Question

Explain how bond energy determines whether a chemical reaction is exothermic or endothermic.

Answer 1

Chemical reactions are driven by changes in the energy stored in chemical bonds. Bond energies (bond dissociation energies) tell you how much energy it takes to break a bond; conversely, forming a bond releases that same amount of energy. Whether a reaction is exothermic or endothermic depends on the balance between the energy required to break the reactant bonds and the energy released when the product bonds form.

Basic rule
- Breaking bonds consumes energy (endothermic step).
- Forming bonds releases energy (exothermic step).
- Net enthalpy change (ΔH) ≈ (sum of energies of bonds broken) − (sum of energies of bonds formed).

If:
- Σ(bonds broken) > Σ(bonds formed) → ΔH > 0 → reaction is endothermic (net energy absorbed).
- Σ(bonds broken) < Σ(bonds formed) → ΔH < 0 → reaction is exothermic (net energy released).

Short numerical example (combustion of methane)
CH4 + 2 O2 → CO2 + 2 H2O
Approximate bond energies: C–H ≈ 413 kJ/mol, O=O ≈ 498 kJ/mol, C=O (in CO2) ≈ 799 kJ/mol, O–H ≈ 463 kJ/mol.

- Bonds broken: 4×C–H + 2×O=O = 4×413 + 2×498 = 2648 kJ
- Bonds formed: 2×(2×C=O) + 4×O–H = 2×(2×799) + 4×463 = 3450 kJ
- ΔH ≈ 2648 − 3450 = −802 kJ (net release → exothermic)

Caveats
- Tabulated bond energies are averages (gas phase) so this is an approximation; Hess’s law and standard enthalpies of formation give more accurate ΔH values.
- Bond-energy accounting gives overall energy change (thermodynamics), not the rate of reaction (kinetics). A reaction can be exothermic but still slow if its activation energy is high.
- Noncovalent interactions, solvation, phase changes, and electronic effects can also affect the actual enthalpy.

In short: a reaction is exothermic when forming the product bonds releases more energy than was needed to break the reactant bonds; it is endothermic when more energy is required to break bonds than is released on bond formation.