How to Study Chemical Energetics for IGCSE Chemistry

Chemical energetics is a short topic in the 0620 syllabus, but it carries reliable marks and connects to several other topics. The concepts are straightforward once you understand the core idea: chemical reactions involve energy changes because breaking bonds requires energy and forming bonds releases energy. Everything else — energy level diagrams, bond energy calculations, activation energy — follows from this single principle.

The core concepts

Exothermic reactions

Energy is released to the surroundings. The temperature of the surroundings increases. The products have less energy than the reactants.

Examples: combustion of fuels, neutralisation of acids, respiration, oxidation of metals.

Endothermic reactions

Energy is absorbed from the surroundings. The temperature of the surroundings decreases. The products have more energy than the reactants.

Examples: thermal decomposition (e.g., CaCO3 → CaO + CO2), photosynthesis, dissolving ammonium nitrate in water.

Study tip: Students confuse “temperature decreases” with “the reaction is giving off cold.” No reaction produces cold. An endothermic reaction absorbs heat from its surroundings, which is why a beaker feels cold to the touch. The reaction is taking energy from your hand.

Energy level diagrams

These diagrams are worth consistent marks and are easy to get right once the format is clear.

Exothermic diagram:

• Reactants are drawn at a higher energy level
• Products are drawn at a lower energy level
• The difference is the energy released (drawn as a downward arrow labelled with the energy change)
• A hump between reactants and products represents the activation energy

Endothermic diagram:

• Reactants are drawn at a lower energy level
• Products are drawn at a higher energy level
• The difference is the energy absorbed (drawn as an upward arrow)
• The activation energy hump still appears

What the examiner looks for:

1. Correct relative positions of reactants and products
2. Labels for “reactants” and “products”
3. An arrow showing the overall energy change, correctly labelled
4. The activation energy hump (if asked for)
5. Axes: y-axis is “energy,” x-axis is “progress of reaction” or “reaction pathway”

Practise drawing both types from memory. Compare your diagram against the textbook or drawing guide. Common errors: forgetting to label the axes, drawing the activation energy arrow from the reactant level to the top of the hump rather than from the reactant level, and putting the energy change arrow in the wrong direction.

Activation energy and catalysts

Activation energy is the minimum energy that colliding particles must have for a reaction to occur. All reactions, including exothermic ones, need activation energy to start. A match provides the activation energy for the combustion of methane; once started, the reaction sustains itself because it is exothermic.

A catalyst lowers the activation energy. On the energy level diagram, the hump is lower, but the overall energy change (the gap between reactants and products) stays the same. This is a frequent exam question: “How does a catalyst increase the rate of reaction?” Answer: it provides an alternative pathway with lower activation energy, so a greater proportion of particles have enough energy to react on collision.

Bond energy calculations (extended)

The principle: energy is needed to break bonds (endothermic process) and energy is released when bonds form (exothermic process).

Method:

1. Draw out the structural formulae of all reactants and products so you can see every bond.
2. List all bonds broken in the reactants. Look up their bond energies from the data table (provided in the exam).
3. Total the energy for bond breaking.
4. List all bonds formed in the products. Look up their bond energies.
5. Total the energy for bond formation.
6. Energy change = energy to break bonds - energy to form bonds.
7. If the answer is positive, the reaction is endothermic. If negative, it is exothermic.

Worked example: combustion of methane CH4 + 2O2 → CO2 + 2H2O

Bonds broken:

• 4 x C-H = 4 x 413 = 1652 kJ
• 2 x O=O = 2 x 498 = 996 kJ
• Total = 2648 kJ

Bonds formed:

• 2 x C=O = 2 x 805 = 1610 kJ
• 4 x O-H = 4 x 464 = 1856 kJ
• Total = 3466 kJ

Energy change = 2648 - 3466 = -818 kJ (negative, so exothermic — correct, combustion releases energy).

Common errors:

• Forgetting to count bonds correctly: two O2 molecules means two O=O bonds, and two H2O molecules means four O-H bonds.
• Using the wrong bond energy value from the table (check you are reading C=O in CO2, not C-O in a single bond context).
• Getting the subtraction the wrong way round. Always: bonds broken minus bonds formed.

Linking energetics to other topics

Chemical energetics connects widely:

• Chemical reactions: activation energy explains why temperature increases the rate (more particles exceed the activation energy). A catalyst lowers the activation energy.
• Acids and bases: neutralisation is exothermic. This is testable with a thermometer in a calorimetry experiment.
• Metals: the thermite reaction (aluminium + iron oxide) is highly exothermic and can be explained using bond energies.
• Organic chemistry: combustion of hydrocarbons is exothermic. Comparison of energy released by different fuels appears on Paper 4.

Study plan

Days 1-2: Definitions and examples. Flashcard exothermic and endothermic definitions with examples. Practise classifying reactions.

Days 3-4: Energy level diagrams. Draw exothermic and endothermic diagrams from memory. Add activation energy and the effect of a catalyst.

Days 5-7: Bond energy calculations. Work through five to six examples covering combustion, hydrogenation, and other reactions. Use the five-step method and check each answer.

Week 2 onward: Past paper questions. Energetics questions are predictable: draw the diagram, calculate the energy change, explain whether the reaction is exothermic or endothermic. See the energetics exam guide for mark scheme patterns.

The topic is small enough to master in two weeks of focused study. Do not leave it late — the bond energy calculation is easy marks once the method is automatic, and easy marks lost if you have never practised it.