Enthalpy change, ΔH — common mistakes

When a question asks for the definition of a standard enthalpy change (e.g., standard enthalpy of combustion), you must mention standard conditions (100 kPa, 298 K) to get full marks. The plimsoll symbol '⦵' is your cue that standard conditions are in play.

In calorimetry calculations, always show your working clearly in steps: calculate q, calculate n, then calculate ΔH. Pay close attention to signs (is it exothermic or endothermic?), units (J vs kJ), and significant figures. A common error is forgetting to use the mass of the solution (m) for q, not just the mass of the solid dissolved.

Most chemical reactions, especially in a school laboratory, are carried out in open containers, exposed to the atmosphere. This means they occur at a constant atmospheric pressure. At constant pressure, the heat transferred is exactly equal to the change in enthalpy (ΔH). If the pressure wasn't constant, we would also have to account for energy used to do work (expanding or contracting against the changing external pressure), which is more complex.

They are related but distinct. Temperature is a measure of the average kinetic energy of the particles in a substance; it tells us how 'hot' or 'cold' something is. Heat is the energy that is transferred from a hotter object to a colder one. A large swimming pool at 25°C contains far more heat energy than a cup of coffee at 90°C, even though the coffee has a higher temperature, simply because there is so much more of it.

This is a key simplifying assumption. For dilute aqueous solutions, we assume that the solution has the same density (1 g cm⁻³) and specific heat capacity (4.18 J g⁻¹ K⁻¹) as pure water. This is a reasonable approximation because the solution is mostly water. In reality, the dissolved substance does slightly alter these properties, which is one source of inaccuracy in simple calorimetry experiments.