Introduction to enthalpy change

When chemical bonds break or form during a reaction, energy is often released or absorbed as heat. This heat energy change, which occurs at constant pressure, is known as the enthalpy change (ΔH).

Enthalpy change is measured in kilojoules per mole (kJ mol^-1).

Exothermic versus endothermic reactions

Exothermic reactions:

• Release heat energy to the surroundings.
• Have a negative enthalpy change (ΔH is negative).
• Temperature of the surroundings increases. For example, the combustion of methane is exothermic:

CH₄ + 2O₂ ➔ CO₂ + 2H₂O, ΔH^⦵_r = $-$890 kJ mol^-1

Endothermic reactions:

• Absorb heat energy from the surroundings.
• Have a positive enthalpy change (ΔH is positive).
• Temperature of the surroundings decreases. For example, the thermal decomposition of calcium carbonate is endothermic:

CaCO₃ ➔ CaO + CO₂, ΔH^⦵_r = +178 kJ mol^-1^

Enthalpy profile diagrams

Enthalpy profile diagrams show how the energy levels of reactants and products change during a reaction.

Key features include:

• The activation energy is the minimum energy needed for a reaction to occur.
• In exothermic reactions, the reactants have a higher energy than the products and the ΔH arrow points downwards
• In endothermic reactions, the reactants have a lower energy than the products and the ΔH arrow points upwards Substances are more stable when they have less enthalpy.

Why standard conditions are used

Actual enthalpy cannot be directly measured, only the change.

Enthalpy changes are measured under standard conditions:

• 298 K temperature
• 101 kPa pressure

The superscript plimsoll (^⦵) is used after ΔH to denote that the enthalpy change refers to standard conditions. For example, ΔH^⦵_r refers to the standard enthalpy change of reaction.

The standard state of a substance is its physical state under standard conditions.

Types of standard enthalpy change

You need to know the definitions for these four types of standard enthalpy change:

• Standard enthalpy change of reaction (ΔH^⦵r) - Enthalpy change when a reaction occurs in the molar quantities shown in the chemical equation, under standard conditions. For example: N_2(g) + 3H_2(g) ➔ 2NH_3(g) ΔH^⦵_r = $-$92 kJ mol^-1.
• Standard enthalpy change of formation (ΔH^⦵f) - Enthalpy change when 1 mole of a compound forms from elements under standard conditions. For example: C_(s) + O_2(g) ➔ CO_2(g) ΔH^⦵_f = $-$394 kJ mol^-1.
• Standard enthalpy change of combustion (ΔH^⦵c) - Enthalpy change when 1 mole of a substance combusts fully in oxygen under standard conditions. For example: CH_4(g) + 2O_2(g) ➔ CO_2(g) + 2H₂O_(l) ΔH^⦵_c = $-$890 kJ mol^-1.
• Standard enthalpy change of neutralisation (ΔH^⦵neut) - Enthalpy change when an acid and alkali neutralise under standard conditions, to form 1 mole of water. For example: HCl_(aq) + NaOH_(aq) ➔ NaCl_(aq) + H₂O_(l) ΔH^⦵_neut = $-$57 kJ mol^-1.

Using standard conditions means these values will be the same regardless of who measures them.