Flashcards · 9701 · AS Topic 6

Electrochemistry (AS Level) flashcards

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Oxidation and reduction in terms of electrons and oxidation number: Oxidation: loss of electrons, increase in oxidation number. Reduction: gain of electrons, decrease in oxidation number.
Oxidising agent and reducing agent: An oxidising agent oxidises another species and is itself reduced (gains electrons). A reducing agent reduces another species and is itself oxidised (loses electrons).
Key oxidation-number rules (with the exceptions): Element 0; simple ion = its charge; F −1; Group 1 +1, Group 2 +2; H +1 (−1 in metal hydrides); O −2 (−1 in peroxides, +2 in OF₂); sum = 0 in a compound, = charge in an ion.
Oxidation number of Mn in MnO₄⁻ and of Cr in Cr₂O₇²⁻: Mn: x + 4(−2) = −1, so +7 (manganate(VII)). Cr: 2x + 7(−2) = −2, so +6 (dichromate(VI)).
Half-equation for MnO₄⁻ → Mn²⁺ in acid: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O
Steps for balancing a half-equation: 1 Balance atoms other than O and H. 2 Balance O with H₂O. 3 Balance H with H⁺. 4 Balance charge with electrons. Combine halves so the electrons cancel.
Disproportionation: A reaction in which the same element is simultaneously oxidised and reduced, e.g. Cl₂ + 2NaOH → NaCl + NaClO + H₂O (Cl: 0 → −1 and 0 → +1).
Balanced equation for MnO₄⁻ oxidising Fe²⁺ in acid: MnO₄⁻ + 5Fe²⁺ + 8H⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O Mn gains 5 electrons (+7 → +2), so 5 Fe²⁺ each lose one (+2 → +3). Charge: +17 on each side.