Lecture 18 summary

We make an analogy between the acid dissociation equation, which as we have seen can be thought of as half of a complete Brønsted-Lowry proton transfer (acid-base) reaction. To extend this analogy, we define a value known as the standard reduction potential, E°, which is a way of ordering half-reactions according to their tendency to favor the reduced species. Just as K_a (or pK_a) provides a quantitative measure of the relative strength of an acid (tendency to act as a proton donor), half-reaction E° values will tell us quantitatively the relative tendency of species that can undergo redox reactions to compete for electrons.

Our goal is to construct an ordered series of reduction half-reactions that will be a useful guide in assessing the strength of various oxidizing or reducing agents, much as we wanted to know the relative strengths of acids and bases. To generate such an ordered list, the voltages produced by various pairs of half-cells must be measured experimentally under defined conditions, and a standard reference half-cell is chosen to set the zero point of the scale, much as sea-level sets the zero reference point for elevation, or the freezing point of water sets the zero point of the Celsius temperature scale. The definition of standard state provides the defined condition for such measurements. The standard state of a substance is its most stable state at a specified temperature, usually 25°C (298 K). Standard states are liquids or solids, while for species in solution or gases, the standard state is defined as 1 M concentration, or 1 bar pressure (close to 1 atm), respectively. The reference half-cell for standard reduction potential measurements, the source of experimentally determined E° values, is the standard hydrogen electrode (abbreviated S.H.E.). See Harris, p.294-295 for details.

In order to apply electrochemical principles and measurements to galvanic cells and redox reactions occurring under arbitrary, non-standard conditions, we need a way of quantifying the effect of variations in the concentrations of all relevant species present in the system. Recall that LeChâtelier's Principle tells us that in a general chemical reaction, either increasing concentrations of reactants or decreasing those of the products will shift the equilibrium of the reaction in favor of the products. The principle of course applies to redox reactions, and a quantitative statement known as the Nernst equation gives us the tool we need to relate measurements made under standard conditions to any arbitrary conditions of reactant and product concentrations.

Ordered redox potentials

The tables below show selected redox half-reactions and their measured standard reduction potentials (in volts, V). By convention, each half reaction is written as a reduction (the reduced species is shown on the product side). The greater the E° value for the half-reaction, the more favorable the reduced species. Hence, with the half-reactions listed in order of highest E° to lowest, the species on the left side of half-reactions at the top of the table are the strongest oxidants, while species on the right side of half-reactions at the bottom of the table are the strongest reductants. Also note the difference between the two tables. Table 2 lists half-reactions according to their biochemical standard reduction potentials, E° '. The principal difference between standard state and biochemical standard state, signified by the prime ('), has to do with pH. When the definition of standard state ("physical chemistry" standard state) given above is applied to [H⁺], we see that 1 M [H⁺], standard state for hydronium ion concentration, corresponds to a pH of 0. This is incompatible with biological and physiological conditions, so for biochemical standard state, the standard state for [H⁺] is redefined to correspond to pH = 7. The difference between standard state and biochemical standard state is also explained in Box 14-1 on p.303 of Harris.