In the Laboratory
Complexation Effect on Redox Potential of Iron(III)-Iron(II) Couple: A Simple Potentiometric Experiment Masood Ahmad Rizvi,* Raashid Maqsood Syed, and Badruddin Khan Department of Chemistry, University of Kashmir, Hazratbal, J&K, India *
[email protected] When a mixture of two or more reducing agents is titrated, a curve with multiple inflection points results provided the reduction potentials of the species are sufficiently different (1). This is comparable to the titration of two acids with different dissociation constants or of two ions forming precipitates of different solubilities with same reagent (2). The effect of complexation on redox potentials is well documented (3-19) and some articles to this effect are published in this Journal (3, 4). Our experiment is a simple potentiometric titration aimed to invoke the underlying concept of the complexation effect on redox potential of a redox couple. Potentiometric titrations have been widely discussed in this Journal (20-24); however, no method is reported for potentiometric titration of metal ions on the basis of the ligand effect. We choose iron(III)-iron(II) couple because it occurs in many natural redox systems. In this experiment, three iron(II) octahedral complexes, iron(II) aqua, [Fe(OH2)6]2þ; iron(II) EDTA, [Fe EDTA]2-; and iron(II) phenanthroline [Fe(o-phen)3]2þ, are combined into mixtures of reducing agents and are oxidized to the corresponding iron(III) complexes. Multiple inflections in the titration curve depict the different redox potential of iron(III)-iron(II) couple in these complexes that can be attributed to the complexation effect. This experiment illustrates the effect of the stabilities of iron(III)-iron(II) complexes with EDTA, water, and 1-10 phenanthroline ligands on redox potential. Experimental Section Reagents All solutions were prepared in double distilled water from the analytical-grade chemicals. A cerium(IV) stock solution, 2.00 10-2 M, was prepared by dissolving 1.27 g of cerium(IV) tetraammonium sulfate, (NH4)4Ce(SO4)4 3 2H2O (J. T. Baker), in 100.00 mL of 1 M sulfuric acid. A commercially supplied iron(II) phenanthroline complex, 2.50 10-2 M [Fe(o-phen)3](SO4) (a redox indicator solution from Qualigens) was used as purchased. Ammonium ferrous sulfate, (NH4)2Fe(SO4)2 3 6H2O ( J. T. Baker) and the disodium salt of EDTA (Merck) were used as solids. Apparatus Potentiometric titrations were performed manually using commercially available platinum and saturated calomel reference electrodes at 25 °C under a nitrogen atmosphere. The titration vessel consisted of a self-designed six-hole rubber stopper (for the microburet, platinum and calomel electrodes, inlet and outlet for nitrogen gas, and a temperature probe) tightly fitted in a tall form of 100.00 mL beaker (22). All potential measurements were 220
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Figure 1. Titration assembly used in the experiment.
performed under nitrogen atmosphere over a magnetic stirrer in a self-fabricated assembly shown in Figure 1. Procedure for Titration of the Iron(II) Complex Mixtures with Cerium(IV) Binary (A) and a ternary (B) mixtures of iron(II) complexes were used in the analysis. Mixture A had a iron(II) aqua complex, [Fe(OH2)6]2þ, and a iron(II) phenanthroline complex, [Fe(ophen)3]2þ, prepared by dissolving 0.16 g of ammonium ferrous sulfate in 16.00 mL of pH 4.00 buffer to which 4.00 mL of 2.5 10-2 M iron(II) phenanthroline complex was added. Ternary mixture B had the iron(II) EDTA complex in addition to mixture A and was prepared (25) by dissolving 0.16 g of ammonium ferrous sulfate and 0.076 g of Na2EDTA in 20.00 mL deareated solution of 4.00 mL of 2.5 10-2 M iron(II) phenanthroline and 16.00 mL of pH 4.00 buffer. The mixtures were titrated with a standard 2 10-2 M cerium(IV) solution and the potentials recorded. A pH variation of mixtures was also established by titrating them in solutions of different initial pH values. Hazards Protective equipment (goggles, lab coat) should be worn while handling corrosive sulfuric acid in the fume hood. Contact
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Vol. 88 No. 2 February 2011 pubs.acs.org/jchemeduc r 2010 American Chemical Society and Division of Chemical Education, Inc. 10.1021/ed100339g Published on Web 11/23/2010
In the Laboratory
Figure 2. Titration of 20 mL solution containing 2 10-2 M iron(II) aqua and 5 10-3 M iron(II) phenanthroline with 2 10-2 M cerium(IV) at pH = 4.0.
Figure 4. The effect of pH on the titration curve of the ternary mixture.
initial portion of the curve prior to the first equivalence point illustrates complexation of iron(II) due to EDTA at pH of 6.50 and 4.00. At pH 0, the effect of EDTA on initial solution potential vanished consistent with ineffective complexation at this low pH value (Figure 4). The results obtained for two mixtures can be explained on the basis of thermochemical cycle (29):
Figure 3. Titration of 20mL solution containing 1 10-2 M iron(II) EDTA, 1 10-2 M iron(II) aqua and 5 10-3 M iron(II) phenanthroline with 2 10-2 M cerium(IV) at pH = 4.0.
with sulfuric acid may cause burns to the skin and eyes. Iron(II) phenanthroline gives red stains on skin and is a mild irritant similar to EDTA. Ammonium cerium(IV) sulfate is harmful by inhalation, irritating to skin and eyes, and toxic if swallowed. Results and Discussion The redox property of a metal in a complex is remarkably different from its free state. In complexed form, the ability of metal ion to get oxidized or reduced is determined by the overall tendency of the complex for such process (4, 6, 17). The titration curve for mixture A (Figure 2) displayed two inflection points indicating presence of two types of reducing agents. An interesting feature is the solution color; the initial red color of iron(II) phenanthroline complex does not change until the first jump and after becomes initially darker and finally blue at the second jump indicating iron(II) phenanthroline complex is not oxidized until all the iron(II) aqua is oxidized. The two inflections correspond to the stoichoimetric amounts of the two complexes that allow the method to be used in quantitative analysis of this binary mixture. The titration curve for mixture B (Figure 3) displayed two inflections and a steep rise in the region of iron(II) EDTA. An interesting feature was the significant lowering of initial solution potential with the addition of EDTA. The steepness of
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where Ecomplex represents the redox potential of the metal complex with a ligand other than water, Eaqua represents the redox potential of [M(OH2)6]3þ-[M(OH2)6]2þ redox couple, βIII is the overall thermodynamic stability constant for the metal(III)-ligand complex and βII is the overall thermodynamic stability constant for the metal(II)-ligand complex. As can be seen from the last equation, complexes that favor the stabilization of bound metal in the 3þ oxidation state will lower the
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reduction potential relative to that of the aqueous form, whereas complexes that stabilize the bound metal in 2þ oxidation state will raise the reduction potential relative to that of the aqueous form The formation constant of iron(III) EDTA is nearly 1010 larger than that of iron(II) EDTA, whereas as formation constant of iron(II) phenanthroline is approximately 105 times larger than that of iron(III) phenanthroline complex (2, 4). Iron(II) corresponds to 3d6 and iron(III) corresponds to 3d5 configurations and phenanthroline is a strong field π acceptor ligand (26, 27). The high formation constant, coupled with lower charge, and t2g6 configuration of octahedral (26, 27) iron(II) phenanthroline complex favors more electron delocalization to the vacant orbitals of ligand than in corresponding iron(III) phenanthroline complex (4). This ability stabilizes iron(II) phenanthroline complex toward oxidation relative to iron(II) aqua and lowers its oxidation potential. The higher formation constant of iron(III) EDTA compared to iron(II) EDTA gives a strong tendency of iron(II) EDTA to oxidize to the corresponding iron(III) EDTA thereby raising its oxidation potential considerably relative to iron(II) aqua and, obviously, iron(II) phenanthroline as well. Conclusion Iron(III) is stabilized by the ligands studied in the order EDTA > OH2 > phenanthroline. This gives the oxidation tendency and oxidation potentials of iron(II)-iron(III) to decrease in the same order, hence, giving multiple inflections in titration curve. The ligand effect can thus tune the potential of a redox couple, which can be explored in design of novel redox systems applicable in titrimetric analysis and help to understand the tuning of potentials in nature (9, 19, 28). The experiment allows students to calculate formal reduction potential of same couple, iron(III)-iron(II), in different complexes (30). Acknowledgment The authors wish to thank Prof. Khaleequz Zaman chairman Department of Chemistry, University of Kashmir and dedicate this article to ideal teacher Mohammad Afaq. Literature Cited 1. Serjeant, E. P. Potentiometry and Potentiometric Titrations. In A Series of Monographs on Analytical Chemistry and Its Applications, Vol. 69; Elving, P. J., Winefordner, J. D., Eds.; John Wiley & Sons: New York, 1984; p 38. 2. Harris, D. C. Quantitative Chemical Analysis, 5th ed.; W. H. Freeman and Company: New York, 1999; pp 161, 313-315, 421. 3. Helsen., J. J. Chem. Educ. 1968, 45, 518.
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4. Ibanez, J. G.; Gonzalez, I.; Cardenas, M. A. J. Chem. Educ. 1988, 65, 173. 5. Ringbom, A. Complexation in Analytical Chemistry; Interscience Publishers: New York, 1963; p 274. 6. Shriver, D. F.; Atkins, P. W. Inorg. Chem., 3rd ed.; Oxford University Press: Oxford, 1999; p 207. 7. Phillips, C. S. G.; Williams, R. J. P. Inorg. Chem., Vol II; Oxford University Press: Oxford, 1966; p 312. 8. Greenwood, N. N.; Earnshaw, A. Chemistry of Elements, 2nd ed.; Butterworth Heinemann: Oxford, 1997; p 1093. 9. Lippard, S. J.; Berg, J. M. Principles of Bioinorganic Chemistry, 2nd Indian Reprint; Panima Publishing Corporation: New Delhi, 2005; p 26. 10. Vydra, F.; Pribil, R. Talanta 1959, 3, 103. 11. Vydra, F.; Pribil, R. Talanta 1960, 5, 44. 12. Krishna, M. N.; Pulla, R. Y. Anal. Chim. Acta 1974, 73, 413. 13. Krishna, M. N.; Pulla, R. Y. Indian J. Chem. 1976, 14A, 721. 14. Umetsu, K.; Itabashi, H.; Satoh, K.; Kawashima, T. Anal. Sci. 1990, 6, 721. 15. Umetsu, K.; Itabashi, H.; Satoh, K.; Kawashima, T. Anal. Sci. 1991, 7, 115. 16. Teshima, N.; Kawashima, T. Bull. Chem. Soc. Jpn. 1996, 69, 1975. 17. Katsumata, H.; Teshima, N.; Kawashima, T. Anal. Sci. 1997, 13, 825. 18. Kuwabara, M.; Katsumata, N.; Teshima, N.; Kurihara, M.; Kawashima, T. Anal. Sci. 1999, 15, 657. 19. Teshima, N.; Katsumata, H.; Kawashima, T. Anal. Sci. 2000, 16, 901. 20. Furman, N. H. J. Chem. Educ. 1926, 3, 932. 21. Goldberg, D. E. J. Chem. Educ. 1962, 39, 328. 22. Braun, R. D. J. Chem. Educ. 1976, 53, 463. 23. Powell, J. R.; Tucker, S. A.; William, E.; Jennifer, A.; Lindsey, H. J. Chem. Educ. 1996, 73, 984. 24. Barreto, M. S.; Medeiros, L. L.; Furtado, P. C. J. Chem. Educ. 2001, 78, 91. 25. Ibanez, J. G.; Clemente, J.; Pastor, J.; Elizabeth, G. Chem. Educator 2000, 5 (226), 230. 26. Cotton, F. A.; Wilkinson, G. Advanced Inorganic Chemistry, 3rd ed.; Wiley: New York, 1978; Chapters 20, 22, 25. 27. Graddon, D. P. An Introduction to Coordination Chemistry, 2nd ed.; Pergamon: Oxford, 1968; Chapter 5. 28. Geiger, K. D. J. Chem. Educ. 1991, 68, 340. 29. Harrington, J. M.; Crumbliss, A. L. BioMetals 2009, 22, 679–689. 30. Harvey, A. E., Jr.; Manning, D. L. J. Chem. Educ. 1951, 28, 527.
Supporting Information Available Student instructions; experimental data; derivations; calculations; report form; instructor notes. This material is available via the Internet at http://pubs.acs.org.
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