.294

54

Xe

Krypton

83.798

36

Kr

Argon

39.948

18

Ar

Neon

20.180

10

Ne

Helium

Updated Spring 2011

Nobelium

(259)

102

No

Ytterbium

173.055

70

Yb

Astatine

(210)

At

85

Iodine

126.904

I

53

Bromine

79.904

Br

35

Chlorine

35.453

Cl

17

Fluorine

18.998

9

F

17 7A

He

2

18 8A

State of Texas Assessments of Academic Readiness Chemistry Practice Tests

TEKS

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STAAR Chemistry

TEKS

Practice Test A

End-of-Course Assessment Chemistry Practice Test A

1 A student performs the following tests to determine the identity of a mineral sample from its physical and chemical properties. Test 1

Microscopic examination of the mineral to see the shape of its crystals

Test 2

Scratch test of the mineral to determine its hardness

Test 3

Rubbing the mineral on a streak plate to determine the color of its powdered form

Test 4

Dropping dilute hydrochloric acid on the mineral to see if bubbles will form

Which of the tests would involve a chemical change in the mineral? A Tests 2 and 4 B Tests 3 and 4 C Tests 1, 2, and 4 D Test 4 only

2 A chemistry student makes careful observations and measurements of a small sample of matter, and determines the following: Appearance

silver solid

Mass

11.85 g

Density

5.9 g/cm3

Melting point

30ºC

The student determines that the unknown substance is gallium (Ga). Which of the following is an extensive property of the gallium sample? A Silver solid B Mass of 11.85 g C Density of 5.9 g/cm3 D Melting point of 30°C

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STAAR Chemistry

Practice Test A

3 The braking systems in modern cars are designed with components utilizing the compressibility properties of solids, liquids, and gases. The compressibility properties of solids, liquids, and gases are — A different for all three states of matter B the same for solids and gases but not for liquids C the same for gases and liquids but not for solids D the same for solids and liquids but not for gases

4 Students in a laboratory are asked to test whether a given sample of liquid matter is a pure substance or a mixture. At their stations, two lab partners make the statements in the following table. Student 1

If we evaporate some of the sample and a residue remains, that means it is a mixture.

Student 2

Whether the sample is a pure substance or a mixture of two volatile substances, no residue will remain if we evaporate it.

Which of the following is correct? A Statement 1 is true and statement 2 is false. B Statement 1 is false and statement 2 is true. C Both statements 1 and 2 are true. D Both statements 1 and 2 are false.

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STAAR Chemistry

Practice Test A

5 The graph below shows a periodic trend. 300

Period 2

200

Period 1

250

Period 4

Period 3

Rb

K

Period 5

Cs

Na Sc

Li

150

Cd Zn Xe

100

50

Kr

Ar Ne

He

10

0

20

30

40

50

60

Atomic number

From the trends shown, which of the variables below is plotted on the vertical axis of the graph? A Ionic radius B Atomic radius C Electronegativity D First ionization energy

6 Compared to other periodic tables in his time, the table developed by Dmitri Mendeleev — A enabled him to predict chemical and physical properties of yet undiscovered elements B arranged elements in order of increasing atomic number C was the only one that arranged elements into groups with similar chemical and physical properties D had all elements arranged in perfect order of increasing atomic mass

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STAAR Chemistry

Practice Test A

7 The table below gives information about a certain element found in the Periodic Table. Element exists as diatomic molecules

High reactivity

Usually gains one electron to form compounds

Rarely found as a free element

The element is most likely a member of which chemical family? A Alkali metals B Alkaline earth metals C Halogens D Transition metals

8 The following two equations represent changes to water. 1.  H2O(l ) → H2O(s) + Energy 2.  H2O(l ) + Energy → 2H2(g) + O2(g) A physical change is represented by — A both 1 and 2 B neither 1 nor 2 C 2 only D 1 only

9 A transparent liquid has uniform color. Under a microscope, no difference in uniformity is apparent. The liquid passes unchanged through a filter paper. Based on this information it may be concluded that the liquid is — A an element B either a pure substance or a homogeneous mixture C either an element or a heterogeneous mixture D a heterogeneous mixture

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STAAR Chemistry

Practice Test A

10 Which of the properties listed in the table below are not usually identified with transition metals? 1. Conduct an electric current 2. Shiny 3. Liquid at room temperature 4. Form several common oxidation states 5. Highly reactive A 4 and 5 B 2 and 4 C 3 and 5 D 1 and 3

11 Some of the Group 1A elements are listed alphabetically in the table below? 1. Cesium 2. Lithium 3. Potassium 4. Rubidium 5. Sodium Arrange the elements in order of increasing electronegativity. List the numbers of the elements from the table above, with the number of the element with lowest electronegativity on the left and the highest electronegativity on the right. Record your answer and fill in the bubbles on your answer document.

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STAAR Chemistry

Practice Test A

12 A liquid has slightly different properties than a solid because in a solid the particles — A are free to flow past one another B are close together C get bigger when the temperature increases D vibrate in place

13 The figure below illustrates the electromagnetic spectrum. Red light, low energy (7.0 × 10 –7 m = 700 nm)

Violet light, high energy (3.8 × 10 –7 m = 380 nm)

Visible light

Frequency 𝛎 (s −1) 3 × 106

3 × 10 8

Radio waves

10 2

1

Radar

3 × 10 10

3 × 10 12

Microwaves

10 –2

3 × 10 14

Infrared

10 –4

3 × 10 16

Ultraviolet

10 –6

10 –8

3 × 10 18

X-rays

10 –10

3 × 10 20

3 × 10 22

Gamma rays

10 –12

10 –14

Wavelength 𝛌 (m)

Using the figure above, determine which of the following choices records types of electromagnetic waves in order of increasing energy? A Infrared, ultraviolet, gamma, visible B Ultraviolet, visible, microwave, radio C Visible, microwave, x-ray, gamma D Radar, infrared, visible, ultraviolet, x-ray

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STAAR Chemistry

Practice Test A

14 Cell phones operate with a range of frequencies between 500 MHz and 4 GHz. Which of the following could be the wavelength of a cell phone electromagnetic wave? A 1000 km B 5 m C 50 cm D 1 cm

15 Ernest Rutherford devised an experiment in 1911 to probe the structure of the atom. Observations and Possible Conclusions from the Rutherford Gold Foil Experiment 1. The reflected alpha particles bounced off of the positively charged nucleus — which proved the existence of the nucleus. 2. The reflected alpha particles bounced off of the atoms in the gold foil — which proved the existence of atoms. 3. Most of the alpha particles went straight through the gold foil undeflected — which proved that atoms are neutral. 4. Most of the alpha particles went straight through the gold foil undeflected — which proved that the positive charge was spread out evenly through the atom. Which of these conclusions did Rutherford draw? A 1 B 2 C 3 D 4

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STAAR Chemistry

Practice Test A

16 Uranium has several naturally occurring isotopes. Isotope

Relative Abundance

Mass (amu)

U-238

99.3%

238.05

U-235

0.7%

235.04

What is the average atomic mass of uranium to the nearest hundredth of an atomic mass unit? Record your answer and fill in the bubbles on your answer document.

17 The electron configuration of phosphorus is [Ne] 3s23p3. Which of the following shows the correct electron dot structure of phosphorus? A P

C P

B Ne

D Ne

18 The element bromine (Br) has 35 electrons. Those electrons are arranged in many different energy levels. Which of the following lists all of the energy levels that would contain electrons for ground state bromine? A 1s, 2s, 2p, 3s, 3p, 4s, 4p B 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p C 1s, 1p, 2s, 2p, 3s, 3p, 4s, 3d, 4p D 1s, 2s, 2p, 3s, 3p, 3d, 4s, 4p, 4d

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STAAR Chemistry

Practice Test A

19 Which equation correctly shows the beta decay of cesium-137? A

137 55  Cs



0 –1 e

+

137 56 Ba

C

137 55 Cs



0 0 γ

B

137 55 Cs



0 –1 e

+

137 54 Xe

D

137 55 Cs



4 2 Hw

137 55 Cs

+

+

133 53 I

20 Radon gas, Rn-222, is produced during alpha decay according to the reaction represented by the following equation. X→

222 46 Rn

+

4 2 He

What element decays to produce radon gas? A Ra-226 B Po-218 C Fr-223 D Th-226

21 A student investigates several different forms of nuclear radiation and makes the following list of characteristics of one type. Characteristics of a Type of Nuclear Radiation No mass Has a negative charge Can be blocked by a sheet of aluminum foil Which type of particle was investigated? A Alpha B Beta C Positron D Gamma

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STAAR Chemistry

Practice Test A

22 A student finds a bottle of HNO3 stored with other highly toxic and corrosive materials. In ancient times, this chemical was called aqua fortis, or “strong water,” because it was used in alchemy to dissolve silver and other metals. What is the IUPAC name of the compound? A Hydrogen nitric oxide B Hydrogen nitrate C Nitrous acid D Nitric acid

23 P2O5 is a very strong desiccator and is used in organic synthesis to dehydrate molecules. What is the name of the compound? A phosphorus pentoxide B phosphorus oxide C diphosphorus tetroxide D diphosphorus pentoxide

24 Aluminum sulfate is used in water purification and wastewater treatment to take colloids out of suspension. It is also found in baking powder and in antiperspirants. What is the chemical formula for aluminum sulfate? A AlS2O4 B AlSO4 C Al(SO4)3 D Al2(SO4)3

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STAAR Chemistry

Practice Test A

25 A student is searching for a bottle containing iron(III) hydroxide. Which formula will be written on the label of the bottle? A Fe3(OH)2 B FeH3O C Fe3OH D Fe(OH)3

26 Gallium sulfide (Ga2S3) is a yellow solid with a distinct rotten egg odor that ignites spontaneously in air. Which of the following is the correct electron dot formula?

A S

B

Ga

S

S 2− Ga3+ 2− S Ga3+ 2− S

Ga

S

C

S 3− Ga2+ 3− S Ga2+ 3− S

S

D S

Ga

Ga

S

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STAAR Chemistry

Practice Test A

27 Formaldehyde (CH2O) is a toxic compound used to preserve organic tissue. Which of the following is the correct electron dot structure for the molecule?

O A H

C

H

B H

C

O

H

C C

H

H

O

H C

O

D H

28 A boron atom has 3 valence electrons. What is the structure of a molecule of boron trichloride (BCl3) according to VSEPR theory? A Pyramidal B Trigonal planar C Tetrahedral D Linear

29 Which of the following corresponds with one mole of particles? A 6.02 × 1023 grams of CO2 × B

1 mol CO2 44 grams

6.02 × 1023 molecules of CO2 3 atoms / molecule

= 1.37 × 1022 mol CO2

= 2.00 × 1023 atoms CO2

C 6.02 × 1023 molecules of CO2 D 6.02 × 1023 molecules of CO2 ×

3 atoms 1 molecule CO2

= 18.06 × 1023 atoms CO2

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STAAR Chemistry

Practice Test A

30 How many molecules are in 45.0 g of C6H6? A 1.04 × 1024 molecules B 3.47 × 1023 molecules C 2.71 × 1025 molecules D 1.34 × 1022 molecules

31 How many ions are present in 50.0 grams of hydrogen cyanide (HCN)? A 3.34 × 1024 ions B 6.50 × 1023 ions C 1.11 × 1024 ions D 2.23 × 1024 ions

32 A compound contains 62.0% carbon, 10.4% hydrogen, and 27.5% oxygen. What is the empirical formula for the compound? A C2H4O B CH3O C C3H6O D CH2O

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STAAR Chemistry

Practice Test A

33 Which coefficient belongs before H2SO4 when the equation shown is balanced? H2SO4 + Fe → Fe2(SO4)3 + H2 A 1 B 2 C 3 D 4

34 Ethane (C2H6) is a colorless gas used as a fuel. Ethane burns in oxygen to produce carbon dioxide and water. The balanced chemical equation that describes this reaction is — A 2C2H6 + 7O2 → 4CO2 + 6H2O B 2C2H6 + 5O2 → 3CO2 + 6H2O C C2H6 + 7O2 → 4CO2 + 3H2O D 3C2H6 + 7O2 → 6CO2 + 9H2O

35 In an experiment, 19.2 g of water is produced when 10.0 g of methane is burned in an excess of oxygen. What is the percentage yield of water, to the nearest tenth of a percent? CH4(g) + 2O2 → CO2(g) + 2H2O(g) Record your answer and fill in the bubbles on your answer document.

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STAAR Chemistry

Practice Test A

36 With a deep breath, human lungs will contain about 5.00 L of air. If the pressure is 1.00 atm, and the temperature of the air in your lungs is 35.0°C, how many millimoles of oxygen will be in the lungs, if the air is 20.95% oxygen? A 198 mmol B 365 mmol C 1.74 × 103 mmol D 41.4 mmol

37 There are two balloons of identical size: one filled with air and the other with helium. Assume that the pressure and temperature in each balloon is the same. Which balloon weighs more and what is the ratio of the masses? Assume that air is 80% nitrogen and 20% oxygen. A The air balloon weighs about 7 times more than the helium balloon. B The air balloon weighs about 4 times more than the helium balloon. C The balloons have exactly the same mass of gas. D The helium balloon weighs about 3 times more than the air balloon.

3 8 The partial pressures of the various gases in our atmosphere can be explained by the kinetic molecular theory. Which of the statements below are correct postulates of the kinetic molecular theory? 1. Gas particles move in random, rapid, straight line motion between collisions. 2. Gases are mostly empty space containing very tiny hard spheres which repeatedly collide. 3. The absolute temperature of the gas is a measure of the total kinetic energy of the particles. 4. In any collision between gas particles, none of the total kinetic energy of the particles is lost. A 1 and 2 B 2 and 4 C 1, 2, and 4 D 1, 3, and 4

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STAAR Chemistry

Practice Test A

39 Many electric cars and hybrid electric vehicles use a regenerative braking system that converts the kinetic energy of the car to chemical energy of the battery. When brakes are applied, the electric motors become electrical generators, recharging chemical energy of the battery as they slow down the vehicle. Right Front Wheel

Power in

Power out

Electric motor

Acceleration Power in from battery; motor turns wheel

Electric motor

Braking Wheel turns motor; power returned to batteries

Comparing the chemical energy input from the battery, which accelerates the vehicle, to the chemical energy returned to the battery when the vehicle slowly brakes to a stop, we would expect the returned energy to be — A equal to the energy input because of the law of conservation of energy B more than the energy input because the tires and generators will heat up C less than the energy input because electrical energy generation is an endothermic process D less than the energy input because some heat will inevitably be lost to the surroundings

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STAAR Chemistry

Practice Test A

40 In the catalytic converter of a car, one of the chemical reactions that occur converts poisonous carbon monoxide to less toxic carbon dioxide. Using average bond energies from the table below, estimate the chemical potential energy change of the reaction described by this equation: 2CO(g) + O2(g) → 2CO2(g) (Note: The carbon-oxygen bond in carbon monoxide is actually a resonance hybrid bond. But for the purposes of this question, assume it is a simple double bond.) Bond

O=O

O–H

C=O

C–O

O–O

Average Bond Energy (kJ/mol)

498

467

745

358

204

A –992 kJ/mol B +992 kJ/mol C –498 kJ/mol D +498 kJ/mol

41 The reaction shown in the graph below could be represented by which of the following equations?

Energy

Reactants

Products

Reaction Coordinate

A Reactants + heat → Products (as an exothermic reaction) B Reactants → Products + heat (as an exothermic reaction) C Reactants + heat → Products (as an endothermic reaction) D Reactants → Products + heat (as an endothermic reaction)

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STAAR Chemistry

Practice Test A

42 On a warm day, a large glass of water is placed in a refrigerator. Calculate the specific heat of the glass using the recorded measurements in the table below, given that the refrigerator removes 39.4 kJ from the glass and water. Substance

Glass

Water

Mass (g)

322

320

Initial temperature (ºC)

28.4

28.4

Final temperature (ºC)

3.9

3.9

?

4.18

Specific heat capacity (J/g·ºC) A 0.840 J/g∙°C B 4.15 J/g∙°C C 0.154 J/g∙°C D 0.819 J/g∙°C

43  A student is given a constant-pressure calorimeter as Constant-Pressure shown below and asked to determine the molar heat Calorimeter of solution of potassium nitrate, KNO3(s). The student records the data shown in the table. Stirrer Foam lid

Water

Nested foam cups

Thermometer

Mass of beaker

106.8 g

Mass Mass of ofbeaker beaker and KNO3(s) Mass of beaker and KNO3(s)

106.8 g g 144.6 144.6 g

Volume of water

200 mL

Initial temperature of water

22.3ºC

Finaltemperature temperature of water Final of water

6.7ºC6.7ºC

Density water Densityofof water

1.001.00 g/mLg/mL

Heat capacity of water

4.18 J/g·ºC

Volume of water

Initial temperature of water

Heat capacity of water

200 mL 22.3ºC

4.18 J/g·ºC

Using the recorded measurements, calculate the molar heat of solution of potassium nitrate to the nearest tenth of a kJ/mol. Record your answer and fill in the bubbles on your answer document. Be sure to indicate whether the value is positive or negative.

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STAAR Chemistry

Practice Test A

44 Which of the following reactions is an acid-base reaction? A 4Al(s) + 3O2(g) → 2Al2O3(s) B Ca(NO3)2(aq) + Na2SO4(aq) → CaSO4(s) + 2NaNO3(aq) C HClO4(aq) + KOH(aq) → KClO4(aq) + H2O(l) D KBr(aq) + AgNO3(aq) → KNO3(aq) + AgBr(s)

45 Use the table below to predict which of the following would be expected if dilute solutions of Na2SO4(aq) and Ba(NO3)2(aq) are mixed.

Solubility Rules for Ionic Compounds in Aqueous Solutions Compounds

Solubility

Exceptions

Salts of alkali metals and ammonia

Soluble

Some lithium compounds

Nitrate salts and chlorate salts

Soluble

Few exceptions

Sulfate salts

Soluble

Compounds of Pb, Ag, Hg, Ba, Sr, and Ca

Chloride salts

Soluble

Compounds of Ag and some compounds of Hg and Pb

Carbonates, phosphates, chromates, sulfides, and hydroxides

Most are insoluble

Compounds of the alkali metals and of ammonia

A No precipitate would be observed. B Both NaNO3(s) and BaSO4(s) would precipitate. C BaSO4(s) would precipitate. D NaNO3(s) would precipitate.

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STAAR Chemistry

Practice Test A

46 A laboratory assistant follows written directions and prepares a solution for an experiment that contains 30.0 g of Na2SO4 in 400 mL of solution. What is the molar concentration of that solution to three decimal places? Record your answer and fill in the bubbles on your answer document.

47 Aqueous solutions of Ag2SO4 and an unknown compound are mixed and two different precipitates form. An equation representing the situation is Ag2SO4(aq) +

(aq) → Ag

(s) +

SO4(s)

Which of the following could be the unknown compound? A PbCO3 B SrCl2 C PbCl2 D Na2CO3

48 A student is asked to determine whether a given solution is an electrolyte. Which of the following pieces of laboratory equipment would be the most useful for that purpose? A pH meter B Conductivity apparatus C Burette and chemical indicator D High grade filter paper and funnel

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STAAR Chemistry

a

49

Practice Test A

a When a tiny crystal ofbthe solute is dropped into a solution, b the crystal appears to grow rapidly

and other crystals appear in the solution. These observations indicate that the solution was —

a

A concentrated

b

B unsaturated C saturated D supersaturated

50 Based on the graph below, an increase in temperature —

B increases the solubility of ionic compounds C causes compounds to dissolve D increases the rate at which substances dissolve

160

Solubility (g/100 g H2O)

A usually affects the solubility of ionic compounds

Solubility Varies With Temperature KNO3

140

NaNO3

120

KBr

100

NH4Cl

80 60

NaCl

40 20

Yb2(SO4)3 0

10

20

30

40

50

60

Temperature (ºC)

70

80

90

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STAAR Chemistry

Practice Test A

51 Based on the graph below, the solubility of oxygen in water increases with —

Oxygen Solubility in Fresh Water

50 40 4 2 1

30 20

Pressure (atm)

Oxygen Solubility (mg/L)

60

10 0

0

5

10

15

20

25

30

35

40

45

50

Temperature (ºC)

A an increase in both pressure and temperature B a decrease in both pressure and temperature C a decrease in pressure and an increase in temperature D an increase in pressure and a decrease in temperature

52 In an oxidation-reduction reaction, which of the following are transferred? A Protons and electrons B Electrons and neutrons C Electrons only D Protons only

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Name

Date

Class

STAAR Chemistry Practice Test A Answer Sheet Answer Sheet 1. 2. 3. 4. 5. 6. 7. 8. 9. 10.

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

24. 25. 26. 27. 28. 29. 30. 31. 32. 33. 34. 35.

11.

12. 13. 14. 15. 16.

17. 18. 19. 20. 21. 22. 23.

0

0

0

0

0

0

0

1

1

1

1

1

1

1

2

2

2

2

2

2

2

3

3

3

3

3

3

3

4

4

4

4

4

4

4

5

5

5

5

5

5

5

6

6

6

6

6

6

6

7

7

7

7

7

7

7

8

8

8

8

8

8

8

9

9

9

9

9

9

9

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

36. 37. 38. 39. 40. 41. 42. 43.

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

0

0

0

0

0

0

0

1

1

1

1

1

1

1

2

2

2

2

2

2

2

3

3

3

3

3

3

3

4

4

4

4

4

4

4

5

5

5

5

5

5

5

6

6

6

6

6

6

6

7

7

7

7

7

7

7

8

8

8

8

8

8

8

9

9

9

9

9

9

9

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

0

0

0

0

0

0

0

1

1

1

1

1

1

1

2

2

2

2

2

2

2

3

3

3

3

3

3

3

4

4

4

4

4

4

4

5

5

5

5

5

5

5

6

6

6

6

6

6

6

0

0

0

0

0

0

0

7

7

7

7

7

7

7

1

1

1

1

1

1

1

8

8

8

8

8

8

8

2

2

2

2

2

2

2

9

9

9

9

9

9

9

3

3

3

3

3

3

3

4

4

4

4

4

4

4

5

5

5

5

5

5

5

6

6

6

6

6

6

B

C

D

A

B

C

D

6 7

7

7

7

7

7

7

A

B

C

D

8

8

8

8

8

8

8

9

9

9

9

9

9

9

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

44. 45.

A

B

C

D

A

B

C

D

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Name

Date

Class

STAAR Chemistry Practice Test A Answer Sheet Answer Sheet 46.

A

B

C

47. 48. 49. 50. 51. 52.

D

0

0

0

0

0

0

0

1

1

1

1

1

1

1

2

2

2

2

2

2

2

3

3

3

3

3

3

3

4

4

4

4

4

4

4

5

5

5

5

5

5

5

6

6

6

6

6

6

6

7

7

7

7

7

7

7

8

8

8

8

8

8

8

9

9

9

9

9

9

9

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

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STAAR Chemistry

TEKS

Practice Test B

End-of-Course Assessment Chemistry Practice Test B

1 A student performs the following tests to determine the identity of a mineral sample. 1. The mineral is heated in chlorine gas to see if it will produce a different substance. 2. The mineral is immersed in an ammonia solution to see if any new color is produced. 3. The mineral is rubbed on a streak plate to determine the color of its powdered form. 4. Dilute hydrochloric acid is placed on the mineral to see if bubbles will form. Which of the above tests were for physical properties of the mineral? A Test 1 B Test 2 C Test 3 D Test 4

2 A sample of sulfur obtained from the crater of a volcano is carefully measured. Which of the following is an intensive property of the sample? A Density of 2.07 g/cm3 B Mass of 3.85 g C Volume of 1.86 cm3 D Temperature of 20°C

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STAAR Chemistry

Practice Test B

3 The diagram below shows the arrangement of atoms in liquid mercury. Which of the following properties of liquid mercury atoms explains why liquid mercury takes the shape of and fills only the bottom of a more voluminous container, but neither solid nor gaseous mercury does?

Liquid mercury atoms — A are practically incompressible B are close together but free to move C have a random arrangement D slightly increase their average separation when heated

4 Students in a laboratory are asked whether a given sample of liquid matter is a pure substance or a mixture. They brainstorm the following suggestions for a test to answer the question. 1. Add NaOH(aq) to see if a precipitate will form. 2. Evaporate it to see if a residue will remain. 3. Test the substance with a pH meter and record the value. 4. Examine it under a microscope to see if it is uniform throughout. 5. Run the liquid through a filter paper. Which of the above tests could yield useful information to answer the question? A 2 only B 2, 4, and 5 C 1, 2, and 5 D 1, 3, and 4

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STAAR Chemistry

Practice Test B

5 The graph below shows a periodic trend. 2500

He

Ne

2000 1500 1000

Ar

N H

500 0

Kr Xe

P

Be

Zn As

Mg Li

Na

10

Rb

K

20

Cd

30

40

Atomic number

Cs

50

60

Which of the variables below is plotted on the vertical axis of the graph? A Ionic radius B Atomic radius C Electronegativity D First ionization energy

6 During the discovery of many new elements in the 1800s, it was found that when the elements were arranged in order of increasing atomic mass — A the gases all seemed to be grouped together B reactivity decreased as the atomic mass increased C density increased as the atomic mass increased D the properties repeated in a regular pattern

7 Halogens and noble gases are alike in that most members of both families — A have relatively low boiling points B form ions with a +1 oxidation state C form ions with a –1 oxidation state D are relatively unreactive

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STAAR Chemistry

Practice Test B

8 A student is asked to identify a substance on the basis of its chemical properties. The table below lists some of the properties of the element bromine (Br). 1. state

liquid

2. color

reddish-brown

3. elemental state

diatomic

4. reactivity

moderately high

5. melting point (°C)

–7

6. boiling point (°C)

59

Which of the properties of bromine is/are chemical properties? A 1 and 3 only B 2 and 4 only C 4 only D 4, 5, and 6 only

9 A forensic technician is examining a fine white powder found at a crime scene. It appears to be of uniform texture and consistency. He finds that upon heating, the substance decomposes, releasing a colorless gas and leaving a black residue. Based on these observations the substance is — A a homogeneous mixture B a heterogeneous mixture C either a heterogeneous mixture or a compound D a compound

10 Group 1A elements differ from Group 2A elements in that Group 1A elements — A form a lower oxidation state B are less reactive C are harder and have higher densities D are found in nature only in compounds

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STAAR Chemistry

Practice Test B

11 Some of the third period elements are listed alphabetically in the table below. 1. aluminum 2. argon 3. chlorine 4. magnesium 5. sodium Arrange the elements in order of increasing atomic radius. List the numbers of the elements from the table above, with the smallest element on the left and the largest on the right. Record your answer and fill in the bubbles on your answer document.

12 The modern Periodic Table has elements with similar properties arranged in — A rows B periods C groups D sequential atomic numbers

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243

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STAAR Chemistry

Practice Test B

13 Several students make the following statements about red light. 1. Red light has a longer wavelength than infrared. 2. Red light has a shorter wavelength than infrared. 3. Red light has a lower frequency than ultraviolet. 4. Red light has a higher frequency than ultraviolet. Based on the diagram below, which statements are correct? Red light, low energy (7.0 × 10 –7 m = 700 nm)

Violet light, high energy (3.8 × 10 –7 m = 380 nm)

Visible light

Frequency 𝛎 (s −1) 3 × 106

3 × 10 8

Radio waves

10 2

1

Radar

3 × 10 10

3 × 10 12

Microwaves

10 –2

10 –4

3 × 10 14

Infrared

3 × 10 16

Ultraviolet

10 –6

10 –8

3 × 10 18

X-rays

10 –10

3 × 10 20

3 × 10 22

Gamma rays

10 –12

10 –14

Wavelength 𝛌 (m)

A 1 and 3 are correct B 2 and 4 are correct C 1 and 4 are correct D 2 and 3 are correct

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STAAR Chemistry

Practice Test B

14 Medical x-rays have wavelengths between 0.1 nm and 0.01 nm. Which of the following could be the frequency of a medical x-ray? A 8 × 1018 Hz B 8 GHz C 8 × 10–19 Hz D 8 Hz

15 Niels Bohr conducted many experiments in 1913 as he developed a new atomic model. Possible Observations that Led to Bohr’s Nuclear Atom Model 1. He observed that atoms will not emit electrons until light of a specific energy shines on them. 2. He observed that alpha particles would be deflected different amounts by the electrons in specific energy levels. 3. He observed that atoms always absorb or emit light with specific energies. 4. He observed different regions or electron clouds around the nucleus that had increased probability of containing electrons. Which of the observations above led Bohr to conclude that each electron orbit has a different energy? A 1 B 2 C 3 D 4

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STAAR Chemistry

Practice Test B

16 Copper has two naturally occurring isotopes. Use the information in the table to calculate the average atomic mass of copper, to the nearest hundredth of an atomic mass unit. Isotope

Relative Abundance

Mass (amu)

Cu-63

69.15%

62.93

Cu-65

30.85%

64.93

Record your answer and fill in the bubbles on your answer document.

17 The electron configuration of fluorine is 1s22s22p5. Which of the following shows the correct electron dot structure of fluorine? A F B

C

F

D F

F

18 Many of the unique properties of tin (Sn) are due to the electron arrangement within the atom. What is the ground state electron configuration of tin? A [Kr] 5s24d105p2 B [Kr] 5s25d105p2 C [Kr] 5s25p2 D [Kr] 5s24d104f145p2

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STAAR Chemistry

Practice Test B

19 What element is produced by beta decay in the following nuclear reaction? 34 15 P

A B C D

→ X + –10 e

33 14 Si 34 16 P 34 14 Si 34 16 S

20 Thorium-234 is produced in an alpha decay reaction as shown in the following reaction. X → 42 He +

234 90 Th

What radioactive element decays to produce Th-234? A Pu-236 B Ra-230 C U-238 D Ta-240

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STAAR Chemistry

Practice Test B

21 Thorium-232 can undergo a nuclear reaction as shown in the equation below. 232 90 Th

+

1 0 n



233 90 Th

+ energy

Without knowing what nuclear decay steps come next, a student predicts whether the process is a fission or fusion reaction. Neutron

235 92 U

236 92 U

Which of the following student’s reasoning is correct? 1. The process is fission, because energy is released. 2. The process is fission, because a neutron is absorbed to start the process, which can cause the nuclei to become unstable. 3. The process is fusion, because the final nuclei is larger than either of the initial nuclei. 4. The process is fusion, because there are no neutrons released which is what starts fission chain reactions. A 1 B 2 C 3 D 4

22 H3PO4 is a clear, colorless, odorless and slightly viscous liquid. It is an ingredient in several kinds of soda beverages because of its sour taste. What is the name of the compound with the chemical formula H3PO4? A Trihydrogen phosphate B Trihydrogen phosphorus tetroxide C Phosphoric acid D Trihydrogen phosphite Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

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STAAR Chemistry

Practice Test B

23 N2O, commonly known as laughing gas, is used in some medical procedures for its anesthetic and analgesic effects. It is also used as an oxidizer to increase the power output of engines in car racing. What is the name of the compound with the chemical formula N2O? A Nitrogen oxide B Dinitrogen oxide C Dinitrogen monoxide D Nitrogen monoxide

24 Ammonium carbonate is used as a smelling salt, which can revive someone who has fainted. What is the chemical formula for ammonium carbonate? A NH4(CO3)2 B NH4CO3 C (NH)2(CO3)4 D (NH4)2CO3

25 Nitrous acid is a weak acid that is found in the lower atmosphere and is one of the components of acid rain. What is the chemical formula for nitrous acid? A HN B HNO3 C HNO2 D HNO4

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STAAR Chemistry

Practice Test B

26 Ammonium fluoride (NH4F) crystals taste salty with a structure very similar to ice. Which of the following is the correct electron dot formula for this compound?

H A N

H

H

H

H

F

C H

N+

F−

H

H H

H B H

N

F

D H

H

H

N

F

H

H

27 Methanol (CH3OH), which is highly toxic to humans, can be used in antifreeze, as a fuel, and solvent. Which of the following electron dot formulas correctly represents the structure of the molecule? H A (CH )+(OH)− 3

C H

C

H

O

H

D C

H

H

H

O

H B H

C

O

H

H

H

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STAAR Chemistry

Practice Test B

28 A student made the following list about the nature of metallic bonding. Metallic Bonding Theory 1. Loosely-held valence electrons can drift freely throughout the metal. 2. Metal cations are arranged in a crystalline lattice. 3. Metallic bonds form from the attraction of the positive metal ions to the negative valence electrons. The property of thermal conductivity is explained by which of the above statements? A 1 B 2 C 3 D 1 and 2

29 Using VSEPR theory, what is the molecular structure of acetylene (C2H2)? A Trigonal planar B Bent C Tetrahedral D Linear

30 How many atoms are in 28.0 grams of copper? A 1.37 × 1024 atoms B 1.07 × 1023 atoms C 2.65 × 1023 atoms D 3.38 × 1020 atoms

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STAAR Chemistry

Practice Test B

31 How many ions are present in 37.0 grams of calcium chloride (CaCl2)? A 1.81 × 1024 ions B 6.02 × 1023 ions C 4.01 × 1023 ions D 2.01 × 1023 ions

32 What is the percent composition of nitrogen in ammonia (NH3)? A 82.22% B 93.33% C 14.00% D 25.00%

33 When this chemical equation is balanced, what is the coefficient of oxygen (O2)? C3H8(g) + O2(g) → H2O(g) + CO2(g) A 2 B 4 C 5 D 10

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STAAR Chemistry

Practice Test B

34 When calcium carbonate reacts with vinegar, bubbles of carbon dioxide form. What is the missing product in the balanced reaction below? CaCO3 + 2CH3COOH → Ca(CH3COO)2 + CO2 + A C2HO B CaO C CO D H2O

35 Natural gas, used for heating homes and in the production of electricity, is a gas mixture that is composed primarily of methane. The equation below describes the combustion of methane. CH4(g) + 2O2(g) → 2H2O(g) + CO2(g) What mass of oxygen, to the nearest gram, is needed to combine with 96.0 grams of methane? Record your answer and fill in the bubbles on your answer document.

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STAAR Chemistry

Practice Test B

36 Some scuba divers use a gas mixture called Nitrox I, to lessen the chance of getting decompression sickness. It contains 32% oxygen and 68% nitrogen. What is the partial pressure of nitrogen, to the nearest tenth of an atmosphere, in a sample of Nitrox I at the pressure and depth in seawater indicated in the diagram?

Pressure = 4.97 atm at 40 m

Record your answer and fill in the bubbles on your answer document.

37 Air is about 80% nitrogen and 20% oxygen. A balloon has a volume of 2.0 L. How many grams of oxygen are in the balloon at standard temperature and pressure? A 0.57 g B 0.29 g C 2.9 g D 13 g

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STAAR Chemistry

Practice Test B

38 Some students are invited to a barbecue. The food is cooked with propane burning in excess oxygen to produce heat, as shown in this equation. C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O(g) + heat If 210 g of propane gas are consumed, what volume of gaseous products is produced at standard temperature and pressure? A 320 L B 747 L C 427 L D 1.28 × 103 L

39 When sodium bicarbonate is added to a beaker of vinegar, a bubbling reaction occurs that raises the temperature of the beaker’s contents. In this reaction, what is being transformed into thermal energy? A Some of the mass of the reactants B Some of the chemical energy of the reactants C Some of the kinetic energy of the reactants D Some of the kinetic energy of the products

40 Two students are trying to understand the direction of heat flow and the extent of the system when food is cooked on a grill. They each have an explanation for the heat flow direction and system. Student 1

When the food is the system, there is an endothermic process occurring.

Student 2

When the grill and the food are the system, there is an exothermic process occurring.

A Only statement 1 is true. B Only statement 2 is true. C Both statements 1 and 2 are true. D Neither statement is true.

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STAAR Chemistry

Practice Test B

41 Elemental iron (Fe) can be reclaimed from iron(III) oxide (Fe2O3) according to the following reaction: Fe2O3(s) + 3CO(g) → 2 Fe(s) + 3CO2(g) + 26.3 kJ When 4.50 moles of iron(III) oxide react, which of the following statements is correct? 1. 26.3 kJ are absorbed in an exothermic reaction 2. 26.3 kJ are released in an endothermic reaction 3. 118 kJ are released in an exothermic reaction 4. 118 kJ are absorbed in an endothermic reaction A 1 B 2 C 3 D 4

42 Glucose is produced by green plants during photosynthesis. Organisms use glucose as a source of energy. When 36.0 g of glucose is metabolized according to the equation below, which choice correctly identifies the amount of energy released or absorbed? C6H12O6 + 6O2 → 6CO2 + 6H2O

ΔH = –2880 kJ

A 2880 kJ released B 2880 kJ absorbed C 576 kJ absorbed D 576 kJ released

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STAAR Chemistry

Practice Test B

43 A student finds that 4.7 kJ of heat are required to raise the temperature of 250.0 g of metal X from 22°C to 63°C. Based on the data in the table, what is the identity of metal X? Metal

Specific Heat (J/(g∙°C)

Aluminum

0.90

Iron

0.46

Mercury

0.14

Silver

0.24

A Aluminum B Iron C Mercury D Silver

4 4 Which of the following reactions is an oxidation-reduction reaction? A 2Ca(s) + O2(g) → 2CaO(s) B Ba(NO3)2(aq) + Na2SO4(aq) → BaSO4(s) + 2NaNO3(aq) C HBrO3(aq) + KOH(aq) → KBrO3(aq) + H2O(l ) D NaBr(aq) + AgNO3(aq) → NaNO3(aq) + AgBr(s)

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STAAR Chemistry

Practice Test B

45 Based on the table below, which of the following solutions should be added to (NH4)3PO4(aq) to produce a precipitate?

Solubility Rules for Ionic Compounds in Aqueous Solutions Compounds

Solubility

Exceptions

Salts of alkali metals and ammonia

Soluble

Some lithium compounds

Nitrate salts and chlorate salts

Soluble

Few exceptions

Sulfate salts

Soluble

Compounds of Pb, Ag, Hg, Ba, Sr, and Ca

Chloride salts

Soluble

Compounds of Ag and some compounds of Hg and Pb

Carbonates, phosphates, chromates, sulfides, and hydroxides

Most are insoluble

Compounds of the alkali metals and of ammonia

A LiCl(aq) B K2CrO4(aq) C Ca(ClO3)2(aq) D Na2SO4(aq)

46 What would be the molar concentration of sugar in a can of soda, if it contained 40.0 g of sucrose (C12H22O11) in 355 mL of solution? Express your answer to three decimal places. Record your answer and fill in the bubbles on your answer document.

47 A student in a laboratory prepares a solution of NaCl(aq) using tap water instead of distilled water as directed. She notices a faint cloudiness of the solution due to a precipitate. This precipitate in the solution may indicate a presence in the tap water of — A sulfate ions B hydroxide ions C calcium or magnesium ions D silver or lead(II) ions

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STAAR Chemistry

Practice Test B

48 A student is asked to determine whether a given solution is unsaturated, saturated, or supersaturated. The student drops a tiny crystal of the solute into the solution and observes carefully for any change in the size of the crystal. Which of the following statements is true? A No observable change means the solution is supersaturated. B A decrease in crystal size means the solution is unsaturated. C A decrease in crystal size means the solution is saturated. D An increase in crystal size means the solution is unsaturated.

49 A solution is classified as an electrolyte if the solution contains — A a dissolved solute B molecular compounds C positive and negative ions D substances with strong forces between their particles

50 Based on the graph to the right, an increase in temperature has the greatest effect on the solubility of compounds containing —

B chloride ions C sodium ions D potassium ions

160

Solubility (g/100 g H2O)

A nitrate ions

Solubility Varies With Temperature KNO3

140

NaNO3

120

KBr

100

NH4Cl

80 60

NaCl

40 20

Yb2(SO4)3 0

10

20

30

40

50

60

Temperature (ºC)

70

80

90

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STAAR Chemistry

Practice Test B

51 Which of the following procedures would increase the rate of dissolution of a gas in a liquid? 1. increase the temperature 2. increase the pressure 3. shake or stir the mixture 4. increase the surface area A 2 and 3 only B 1 and 2 only C 1, 2, and 3 D 2, 3, and 4

52 In an acid-base reaction, which of the following are transferred? A Protons and electrons B Electrons and neutrons C Protons only D Electrons only

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Name

Date

Class

STAAR Chemistry Practice Test B Answer Sheet Answer Sheet 1. 2. 3. 4. 5. 6. 7. 8. 9. 10.

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

24. 25. 26. 27. 28. 29. 30. 31. 32. 33. 34. 35.

11.

12. 13. 14. 15. 16.

17. 18. 19. 20. 21. 22. 23.

0

0

0

0

0

0

0

1

1

1

1

1

1

1

2

2

2

2

2

2

2

3

3

3

3

3

3

3

4

4

4

4

4

4

4

5

5

5

5

5

5

5

6

6

6

6

6

6

6

7

7

7

7

7

7

7

8

8

8

8

8

8

8

9

9

9

9

9

9

9

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

36.

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

0

0

0

0

0

0

0

1

1

1

1

1

1

1

2

2

2

2

2

2

2

3

3

3

3

3

3

3

4

4

4

4

4

4

4

5

5

5

5

5

5

5

6

6

6

6

6

6

6

7

7

7

7

7

7

7

8

8

8

8

8

8

8

9

9

9

9

9

9

9

A

B

C

D

0

0

0

0

0

0

0

1

1

1

1

1

1

1

2

2

2

2

2

2

2

3

3

3

3

3

3

3

0

0

0

0

0

0

0

4

4

4

4

4

4

4

1

1

1

1

1

1

1

5

5

5

5

5

5

5

2

2

2

2

2

2

2

6

6

6

6

6

6

6

3

3

3

3

3

3

3

7

7

7

7

7

7

7

4

4

4

4

4

4

4

8

8

8

8

8

8

8

5

5

5

5

5

5

5

9

9

9

9

9

9

9

6

6

6

6

6

6

6

7

7

7

7

7

7

7

8

8

8

8

8

8

8

9

9

9

9

9

9

9

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

37. 38. 39. 40. 41. 42. 43. 44. 45.

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

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Name

Date

Class

STAAR Chemistry Practice Test B Answer Sheet Answer Sheet 46.

A

B

C

47. 48. 49. 50. 51. 52.

D

0

0

0

0

0

0

0

1

1

1

1

1

1

1

2

2

2

2

2

2

2

3

3

3

3

3

3

3

4

4

4

4

4

4

4

5

5

5

5

5

5

5

6

6

6

6

6

6

6

7

7

7

7

7

7

7

8

8

8

8

8

8

8

9

9

9

9

9

9

9

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

A

B

C

D

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