R1.1 Measuring enthalpy changesIB Chemistry SL: Revision notes

In energetics the system is the reacting chemicals and the surroundings are everything else (solvent, container, air). Total energy is conserved: energy lost by the system is gained by the surroundings, and vice versa.

Temperature is a measure of the average kinetic energy of the particles (units °C or K). Heat is energy transferred because of a temperature difference (units J). A large bath at 40 °C contains more thermal energy than a cup at 60 °C, but its temperature is lower.

In an exothermic reaction energy is transferred from the system to the surroundings, so the temperature of the surroundings rises and ΔH is negative (combustion, neutralisation, the hand-warmer oxidation of iron). In an endothermic reaction energy is transferred from the surroundings to the system, so the temperature falls and ΔH is positive (dissolving ammonium nitrate, thermal decomposition of CaCO₃, photosynthesis).

Whether a reaction is exo- or endothermic depends on the relative stability of reactants and products. If the products are more stable (lower enthalpy), the difference is released: exothermic. If they are less stable (higher enthalpy), energy must be absorbed: endothermic.

An energy profile plots enthalpy against the progress of reaction. Exothermic: products below reactants, ΔH arrow down. Endothermic: products above reactants, ΔH arrow up. Both show a hump whose height above the reactants is the activation energy, Ea; in an endothermic profile Ea must be larger than ΔH.

The enthalpy change, ΔH, is the heat transferred at constant pressure. The standard enthalpy change, ΔH⦵, refers to standard conditions: 100 kPa pressure, substances in their standard states (their normal physical state under these conditions) and solutions at 1 mol dm⁻³; a temperature is usually quoted as 298 K. ΔH is given per mole of the reaction as written, in kJ mol⁻¹, always with its sign.

1. Q = mcΔT: m = mass of the substance being heated (usually the water or solution, assuming 1.00 g cm⁻³), c = specific heat capacity (4.18 J g⁻¹ K⁻¹ for water), ΔT = temperature change (same in °C and K).
2. Find the amount (mol) of the limiting reactant, or of fuel burned.
3. ΔH = −Q ÷ n, converting J to kJ. Negative if the temperature rose.

Example: 50.0 cm³ of 1.00 mol dm⁻³ HCl + 50.0 cm³ of 1.00 mol dm⁻³ NaOH, ΔT = 6.6 K: Q = 100.0 × 4.18 × 6.6 = 2.76 kJ; n = 0.0500 mol; ΔH = −55.2 kJ mol⁻¹.

Calorimetry values are usually less exothermic than data booklet values because of heat loss to the surroundings and heat absorbed by the container and thermometer. For fuels, incomplete combustion (soot) and evaporation of the fuel from the wick also lower the result. Calculations assume the solution has the density and specific heat capacity of water. Insulating the cup, using a lid and stirring reduce the errors. Report percentage error = |experimental − accepted| ÷ |accepted| × 100.