Chemical energetics deals with energy changes in reactions. All reactions involve breaking bonds (endothermic) and making bonds (exothermic) — the balance determines whether a reaction is overall exothermic or endothermic.
Exothermic: releases energy to surroundings (temperature rises). Examples: combustion, neutralisation, oxidation. Endothermic: absorbs energy from surroundings (temperature falls). Examples: thermal decomposition, photosynthesis, dissolving ammonium nitrate. Energy changes measured by temperature change of surrounding water.
Energy profile diagrams: show reactants, products, and activation energy (minimum energy to start reaction). Exothermic: products lower than reactants. Endothermic: products higher. Bond energy (Supplement): energy to break 1 mole of a bond. Breaking bonds = endothermic. Making bonds = exothermic. ΔH = Σ(bonds broken) − Σ(bonds made). Negative ΔH = exothermic.
Fossil fuels (coal, oil, gas): combustion releases CO₂ and H₂O. Environmental issues: greenhouse effect, acid rain (SO₂). Hydrogen fuel: H₂ + ½O₂ → H₂O. Only produces water — clean fuel. Used in fuel cells (hydrogen + oxygen → electricity + water). Challenges: storage (flammable gas), production (currently from fossil fuels), infrastructure.
Even though exothermic reactions release energy overall, the existing bonds in reactants must be broken first — and breaking bonds always requires energy. The activation energy is the minimum energy needed to start breaking these bonds and initiate the reaction. Once started, the energy released from forming new bonds is greater than the activation energy, so the reaction is self-sustaining. Catalysts lower the activation energy by providing an alternative reaction pathway.
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