Enthalpy Changes and Calorimetry

This lesson covers the key enthalpy change definitions required for AQA A-Level Chemistry, standard conditions, and the experimental determination of enthalpy changes by calorimetry (specification section 3.1.4).

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Enthalpy (H) and Enthalpy Changes (ΔH)

Key Definition: The enthalpy change (ΔH) of a reaction is the heat energy transferred at constant pressure.

• If ΔH is negative, the reaction is exothermic (releases heat to surroundings, temperature of surroundings increases)
• If ΔH is positive, the reaction is endothermic (absorbs heat from surroundings, temperature of surroundings decreases)

We cannot measure absolute enthalpy, only changes in enthalpy.

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Standard Conditions

Standard enthalpy changes are measured under standard conditions:

• Temperature: 298 K (25°C)
• Pressure: 100 kPa (approximately 1 atm)
• Concentration: 1.00 mol dm⁻³ (for solutions)
• Substances in their standard states (the most stable form at 298 K and 100 kPa)

The symbol for a standard enthalpy change is ΔH° (with the plimsoll sign °).

Standard States

Element	Standard State at 298 K
Carbon	Graphite (not diamond)
Oxygen	O₂(g)
Sulfur	S₈(s) (rhombic)
Phosphorus	P₄(s) (white)
Bromine	Br₂(l)
Mercury	Hg(l)
Iron	Fe(s)

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Key Enthalpy Change Definitions

Standard Enthalpy of Combustion (ΔcH°)

The enthalpy change when one mole of a substance is completely burned in excess oxygen under standard conditions, with all reactants and products in their standard states.

Example: CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l) ΔcH° = −890 kJ mol⁻¹

Key points:

• Always one mole of the substance being burned
• Always exothermic (negative ΔH)
• Products must include CO₂(g) and H₂O(l) — not H₂O(g)

Standard Enthalpy of Formation (ΔfH°)

The enthalpy change when one mole of a compound is formed from its elements in their standard states under standard conditions.

Example: C(graphite) + 2H₂(g) → CH₄(g) ΔfH° = −74.8 kJ mol⁻¹

Key points:

• Always one mole of product
• Elements must be in their standard states
• ΔfH° of any element in its standard state is zero by definition
• Can be exothermic or endothermic

Standard Enthalpy of Neutralisation (ΔneutH°)

The enthalpy change when an acid and a base react to form one mole of water under standard conditions.

Example: HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l) ΔneutH° = −57.1 kJ mol⁻¹

Key points:

• Always produces one mole of H₂O
• Always exothermic (negative ΔH)
• For strong acid + strong base, ΔneutH° ≈ −57 kJ mol⁻¹ regardless of which acid/base
• For weak acids or bases, the value is less exothermic because energy is needed to fully ionise the weak acid/base

Standard Enthalpy of Atomisation (ΔatH°)

The enthalpy change when one mole of gaseous atoms is formed from the element in its standard state.

Examples:

• ½Cl₂(g) → Cl(g) ΔatH° = +121 kJ mol⁻¹
• Na(s) → Na(g) ΔatH° = +107 kJ mol⁻¹
• C(graphite) → C(g) ΔatH° = +717 kJ mol⁻¹

Key points:

• Always one mole of gaseous atoms formed
• Always endothermic (positive ΔH) — bonds/interactions must be broken
• Note that for diatomic elements, the equation involves ½ mole of the element

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Calorimetry: Measuring Enthalpy Changes

Calorimetry uses the equation:

q = mcΔT

Where:

• q = heat energy transferred (J)
• m = mass of the substance being heated (g) — usually water or solution
• c = specific heat capacity (J g⁻¹ K⁻¹) — for water, c = 4.18 J g⁻¹ K⁻¹
• ΔT = temperature change (K or °C — the size of the change is the same)

Worked Example 1: Combustion Calorimetry

Question: 0.500 g of ethanol (C₂H₅OH, Mr = 46.0) was burned and the heat produced raised the temperature of 200 g of water by 13.5°C. Calculate the enthalpy of combustion of ethanol.

Step 1: Calculate heat transferred to water. q = mcΔT = 200 × 4.18 × 13.5 = 11 286 J = 11.286 kJ

Step 2: Calculate moles of ethanol burned. n = mass / Mr = 0.500 / 46.0 = 0.01087 mol