Oxidation-Reduction Reactions

Introduction

Recall that chemical reactions are all about electrons. In redox (an often used shorthand term for oxidation-reduction) reactions, one is concerned with how electrons are transferred from one species (element, ion, compound) to another. The law of conservation of charge stipulates that if one species loses electrons, there must be another species that gains the electrons. The term oxidation refers to any process that loses electrons. This is a tricky word for some because oxidation sounds like it might only involve oxygen. In truth, oxygen is an important element in redox reactions, and because of its large electronegativity, causes many other elements to give it electrons. Since oxygen often "forces" other elements to lose electrons, the process of losing electrons became known as oxidation. Reduction is a similarly troubling word, since it refers to a process that gains electrons. In this case, the word "reduce" has nothing to do with the quantity of electrons, but rather the change in charge. Since the addition of electrons will add more negative charges to a species, it is its charge that is reduced.

Assigning Oxidation Numbers

Oxidation numbers are the charges of the individual elements that make up an ion or compound. The sum of all the individual oxidation numbers must equal the charge of the ion, or zero if it is a compound. In general, the first element in a formula will be positive and the last will be negative. Many elements have other more specific guidelines to follow:

Free state elements (those that are uncombined with any other element, such as Ni or I2) are zero.
The metals of group 1 are +1.
Hydrogen is usually +1, but when it appears as a hydride (a metal-hydrogen compound), it will be a -1.
The elements of group 2 are +2.
The transition elements, lanthanides, and actinides will vary between +1 and +7, although each element may only have specific possibilities.
When appearing last in a formula, the halogens are -1. Fluorine will never appear first because of its electronegativity.
When appearing last in a formula, the chalcogens are -2.
Oxygen can form peroxides such as H2O2, and in such a case it is -1.
Nitrogen and phosphorus are -3 when appearing last in a formula.
The metals under the "staircase" (groups 13-16), like the transition metals can vary, but are always positive.
Assign the oxidation number for the negative charge first. Use it to deduce the positive charge if it is an element that can exhibit multiple charges.

Example Reaction

Consider the reaction shown below:

This is an example of a synthesis reaction, but it also a redox reaction. The reason is because the two elements involved, iron and oxygen, exchange electrons to produce iron(III) oxide, the product. The reactants are free state elements, so their oxidation numbers are each zero. However, in the compound Fe2O3 the iron and oxygen do not have a charge of zero. Based on oxidation number assignment rules, it can be assumed that oxygen is a -2. Since there are three of them, the total charge of the oxygens is -6. The two iron ions must balance that and total a +6. In order to do this, each iron must be a +3.