Measuring enthalpy changesEdexcel A-Level Chemistry: Revision notes

To find an enthalpy change experimentally, the heat energy transferred to (or from) a known mass of water or solution is calculated:

Q = m × c × ΔT

• Q = energy transferred in joules
• m = mass of water or solution in grams (1 cm³ of dilute solution is taken as 1 g)
• c = specific heat capacity, 4.18 J g⁻¹ K⁻¹ for water
• ΔT = temperature change in K (the same size as in °C)

Then ΔH = −Q / n, with Q converted to kJ and n the amount in moles of the reactant that is limiting or burned. The sign is negative if the temperature rises (exothermic) and positive if it falls.

For neutralisation or displacement, mix the solutions in an insulated polystyrene cup with a lid and stir. Record the temperature before mixing, then at regular intervals after.

Plot a temperature–time graph and extrapolate the cooling line back to the time of mixing to find the maximum temperature change, correcting for heat loss.

Assumptions: the solution has the density and specific heat capacity of water, and no heat is lost to the surroundings or absorbed by the cup.

A spirit burner heats water in a metal can. Weigh the burner before and after and measure the water temperature change.

Worked example: 0.60 g of propan-1-ol (Mr = 60.0) heats 150 g of water by 28.0 K.

• Q = 150 × 4.18 × 28.0 = 17 556 J = 17.6 kJ
• n = 0.60 / 60.0 = 0.0100 mol
• ΔcH = −17.6 / 0.0100 = −1760 kJ mol⁻¹

The data book value is about −2021 kJ mol⁻¹, so the experiment under-estimates the energy released.

Experimental combustion values are less exothermic than data book values because of:

• Heat loss to the surroundings, so use a draught shield, lid or insulation
• Heat absorbed by the can and thermometer, so include the heat capacity of the can or calibrate
• Incomplete combustion (soot, carbon monoxide), so less energy is released
• Evaporation of fuel from the wick, so use a lid
• Non-standard conditions, such as water produced as steam rather than liquid

For solutions, errors come from heat loss, the assumptions about density and heat capacity, and the thermometer uncertainty (percentage uncertainty is larger when ΔT is small).

Some enthalpy changes cannot be measured directly because the reaction is too slow, incomplete or does not stop at the wanted product, for example forming carbon monoxide from carbon and oxygen. They are found indirectly from other reactions that can be measured by experiment.