Redox Reactions
 
 
 
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REVIEW

Redox reactions, also known as oxidation–reduction reactions, are electron transfer reactions consisting of two parts. Each of these two parts is called a “half reaction.” The oxidation half reaction donates electrons that are accepted by the reduction half reaction.

The simple mnemonic “OILRIG” can be used to define and remember redox reactions:

Oxidation Is Loss (of electrons)
Reduction Is Gain (of electrons)

It is important to be able to identify which compounds are oxidized and which are reduced in a redox reaction. There are several simple rules for determining the oxidation states of compounds:

  1. All uncharged elements and compounds have an oxidation state of zero. Examples are Zn, H2, O2, and KMnO4.
  2. All charged elements and compounds have an oxidation state equal to their charge.
  3. Oxygen in a compound almost always has an oxidation state of –2 .
  4. Hydrogen in a compound almost always has an oxidation state of +1.
  5. Some elements always have the same oxidation states when they are in a compound. These include “group 1” elements such as H, Li, Na, K (always +1), “group 2” elements such as Mg and Ca (always+2), and “group 7” elements such as F and Cl (always –1).
The higher the E°' value, the stronger the tendency for the compound to gain electrons.

An E°' value, known as the reduction potential, measures the tendency of a substance to gain electrons. A higher E°' value indicates a stronger tendency to gain electrons. When two half reactions are paired in a redox reaction, the half reaction with the higher E°' will act as the reduction reaction and the half reaction with the lower E°' will act as the oxidation reaction.

The tendency for a reaction to proceed in the direction it is written can be determined from DE°'. A positive DE°' indicates that a reaction will proceed in the direction it is written. The DE°' can be calculated from the formula

DE°' = (E°' from reduction reaction) – (E°' from oxidation reaction)

The DE°' value is related to DG°', the Gibbs Free Energy change of a reaction, by the formula

DG°' = – nF(DE°')
where

DG°' is the Gibbs free energy change of a system (in Joules)
DE°' is the change in reduction potential of the redox reaction (in Volts)
n is the number of electrons transferred in the reaction (in moles)
F is Faraday’s constant (96,500 Joules/Volt/mol)

A spontaneous redox process is indicated by either a positive DE°' or a negative DG°'.

You have completed this exercise.