The standard reduction potential, E^'0, of any redox pair is...

The standard reduction potential, $E^{\prime 0},$ of any redox pair is defined for the half-cell reaction: Oxidizing agent $+n$ electrons $\rightarrow$ reducing agent The $E^{\prime \circ}$ values for the $\mathrm{NAD}^{+} / \mathrm{NADH}$ and pyruvate/lactate conjugate redox pairs are -0.32 V and $-0.19 \mathrm{V},$ respectively. (a) Which redox pair has the greater tendency to lose electrons? Explain. (b) Which pair is the stronger oxidizing agent? Explain. (c) Beginning with $1 \mathrm{M}$ concentrations of each reactant and product at $\mathrm{pH} 7$ and $25^{\circ} \mathrm{C}$ in which direction will the following reaction proceed? $$\text { Pyruvate }+\mathrm{NADH}+\mathrm{H}^{+} \rightleftharpoons \text { lactate }+\mathrm{NAD}^{+}$$ (d) What is the standard free-energy change $\left(\Delta G^{\prime \prime}\right)$ for the conversion of pyruvate to lactate? (c) What is the equilibrium constant $\left(K_{e q}^{\prime}\right)$ for this reaction?

The standard reduction potential, $E^{\prime 0},$ of any redox pair is defined for the half-cell reaction: Oxidizing agent $+n$ electrons $\rightarrow$ reducing agent
The $E^{\prime \circ}$ values for the $\mathrm{NAD}^{+} / \mathrm{NADH}$ and pyruvate/lactate conjugate redox pairs are -0.32 V and $-0.19 \mathrm{V},$ respectively.
(a) Which redox pair has the greater tendency to lose electrons? Explain.
(b) Which pair is the stronger oxidizing agent? Explain.
(c) Beginning with $1 \mathrm{M}$ concentrations of each reactant and product at $\mathrm{pH} 7$ and $25^{\circ} \mathrm{C}$ in which direction will the following reaction proceed? $$\text { Pyruvate }+\mathrm{NADH}+\mathrm{H}^{+} \rightleftharpoons \text { lactate }+\mathrm{NAD}^{+}$$
(d) What is the standard free-energy change $\left(\Delta G^{\prime \prime}\right)$ for the conversion of pyruvate to lactate?
(c) What is the equilibrium constant $\left(K_{e q}^{\prime}\right)$ for this reaction?

Key Concepts

Relationship Between Free Energy and Reduction Potential

The free energy change (?G) in a redox reaction is directly linked to the standard reduction potential through the equation ?G = -nFE, where n is the number of electrons transferred and F is Faraday's constant. A more positive reduction potential leads to a more negative ?G, which indicates a more spontaneous reaction under standard conditions.

Equilibrium Constant and Gibbs Free Energy

The equilibrium constant (K_eq) for a redox reaction is connected to the change in standard free energy (?G°) by the relationship ?G° = -RT ln K_eq, where R is the gas constant and T is the temperature in Kelvin. This equation helps predict the position of equilibrium in a reaction, with a larger K_eq corresponding to a more thermodynamically favored reaction.

Standard Reduction Potential

Standard reduction potential measures the tendency of a chemical species to acquire electrons and be reduced. It is defined under standard conditions for half-cell reactions. A higher (more positive) value indicates a greater tendency to be reduced, whereas a lower (more negative) value implies that the species is less likely to gain electrons and more inclined to lose them when compared to another redox pair.

Oxidizing and Reducing Agents

An oxidizing agent is a substance that gains electrons in a redox reaction and is reduced, while a reducing agent loses electrons and is oxidized. The strength of these agents is associated with their standard reduction potentials: a strong oxidizing agent has a high (positive) reduction potential, and a strong reducing agent has a low (negative) reduction potential, reflecting the tendency to undergo oxidation.

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