Energetics: Enthalpy Change, Calorimetry, Hess's Law and Bond Enthalpies

This guide covers 3.1.4, Energetics, from the Physical chemistry section of OxfordAQA International AS and A-level Chemistry 9620. It follows Molecular shapes, bond polarity and intermolecular forces and covers how the energy changes of a reaction can be measured, defined and calculated.

Enthalpy change

You need to know that reactions can be endothermic (absorbing heat) or exothermic (releasing heat), and that enthalpy change (ΔH) is the heat energy change measured under conditions of constant pressure. Standard enthalpy changes refer specifically to standard conditions – 100 kPa and a stated temperature, conventionally 298 K, shown as ΔH°₂₉₈. You need to be able to define two standard enthalpy changes precisely: standard enthalpy of combustion (ΔcH°), the enthalpy change when one mole of a substance is completely burned in oxygen under standard conditions, and standard enthalpy of formation (ΔfH°), the enthalpy change when one mole of a compound is formed from its constituent elements in their standard states under standard conditions.

Calorimetry

You need to use the equation q = mcΔT, where q is heat change, m is the mass of substance undergoing temperature change ΔT, and c is specific heat capacity, to calculate the molar enthalpy change for a reaction and in related calculations. As with the Avogadro constant and gas constant elsewhere in this specification, you will not be expected to recall the value of specific heat capacity – it will be provided in any question that requires it. This subtopic carries Required Practical 2: measuring an enthalpy change, typically by combining q = mcΔT with a measured temperature change to find the heat released or absorbed, then converting to a molar quantity using the moles of substance reacted.

Hess’s law

You need to use Hess’s law to perform calculations, including calculating enthalpy changes for reactions from enthalpies of combustion or from enthalpies of formation. Hess’s law states that the total enthalpy change for a reaction is independent of the route taken between reactants and products, as long as the initial and final conditions are the same – meaning an enthalpy change that cannot be measured directly can be calculated indirectly via an alternative route of known enthalpy changes, typically drawn as an energy cycle.

Bond enthalpies

You need to define mean bond enthalpy, and use mean bond enthalpies to calculate an approximate value of ΔH for reactions in the gaseous phase. A specific and frequently examined point of understanding is required here: you need to be able to explain why values calculated from mean bond enthalpies differ from those determined using Hess’s law. Mean bond enthalpies are average values taken across many different molecules containing a given bond type, so a specific bond in a specific molecule may have a slightly different actual bond enthalpy than the mean value used in the calculation – this makes the bond-enthalpy method only an approximation, whereas a Hess’s law calculation using measured enthalpies of formation or combustion is specific to the actual compounds involved in that particular reaction, rather than an average taken across many different molecules.

A worked example

To find the enthalpy of formation of methane using Hess’s law from enthalpies of combustion: construct a cycle with the direct formation reaction (C(s) + 2H₂(g) → CH₄(g)) on one side, and an indirect route via complete combustion of the elements and of methane on the other. Since Hess’s law states both routes must give the same overall enthalpy change, ΔfH(CH₄) = [ΔcH(C) + 2 × ΔcH(H₂)] − ΔcH(CH₄). This general pattern – sum of combustion enthalpies of reactants minus combustion enthalpy of the product – applies to any enthalpy of formation calculated from combustion data, and is worth learning as a reusable structure rather than re-deriving the cycle from scratch every time.

Why energetics matters beyond this subtopic

The enthalpy concepts introduced here are not confined to Physical chemistry in isolation. Standard enthalpies of formation and combustion reappear directly whenever a later part of the course asks you to evaluate the energetics of a specific reaction type, and the Hess’s law cycle method generalises to any situation where a wanted enthalpy change cannot be measured directly but can be constructed from known steps. The distinction between a Hess’s law value specific to the compounds actually involved and an approximate, averaged bond-enthalpy value is also a recurring theme in how this specification treats experimental versus theoretical data more broadly – recognising which type of value a question is asking you to calculate, and why the two methods can disagree, is a transferable skill rather than a one-off fact about energetics specifically.

How to approach it

Learn the precise wording of the standard enthalpy of combustion and formation definitions, including “one mole,” “standard conditions,” and (for formation) “from its constituent elements in their standard states” – definition questions are marked closely against this exact phrasing. Practise drawing Hess’s law energy cycles as a first step in any indirect enthalpy calculation, since a correctly drawn cycle makes the arithmetic almost automatic, while attempting the calculation without one is a common source of sign errors. For bond enthalpy calculations, remember the reaction must be in the gaseous phase for mean bond enthalpies to apply directly, and always be ready to explain the Hess’s-law-versus-bond-enthalpy discrepancy in terms of mean values being averaged across many molecules rather than measured for the specific bond in the specific reaction being studied.

Official syllabus

OxfordAQA, International AS and A-level Chemistry (9620) specification, Version 4.3, for International AS and A-level exams May/June 2020 onwards: official specification PDF, section 3.1.4 “Energetics”. Verified 2026-09-02.