Learning Objectives
- Understand oxidation, reduction, and redox reactions
- Learn to assign oxidation numbers (states)
- Balance redox equations using ion-electron and oxidation number methods
- Classify redox reactions into different types
- Understand electrochemical cells and their relation to redox
Key Concepts
Oxidation and Reduction
Classical definition: Oxidation = gain of oxygen / loss of hydrogen. Reduction = loss of oxygen / gain of hydrogen.
Electronic concept: Oxidation = loss of electrons. Reduction = gain of electrons.
Oxidation number concept: Oxidation = increase in oxidation number. Reduction = decrease in oxidation number.
Oxidation and reduction always occur simultaneously (redox reactions). The substance oxidised is the reducing agent; the substance reduced is the oxidising agent.
Rules for Assigning Oxidation Numbers
- Free elements: 0 (e.g., Na, O₂, S₈).
- Monoatomic ions: equal to charge (Na⁺ = +1, Cl⁻ = -1).
- Hydrogen: +1 (except metal hydrides: -1, e.g., NaH).
- Oxygen: -2 (except peroxides: -1, superoxides: -½, OF₂: +2).
- Fluorine: always -1.
- Sum of oxidation numbers = 0 (neutral molecule) or charge (for ions).
Types of Redox Reactions
- Combination: A + B → AB (e.g., 2Mg + O₂ → 2MgO)
- Decomposition: AB → A + B (e.g., 2H₂O → 2H₂ + O₂)
- Displacement: A + BC → AC + B (e.g., Zn + CuSO₄ → ZnSO₄ + Cu)
- Disproportionation: Same element is both oxidised and reduced (e.g., 2H₂O₂ → 2H₂O + O₂, where O goes from -1 to -2 and 0).
- Comproportionation: Same element in two different oxidation states combine to form a single oxidation state.
Balancing Redox Equations
Oxidation Number Method:
- Assign oxidation numbers to all atoms.
- Identify atoms whose oxidation numbers change.
- Calculate increase and decrease in oxidation number.
- Equalise total increase and decrease.
- Balance remaining atoms (H and O).
Ion-Electron (Half-Reaction) Method:
- Split into oxidation and reduction half-reactions.
- Balance atoms other than H and O.
- Balance O by adding H₂O (acidic) or OH⁻ (basic).
- Balance H by adding H⁺ (acidic) or H₂O (basic).
- Balance charge by adding electrons.
- Equalise electrons and add half-reactions.
Redox Reactions and Electrode Processes
Electrochemical cells convert chemical energy to electrical energy (galvanic cells) or vice versa (electrolytic cells).
Oxidation occurs at anode (-ve in galvanic). Reduction at cathode (+ve in galvanic).
Activity series determines which metal displaces another from solution.
Summary
Redox reactions involve simultaneous oxidation and reduction, tracked by changes in oxidation numbers. The oxidising agent is reduced (gains electrons); the reducing agent is oxidised (loses electrons). Redox equations are balanced by the oxidation number method or ion-electron method. Disproportionation involves the same element being both oxidised and reduced. Electrochemical applications connect redox chemistry to electricity.
Important Terms
- Oxidation Number: Hypothetical charge on an atom if all bonds were ionic
- Oxidising Agent: Accepts electrons, gets reduced
- Reducing Agent: Donates electrons, gets oxidised
- Disproportionation: Same species undergoes both oxidation and reduction
- Half-Reaction: Oxidation or reduction part of a redox reaction
Quick Revision
- Oxidation: loss of electrons, increase in oxidation number
- Reduction: gain of electrons, decrease in oxidation number
- O.N. rules: H = +1 (except hydrides -1), O = -2 (except peroxides -1)
- Oxidising agent gets reduced; Reducing agent gets oxidised
- Balance by: oxidation number method or ion-electron method
- Disproportionation: H₂O₂, Cl₂ in NaOH, P₄ in NaOH