5.3. Measurement of ∆U and ∆H: Calorimetry

We can measure energy changes associated with chemical or physical processes by an experimental technique called calorimetry. In calorimetry, the process is carried out in a vessel called calorimeter, which is immersed in a known volume of a liquid. Knowing the heat capacity of the liquid in which the calorimeter is immersed and the heat capacity of calorimeter, it is possible to determine the heat evolved in the process by measuring temperature changes. Measurements are made under two different conditions: i) at constant volume, qV ii) at constant pressure, qp.

Calorimetry is a method used to measure the heat involved in chemical reactions or physical changes. A calorimeter is a device that helps us monitor these changes. It works by immersing a calorimeter in a liquid that has a known heat capacity. This allows us to measure any temperature changes that occur when a reaction happens, helping us quantify how much energy is either absorbed or released by the system in the form of heat. We commonly perform these measurements either at a constant volume (where the volume doesn't change) or at constant pressure (like in most real-life scenarios).

Imagine boiling water to make a cup of tea. You place a pot on the stove (the calorimeter) filled with water (the liquid). As you heat the pot, it absorbs energy from the stove, causing the temperature of the water to rise. By measuring how much the temperature increases and knowing the amount of water, you can calculate how much energy was used to heat the water, just like calorimetry helps scientists measure energy changes in reactions.

For chemical reactions, heat absorbed at constant volume is measured in a bomb calorimeter. Here, a steel vessel (the bomb) is immersed in a water bath. The whole device is called calorimeter. The steel vessel is immersed in water bath to ensure that no heat is lost to the surroundings. A combustible substance is burnt in pure dioxygen supplied in the steel bomb. Heat evolved during the reaction is transferred to the water around the bomb and its temperature is monitored. Since the bomb calorimeter is sealed, its volume does not change; i.e., the energy changes associated with reactions are measured at constant volume. Under these conditions, no work is done as the reaction is carried out at constant volume in the bomb calorimeter. Even for reactions involving gases, there is no work done as ∆V = 0. Temperature change of the calorimeter produced by the completed reaction is then converted to qV, by using the known heat capacity of the calorimeter with the help of equation.

In a bomb calorimeter, we measure the heat energy changes during a reaction at constant volume. This apparatus consists of a steel container that holds the reacting materials and is submerged in water to maintain an accurate reading of energy transfer. When a combustible substance is burned in the calorimeter, the energy released is absorbed by the surrounding water, which raises its temperature. By measuring this temperature change and knowing the heat capacity of the system, we can quantify the heat released during the chemical reaction, using the equation qV. Notably, during this reaction, the volume remains unchanged; therefore, the work done (∆V) is zero, simplifying our calculations.

Consider a pressure cooker that can retain heat without allowing steam to escape, similar to a bomb calorimeter. When food is cooked inside, it uses heat (energy) to change the state of the food without any steam escaping. Just like how the calorimeter measures heat transfer without losing energy to the surroundings, the pressure cooker effectively keeps the heat within, making the cooking process efficient.

Measurement of heat change at constant pressure (generally under atmospheric pressure) can be done in a calorimeter. We know that ∆H = qp (at constant p) and, therefore, heat absorbed or evolved, qp at constant pressure is also called the heat of reaction or enthalpy of reaction, ∆rH.

To measure the heat change at constant pressure, we utilize a calorimeter designed for that purpose, often referred to as a constant-pressure calorimeter. In this scenario, we measure the heat absorbed or released during a chemical reaction, known as enthalpy change (∆H). The relationship here signifies that at constant pressure, the heat exchanged with the surroundings (q) is equivalent to the change in enthalpy (∆H). This is an important aspect of thermodynamic principles, allowing chemists to find out how heat affects reactions under natural atmospheric conditions.

Think about making soup in a pot over the stove. The heat from the stove keeps the pressure inside the pot constant while the soup simmers. During cooking, the soup absorbs heat, which increases its temperature and changes its state as it cooks - this is similar to measuring enthalpy change in a calorimeter. By tasting the soup at different times, you can actually feel the heat change, just as the calorimeter measures it.

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