What you'll learn
Energy changes occur in all chemical reactions. Understanding whether reactions release or absorb energy is essential for explaining everyday phenomena from hand warmers to photosynthesis. This guide covers the AQA GCSE Chemistry specification on energy changes, including reaction profiles, activation energy, and calculating energy changes using bond energies.
Key terms and definitions
Exothermic reaction — a reaction that transfers energy to the surroundings, usually as heat, shown by a temperature rise
Endothermic reaction — a reaction that takes in energy from the surroundings, usually as heat, shown by a temperature decrease
Activation energy — the minimum amount of energy required for a reaction to occur
Reaction profile — a diagram showing the relative energies of reactants and products, and how energy changes during a reaction
Bond energy — the energy required to break one mole of a particular covalent bond in the gaseous state
Enthalpy change — the overall energy change in a reaction, given the symbol ΔH
Calorimetry — an experimental technique used to measure the energy change in a chemical reaction
Chemical bond — a force of attraction between atoms, requiring energy to break and releasing energy when formed
Core concepts
Exothermic and endothermic reactions
Exothermic reactions
Exothermic reactions release energy to the surroundings. The energy released comes from the formation of new chemical bonds in the products.
Key characteristics:
- Temperature of surroundings increases
- Energy is transferred from the reaction to the surroundings
- ΔH is negative (by convention)
- Products have less energy than reactants
Common examples:
- Combustion reactions (burning fuels)
- Neutralisation reactions (acid + alkali)
- Oxidation reactions (rusting, respiration)
- Self-heating cans and hand warmers
- Displacement reactions (adding zinc to copper sulfate solution)
Endothermic reactions
Endothermic reactions absorb energy from the surroundings. Energy is required to break the bonds in the reactants, and more energy is absorbed than is released when new bonds form.
Key characteristics:
- Temperature of surroundings decreases
- Energy is transferred from the surroundings to the reaction
- ΔH is positive (by convention)
- Products have more energy than reactants
Common examples:
- Thermal decomposition (heating calcium carbonate)
- Photosynthesis
- Dissolving ammonium nitrate in water
- Sports injury cold packs
- Electrolysis
Reaction profiles
Reaction profiles are diagrams that show the energy changes during a chemical reaction. They plot energy on the y-axis against progress of reaction on the x-axis.
Exothermic reaction profile
In an exothermic reaction profile:
- Reactants start at a higher energy level than products
- The curve rises to a peak (activation energy barrier)
- The curve falls to the products level
- The difference between reactants and products is the energy released (ΔH is negative)
- The activation energy is shown from the reactant level to the peak
Endothermic reaction profile
In an endothermic reaction profile:
- Reactants start at a lower energy level than products
- The curve rises to a peak (activation energy barrier)
- The curve remains at the higher products level
- The difference between reactants and products is the energy absorbed (ΔH is positive)
- The activation energy is shown from the reactant level to the peak
The importance of activation energy
All reactions require activation energy to start. This energy is needed to break the initial bonds in the reactants. Even exothermic reactions need activation energy to get started.
Factors affecting activation energy:
- Catalysts lower the activation energy, making reactions faster without being used up
- Higher temperatures don't change the activation energy but give particles more energy to overcome it
- Breaking stronger bonds requires more activation energy
Energy changes in chemical reactions
Bond breaking and bond making
Understanding energy changes requires knowledge of what happens to bonds:
Bond breaking:
- Requires energy (endothermic process)
- Energy must be supplied to overcome the attraction between atoms
- Different bonds require different amounts of energy to break
Bond making:
- Releases energy (exothermic process)
- Energy is released as new attractions form between atoms
- Different bonds release different amounts of energy when formed
Overall energy change
The overall energy change in a reaction depends on the balance between energy required to break bonds and energy released when new bonds form:
If more energy is released making bonds than is required to break bonds:
- The reaction is exothermic overall
- Temperature increases
- ΔH is negative
If more energy is required to break bonds than is released making bonds:
- The reaction is endothermic overall
- Temperature decreases
- ΔH is positive
Calculating energy changes
Using bond energies
The energy change of a reaction can be calculated using bond energies:
Energy change = Energy required to break bonds - Energy released making bonds
Or: ΔH = Σ(bonds broken) - Σ(bonds made)
Method:
- Write out the balanced equation
- Draw out the structural formulae showing all bonds
- Calculate total energy needed to break all bonds in reactants
- Calculate total energy released making all bonds in products
- Subtract energy released from energy required
- State whether the reaction is exothermic (negative ΔH) or endothermic (positive ΔH)
Important points about bond energy calculations
- Bond energies are always positive values (energy required to break bonds)
- Bond energies are average values taken from many compounds
- All species must be in gaseous state (bond energies are defined for gaseous atoms)
- The same type of bond releases the same energy when formed as required to break it
- More accurate calculations would use standard enthalpy changes of formation
Measuring energy changes (Required Practical)
Simple calorimetry experiments
Students must be able to perform simple calorimetry experiments to measure temperature changes.
Typical method for combustion:
- Measure a known volume of water (e.g., 100 cm³) into a copper calorimeter or metal can
- Measure the initial temperature of the water
- Weigh the spirit burner containing the fuel
- Place the burner under the calorimeter and light it
- Stir the water constantly
- After a set time (or temperature rise), extinguish the flame
- Record the highest temperature reached
- Reweigh the spirit burner to find mass of fuel burned
For dissolving or neutralisation reactions:
- Measure volumes of solutions using measuring cylinders
- Measure initial temperature
- Mix solutions in a polystyrene cup (good insulator)
- Stir and record maximum or minimum temperature reached
Calculating energy transferred
Q = m × c × ΔT
Where:
- Q = energy transferred (J)
- m = mass of water or solution (g)
- c = specific heat capacity (4.2 J/g°C for water)
- ΔT = temperature change (°C)
Key assumptions:
- All energy released heats the water
- No energy lost to surroundings
- The solution has the same specific heat capacity as water (4.2 J/g°C)
- The density of the solution is 1 g/cm³ (so 1 cm³ = 1 g)
Sources of error
- Energy lost to surroundings (largest source)
- Energy absorbed by the calorimeter/container
- Incomplete combustion
- Heat lost to the air
- Evaporation of water
- Non-standard conditions
Improvements:
- Use a lid to reduce energy loss
- Use insulation around the calorimeter
- Use a draught shield
- Repeat and calculate a mean
- Use a larger temperature change
Worked examples
Example 1: Calculating energy change using bond energies
Question: Hydrogen reacts with chlorine to form hydrogen chloride: H₂ + Cl₂ → 2HCl
Use the bond energies below to calculate the energy change for this reaction. State whether the reaction is exothermic or endothermic.
Bond energies: H-H = 436 kJ/mol, Cl-Cl = 243 kJ/mol, H-Cl = 432 kJ/mol
[4 marks]
Answer:
Energy required to break bonds:
- 1 × H-H = 436 kJ/mol
- 1 × Cl-Cl = 243 kJ/mol
- Total = 679 kJ/mol [1 mark for correct calculation of bonds broken]
Energy released making bonds:
- 2 × H-Cl = 2 × 432 = 864 kJ/mol [1 mark for correct calculation of bonds made]
Energy change = 679 - 864 = -185 kJ/mol [1 mark for correct calculation]
The reaction is exothermic because energy change is negative / more energy is released than absorbed [1 mark for correct conclusion with explanation]
Example 2: Calorimetry calculation
Question: A student added 50 cm³ of hydrochloric acid to 50 cm³ of sodium hydroxide solution in a polystyrene cup. The temperature rose from 19.5°C to 26.2°C.
Calculate the energy released in this neutralisation reaction. (Specific heat capacity of water = 4.2 J/g°C. Assume the density of the solution is 1 g/cm³)
[3 marks]
Answer:
Temperature change: 26.2 - 19.5 = 6.7°C [1 mark]
Total volume = 50 + 50 = 100 cm³ Mass = 100 g (since density = 1 g/cm³) [1 mark]
Q = m × c × ΔT Q = 100 × 4.2 × 6.7 Q = 2814 J (or 2.814 kJ) [1 mark]
Example 3: Identifying reaction types from temperature changes
Question: A student carried out four reactions and measured the temperature change. Identify which reactions are exothermic and which are endothermic.
| Reaction | Temperature change |
|---|---|
| A | +8.5°C |
| B | -4.2°C |
| C | +15.3°C |
| D | -2.1°C |
[2 marks]
Answer:
Exothermic: A and C (temperature increases) [1 mark]
Endothermic: B and D (temperature decreases) [1 mark]
Common mistakes and how to avoid them
Confusing the sign of energy changes: Remember that exothermic reactions have negative ΔH values, even though they release energy and feel hot. The negative sign indicates energy leaving the system.
Getting bond breaking and bond making backwards: Bond breaking always requires energy (endothermic). Bond making always releases energy (exothermic). Write this clearly in calculations.
Forgetting to multiply by the number of bonds: In the equation 2HCl, there are TWO H-Cl bonds being formed. Count carefully using structural formulae.
Using the wrong units in calorimetry calculations: Ensure mass is in grams, specific heat capacity in J/g°C, and temperature in °C. Convert your final answer to kJ if required (divide by 1000).
Misinterpreting activation energy: Activation energy is the minimum energy needed for a reaction to occur, not the overall energy change. It's shown from reactants to the peak of the reaction profile, not from reactants to products.
Assuming all energy heats the solution: In calorimetry experiments, significant energy is lost to surroundings. Always mention this as a source of error and suggest insulation as an improvement.
Exam technique for Energy Changes
Command word "calculate": Show all working clearly. Write the formula, substitute values, and give the answer with correct units. Marks are often available for method even if the final answer is incorrect.
Drawing reaction profiles: Use a ruler for axes and label them clearly (Energy on y-axis, Progress of reaction on x-axis). Mark reactants, products, and activation energy clearly with labels and arrows. Show whether the reaction is exothermic or endothermic.
Extended response questions: Questions worth 6 marks may ask you to explain energy changes in reactions. Structure your answer to include: bond breaking (endothermic), bond making (exothermic), comparison of energy required vs. released, and conclusion about overall energy change.
Practical questions: Expect questions on the required practical. Be ready to identify control variables, suggest improvements to reduce errors, explain why polystyrene cups are used, and calculate energy changes from results.
Quick revision summary
Energy changes occur in all reactions. Exothermic reactions release energy (ΔH negative), causing temperature increases; endothermic reactions absorb energy (ΔH positive), causing temperature decreases. All reactions need activation energy to start. Energy changes can be calculated using bond energies: energy change equals energy needed to break bonds minus energy released making bonds. Calorimetry experiments measure energy changes using Q = mcΔT, though energy losses to surroundings affect accuracy. Reaction profiles illustrate energy changes graphically.