Defining Oxidation-Reduction Reactions
Electrochemistry is the branch of science that studies chemical reactions that cause electrons to move, resulting in the transfer of electric charge. Reactions in which electrons move from one element to another are called oxidation-reduction reactions, commonly referred to as redox reactions. In many chemical reactions, the enthalpy change involves a transfer of heat into or out of the system. Redox reactions can be set up so that the enthalpy change in the reaction involves a flow of electrons instead of an exchange of heat.
In redox reactions, an element that loses one or more electrons is said to be oxidized, while an element that gains one or more electrons is said to be reduced. A helpful way to remember the relationship is through this phrase: LEO (Loses Electrons—Oxidation) goes GER (Gains Electrons—Reduction). The species that is oxidized is also called the reducing agent, or reductant, while the species that is reduced is also called the oxidizing agent, or oxidant.For example, the Haber-Bosch process can be shown by the following redox reaction.
Oxidation Number Change in a Redox Reaction
Balancing Oxidation-Reduction Reactions
Balancing a redox reaction is similar to balancing other chemical reactions, except that the electron transfer must be considered. There are eight steps to follow when balancing redox reactions. These eight steps have variations depending on whether the reaction occurs in a neutral, acidic, or basic environment. Acidic conditions have excess hydrogen ions (H+). Basic conditions have excess hydroxide ions (OH–). Neutral conditions do not have excess hydrogen (H+) or hydroxide ions (OH–).
1. Divide the reaction into its two half-reactions.
2. Balance all elements in the half-reaction other than oxygen and hydrogen.
3. Balance oxygen by adding the appropriate number of water (H2O) molecules to the other side of the reaction. (If the condition is neutral, skip this step.)
4. Balance hydrogen (including those added in step 3) by adding hydrogen ions (H+), also known as protons, to the opposite side of the reaction. (If the condition is neutral, skip this step.)
5. Balance the charges on both sides of the half-reactions by adding electrons (e–) to the side with the greater positive charge.
6. Multiply each half-reaction by the smallest whole number required to equalize the electrons gained by reduction and lost by oxidation.
7. Add the half-reactions together to form a complete reaction and cancel any common terms. (If the condition is basic, go on to step 8. Otherwise stop at this step.)
8. Add hydroxide ions (OH–) to turn remaining hydrogen ions (H+) to water molecules, and cancel any common terms on either side of the reaction.