Redox Reactions: Oxidation and Reduction

Introduction

Redox reactions, short for reduction-oxidation reactions, are chemical reactions that involve the transfer of electrons between two species. These reactions are fundamental in chemistry and occur in numerous everyday processes, from respiration to corrosion.

In a redox reaction, one species loses electrons (oxidation) and another species gains electrons (reduction). These two processes always happen together; you cannot have one without the other.

Oxidation: Loss of Electrons

Oxidation is the process of a substance losing electrons. It's a fundamental change in the electronic structure of an atom, ion, or molecule.

A helpful mnemonic to remember oxidation is: OIL - Oxidation Is Loss.

Here's a more detailed explanation:

• Oxidation always involves an increase in the oxidation state of a substance.
• Metals typically undergo oxidation, losing electrons to form positive ions (cations).
• Non-metals typically undergo oxidation, gaining electrons to form negative ions (anions).

Reduction: Gain of Electrons

Reduction is the process of a substance gaining electrons. It's the reverse process of oxidation.

A helpful mnemonic to remember reduction is: RIG - Reduction Is Gain.

Here's a more detailed explanation:

• Reduction always involves a decrease in the oxidation state of a substance.
• Non-metals typically undergo reduction, gaining electrons to form negative ions (anions).
• Metals typically undergo reduction, gaining electrons to form positive ions (cations).

Oxidation States

Oxidation states (also known as oxidation numbers) are a way of keeping track of the number of electrons that an atom has gained, lost, or shared in a chemical compound. They are used to identify whether oxidation or reduction has occurred.

The rules for assigning oxidation states are complex, but here are a few basic ones:

• The oxidation state of an atom in its elemental form is 0.
• The oxidation state of a monatomic ion is equal to its charge (e.g., Na⁺ is +1, Cl^- is -1).
• Oxygen is usually -2 (except in peroxides like H₂O₂ where it's -1).
• Hydrogen is usually +1 (except in metal hydrides like NaH where it's -1).
• Group 1 metals are always +1.
• Group 2 metals are always +2.

Examples of Oxidation and Reduction

Here are some examples to illustrate oxidation and reduction:

Reaction	Oxidation	Reduction
$2Na(s) + Cl_2(g) \rightarrow 2NaCl(s)$	Sodium loses an electron to become Na+ ($Na \rightarrow Na^+ + e^-$)	Chlorine gains an electron to become Cl- ($Cl + e^- \rightarrow Cl^-$)
$Fe(s) + O_2(g) \rightarrow Fe_2O_3(s)$	Iron loses electrons to become Fe3+ ($Fe \rightarrow Fe^{3+} + 3e^-$)	Oxygen gains electrons to become O2- ($O_2 + 4e^- \rightarrow 2O^{2-}$ )
$Zn(s) + Cu^2+(aq) \rightarrow Zn^2+(aq) + Cu(s)$	Zinc loses electrons to become Zn2+ ($Zn \rightarrow Zn^{2+} + 2e^-$)	Copper(II) ion gains electrons to become Copper metal ($Cu^{2+} + 2e^- \rightarrow Cu$)

Summary

Understanding oxidation and reduction is crucial for comprehending many chemical processes. Remember the acronyms OIL and RIG to help you recall the definitions. Oxidation is the loss of electrons, and reduction is the gain of electrons. These processes always occur together in redox reactions.