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Enthalpy (H) is the heat content of a system measured at constant pressure; the enthalpy change ΔH of a reaction is the heat exchanged with the surroundings when the reaction proceeds at constant pressure. This lesson defines the five standard enthalpy changes you must know for AQA A-Level Chemistry — combustion (ΔcH°), formation (ΔfH°), neutralisation (ΔneutH°), reaction (ΔrH°) and atomisation (ΔatH°) — sets out the standard conditions (100 kPa and a stated temperature, usually 298 K, with substances in their standard states), distinguishes exothermic (ΔH < 0) from endothermic (ΔH > 0) processes using energy-level diagrams, and develops calorimetry as the standard experimental route to measuring ΔH via q = mcΔT. The lesson anchors Required Practical 2 (measurement of an enthalpy change by simple calorimetry); the full RP2 walkthrough — apparatus, cooling-curve extrapolation, uncertainty budget — is treated in lesson 7 of this course.
Spec mapping (AQA 7405): This lesson maps to §3.1.4 (energetics) of the AS specification and connects forward to §3.1.8 (thermodynamics, A2), where Born-Haber cycles and entropy/Gibbs analyses build directly on the same standard-enthalpy framework. It is the foundation for lesson 2 (Hess's law — alternative routes to ΔH when direct measurement is impossible) and lesson 7 (the RP2 deep dive). It depends on §3.1.2.5 (stoichiometry and the mole concept) for the per-mole conversion of measured heat. Refer to the official AQA specification document for the exact wording of each section.
Assessment objectives: Definitions of the five standard enthalpy changes — with their conditions of 100 kPa, stated temperature (usually 298 K), substances in their standard states, and the specific heat capacity of water c = 4.18 J g⁻¹ K⁻¹ — are AO1 recall items. Calorimetry calculations using q = mcΔT, and the conversion of q into ΔH per mole of limiting reagent or fuel, are AO2 staples on Paper 2. AO3 reasoning is tested by evaluating sources of error in calorimetry (heat loss, incomplete combustion, calorimeter heat capacity, evaporation) and explaining why simple-calorimetry values for ΔcH° are typically less exothermic than literature values — a textbook example of the cooling-curve extrapolation technique introduced in lesson 7.
The enthalpy H of a system is its internal energy U plus the product of pressure and volume, H = U + pV. At constant pressure (the normal laboratory condition for solution chemistry, since the atmosphere is effectively a constant-pressure reservoir), the heat q exchanged with the surroundings equals the change in enthalpy: q_p = ΔH. We never measure absolute H — only differences ΔH between defined initial and final states.
The sign convention is from the system's point of view:
Key Point: The surroundings get hotter in an exothermic reaction, but ΔH is negative because we report the change for the system. A common Grade-C error is to write "exothermic, ΔH = +" because temperature went up.
Standard enthalpy changes are measured under standard conditions:
The symbol for a standard enthalpy change is ΔH° (with the plimsoll sign °).
| 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) |
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:
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:
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:
The enthalpy change when one mole of gaseous atoms is formed from the element in its standard state.
Examples:
Key points:
Calorimetry uses the equation:
q = mcΔT
Where:
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
Step 3: Calculate ΔH per mole. ΔH = −q / n = −11.286 / 0.01087 = −1038 kJ mol⁻¹
Step 4: State the sign and compare. ΔcH° = −1038 kJ mol⁻¹ (experimental) Literature value = −1367 kJ mol⁻¹
The experimental value is less exothermic because of heat loss to the surroundings, incomplete combustion, and heat absorbed by the calorimeter.
Question: 50.0 cm³ of 1.00 mol dm⁻³ HCl was mixed with 50.0 cm³ of 1.00 mol dm⁻³ NaOH. The temperature rose by 6.8°C. Calculate the enthalpy of neutralisation. Assume the density of the solution is 1.00 g cm⁻³ and c = 4.18 J g⁻¹ K⁻¹.
Step 1: Total volume = 50.0 + 50.0 = 100.0 cm³ Mass of solution = 100.0 g (assuming density = 1.00 g cm⁻³)
Step 2: q = mcΔT = 100.0 × 4.18 × 6.8 = 2842 J = 2.842 kJ
Step 3: Moles of water formed = 0.0500 mol (limiting reagent: n = c × V = 1.00 × 50.0/1000)
Step 4: ΔneutH = −q / n = −2.842 / 0.0500 = −56.8 kJ mol⁻¹
This is close to the accepted value of −57.1 kJ mol⁻¹.
Question: When 5.00 g of ammonium nitrate (NH₄NO₃, Mr = 80.0) was dissolved in 50.0 g of water, the temperature dropped by 5.0°C. Calculate the enthalpy of solution.
Step 1: q = mcΔT = 50.0 × 4.18 × 5.0 = 1045 J = 1.045 kJ
Step 2: n = 5.00 / 80.0 = 0.0625 mol
Step 3: Temperature dropped, so the process is endothermic. ΔsolH = +q / n = +1.045 / 0.0625 = +16.7 kJ mol⁻¹
Exam Tip: If the temperature increases, the reaction is exothermic (ΔH is negative). If the temperature decreases, the reaction is endothermic (ΔH is positive). Always include the correct sign in your answer.
| Source of Error | Effect |
|---|---|
| Heat loss to surroundings | Measured ΔT is too small → calculated |
| Incomplete combustion | Less heat produced → calculated |
| Heat absorbed by calorimeter | Less heat transferred to water → calculated |
| Evaporation of water | Some heat used for evaporation → measured ΔT is too small |
| Assuming c = 4.18 J g⁻¹ K⁻¹ | Solutions may have a slightly different specific heat capacity |
| Assuming density = 1.00 g cm⁻³ | Dilute solutions have density close to but not exactly 1.00 |
Energy
|
| Reactants
| ___________
| \ Ea (activation energy)
| \ /
| \/
| \
| \___________ Products
| ΔH (negative)
|________________________________
Progress of reaction
The products are at a lower energy level than the reactants. ΔH is negative.
Energy
|
| ___________ Products
| / ΔH (positive)
| /
| /\
| / \ Ea
| __________
| Reactants
|________________________________
Progress of reaction
The products are at a higher energy level than the reactants. ΔH is positive.
Exam Tip: In enthalpy profile diagrams, ΔH is the difference between the energy of the products and the energy of the reactants. The activation energy (Ea) is the minimum energy required for the reaction to occur — it is always positive.