Key Points

Redox reactions are chemical reactions in which oxidation and reduction occur simultaneously. The term 'redox' is a portmanteau of reduction and oxidation.

Classically, oxidation is defined as the addition of oxygen or another electronegative element to a substance, or the removal of hydrogen or an electropositive element from it.

Reduction is defined as the removal of oxygen or an electronegative element from a substance, or the addition of hydrogen or an electropositive element to it.

In terms of electron transfer, oxidation is the loss of electrons by a species (LEO: Loss of Electrons is Oxidation), while reduction is the gain of electrons (GER: Gain of Electrons is Reduction).

An oxidizing agent (oxidant) is a substance that accepts electrons and gets reduced. A reducing agent (reductant) is a substance that donates electrons and gets oxidized.

The oxidation number represents the hypothetical charge an atom would have if all bonds to atoms of different elements were 100% ionic. It is used to track electron shifts in reactions.

A redox reaction can be identified by a change in oxidation numbers. Oxidation involves an increase in oxidation number, while reduction involves a decrease in oxidation number.

For a free element, the oxidation number is 0. For a monatomic ion, it equals the ion's charge. Oxygen is typically -2 (except in peroxides like $\text{H}_2\text{O}_2$ where it is -1), and Hydrogen is typically +1 (except in metal hydrides like $\text{NaH}$ where it is -1).

Stock notation indicates the oxidation state of a metal in a compound using a Roman numeral in parentheses after the metal's symbol. For instance, Iron(III) oxide is written as $\text{Fe}_2\text{O}_3$.

Redox reactions are classified into four main types: combination, decomposition, displacement, and disproportionation reactions.

A combination redox reaction involves elements combining, like $2\text{Mg(s)} + \text{O}_2\text{(g)} \rightarrow 2\text{MgO(s)}$. A decomposition redox reaction is the opposite, like $2\text{H}_2\text{O(l)} \rightarrow 2\text{H}_2\text{(g)} + \text{O}_2\text{(g)}$.

In a displacement reaction, an atom or ion in a compound is replaced by one from another element. An example is metal displacement: $\text{Zn(s)} + \text{CuSO}_4\text{(aq)} \rightarrow \text{ZnSO}_4\text{(aq)} + \text{Cu(s)}$.

This is a special type of redox reaction where an element in one oxidation state is simultaneously oxidized and reduced. For example, in $2\text{H}_2\text{O}_2 \rightarrow 2\text{H}_2\text{O} + \text{O}_2$, oxygen is both oxidized (from -1 to 0) and reduced (from -1 to -2).

This method involves splitting the reaction into oxidation and reduction half-reactions. Each half is balanced for atoms and charge, electrons are equalized, and then the half-reactions are added together.

A redox couple consists of the oxidized and reduced forms of a substance involved in a half-reaction, written as Oxidized form/Reduced form (e.g., $\text{Zn}^{2+}/\text{Zn}$). These form the basis of electrochemical cells.

The standard electrode potential ($E^{\ominus}$) measures the tendency of a species to be reduced. A more positive $E^{\ominus}$ value indicates a stronger oxidizing agent, while a more negative value indicates a stronger reducing agent.