Understanding energy changes in chemical reactions is a fundamental part of your AQA GCSE Chemistry course. These notes cover everything you need to know about exothermic and endothermic reactions, activation energy, and bond energy calculations.

Exothermic and Endothermic Reactions

In any chemical reaction, energy is conserved. It cannot be created or destroyed, only transferred from one form to another. This transfer is what defines a reaction as either exothermic or endothermic.

What is an Exothermic Reaction? An exothermic reaction is one that transfers energy to the surroundings, usually in the form of heat, causing the temperature of the surroundings to increase. Think of it as an ‘exit’ of heat from the reaction. In these reactions, the product molecules have less energy stored in their chemical bonds than the reactant molecules did.

• Key Examples: Common examples include combustion (burning fuels), neutralisation reactions between acids and alkalis, and many oxidation reactions.

• Everyday Uses: You’ll find exothermic reactions in everyday items like self-heating cans and reusable hand warmers.

What is an Endothermic Reaction? An endothermic reaction is the opposite. It takes in energy from the surroundings, causing the temperature of the surroundings to decrease. Think of it as heat ‘entering’ the reaction. In this case, the product molecules have more energy stored in their chemical bonds than the reactants.

• Key Examples: A key example is thermal decomposition. The reaction between citric acid and sodium hydrogencarbonate is another common one.

• Everyday Uses: Instant cold packs used for sports injuries are a perfect example of an endothermic reaction in action.

Activation Energy and Reaction Profiles

For any reaction to occur, the reacting particles must collide with enough energy. The activation energy is the minimum amount of energy needed to start a reaction by breaking the existing bonds in the reactants.

Reaction profiles are diagrams that show the energy changes during a reaction.

• For an exothermic reaction, the profile shows the products at a lower energy level than the reactants.

• For an endothermic reaction, the profile shows the products at a higher energy level than the reactants. A catalystworks by providing an alternative reaction pathway with a lower activation energy, which speeds up the rate of the reaction without being used up.

Bond Energy Calculations (Higher Tier Only)

Chemical reactions involve two key stages: breaking bonds and making bonds.

• Bond Breaking: Energy is required to be put in to break the chemical bonds in the reactants. This is an endothermic process.

• Bond Making: Energy is released when new chemical bonds are formed in the products. This is an exothermic process.

The overall energy change of a reaction can be calculated by subtracting the energy released from the energy supplied:Overall Energy Change = Energy required to break bonds – Energy released by forming bonds

• If the result is negative, more energy was released than supplied, and the reaction is exothermic.

• If the result is positive, more energy was supplied than released, and the reaction is endothermic.

Required Practical 4: Investigating Temperature Changes

In this practical, you investigate the variables affecting temperature changes in reacting solutions, such as neutralisation. When an acid and alkali react, the temperature rises because neutralisation is exothermic. By adding the acid to the alkali bit by bit and measuring the temperature, you can find the exact point of neutralisation, which is the point where the maximum temperature is reached. A major source of error is heat loss to the surroundings, which can be minimised by using a polystyrene cup and a lid.

Exam tip – When asked to draw a reaction profile, make sure you label both axes correctly (‘Energy’ on the y-axis and ‘Progress of reaction’ on the x-axis). Clearly label the reactants and products, the activation energy (Ea​), and the overall energy change ($\Delta H$). For exothermic reactions, the products must be shown at a lower energy level than the reactants, and for endothermic, they must be higher.

Reaction profiles, also known as energy level diagrams, are essential for your AQA GCSE Chemistry exam. They are graphs that show the energy changes happening during a chemical reaction, helping you visualise the difference between exothermic and endothermic processes.

What are the Key Parts of a Reaction Profile?

Every reaction profile plots Energy on the y-axis against the Progress of the reaction on the x-axis. You must be able to label three key features:

• Reactants and Products: The energy levels of the substances you start with (reactants) and end with (products).

• Activation Energy (Ea​): This is the peak of the curve on the diagram. It represents the minimum amount of energy required to start a reaction by breaking the bonds in the reactants. According to collision theory, particles must collide with at least this much energy for a reaction to occur.

• Overall Energy Change ($\Delta H$): This is the difference in energy between the reactants and the products. It shows how much energy is taken in or given out overall.

Reaction Profile for an Exothermic Reaction

In an exothermic reaction, energy is released to the surroundings. This is shown on a reaction profile in a specific way:

• The products are at a lower energy level than the reactants.

• This means the overall energy change ($\Delta H$) is negative, as energy has been lost from the reaction to the surroundings.

• The activation energy is the ‘hump’ that must be overcome for the reaction to start.

Reaction Profile for an Endothermic Reaction

In an endothermic reaction, energy is taken in from the surroundings. The reaction profile reflects this:

• The products are at a higher energy level than the reactants.

• This means the overall energy change ($\Delta H$) is positive, as energy has been absorbed by the reaction from the surroundings.

• Again, the activation energy is the initial energy ‘hump’ required to get the reaction going.

How do Catalysts Affect Reaction Profiles?

A catalyst speeds up a chemical reaction without being used up. It does this by providing an alternative reaction pathway with a lower activation energy.

• On a reaction profile, a catalysed reaction will have a lower peak.

• Crucially, a catalyst does not change the energy of the reactants or products. Therefore, the overall energy change ($\Delta H$) remains exactly the same.

Exam tip – When drawing a reaction profile in your exam, always use a ruler for the axes and label them clearly: “Energy” and “Progress of reaction”. Use a smooth, single curve to show the reaction pathway. Make sure your arrows for activation energy (Ea​) and overall energy change ($\Delta H$) are distinct and start and end at the correct energy levels to avoid losing easy marks.

This topic is for Higher Tier students only and covers how to calculate the energy change in a reaction using bond energies, a key part of Topic 5, Energy Changes.

Bond Breaking and Bond Making

Every chemical reaction involves two main processes: breaking existing bonds in the reactants and forming new bonds in the products.

• Bond Breaking: Energy must be supplied to break chemical bonds. This is an endothermic process. The initial energy required to start breaking these bonds is the activation energy.

• Bond Making: Energy is released when new chemical bonds are formed. This is an exothermic process.

How to Calculate the Overall Energy Change

The overall energy change of a reaction is the difference between the energy taken in to break bonds and the energy given out when making bonds. For your exam, you will be given the specific bond energies required for the calculation.

The formula you need to use is: Overall Energy Change = Sum of energy needed to break bonds − Sum of energy released from making bonds

Interpreting the Result

The sign of your final answer tells you whether the reaction is exothermic or endothermic.

• A negative energy change means the reaction is exothermic. More energy is released when forming new bonds than was needed to break the original bonds.

• A positive energy change means the reaction is endothermic. More energy was needed to break the bonds in the reactants than was released when forming the new bonds in the products.

Worked Example: Hydrogen + Oxygen Reaction

Let’s calculate the energy change for the reaction: 2H₂ + O₂ ⟶ 2H₂O

You would be given these bond energies:

• H−H = 436 kJ/mol

• O=O = 498 kJ/mol

• O−H = 464 kJ/mol

Step 1: Calculate the energy needed to break bonds (Reactants)

• You are breaking two H−H bonds: 2 × 436 = 872 kJ

• You are breaking one O=O bond: 1 × 498 = 498 kJ

• Total energy in = 872 + 498 = 1370 kJ

Step 2: Calculate the energy released from making bonds (Products)

• You are making two molecules of H₂O (water).

• Each water molecule has two O−H bonds, so you are forming a total of four O−H bonds.

• Total energy out = 4 × 464 = 1856 kJ

Step 3: Calculate the overall energy change

• Energy change = Energy in − Energy out

• Energy change = 1370 − 1856 = −486 kJ/mol

The final answer is negative, so the reaction is exothermic.

Exam tip – A very common mistake in bond energy calculations is miscounting the number of bonds, especially in the products. For a reaction making 2H₂O, remember that’s two separate water molecules, and each one has two O-H bonds, giving a total of four O-H bonds. Always draw out the molecules if you’re unsure, so you can count the bonds accurately before you start calculating.

This revision guide is for Chemistry Only (Triple Science) students, covering Topic 4.5.2, Chemical Cells and Fuel Cells, from the AQA GCSE specification.

Chemical Cells and Batteries

Chemical cells are devices that convert chemical energy into electrical energy through chemical reactions.

How do Simple Chemical Cells Work?

A simple cell can be made by connecting two different metals as electrodes and placing them in an electrolyte. An electrolyte is a solution that can conduct electricity.

• The reaction between the chemicals in the cell produces electricity and generates a voltage, or potential difference.

• The size of the voltage produced depends on factors like the type of metals used for the electrodes and the choice of electrolyte.

What is a Battery?

A battery is simply two or more cells connected in series. This is done to produce a greater overall voltage than a single cell can provide.

Rechargeable vs Non-Rechargeable Cells

• Non-Rechargeable Cells: In devices like alkaline batteries, the chemical reactions are irreversible. They stop producing electricity when one of the reactants is completely used up and cannot be recharged.

• Rechargeable Cells: In these batteries, the chemical reactions can be reversed by applying an external electrical current. This process recharges the cell, allowing it to be used again.

Fuel Cells

Fuel cells are an alternative to conventional cells and batteries, also producing electricity from a chemical reaction.

How do Fuel Cells Work?

Unlike a normal cell, a fuel cell is continuously supplied with a fuel (like hydrogen) and oxygen (or air) from an external source. Inside the cell, the fuel is oxidised electrochemically, which means it reacts without burning to produce a voltage and generate electrical energy.

The Hydrogen Fuel Cell

The most common type is the hydrogen fuel cell.

• Overall Reaction: Hydrogen is oxidised by oxygen to produce just one product: water.

  • Hydrogen + Oxygen ⟶ Water

  • 2H₂ + O₂ ⟶ 2H₂O

• Advantages: Compared to rechargeable batteries, hydrogen fuel cells are often more efficient, produce less pollution (only water is made), and can be more compact and lightweight.

Half Equations (Higher Tier Only)

For the hydrogen fuel cell, you need to know the half-equations that occur at each electrode.

• At the negative electrode: 2H₂ ⟶ 4H⁺ + 4e⁻

• At the positive electrode: O₂ + 4H⁺ + 4e⁻ ⟶ 2H₂O

Exam tip – When comparing hydrogen fuel cells and rechargeable batteries, think about the practical applications. For example, a key advantage of fuel cells for cars is that they can be refuelled quickly, much like a petrol car, whereas rechargeable batteries can take several hours to charge. However, a disadvantage is the difficulty in safely storing hydrogen fuel under high pressure.

A key topic for Triple Science (Chemistry Only) students studying AQA GCSE Chemistry (Topic 4.5.2).

How Do Fuel Cells Work?

Fuel cells are electrochemical devices that convert chemical energy directly into electrical energy. Unlike conventional batteries, they are supplied with a continuous external source of fuel, such as hydrogen, and oxygen from the air.

Inside the cell, the fuel is oxidised electrochemically, which means it reacts to lose electrons without burning. This process creates a potential difference (voltage), generating an electric current that can power a device, like a car.

The Hydrogen Fuel Cell

The most common example is the hydrogen fuel cell.

• Overall Reaction: Hydrogen and oxygen react to produce only one product: water. Hydrogen + Oxygen ⟶ Water

• Energy Transfer: This reaction releases electrical energy to power the vehicle.

• Half Equations (Higher Tier Only): Higher Tier students must know the half-equations for the reactions at the electrodes.

Evaluating Fuel Cells vs. Rechargeable Batteries

You need to be able to compare the advantages and disadvantages of hydrogen fuel cells against rechargeable batteries.

Advantages of Hydrogen Fuel Cells:

• High Efficiency: They are very efficient at converting chemical energy to electrical energy.

• Low Pollution: The only product is water, so there are no polluting emissions at the point of use.

• Lightweight and Compact: They can be smaller and lighter than batteries that produce the same amount of power.

• Reliable: Having no moving parts makes them very reliable.

Infrastructure Challenges: A major drawback is the lack of infrastructure. Cars powered by fuel cells require a network of hydrogen filling stations, which are currently very rare compared to petrol stations. This is partly due to low demand and the significant time and investment required to transition from a petrol-based system.

Exam tip – For evaluation questions, always provide a balanced argument. While hydrogen fuel cells are clean at the point of use, you could mention that producing hydrogen often requires energy from burning fossil fuels, which creates emissions. Also, contrast the quick refuelling time of a fuel cell vehicle with the often lengthy recharging time required for an electric car with a rechargeable battery.