Exothermic and Endothermic Reactions for the ESAT

The Nature of Enthalpy Change

In chemistry, every substance has a certain amount of chemical energy stored within its bonds and structure. This total heat content of a system at constant pressure is called enthalpy, represented by the symbol $H$. While the absolute enthalpy of a substance cannot be measured directly, chemists can measure the energy exchanged during a reaction. This is known as the enthalpy change, denoted as $\Delta H$. The enthalpy change of a reaction is defined as the difference between the enthalpy of the products and the enthalpy of the reactants: $\Delta H = H_{products} - H_{reactants}$.

Exothermic Reactions

An exothermic reaction is one in which energy is transferred from the chemical system to the surroundings. Because the system loses energy, the products have less enthalpy than the reactants. Consequently, the value of $\Delta H$ is always negative ($\Delta H < 0$). In the laboratory, the most common indicator of an exothermic reaction is an increase in the temperature of the reaction mixture or the surrounding environment. Common examples of exothermic processes include the combustion of fuels, such as burning methane in oxygen, and the neutralisation of an acid with an alkali. When bonds are formed in the products, more energy is released than was required to break the bonds in the reactants.

Endothermic Reactions

An endothermic reaction is one in which energy is absorbed by the chemical system from its surroundings. In this case, the products have more enthalpy than the reactants, meaning the system has gained energy. Therefore, the value of $\Delta H$ is always positive ($\Delta H > 0$). These reactions typically result in a decrease in the temperature of the surroundings as thermal energy is converted into chemical potential energy. Examples of endothermic reactions include thermal decomposition, such as heating calcium carbonate to form calcium oxide and carbon dioxide, and the process of photosynthesis. These reactions cannot proceed without a constant input of energy.

Reaction Profiles and Activation Energy

Reaction profiles are used to visualise the energy changes in a reaction. On a graph where the $y$ axis is enthalpy and the $x$ axis is the progress of reaction, an exothermic profile shows the product level below the reactant level. Conversely, an endothermic profile shows the product level above the reactant level. All reactions, whether exothermic or endothermic, require an initial input of energy to begin the process of breaking bonds. This minimum energy required is called the activation energy ($E_a$). On a reaction profile, this is represented by the energy peak or hump between the reactants and products.

Bond Energies

The enthalpy change of a reaction is fundamentally determined by the energies required to break bonds and the energy released when new bonds form. Breaking a chemical bond is an endothermic process: energy must be supplied to overcome the attractive forces between atoms. Making a chemical bond is an exothermic process: energy is released as atoms move to a more stable, lower energy state. By using average bond enthalpies, we can calculate $\Delta H$ using the following method:

1. Calculate the total energy required to break all bonds in the reactants.

2. Calculate the total energy released when all new bonds in the products are formed.

3. Use the formula: $\Delta H = \text{Total energy of bonds broken} - \text{Total energy of bonds made}$.

Worked Example

Calculating Enthalpy Change. Consider the reaction between hydrogen and chlorine to form hydrogen chloride: $H_2 + Cl_2 \rightarrow 2HCl$. Given bond enthalpies: $H-H = 436 \text{ kJ mol}^{-1}$, $Cl-Cl = 242 \text{ kJ mol}^{-1}$, and $H-Cl = 431 \text{ kJ mol}^{-1}$. Step 1: Energy to break bonds = $436 + 242 = 678 \text{ kJ mol}^{-1}$. Step 2: Energy released making bonds = $2 \times 431 = 862 \text{ kJ mol}^{-1}$. Step 3: $\Delta H = 678 - 862 = -184 \text{ kJ mol}^{-1}$. Since $\Delta H$ is negative, this reaction is exothermic.