Exothermic & Endothermic Reactions

Key Definitions

  • Exothermic reaction: releases energy to surroundings; products have lower energy than reactants; \Delta H < 0.
  • Endothermic reaction: absorbs energy from surroundings; products have higher energy than reactants; \Delta H > 0.
  • Activation energy EaE_a: minimum energy required to start any reaction.

Energy Level Diagrams

  • Exothermic: reactant line higher than product line; downward arrow shows ΔH\Delta H (negative).
  • Endothermic: product line higher than reactant line; upward arrow shows ΔH\Delta H (positive).
  • The peak above reactants = transition state; height = EaE_a.

Representative Equations

  • Exothermic (general): reactants→products+energy\text{reactants} \rightarrow \text{products} + \text{energy}.
  • Endothermic (general): reactants+energy→products\text{reactants} + \text{energy} \rightarrow \text{products}.
  • Example exothermic: Fe(s)+S(s)→FeS(s)ΔH=−100  kJ mol−1.\text{Fe}(s) + \text{S}(s) \rightarrow \text{FeS}(s)\quad \Delta H = -100\;\text{kJ mol}^{-1}.
  • Example endothermic: CaCO<em>3(s)→CaO(s)+CO</em>2(g)ΔH=+178  kJ mol−1.\text{CaCO}<em>3(s) \rightarrow \text{CaO}(s) + \text{CO}</em>2(g)\quad \Delta H = +178\;\text{kJ mol}^{-1}.

Bond Energies (kJ mol−1^{-1})

  • H–H=436\text{H–H}=436, Cl–Cl=242\text{Cl–Cl}=242, H–Cl=431\text{H–Cl}=431, C–C=346\text{C–C}=346, C=C=612\text{C=C}=612, C–H=413\text{C–H}=413, C–O=358\text{C–O}=358, O=O=498\text{O=O}=498, N  =  N=946\text{N\;=\;N}=946, N–H=391\text{N–H}=391.

Calculating ΔH\Delta H

  • Formula: ΔH=∑(energy to break bonds)−∑(energy released when bonds form)\Delta H = \sum \text{(energy to break bonds)} - \sum \text{(energy released when bonds form)}.
  • Sign rules:
    • \Delta H < 0 ⇒ exothermic (temperature rises).
    • \Delta H > 0 ⇒ endothermic (temperature falls/needs heating).
  • Worked example: H<em>2+Cl</em>2→2HCl\text{H}<em>2 + \text{Cl}</em>2 \rightarrow 2\text{HCl}
    • Energy in: 1×436+1×242=678  kJ1\times436 + 1\times242 = 678\;\text{kJ}.
    • Energy out: 2×431=862  kJ2\times431 = 862\;\text{kJ}.
    • ΔH=678−862=−184  kJ\Delta H = 678 - 862 = -184\;\text{kJ} (exothermic).

Quick Recall Rules

  • "Exo = Exit": energy exits system; negative ΔH\Delta H.
  • "Endo = Into": energy enters system; positive ΔH\Delta H.
  • Breaking bonds absorbs energy; making bonds releases energy.
  • Compare energy absorbed vs released to classify reaction quickly.