What is Chemical Energetics in the context of Cambridge IGCSE Chemistry?

Chemical Energetics, in the context of Cambridge IGCSE Chemistry, refers to the study of energy changes that occur during chemical reactions. It involves understanding how energy is absorbed or released, the concept of enthalpy, and the role of exothermic and endothermic reactions in chemical processes.

How do exothermic and endothermic reactions differ in Chemical Energetics?

In Chemical Energetics, exothermic reactions are those that release energy to the surroundings, usually in the form of heat, resulting in an increase in temperature. Conversely, endothermic reactions absorb energy from the surroundings, leading to a decrease in temperature. These concepts are crucial for understanding energy transfer in chemical reactions at the Cambridge IGCSE level.

Why is the concept of enthalpy important in Chemical Energetics?

Enthalpy is a key concept in Chemical Energetics because it represents the total heat content of a system. It helps in quantifying the energy changes during chemical reactions. Understanding enthalpy changes allows students to predict whether a reaction is exothermic or endothermic, which is essential for Cambridge IGCSE Chemistry students.

What role do bond energies play in Chemical Energetics?

Bond energies are crucial in Chemical Energetics as they provide insight into the energy required to break chemical bonds and the energy released when new bonds form. By comparing the bond energies of reactants and products, students can determine the overall energy change in a reaction, which is a fundamental aspect of the Cambridge IGCSE Chemistry curriculum.

How is the concept of activation energy relevant to Chemical Energetics?

Activation energy is the minimum energy required for a chemical reaction to occur. In Chemical Energetics, understanding activation energy helps explain why certain reactions occur spontaneously while others require external energy input. This concept is important for Cambridge IGCSE Chemistry students to comprehend how energy barriers affect reaction rates.

Why is "Writing $\Delta H = +$ for an exothermic reaction" a common mistake in Chemical Energetics?

Why it happens: Forgetting the convention. How to avoid it: Exothermic releases energy → products are LOWER → H < 0 (NEGATIVE). Source: 0620/42 — recurring

Why is "Using 'bonds made − bonds broken'" a common mistake in Chemical Energetics?

Why it happens: Inverting the formula. How to avoid it: H = broken − made. (Energy needed to break minus energy released on making.)

Why is "Counting only one bond per molecule" a common mistake in Chemical Energetics?

Why it happens: Missing equation coefficients. How to avoid it: Multiply each bond energy by the number of those bonds in the balanced equation.

Why is "Saying temperature rise = endothermic" a common mistake in Chemical Energetics?

Why it happens: Confusing direction. How to avoid it: Surroundings WARM UP → energy released by reaction → EXOTHERMIC.

Chemical EnergeticsCambridge IGCSE Chemistry (0620)

Chemical Energetics

Q: Why is "Using 'bonds made − bonds broken'" a common mistake in Chemical Energetics?

Why it happens: Inverting the formula. How to avoid it: H = broken − made. (Energy needed to break minus energy released on making.)

Q: Why is "Counting only one bond per molecule" a common mistake in Chemical Energetics?

Why it happens: Missing equation coefficients. How to avoid it: Multiply each bond energy by the number of those bonds in the balanced equation.

Q: Why is "Saying temperature rise = endothermic" a common mistake in Chemical Energetics?

Why it happens: Confusing direction. How to avoid it: Surroundings WARM UP → energy released by reaction → EXOTHERMIC.

Work through the notes, try the practice questions, then take the quiz. The report tells you exactly what to revise next.

PreviousNext

Learn this Topic

Start with the slides for the quick version, then go deeper with the full study notes.

Short Study Notes in the form of Slides

Read the notes first. If the method in a worked example clicks, you're ready for the questions.

Page 1 / 0

Detailed Notes

Full prose, callouts and a recap — built for A* mastery, not just a quick scan.

Take these study notes with you

Download a branded PDF — full prose, callouts, recap and memorise list for Chemical Energetics, ready to print or save offline.

Chemical Energetics — Cambridge IGCSE 0620 Chemistry Extended (2026)

Exothermic vs endothermic reactions, energy-level diagrams, and (Extended) calculating $\Delta H$ from bond energies. The energy bookkeeping of chemistry.

What you’ll learn

Mapped to the Cambridge IGCSE 0620 syllabus (2026-2028).

6.1 — Distinguish between exothermic and endothermic reactions.
6.1 — Draw energy level diagrams labelling $\Delta H$ and $E_a$ .
6.2 — Calculate $\Delta H$ from bond energies (Extended).

Exothermic and endothermic reactions

Exo: heat OUT, $\Delta H$ negative. Endo: heat IN, $\Delta H$ positive.

Exothermic. Releases energy (usually as heat) to the surroundings. The reaction mixture gets HOTTER.

$\Delta H$ is NEGATIVE (energy LEFT the system).
Common: combustion, neutralisation, oxidation of metals.
Examples: $\text{CH}_{4} + 2\text{O}_{2} \to \text{CO}_{2} + 2\text{H}_{2}\text{O}$ , $\text{HCl} + \text{NaOH} \to \text{NaCl} + \text{H}_{2}\text{O}$ .

Endothermic. Absorbs energy from the surroundings. The reaction mixture gets COLDER.

$\Delta H$ is POSITIVE (energy ENTERED the system).
Common: thermal decomposition, photosynthesis, dissolving certain salts (e.g. ammonium nitrate in water).
Examples: $\text{CaCO}_{3} \to \text{CaO} + \text{CO}_{2}$ .

Sign convention. $\Delta H$ is the ENTHALPY CHANGE — energy of products minus energy of reactants.

Worked qualitative. A student dissolves ammonium nitrate in water; the temperature drops from $20\degree\text{C}$ to $5\degree\text{C}$ . Endothermic; $\Delta H$ positive.

Exo: heat OUT, $\Delta H < 0$ , mixture warms.
Endo: heat IN, $\Delta H > 0$ , mixture cools.
$\Delta H = E(\text{products}) - E(\text{reactants})$ .
Common exo: combustion, neutralisation.
Common endo: thermal decomposition.

Energy level diagrams

Plot energy on y-axis. Reactants on left, products on right. $\Delta H$ between them. $E_a$ is the hump in the middle.

Drawing the diagram.

Draw a horizontal line for the reactants (energy level).
Draw a horizontal line for the products at a higher (endo) or lower (exo) energy.
Connect with a curve that rises to a peak and falls — that peak is the transition state.
The vertical distance from reactants to peak = activation energy $E_a$ .
The vertical distance from reactants to products = $\Delta H$ (negative for exo, positive for endo).

Activation energy. The minimum energy that colliding particles need for a successful reaction. Provides the energy to break starting bonds (the 'bond-breaking peak').

Catalysts lower $E_a$ — they provide an alternative reaction pathway with a smaller hump. They DON'T change $\Delta H$ (the products and reactants have the same energy as before).

Exothermic: products below reactants. Endothermic: products above. The hump is the activation energy.

Worked qualitative. Adding a catalyst to a reaction doubles the rate. What does this look like on the energy diagram? Same reactants and products (same energy levels). The hump is LOWER (smaller $E_a$ ). $\Delta H$ unchanged.

Tip. Always label the energy axis ("energy" or " $E$ "), the horizontal axis ("reaction progress"), and clearly mark $\Delta H$ and $E_a$ with arrows.

Reactants → transition state → products.
Hump = activation energy.
$\Delta H$ = level change, reactants → products.
Catalyst lowers $E_a$ , $\Delta H$ unchanged.

Bond energies and $\Delta H$ (Extended)

Energy IN to break bonds; energy OUT when bonds form. Net = $\Delta H$ .

Bond energy. The energy needed to break ONE MOLE of a particular bond in the gas phase. Typical values in $\text{kJ/mol}$ :

C–H: $413$
C=O: $805$
O=O: $498$
O–H: $464$
C–C: $347$
C=C: $612$

Calculating $\Delta H$ . $\Delta H = \sum E(\text{bonds broken}) - \sum E(\text{bonds made}).$

Breaking bonds: ENERGY IN (positive contribution). Making bonds: ENERGY OUT (subtracted, but the formed-bond energies appear as negative contributions to $\Delta H$ ).

Energy is absorbed to break bonds and released to make them; ΔH is the net difference.

Worked. Combustion of methane: $\text{CH}_{4} + 2\text{O}_{2} \to \text{CO}_{2} + 2\text{H}_{2}\text{O}$ .

Bonds broken in reactants: $4 \times \text{C–H} + 2 \times \text{O=O} = 4 \times 413 + 2 \times 498 = 2648\,\text{kJ}$ .
Bonds made in products: $2 \times \text{C=O} + 4 \times \text{O–H} = 2 \times 805 + 4 \times 464 = 3466\,\text{kJ}$ .
$\Delta H = 2648 - 3466 = -818\,\text{kJ/mol}$ .

Negative answer → exothermic. That's why methane combustion releases heat — the bonds in CO₂ and H₂O are stronger (release more energy) than the bonds in CH₄ and O₂ (absorbed less).

Cambridge tip. Be very careful with signs. $\Delta H$ NEGATIVE = exothermic. POSITIVE = endothermic.

Bond energy: kJ/mol to BREAK a bond.
Breaking: energy IN. Making: energy OUT.
$\Delta H = \Sigma(\text{broken}) - \Sigma(\text{made})$ .
Negative $\Delta H$ → exothermic.

How it’s examined

Energetics appears every Paper 4 (6-8 marks): label an energy diagram, distinguish exo/endo, calculate $\Delta H$ from bond energies. Examiner reports flag sign errors in $\Delta H$ calculations and forgetting to multiply by bond counts (e.g. CH₄ has 4 C-H bonds).

Sources: Cambridge IGCSE Chemistry 0620 syllabus 2026-2028 (6.1-6.2); 0620/42 Oct/Nov 2024 — Q10 (bond energies); 0620 Examiner Reports 2022-2024. Last reviewed 2026-05-06.

Master this Topic

Worked examples, formulae, definitions and the mistakes examiners flag — everything you need to push from a pass to an A*.

Take this whole topic with you

Download a branded revision sheet — worked examples, formulae, definitions and common mistakes for Chemical Energetics, ready to print or save as PDF.

Step-by-step worked examples — Chemical Energetics

Step-by-step solutions to past-paper-style questions on chemical energetics, written exactly the way a tutor would explain them at the board.

1Classify reactions
Core• exo/endo
Question
Classify each as exothermic or endothermic: combustion of methane, photosynthesis, neutralisation of HCl + NaOH, thermal decomposition of CaCO $_{3}$ .
Step-by-step solution
1. Step 1
  Exothermic: gives out heat (surroundings warm up). Endothermic: takes in heat (surroundings cool).
Answer
Exo: combustion of methane, neutralisation. Endo: photosynthesis, thermal decomposition.
2Draw an energy-level diagram (exothermic)
Core• diagram
Question
Sketch the energy-level diagram for the combustion of methane, marking $\Delta H$ and the activation energy $E_{a}$ .
Step-by-step solution
1. Step 1
  Reactants higher than products → exothermic.
2. Step 2
  $\Delta H$ = products − reactants → NEGATIVE.
3. Step 3
  $E_{a}$ = peak height above reactant level.
Answer
Reactants high → peak (transition state) → products low. $\Delta H < 0$ .
3Use bond energies to find $\Delta H$
Extended• Adapted from 0620/42 May/Jun 2024 Q14• bond energy
Question
Use bond energies to estimate $\Delta H$ for $\text{H}_{2} + \text{Cl}_{2} \to 2\text{HCl}$ . Given: $E(\text{H-H}) = 436$ , $E(\text{Cl-Cl}) = 242$ , $E(\text{H-Cl}) = 431\,\text{kJ/mol}$ .
Step-by-step solution
1. Step 1
  Bonds broken (energy IN).
  $436 + 242 = 678$
2. Step 2
  Bonds made (energy OUT).
  $2 \times 431 = 862$
3. Step 3
  $\Delta H$ = bonds broken − bonds made.
  $\Delta H = 678 - 862 = -184\,\text{kJ/mol}$
Answer
$\Delta H = -184\,\text{kJ/mol}$ (exothermic)
4Energy from temperature rise
Extended• calorimetry
Question
When $50\,\text{cm}^{3}$ of HCl reacts with $50\,\text{cm}^{3}$ of NaOH, the temperature rises by $6.5\,\degree\text{C}$ . Find the energy released. ( $c = 4.2\,\text{J/(g\,K)}$ .)
Step-by-step solution
1. Step 1
  Total mass = $50 + 50 = 100\,\text{g}$ (assuming density of water).
2. Step 2
  $E = mc\Delta T$ .
  $E = 100 \times 4.2 \times 6.5 = 2730\,\text{J}$
Answer
$2730\,\text{J} \approx 2.73\,\text{kJ}$

Key Formulae — Chemical Energetics

The formulae you need to memorise for chemical energetics on the Cambridge IGCSE 0620 paper, with every variable defined in plain English and a note on when to use it.

$\Delta H$ from bond energies
$\Delta H = \Sigma(\text{bonds broken}) - \Sigma(\text{bonds made})$
When to use
When bond energies (in kJ/mol) are given for all reactants and products.
Energy from temperature change
$E = m c \Delta T$
$c$
specific heat capacity (J/(g·K))
When to use
Calorimetry / thermochemistry experiments.

Key Definitions and Keywords — Chemical Energetics

Definitions to memorise and the exact keywords mark schemes credit for chemical energetics answers — sharpened from recent examiner reports for the 2026 0620 sitting.

Exothermic reaction
Examiner keyword
A reaction that transfers energy TO the surroundings (temperature rises). $\Delta H$ negative.
Endothermic reaction
Examiner keyword
A reaction that takes in energy FROM the surroundings (temperature drops). $\Delta H$ positive.
Activation energy ( $E_{a}$ )
Examiner keyword
Minimum energy reactants must have for a reaction to occur.
Enthalpy change ( $\Delta H$ )
Examiner keyword
Energy change at constant pressure. $\Delta H = \text{products} - \text{reactants}$ .

Common Mistakes and Misconceptions — Chemical Energetics

The traps other students keep falling into on chemical energetics questions — taken from recent Cambridge IGCSE 0620 examiner reports and mark schemes — and how to avoid them.

✕Writing $\Delta H = +$ for an exothermic reaction
0620/42 — recurring
Why it happens
Forgetting the convention.
How to avoid it
Exothermic releases energy → products are LOWER → $\Delta H < 0$ (NEGATIVE).
✕Using 'bonds made − bonds broken'
Why it happens
Inverting the formula.
How to avoid it
$\Delta H =$ broken − made. (Energy needed to break minus energy released on making.)
✕Counting only one bond per molecule
Why it happens
Missing equation coefficients.
How to avoid it
Multiply each bond energy by the number of those bonds in the balanced equation.
✕Saying temperature rise = endothermic
Why it happens
Confusing direction.
How to avoid it
Surroundings WARM UP → energy released by reaction → EXOTHERMIC.

Practice questions

Exam-style questions with step-by-step worked solutions. Try one before checking the method.

Past paper style quiz

Get a report showing which sub-topics you've nailed and which ones still need work.

Video lesson

Short walkthrough of the concepts students most often get stuck on.

Chemical Energetics — frequently asked questions

The things students keep getting wrong in this sub-topic, answered.

Chemical Energetics

Short Study Notes in the form of Slides

Detailed Notes

What you’ll learn

How it’s examined

1Classify reactions

2Draw an energy-level diagram (exothermic)

3Use bond energies to find ΔH\Delta HΔH

4Energy from temperature rise

ΔH\Delta HΔH from bond energies

Energy from temperature change

Exothermic reaction

Endothermic reaction

Activation energy (EaE_{a}Ea​)

Enthalpy change (ΔH\Delta HΔH)

✕Writing ΔH=+\Delta H = +ΔH=+ for an exothermic reaction

✕Using 'bonds made − bonds broken'

✕Counting only one bond per molecule

✕Saying temperature rise = endothermic

Practice questions

Past paper style quiz

Assess your understanding

Video lesson

Chemical Energetics — frequently asked questions

3Use bond energies to find $\Delta H$

$\Delta H$ from bond energies

Activation energy ( $E_{a}$ )

Enthalpy change ( $\Delta H$ )

✕Writing $\Delta H = +$ for an exothermic reaction