Enthalpy Definitions Revision Notes | A Level AQA Chemistry

1. Why Precise Definitions Matter in Thermodynamics

Every enthalpy definition in this topic is a precisely worded instruction. Change one word — swap 'gaseous' for 'aqueous', or drop 'one mole' — and you are describing a completely different energy change. Examiners know this, and they test it deliberately.

This lesson builds a toolkit of seven enthalpy definitions that underpin Born-Haber cycles, lattice enthalpy calculations, and enthalpy of solution. Each definition has a specific starting state, a specific ending state, and a specific quantity. Nail those three things and the marks follow.

2. Standard Enthalpy of Formation

Picture building a compound from scratch using only the raw elemental ingredients. The enthalpy change for that construction process — starting from elements in their standard states and ending with exactly one mole of the compound — is the standard enthalpy of formation, symbol $\Delta_f H^\circ$.

The standard state of an element is its most stable physical form at 298 K and 100 kPa. Carbon's standard state is graphite, not diamond. Bromine's standard state is liquid, $\text{Br}_2\text{(l)}$, not gaseous. Getting the standard state wrong changes the equation and therefore the enthalpy value.

By definition, the standard enthalpy of formation of any element in its standard state is zero. This is a reference point, not a measured value — it is the baseline from which all formation enthalpies are measured.

3. Enthalpy of Atomisation and Ionisation Energy

To build a Born-Haber cycle, you need to convert solid metals and diatomic gases into individual gaseous atoms and then into gaseous ions. Two definitions handle this conversion: enthalpy of atomisation and ionisation energy.

Once you have gaseous atoms, the next step in a Born-Haber cycle is removing electrons to form positive ions. That process is described by ionisation energy.

4. Bond Enthalpy and Electron Affinity

Bond enthalpy and electron affinity sit on opposite sides of the energy picture: one describes breaking covalent bonds in the gas phase, the other describes adding electrons to gaseous atoms or ions.

Now consider the non-metal side of a Born-Haber cycle. After atomisation produces gaseous non-metal atoms, electrons must be added to form negative ions. This is described by electron affinity.

This sign difference between first and second electron affinities is a favourite examiner trap. For oxygen forming $\text{O}^{2-}$, the first EA is exothermic but the second EA is endothermic — yet the lattice enthalpy of formation of MgO is large enough to make the overall cycle energetically feasible.

5. Lattice Enthalpy — Two Conventions, One Concept

Lattice enthalpy is the energy term that makes ionic compounds stable. It describes the strength of the ionic lattice — the electrostatic forces holding the array of positive and negative ions together. Here is where you must be careful: AQA recognises two valid definitions, and they have opposite signs.

The two values are numerically identical but opposite in sign. If the lattice enthalpy of formation of NaCl is $-787$ kJ mol$^{-1}$, then the lattice enthalpy of dissociation is $+787$ kJ mol$^{-1}$. When you use a Born-Haber cycle, you must be consistent: choose one convention and stick to it throughout.

A more negative (more exothermic) lattice enthalpy of formation indicates a stronger ionic lattice. Lattice enthalpy becomes more exothermic as ionic charge increases and as ionic radius decreases, because both factors increase the electrostatic attraction between ions.

6. Enthalpy of Hydration

When an ionic compound dissolves in water, the gaseous ions released from the lattice become surrounded by water molecules. The energy released during this process is the enthalpy of hydration. It connects directly to lattice enthalpy in the enthalpy of solution cycle.

The enthalpy of hydration becomes more exothermic as ionic charge increases and as ionic radius decreases — exactly the same trend as lattice enthalpy. A small, highly charged ion polarises water molecules more strongly, forming stronger ion-dipole interactions and releasing more energy.

The enthalpy of solution links lattice enthalpy and enthalpy of hydration via a Hess's Law cycle: $\Delta_{sol} H^\circ = \Delta_{latt} H^\circ \text{(dissociation)} + \Delta_{hyd} H^\circ \text{(cation)} + \Delta_{hyd} H^\circ \text{(anion)}$. Whether dissolving is exothermic or endothermic overall depends on the relative magnitudes of these terms.

7. Comparing All Seven Definitions

You now have all seven definitions. Before moving to exam technique, it helps to see them side by side — particularly their starting states, ending states, and typical sign. This is the kind of comparison an examiner might ask you to make directly.

Notice the pattern: any process that forms bonds or attractions (lattice formation, hydration, first electron affinity) releases energy and is exothermic. Any process that breaks bonds or separates charges (atomisation, ionisation, bond breaking, lattice dissociation, second electron affinity) requires energy and is endothermic.

8. Exam Technique and Common Pitfalls

Definition questions at A-level are unforgiving. A mark scheme will award marks for specific key phrases, and missing one phrase loses the mark even if the rest of the answer is correct. The three most tested phrases across all seven definitions are: one mole, gaseous (for the correct state), and standard conditions.

Every definition in this topic is a precisely bounded energy change: a fixed starting state, a fixed ending state, and exactly one mole of the specified species. Mastering these boundaries — and the sign conventions that follow from them — gives you the foundation for every Born-Haber cycle and enthalpy of solution calculation at A-level.