CHAPTER 5

BIOELECTROCHEMISTRY

 

5.1 Introduction

The roots of electrochemistry can be traced to Egyptian and Roman physicians who used the discharges of electric eels as a method of treatment. One of the first bioelectrical observations was made when Galvani’s assistant touched the crural nerve of a dissected frog with his scalpel, and the frog’s limbs contracted. In 1890, Ostwald recorded observations of potential differences across semipermeable membranes.1

In 1902, Bernstein derived the fundamental relationship between the differing alkali-ion concentrations which are found to be maintained across a membrane (Figure 41).2,3 Typically, a biological cell contains 25 times more K+ inside than is on the outside. The Na+-K+ pump is orientated so that it pumps Na+ out of the cell and K+ into the cell. ATP located on the inside of the pump drives the system. Bernstein applied the only approach known then to the "theory of electrochemical potentials" Nernst’s theory of electrode potential differences for systems at thermodynamic equilibrium. Decades of observations concerning membranes potentials followed, and bioelectrochemisity developed as an integral facet of the biomedical sciences (Table 3).4

A surface may acquire a charge in four ways:

1. imposition of a potential difference from an external potential source

2. adsorption of ions on a solid surface

3. electron-transfer between a metallic conductor and the solution

4. for micelles, biological macromolecules and membranes, ionization of functional groups such as carboxylate, phosphate, or amino acids.

5.2 Semiconductors

Optically transparent semiconductor electrodes of doped tin oxide and indium oxide have been used successfully in the study of biological electron-transfer reactions.5

The electronic structure of semiconductors is depicted in Figure 42. Electrons may be present at energy levels within the conduction or valance bands; however, electrons will not be found at energy levels within the band gap (Eg). The valence band is usually completely filled with electrons, and thermal fluctuations within intrinsic semiconductors tend to populate the typically vacant conduction band with electrons. The Fermi level, located in the middle of the band gap, is defined as the electronic energy level with 50% probability of being occupied. Ed represents electron donors in n-type semiconductors and Ea represents electron acceptors in p-type semiconductors. The charge and potential distribution in the semiconductor/aqueous electrolyte interfacial region is dependent upon factors including: electrolyte concentration, interactions with water, and pH. The potential of zero charge (PZC) is the unique value of electrode potential at which excess charge is zero. PZC data depends on the chemical nature of electrode surfaces and wide variations in pzc values for electrodes of the same metal in identical solutions have been reported.6 The structure and potential profile at the semiconductor/electrolyte interface can have a significant impact on the analysis and interpretation of electrobioanalytical experiments.

5.3 Biological Membranes

The biological cell is enclosed by a membrane, the thickness of which is ~70-115Å. This membrane consists of a bimolecular layer of phospholipids. The fluid mosaic model for the overall organization of biological membranes was proposed in 1972 by Jonathan Singer and Garth Nicolson (Figure 43).6,7 Oriented globular proteins and lipids are distributed throughout this membrane in a bilayer arrangement with the alkyl chains pointing inward and the lipid head groups pointing outward. This lipid bilayer serves as a solvent for integral membrane proteins and as a permeability barrier. Membrane proteins have the capability of lateral diffusion through the lipid matrix unless otherwise influenced by special interactions. Furthermore, membranes are structurally and functionally asymmetric.8

Figure 44 illustrates a membrane model with two primary types of proteins: integral and peripheral.5 Peripheral proteins are weakly associated with the biological membrane. On the other hand, integral proteins are embedded in the protein and may have protruding ionic regions. An electrical double-layer is created in aqueous solution group as a result. The behavior of charged species near the membrane surface is analogous to the Helmholtz region at the electrode/aqueous electrolyte interface.

5.4 Stability of a Redox Protein

The stability of a redox protein is governed by the following characteristics: interaction of the metal ion with coordination ligands; dipoles surrounding and within the protein; buried charges; surface charges; ions in solution; and ligands which bind to the protein (Figure 45).9

Interpretation of D Hstab and D Sstab contributions are not clear.

D G = - nFE

The Gibbs free energy, D G, is the thermodynamic criterion for determining the spontaneity of a reaction. When D G is negative, the reaction is spontaneous and when D G is positive the reaction is non-spontaneous. This equation relates D G and the standard electromotive force, E, of a reaction (the difference in potential between two half cells.)

5.5 Reduction Potentials

The tendency of a redox couple to donate or accept electrons is given by the redox potential (E) and is measured as a reduction potential with reference to a standard hydrogen electrode. Figure 46 is an illustration of the formal reduction potentials of the participants in the oxidative phosphorylation chain.9 Cytochrome c has a formal potential of +260 mV vs NHE. The value of E exhibited by an individual redox couple is a reflection of the relative stability of the reduced and oxidized states. Any factor which acts to stabilize the oxidized form makes the redox couple a better electron donor and results in a more negative redox potential. Conversely, any factor which acts to stabilize the reduced form makes the couple a better electron acceptor and gives rise to a more positive redox potential (Figure 47).10 Experimental contributions including temperature, pH, ionic strength, type of electrolyte, and type of buffer also tend to affect the formal potential of a redox couple.11, 12, 13, 14

5.6 Cytochrome c

The study of heterogeneous electron-transfer reactions involves the study of electron transfer between a solid protein and a substance in solution. Cytochrome c is the most extensively studied biological redox molecule. In some elegant experiments conducted in 1977, voltammograms were obtained of horse heart cytochrome c using a Sn-doped In2O3 electrode.15 Proteins react very specifically with the electrode surface. Furthermore, these proteins decompose when interacting with normal metal electrodes like platinum. A combination of electrostatic and chemical interactions in the electrode/solution interface determine the rate of heterogeneous electron-transfer. The cytochrome c voltammograms indicated diffusion-controlled electrochemistry because the peak separation was about 60 mV, and the faradaic current was proportional to the v1/2. The rate constant was later estimated to be about 50s-1 which is "fast" considering that cytochrome c is an enormous protein on the scale of normal redox ions. In other words, electron transfer to cytochrome c would be expected to be slow compared to ions having radii 15-20 times smaller.

Figure 48 is a typical cyclic voltammogram of horse heart cytochrome c at a gold electrode, modified by adsorption with 4,4’-bipyridyl. The peak separation is approximately 60 mV, and faradaic currents vary linearly with (scan rate) 1/2 indicative of electrochemical reversibility and planar diffusion of cytochrome c to and from the electrode surface. An illustration of the electrode-protein complex for the idealized limiting case of a planar surface for the cyclic voltammogram in Figure 48 is depicted in Figure 49.16 In this schematic, bonding between 4,4’-bipyridyl and the lysine groups of cytochrome c facilitates stabilized electron-transfer between the gold-bipyridyl electrode and the heme group in the protein.

Taniguchi et al.17 later showed that bases like 4-pyridyl disulfide were good modifiers which promoted electron-transfer with cytochrome c. Subsequent studies of the interfacial electrochemistry of cytochrome c at tin oxide, indium oxide, gold, and platinum electrodes revealed that electron transfer in cytochrome c is directly dependent on the state of the electrode surface.18 Furthermore, minimal activity has been associated with a ‘fully hydrated’ surface.

5.7 Myoglobin

The electrochemical behavior of myoglobin, studied particularly at bare metallic electrodes including mercury,19 gold,20 platinum,21 and silver22 has been characterized as a quasi-reversible response with reaction rates of the oxidation being much slower than the reduction. Electron transfer in myoglobin involves the sixth coordination position of the heme iron and proceeds via a plane perpendicular to the heme ring.23 Moreover, the site of electron transfer is buried with respect to the protein surface (Figure 50).24, 25, 26, 27

Myoglobin/oxygen ligand-binding reactions28, 29, 30, 31, 32, 33 have been studied directly at indium oxide transparent electrodes (Figure 51).34 The reaction mechanism is an EC mechanism (an electrode reaction followed by a chemical reaction that consumes the product of the electrode reactions) as shown below:

Electrode: Mb(III) + e- Û Mb(II)

Solution: Mb(II) + O2 Û Mb(II)O2

where Mb(III) and Mb(II) are the oxidized (metmyoglobin) and reduced forms of myoglobin.35 Myoglobin only binds to oxygen when in the reduced state.

5.8 References

1. Ostwald, W. Phys. Chem. 1890, 6, 71.

2. Bernstein, J. Pfleugers Arch. Ges. Physiol. Menschen Tiere 1902, 92, 521.

3. Stryer, L. Biochemistry. New York: W.H. Freeman and Company, 1988.

4. Srinivasan, S.; Cahen, G.L.; Stoner, G.E. Electrochemistry: The Past 30 Years, The Next 30 Years. New York: Plenum Press, 1975.

5. Bowden, E.F.; Hawkridge, F.M.; Blount, H.N. Comprehensive Treatise of Electroanalytical Chemistry. Vol. 10 (S. Srinvasan, Y.A. Chizmadshev, J.O’M. Bockris, B.E. Conway, E. Yeager, Eds.) New York: Plenum Press, 1985.

6. Singer, S.J.; Nicolson, G.L. Science. 1972, 175, 723-731.

7. Koryta, J.; Dvorak, J. Principles of Electrochemistry. New York: John Wiley & Sons, 1987.

8. Bockris, J. O’M.; Khan, S.U.M. Surface Electrochemistry: A Molecular Level Approach. New York: Plenum Press, 1993.

9. Pettigrew, G.W.; Moore, G.R. Cytochromes c: Evolutionary, Structural and Physiochemical Aspects. Berlin: Springer-Verlag, 1990.

10. Bard, A.J.; Faulkner, L.R. Electrochemical Methods: Fundamentals and Applications. New York: John Wiley & Sons, 1980.

11. Wilson, M.T.; Greenwood, C. Eur. J. Biochem. 1971, 22, 11-18.

12. Margalit, R.; Schejter, A. Eur. J. Biochem. 1973, 32, 492-499.

13. Koller, K.B.; Hawkridge, F.M. J. Electroanal. Chem. 1988, 239, 291-306.

14. Sun, S.; Reed, D.E.; Hawkridge, F.M. Redox Chemistry and Interfacial Behavior of Biological Molecules. (G. Dryhurst and K. Niki Eds.) New York: Plenum Publishing Corporation, 1988.

15. Yeh, P.; Kuwana, T. Chem. Lett. 1977, 93, 1145.

16. Armstrong, F.A.; Hill, H.A.O.; Walton, N.J. Acc. Chem. Res. 1988, 21, 407-413.

17. Taniguchi, I.; Toyosawa, K.; Yamiguchi, H.; Yasukouchi, K.J. J. Chem. Soc., Chem. Commun. 1982, 1032-1033.

18. Bowden, E.F.; Hawkridge, F.M.; Blount, H.N. J. Electroanal. Chem. 1984, 161, 355-376.

19. Scheller, F.; Janchen, M.; Lampe, J.; Prumke, H.J.; Blanck, J.; Palecek, E. Biochim. Biophys. Acta. 1975, 412, 157-167.

20. Stargardt, J.F.; Hawkridge, F.M.; Landrum, H.L. Anal. Chem. 1978, 50, 930-932.

21. Song, S.; Dong, S. Bioelectrochem. Bioenerg. 1988, 19, 337-346.

22. Cotton, T.M.; Schultz, S.G.; van Duyne, R.P. J. Am. Chem. Soc. 1980, 102, 7960- 7962.

23. Castner, J.E.; Hawkridge, F.M. J. Electroanal. Chem. 1983, 143, 217-232.

24. Stellwagen, E. Nature. 1978, 275, 73-74.

25. Schlereth, D.D.; Mantele, W. Biochemistry. 1992, 31, 7494-7502.

26. Austin, R.H.; Beeson, K.W.; Eisenstein, L.; Frauenfelder, H.; Gunsalus, I.C. Biochemistry. 1975, 14, 5355-5372.

27. Lambright, D.G.; Balasubramanian, S.; Decatur, S.M.; Boxer, S.G. Biochemistry. 1994, 33, 5518-5525.

28. Gibson, Q.H.; Regan, R.; Olson, J.S.; Carver, T.E.; Dixon, B.; Pohajdak, B.; Sharma, P.K.; Vinogradov, S.N. J. Biol. Chem. 1993, 268, 16993-16998.

29. Nienhaus, G.U.; Mourant, J.R.; Chu, K.; Frauenfelder, H. Biochemistry. 1994, 33, 13413-13430.

30. Miller, L.M.; Chance, M.R. Biochemistry. 1995, 34, 10170-10179.

31. Brantley, R.E.; Smerdon, S.J.; Wilkinson, A.J.; Singleton, E.W.; Olson, J.S. J. Biol. Chem. 1993, 268, 6995-7010.

32. Dong, A.; Huang, P.; Caughey, B.; Caughey, W.S. Arc. Biochem. Biophys. 1995, 316, 893-898.

33. Antonini, E.; Brunori, M.; Hemoglobin and Myoglobin In Their Reactions With Ligands. New York: North Holland Publishing Company, 1971.

34. King, B.C.; Hawkridge, F.M. J. Electroanal. Chem. 1987, 237, 81-92.

35. King, B.C.; Hawkridge, F.M. Talanta. 1989, 36, 331-334.