Chapter 16 Acids and Bases

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Over time the Arrhenius concept of acids and bases came to be stated in the following way: • An acid is a substance that, when dissolved in water, increases the.

OBJECTIVES To familiarize with Acid-Base Equilibria Theories;  To become familiar with conjugated acid and base pairs;  To introduce acid and bases strength 

TEXTBOOK: Brown, Lemay & Bursten, Chemistry: The Central Science, 10th Ed. (Chapter 16)

INTRODUCTION (1) • Acids and bases are among the most important substances in chemistry, because they affect our daily lives in innumerable ways. • Not only are they present in our foods, but acids and bases are also crucial components of living systems, such as the amino acids that are used to synthesize proteins and the nucleic acids that code genetic information. • The impact of acids and bases depends not only on the type of acid or base, but also on how much is present.

A circle of shiny pennies is created by the reaction between the citric acid of the lemon and the tarnish on the surface of the copper.


• The application of acid–base chemistry has also had critical roles in shaping modern society, including such human-driven activities as industrial manufacturing, the creation of advanced pharmaceuticals, and many aspects of the environment. • The time required for a metal object immersed in water to corrode, the ability of an aquatic environment to support fish and plant life, the fate of pollutants washed out of the air by rain, and even the rates of reactions that maintain our lives all critically depend on the acidity or basicity of solutions.

INTRODUCTION (3) • From the earliest days of experimental chemistry, scientists have recognized acids and bases by their characteristic properties. • In this chapter we will take a closer look at how acids and bases are identified and characterized. In doing so, we will consider their behavior both in terms of their structure and bonding and in terms of the chemical equilibria in which they participate.

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Hydrochloric acid – HCl Nitric Acid – HNO3 Sulfuric Acid – H2SO4 Phosphoric Acid – H3PO4 In general acids taste sour.


Sodium Hydroxide – NaOH Calcium Hydroxide – Ca(OH)2 Ammonia – NH3 Sodium Carbonate – Na2CO3 In general bases taste bitter. They are slippery.

ACID- BASE EQUILIBRIA THEORIES There are three main Theories for acid-base equilibria: 1.Arrhenius (1880): Acid is who increases H+; 2.Brønsted-Lowry (1923): Acid is a proton donor; 3.Lewis (1923): Acid is a electronpair acceptor. Go to conjugate acids and bases pairs

ARRHENIUS THEORY (1) • This is the earliest or classical acid-base definition, which classifies these substances by their behavior in water. • The Swedish Chemist Svante Arrhenius (1859–1927) started from the fact that it is evident that all acids contain hydrogen but not all hydrogencontaining substances are acids. • Arrhenius defined acids as substances that produce H+ ions in water and bases as substances that produce OH- ions in water.

ARRHENIUS THEORY (2) Over time the Arrhenius concept of acids and bases came to be stated in the following way: • An acid is a substance that, when dissolved in water, increases the concentration of H+ ions. • A base is a substance that, when dissolved in water, increases the concentration of OH- ions. Hydrogen Chloride gas, which is highly soluble in water, is an example of an Arrhenius acid.

ARRHENIUS THEORY (3) When Hydrogen Chloride dissolves in water, HCl(g) produces hydrated H+ and Cl- ions:  HCl( g )  H( aq ) H 2O

  Cl( aq )

• The aqueous solution of HCl is known as hydrochloric acid. • Concentrated hydrochloric acid is about 37% HCl by mass and is 12 Mol/L in HCl. In general: H 2O   HX ( l ,g )  H( aq )  X ( aq )

ARRHENIUS THEORY (4) • Sodium hydroxide is an Arrhenius base; because NaOH is an ionic compound, and it dissociates into Na+ and OH- ions when it dissolves in water, thereby increasing the concentration of OH- ions in the solution. See below: H 2O   NaOH( s )  Na( aq )  OH( aq ) • In general: H 2O   XOH( s )  X ( aq )  OH( aq )

SUMMARIZING ARRHENIUS THEORY (1)  Acids: Substances that, when dissolved in water, increases the concentration of hydrogen ions.  Bases: Substances that, when dissolved in water, increases the concentration of hydroxide ions. Neutralization: is the reaction of an H+ (H3O+) ion from the acid and the OH- ion from the base to form water (H2O).

SUMMARIZING ARRHENIUS THEORY (2) Acid is a substance that produces H+ (H3O+) in water:

Base is a substance that produces OH- in water:

Back to Acid- Base Equilibria Theories

BRØNSTED–LOWRY THEORY (1) • The Arrhenius concept of acids and bases, while useful, is rather limited. For one thing, it is restricted to aqueous solutions. • Brønsted–Lowry concept is based on the fact that acid–base reactions involve the transfer of H+ ions (protons) from one substance to another. • To understand this definition better, we need to examine the behavior of the H+ ion in water more closely.

In 1923 the Danish Chemist Brønsted (1879–1947) and the English Chemist Lowry (1874–1936) independently proposed a more general definition of acids and bases.

BRØNSTED–LOWRY THEORY (2) • We might at first imagine that ionization of HCl in water produces just H+ and Cl-. • A hydrogen ion is no more than a bare proton: a very small particle with a positive charge. • As such, an H+ ion interacts strongly with any source of electron density, such as the nonbonding electron pairs on the oxygen atoms of water molecules. • For example, the interaction of a proton with water forms the hydronium ion (H3O+):  H( aq )

  H2O  H3O( aq )


• The transfer of a proton always involves both an acid (donor) and a base (acceptor). • In other words, a substance can function as an acid only if another substance simultaneously behaves as a base. • To be a Brønsted – Lowry acid, a molecule or ion must have a hydrogen atom it can lose as an H+ ion (proton). • To be a Brønsted – Lowry base, a molecule or ion must have a nonbonding pair of electrons it can use to bind the H+ ion.


Let’s consider an example that compares the relationship between the Arrhenius and Brønsted – Lowry definitions of acids and bases: an aqueous solution of ammonia, in which we have the equilibrium:  NH3( aq )  H2O( l )  NH4 ( aq )

  OH( aq )

• Ammonia is a Brønsted – Lowry base because it accepts a proton from H2O. NH3 is also an Arrhenius base because adding it to water leads to an increase in the concentration of OH-

BRØNSTED–LOWRY THEORY (5) MONOPROTIC VS POLIPROTIC ACIDS: • Monoprotic acids are those that can donate only one proton (Ex.: HCl). • Some acids can donate more than one proton to the solution. • Thus a diprotic acid has two protons such as: H2S and H2SO4. • While a common triprotic acid has three acidic protons that can be donated like: H3PO4.


Proton donors Proton acceptors

So, a Brønsted–Lowry acid: • Must have a removable (acidic) proton. Then, a Brønsted–Lowry base: • Must have a pair of nonbonding electrons.

SUMMARIZING BRØNSTED– LOWRY THEORY (2) A Brønsted-Lowry acid is a proton donor. A Brønsted-Lowry base is a proton acceptor.



conjugate conjugate acid base Back to Acid- Base Equilibria Theories

LEWIS THEORY (1) • Lewis suggested another way of looking at the reaction between H+ and OH- ions. • In the Brønsted-Lowry Theory, the OHion is the active species in this reaction it accepts an H+ to form a covalent bond. • In the Lewis Theory, the H+ ion is the active species it accepts a pair of electrons from the OH- ion to form a covalent bond. • This Theory, states that bases donate pairs of electrons and acids accept pairs of electrons.

Gilbert Newton Lewis (1875 – 1946) was an American Physical-Chemist known for the his concept of electron pairs. Lewis successfully contributed to Thermodynamics, and is also known for his concept of acids and bases.

LEWIS THEORY (2) Acids are electron-pair acceptors. Bases are electron-pair donors. F +










adduct M(H2O)42+(aq)

+ M2+




SUMMARIZING LEWIS THEORY • A Lewis acid is therefore any substance, such as the H+ ion, that can accept a pair of nonbonding electrons. In other words, a Lewis acid is an electron-pair acceptor. • A Lewis base is any substance, such as the OH- ion, that can donate a pair of nonbonding electrons. A Lewis base is therefore an electron-pair donor. • The principal advantage of the Lewis Theory is the way it expands the number of acids and therefore the number of acid-base reactions. Back to Acid- Base Equilibria Theories

CONJUGATE ACID-BASE PAIRS (1) What Happens when an acid dissolves in water? • Water acts as a Brønsted–Lowry base and abstracts a proton (H+) from the acid. • As a result, the conjugate base of the acid and a hydronium ion are formed.

CONJUGATE ACID-BASE PAIRS (2) • From the Latin word conjugare, meaning “to join together.” • Reactions between acids and bases always yield their conjugate bases and acids, E.g.:

CONJUGATED ACID-BASE PAIRS AN EXAMPLE TO CONSIDER (1) 1. What is the conjugate base of HClO4, H2S and PH4+? 2. What is the conjugate acid of CN- , SO42and H2O? Solution: • Analyze: We are asked to give the conjugate base for several acids and the conjugate acid for several bases. • Plan: The conjugate base of a substance is simply the parent substance minus one proton, and the conjugate acid of a substance is the parent substance plus one proton.

CONJUGATED ACID-BASE PAIRS AN EXAMPLE TO CONSIDER (2) 1. What is the conjugate base of HClO4, H2S and PH4+? 2. What is the conjugate acid of CN- , SO42and H2O? Solution: 1. If we remove a proton from HClO4, we obtain ClO4- , which is its conjugate base. The other conjugate bases are HS- and PH3. 2. If we add a proton to CN- , we get HCN, its conjugate acid. The other conjugate acids are HSO4- and , H3O+.


• In any acid–base equilibrium, both the forward reaction (to the right) and the reverse reaction (to the left) involve proton transfer. • In the forward reaction, HA donates a proton to H2O. Therefore, HA is the Brønsted–Lowry acid and H2O is the Brønsted–Lowry base. In the reverse reaction, the H3O+ ion donates a proton to the A- ion, so H3O+ is the acid and A- is the base. • When the acid HA donates a proton, it leaves behind a substance (A-) that can act as a base. Likewise, when H2O acts as a base, it generates H3O+, which can act as an acid.

ACID AND BASE STRENGTH (1) Strong acids are completely dissociated in water.  Their conjugate bases are quite weak. Weak acids only dissociate partially in water.  Their conjugate bases are weak bases.

ACID AND BASE STRENGTH (2) Substances with negligible acidity do not dissociate in water.  Their conjugate bases are exceedingly strong.

ACID AND BASE STRENGTH (3) In any acid-base reaction, the equilibrium will favor the reaction that moves the proton to the stronger base: HCl(aq) + H2O(l) D H3O+(aq) + Cl−(aq)

• H2O is a much stronger base than Cl−, so the equilibrium lies so far to the right K is not measured (K>>1).

ACID AND BASE STRENGTH (4) For the following example: C2H4O2(aq)+H2O(l) D H3O+(aq)+C2H3O2−(aq)

• Acetate is a stronger base than H2O, so the equilibrium favors the left side (K