Question

When you reverse a thermochemical equation, why must you change the sign of \(\Delta H ?(\text { Chapter } 15)\)

Short answer

When reversing a thermochemical equation, the sign of ΔH must be changed because the roles of reactants and products switch, resulting in the reversal of heat flow direction. This change affects the enthalpy, requiring the sign of ΔH to be reversed to accurately represent the endothermic or exothermic nature of the reversed reaction.

Step-by-step solution

Understanding Thermochemical Equations

A thermochemical equation is a balanced chemical equation that includes the enthalpy change, ΔH, for the reaction. The enthalpy change represents the heat absorbed or released during a chemical reaction at constant pressure.

Defining Enthalpy Change and its Significance

The enthalpy change (ΔH) is the difference in enthalpies between the products and the reactants in a chemical reaction. ΔH can be either positive or negative. - If ΔH is positive, it indicates that the reaction is endothermic, meaning that heat is absorbed from the surroundings. - If ΔH is negative, it indicates that the reaction is exothermic, meaning that heat is released to the surroundings.

The Effect of Reversing the Thermochemical Equation

When a thermochemical equation is reversed, the roles of the reactants and the products are switched, and the direction of heat flow is also reversed. This change affects the enthalpy change (ΔH). If the initial reaction was endothermic (ΔH > 0), the reversed reaction would be exothermic (ΔH < 0) because the heat absorbed by the initial reaction would be released in the reversed reaction. Similarly, if the initial reaction was exothermic (ΔH < 0), the reversed reaction would be endothermic (ΔH > 0) because the heat released by the initial reaction would be absorbed in the reversed reaction.

Why Changing the Sign of ΔH is Necessary

Since reversing the thermochemical equation switches the roles of reactants and products and alters the direction of heat flow, it is necessary to change the sign of ΔH because: 1. The enthalpy change is essentially the difference in enthalpy between the products and reactants. Reversing the thermochemical equation results in switching the roles of reactants and products, hence changing the sign of ΔH. 2. Changing the sign of ΔH ensures that the endothermic or exothermic nature of the reversed reaction is accurately represented. So in conclusion, when a thermochemical equation is reversed, the sign of ΔH must be changed to accurately represent the enthalpy change and the endothermic or exothermic nature of the reversed reaction.

Key concepts

Enthalpy Change

Enthalpy change, symbolized as \( \Delta H \), is a crucial part of understanding chemical reactions. At its core, it represents the difference in energy between the reactants and the products of a reaction, expressed in terms of heat. This energy change occurs at constant pressure. Enthalpy change is not just about numbers; it tells us whether a reaction absorbs or releases energy as heat. It acts as an energy debit or credit for a chemical equation.

Depending on the circumstances, \( \Delta H \) can be either positive or negative. A positive \( \Delta H \) implies that the reaction absorbs heat from its surroundings, while a negative \( \Delta H \) indicates that the reaction releases heat. This understanding helps us categorize reactions and predict energy exchanges in chemical processes.

Endothermic Reactions

An endothermic reaction is one where the system absorbs energy from its surroundings in the form of heat. When the enthalpy change \( \Delta H \) is positive, it highlights an endothermic process.

In these reactions, the energy needed to break the bonds in the reactants is greater than the energy released when new bonds form in the products. This is why they feel cold to the touch, as they absorb heat.

Examples of endothermic reactions include:

• Photosynthesis, where plants absorb sunlight to convert carbon dioxide and water into glucose and oxygen.
• Melting ice, which absorbs heat from the environment to change state from solid to liquid.

These reactions require a continual input of energy to proceed.

Exothermic Reactions

Exothermic reactions are those where energy is released into the surroundings as heat. Here, the enthalpy change \( \Delta H \) is negative. This means that more energy is released when the products form than the energy required to break the bonds of the reactants.

This release of heat often makes the reaction environment feel warmer.

Common examples of exothermic reactions include:

• Combustion, like burning wood or fossil fuels, releasing heat and light.
• Respiration in our bodies, which releases energy from food to power various biological processes.

These reactions are self-sustaining once started, as the energy released can fuel further reactions.