Revision notes for OCR GCSE Chemistry Energetics. Open the guide for explanations and worked examples. Written against the OCR GCSE Chemistry (J248) specification, so the content matches what's examinable rather than general Chemistry background.
Energetics
What you'll learn
- How to tell whether a reaction is exothermic or endothermic from temperature changes.
- How to draw and label reaction profile diagrams.
- What activation energy means and why reactions may need a spark or heating to start.
- How to calculate reaction energy changes from bond energies — this bit is Higher Tier only.
The big idea: reactions transfer energy
Chemical reactions involve energy being transferred between the reacting chemicals and their surroundings.
The surroundings means everything outside the reacting chemicals: the test tube, solution, thermometer, air, your hand, and so on. In school experiments, we usually judge the energy transfer by measuring the temperature change of the reaction mixture or solution.
Key Idea
Energy is transferred, not used up
In energetics, avoid saying energy is “lost” or “used up”. Energy is transferred to the surroundings or transferred from the surroundings.
Exothermic and endothermic reactions
Definition
Exothermic and endothermic reactions
- An exothermic reaction transfers energy to the surroundings, so the temperature of the surroundings increases.
- An endothermic reaction transfers energy from the surroundings, so the temperature of the surroundings decreases.
Exothermic reactions
Common exothermic reactions include:
- combustion, such as burning methane;
- neutralisation between an acid and an alkali;
- many oxidation reactions.
For example:
CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g)
Combustion is exothermic because it transfers energy to the surroundings as heat and often light.
Endothermic reactions
Common endothermic processes include:
- some thermal decomposition reactions;
- reactions such as citric acid reacting with sodium hydrogencarbonate;
- photosynthesis overall, because it requires energy from light.
For example, thermal decomposition of calcium carbonate requires heating:
CaCO₃(s) → CaO(s) + CO₂(g)
Using temperature change
In a practical, you may measure the starting temperature, mix the reactants, then record the highest or lowest temperature reached. The temperature change is:
ΔT=Tfinal−Tinitial\Delta T = T_{\text{final}} - T_{\text{initial}}ΔT=Tfinal−TinitialIf ΔT\Delta TΔT is positive, the surroundings got warmer, so the reaction is exothermic. If ΔT\Delta TΔT is negative, the surroundings got cooler, so the reaction is endothermic.
Example
Classifying a reaction from temperature data
A student mixes two solutions. The starting temperature is 21.0 °C. The lowest temperature reached is 16.5 °C. Decide whether the reaction is exothermic or endothermic.
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Calculate the temperature change using ΔT=Tfinal−Tinitial\Delta T = T_{\text{final}} - T_{\text{initial}}ΔT=Tfinal−Tinitial:
ΔT=16.5−21.0=−4.5 ∘C\Delta T = 16.5 - 21.0 = -4.5\ ^\circ\text{C}ΔT=16.5−21.0=−4.5 ∘C -
Interpret the sign. A negative temperature change means the surroundings became cooler.
-
Therefore, energy was transferred from the surroundings into the reacting chemicals, so the reaction is endothermic.
Common Mistake
Judging by whether you heat it first
A candle flame is exothermic, even though you need a match to start it. The match supplies activation energy; once burning, the reaction transfers energy to the surroundings.
Reaction profile diagrams
A reaction profile is a graph showing how the energy changes as a reaction goes from reactants to products.
The vertical axis is energy. The horizontal axis is progress of reaction, which just means how far through the reaction you are. It is not time.
Exothermic reaction profiles
In an exothermic reaction:
- reactants are at a higher energy level than products;
- products are more stable because they have lower energy;
- the energy change arrow points downwards;
- the overall energy change is negative.
Endothermic reaction profiles
In an endothermic reaction:
- reactants are at a lower energy level than products;
- products have more energy than reactants;
- the energy change arrow points upwards;
- the overall energy change is positive.
Definition
Energy change
The energy change is the difference between the energy of the products and the energy of the reactants. On a reaction profile, it is shown by a vertical arrow between the reactant and product energy levels.
Example
Interpreting a reaction profile
A reaction profile shows reactants at 200 kJ/mol and products at 80 kJ/mol. The peak of the curve is at 320 kJ/mol. Identify the reaction type and calculate the activation energy.
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Compare products with reactants. Products are lower energy than reactants, so energy has been transferred to the surroundings and the reaction is exothermic.
-
Calculate the activation energy from the reactants to the peak:
Ea=320−200=120 kJ/molE_{\text{a}} = 320 - 200 = 120\ \text{kJ/mol}Ea=320−200=120 kJ/mol -
Calculate the overall energy change:
ΔE=80−200=−120 kJ/mol\Delta E = 80 - 200 = -120\ \text{kJ/mol}ΔE=80−200=−120 kJ/molThe negative sign matches an exothermic reaction.
Activation energy
Definition
Activation energy
Activation energy, often written as EaE_{\text{a}}Ea, is the minimum energy needed for a reaction to start.
Reacting particles must collide with enough energy to break some existing bonds and begin forming new ones. If the particles do not have enough energy, they may collide but not react.
On a reaction profile, activation energy is the energy gap from the reactants up to the top of the curve. The top of the curve represents the highest-energy stage during the reaction.
Tip
Drawing reaction profiles
Always include four labels if asked for a reaction profile: reactants, products, activation energy, and energy change. Then check whether products are lower for exothermic or higher for endothermic.
Why bonds matter
Chemical reactions involve breaking bonds in the reactants and making new bonds in the products.
Key Idea
Bond breaking and bond making
Breaking bonds requires energy. Making bonds releases energy.
This is one of the most important ideas in this topic. It is also a common source of mistakes: energy is not released when bonds break. Energy is needed to pull bonded atoms apart.
Linking bonds to exothermic and endothermic reactions
A reaction is exothermic if more energy is released when new bonds form than is absorbed when old bonds break.
A reaction is endothermic if more energy is absorbed breaking old bonds than is released making new bonds.
Common Mistake
Getting bond energy backwards
Do not write “breaking bonds releases energy”. At GCSE, use this rule every time: break = take in energy, make = give out energy.
Calculating energy changes from bond energies
This part is Higher Tier only.
A bond energy is the energy needed to break one mole of a particular type of bond. Bond energies are measured in kilojoules per mole, kJ/mol. In GCSE questions, the bond energy values are normally given to you.
The calculation is:
ΔE=∑Ebonds broken−∑Ebonds made\Delta E = \sum E_{\text{bonds broken}} - \sum E_{\text{bonds made}}ΔE=∑Ebonds broken−∑Ebonds madeSo:
- breaking bonds adds energy to the calculation;
- making bonds subtracts energy because energy is released;
- a negative answer means exothermic;
- a positive answer means endothermic.
Example
Calculating energy change using bond energies
Calculate the energy change for the combustion of methane:
CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g)
Use these bond energies: C–H = 413 kJ/mol, O=O = 498 kJ/mol, C=O = 805 kJ/mol, O–H = 464 kJ/mol.
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Count the bonds broken in the reactants. In CH₄ there are 4 C–H bonds. In 2O₂ there are 2 O=O bonds.
Ebroken=(4×413)+(2×498)E_{\text{broken}} = \left(4 \times 413\right) + \left(2 \times 498\right)Ebroken=(4×413)+(2×498) Ebroken=1652+996=2648 kJ/molE_{\text{broken}} = 1652 + 996 = 2648\ \text{kJ/mol}Ebroken=1652+996=2648 kJ/mol -
Count the bonds made in the products. In CO₂ there are 2 C=O bonds. In 2H₂O there are 4 O–H bonds.
Emade=(2×805)+(4×464)E_{\text{made}} = \left(2 \times 805\right) + \left(4 \times 464\right)Emade=(2×805)+(4×464) Emade=1610+1856=3466 kJ/molE_{\text{made}} = 1610 + 1856 = 3466\ \text{kJ/mol}Emade=1610+1856=3466 kJ/mol -
Substitute into the bond energy equation.
ΔE=2648−3466=−818 kJ/mol\Delta E = 2648 - 3466 = -818\ \text{kJ/mol}ΔE=2648−3466=−818 kJ/mol -
Interpret the result. The answer is negative, so the reaction is exothermic.
Common Mistake
Average bond energies are approximate
Bond-energy calculations give an estimate because the values are average bond energies. For GCSE, use the numbers provided in the question and do not worry about where the averages come from.
Practical note: measuring temperature changes
In a school experiment, you might measure the temperature change when two solutions react. A simple method is:
- Put a measured volume of one solution into an insulated cup.
- Measure its initial temperature.
- Add the second reactant, stir, and record the highest or lowest temperature reached.
- Calculate the temperature change.
The insulation reduces unwanted energy transfer to or from the room, making the result more reliable.
Tip
Temperature-change sanity check
If the temperature rises, the reaction is exothermic. If the temperature falls, the reaction is endothermic. Always describe the temperature change of the surroundings, not just “the reaction”.
Exam technique
In the exam
- For temperature data, calculate ΔT\Delta TΔT first, then link the sign to energy transfer: warmer surroundings means exothermic, cooler surroundings means endothermic.
- For reaction profiles, check whether products are above or below reactants before labelling the reaction type.
- For Higher Tier bond calculations, remember: total energy for bonds broken minus total energy for bonds made.
Self review
Check yourself
- Why is burning a candle exothermic even though a flame is needed to start it?
- On a reaction profile, where is the activation energy shown?
- In a bond energy calculation, why are the energies of bonds made subtracted?
Recap questions
Test yourself with 5 quick questions on this guide. Answer them all correctly to complete it.
Chemical reactions
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