Introduction to Thermodynamics: Energy, Heat, and Enthalpy

Energy and Heat

In thermodynamics, energy is the capacity to do work or produce heat. One of the most important distinctions to make is between internal energy and heat.

Heat ($q$) is the transfer of energy between a system and its surroundings due to a temperature difference. It is important to remember that a system does not "contain" heat; rather, heat is the energy in transit. Once the energy is inside the system, it is referred to as internal energy.

When a system absorbs heat, we say the process is endothermic. When a system releases heat to its surroundings, the process is exothermic.

Understanding Enthalpy ($H$)

One of the most practical concepts in thermochemistry is enthalpy ($H$). Enthalpy is a measurement of the total heat content of a system. While measuring total enthalpy is difficult, scientists focus on the change in enthalpy ($\Delta H$) during a process, which occurs at constant pressure.

The change in enthalpy is calculated using the following relationship: $\Delta H = H_{\text{products}} - H_{\text{reactants}}$

Exothermic vs. Endothermic Enthalpy Changes

The sign of $\Delta H$ tells us everything we need to know about the direction of heat flow:

1. Exothermic ($\Delta H < 0$): The system loses heat to the surroundings. The products have less enthalpy than the reactants.
   • Example: Combustion (burning wood).
2. Endothermic ($\Delta H > 0$): The system absorbs heat from the surroundings. The products have more enthalpy than the reactants.
   • Example: Melting ice.

Worked Example: Calculating Enthalpy Change

Let's apply these concepts to a simple scenario. Suppose you are studying a chemical reaction where the reactants have an enthalpy of $150\text{ kJ/mol}$ and the products have an enthalpy of $50\text{ kJ/mol}$.

Step 1: Identify the values. $H_{\text{reactants}} = 150\text{ kJ/mol}$ $H_{\text{products}} = 50\text{ kJ/mol}$

Step 2: Apply the enthalpy change formula. $\Delta H = H_{\text{products}} - H_{\text{reactants}}$ $\Delta H = 50\text{ kJ/mol} - 150\text{ kJ/mol}$ $\Delta H = -100\text{ kJ/mol}$

Step 3: Interpret the result. Because $\Delta H$ is negative ($-100\text{ kJ/mol}$), the reaction is exothermic. This means $100\text{ kJ}$ of energy was released into the surroundings for every mole of reactant consumed.

Summary Checklist:

• Is the system open, closed, or isolated?
• Is heat entering (endothermic) or leaving (exothermic)?
• Is $\Delta H$ positive or negative?