AQA · A-Level · Chemistry
Course version: 7405 · Version from Sept 2016
Students should be able to: define the term mean bond enthalpy use mean…
Full specification wording
Bond enthalpies: Students should be able to: define the term mean bond enthalpy use mean bond enthalpies to calculate an approximate value of ∆H for reactions in the gaseous phase explain why values from mean bond enthalpy calculations differ from those determined using Hess’s law.
Mean bond enthalpy is the average enthalpy change to break one mole of a specified covalent bond in the gaseous state, averaged across many different compounds.
Checked against the AQA 7405 specification (opens in a new tab) on .
- Point 2 of 2
- 4 key terms
Study notes
From the checked teaching notes for this specification point.
The idea
Mean bond enthalpy is the average enthalpy change to break one mole of a specified covalent bond in the gaseous state, averaged across many different compounds. To use these values, treat the reaction as pulling reactant molecules apart into atoms and reassembling them: ∆H = Σ(bonds broken) − Σ(bonds made). For example, calculating ∆H for the combustion of methane by counting all C-H, O=O, C=O, and O-H bonds. Count each bond type from displayed formulae and multiply by balancing numbers. The answer often disagrees with a Hess's law value because tabulated figures are averages, not exact values for specific molecules, and the method assumes all species are gaseous, ignoring intermolecular forces in liquids or solids.
What earns the marks
- Defining mean bond enthalpy as the energy to break one mole of a covalent bond in the gas phase, averaged across compounds.
- The correct expression, total bonds broken minus total bonds made.
- Counting every bond in each species and multiplying by the balancing numbers.
- Explaining a difference from a Hess's law value as averaging over compounds, or as data applying only to the gaseous state.
Common mistakes and how to put them right
Reversing the expression to bonds made minus bonds broken, which gives the right size and the wrong sign.
Missing the unchanged bonds in a spectator part of the molecule, or double counting the C−H bonds in a large molecule.
Explaining the discrepancy as experimental error rather than as a consequence of using averaged, gas phase data.
Exam tips
- Draw the displayed formulae of everything in the equation and tick off bonds as you count them.
- When asked why the value differs from one found by Hess's law, give both reasons: averaging, and the assumption that all species are gaseous.
Try the lesson
The start of the Bond enthalpies: Students should be able to: define the term mean bond enthalpy use mean bond enthalpies to calculate an approximate value of ∆H for reactions in the gaseous phase explain why values from mean bond enthalpy calculations differ from those determined using Hess’s law. lesson, made from checked teaching notes for this specification point.
Have a go first
Which statement defines the term mean bond enthalpy?
Listen
Mean bond enthalpy is the average enthalpy change to break one mole of a specified covalent bond in the gaseous state, averaged across many different compounds. To use these values, treat the reaction as pulling reactant molecules apart into atoms and reassembling them: ∆H = Σ(bonds broken) − Σ(bonds made). For example, calculating ∆H for the combustion of methane by counting all C-H, O=O, C=O, and O-H bonds. Count each bond type from displayed formulae and multiply by balancing numbers. The answer often disagrees with a Hess's law value because tabulated figures are averages, not exact values for specific molecules, and the method assumes all species are gaseous, ignoring intermolecular forces in liquids or solids. A strong answer earns marks for these points. defining mean bond enthalpy as the energy to break one mole of a covalent bond in the gas phase, averaged across compounds. the correct expression, total bonds broken minus total bonds made. counting every bond in each species and multiplying by the balancing numbers. explaining a difference from a Hess's law value as averaging over compounds, or as data applying only to the gaseous state. Exam tip: Draw the displayed formulae of everything in the equation and tick off bonds as you count them.
Key words
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mean bond enthalpy
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bonds broken minus bonds made
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gaseous phase assumption
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averaged data
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See every definition
- mean bond enthalpy
- The average enthalpy change to break one mole of a specified covalent bond in the gaseous state, averaged across many different compounds.
- bonds broken minus bonds made
- The expression used to calculate an approximate ∆H, taking the total bonds broken minus the total bonds made.
- gaseous phase assumption
- The method assumes all species are gaseous, ignoring intermolecular forces in liquids or solids.
- averaged data
- Tabulated figures are averages, not exact values for specific molecules, so the answer often disagrees with a Hess's law value.
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Test your understanding of Bond enthalpies: Students should be able to: define the term mean bond enthalpy use mean bond enthalpies to calculate an approximate value of ∆H for reactions in the gaseous phase explain why values from mean bond enthalpy calculations differ from those determined using Hess’s law.
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