Endothermic Reaction

Summary: An endothermic reaction is a chemical reaction that absorbs thermal energy from its surroundings, causing a decrease in the temperature of the surroundings and producing products with a higher total energy than the reactants. The enthalpy change (ΔH) for an endothermic reaction is always positive. Tags: endothermic enthalpy energetics igcse-chemistry thermochemistry reaction-pathway bond-energy Created: 2026-07-18

An endothermic reaction is a chemical reaction in which the reacting system takes in thermal energy from its surroundings, and this net absorption of energy results in a drop in the temperature of the surroundings that can be measured with a thermometer. Because energy is gained by the chemical system, the products possess a greater total energy than the reactants, and the enthalpy change of the reaction, ΔH, is given a positive sign (ΔH > 0) to indicate that heat has entered the system from the environment. On a Reaction Pathway Diagram, endothermic reactions are represented by a curve where the products sit at a higher energy level than the reactants, with the vertical arrow pointing upward to show the net gain in energy. At the molecular level, endothermic change is intimately linked to Bond Breaking: breaking chemical bonds always requires an input of energy, so any reaction in which the total energy absorbed to break bonds in the reactants exceeds the total energy released when new bonds form in the products will be endothermic overall. Familiar examples include the thermal decomposition of calcium carbonate into calcium oxide and carbon dioxide, the dissolution of ammonium nitrate in water (used in instant cold packs), and the process of photosynthesis, in which plants absorb sunlight to convert carbon dioxide and water into glucose and oxygen.

Energy Changes in Endothermic Reactions

In an endothermic reaction, energy is transferred from the surroundings into the chemical system. This has several measurable consequences:

• The temperature of the surroundings decreases (the reaction vessel feels cold to the touch).
• The enthalpy change (ΔH) is positive: ΔH > 0.
• The products contain more stored chemical energy than the reactants.

This is the opposite of an Exothermic Reaction, where energy is released to the surroundings and ΔH is negative.

Reaction Pathway Diagram

A reaction pathway diagram (also called an energy profile or enthalpy level diagram) for an endothermic reaction shows:

Energy
  ^
  |        Products
  |        /\
  |       /  \
  |      /    \  ← Activation energy (Ea)
  |     /      \
  |    /        \
  |   /          \
  |  /   Reactants \
  | /______________\
  +------------------------> Progress of reaction
  
  ΔH = positive (upward arrow)

Key features:

• The products are at a higher energy level than the reactants.
• The vertical distance between reactants and products represents ΔH, the enthalpy change of the reaction — positive for endothermic reactions.
• The peak of the curve represents the transition state; the energy required to reach this peak from the reactants is the activation energy (E₁).
• Even though the overall reaction is endothermic, an activation energy must still be supplied to initiate the reaction (e.g. heating calcium carbonate to begin thermal decomposition).

For comparison, see the Reaction Pathway Diagram page, which covers both endothermic and exothermic profiles.

Bond Breaking and Bond Making

Every chemical reaction involves two competing energy processes:

A reaction is endothermic overall when:

Energy absorbed to break bonds > Energy released when making bonds

This means the bonds in the reactants are, on average, stronger than the bonds in the products — more energy must be put in to pull the reactant atoms apart than is given back when the product bonds form. For a worked calculation method, see Bond Breaking and Enthalpy Change.

Worked Example

Consider the thermal decomposition of calcium carbonate:

CaCO₃(s) → CaO(s) + CO₂(g)    ΔH = +178 kJ mol⁻¹

The positive ΔH value confirms the reaction is endothermic. For every mole of CaCO₃ decomposed, 178 kJ of energy must be absorbed from the surroundings. This is why the reaction is carried out in a lime kiln at high temperatures (~900–1000 °C) — the continuous supply of thermal energy drives the decomposition forward.

Common Examples of Endothermic Reactions

1. Thermal Decomposition of Metal Carbonates

When a metal carbonate is heated, it breaks down into a metal oxide and carbon dioxide:

CaCO₃(s) → CaO(s) + CO₂(g) CuCO₃(s) → CuO(s) + CO₂(g)

These reactions require strong, continuous heating. The carbon dioxide produced can be confirmed by bubbling the gas through limewater, which turns milky.

2. Photosynthesis

Photosynthesis is the most important endothermic reaction on Earth. Plants absorb light energy from the Sun (the surroundings) to drive the synthesis of glucose:

6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂

Light energy (mainly from the Sun) is absorbed and stored as chemical energy in glucose molecules. Without this endothermic process, life as we know it could not exist — the chemical energy in glucose ultimately fuels nearly all food chains.

IGCSE note: The syllabus treats light as a form of energy absorbed from the surroundings, making photosynthesis a clear example of an endothermic process.

3. Dissolving Ammonium Nitrate in Water

When ammonium nitrate (NH₄NO₃) dissolves in water, the temperature of the solution drops noticeably:

NH₄NO₃(s) + aq → NH₄⁺(aq) + NO₃⁻(aq)    ΔH = +25.7 kJ mol⁻¹

This is the principle behind instant cold packs used in first aid: an inner pouch of water is broken inside an outer pouch containing ammonium nitrate; the dissolving process absorbs heat from the injured area, providing cooling relief. Other salts such as potassium chloride also produce a cooling effect upon dissolution, though ammonium nitrate is the most commonly used.

4. Reaction of Citric Acid with Sodium Hydrogen Carbonate

Mixing citric acid solution with sodium hydrogen carbonate (sodium bicarbonate) produces a measurable temperature drop:

H⁺(aq) + HCO₃⁻(aq) → H₂O(l) + CO₂(g)

The effervescence (CO₂ gas) is visible, and the container feels cold. This is a common laboratory demonstration of an endothermic reaction.

5. Cracking of Hydrocarbons

The thermal decomposition (cracking) of long-chain alkanes into shorter, more useful alkanes and alkenes is endothermic:

C₁₀H₂₂ → C₈H₁₈ + C₂H₄ (decane → octane + ethene)

Cracking requires high temperatures (~600–700 °C) and a catalyst (catalytic cracking) or steam (steam cracking) to supply the necessary energy.

Measuring Endothermic Changes in the Lab

A simple calorimetry experiment can demonstrate an endothermic reaction:

1. Place a known volume of solution (e.g. citric acid) in a polystyrene cup (a simple calorimeter).
2. Record the initial temperature with a thermometer.
3. Add a known mass of the solid reactant (e.g. sodium hydrogen carbonate).
4. Stir continuously and record the lowest temperature reached.
5. Calculate the temperature change: ΔT = Tₓ₂₊₊₎ − Tₓ₍₍₍₍₃₊ (this will be negative for an endothermic reaction).
6. Use q = mcΔT to calculate the heat energy absorbed, remembering that ΔT is negative, so q will be negative — indicating heat absorbed by the system from the surroundings.

The polystyrene cup acts as insulation to minimise heat exchange with the external environment, improving the accuracy of the measurement.

Endothermic vs. Exothermic: Quick Comparison

Common Exam Question Types

1. Interpreting a reaction pathway diagram: Given a diagram, state whether the reaction is endothermic or exothermic by comparing the relative energy levels of reactants and products.

2. Explaining a temperature drop: “The temperature of the solution decreased. Explain why this means the reaction is endothermic.” Answer: The system absorbed energy from the surroundings, so the surroundings (solution) lost thermal energy and became cooler.

3. ΔH sign convention: “An endothermic reaction has ΔH = +45 kJ/mol. What does the positive sign indicate?” Answer: Energy is taken in (absorbed) by the reaction from the surroundings.

4. Bond energy calculations: Given bond energy data, calculate the overall ΔH and state whether the reaction is endothermic or exothermic. If the sum of bond energies broken exceeds the sum of bond energies made, the reaction is endothermic.

5. Identifying endothermic processes from data: “Substance X is heated strongly and decomposes. The mass decreases and a gas that turns limewater milky is given off. The reaction vessel must be heated continuously for the reaction to proceed.” This describes an endothermic thermal decomposition.

Sources

• Cambridge IGCSE Chemistry 0620 Syllabus, Section 5.1: Energetics of a Reaction
• Cambridge IGCSE Chemistry Coursebook (Jones & Jones, 4th Edition), Chapter 5: Chemical Energetics
• Cambridge IGCSE Chemistry Practical Workbook, Experiment 5.1: Measuring Enthalpy Changes
• Complete Chemistry for Cambridge IGCSE (Gallagher & Ingram, 3rd Edition), Chapter 9: Energy Changes and Reversible Reactions
• Royal Society of Chemistry (RSC), “Exothermic and Endothermic Reactions,” Education in Chemistry
• AQA GCSE Chemistry 8462, Topic 5.5: Energy Changes

Common Misconceptions