CHEM 445 / BIOL 445
Biochemistry II

J. D. Cronk
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2. header

Lecture 2. Redox potentials and free energy.

Friday 19 January 2007

More on membrane-associated proteins. Redox potentials and free energy. Biochemical standard reduction potentials. The Nernst equation and calculation of free energy changes for redox reactions from the corresponding changes in reduction potential. Overview of electron transport chain.

Reading: BTS6, Ch.12, pp.336-347; Ch.18, pp.506-509. Recommended: Ch.13

 

2. Summary

Lecture 2 Summary

In today's class we discuss the free energy associated with an electrochemical gradient, as well as that associated with redox reactions. In the latter case, for each redox couple (an oxidant and its conjugate reductant) we measure electrochemically a quantity called the standard reduction potential and use a form of the Nernst equation to calculate standard free energy change (DG°') from a difference in standard reduction potential, DE°'.

Reduction potentials and free energy

The electron transfer reactions that constitute the electron transport chain are by definition oxidation-reduction ("redox") reactions. We can distinguish the species oxidized and the species reduced in redox reactions, and each reaction can be thought of as being composed of two "half-reactions", one a reduction in which electron(s) are a reactant, and the other in which electron(s) are a product. These are not real reactions, but they are useful in analysis of redox reactions. For a reduction half-reaction, a standard reduction potential Eo can be assigned, relative to a reference reaction, for reactants and products under standard conditions. The more positive Eo is for a given half-reaction, the greater the relative tendency toward the reduced products. This suggests a relationship between reduction potential and free energy.

 
Selected standard (biochemical) reduction potentials   As with free energy, the important quantities connected with reduction potentials are not absolute values, but differences. By summing two half-reactions and taking the corresponding algebraic sum of Eos, we can say whether the resulting net redox reaction can occur spontaneously. In the table, the half-reaction with the most positive Eo value is the reduction of O2 to H2O. This means that in combination with any of the other half-reactions, O2 will be the species reduced, and the other half-reaction will occur in the reverse direction.
As a particularly pertinent example, we obtain the change in standard reduction potential, DEo' for the redox reaction representing the reduction of O2 by NADH. (In the table above, V represents volts, and - as for free energy - the prime denotes a biochemical standard state where the pH = 7.0. The definition of the biochemical standard state can be found here.)
Redox reaction equation: NADH reduction of molecular oxygen to water, and the corresponding change in biochemical standard reduction potential
A positive value for DEo' represents a reaction that can proceed spontaneously, i.e. DGo' < 0. In fact, there is a quantitative relationship between DEo' and DGo', which can be worked out from the Nernst equation:
[ A review of redox chemistry, showing how the relation between K'eq and DEo' arises can be found here ]
Nernst equation showing that free energy change for a redox reaction is proportional to the change in reduction potential
This equation states that the standard free energy change for a redox reaction is equal to the negative of the product of the change in standard reduction potential times n, the number of electrons transferred in thereaction, times a constant F, which is the Faraday. F is equal to 23.06 kcal-mol-1-V-1. Thus, the free energy available under (biochemical) standard conditions from the reduction of O2 by NADH is  – (2)(23.06)(1.14)  =  – 52.6 kcal/mol.
 

Study questions

  • Describe how biochemical standard state differs from physical chemistry standard state.
  • Use a table of half-reactions to write complete redox reactions and use E'o values to determine DE'° and whether the products or reactants would be favored at equilibrium.
  • The Nernst equation, its meaning and applications:
    • Derive DG'o and the equilbrium constant K'eq for a redox reaction from its DE'° value.
    • Calculate the biochemical reduction potential for a redox reaction under arbitrary conditions

Page updated 01-06-07

References

  1. Berg, Tymoczko, and Stryer. Biochemistry (BTS): 6th edition (2007, Freeman) Ch.18, pp.506-509.
 
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