Overview

What you'll learn

Distinguish exo and endo. Know the sign of ΔH, draw energy-level diagrams, and give examples of each type of reaction.

Use bond energies. Apply ΔH = (energy to break LHS bonds) − (energy to form RHS bonds) to calculate an enthalpy change.

Read energy profiles. Identify reactants, products, activation energy and ΔH on an energy profile diagram.

Explain the catalyst effect. A catalyst lowers Eₐ but does not change ΔH — and you should be able to say why.

Tutor's Insight

"Small topic. Tight rules — drill them once and these are free marks."
Energetics is one of the shorter topics, but the questions are very predictable: MCQ for the sign of ΔH, structured questions for the bond-energy calculations. Predict whether a reaction is exothermic or endothermic from its bond changes, compute ΔH from a bond-energy table, and read an energy profile to explain what a catalyst does. One equation does the heavy lifting — broken minus formed — so once you have drilled it, these become reliable marks.

9.1 Energy Changes

9.1 Energy Changes

Syllabus asks
  • Describe the meaning of exothermic and endothermic reactions, and the transfer of heat energy between the system and the surroundings.
  • State the sign of the enthalpy change ΔH for exothermic (negative) and endothermic (positive) reactions, and represent them on energy-level diagrams.

9.1 Energy Changes

System, Surroundings & ΔH

The System
The chemicals reacting
The substances in the test tube or flask — the reactants and products. Their bonds are where chemical energy is stored.
The Surroundings
Everything else
The water bath, the flask, the air. Whatever the reaction can swap heat energy with — measured as a temperature change.
ΔH
The enthalpy change
ΔH = (energy of products) − (energy of reactants). Negative → energy left the system (exothermic). Positive → energy entered (endothermic).
Heat moves; energy is conserved. If the system loses energy, the surroundings gain it — an exothermic reaction warms the flask; an endothermic one cools it. Always describe both sides: the chemical energy of the bonds and the heat flowing in or out.

Free Notes · O-Level Pure Chemistry

Read the full chapter

Exothermic vs endothermic, energy-level diagrams, bond breaking and forming, the bond-energy table, a worked ΔH calculation, energy profiles, catalysts and 4 worked exam questions.

9.1 Energy Changes

Exothermic vs Endothermic

Exothermic
Releases heat to the surroundings
ΔH is negative. Products at lower energy than reactants.

Examples:
• combustion of fuels
• neutralisation (acid + alkali)
• reactive metal + water / acid
• respiration
• freezing, condensing (state changes)
Endothermic
Absorbs heat from the surroundings
ΔH is positive. Products at higher energy than reactants.

Examples:
• photosynthesis
• thermal decomposition (e.g. CuCO₃ → CuO + CO₂)
• dissolving certain salts (NH₄NO₃ in water)
• melting, boiling (state changes)
Sign convention: negative ΔH → exothermic (heat out, flask warms). Positive ΔH → endothermic (heat in, flask cools).

9.1 Energy Changes

Energy-Level Diagrams

Exothermic · ΔH < 0
Reactants high, products low
Reactants start at a higher energy level; products drop to a lower level. The downward step from reactants to products is ΔH, which is negative because energy was released.
Endothermic · ΔH > 0
Reactants low, products high
Reactants start at a lower energy level; products rise to a higher level. The upward step from reactants to products is ΔH, which is positive because energy was absorbed.
Exam Habit

Label everything — Always label the axes (energy vs reaction progress), the reactants, the products and ΔH. The arrow points from reactants to products, and its direction shows the sign of ΔH.

9.2 Bond Energy & Energy Profiles

9.2 Bond Energy & Energy Profiles

Syllabus asks
  • Explain that bond breaking is endothermic and bond forming is exothermic, and use bond energies to calculate the enthalpy change of a reaction.
  • Interpret energy profile diagrams, identify the activation energy Eₐ and ΔH, and describe the effect of a catalyst on the activation energy.

9.2 Bond Energy & Energy Profiles

Breaking vs Forming Bonds

Bond Breaking
Endothermic
It takes energy to pull two atoms apart against their attraction. Energy is absorbed from the surroundings — the system gains it.
Bond Forming
Exothermic
When atoms come together, the bond they form is at a lower energy than the separate atoms. The leftover energy is released to the surroundings.
Reading the result
Broken minus formed

Result negative → forming releases more than breaking absorbed → exothermic.


Result positive → forming releases less than breaking absorbed → endothermic.

The Master Equation
ΔH =
energy to break LHS bonds − energy to form RHS bonds
Add up every bond broken in the reactants, subtract every bond formed in the products.

9.2 Bond Energy & Energy Profiles

Bond Energy Table

Bond Type Bond energy (kJ/mol)
H–H single 436
O=O double 496
N≡N triple 944
O–H single 460
C–C single 350
C=C double 610
Pattern: more shared pairs → stronger bond → higher energy. C–C (350) < C=C (610), and triple-bonded N≡N (944) is one of the strongest bonds at O-level.

9.2 Bond Energy & Energy Profiles

From Bonds to ΔH, in Four Steps

The Reaction

H₂ (g) + Cl₂ (g) → 2 HCl (g)

Bonds in reactants: 1 × H–H + 1 × Cl–Cl

Bonds in products: 2 × H–Cl


Bond energies (kJ/mol):
H–H = 436  ·  Cl–Cl = 242  ·  H–Cl = 431

Worked Answer
ΔH = −184 kJ/mol (exothermic)
  1. List the bonds broken (LHS) and formed (RHS).
  2. Energy to break the reactants: 436 + 242 = 678 kJ/mol.
  3. Energy released forming the products: 2 × 431 = 862 kJ/mol.
  4. ΔH = 678 − 862 = −184 kJ/mol. Negative → exothermic.

9.2 Bond Energy & Energy Profiles

Energy Profile Diagrams

Activation Energy (Eₐ)
The minimum energy for an effective collision
On a profile diagram, Eₐ is the height of the hump above the reactants. The particles need ≥ Eₐ to react — and they need to collide in the right orientation.
Reading the Diagram
Hump for Eₐ, step for ΔH
• Reactants on the left, products on the right.
• Eₐ = top of the hump − reactant level.
• ΔH = product level − reactant level (signed).

9.2 Bond Energy & Energy Profiles

Activation Energy & Catalysts

What a catalyst does
Lowers Eₐ
The hump gets shorter, so more collisions have enough energy — the reaction speeds up.
What it does not do
Does not change ΔH
The reactants and products sit at the same energies as before, so the step between them is unchanged.
Why it matters

A catalyst lowers Eₐ but does not change ΔH. This is a favourite MCQ trap: a shorter hump on the profile does not mean a different enthalpy change — only the reactants and products decide ΔH, and the catalyst leaves them exactly where they were.

Practice

Exam-style questions

Question 01  ·  ACS(BR) 2024 Prelim
The Question
MCQ

Combustion of candle wax

The conversion of candle wax (C₃₀H₆₂) into CO₂ and H₂O proceeds in four stages:

  • 1  — C₃₀H₆₂ (s) → C₃₀H₆₂ (l)
  • 2  — C₃₀H₆₂ (l) → C₃₀H₆₂ (g)
  • 3  — 2C₃₀H₆₂ (g) + 90O₂ (g) → 60CO₂ (g) + 62H₂O (g)
  • 4  — 60CO₂ (g) + 62H₂O (g) → 60CO₂ (g) + 62H₂O (l)

Which stages are exothermic?

A 1 and 3    B 1 and 4    C 2 and 3    D 3 and 4

Worked Answer
D — 3 and 4
  1. Stage 1: solid → liquid (melting). Particles overcome attractions → absorbs heat → endothermic.
  2. Stage 2: liquid → gas (vaporising). Endothermic, for the same reason.
  3. Stage 3: combustion. Stronger bonds form in CO₂ and H₂O than were in C₃₀H₆₂ and O₂ → releases heat → exothermic.
  4. Stage 4: gas → liquid (condensing). Bonds form between H₂O molecules → exothermic. Answer D.
Question 02  ·  ACS(BR) 2024 Prelim
The Question
MCQ

Burning hydrogen

Hydrogen burns in air to form water:

2H₂ (g) + O₂ (g) → 2H₂O (l)

Which statement about the enthalpy change ΔH is correct?

  • A  ΔH negative — energy to break bonds is greater than energy to form bonds
  • B  ΔH negative — energy to break bonds is less than energy to form bonds
  • C  ΔH positive — energy to break bonds is greater than energy to form bonds
  • D  ΔH positive — energy to break bonds is less than energy to form bonds
Worked Answer
B
  1. Combustion of hydrogen is exothermic — that rules out C and D.
  2. Exothermic means net energy is released → bond forming releases MORE than bond breaking absorbed.
  3. i.e. energy to break < energy to form.
  4. Answer B.
Question 03  ·  CCHM 2024 Prelim
The Question
MCQ

Hydrogenation of ethyne

Ethyne reacts with hydrogen to form ethane:

C₂H₂ + 2H₂ → C₂H₆

Bond energies (kJ/mol):
C–H = 413  ·  C–C = 347  ·  C≡C = 839  ·  H–H = 432

What is ΔH for the reaction?

A +51    B −176    C −296    D −728 kJ/mol

Worked Answer
C — −296 kJ/mol
  1. Break (LHS): C₂H₂ has 2 C–H + 1 C≡C; 2H₂ has 2 H–H. So 2(413) + 839 + 2(432) = 2529 kJ.
  2. Form (RHS): C₂H₆ has 6 C–H + 1 C–C. So 6(413) + 347 = 2825 kJ.
  3. ΔH = broken − formed = 2529 − 2825 = −296 kJ/mol.
  4. Negative → exothermic. Answer C.
Question 04
The Question
Structured

Forming hydrogen chloride, and the catalyst

The formation of hydrogen chloride is exothermic:

H₂ (g) + Cl₂ (g) → 2HCl (g)    ΔH = −184 kJ/mol

(a) State, with a reason, whether the products or the reactants are at the higher energy.
(b) A catalyst is added. State the effect on the activation energy and on ΔH.

Worked Answer
Exothermic → products lower
  1. (a) The reactants are at the higher energy. ΔH is negative (exothermic), so the products sit at a lower energy level than the reactants — energy was released to the surroundings.
  2. (b) The catalyst lowers the activation energy Eₐ, so more collisions are effective and the reaction speeds up.
  3. (b) The catalyst does not change ΔH — the reactants and products stay at the same energies, so ΔH stays at −184 kJ/mol.

Frequently Asked Questions

Chemical Energetics — FAQ

What is the difference between an exothermic and an endothermic reaction?
An exothermic reaction releases heat to the surroundings, so ΔH is negative and the products sit at a lower energy than the reactants — examples include combustion, neutralisation and respiration. An endothermic reaction absorbs heat from the surroundings, so ΔH is positive and the products sit at a higher energy than the reactants — examples include photosynthesis and thermal decomposition. An exothermic reaction warms the flask; an endothermic one cools it.
What does the sign of ΔH tell you?
ΔH = (energy of products) − (energy of reactants). A negative ΔH means energy left the system, so the reaction is exothermic. A positive ΔH means energy entered the system, so the reaction is endothermic. Remember: negative is exothermic, positive is endothermic.
Why is bond breaking endothermic and bond forming exothermic?
Bond breaking is endothermic because it takes energy to pull two atoms apart against their attraction, so energy is absorbed from the surroundings. Bond forming is exothermic because when atoms come together the bond they form is at a lower energy than the separate atoms, so the leftover energy is released to the surroundings.
How do you calculate ΔH from bond energies?
Use ΔH = (energy to break all the bonds in the reactants) − (energy to form all the bonds in the products). Add up the bond energies on each side using the bond-energy table, then subtract. If the result is negative the reaction is exothermic; if positive it is endothermic.
What is activation energy on an energy profile diagram?
Activation energy (Eₐ) is the minimum energy needed for an effective collision. On an energy profile diagram it is the height of the hump above the reactant level. Particles must have at least Eₐ and collide in the correct orientation to react.
How does a catalyst change the energy profile?
A catalyst lowers the activation energy (Eₐ) but does not change ΔH. The reactants and products stay at the same energies, so the step between them is unchanged — only the hump gets shorter. This means more collisions have enough energy, so the reaction speeds up.

O-Level Pure Chemistry  ·  Syllabus 6092  ·  Topic 9 of 12  ·  © 2026 Overmugged. For personal study use only.