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*Revision Materials and Past Papers* 2.1.1 Atomic structure and isotopes 2.1.2 Compounds, formulae and equations 2.1.3 Amount of substance 2.1.4 Acids 2.1.5 Redox 2.2.1 Electron structure 2.2.2 Bonding and structure 3.1.1 Periodicity 3.1.2 Group 2 3.1.3 The halogens 3.1.4 Qualitative analysis 3.2.1 Enthalpy 3.2.2 Reaction Rates 3.2.3 Chemical equilibrium 4.1 Basic concepts and hydrocarbons 4.1.2 Alkanes 4.1.3 Alkenes 4.2.1 Alcohols 4.2.2 Haloalkanes 4.2.3 Organic synthesis 4.2.4 Analytical techniques 5.1.1 How fast? 5.1.2 How far? 5.1.3 Acids, bases and buffers 5.2.1 Lattice enthalpy 5.2.2 Enthalpy and entropy 5.2.3 Redox and electrode potentials 5.3.1 Transition elements 5.3.2 Qualitative analysis 6.1.1 Aromatic compounds 6.1.2 Carbonyl compounds 6.1.3 Carboxylic acids and esters 6.2.1 Amines 6.2.2 Amino acids, amides and chirality 6.2.3 Polyesters and polyamides 6.2.4 Carbon–carbon bond formation 6.2.5 Organic synthesis 6.3.1 Chromatography and qualitative analysis 6.3.2 Spectroscopy Required Practicals

3.2.1 Enthalpy

Bond enthalpiesEnthalpy changesHess’ law and enthalpy cycles

Bond Enthalpies

Specification Reference 3.2.1 (f)

Quick Notes

  • Bond Enthalpy: Energy to break 1 mole of bonds in the gaseous state.
    • Breaking bonds is endothermic (ΔH is positive, energy absorbed).
    • Making bonds is exothermic (ΔH is negative, energy released).
  • Mean bond enthalpy: average energy required to break a particular type of bond across different compounds.
  • Use mean bond enthalpies to estimate enthalpy changes:
    • ΔH = Σ(Bond enthalpies of bonds broken) − Σ(Bond enthalpies of bonds formed)

Full Notes

Bond Enthalpies

Average (or mean) bond enthalpy is the average energy required to break 1 mole worth of a given bond type in gaseous molecules.

They are calculated using different molecules that have that bond type in.

Example: The C–H bond has a mean bond enthalpy of +412 kJ mol−1. However, the exact bond enthalpy of a C–H bond will vary slightly depending on the specific molecule it is in.

Breaking bonds requires energy = Endothermic process (ΔH is positive).
Making bonds releases energy = Exothermic process (ΔH is negative).

Example:

OCR (A) A-Level Chemistry diagram showing bond enthalpy of H–H bond breaking and forming.

Breaking H–H bond: H2(g) → 2H(g), ΔH = +436 kJ mol−1
Forming H–H bond: 2H(g) → H2(g), ΔH = −436 kJ mol−1

Calculating Enthalpy Change Using Bond Enthalpies

The enthalpy change of a reaction can be estimated using:

OCR (A) A-Level Chemistry formula for enthalpy change calculation using bond enthalpies.

Where:

Note: Calculations using mean bond enthalpies are estimates and may differ from experimental values.

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Matt’s exam tip

Remember bond enthalpies are for substances in the gaseous state. Always check that all species are in gaseous phase when doing bond enthalpy calculations – sometimes enthalpy of vaporisation must be used first.


Worked Example

Calculate the enthalpy change (ΔH) for the combustion of methane, using given bond enthalpies:

  • C–H = +412 kJ mol−1
  • O=O = +498 kJ mol−1
  • C=O = +805 kJ mol−1
  • O–H = +463 kJ mol−1
OCR (A) A-Level Chemistry worked example showing combustion of methane with bond enthalpy calculations.
  1. Bonds Broken (Reactants - Energy Absorbed)
    Bonds in CH4: 4 × C–H = 4 × 412 = 1648 kJ
    Bonds in O2: 2 × O=O = 2 × 498 = 996 kJ
    Total energy to break bonds = 1648 + 996 = 2644 kJ
  2. Bonds Formed (Products - Energy Released)
    Bonds in CO2: 2 × C=O = 2 × 805 = 1610 kJ
    Bonds in H2O: 4 × O–H = 4 × 463 = 1852 kJ
    Total energy released = 1610 + 1852 = 3462 kJ
  3. Calculate Enthalpy Change
    ΔH = Bonds broken − Bonds formed
    ΔH = 2644 − 3462
    ΔH = −818 kJ mol−1 (exothermic reaction)

Summary