Inorganic Chemistry (Section 2) covers the patterns and reactions of key groups in the periodic table, the reactivity series, metal extraction, acid-base chemistry, and qualitative analysis.
Group 1 (alkali metals): soft, low density, react with water → metal hydroxide + hydrogen. Reactivity increases down the group (easier to lose outer electron as atom gets larger). Group 7 (halogens): diatomic molecules. Reactivity decreases down (harder to gain electron as atom gets larger). Halogen displacement: more reactive halogen displaces less reactive from solution. Transition metals: form coloured compounds, variable oxidation states, catalysts. Reactivity series: K, Na, Ca, Mg, Al, C, Zn, Fe, H, Cu, Ag, Au. Displacement: more reactive metal displaces less reactive from compound. Extraction: reactive metals by electrolysis (e.g., Al from bauxite), less reactive by reduction with carbon (e.g., Fe in blast furnace).
Acids: produce H⁺ ions in solution. pH < 7. Bases: react with acids to form salt + water (neutralisation). Alkali: soluble base, produces OH⁻ ions. pH > 7. Reactions: acid + metal → salt + H₂; acid + base → salt + water; acid + carbonate → salt + water + CO₂. HCl forms chlorides, H₂SO₄ forms sulfates, HNO₃ forms nitrates. Salt preparation: insoluble salts by precipitation (mixing solutions); soluble salts by adding excess insoluble reactant then filtering. Tests: flame tests (Li red, Na yellow, K lilac, Cu green). NaOH to precipitate metal hydroxides. Test gases: H₂ (squeaky pop), O₂ (relights splint), CO₂ (limewater turns milky), Cl₂ (bleaches litmus).
Very reactive metals (K, Na, Ca, Mg, Al) have a stronger attraction to their non-metal partners than carbon does. Carbon can only reduce oxides of metals below it in the reactivity series. Because aluminium, for instance, is above carbon, carbon cannot remove the oxygen from aluminium oxide — there\'s not enough energy. Instead, electrolysis is used: the ionic compound is melted (or dissolved) so ions are free to move, then electricity forces the positive metal ions to the cathode where they gain electrons (reduction). This requires a lot of electrical energy, making it expensive — which is why aluminium and sodium are more costly to produce than iron.
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