Redox processes
Section: 6 Electrochemistry | Syllabus: Cambridge AS Level Physics 9702
Oxidation and Reduction - Electron Transfer Redox reactions involve the simultaneous transfer of electrons between species. A useful memory aid is OIL RIG : OIL RIG O xidation I s L oss (of electrons) | R eduction I s G ain (of electrons) Oxidation Loss of electrons by a species, or an increase in oxidation number.
Reduction Gain of electrons by a species, or a decrease in oxidation number. Redox Reaction A reaction in which both oxidation and reduction occur simultaneously - electrons are transferred from one species to another.
Example: Zinc displacing copper Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s) Zn → Zn²⁺ + 2e⁻ (oxidation - loses electrons) Cu²⁺ + 2e⁻ → Cu (reduction - gains electrons) Oxidising Agents and Reducing Agents Oxidising Agent A substance that oxidises another species by accepting electrons from it.
The oxidising agent is itself reduced in the process. Reducing Agent A substance that reduces another species by donating electrons to it. The reducing agent is itself oxidised in the process. Species What it does to electrons What happens to it Change in oxidation number Oxidising agent Accepts electrons Is reduced Decreases Reducing agent Donates electrons Is oxidised Increases Identifying oxidising and reducing agents Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s) Zn is the reducing agent - it donates electrons and is oxidised (0 → +2).
Cu²⁺ is the oxidising agent - it accepts electrons and is reduced (+2 → 0). Oxidation Numbers (Oxidation States) Oxidation numbers are a bookkeeping tool to track electron transfer in redox reactions.
They are the charge an atom would have if the compound were fully ionic. Rules for Assigning Oxidation Numbers Uncombined elements have oxidation number 0 (e.g. Na, O₂, Cl₂, Fe). Simple monatomic ions have oxidation number equal to their charge (e.g.
Na⁺ = +1, Cl⁻ = −1, Fe³⁺ = +3). In compounds, oxygen is usually −2 (exception: in peroxides, e.g. H₂O₂, O = −1; in OF₂, O = +2). In compounds, hydrogen is usually +1 (exception: in metal hydrides, e.g.
NaH, H = −1). Fluorine is always −1 (most electronegative element). The sum of oxidation numbers in a neutral compound = 0 . The sum of oxidation numbers in a polyatomic ion = the charge on the ion . Oxidation number of S in H₂SO₄ H₂SO₄ is neutral: sum = 0 2(+1) + S + 4(−2) = 0 2 + S − 8 = 0 → S = +6 Oxidation number of Cr in Cr₂O₇²⁻ Sum of oxidation numbers = −2 (charge on ion) 2(Cr) + 7(−2) = −2 2Cr − 14 = −2 → 2Cr = 12 → Cr = +6 Oxidation number of Mn in MnO₄⁻ Mn + 4(−2) = −1 Mn − 8 = −1 → Mn = +7 Figure 6.1: Oxidation Number Rules - Summary Card (Table/card format listing all seven rules with examples: element = 0; monatomic ion = charge; O = −2 (−1 in peroxides); H = +1 (−1 in hydrides); F = −1; compound sum = 0; ion sum = charge.
Include one example for each rule.) Roman Numerals for Oxidation Numbers When a metal can have more than one oxidation state, a Roman numeral in brackets after the element name specifies the oxidation state in that particular compound.
Name Roman numeral Ion Example compound Iron(II) Fe(II) Fe²⁺ FeCl₂ (iron(II) chloride) Iron(III) Fe(III) Fe³⁺ FeCl₃ (iron(III) chloride) Copper(I) Cu(I) Cu⁺ Cu₂O (copper(I) oxide) Copper(II) Cu(II) Cu²⁺ CuO (copper(II) oxide) Manganese(VII) Mn(VII) Mn in MnO₄⁻ KMnO₄ (potassium manganate(VII)) Chromium(VI) Cr(VI) Cr in Cr₂O₇²⁻ K₂Cr₂O₇ (potassium dichromate(VI)) Key Point The Roman numeral gives the oxidation number of the metal only - not the overall charge of the compound.
Cu₂O: each Cu is +1 (copper(I)), even though the compound is neutral overall (2 × +1 balances 1 × −2 from O). Identifying Redox Reactions Using Oxidation Numbers A reaction is a redox reaction if one or more elements change oxidation number.
Non-redox reactions show no change in oxidation numbers. Is this a redox reaction? 2Mg + O₂ → 2MgO Mg: 0 → +2 (increase → oxidation) | O: 0 → −2 (decrease → reduction) Yes - redox reaction . Mg is oxidised; O₂ is reduced.
Is this a redox reaction? NaOH + HCl → NaCl + H₂O Na: +1 → +1, O: −2 → −2, H: +1 → +1, Cl: −1 → −1 No change in any oxidation numbers → not a redox reaction (acid–base neutralisation). Figure 6.2: Tracking Oxidation Number Changes in a Redox Equation (Balanced equation: 2Mg + O₂ → 2MgO.
Above the equation, show oxidation numbers for each atom: Mg(0), O(0) on left; Mg(+2), O(−2) on right. Draw curved arrows above Mg showing "+2" increase (oxidation) and below O showing "−2" decrease (reduction).
Label "oxidised" and "reduced" with brackets.) Using Oxidation Numbers to Balance Equations Changes in oxidation numbers must balance - total increase in oxidation number must equal total decrease. This principle helps balance redox equations.
Step 1: Assign oxidation numbers to all atoms. Step 2: Identify which atoms are oxidised and which are reduced. Step 3: Calculate the change in oxidation number for each. Step 4: Balance the changes - multiply species so that total increase = total decrease.
Step 5: Balance remaining atoms (O, H, charge) using H₂O a…
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