Calorimetry — Chemistry, 14–17
Calorimetry uses measured temperature changes to find how much energy a reaction transfers. It turns a warm or cool solution into chemical evidence.
What calorimetry measures
When a reaction happens in a solution, energy may move into or out of the liquid. A thermometer records the liquid’s temperature change, and that change lets you estimate the transferred heat. The apparatus is simple, but the measurement is only useful if you know what mass, heat capacity and temperature change belong in the calculation.
Why measure heat
You cannot usually watch energy moving between chemical bonds. Early chemists needed a practical way to compare fuels, foods and reactions, so they measured what happened to a known amount of water. Calorimetry solves that problem: temperature is visible and measurable, while microscopic energy transfers are not.
A simple calculation
A reaction warms 100 g of water from 20.0 °C to 27.0 °C. The change is 7.0 °C. Using q = mcΔT and c = 4.18 J g⁻¹ °C⁻¹, q = 100 × 4.18 × 7.0 = 2926 J, or 2.93 kJ, absorbed by the water. If the reaction caused this warming, it released about 2.93 kJ.
The heat did not vanish
A common mistake is to use only the temperature rise and ignore the cup, thermometer or heat escaping to the air. That seems reasonable because the water is what you can see changing. In reality, the water often records only part of the transfer, so a simple school calorimeter usually underestimates the reaction’s energy.
Where it is used
Food labels use calorimetry to estimate the energy released when food is completely burned. Engineers also test fuels, batteries and building materials this way. The laboratory measurement is not the same as how your body digests food, so a label gives a useful estimate, not a perfect personal prediction.
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