Bond enthalpies – calculating ΔH of a reaction

What it shows

Breaking a chemical bond always takes in energy, and forming one always releases energy. To estimate the enthalpy change of a gas-phase reaction, imagine two steps: break every bond in the reactants to give free gaseous atoms, then join the atoms into the products. Then ΔH = Σ(bond enthalpies of bonds broken) − Σ(bond enthalpies of bonds formed). If the bonds formed are stronger overall, ΔH is negative and the reaction is exothermic. Because mean bond enthalpies are averages over many compounds, the result differs slightly from the value obtained from standard enthalpies of formation.

How to use

Choose a reaction button, then press Play or Next step to break and form one bond at a time; you can also click the picture. The energy diagram climbs by one bond enthalpy per bond broken and falls per bond formed. Change Speed, or press Reverse reaction to see ΔH change sign. With the quiz option, the totals are hidden: calculate ΔH from the table, type it in Your ΔH and press Check.

Parameters you can change

• Reaction CH₄ + 2O₂ → CO₂ + 2H₂O, H₂ + Cl₂ → 2HCl, 2H₂ + O₂ → 2H₂O, N₂ + 3H₂ → 2NH₃, H₂ + Br₂ → 2HBr, H₂ + F₂ → 2HF, C₂H₄ + H₂ → C₂H₆, CH₄ + Cl₂ → CH₃Cl + HCl, 2HI → H₂ + I₂
• Use the reverse reaction
• Animation speed 0.5–3 steps/s
• Students calculate (hide answer)

Questions to explore

1. Why does breaking a bond always absorb energy while forming a bond always releases it?
2. Why does the bond-enthalpy ΔH for C₂H₄ + H₂ differ more from the data-book value than the others?
3. The F–F bond is very weak. How does this help explain why H₂ + F₂ releases so much energy?