Bond Energy Calculations

Summary: Bond energy is the energy required to break 1 mole of bonds in the gaseous state. ΔH = Σ(bond energies of bonds broken) − Σ(bond energies of bonds formed). Bond energy calculations give approximate values because average bond energies are used. Bond breaking is endothermic; bond making is exothermic. Tags: igcse chemistry energetics bond-energy Created: 2026-07-14 Last Updated: 2026-07-16

What Is Bond Energy?

Bond energy (also called bond enthalpy or bond dissociation energy) is defined as:

The energy required to break 1 mole of a particular covalent bond in the gaseous state.

Bond energies are always given as positive values (in kJ/mol) because energy must be supplied to break a bond.

Example: The H-H bond energy is +436 kJ/mol.

H2(g) -> 2H(g)     ΔH = +436 kJ/mol

This means 436 kJ of energy is needed to break every mole of H-H bonds.

Key bond energies (approximate values in kJ/mol):

Bond energy values do not need to be memorised — they are always provided in questions. However, common values (H-H = 436, Cl-Cl = 242, H-Cl = 431) appear frequently.

Bond Breaking and Bond Making

Every chemical reaction involves two processes:

1. Breaking bonds in the reactants — energy is absorbed (endothermic, positive value)
2. Making bonds in the products — energy is released (exothermic, negative value)

The overall enthalpy change (ΔH) depends on which is greater — the energy absorbed breaking bonds or the energy released making bonds.

The Formula

ΔH = Σ(bond energies of bonds broken) − Σ(bond energies of bonds formed)

Where:

• Σ = sum of
• Bonds broken = all bonds in the reactants that are broken
• Bonds formed = all new bonds in the products that are formed

Interpretation of the result:

• If ΔH is negative: the reaction is exothermic (more energy released making bonds than absorbed breaking them)
• If ΔH is positive: the reaction is endothermic (more energy absorbed breaking bonds than released making them)

Method for Bond Energy Calculations

Step 1: Write the balanced chemical equation (using displayed formulae / structural formulae if helpful).

Step 2: List all the bonds in the reactants and add up their bond energies — this is the energy absorbed breaking bonds.

Step 3: List all the bonds in the products and add up their bond energies — this is the energy released making bonds.

Step 4: Apply the formula: ΔH = total energy to break bonds − total energy released making bonds.

Step 5: State the sign of ΔH and whether the reaction is exothermic or endothermic.

Example 1: H2 + Cl2 → 2HCl

Step 1: Write the equation with structural formulae.

H-H + Cl-Cl -> 2 × H-Cl

Step 2: Bonds broken (reactants):

Step 3: Bonds formed (products):

Step 4: Apply the formula.

ΔH = Σ(bonds broken) − Σ(bonds formed)
ΔH = 678 − 862
ΔH = −184 kJ/mol

Step 5: ΔH is negative, so the reaction is exothermic. 184 kJ of energy is released per mole of reaction (per mole of H2 reacted, or per 2 moles of HCl formed).

Example 2: Combustion of Methane (CH4)

Reaction: CH4 + 2O2 → CO2 + 2H2O

Step 1: Draw the structural formulae (showing all bonds).

    H
    |
H — C — H    +    2 × (O = O)    ->    O = C = O    +    2 × (H — O — H)
    |
    H

Step 2: Bonds broken (reactants):

Bond energy values used: C-H = 413, O=O = 498.

Step 3: Bonds formed (products):

Bond energy values used: C=O = 799, O-H = 464.

Step 4: Apply the formula.

ΔH = 2648 − 3454
ΔH = −806 kJ/mol

Step 5: ΔH is negative, so the combustion of methane is highly exothermic. This is why methane is used as a fuel.

Example 3: 2H2 + O2 → 2H2O

Reaction: 2H2 + O2 → 2H2O

Step 1: Structural formulae.

2 × (H-H) + (O=O) -> 2 × (H-O-H) = 4 × (O-H) bonds

Step 2: Bonds broken:

Step 3: Bonds formed:

Step 4:

ΔH = 1370 − 1856
ΔH = −486 kJ/mol

Step 5: The reaction is exothermic (ΔH negative).

Each H2O molecule has two O-H bonds. With 2H2O, that makes 4 O-H bonds in total. Missing this doubling is a common error.

Why Bond Energy Calculations Are Approximate

Bond energy calculations using average bond energies give approximate values, not exact values. There are two reasons for this:

1. Average bond energies: The bond energy values given in data tables are average values calculated from many different compounds. For example, the C-H bond energy of 413 kJ/mol is the average C-H bond energy across all molecules containing C-H bonds (methane, ethane, propane, etc.). In reality, a C-H bond in methane has a slightly different energy from a C-H bond in ethane.

2. Different molecular environments: The same type of bond can have different energies depending on what other atoms are attached nearby. For instance, a C-H bond next to a C=O group will have a slightly different bond energy from a C-H bond in an alkane.

Further examples:

Example 4: Simple Bond Energy Calculation Use the bond energy data to calculate ΔH for the reaction: H2 + Br2 → 2HBr. Bond energies (kJ/mol): H-H = 436, Br-Br = 193, H-Br = 366.

• Bonds broken: 1 × H-H (436) + 1 × Br-Br (193) = 629 kJ
• Bonds formed: 2 × H-Br (366) = 732 kJ
• ΔH = 629 − 732 = −103 kJ/mol (exothermic)

Example 5: Identify Exothermic or Endothermic A reaction has the following bond energies: Bonds broken: 1430 kJ; Bonds formed: 1180 kJ. Is the reaction exothermic or endothermic?

• ΔH = 1430 − 1180 = +250 kJ/mol
• The reaction is endothermic because ΔH is positive (more energy is absorbed breaking bonds than is released forming bonds)

Key Points

• Bond energy = energy required to break 1 mole of bonds in the gaseous state (kJ/mol)
• Bond breaking is endothermic (absorbs energy, positive value)
• Bond making is exothermic (releases energy, negative value)
• ΔH = Σ(bond energies of bonds broken) − Σ(bond energies of bonds formed)
• ΔH negative = exothermic; ΔH positive = endothermic
• Bond energy calculations give approximate values because:
  1. Average bond energies are used (not specific to that molecule)
  2. Bond energy varies with molecular environment
• Always draw structural formulae to accurately count each type of bond
• Count bonds carefully — each H2O has two O-H bonds; CO2 has two C=O bonds

Key Concepts from Past Papers

• Bond energy: the energy required to break one mole of a particular covalent bond in the gaseous state
• ΔH = bonds broken − bonds formed
• Bond energy values are average values / vary with molecular environment
• Exothermic because energy level of reactants above energy level of products
• Endothermic because heating is needed

Keywords from Past Papers

exothermic, because, energy, level, endothermic, heating, needed, reactants, products, arrow, going, downwards

Related Notes

• Energetics of Reactions
• Rates of Reaction
• IGCSE-Chem-Index

Sources

• OpenStax Chemistry 2e — [Chapter 7: Chemical Bonding and Molecular Geometry (Bond Energies)], Rice University (free, CC BY 4.0)
• BBC Bitesize GCSE Chemistry — [Bond Energy Calculations], BBC (free educational resource)
• Cambridge IGCSE Chemistry 0620 — Syllabus Section 5: Chemical Energetics (Bond Energies), Cambridge Assessment International Education
• CK-12 Chemistry for High School — [Chapter 17: Bond Enthalpy], CK-12 Foundation (free, CC BY-NC 3.0)

Past Paper Sources

• 0620/32 May/June 2018: Q88(d)(ii) (1m)
• 0620/33 October/November 2016: Q55(a)(i) (1m)
• 0971/32 May/June 2022: Q44(d)(i) (1m)

Common Misconceptions