Redox reaction: A reaction in which one reactant is oxidised and another is reduced.

Oxidation: a process in which a species loses electrons
Example:

Reduction: a process in which a species gains electrons
Example:

Some background information

In a redox reaction one reactant called the oxidant (older term oxidising agent) oxidises a second reactant called the reductant (older term reducing agent) and is itself reduced.

Example:

More about the jargon

Half−reactions: Redox reactions may be separated into two half−reactions, one involving oxidation (loss of electrons) and the other reduction (gain of electrons).

For the above two examples:

In a balanced half−reaction equation the number of atoms of each element and the total charge must be the same on each side of the equation.

Balanced redox equation: One in which there are no free electrons on either side of the reaction equation. If two half−equations (one oxidation and one reduction) with different numbers of free electrons are added to give an overall redox equation, the equations must be multiplied by integers so that the electrons cancel on addition.

Example:
For the reaction

Oxidation number (oxidation state): A concept central to redox chemistry.

Electronegativity: A measure of the power of an atom to attract electrons to itself when it is part of a compound. The order of electronegativity of common elements is:

F > O > Cl > N > Br > I > S > C > H >P > B > Si

Rules for oxidation number (state):
The oxidation number of an atom in a substance may be determined from the following rules:

1. The oxidation number of an atom in an element is zero.
2. The oxidation number of an atom in a monoatomic ion equals the charge on the ion.
3. Oxygen has the oxidation number of −2, except in peroxides (−1) and when bound to fluorine (+2).
4. Hydrogen has an oxidation number of +1 except in metal hydrides(−1).
5. Halogens have oxidation number −1 except in oxygen−halogen species.
6. The sum of the oxidation numbers of the atoms in a polyatomic species equals the charge on that species.

When an atom in a species has an oxidation number x it is said to be in the x oxidation state.

The above rules, which usually give the correct number, are based on the general concept "the oxidation number of the atom in a species is the charge it would have in the most probable ionic formulation of that species". Putting this another way, it is assumed that all the electrons of a particular bond reside on the atom with the greater electronegativity. So an overall rule to determine oxidation numbers is: write the Lewis structure of the species and count the number of valence electrons on each atom assuming the bonding electrons reside on the more electronegative atom, and compare the number with that of the neutral atom; the difference is the oxidation number.
Example:
HCl; H−Cl; Cl > H; H has 0 electrons; Cl has 8 valence electrons: H, +1; Cl, −1
SiH₄; H−SiH₃; H>Si, each H has 2 valence electrons, Si has 0 valence electrons: H, −1; Si, +4.

The concept of oxidation state plays a major role in the classification of inorganic compounds, but understanding of its significance and usefulness comes only with experience.

A Roman numeral is often used in the name of a species to indicate the oxidation state.
Example: The compounds CrCl₃ and Na₂CrO₄ are called chromium(III) chloride and sodium tetraoxochromate(VI) respectively. See section 13.

In redox reactions the oxidation number of an atom in the oxidant decreases, while that of an atom in the reductant increases.
Example: In the examples above the oxidation state of zinc has increased from 0 to +2, that of bromine has increased from −1 to 0, that of chlorine has decreased from 0 to −1, and that of silver has decreased from +1 to 0.

Rules for balancing redox equations: The ability to write balanced redox equations is an essential skill. This can be done by following a simple set of "book−keeping" rules:

1. Write down the formula for one of the reactants on the left and the formula for its product on the right of an arrow.
2. Balance all elements other than oxygen and hydrogen.
3. Balance oxygens by adding H₂O to the appropriate side.
4. Balance hydrogens by adding H⁺ to the appropriate side.
5. Balance charge by adding electrons to the appropriate side.
   This gives a balanced half−equation. If the electrons are on the right−hand side, the reactant is the reductant (lost electrons and therefore has been oxidised). If the electrons are on the left−hand side the reactant is the oxidant (has gained electrons and therefore has been reduced).
6. Repeat rules 1−5 for the other reactant.
7. Multiply the two half equations by integers so that the number of electrons shown in each half−equation is the same.
8. Add the two half−equations, cancelling equal number of species that occur on both sides.
   The result is the balanced redox equation.

Example: Balance the following redox equation: SO₂ + Cr₂O₇²¯ SO₄²¯ + Cr³⁺

	reactant 1	reactant 2
1.	SO2 SO42¯	Cr2O72¯ Cr3+
2.	SO2 SO42¯	Cr2O72¯ 2Cr3+
3.	2H2O + SO2 SO42¯	Cr2O72¯ 2Cr3+ + 7H2O
4.	2H2O + SO2 SO42¯ + 4H+	14H+ + Cr2O72¯ 2Cr3+ + 7H2O
5.	2H2O + SO2 SO42¯ + 4H+ +2e¯	6e¯ + 14H+ + Cr2O72¯ 2Cr3+ + 7H2O
	(Therefore SO2 is the reductant)	(Therefore Cr2O72¯ is the oxidant)
	Now multiply eqn 5 (Reactant 1) by 3
7.	6H2O + 3SO2 3SO42¯ + 12+ +6e¯	6e¯ + 14H+ + Cr2O72¯ 2Cr3+ + 7H2O

finally:
2H⁺ + 3SO₂ + Cr₂O₇²¯ 3SO₄²¯ + 2Cr³⁺ + H₂O

A useful check when balancing redox equations is to note that the number of electrons in each half−reaction is the same as the changes in oxidation numbers.