Energy in Reactions
Year 9 ⚗️ Chemical Reactions Distinguish exothermic and endothermic; use bond energies to calculate ΔH.
🌡️ Exothermic & Endothermic
Reactions involve energy changes (enthalpy change, ΔH).
Energy Changes
$$\text{Exothermic: } \Delta H < 0 \quad \text{(releases energy, temperature rises)}$$
$$\text{Endothermic: } \Delta H > 0 \quad \text{(absorbs energy, temperature falls)}$$Exothermic: combustion, neutralisation, respiration, hand warmers.
Endothermic: thermal decomposition, dissolving NH₄NO₃, photosynthesis, cold packs.
Endothermic: thermal decomposition, dissolving NH₄NO₃, photosynthesis, cold packs.
🔗 Bond Energies
Energy is needed to break bonds (endothermic step); energy is released when bonds form (exothermic step).
ΔH from Bond Energies
$$\Delta H = \sum E(\text{bonds broken}) - \sum E(\text{bonds formed})$$H₂ + Cl₂ → 2HCl:
Break: H-H (436) + Cl-Cl (242) = 678 kJ
Form: 2 × H-Cl (2×431) = 862 kJ
ΔH = 678 − 862 = −184 kJ/mol (exothermic)
Break: H-H (436) + Cl-Cl (242) = 678 kJ
Form: 2 × H-Cl (2×431) = 862 kJ
ΔH = 678 − 862 = −184 kJ/mol (exothermic)
📈 Energy Profile Diagrams
An energy profile shows the energy of reactants, transition state and products.
The peak = transition state. Ea = peak energy − reactants energy. A catalyst lowers the peak (lower Ea), making the reaction faster.
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Interactive Demonstration — Energy in Reactions
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