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4.3.3Chemical Reactions

Assigning Oxidation Numbers — AP CHEMISTRY Revision Notes & Topical Questions

5 min read 5 parts Lesson 17 of 18 Unit quiz

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Curriculum Equation Key قوانين الرقم الهيدروجيني والاتزان الأيوني - pH & Acid-Base Equilibrium FormulasFast Exam Solution
pH = -log[H⁺] | pOH = -log[OH⁻] | pH + pOH = 14 | [H⁺][OH⁻] = 10⁻¹⁴
Word Equation & Conditions: الرقم الهيدروجيني = -لوغاريتم تركيز أيون الهيدروجين | pH + pOH = 14
Location & Specifications: الاتزان الأيوني، كيمياء 3 ثانوي، التحصيلي، و AP Chemistry.
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Part 1 of 5What an Oxidation Number Is

An oxidation number is a bookkeeping charge showing where electrons would sit if every bond were treated as fully ionic.

What an Oxidation Number Is
Bookkeeping device
  • Bookkeeping device It tracks electron transfer even in covalent species, where no atom actually carries the full charge.
Written with the sign first
  • Written with the sign first Oxidation numbers are written +2 or −1, whereas ionic charges are written 2+ or 1−.
Increase means oxidation
  • Increase means oxidation An atom whose oxidation number rises has lost electrons, or electron density, and has been oxidised.
Decrease means reduction
  • Decrease means reduction An atom whose oxidation number falls has gained electrons, or electron density, and has been reduced.

Part 2 of 5The Rules in Priority Order

Apply the rules from the top down; when two rules clash, the one listed earlier wins.

1
An atom in an uncombined element, such as Cu, O₂ or S₈, has oxidation number 0.
2
A monatomic ion has an oxidation number equal to its charge, so Mg²⁺ is +2 and Cl⁻ is −1.
3
The oxidation numbers in a neutral compound add up to 0, and in a polyatomic ion they add up to the ion's charge.
4
Group 1 metals are +1 and Group 2 metals are +2 in their compounds; fluorine is always −1.
5
Hydrogen is usually +1 (but −1 in metal hydrides such as NaH) and oxygen is usually −2 (but −1 in peroxides such as H₂O₂).
6
The remaining element is found by subtraction, so that the total equals the required sum.

Part 3 of 5Using the Numbers

Once numbers are assigned, comparing reactants with products identifies the redox process.

Oxidising agent
  • Oxidising agent The species containing the atom that is reduced; it takes electrons from another species.
Reducing agent
  • Reducing agent The species containing the atom that is oxidised; it supplies electrons to another species.
Not a redox reaction
  • Not a redox reaction If no atom changes its oxidation number, the reaction is not redox, as in precipitation or simple acid-base reactions.
Variable metals
  • Variable metals Transition metals have several values, for example iron is +2 in FeCl₂ and +3 in FeCl₃.
Fractional values
  • Fractional values An average such as +8/3 for iron in Fe₃O₄ can occur when atoms of one element are in different environments.
SpeciesAtomOxidation number
MnO₄⁻Mn+7
Cr₂O₇²⁻Cr+6
NH₄⁺N−3
H₂O₂O−1

Part 4 of 5Worked Example

Question Find the oxidation number of manganese in MnO₄⁻, then state what happens to it in MnO₄⁻ → Mn²⁺.

  1. Oxygen is −2, so the four oxygen atoms contribute 4 × (−2) = −8.
  2. The ion has charge −1, so Mn + (−8) = −1, giving Mn = +7.
  3. In Mn²⁺ the oxidation number is +2, so manganese changes from +7 to +2, a decrease of 5.

Answer

  • Mn is +7 in MnO₄⁻;
    • it is reduced, gaining 5 electrons per ion, so MnO₄⁻ is the oxidising agent

Part 5 of 5Exam Checklist & Tips

4.3.3 – Key Points for the Exam

  • Oxidation number The hypothetical charge an atom would have if all bonding electrons were assigned to the more electronegative atom.
  • Oxidation A process in which an atom's oxidation number increases because electrons are lost.
  • Reduction A process in which an atom's oxidation number decreases because electrons are gained.
  • Oxidising agent The reactant that is reduced and causes another species to be oxidised.
  • Reducing agent The reactant that is oxidised and causes another species to be reduced.

Common Mistakes

  • Confusing oxidation number with ionic charge and writing the sign after the number.
  • Giving oxygen −2 in peroxides such as H₂O₂, where it is −1.
  • Forgetting that the sum must equal the ion's charge, not zero, for polyatomic ions.
  • Naming the species being oxidised as the oxidising agent, when it is actually the reducing agent.

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