Calculating energy changes using bond energies

Energy changes • Exothermic and endothermic reactions

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Why is ΔH calculated as broken minus formed, not formed minus broken?

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Because energy input is required to break bonds (positive), and energy release from forming bonds reduces the net input; broken minus formed yields the net energy absorbed.

Key concepts

What you'll likely be quizzed about

Definition of bond energy

Bond energy represents the average energy required to break one mole of a particular bond in gaseous molecules, measured in kilojoules per mole (kJ mol−1). Bond energies are tabulated values that represent averages across different molecules and bonding environments. Bond energy values apply as approximations and do not capture variations caused by molecular context, phase, or temperature.

Energy transfer during bond breaking and bond formation

Breaking chemical bonds requires an input of energy because atoms move from a lower-energy bonded state to higher-energy separated atoms. Forming chemical bonds releases energy because atoms move from higher-energy separated states to a lower-energy bonded state. The net energy transferred during a reaction equals the total energy absorbed to break bonds minus the total energy released on bond formation.

Step-by-step calculation method

Step 1: Write a balanced equation and draw structures to identify all bonds broken and formed. Step 2: Use a bond energy table to find the average bond energy for each bond type present, expressed in kJ mol−1. Step 3: Multiply each bond energy by the number of those bonds broken or formed and sum separately. Step 4: Calculate the enthalpy change using ΔH = Σ(bond energies broken) − Σ(bond energies formed). A positive ΔH indicates net energy absorption (endothermic). A negative ΔH indicates net energy release (exothermic).

Worked numerical example

Example: H–H + Cl–Cl → 2 H–Cl. Average bond energies: H–H = 436 kJ mol−1, Cl–Cl = 243 kJ mol−1, H–Cl = 432 kJ mol−1. Energy absorbed = 436 + 243 = 679 kJ mol−1. Energy released = 2 × 432 = 864 kJ mol−1. ΔH = 679 − 864 = −185 kJ mol−1. The negative sign indicates an overall exothermic reaction with 185 kJ mol−1 released.

Limitations and accuracy of bond energy calculations

Tabulated bond energies are averages and do not account for differences in molecular environment, bond angles, resonance, or intermolecular interactions. Calculations assume gaseous reactants and products; values differ for condensed phases. Calculated enthalpy changes give useful estimates but can differ from experimental enthalpies by tens of kJ mol−1. High-accuracy work requires experimental enthalpies or calculations that use formation enthalpies or quantum-chemical methods.

Key notes

Important points to keep in mind

Always list every bond broken and every bond formed from full molecular structures.

Use ΔH = Σ(bond energies broken) − Σ(bond energies formed) and include correct multipliers for repeated bonds.

Report the answer in kJ mol−1 and include the sign to indicate exothermic or endothermic behaviour.

Treat tabulated bond energies as averages; expect deviations from experimental enthalpies.

Avoid applying covalent bond-energy values to ionic lattices or metallic bonds; use appropriate data sources.