What is a calorimeter & what are its limitations?

Chemists often need to know how much heat energy is released or absorbed by a particular reaction. This measurement helps them to understand more about why the reaction occurs and make useful predictions.

Calorimeters are instruments that measure the amount of heat released or absorbed by the contents during a reaction. It's easy to make a simple calorimeter, although the instruments used in labs are typically more precise.


Basically, a calorimeter measures the change in temperature of the calorimeter and its contents. If the calorimeter has been calibrated, the chemist will already have a number called the calorimeter constant, which tells her how much the temperature of the calorimeter changes per amount of heat added. Using this information and the mass of the reactants, the chemist can determine how much heat was released or absorbed. It's important that the calorimeter minimise the rate of heat loss to the outside, since rapid heat loss to the surrounding air would skew the results.


It's easy to make a simple calorimeter yourself; all you need are two styrofoam coffee cups, a thermometer or a lid. This coffee-cup calorimeter is surprisingly reliable and thus is a common feature of undergraduate chemistry labs. More sophisticated instruments are found in physical chemistry laboratories, where chemists make use of so-called "bomb calorimeters." In these devices, the reactants are isolated in a sealed chamber called the bomb; an electrical spark ignites them, and the change in temperature can be used to determine the heat lost or gained.


To calibrate a calorimeter, you use a process that transfers a known amount of heat -- measuring the temperature of some hot and cold water, for example, then mixing them in your coffee-cup calorimeter. Next, you measure the temperature over time and use linear regression to calculate the "final temperature" of the calorimeter and its contents. Subtracting the heat gained by the cold water from the heat lost by the hot water yields the heat gained by the calorimeter. Dividing this figure by the temperature change of the calorimeter gives its calorimeter constant, which can then be used in subsequent experiments.


No calorimeter is perfect; they lose heat to their surroundings. Although bomb calorimeters in laboratories are insulated to minimise these loses, heat loss cannot, of course, be eradicated entirely. Moreover, the reactants in the calorimeter may not be well-mixed, which leads to uneven heating and another possible source of error in your measurements.

Aside from possible sources of error, another limitation involves the kinds of reactions you can study. You might want, for example, to know how much heat is released through the decomposition of TNT. This kind of reaction would be impossible to study in a coffee-cup calorimeter; however, and might not even be practical in a bomb calorimeter. Alternatively, a reaction may take place very slowly -- like the oxidation of iron to form rust. This kind of reaction would be very difficult to study with a calorimeter.