Energy Changes in Chemical Reactions (Grade 11 NSC Matric Physical Sciences): Revision Notes

Energy Changes in Chemical Reactions

Chemical reactions are all around us, from the burning campfire that provides warmth to the food we digest for energy. One of the most important aspects of chemical reactions is that they all involve energy changes. Understanding these energy changes helps us predict how reactions behave and why they occur.

What causes energy changes in chemical reactions

Energy changes in chemical reactions happen because of what occurs to the chemical bonds during the reaction. When a reaction takes place, existing bonds in the reactant molecules must break, and new bonds form in the product molecules.

Let's examine this using a simple example. When hydrogen gas reacts with oxygen gas to form water:

$2\text{H}_2(g) + \text{O}_2(g) → 2\text{H}_2\text{O}(g)$

During this reaction, the bonds between hydrogen atoms in H₂ molecules must break, and the bonds between oxygen atoms in O₂ molecules must also break. At the same time, new bonds form between hydrogen and oxygen atoms to create water molecules.

Bond energy

Understanding bond strength is fundamental to predicting energy changes in reactions. Bond energy is a measure of bond strength in a chemical bond. It tells us exactly how much energy is needed to break the bond between two atoms. Bond energies are measured in kilojoules per mole (kJ·mol⁻¹).

The energy needed to break a bond is also called bond dissociation energy. When we look at the overall energy change in a reaction, we need to consider both the energy required to break existing bonds and the energy released when new bonds form.

Enthalpy

When we study the total energy of a chemical reaction system, we use a concept called enthalpy. Enthalpy is a measure of the total energy of a chemical system at a given pressure, represented by the symbol H.

Enthalpy helps us understand the overall energy changes that occur during reactions. As we learn about different types of reactions, the concept of enthalpy becomes essential for understanding why some reactions release energy while others absorb it.

Exothermic reactions

An exothermic reaction is one that releases energy in the form of heat or light. In these reactions, the energy needed to break the bonds in the reactants is less than the energy released when new bonds form in the products.

We can represent exothermic reactions using the general formula: Reactants → Products + Energy

Another way to understand exothermic reactions is that the energy of the products is less than the energy of the reactants because energy has been released during the reaction. When energy is released, the surroundings (like the air around the reaction) become warmer.

Common examples of exothermic reactions include:

Combustion reactions: When fuels burn, they release large amounts of energy. For example, when petrol burns: $\text{C}_8\text{H}_{18}(l) + 25\text{O}_2(g) → 16\text{CO}_2(g) + 18\text{H}_2\text{O}(g) + \text{heat}$

This is why we use fuels for energy - the combustion reactions release energy that we can use for heating, electricity, and powering vehicles.

Respiration: The chemical reaction in our bodies that produces energy from food: $\text{C}_6\text{H}_{12}\text{O}_6(s) + 6\text{O}_2(g) → 6\text{CO}_2(g) + 6\text{H}_2\text{O}(l) + \text{energy}$

This reaction allows our cells to function by providing the energy they need.

Endothermic reactions

An endothermic reaction is one that absorbs energy in the form of heat or light. In these reactions, the energy needed to break bonds in the reactants is greater than the energy released when new bonds form in the products.

We can represent endothermic reactions using the general formula: Reactants + Energy → Products

In endothermic reactions, the energy of the products is greater than the energy of the reactants because energy has been absorbed during the reaction. When energy is absorbed, the surroundings become cooler.

Common examples of endothermic reactions include:

Photosynthesis: Plants use energy from sunlight to convert carbon dioxide and water into glucose: $6\text{CO}_2(g) + 6\text{H}_2\text{O}(l) + \text{energy} → \text{C}_6\text{H}_{12}\text{O}_6(s) + 6\text{O}_2(g)$

Thermal decomposition of limestone: In industry, limestone is heated to very high temperatures to break it down: $\text{CaCO}_3(s) → \text{CaO}(s) + \text{CO}_2(g)$

This reaction requires continuous heating because it absorbs energy.

Heat of reaction and enthalpy change

The difference in energy between the reactants and products is called the heat of the reaction. This is also referred to as the enthalpy change of the system, represented by the symbol ΔH (delta H).

Worked example 1: investigating an endothermic reaction

Worked example 2: investigating an exothermic reaction

Identifying reaction types through temperature changes

Exam tips for energy changes