Chemical reactions that involve transfer of electrons are called redox (reduction-oxidation) reactions. One species is reduced (gains electrons, its oxidation number decreases) and another species is oxidised (loses electrons, its oxidation number increases).
A reaction can be classified as redox (or not) by determining the oxidation numbers of all elements in the reaction equation. If a change in oxidation number occurs, the reaction is redox.

Rules for determining oxidation numbers (in order):

• If the species only has one element, use the rule that total oxidation number is equal the charge on the species
• Fluorine has an oxidation number of −1 in all its compounds
• Hydrogen has an oxidation number of +1 except in metal hydrides (-1)
• Oxygen has an oxidation number of −2 except in hydrogen peroxide (−1)
• Alkali metals (group I metals) are +1 and alkaline earth metals (group II metals) are +2
• Starting with the most electronegative, complete the rest such that the total oxidation number is equal to the total charge of the species

Redox numbers practice here.

Redox equations can be balanced as half-equations and then combined to produce a full equation. When balancing a half-equation:

1. Balance elements that are already on both sides
2. Balance the oxygen by adding water
3. Balance the hydrogen by adding hydrogen ions
4. Balance the charge by adding electrons

To combine the half-equations, add enough of each so the electrons are equal on both sides. Then cancel anything that is the same on both sides.

Half-equation method practice here.

Elements can be listed in order of how easily they are oxidised (or reduced). This ordered list is called a reactivity series; the most easily oxidised (most difficult to reduce) elements are at the top and the most easily reduced (most difficult to oxidise) elements are at the bottom.
Species that are easily oxidised are strong reducing agents, and species that are easily reduced are strong oxidising agents.

Single displacement is a reaction in which a more reactive metal replaces ions of a less reactive metal in a compound or solution. The metal loses electrons to become ions, and the ions gain electrons to become metal.
If the more reactive metal is in the form of ions already (oxidised form), or the less reactive metal is in the form of metal already (reduced form), no reaction will occur.

Redox reaction prediction practice here.

An electrochemical cell either produces electricity by a chemical reaction or consumes electricity to cause a chemical reaction.

Galvanic (voltaic) cells produce an electric current from a spontaneous redox reaction. The electrons will flow from the more easily oxidised electrode (anode) to the more easily reduced electrode (cathode).

YouTube video about galvanic cells here.

Electrolytic cells use an electric current from an external source to force a non-spontaneous redox reaction. During electrolysis of a molten or aqueous salt, the nonmetal ion or water will be oxidised at the anode, and the metal ion or water will be reduced at the cathode. Use an electrochemical series to determine which species are more easily oxidised or reduced.

YouTube video about electrolytic cells here.

Electrolysis product practice here.