Ch 1 introduces chemical reactions, how to write and balance chemical equations, and classifies reactions into combination, decomposition, displacement, double displacement, and redox reactions.
A chemical equation represents a reaction using symbols/formulae. Reactants → Products. Balanced equation: same number of atoms of each element on both sides (law of conservation of mass). States indicated as (s), (l), (g), (aq). Conditions shown above/below the arrow.
Combination: 2Mg + O₂ → 2MgO. Decomposition: 2FeSO₄ → Fe₂O₃ + SO₂ + SO₃ (by heat/light/electricity). Displacement: Fe + CuSO₄ → FeSO₄ + Cu (more reactive displaces less reactive). Double displacement: Na₂SO₄ + BaCl₂ → BaSO₄↓ + 2NaCl (exchange of ions, often forms precipitate).
Oxidation: gain of oxygen or loss of hydrogen. Reduction: loss of oxygen or gain of hydrogen. Redox: both happen simultaneously. Example: CuO + H₂ → Cu + H₂O (CuO is reduced, H₂ is oxidised). Rusting (Fe + O₂ + H₂O → rust) and rancidity (oxidation of fats) are everyday examples.
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A redox reaction is one where both oxidation and reduction happen simultaneously. One substance is oxidised (gains oxygen or loses electrons) while another is reduced (loses oxygen or gains electrons). Example: CuO + H₂ → Cu + H₂O. Here CuO is reduced to Cu and H₂ is oxidised to H₂O.
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