Bond enthalpies: Mean bond enthalpy.

The idea

A mean bond enthalpy is the enthalpy change when one mole of a stated covalent bond is broken in the gaseous state, averaged over a range of different compounds containing that bond. The two halves both matter. Gaseous state, because breaking a bond in a liquid would also have to pay for the intermolecular forces. Averaged, because the O−H bond in water is not quite the same strength as the O−H bond in ethanol, so the tabulated value is a compromise that fits neither exactly. Bond breaking always takes energy in, so mean bond enthalpies are always positive and examiners refuse any definition that says energy released. The averaging is also why any ∆H calculated from them is only approximate.

What earns the marks

• Enthalpy change to break one mole of bonds.
• Specifying the gaseous state.
• Averaged over a range of different compounds or molecules.
• Recognising the value as positive, since breaking a bond requires energy.

Common mistakes and how to put them right

• Defining it as the energy released when a bond is formed, which examiners explicitly refuse.

• Omitting the gaseous state, so the definition would include the energy to vaporise the substance.

• Omitting the averaging, which is the point that makes it a mean bond enthalpy rather than a bond dissociation enthalpy.

• Writing per mole of molecules rather than per mole of bonds.

• Assuming the tabulated value applies exactly to the compound in the question.

Exam tips

• Learn the definition as four chunks: enthalpy change, to break one mole of bonds, in the gaseous state, averaged over different compounds.
• Heat energy change is accepted in place of enthalpy change, but enthalpy released is never accepted.
• If the question names a specific molecule, still say averaged over a range of compounds; the average is the defining feature.