8.2.1: Fundamentals of Biological Redox Potential
- Page ID
- 207721
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In a redox (reduction-oxidation) reaction, electrons are transferred from an electron donor (or electron "source") to an electron acceptor (or electron “sink”). The electron source becomes oxidized while the electron sink becomes reduced. The tendency for molecules to gain electrons is measured as the reduction potential. The opposite, oxidation potential, can also be used to discuss redox reactions. However, in contemporary practice, we use reduction potential values and refer to them as redox potential (Eº) values. These values are related to thermodynamics and to the values of K and Gibbs free energy.
Standard reduction potential
A redox reaction can be viewed as two reduction half-reactions. The potential of the reduction half-reaction is used to determine whether a redox reaction is favorable or not. Each reduction half reaction is written with the thing being reduced (electron "sink") on the left and the thing being oxidized (electron "source") on the right.
For example the following redox reaction,
\(Zn + Cu^{2+} \rightarrow Zn^{2+} + Cu\)
can be viewed as two reduction half-reactions:
\(Zn^{2+} + 2e^- \rightarrow Zn\) Eº = − 0.76181 Volts
\(Cu^{2+} + 2e^- \rightarrow Cu\) Eº = + 0.3372 Volts
Spontaneous reactions have a Eºoverall > 0 according to the following equation:
Eºoverall = Eºthing being reduced - Eºthing being oxidized
In the equation above, Eº is the standard reduction potential; the reduction potential of the half-reaction at standard conditions.3 Note that standard conditions are those at standard temperature and pressure, and at pH = 0. These are not biological conditions!
Calculate the Eºoverall for the reaction of \(Zn\) and \(Cu^{2+}\). Is the reaction spontaneous at standard conditions?
- Answer
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Eºoverall = EºCu - EºZn
Eºoverall = (+ 0.337) - (-0.7618) = 1.099 V
The Eºoverall calculated at standard conditions is a positive value, thus this reaction is spontaneous.
Reduction potential at pH 7
Sources:
- Vanýsek, Petr (2012). "Electrochemical Series". In Haynes, William M. (ed.). Handbook of Chemistry and Physics: 93rd Edition. Chemical Rubber Company. pp. 5–80. ISBN 9781439880494.
- Bard, A. J., Parsons, R., and Jordan, J. (1985). Standard Potentials in Aqueous Solutions (Marcel Dekker, New York)
- Wikipedia, Standard electrode potential (data page). accessed Feb 23, 2010 at https://en.Wikipedia.org/wiki/Standa...te_note-Bar-10