Chemistry 401

Oxidation/Reduction Chemistry

Transfer of electrons from oxidized species (loses electrons) to reduced species (gains electrons)

Connections between structure/bonding and reactivity are less obvious than in acid/base chemistry.

Thermodynamic Considerations

The tendency for an electron transfer reaction to occur is embodied in the potential, E, which is directly related to the Gibb's energy, ΔG

ΔG = –nFE

F is Faraday's constant = 96485 Coulombs/mole

n = the number of moles of electrons transferred in the reaction

E is typically more convenient because it can be measured readily using electrochemical techniques.

At standard conditions (1 bar pressure, 25 °C, 1 mole reactant, pH = 0), potentials are denoted E° and are further partitioned (arbitrarily) into reduction potentials for each reacting species:

E° = E°_red(reduced species) – E°_red(oxidized species)

Reduction potentials are tabulated in a variety of formats.

At nonstandard conditions, the Nernst equation is used to find a reaction potential:

Q is the reaction quotient (mass action expression)

The Nernst equation can also be applied to half-reactions:

Q_red is the mass action expression of the reduction half-reaction, ignoring the electrons.

Q_red is only a relative value since it is based on the arbitrary partitioning done to obtain E°_red

This explains the strong pH dependence of many redox reactions:

O₂(g) + 4 H⁺(aq) + 4 e^– → 2 H₂O(l)

at pH = 0 (standard conditions) E°_red = +1.23 V

at other pH values:

assuming an oxygen pressure of 1 bar