R3.2 Electron transfer reactionsIB Chemistry SL: Revision notes
Section 1
Oxidation, reduction and oxidation states
Oxidation is loss of electrons, an increase in oxidation state, gain of oxygen or loss of hydrogen; reduction is the reverse. Rules for oxidation states: elements are 0; the sum equals the charge on the species; O is usually −2 (−1 in peroxides), H usually +1 (−1 in metal hydrides), F always −1.
The oxidising agent is reduced (it takes electrons); the reducing agent is oxidised. In Cr₂O₇²⁻, 2x − 14 = −2 so Cr is +6.
OIL RIG: oxidation is loss, reduction is gain (of electrons).
Section 2
Half-equations in acidic solution
Half-equations show electrons lost or gained. To balance in acid: (1) balance the main atom; (2) add H₂O to balance O; (3) add H⁺ to balance H; (4) add e⁻ to balance charge. Example: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O. To combine half-equations, multiply so the electrons cancel: MnO₄⁻ + 8H⁺ + 5Fe²⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O.
Section 3
Reactivity of metals and halogens
Down group 1 and group 2, metals are oxidised more easily (outer electron further from the nucleus, more shielded). Down group 17, halogens are reduced less easily, so oxidising strength falls F₂ > Cl₂ > Br₂ > I₂: Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂, but I₂ does not oxidise Br⁻.
A metal reduces the ions of any less reactive metal: Zn + Cu²⁺ → Zn²⁺ + Cu. Counting displacements ranks the metals. Metals above hydrogen react with dilute acids to release hydrogen: Mg + 2HCl → MgCl₂ + H₂; Zn + H₂SO₄ → ZnSO₄ + H₂. Copper and silver do not.
Section 4
Voltaic (primary) and secondary cells
In every electrochemical cell, oxidation occurs at the anode and reduction at the cathode. A voltaic cell converts chemical energy from a spontaneous redox reaction into electrical energy: the more reactive metal is the anode (negative). Electrons flow through the wire from anode to cathode; in the salt bridge, anions move towards the anode and cations towards the cathode to keep each solution neutral.
Secondary (rechargeable) cells such as lead–acid and lithium-ion can be recharged: electrical energy drives the discharge reactions in reverse. To get the charging reactions, reverse each discharge half-equation.
Electrons never travel through the salt bridge or the solution — only through the external wire.
Section 5
Electrolytic cells and molten salts
An electrolytic cell uses electrical energy to drive a non-spontaneous reaction. The anode is positive and the cathode negative. Current is carried by electrons in the external circuit and by mobile ions in the molten or dissolved electrolyte; a solid ionic compound does not conduct because its ions are fixed.
Molten salt: cations are reduced at the cathode, anions oxidised at the anode. Molten PbBr₂: cathode Pb²⁺ + 2e⁻ → Pb(l); anode 2Br⁻ → Br₂(g) + 2e⁻.
Section 6
Redox in organic chemistry
With acidified K₂Cr₂O₇ (orange → green): primary alcohols → aldehyde (distil off) → carboxylic acid (heat under reflux); secondary alcohols → ketones; tertiary alcohols are not oxidised.
Reduction: carboxylic acid → aldehyde → primary alcohol using LiAlH₄; ketone → secondary alcohol using NaBH₄ (or LiAlH₄). Write the reducing agent as [H]: CH₃COCH₃ + 2[H] → CH₃CH(OH)CH₃.
Adding hydrogen (Ni catalyst, heat) to alkenes and alkynes lowers the degree of unsaturation: alkyne → alkene → alkane.
Must know
- Oxidation = electron loss / oxidation state increase; the oxidising agent is reduced.
- Balance half-equations with H₂O, H⁺ and e⁻.
- Metals: more easily oxidised down a group; halogens: less easily reduced down the group.
- Anode = oxidation (negative in voltaic, positive in electrolytic).
- Charging a secondary cell reverses the discharge reactions.
- 1° alcohol → aldehyde → carboxylic acid; 2° alcohol → ketone; reductions reverse these.
That's the notes covered.
Carry on to the next subtopic.