Oxidation and Reduction

1. Introduction

Oxidation and reduction together are called redox reactions and are among the most important reactions in chemistry (combustion, corrosion, respiration, batteries, and metal extraction all involve redox). The two processes always occur together: whenever one species is oxidized, another is reduced.

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2. General and Electronic Concepts of Oxidation and Reduction

2.1 Classical (oxygen/hydrogen) concept

• Oxidation = addition of oxygen or removal of hydrogen.
• Reduction = removal of oxygen or addition of hydrogen.

For example, in

$$ \ce{CuO + H2 -> Cu + H2O} $$

, $\ce{CuO}$ loses oxygen (reduced) and $\ce{H2}$ gains oxygen (oxidized).

2.2 Electronic concept (the modern, general definition)

• Oxidation = loss of electrons.
• Reduction = gain of electrons.

A useful memory aid is OIL RIG: Oxidation Is Loss, Reduction Is Gain (of electrons).

$$ \begin{aligned} \ce{Zn -> Zn^2+ + 2e-} \quad (\text{oxidation, loses electrons}) \ \ce{Cu^2+ + 2e- -> Cu} \quad (\text{reduction, gains electrons}) \end{aligned} $$

2.3 Oxidizing and reducing agents

• An oxidizing agent oxidizes another species and is itself reduced (it gains electrons). Examples: $\ce{O2}$, $\ce{KMnO4}$, $\ce{K2Cr2O7}$, $\ce{Cl2}$.
• A reducing agent reduces another species and is itself oxidized (it loses electrons). Examples: $\ce{H2}$, $\ce{C}$, $\ce{Zn}$, $\ce{H2S}$.

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3. Oxidation Number and Rules for Assigning It

The oxidation number (oxidation state) is the apparent charge an atom would carry if all the bonds in the compound were assumed to be completely ionic. It provides a bookkeeping method for tracking electron transfer.

Rules (apply in order of priority):

1. The oxidation number of a free element is zero (e.g. $\ce{Na}$, $\ce{O2}$, $\ce{P4}$ all = 0).
2. For a monatomic ion, it equals the charge on the ion ($\ce{Na+}$ = $+1$, $\ce{S^2-}$ = $-2$).
3. Oxygen is usually $-2$ (except $-1$ in peroxides like $\ce{H2O2}$, and $+2$ in $\ce{OF2}$).
4. Hydrogen is usually $+1$ (except $-1$ in metal hydrides like $\ce{NaH}$).
5. Fluorine is always $-1$; other halogens are $-1$ except when combined with oxygen or a lighter halogen.
6. The sum of oxidation numbers equals zero for a neutral molecule and equals the charge for a polyatomic ion.