Question

(a) Why is the change in enthalpy usually easier to measure than the change in internal energy? (b) For a given process at constant pressure, \(\Delta H\) is negative. Is the process endothermic or exothermic?

Short answer

(a) Measuring the change in enthalpy (\(\Delta H\)) is usually easier than the change in internal energy (\(\Delta U\)) because it can be measured directly through a calorimeter at constant pressure, while measuring \(\Delta U\) requires assessing changes in kinetic and potential energy of the system's molecules, which is challenging to do in a typical laboratory. (b) If the change in enthalpy at constant pressure, \(\Delta H\), is negative for a given process, the process is exothermic, as it releases heat to the surroundings.

Step-by-step solution

(a) Comparison of measuring enthalpy and internal energy changes

Enthalpy (H) is a thermodynamic property of a system, defined as the sum of its internal energy (U) and the product of pressure (P) and volume (V), that is, \(H = U + PV\). On the other hand, internal energy is the complete energy content of the molecules in the system, like kinetic and potential energy. The change in enthalpy (\(\Delta H\)) for a process, such as a chemical reaction, can be measured easily by carrying out the process at constant pressure in a device known as a calorimeter. The heat released or absorbed by the system corresponds directly to the enthalpy change. On the contrary, the change in internal energy (\(\Delta U\)) is challenging to measure directly because it involves measuring changes in both kinetic and potential energy of the system's molecules, which are not accessible by measurements in a typical laboratory. Therefore, measuring the change in enthalpy (\(\Delta H\)) is generally easier than the change in internal energy (\(\Delta U\)).

(b) Determining the endothermic or exothermic nature of a process

A process is endothermic if it absorbs heat (energy) from the surroundings, and it is exothermic if it releases heat (energy) to the surroundings. The change in enthalpy (\(\Delta H\)) can be utilized to determine the nature of the process. If \(\Delta H\) is negative, it indicates that the system releases heat to the surroundings; hence the process is exothermic. Conversely, if \(\Delta H\) is positive, it shows that the system absorbs heat from the surroundings, making the process endothermic. Given that the change in enthalpy at constant pressure, \(\Delta H\), is negative for the given process, the process is exothermic.

Key concepts

Measuring Enthalpy vs Internal Energy

Understanding the difference between measuring enthalpy and internal energy is fundamental in thermodynamics. Enthalpy is defined by the equation \( H = U + PV \), where \( H \) stands for enthalpy, \( U \) for internal energy, and \( PV \) represents the product of pressure and volume. This makes enthalpy a state function that is much easier to measure, especially under constant pressure conditions.

Calorimetry, the measurement of heat transfer, directly determines the change in enthalpy (\( \Delta H \)). Since most chemical reactions and physical changes occur at constant pressure, this is a practical approach for laboratories. In contrast, internal energy includes both kinetic and potential energy at the molecular level, where direct measurement is not straightforward without advanced techniques. This simplicity in measurement is why enthalpy change becomes a more accessible and commonly used parameter than internal energy change for studying thermodynamic processes.

Calorimetry in Chemistry

Calorimetry is a pivotal technique in chemistry used to measure the heat effects of chemical reactions, phase changes, and other physical changes. A device called a calorimeter is employed to ascertain this thermal transfer. In its essence, calorimetry is the process of correlating the amount of heat absorbed or released by a system to the change in enthalpy (\( \Delta H \)).

During a calorimetric experiment, the system will exchange heat with its surroundings, while the change in temperature is closely monitored. For many reactions occurring at constant pressure, this heat exchange measured by the calorimeter corresponds to \( \Delta H \) of the system. Due to its practicality, calorimetry is often used for determining the thermal properties of substances and quantifying the energy changes during various processes, making it a cornerstone of chemical thermodynamics.

Endothermic and Exothermic Processes

Chemical reactions and physical changes can either absorb heat from or release heat to their surrounding environment. When a process absorbs heat, it is labeled as endothermic, characterized by a positive change in enthalpy (\( \Delta H > 0 \)). Such processes may include the melting of ice or the evaporation of water.

Conversely, exothermic processes are those that release heat, indicated by a negative \( \Delta H \) value (\( \Delta H < 0 \)). Common examples are the combustion of fuels or the freezing of water. Understanding whether a process is endothermic or exothermic is crucial because it allows scientists to predict the energy flow and design processes or reactions accordingly for industrial, environmental, and research applications.