4.2.1Additional lessons
Bond Enthalpy Calculations — IB Chemistry SL Revision Notes & Topical Questions
5 min read 6 parts Lesson 51 of 98
Part 1 of 6What Bond Enthalpy Measures
Bond enthalpy lets you estimate the energy change of a reaction before you perform it.
Definition
- Bond enthalpy is the energy needed to break one mole of a particular covalent bond in the gaseous state.
Always positive
- Breaking bonds requires energy,
- so every bond enthalpy is a positive value in kJ mol⁻¹,
- and making the same bond releases the same amount.
Average values
- Tabulated values are averages over many different molecules, so calculations based on them give only approximate results.
Strength trend
- Stronger bonds have larger values, and a double bond is stronger than a single bond between the same two atoms.
Part 2 of 6The Calculation Method
The energy change is the difference between the energy put in and the energy given out.
1
Draw out the structures of every reactant and product so that all the bonds to be broken or made can be seen and counted.
2
List each type of bond in the reactants and multiply its enthalpy by the number of such bonds to find the total energy absorbed.
3
List each type of bond in the products and multiply likewise to find the total energy released when they form.
4
Subtract: ΔH = (sum of bonds broken) − (sum of bonds formed).
5
Give the answer in kJ mol⁻¹ with its sign, negative if more energy is released than absorbed.
Part 3 of 6Limits of the Method
Limits of the Method
Gas phase only
- The data apply to gaseous species, so liquid or solid substances must first be converted to gases for the calculation to be valid.
Average values
- Real bond enthalpies vary slightly with the molecule, so answers differ from experimental values obtained by calorimetry.
State of water
- Using bond enthalpies gives ΔH for water as a gas, which is less exothermic than the value for liquid water.
Use
- The method is most useful for estimating the energy of reactions for which no experimental data are available.
| Bond | Average bond enthalpy / kJ mol⁻¹ |
|---|---|
| C–H | 414 |
| O=O | 498 |
| C=O (in CO₂) | 804 |
| O–H | 463 |
Part 4 of 6Choosing Between Methods
Choosing Between Methods
Bond enthalpies
- Best when all species are gaseous molecules and only an estimate is needed.
Calorimetry
- Gives a measured value for one reaction but suffers from heat loss to the surroundings.
Hess cycles
- Use formation or combustion data and give more accurate values than average bond data.
Consistency check
- Strongly exothermic reactions form stronger bonds than they break, which is a useful sense check on any answer.
Part 5 of 6Worked Example
Question Estimate ΔH for CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g) using the bond enthalpies in the table.
- Bonds broken4 C–H and 2 O=O, giving 4 × 414 + 2 × 498 = 1656 + 996 = 2652 kJ.
- Bonds formed2 C=O and 4 O–H, giving 2 × 804 + 4 × 463 = 1608 + 1852 = 3460 kJ.
- ΔH = 2652 − 3460 = −808 kJ mol⁻¹.
Answer ΔH ≈ −808 kJ mol⁻¹, an exothermic reaction.
Part 6 of 6Exam Checklist & Tips
4.2.1 – Key Points for the Exam
- Bond enthalpy The energy required to break one mole of a given covalent bond in the gas phase.
- Average bond enthalpy The mean value of a bond enthalpy taken over a range of compounds.
- Bonds broken Bonds in the reactants, which absorb energy.
- Bonds formed Bonds in the products, which release energy.
Common Mistakes
- Reversing the subtraction and obtaining the wrong sign; it is broken minus formed.
- Forgetting to multiply a bond enthalpy by the number of such bonds in the molecule.
- Applying bond enthalpies to liquids or solids without allowing for the energy of changes of state.
- Counting a double bond as a single bond, or missing bonds in the structural formula.
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