Synthesis and Reactivity in Inorganic, Metal-Organic

This article was downloaded by: [Yahia Z. Hamada] On: 23 October 2013, At: 11:34 Publisher: Taylor & Francis Informa Ltd Registered in England and Wales Registered Number: 1072954 Registered office: Mortimer House, 37-41 Mortimer Street, London W1T 3JH, UK

Synthesis and Reactivity in Inorganic, Metal-Organic, and Nano-Metal Chemistry Publication details, including instructions for authors and subscription information: http://www.tandfonline.com/loi/lsrt20

Interaction of Malate and Lactate with Chromium(III) and Iron(III) in Aqueous Solutions a

a

Yahia Z. Hamada , Brandon Carlson & Joseph Dangberg

a

a

Department of Physical Sciences and Mathematics , Wayne State College , Wayne, NE, USA Published online: 15 Feb 2005.

To cite this article: Yahia Z. Hamada , Brandon Carlson & Joseph Dangberg (2005) Interaction of Malate and Lactate with Chromium(III) and Iron(III) in Aqueous Solutions, Synthesis and Reactivity in Inorganic, Metal-Organic, and Nano-Metal Chemistry, 35:7, 515-522, DOI: 10.1080/15533170500198887 To link to this article: http://dx.doi.org/10.1080/15533170500198887

PLEASE SCROLL DOWN FOR ARTICLE Taylor & Francis makes every effort to ensure the accuracy of all the information (the “Content”) contained in the publications on our platform. However, Taylor & Francis, our agents, and our licensors make no representations or warranties whatsoever as to the accuracy, completeness, or suitability for any purpose of the Content. Any opinions and views expressed in this publication are the opinions and views of the authors, and are not the views of or endorsed by Taylor & Francis. The accuracy of the Content should not be relied upon and should be independently verified with primary sources of information. Taylor and Francis shall not be liable for any losses, actions, claims, proceedings, demands, costs, expenses, damages, and other liabilities whatsoever or howsoever caused arising directly or indirectly in connection with, in relation to or arising out of the use of the Content. This article may be used for research, teaching, and private study purposes. Any substantial or systematic reproduction, redistribution, reselling, loan, sub-licensing, systematic supply, or distribution in any form to anyone is expressly forbidden. Terms & Conditions of access and use can be found at http:// www.tandfonline.com/page/terms-and-conditions

Synthesis and Reactivity in Inorganic, Metal-Organic and Nano-Metal Chemistry, 35:515–522, 2005 Copyright # 2005 Taylor & Francis, Inc. ISSN: 0094-5714 print/1532-2440 online DOI: 10.1080/15533170500198887

Interaction of Malate and Lactate with Chromium(III) and Iron(III) in Aqueous Solutions Yahia Z. Hamada, Brandon Carlson, and Joseph Dangberg

Downloaded by [Yahia Z. Hamada] at 11:34 23 October 2013

Department of Physical Sciences and Mathematics, Wayne State College, Wayne, NE, USA

Malate is generated from the hydration of fumarate by the action of fumarase in the citric acid cycle. Lactate is one of the most important products of glucose catabolism. From our ongoing efforts to study and understand the interaction and behavior of many hydroxycarboxylates with variety of metal ions (Hamada et al., 2003a; 2003b), we are reporting the interaction and behavior of malate and lactate with the Cr31 and the Fe31 ions using the potentiometric method in aqueous solutions at 25 8 C. It appeared that the potentiometric titration graphs are different for the different titration systems; these differences are discussed. It also appeared that the malate as a chelating ligand solubilized both Cr31 and Fe31 at the physiological pH. In contrast to malate, the lactate ligand did not solubilize either Cr31 or Fe31 at the physiological pH. The Cr31 and the Fe31 lactate systems form the simple one-to-one metal chelates because only one proton was observed in all metal to lactate molar ratios with both the Cr31 and the Fe31 metal ions, even when we increased the lactate to metal ratio to 5 fold excess, i.e., in the 1 : 5 titration systems. Keywords

malate, lactate, chromium, iron, ternary complexes

INTRODUCTION Malate Malate is ubiquitous in the biological environment whether it is cytosolic, glyoxysomal, or mitochondrial. It is the product of the conversion of fumarate via the action of fumarase in the citric acid cycle (mitochondrial). It is the main end product of the glyoxalate pathway (glyoxysomal). It is part of the urea cycle (cytosolic). It is also part of the malate-aspartate Received 13 May 2004; accepted 8 May 2005. The authors would like to thank the Nebraska Academy of Science for the research grants awarded to Y. Z. Hamada over the last four years to purchase most of the chemicals and equipment used in this study. YZH also would like to thank the new initiative grant from WSC, and WSC foundation annual grant. We would like to give special thanks to Professor Mark Hammer of WSC for reading the manuscript and his useful suggestions. Address correspondence to Yahia Z. Hamada, LeMoyne – Owen College, 807 Walker Ave., Memphis, TN 38126, USA. E-mail: [email protected]

shuttle in which cytosolic oxaloacetate is reduced to malate for transport into the mitochondrion (Voet et al., 1999). Lactate Lactic acid, an intermediate in carbohydrate metabolism (Voet et al., 1999), is derived primarily from muscles, brain, and erythrocytes. During vigorous exercise, lactate concentration may increase significantly from an average concentration of 0.9 mM to more than 20 mM within 10 seconds (Burtis and Ashwood, 1994). There is no accepted concentration for lactate for the diagnosis of the so-called lactate acidosis disorder, but lactate concentration exceeding 5 mM and at pH less than 7.25 indicates significant lactate acidosis (Burtis and Ashwood, 1994). In humans, the normal L(þ) lactate concentration in whole blood is in the range of 0.5 mM to 2.2 mM, depending on whether it is arterial or venous blood. In children ages 6 months to 3 years, the plasma concentration of D( –) lactate is in the range of 0 mM to 0.3 mM (Rosenthal and Pesca, 1985). Much of the lactate produced is exported to the liver, where it is reused to synthesize glucose (Voet et al., 1999). Very recently, an article reported an 11-month-old girl with B-cell leukemia/lymphoma who developed severe lactate acidosis (Svahn et al., 2003). Studying the interaction of lactate with the biologically essential metal ions such as Al(III), Fe(III) and Cr(III) is one of our interests (Scheme 1). Chromium Cr3þ is an essential trace metal necessary for the formation of the so-called ‘low-molecular-weight chromium-binding’ (LMWCr). The LMWCr has been suggested to be an organic low-molecular weight complex that contains four Cr3þions and an oligopeptide composed of glycine, cysteine, aspartic acid, glutamic acid with the carboxylate comprising more than half of the total amino acid residues (Vincent, 2001). The isolation and characterization of LMWCr has not been achieved, and thus its precise structure is still somewhat controversial. The site and pathway of LMWCr biosynthesis are unknown. The site and mechanism of intestinal chromium absorption in humans have not been determined (Burtis and Ashwood, 1994), although there is a wealth of literature

515

516

Y. Z. HAMADA ET AL.


SCH. 1. Hydroxycarboxylic acids used in this study.

available, a considerable amount of experimental and theoretical calculations are still needed in this area (Vincent, 2000; 2001; Sun et al., 2000; Buffie et al., 2001; Thompson and Connick, 1981; Finholt et al., 1981; Stunzi and Marty, 1983; Levina et al., 2001; Chang et al., 1983; Green et al., 1984; Broderic and Legg, 1985). EXPERIMENTAL SECTION Materials and Equipment All solutions were prepared using reagent grade D ,L -lactic acid, formula weight ¼ 90.08 g/mol (d ¼ 1.20 g/ml), malic acid, formula weight ¼ 134.09 g/mol, iron nitrate nonahydrate, Fe(NO3)3 . 9H2O, formula weight ¼ 404.00 g/mol, and chromium nitrate nona-hydrate, Cr(NO3)3 . 9H2O, formula weight 400.15 g/mol using deionized water. The KBr used for the IR experiments was IR-grade. The pHs were measured using Hanna Instrument HI 8314 membrane pH meter with a combination glass electrode. No other salts were added to adjust the ionic strength of the solutions. The UV-Vis spectra were measured in quartz cuvettes using HP 8421A single beam diode array spectrophotometer. The IR spectra were obtained using the Bowmann B-100 IR spectrophotometer as KBr pellets. Preparation of the Potentiometric Titration Solutions In any metal-ligand potentiometric titrations, the NaOH solution was the titrant. The NaOH solutions were prepared from NaOH laboratory grade pellets. The NaOH solutions were standardized using primary standard potassium hydrogen phthalate (KHP) that was purchased from Fisher Chemical Co. Before any KHP titration, the KHP was dried at 110oC for 24 hours, and then kept in a desiccator. A stock indicator solution of 0.20% phenolphthalein in 90% ethanol was prepared from reagent grade phenolphthalein. Nine KHP titrations were performed and averaged to calculate the exact NaOH concentration. We performed t-test and Q-test for the numbers generated. Potentiometric Titrations The potentiometric titration solutions were contained in a 250 mL beaker equipped with a magnetic stirring bar. The beaker was covered with a custom-made Teflon cover. The total concentration of the metal ions was always 2.50 mM. In a typical titration, the ligand was added first (lactic acid or malic acid), then the metal ion solution (Cr3þ or Fe3þ) was

added second, followed by the addition of enough water to take the total volume to 100.00 mL. This sequence of ligand addition followed by metal addition insured minimum metal ion hydrolysis at the start of the titration. No other salts were added to adjust the ionic strength of the solutions. Before each titration, the titration solution mixtures were allowed to stand for 25 – 30 minutes for complete equilibration. The NaOH titrant was added in the 100 mL increments using Eppendorf micropipette. We set the time intervals between the additions of the NaOH solution to three minutes, which was sufficient enough to get each of the pH values stabilized. Three runs for each metal-ligand titration were performed to ensure reproducible data. Aqueous Solution Synthesis of the Chromium-Lactate Complexes The following procedure was carried out as an attempt to grow crystals of the Cr3þ/lactic system in 1 : 3 molar ratio. A sample of 50.00 mL of 0.04713 M Cr3þ and 0.7497 g lactic acid along with 10.00 mL deionized water were placed in a 250 mL beaker that was sealed and left to stabilize for 120 hours. The initial pH after complete stabilization appeared to be 2.10. Freshly prepared and standardized 0.1000 M NaOH was added to deprotonate the carboxylate group in order to enhance the metal binding. We noted no precipitate formation as the pH increased from 2.10 to 4.24. The solution had a bluish color in the pH range of 2.10 to 4.24. The Cr3þ/lactic solution was placed in a 250 mL round bottom flask and most of the water was removed at 708C for 20 hours under reduced pressure. The remaining solution was placed on an evaporating dish and left to evaporate completely in a dark cabinet. The synthesized micro-crystals were not suitable to be sent to an X-ray lab to solve their crystal structure. A small amount of the micro-crystals grown on the evaporating dish described above were scraped and its weight was measured to be 1.0 mg. Using a clean spatula, the sample was placed into a mortar with enough IR grade KBr to reach 20 : 1 ratio of KBr to Cr3þ/lactic complex. After being uniformly mixed into a fine powder, the mixture was placed into the compartment of the pellet press and a semi-clear KBr pellet was prepared. We repeated this process twice. In each case we observed identical IR spectra. RESULTS AND DISCUSSION Cr31-Malate System Figure 1 shows the family of titration curves for the Cr3þmalate system in different molar ratios. Table 1 is the summary of these curves. The free malic acid curve was generated by titrating 0.50 mmol of free malic acid. It is clear from Figure 1 and Table 1 that malic acid is a di-protic acid, releasing a net of two protons from the two carboxylate groups. This allowed us to define malate as an H2L ligand. No reliable data have been presented in either Martell’s


INTERACTION OF MALATE AND LACTATE

FIG. 1. pH titration curves of free malic acid and Cr3þ/malic acid in 1 : 1, 1 : 2, 1 : 3, 1 : 4, 1 : 5, and 1 : 6 molar ratios with 0.4922 M NaOH. The free malic acid released 2.05 + 0.13 equivalents which represents two protons. The equivalent is defined as the milli-moles of NaOH per milli-moles of Cr3þ.

TABLE 1 Titration data for the Cr3þ-malate, 258C, [Cr3þ] ¼ 2.50 mM M:L ratio

mL NaOH

Eq.a

Comments

0 : 1b 1:1

2.00 2.40

2.05 + 0.13 4.70 + 0.15

1:2

3.20

6.30 + 0.10

1:3

4.40

8.70 + 0.10

1:4

5.40

10.60 + 0.05

1:5

6.40

12.60 + 0.05

1:6

7.40

14.60 + 0.05

Malic is a di-protic acid More than 4.0 Hþ released Trimeric complex may be present More than 8.0 Hþ released Trimeric complex may be present Trimeric complex may be present Trimeric complex may be present

a

The equivalent (Eq.) is defined as the number of milli-moles of NaOH titrant per milli-moles of malate. In the presence of Cr3þ, or any other metal in that regard, the equivalent is defined as the number of milli-moles NaOH per milli-moles Cr3þ. b 0.50 milli-moles malic acid was used to generate the free acid curve. Three runs generated 1.97, 2.00, and 2.20 equivalents or 2.05 + 0.13 protons from the carboxylates of malic acid.

517

data-base or in the literature regarding the Cr3þ-malate system (Martell and Smith, 2001). From the 1 : 1 titration system, it appeared that 4.70 + 0.15 equivalents of base are used to terminate the sharp inflection point formed. The equivalent point is defined as the number of milli-moles of NaOH per milli-moles of Cr3þ. This indicated that at least four protons have been released from the Cr3þ/ malate complex or complexes. This further indicated that the simple one to one metal chelate (at which one malate ligand is binding to one chromium ion via the two carboxylates only) did not form. The hydroxyl group has to be deprotonated to account for an extra proton in addition to another hydrolytic proton from the aqua ligand. The simplest scheme expected is the appearance of an inflection with two protons in the event of the formation of the one to one metal chelate. That was not the case. From surveying the literature of the hydroxy carboxylates and the notation of the chelate effect (Hamada et al., 2003a; 2003b; Douglas et al., 1994; Kettle, 1996), we are predicting that upon the chelation of the carboxylate(s) of the malic acid, the chelate effect will be dominant and force the hydroxy group to participate in the metal chelation, forming the stable fused five and six-membered chelating rings (Douglas et al., 1994; Kettle, 1996). Based on the above mentioned observations, we are proposing the appearance of a ternary hydroxo chromium malate complex [Cr3þ (Mal32) (OH2)]2 at which the (Mal32) is the fully deprotonated malate in addition to the loss of the hydroxyl proton. This indicates that the fully deprotonated malate is in its tri-anionic state. These Cr3þ/malate potentiometric titration plots shed some light on this system. For the 1 : 2 titration system we have overlaid two of the 1 : 2 titration plots to further ensure data reproducibility (see Figure 1). There is a significant possibility that this ternary hydroxo complex may form a dinuclear species in which the hydroxide ion acts as a bridge. Fe31-Malate System Figure 2 shows the family of titration curves for the Fe3þmalate system. Table 2 is the summary of these curves. The Fe3þ-malate system is similar to the Cr3þ-malate system. We are predicting that the two systems will generate the same kind of species or group of species. It is noteworthy to mention that there was no precipitation for either the Cr3þ/ malate or the Fe3þ/malate systems. The malate as a chelating ligand kept both the Cr3þ and Fe3þ ions in solution throughout the entire pH range of pH 2.50 to 11.50. In other words, malate solubilized both metal ions, as did the most famous biologically important hydroxyl carboxylate, citrate (Martell and Smith, 2001; Hamada et al., 2003a; 2003b). From the 1 : 2 titration system, the appearance of the inflection at 6.30 + 0.05 equivalents makes sense because of the presence of an extra mole of free malate that released a net of two protons from the two carboxylates plus four protons from the formation of [Fe3þ (Mal3-) (OH2)]2. Likewise, from the 1 : 3 titration system, the appearance of the inflection at 8.70 + 0.10 equivalents makes sense because of the

518

Y. Z. HAMADA ET AL.


kind of lower oligomer(s) will be formed in this milli-molar concentration and under the described experimental conditions. This could account for the fractions of the equivalents observed (Martell and Smith, 2001; Hamada et al. 2003a; 2003b). For example, the 1 : 2 titration system showing an inflection at 6.30 + 0.05 equivalents instead of 6.00 suggests that the dominant metal chelate present in solutions might be the trinuclear species (Martell and Smith, 2001; Hamada et al., 2003b; Feng et al., 1990). At this point, we cannot tell the exact structure to this tri-nuclear species. Martell’s database showed, for the Fe3þ/malate system, the presence of a dimeric [Fe2 (Mal)2] and a trimeric [Fe3 (Mal)4] for the ferric ion (Martell and Smith, 2001).

FIG. 2. pH titration curves of free malic acid and Fe3þ/malic acid in 1 : 2, 1 : 3, 1 : 4, 1 : 5, and 1 : 6 molar ratios with 0.4922 M NaOH, [Fe3þ] ¼ 2.5 1023 M. The equivalent points were at 6.3, 8.7, 10.24, 11.81, and 14.17 equivalents, see Figure 1 and text for the definition of the equivalent.

presence of two extra moles of free malate that released a net of four protons from the two carboxylates of each malate, plus four protons from [Fe3þ (Mal32) (OH2)]2. A ninth proton can easily be released from the aqua ligand, which is not that unusual for most trivalent metal ions in aqueous solutions (Martell and Smith, 2001). It might be expected that some

Cr31-Lactate System Figure 3 shows the family of titration curves for the Cr3þlactate system. Table 3 is the summary of these curves. The free lactic acid curve was generated by titration 0.266 millimoles of the acid. It is clear from Figure 3 and Table 3 that the acid is a mono-protic acid (releasing a net of one proton from the single carboxylate group present). From the 1 : 1, 1 : 2, 1 : 3, 1 : 4, and 1 : 5 titration systems, it appears that one proton has been released. This indicates that the simple one to one metal chelate (at which one lactate ligand is binding to one chromium ion) is formed. The alcoholic hydroxyl group did not release its protons because the number of equivalents never exceeded two protons. It is worth mentioning that the solution had a bluish green color throughout the titration, and around pH 6.50 there was turbidity

TABLE 2 Titration data for the Fe3þ-malate, 258C, [Fe3þ] ¼ 2.50 mM M:L ratio

mL NaOH

Eq.a

Comments

0 : 1b 1:1 1:2

2.00 N/A 3.20

2.05 + 0.13 N/A 6.30 + 0.05

1:3

4.40

8.70 + 0.10

1:4

5.20

10.24 + 0.10

1:5

6.00

11.81 + 0.10

1:6

7.20

14.17 + 0.10

Malic is a di-protic acid N/A Trimeric complex may be present More than 8.0 Hþ released More than 10 Hþ released More than 11 Hþ released A little more than 14 Hþ released

a

See Table 1 and Figure 1 for the definition of the equivalents. See Table 1 for details.

b

FIG. 3. pH titration curves of free lactic acid and Cr3þ/lactic acid in 1 : 1, 1 : 2, 1 : 3, 1 : 4, and 1 : 5 molar ratios with 0.4922 M NaOH. [Cr3þ] ¼ 2.5 mM. The free lactic acid released one proton off the carboxylic acid moiety. Turbidity started around pH 6.0.


519


TABLE 3 Titration data for the Cr3þ-lactate, 258C, [Cr3þ] ¼ 2.50 mM M:L ratio

mL NaOH

Eq.a

Comments

0 : 1b

0.50

0.93 + 0.01

1:1 1:2 1:3

1.80 2.00 2.40

0.88 + 0.05 0.98 + 0.05 1.18 + 0.07

1:4

2.80

1.38 + 0.05

1:5

3.20

1.58 + 0.08

Lactic is a monoprotic acid 1 Hþ released or less 1 Hþ released A little more than 1.0 Hþ released More than 1.0 Hþ released 1.5 Hþ released but not 2.0 Hþ

a

The equivalent is defined as the number of milli-moles of NaOH titrant per milli-moles of lactate. In the presence of Cr3þ the equivalent is defined as the number of milli-moles NaOH per milli-moles Cr3þ. b 20 mL (0.266 milli-moles) lactic acid was used to generate the free acid curve, which corresponds to one proton.

or a solid material that precipitated out. In another words, the lactate as a chelating agent did not solubilize Cr3þ above pH 6.5. That is in contrast to the Cr3þ/malate system. In contrast to these observations, all solutions in Equation (1) are colorless. Lactic acid (HL)(aq) þ NaOH(aq) ! Sodium lactate (NaL, colorless)(aq) þ H2 O

ð1Þ

UV-Vis Absorption Spectra of the Cr31/Lactate System Figure 4 shows the UV-Vis absorption spectra for the Cr3þ/ lactic acid system. The total chromium concentration is 19.28 mM and the total lactic acid concentration is 449.30 mM. These concentrations gave a Cr3þ to lactic acid ratio of 1 to 23. We increased the total Cr3þ concentration to observe decent absorption peaks in the acidic region between 0.20 and 0.80 absorption units. We also selected this high lactate to Cr3þ ratio to minimize the oligomerization and the precipitation process. Regardless of all of these precautions, the solution precipitated above pH 5.10. The initial pH of the mixture was 1.75 and the final pH recorded was 9.40. Typical Cr3þ d-d electronic transition peaks were observed from the family of titration curves shown in Figure 4. The peaks at 400 nm and 580 nm were observed as expected from the literature (Thompson and Connick, 1981; Chang et al., 1983; Hamada et al., 2003a; Douglas et al., 1994; Kettle, 1996). The same peaks were observed for the Cr3þcitrate system (Hamada et al., 2003a), the Cr3þ-nicotinic acid system (Chang et al., 1983), and even the Cr3þ aqua system (Thompson and Connick, 1981; Douglas et al., 1994; Kettle, 1996). Between pH 5.10 and pH 9.40, the pH values were

FIG. 4. UV-Vis absorption spectra of Cr3þ/lactic acid in 1 : 23 mole ratio. [Cr3þ] ¼ 19.28 mM, [Lac] ¼ 449.3 mM. The pH range was from 1.75 to 9.40. There was a precipitate above pH 5.10 and the pH reading became erratic above that value.

erratic and unstable, probably because this is where the inflection point is expected to appear as it has been shown from the pH-titrations, or because of the formation of a precipitate. There is also the possibility of competition of the titrant hydroxide ligand with the lactate as metal chelate. At any given pH value the spectra have rising absorbance into the shorter UV range. This observation was true for the Cr3þ/citrate as well (Hamada et al., 2003a). By examining Figure 4 more closely, we can see that the lower energy electronic transition, 4A2 g ! 4T2 g, which represents the absorption peak at 580 nm, had little change as the pH was increased from 1.75 to 9.40. On the other hand, for the higher energy transition, 4A2 g ! 4T1 g, which represents the absorption peak at 400 nm, a red shift for all peaks between pH 1.75 to pH 4.73 followed by a blue shift between pH 4.73 and 9.40 can be observed. It is worth noting that the absorbance increased as the pH increased until the erratic pH behavior occurred. Fe31-Lactate System Figure 5 shows the family of titration curves for the Fe3þlactate system. Table 4 is the summary of these curves. From the 1 : 1, 1 : 2, 1 : 3, 1 : 4, and 1 : 5 titration systems, it appeared again that only one proton has been released. This indicates that the simple one to one metal chelate is formed. The alcoholic hydroxyl group did not release its protons because the number of equivalents never reached two protons. The solution had a yellowish green color between pH 2.00 to 3.50, which changed to a reddish bronze color, and eventually, around pH 4.10, there was turbidity. The turbidity in the iron system is best described as a very heavy precipitation when compared to the turbidity of the Cr3þ-lactate system. It seems that lactate is a weak metal binder as one might expect from the lack of another chelating functionality

520

Y. Z. HAMADA ET AL.


TABLE 5 Selected stability constants to show that lactate is a weak metal binder compared to other hydroxycarboxylates

FIG. 5. pH titration curves of free lactic acid and Fe3þ/lactic acid in 1 : 1, 1 : 2, 1 : 3, 1 : 4, and 1 : 5 molar ratios with 0.4922 M NaOH. [Fe3þ] ¼ 2.5 mM. Turbidity started around pH 4.5. The equivalent points were at 0.84, 0.98, 1.13, 1.23, and 1.38 equivalents but never two equivalents.

that is present in malate or in citrate. See Table 5 for stability constants and pKa values taken from the literature that support this contention (Martell and Smith, 2001; Hamada et al., 2003b). IR Spectroscopy The IR spectrum for the bluish green Cr3þ/lactate complex or complexes was compared to that of the free lactic acid obtained from the National Institute of Standards and TABLE 4 Titration data for the Fe3þ-lactate, 258C, [Fe3þ] ¼ 2.50 mM M:L ratio

mL NaOH

Eq.a

0 : 1b

0.50

0.93 + 0.01

1:1 1:2 1:3

1.70 2.00 2.30

0.84 + 0.05 0.98 + 0.05 1.13 + 0.03

1:4

2.50

1.23 + 0.05

1:5

2.80

1.38 + 0.07

a,b

Comments Lactic is a monoprotic acid 1 Hþ released 1 Hþ released A little more than 1.0 Hþ released More than 1.0 Hþ released More than 1.0 Hþ released but not 2.0 Hþ

See Table 3 for definition of the equivalents.

Metal/ligand

Lactate

Malate

Citrate

pKa1 pKa2 pKa3 Log K for the 1 : 1 Mg2þ complexa Co2þ Ni2þ Cu2þ Al3þ Cr3þ Fe3þ

3.66 — — 0.93

4.68 3.24 — 1.71

5.64 4.35 2.90 3.43

1.38 1.64 2.54 2.36 — —

2.86 3.17 3.64 4.60 — 7.1

4.90 5.18 — 8.04b — 11.19

a pKa’s and stability constant (log K values for the 1 : 1 metal complexes) taken from (Martell and Smith; 2001). b see (Hamada et al., 2003b) for more confirmation. The missing metal-ligand stability constants are due to the fact that no reliable stability constants have been reported in the literature.

Technology (NIST) (NIST, 2004). The free lactic acid showed a sharp and strong band at 1150 cm21 for the C-O stretching, a sharp and strong band at 1795 cm21 for the carbonyl C55O of the carboxylate, and a medium band between 3500 – 3600 cm21 for the OH groups averaged for both the carboxylate and the alcoholic groups. These three main bands were observed in the Cr3þ/lactic acid complex synthesized, but all were shifted to lower energy bands due to lactate binding to Cr3þ. The bond stretching frequencies for the Cr3þ/lactate appeared at 1131 cm21, 1736 cm21, and broad 2426 cm21, respectively. These IR data confirm only that lactate is binding to the metal ion. We are not able to confirm whether the lactate is mono or a bi-dentate based solely on these IR data. The IR also cannot tell us if the oxygen of the carbonyl is bound to the metal ion or if it is the oxygen of the acid hydroxyl group that is bound to the metal ion. In any event, the IR data provide further confirmation that the solid colored material precipitate was not the pure metal hydroxide, but rather the metal/lactate chelate(s). Furthermore, the IR proved that the solid colored material precipitate was not the sodium lactate [see Equation (1)]. The microcrystalline product that was isolated and further used in the IR investigations is definitely for the Cr(III)lactate species and is not for sodium lactate. This is true for three reasons. (1) We already showed that the sodium lactate was shown to be a colorless material from the first titration curves for the free lactic acid in Figures 3 and 5. No precipitation whatsoever happened when we added NaOH to lactic acid alone, and lactic acid has changed into aqueous sodium lactate, according to Equation (1), that is in contrast to the


TABLE 6 Some IR bands in cm for the free lactic acid,a Cr3þ/lactate, Cr(NO3)3 . 9H2O, and Fe(NO3)3.9H2O 21

Free lactic acid

Cr3þ/ lactate

1131, C – O 1375 1795, C55O 1736, C55O 3500– 3600, 2426, O –H O–H

Cr(NO3)3 . 9H2O Fe(NO3)3 . 9H2O

1150, C –O

1385

1383

3400, H2O

3400, H2O

a


Bands for the free lactic acid were taken from (NIST, 2004).

Cr3þ/lactate complex. (2) The isolated material had a bluish green color that is not related to sodium lactate because it is colorless. (3) The Cr3þ/lactate microcrystalline material had the metal signature band from the IR spectrum for the Cr3þ/ lactate complex that appeared at 1375 cm21. See Table 6 for more details. All of the IR bands are shifted to lower frequencies compared to those of free lactic acid, denoting the changes in the vibrational state of lactate upon complexation to chromium. CONCLUSION Malate as a chelating ligand solubilized both Cr3þ and Fe3þ at physiological pH. In contrast to malate, the lactate ligand did not solubilize either Cr3þ or Fe3þ at physiological pH. The Cr3þ and the Fe3þ malate systems formed at least a hydrolytic metal chelate or the ternary hydroxo ligand chelate(s), or perhaps higher oligomer. This prediction is based on the number of protons observed, which was four and above and never two exactly. If one might suggest that the two carboxylates of the malate are the only chelating functional groups to be bound to the metal ion, let he/she come up with an intelligent explanation on why malate will favor the sevenmembered chelating ring over the more stable five and sixmembered chelating rings proposed in Scheme 2. The chemistry literature does not contain the UV-Vis spectra, the IR spectra, and the potentiometric titrations of the Cr3þ/lactate system, or the potentiometric titrations of

521

the Cr3þ/malate system. By examining Figure 4 for the lower energy peak at 580 nm, which corresponds to the 4 A2 g ! 4T2 g energy transition, the corresponding 10Dq value is 17,241 cm21. These values are in an agreement with the literature values (Douglas et al., 1994; Kettle, 1996; NIST, 2004; Huheey et al., 1993; Shriver and Atkins, 1999). At pH 4.73 and from the high energy peak at 400 nm using Beer’s law (A ¼ 1 c l), in which (A) is absorbance, (1) is the molar extinction coefficient, (c) is the molar concentration, and (l) is the optical path length, one might calculate (1) for the Cr3þ/lactate system at 400 nm as 1400 ¼ 25 M21 cm21. At pH 4.73 and from the lower energy peak at 580 nm, 1580 ¼ 18 M21 cm21. We hope that these new data will shed some light on the identity of the species formed. For the 1 : 2 titration systems in both the Cr3þ and Fe3þ malate systems, there is an extra mole of malate in these systems. There are two proposals for the observation of a sharp inflection at 6.30 equivalents of titrant: first, the formation of a tri-nuclear species as the predominant species at which there are three metal ions per species that release a net of 1/3 of a proton per metal ion, this has been confirmed in Martell’s database and has been observed the literature (Martell and Smith, 2001; Hamada et al., 2003b; Feng et al., 1990). Second, there is a possibility for the formation of a ternary metal-chelate complex [M3þ (Mal3-) (OH2)]2 in which a net of 4.0 protons are released in addition to the extra two carboxylate protons from the extra free malate. Additional detailed speciation studies are needed to sort out which proposal will be more reliable. The Cr3þ and the Fe3þ lactate systems formed the simple one-to-one metal chelates because only one proton was observed in all metal to lactate molar ratios with both the Cr3þ and the Fe3þ metal ions, even when we increased the lactate to metal ratio to 5 fold excess, i.e., in the 1 : 5 titration systems. It seems that the presence of one additional carboxylate enhanced the solubilizing effect of malate over lactate, which has only one carboxylate and one alcoholic hydroxyl. The presence of an additional carboxylate also helped to generate a complex or complexes with anionic net charge(s), which helped the metal-malate to be aqueous soluble throughout the entire pH range. In both the malate and the lactate systems, the chelate effect is predominant and the carboxyl(s) and the hydroxyl group have participated in the metal chelation that has been observed extensively in the literature (Feng et al., 1990; Quiros et al., 1992; Martell and Motekaitis, 1984; Salifoglou et al., 1998; Zhao-Hui et al., 1997; Salifoglou et al., 2003; Kiss et al., 1994; Ohman and Marlund, 1990). REFERENCES

SCH. 2. The proposed binding modes of lactate and malate for Cr3þ and Fe3þmetal ions in aqueous solutions.

Broderic, W. E.; Legg, J. I. Synthesis of the first stable nitrogencoordinated nicotinic acid – chromium(III) complexes: cis- and trans-H[Cr(mal)2(nic-N)2]. Inorg. Chem. 1985, 24, 3724– 3725. Buffie, J. C.; Juliet, E.; Dion, D. D. H.; Chakov, N. E.; Nettles, H. S.; Vincent, J. B. The trail of chromium(III) in vivo from the blood to


522

Y. Z. HAMADA ET AL.

the urine: the roles of transferrin and chromodulin. J. Biol. Inorg. Chem. 2001, 6, 608– 617. Burtis, C. A.; Ashwood, E. R. Tietz Textbook of Clinical Chemistry, 2nd ed.; Burtis, C. A., Ashwood, E. R., eds.; Philadelphia, PA: Saunders, 1994. Chang, J. C.; Gerdom, L. E.; Baenziger, N. C.; Goff, H. M. Synthesis and molecular structure determination of carboxy-bound nicotinic acid (Niacin) complexes of chromium. Inorg. Chem. 1983, 22, 1739 –1744. Douglas, B.; MacDaniel, D.; Alexander, J. Concepts and Models of Inorganic Chemistry, 3rd ed.; John Wiley & Sons Inc.: New York, 1994. Feng, T. L.; Gurian, P. L.; Healy, M. D.; Barron, A. R. Aluminum citrate: Isolation and structural characterization of stable trinuclear complex. Inorg. Chem. 1990, 29, 408– 411. Finholt, J. E.; Thompson, M.; Connick, R. Hydrolytic polymerization of chromium(III). A trimeric species. Inorg. Chem. 1981, 20, 4151 –4155. Green, C. A.; Bianchini, R. J.; Legg, J. I. Characterization of a stable chromium(III)-nicotinic acid complex by deuteron NMR. Inorg. Chem. 1984, 23, 2713– 2715. Hamada, Y. Z.; Carlson, B. L.; Shank, J. T. Potentiometric and UV-vis spectroscopy studies of citrate with the hexaquo Fe3þ and Cr3þ metal ions. Syn. Reac. Inorg. Metal-Org. Chem. 2003a, 33 (8), 1425 –1440. Hamada, Y. Z.; Zhepeng, W.; Harris, W. R. Competition between transferrin and serum ligands citrate and phosphate for the binding of aluminum. Inorg. Chem. 2003b, 42, 3262– 3273. Huheey, J.; Keiter, E.; Keiter, R. Inorganic Chemistry Principals of Structure and Reactivity, 4th ed.; New York, NY: Harper Collins College Publishers, 1993. Kettle, S. F. A. Physical Inorganic Chemistry, A Coordination Chemistry Approach, Spektrum; University Science Book: Sausalito, CA, 1996. Kiss, T.; Sovago, I.; Martin, B. R.; Pursiainen, J. Ternary complex formation between Al(III)-Adenosine-50 -phosphate and carboxylic acid derivatives. Journal of Inorganic Biochemistry 1994, 55, 53 – 65, and references therein. Levina, A.; Lay, P. A.; Dixon, N. E. Disproportionation of a model chromium(V) complex causes extensive chromium(III)-DNA binding in vitro. Chem. Res. Toxicol. 2001, 14, 946–950 and references therein. Martell, A. E.; Motekaitis, R. J. Complexes of aluminum(III) with hydroxy carboxylic acids. Inorg. Chem. 1984, 23, 18 – 23. Martell, A. E.; Smith, R. M. Critical Selected Stability Constants of Metal Complexes Database, Version 6.0 for Windows; National Institute of Standards and Technology, April 2001. National Institute of Standards and Technology (NIST), http:// webbook.nist.gov/chemistry, March 1, 2004.

Ohman, L.-O.; Marlund, E. Equilibrium and structural studies of silicon(IV) and aluminum(III) in aqueous solutions. 24 A Potentiometric and 27Al NMR study of polynuclear aluminum(III) hydroxo complexes with lactic acid. Acta Chem. Scand 1990, 44, 228– 234. Quiros, M.; Goodgame, D. M.; Williams, D. J. Crystal structure and EPR spectrum of bispyridinium bis(citrato)chromium(III) tetrahydrate. Polyhedron 1992, 11 (11), 1343– 1348. Rosenthal, P.; Pesca, M. A. Long-term monitoring of D-Lactic acidosis in a child. J. Pediatr Gastroenterol Nutr. 1985, 4, 674– 676. Salifoglou, A.; Matzapetakis, M.; Raptopoulou, C. P.; Tsohos, A.; Papaefthymiou, V.; Moon, N. Synthesis, spectroscopic and structural characterization of the first mononuclear, water soluble iron-citrate complex, (NH4)5Fe(cit)2 . 2H2O, J. Am. Chem. Soc. 1998, 120, 13266– 13267. Salifoglou, A.; Raptopoulou, C. P.; Dakanali, M.; Terzis, A.; Lakatos, A.; Banyai, I.; Kiss, T. A novel dinuclear species in the aqueous distribution of aluminum in the presence of citrate. Inorg. Chem. 2003, 42, 252– 254. Shriver, D.; Atkins, P. Inorganic Chemistry, 3rd edn.; W. H. Freeman and Company: New York, 1999. Stunzi, H; Marty, W. Early stages of the hydrolysis of chromium(III) in aqueous solution. Characterization of a tetrameric species. Inorg. Chem. 1983, 22, 2145– 2150. Sun, Y.; Ramirez, J.; Woski, S. A.; Vincent, J. B. The binding of trivalent chromium to low-molecular-weight chromium-binding substance (LMWCr) and transfer of chromium from transferrin and chromium picolinate to LMWCr. J. Bioinorg. Chem. 2000, 5, 129– 136. Svahn, J.; Schiaffino, M.; Caruso, U.; Calvillo, M.; Minniti, G.; Dufour, C. Severe lactic acidosis due to thiamine deficiency in patient with B-cell leukemia/lymphoma on total parental nutritio during high dose methotrexate therapy. Journal of Pediatric Hematology/Oncology 2003, 25 (12), 965– 968. Thompson, M.; Connick, R. Hydrolytic polymerization of chromium(III). Two dimeric species. Inorg. Chem. 1981, 20, 2279–2285. Vincent, J. B. Elucidating a biological role for chromium at a molecular level. Acc, of Chemical Research 2000, 33, 503– 510. Vincent, J. B. The bioinorganic chemistry of chromium(III). Polyhedron 2001, 20, 1– 26, references therein. Voet; Voet; Pratt. Fundamentals of Biochmistry; John Wiley & Sons, Inc.: New York, 1999. Zhao-Hui, Z.; Yi-Ji, L.; Hong-Bin, Z.; Guo-Dong, L.; Khi-Rui, T. Syntheses, structures and spectroscopic properties of nickel(II) citrate complexes, (NH4)2[Ni(Hcit)(H2O)2]2 . 2H2O and (NH4)4 [Ni(Hcit)2] . 2H2O. J. Coord. Chem. 1997, 42, 131– 141.