calorimetry
/ kal-uh-RIM-uh-tree /
Calorimetry is the art of measuring heat by watching a temperature change. You can't see heat directly, so the trick is to let the heat flow into a known amount of water (or other material) and read how much its temperature rises. From that rise, and the water's known heat capacity, you work backward to how much heat was released or absorbed. It is energy-measurement by careful eavesdropping on a thermometer.
The instrument that does this is a calorimeter — at its simplest, an insulated cup of water with a thermometer; at its most precise, a sealed steel "bomb" submerged in a stirred water bath. A reaction or a burning sample is run inside, and the surrounding water faithfully soaks up the heat. Because the container is insulated, almost none escapes, so the temperature change reports the energy honestly.
Calorimetry is how the enthalpy and energy values printed in tables are actually pinned down — the kilojoules-per-mole behind a fuel's energy content or a food's calories were each measured this way. It turns the invisible flow of heat into a number you can trust and reuse.
To find the energy in a peanut, burn it under a small can of water and measure how much the water warms. The food calories on a nutrition label come from exactly this kind of bomb-calorimetry measurement.
Burn the sample, measure the water's temperature rise, and the heat released falls out.
A constant-volume "bomb" calorimeter measures the change in internal energy ΔU, because the sealed rigid vessel does no expansion work. A constant-pressure (coffee-cup) calorimeter measures the enthalpy change ΔH. The two differ by the small pressure–volume work, which matters most when gases are produced or consumed.