Electrochemistry links chemical reactions with electricity. It covers electrolysis (using electricity to decompose compounds), the reactivity series, and redox chemistry.
Electrolysis: decomposition of an ionic compound by passing electric current through it (molten or dissolved). Electrolyte: ionic compound that conducts. Anode: positive electrode (attracts anions/negative ions). Cathode: negative electrode (attracts cations/positive ions). At cathode: cations gain electrons (reduction). At anode: anions lose electrons (oxidation). OILRIG: Oxidation Is Loss, Reduction Is Gain.
In aqueous solutions, water ions compete: H⁺ and OH⁻. At cathode: hydrogen produced unless metal is less reactive than hydrogen (then metal deposited). At anode: if halide present → halogen produced; otherwise → oxygen from OH⁻. Example: electrolysis of brine (NaCl solution) → H₂ at cathode, Cl₂ at anode, NaOH left in solution. Industrial importance: aluminium extraction (electrolysis of Al₂O₃ in cryolite).
Reactivity series: K, Na, Ca, Mg, Al, C, Zn, Fe, H, Cu, Ag, Au. More reactive metals: react with water/acid, lose electrons easily. Extraction: reactive metals by electrolysis, less reactive by reduction with carbon, least reactive found native. Displacement: more reactive metal displaces less reactive from solution. Redox: oxidation = loss of electrons / gain of oxygen; reduction = gain of electrons / loss of oxygen. Oxidation and reduction always occur together.
Aluminium is more reactive than carbon in the reactivity series, so carbon cannot reduce aluminium oxide. Electrolysis of molten aluminium oxide (dissolved in cryolite to lower the melting point from ~2050°C to ~950°C) provides the energy needed to decompose Al₂O₃. At the cathode: Al³⁺ + 3e⁻ → Al. At the anode: 2O²⁻ → O₂ + 4e⁻ (oxygen produced reacts with carbon anodes, which must be replaced regularly). This makes aluminium extraction energy-intensive and expensive.
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