Calorimetry

Calorimetry is the experimental measurement of heat changes, using apparatus such as a copper can or an insulated polystyrene cup with a thermometer, and applying Q = mcθ to the temperature change. Malay: kalorimetri · Chinese: 量热法.

The idea is to let the heat of a reaction pass into a known mass of water or solution and read the temperature change, from which Q = mcθ gives the heat energy. The choice of apparatus depends on the reaction: a polystyrene cup insulates a neutralisation, displacement or precipitation in solution, while a copper can holds the water heated by a burning fuel because copper conducts heat well.

Example. To measure a heat of neutralisation, you place acid in a polystyrene cup, record its temperature, add the alkali, stir, and read the highest temperature reached. The rise, the total volume as mass, and the specific heat capacity go into Q = mcθ; dividing by the moles of water formed gives the result per mole.

Confusion to avoid. Calorimetry results are always a little off the accepted values because some heat escapes to the surroundings and, for combustion, some fuel burns incompletely. Do not treat the experimental value as exact; part of the reasoning marks come from explaining these heat losses and how insulation reduces them.

Calorimetry is the practical thread that ties the whole chapter together, and SPM Chemistry may test it either as a calculation or as an experiment you must describe and evaluate. The main assumptions in every school calorimetry experiment are that all the heat of the reaction passes into the solution or water alone, that no heat is lost to the surroundings, and that the solution behaves like pure water; stating these assumptions shows you understand the limits of the method. Covering the container, stirring evenly, and reading the highest temperature quickly are the practical steps that reduce error, and they are worth naming whenever you are asked to improve the experiment.