Enthalpy Change

Summary: Enthalpy change (ΔH) is the amount of thermal energy transferred between a chemical reaction and its surroundings at constant pressure, measured in kJ/mol. Reactions with a negative ΔH release energy to the surroundings (exothermic), while those with a positive ΔH absorb energy from the surroundings (endothermic). ΔH can be calculated from bond energies, measured experimentally by calorimetry, and visualised on reaction pathway diagrams.

Tags: igcse chemistry enthalpy energetics exothermic endothermic bond-energy reaction-pathway-diagram delta-h

Created: 2026-07-13

Enthalpy change, denoted ΔH, is the heat energy transferred between a chemical system and its surroundings when a reaction is carried out at constant pressure — the most common condition in a school laboratory. It is a quantitative measure of the difference in energy content between the products and the reactants, expressed in kilojoules per mole (kJ/mol), and its sign tells us the direction of heat flow: a negative ΔH means the reaction is exothermic (energy is released, usually as heat, causing the surroundings to warm up), whereas a positive ΔH means the reaction is endothermic (energy is absorbed from the surroundings, causing them to cool down). Every reaction involves two competing energy processes — breaking bonds in the reactants (which requires an input of energy and is therefore endothermic) and making bonds in the products (which releases energy and is therefore exothermic) — and ΔH is simply the net balance of these two. Because bond energies are known and tabulated (as average values in kJ/mol), ΔH for a reaction can be predicted theoretically without ever performing the experiment, using the formula ΔH = Σ(bond energies broken) − Σ(bond energies formed). On a Reaction Pathway Diagram, ΔH appears as the vertical arrow between the energy levels of reactants and products, while the hump preceding it represents the Activation Energy, the minimum energy barrier that must be overcome for the reaction to begin.

Definition and Sign Convention

Enthalpy (H) is the total heat content of a chemical system at constant pressure. While the absolute enthalpy of a substance cannot be measured directly, the enthalpy change (ΔH) — the difference between the enthalpy of products and reactants — can be measured or calculated:

[ \Delta H = H_{\text{products}} - H_{\text{reactants}} ]

The sign of ΔH tells us what type of reaction has occurred:

Units

ΔH is expressed in kilojoules per mole (kJ/mol) — the “per mole” refers to one mole of the reaction as written in the balanced equation, not necessarily one mole of a particular reactant or product. For example, the combustion of methane:

[ \text{CH}_4(g) + 2\text{O}_2(g) \rightarrow \text{CO}_2(g) + 2\text{H}_2\text{O}(l) \quad \Delta H = -890 \text{ kJ/mol} ]

Here, -890 kJ/mol means that 890 kJ of energy is released when 1 mole of CH₄ reacts with 2 moles of O₂ to give 1 mole of CO₂ and 2 moles of H₂O.

Bond Breaking and Bond Making

Every chemical reaction involves two fundamental processes that underpin the overall enthalpy change:

Bond Breaking is Endothermic

Energy must be supplied to break a chemical bond, pulling atoms apart against the attractive forces that hold them together. This energy input is the bond energy (also called bond dissociation energy), defined as the energy required to break one mole of a given covalent bond in the gaseous state. Because energy is absorbed, bond breaking contributes a positive term to ΔH.

[ \text{Energy absorbed to break bonds} = \sum (\text{bond energies of all bonds in reactants}) ]

Bond Making is Exothermic

When new bonds form in the products, energy is released as the atoms settle into a more stable, lower-energy arrangement. The magnitude of energy released equals the bond energy of the newly formed bond. Because energy is released, bond making contributes a negative term to ΔH.

[ \text{Energy released forming bonds} = -\sum (\text{bond energies of all bonds in products}) ]

Net Enthalpy Change

The overall ΔH is the balance of these two opposing contributions:

[ \Delta H = \sum (\text{bond energies broken}) - \sum (\text{bond energies formed}) ]

• If more energy is released in bond making than is absorbed in bond breaking, the reaction is exothermic (ΔH is negative).
• If more energy is absorbed in bond breaking than is released in bond making, the reaction is endothermic (ΔH is positive).

This net view explains why chemical reactions are rarely ever purely endothermic or exothermic on a bond-by-bond basis — every reaction involves both endothermic and exothermic steps simultaneously.

Bond Energy Calculations

Bond energy calculations allow ΔH to be determined purely from tabulated average bond energies, without the need for experiment. This approach works for any reaction where bond energies for all bonds in reactants and products are known.

Key Steps

1. Write the balanced equation and draw out the displayed (structural) formula of each molecule to identify every bond present.
2. Calculate the total energy required to break all bonds in the reactants: sum the bond energies of every bond that must be broken.
3. Calculate the total energy released when all bonds in the products are formed: sum the bond energies of every bond that is made.
4. Apply the formula: ΔH = (total energy to break bonds) − (total energy released making bonds).

Worked Example: Combustion of Methane

[ \text{CH}_4(g) + 2\text{O}_2(g) \rightarrow \text{CO}_2(g) + 2\text{H}_2\text{O}(g) ]

Bond energies (kJ/mol): C–H = 413, O=O = 498, C=O = 799, O–H = 464

Bonds broken (reactants):

Bonds formed (products):

[ \Delta H = 2648 - 3454 = -806 \text{ kJ/mol} ]

The negative value confirms the reaction is exothermic. (Note: bond energy calculations using average values give ΔH ≈ −806 kJ/mol, while experimental measurements give −890 kJ/mol. The discrepancy arises because average bond energies are used rather than the specific bond energies for the exact molecular environment.)

Worked Example: Formation of Hydrogen Iodide

[ \text{H}_2(g) + \text{I}_2(g) \rightarrow 2\text{HI}(g) ]

Bond energies (kJ/mol): H–H = 436, I–I = 151, H–I = 299

Bonds broken: 1 × H–H (436) + 1 × I–I (151) = 587 kJ/mol Bonds formed: 2 × H–I = 2 × 299 = 598 kJ/mol

[ \Delta H = 587 - 598 = -11 \text{ kJ/mol} ]

The reaction is slightly exothermic; the energy released forming two H–I bonds very slightly exceeds the energy required to break one H–H and one I–I bond.

Reaction Pathway Diagrams

A Reaction Pathway Diagram (also called an energy profile diagram) is a graphical representation of the energy changes during a chemical reaction. It plots energy (enthalpy) on the vertical axis against the progress of the reaction (sometimes labelled “reaction coordinate” or “extent of reaction”) on the horizontal axis.

Features of the Diagram

• Reactants are shown at a certain energy level on the left.
• Products are shown at a certain energy level on the right.
• ΔH is the vertical difference between the energy levels of products and reactants.
• Activation energy (Eₐ) is the energy barrier (the “hump”) between reactants and products — the minimum energy colliding particles must possess for a reaction to occur.
• The peak of the curve represents the transition state (or activated complex), where bonds are partially broken and partially formed.

Exothermic Reaction Pathway

Energy
  |    Transition state
  |       /\
  |      /  \
  |     /    \    Eₐ
  |    /      \
  |   /Reactant\
  |  /          \___________ Products
  |  |----------| ΔH (negative)
  |___________________________ Progress of reaction

In an exothermic reaction, products sit at a lower energy level than reactants. The vertical drop from reactants to products is the negative ΔH — energy has been transferred to the surroundings.

Endothermic Reaction Pathway

Energy
  |                 /\
  |                /  \
  |      Products /    \
  |     _________/      \ Eₐ
  |    |----------|      \
  |    | ΔH (+)  |       \
  |    |         |        \
  |   Reactants  |         \
  |___________________________ Progress of reaction

In an endothermic reaction, products sit at a higher energy level than reactants. The vertical rise from reactants to products is the positive ΔH — energy has been absorbed from the surroundings.

Reading ΔH from a Diagram

ΔH is always the vertical arrow between the reactant energy level and the product energy level. It is independent of the activation energy — a reaction can have a large activation energy and still be strongly exothermic, or a small activation energy and be endothermic. The two quantities are unrelated: Eₐ determines the rate of the reaction, while ΔH determines the heat transfer with the surroundings.

Measuring Enthalpy Change: Calorimetry

In the laboratory, ΔH is measured experimentally using a calorimeter — typically a polystyrene cup (to minimise heat loss) with a thermometer, placed inside a beaker for support. The method is:

1. Measure a known volume and concentration of one reactant into the cup; record its initial temperature.
2. Add the second reactant, stir, and record the highest (exothermic) or lowest (endothermic) temperature reached.
3. Calculate the temperature change: ΔT = T(final) − T(initial).
4. Calculate the heat energy transferred using:

[ q = mc\Delta T ]

where:

• q = heat energy (J)
• m = mass of the solution (g) — approximated by assuming the density of a dilute aqueous solution is 1 g/cm³, so volume in cm³ equals mass in g
• c = specific heat capacity of water (4.2 J/g°C)
• ΔT = temperature change (°C)

5. Convert q to kJ and divide by the number of moles of the limiting reactant to obtain ΔH in kJ/mol:

[ \Delta H = \frac{-q}{n} \quad \text{(negative sign accounts for whether temperature rose or fell)} ]

If the temperature increased (exothermic), ΔH is negative. If the temperature decreased (endothermic), ΔH is positive.

Sources of Error in Calorimetry

Examples of Enthalpy Changes

Exothermic Reactions (ΔH negative)

Endothermic Reactions (ΔH positive)

Related Pages

• Exothermic Reaction — detailed treatment of reactions that release energy
• Endothermic Reaction — detailed treatment of reactions that absorb energy
• Bond Energy Calculations — step-by-step worked examples and practice problems
• Bond Breaking and Bond Making — the two processes that determine ΔH
• Reaction Pathway Diagram — drawing and interpreting energy profile diagrams
• Activation Energy — the energy barrier that controls reaction rate
• Calorimetry — experimental methods for measuring enthalpy changes
• Combustion — a key class of exothermic reactions
• Photosynthesis — the archetypal endothermic reaction
• Bond Energy — tabulated average bond energies
• Energetics — broader topic overview for IGCSE Chemistry
• Mole — essential for converting measured quantities to molar enthalpy changes

Sources

• Cambridge IGCSE Chemistry 0620 Syllabus, Section 5.1: Exothermic and Endothermic Reactions
• Harwood, R. & Lodge, I., Cambridge IGCSE Chemistry Coursebook, 5th Edition, Cambridge University Press, 2021.
• Gallagher, R. & Ingram, P., Complete Chemistry for Cambridge IGCSE, 3rd Edition, Oxford University Press, 2016.
• Clegg, A. & Renshaw, J., Essential Chemistry for Cambridge IGCSE, 2nd Edition, Oxford University Press, 2018.
• Petrucci, R. H. et al., General Chemistry: Principles and Modern Applications, 11th Edition, Pearson, 2017.
• Cambridge Assessment International Education, Cambridge IGCSE Chemistry 0620 Learner Guide, 2023.

Common Misconceptions