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4.9Chemical Reactions

Oxidation-Reduction (Redox) Reactions — AP CHEMISTRY Revision Notes & Topical Questions

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Lesson summary

4.9 Oxidation-Reduction (Redox) Reactions — Notes — AP CHEMISTRY notes: Assign oxidation numbers, identify oxidation and reduction, and balance redox equations using half-reactions.Oxidation = loss of electrons, with an increase in oxidation number.Reduction = gain of electrons, …

  • Disproportionation
  • Reduction = gain of electrons, with a decrease in oxidation number
  • Oxidation and reduction always occur together, so electrons lost equal electrons gained
  • The sum equals the overall charge, so S in SO42− is +6
  • Roman numerals show oxidation number, for example iron(III)
  • Divide the equation into oxidation and reduction half-reactions

Section 1 of 10Oxidation and reduction

Oxidation
Oxidation loss of electrons, with an increase in oxidation number.
Reduction
Reduction gain of electrons, with a decrease in oxidation number.
Oxidation and reduction always occur together, so electrons lost equal electrons gained.
The oxidising agent is the species reduced and the reducing agent is the species oxidised, though these terms are not assessed on the AP exam.

Section 2 of 10Oxidation numbers

RuleOxidation numberExample
Free element0Na, O2, Cl2
Monatomic ionIts chargeNa+ is +1
Fluorine−1HF
Oxygen−2 (peroxides −1)H2O
Hydrogen+1 (metal hydrides −1)HCl
Group 1 and Group 2+1 and +2NaCl, MgO
Oxidation numbers
The sum equals the overall charge, so S in SO42− is +6.
Roman numerals show oxidation number, for example iron(III).

Section 3 of 10Half-reactions

  • A half-reaction shows only the oxidation or only the reduction, with electrons written explicitly, such as Fe2+ → Fe3+ + e−.

Section 4 of 10Balancing in acidic solution

1
Divide the equation into oxidation and reduction half-reactions.
2
Balance atoms other than O and H.
3
Add H2O to balance O.
4
Add H+ to balance H.
5
Add e− to balance charge.
6
Multiply the half-reactions so the electrons cancel.
7
Add them and simplify.
  • Example
    • MnO4− + 8 H+ + 5 e− → Mn2+ + 4 H2O, in which Mn goes from +7 to +2.

Section 5 of 10Common redox reactions

  • Combustion.
  • Metals reacting with acids, such as Mg + 2 H+ → Mg2+ + H2.
  • Single replacement of a metal ion by a more reactive metal.
  • Disproportionation, such as Cl2 + 2 NaOH → NaCl + NaClO + H2O, where one element is both oxidised and reduced.

Section 6 of 10Redox titrations

  • The balanced equation from electron transfer gives the mole ratio used to find the unknown concentration.

Section 7 of 10Electrons in cells

  • In an electrochemical cell electrons flow from the anode to the cathode in the external circuit.

Section 8 of 10More from the specification

Ionic equation
  • An equation showing only the ions that react.
Carbon dioxide
  • A gas that turns limewater cloudy.
Acid-base reaction
  • Reaction involving proton transfer.
Neutralization
  • Acid + base gives salt + water.
Precipitation
  • Forming an insoluble solid when two solutions are mixed.
Chloride
  • Salt of hydrochloric acid.

Section 9 of 10Key terms

TermMeaning
OxidationLoss of electrons.
ReductionGain of electrons.
Oxidation numberCharge an atom would have if shared electrons went to the more attractive atom.
Half-reactionEquation showing only oxidation or only reduction.
DisproportionationReaction in which one species is both oxidised and reduced.
AnodeElectrode where oxidation occurs.

Section 10 of 10Exam technique

1
Mixing up oxidation (loss) and reduction (gain)
2
Forgetting that the sum of oxidation numbers equals the overall charge
3
Not equalising the electrons before adding half-reactions
4
Giving oxygen −2 in peroxides — it is −1

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