Enthalpy Change from Bond Energies Worked Examples Cheat Sheet

This cheat sheet covers how to estimate enthalpy change from average bond energies in chemical reactions. Students need it because bond energy calculations are a common way to connect molecular structure with energy changes. It is especially useful for checking whether a reaction is exothermic or endothermic.

Worked examples help students organize bonds broken and bonds formed without losing signs.

Key Facts

The main formula is $\Delta H_{\text{rxn}} = \sum E_{\text{bonds broken}} - \sum E_{\text{bonds formed}}$ .
Breaking bonds requires energy, so bonds broken are counted as positive energy changes.
Forming bonds releases energy, so bonds formed are subtracted in the formula.
A negative value of $\Delta H_{\text{rxn}}$ means the reaction is exothermic and releases heat.
A positive value of $\Delta H_{\text{rxn}}$ means the reaction is endothermic and absorbs heat.
Bond energy values are usually measured in $\text{kJ mol}^{-1}$ , so the final $\Delta H$ is often reported in $\text{kJ mol}^{-1}$ of reaction as written.
Coefficients in a balanced equation multiply every bond count in that substance.
Average bond energies give estimates because real bond strengths depend on the molecule and its surroundings.

Vocabulary

Enthalpy change: The heat energy change of a reaction at constant pressure, written as $\Delta H$ .
Bond enthalpy: The energy needed to break one mole of a specified covalent bond in the gas phase.
Bonds broken: The bonds in the reactants that must be separated before new products can form.
Bonds formed: The bonds made in the products, which release energy as atoms become more stable.
Exothermic reaction: A reaction with $\Delta H < 0$ because more energy is released by forming bonds than is absorbed breaking bonds.
Endothermic reaction: A reaction with $\Delta H > 0$ because more energy is absorbed breaking bonds than is released forming bonds.

Common Mistakes to Avoid

Adding bonds formed instead of subtracting them is wrong because bond formation releases energy and lowers $\Delta H$ .
Forgetting to multiply by coefficients is wrong because the balanced equation tells how many molecules and therefore how many bonds are involved.
Counting atoms instead of bonds is wrong because bond energy calculations depend on the number and type of covalent bonds, not just the formula.
Using an unbalanced equation is wrong because the bond totals will not represent the actual reaction stoichiometry.
Reporting the wrong sign is wrong because $\Delta H < 0$ means exothermic and $\Delta H > 0$ means endothermic.

Practice Questions

1 For $\text{H}_2 + \text{Cl}_2 \rightarrow 2\text{HCl}$ , estimate $\Delta H$ using $E_{\text{H-H}} = 436\,\text{kJ mol}^{-1}$ , $E_{\text{Cl-Cl}} = 243\,\text{kJ mol}^{-1}$ , and $E_{\text{H-Cl}} = 431\,\text{kJ mol}^{-1}$ .
2 For $\text{CH}_4 + 2\text{O}_2 \rightarrow \text{CO}_2 + 2\text{H}_2\text{O}$ , count the bonds broken and formed before calculating $\Delta H$ .
3 Estimate $\Delta H$ for $\text{N}_2 + 3\text{H}_2 \rightarrow 2\text{NH}_3$ using $E_{\text{N}\equiv\text{N}} = 945\,\text{kJ mol}^{-1}$ , $E_{\text{H-H}} = 436\,\text{kJ mol}^{-1}$ , and $E_{\text{N-H}} = 391\,\text{kJ mol}^{-1}$ .
4 Explain why average bond energies give approximate values rather than exact experimental enthalpy changes.

Sign in to save

Sign in to save

Enthalpy Change from Bond Energies Worked Examples Cheat Sheet

Related Tools

Related Labs

Related Worksheets

Related Infographics

Study as Flashcards

Key Facts

Vocabulary

Common Mistakes to Avoid

Practice Questions