calorimetry
Calorimetry is the art of measuring heat by carefully tracking temperature changes. Since heat itself is invisible, we cannot read it off directly; instead we let heat flow between objects inside a well-insulated container and use the temperature changes to work out how much energy moved. It is essentially heat accounting.
The method rests on conservation of energy. In an isolated system, the heat lost by the hot objects equals the heat gained by the cold ones, so all the Q = m c ΔT terms must sum to zero. The insulated container used to enforce this isolation is called a calorimeter. By mixing substances and measuring the final equilibrium temperature, you can solve for an unknown specific heat, a latent heat, or the energy released by a chemical reaction or by burning food.
A classic experiment: drop a hot metal block of known mass into cool water and wait for thermal equilibrium. Because the heat the metal loses equals the heat the water gains, one measurement of the final temperature lets you calculate the metal's specific heat. Real calorimeters are never perfectly insulated, so careful work must account for small heat leaks to the surroundings.
Drop a 0.1 kg metal block at 100 C into 0.2 kg of water at 20 C; when they settle at, say, 25 C, setting the metal's heat loss equal to the water's heat gain lets you solve for the metal's specific heat.
In an insulated calorimeter, heat lost by the hot object equals heat gained by the cold one.
Calorimetry assumes the container is perfectly isolated so no heat escapes; in reality, small leaks to the surroundings must be corrected for.