Redox Processes: Electron Transfer and Changes in Oxidation Number (Oxidation State)

2026 Syllabus Objectives

By the end of this topic, you should be able to:

  1. Calculate oxidation numbers of elements in compounds and ions
  2. Use changes in oxidation numbers to help balance chemical equations
  3. Explain and use the terms redox, oxidation, reduction and disproportionation in terms of electron transfer and changes in oxidation number
  4. Explain and use the terms oxidising agent and reducing agent
  5. Use a Roman numeral to indicate the magnitude of the oxidation number of an element

1. What is an Oxidation Number?

https://cambridgeassistant.com

An oxidation number (also called oxidation state) is a number we give to each atom in a compound or ion. It helps us keep track of how many electrons an atom has "control of" compared to when it's on its own.

  • A positive oxidation number means the atom has lost control of electrons (lost electrons)
  • A negative oxidation number means the atom has gained control of electrons (gained electrons)
  • An oxidation number of zero means the atom hasn't lost or gained control of any electrons

Think of it like a score that shows whether an atom is "winning" or "losing" electrons in a chemical battle.

2. Rules for Calculating Oxidation Numbers

https://cambridgeassistant.com

To work out oxidation numbers, follow these rules in order:

Rule 1: Uncombined Elements

Any element on its own (not bonded to anything else) has an oxidation number of 0.

Examples:

  • H₂ → each H is 0
  • Zn → oxidation number is 0
  • O₂ → each O is 0

Rule 2: Fixed Oxidation Numbers

Some atoms almost always have the same oxidation number:

  • Group 1 elements (Li, Na, K, etc.) are always +1
  • Group 2 elements (Mg, Ca, etc.) are always +2
  • Group 3 elements (Al) are always +3
  • Fluorine is always -1
  • Hydrogen is usually +1 (except in metal hydrides like NaH, where it's -1)
  • Oxygen is usually -2 (except in peroxides like H₂O₂ where it's -1, and in F₂O where it's +2)

Rule 3: Simple Ions

The oxidation number of a simple ion (single atom) equals its charge.

Examples:

  • Na⁺ → oxidation number = +1
  • Cl⁻ → oxidation number = -1
  • Fe³⁺ → oxidation number = +3
  • O²⁻ → oxidation number = -2

Rule 4: Compounds

The sum of all oxidation numbers in a neutral compound = 0.

Example: NaCl

  • Na = +1
  • Cl = -1
  • Sum = +1 + (-1) = 0 ✓

Rule 5: Complex Ions

The sum of all oxidation numbers in a complex ion = the charge on the ion.

Example: SO₄²⁻

  • Each O = -2
  • 4 oxygen atoms = 4 × (-2) = -8
  • For the whole ion to have a charge of -2:
  • S + (-8) = -2
  • S = +6

Rule 6: Electronegativity

The more electronegative element (the one that pulls electrons harder) gets the negative oxidation number.

Example: F₂O

  • Fluorine is more electronegative than oxygen
  • Each F = -1 (2 fluorines = -2)
  • So O must be +2 (to make the sum zero)

Sign in to view full notes