Bond Energy Calculations (Higher)

This lesson covers bond energy calculations, which is a Higher tier topic in the Edexcel GCSE Chemistry (1CH0) specification. You will learn why bond breaking is endothermic and bond making is exothermic, how to use bond energy data to calculate the overall energy change of a reaction, and how to determine whether a reaction is exothermic or endothermic from the calculation. This is a very commonly examined topic at Higher tier.

Note: This topic is assessed at Higher tier only. Foundation tier students do not need to carry out bond energy calculations, but understanding the general principles of bond breaking and bond making is useful for all students.

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Bond Breaking and Bond Making

During a chemical reaction:

1. Bonds in the reactants must be broken — this requires energy to be put in (it is an endothermic process).
2. New bonds in the products are formed — this releases energy (it is an exothermic process).

The overall energy change of a reaction depends on the balance between these two processes.

Process	Energy change	Type
Breaking bonds	Energy is taken in	Endothermic
Making bonds	Energy is released	Exothermic

Exam tip: A very common mistake is to say that breaking bonds releases energy. It does not. Breaking bonds always requires energy. Only forming new bonds releases energy. This is one of the most important facts in the energy changes topic.

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What Is Bond Energy?

The bond energy (also called bond enthalpy) is the amount of energy needed to break one mole of a particular covalent bond in a gaseous molecule. It is measured in kJ/mol.

• A high bond energy means the bond is strong and requires a lot of energy to break.
• A low bond energy means the bond is weak and requires less energy to break.
• The same amount of energy is released when the same type of bond is formed.

Common Bond Energies

Bond	Bond energy (kJ/mol)
H—H	436
O=O	498
O—H	463
C—H	413
C=O (in CO₂)	805
C—C	347
C=C	614
C—O	358
N≡N	945
N—H	391
Cl—Cl	243
H—Cl	432
C—Cl	346

Exam tip: You do not need to memorise bond energies — they will be given to you in the exam. However, you must know how to use them in calculations. Practise with as many examples as possible.

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The Overall Energy Change Formula

The overall energy change of a reaction is calculated as:

Overall energy change = total energy to break bonds in reactants − total energy released forming bonds in products

Or more concisely:

ΔH = energy in (breaking bonds) − energy out (making bonds)

Interpreting the Result

Result	Meaning	Reaction type
Energy in > energy out → ΔH is positive	More energy is needed to break bonds than is released by making bonds	Endothermic
Energy in < energy out → ΔH is negative	Less energy is needed to break bonds than is released by making bonds	Exothermic

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Step-by-Step Calculation Method

Follow these steps for every bond energy calculation:

flowchart TD
    A[Step 1: Write out the balanced equation with structural formulas] --> B[Step 2: List ALL bonds broken in the REACTANTS]
    B --> C[Step 3: Calculate the TOTAL energy to break all reactant bonds]
    C --> D[Step 4: List ALL bonds formed in the PRODUCTS]
    D --> E[Step 5: Calculate the TOTAL energy released forming all product bonds]
    E --> F[Step 6: Calculate ΔH = energy in - energy out]
    F --> G{Is ΔH positive or negative?}
    G -->|Positive| H[Endothermic reaction]
    G -->|Negative| I[Exothermic reaction]
    style H fill:#99ccff,stroke:#336699
    style I fill:#ff9999,stroke:#cc0000

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Worked Example 1: Combustion of Hydrogen

Equation: 2H₂ + O₂ → 2H₂O

Step 1: Write the structural formulas:

• Reactants: H—H (×2) and O=O (×1)
• Products: H—O—H (×2), which means O—H bonds (×4)

Step 2 & 3: Calculate energy to break bonds in reactants:

Bond	Number	Bond energy (kJ/mol)	Total (kJ)
H—H	2	436	2 × 436 = 872
O=O	1	498	1 × 498 = 498
		Total energy in	1370

Step 4 & 5: Calculate energy released forming bonds in products:

Bond	Number	Bond energy (kJ/mol)	Total (kJ)
O—H	4	463	4 × 463 = 1852
		Total energy out	1852

Step 6: Calculate overall energy change:

ΔH = energy in − energy out

ΔH = 1370 − 1852

ΔH = −482 kJ/mol

Conclusion: ΔH is negative, so the reaction is exothermic. This makes sense because the combustion of hydrogen is a highly exothermic reaction.

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Worked Example 2: Combustion of Methane

Equation: CH₄ + 2O₂ → CO₂ + 2H₂O

Step 1: Structural formulas:

• Reactants: CH₄ has 4 × C—H bonds; O₂ has O=O bonds (×2)
• Products: CO₂ has 2 × C=O bonds; H₂O has 2 × O—H bonds (×2 molecules = 4 × O—H)

Step 2 & 3: Energy to break bonds in reactants:

Bond	Number	Bond energy (kJ/mol)	Total (kJ)
C—H	4	413	4 × 413 = 1652
O=O	2	498	2 × 498 = 996
		Total energy in	2648

Step 4 & 5: Energy released forming bonds in products:

Bond	Number	Bond energy (kJ/mol)	Total (kJ)
C=O	2	805	2 × 805 = 1610
O—H	4	463	4 × 463 = 1852
		Total energy out	3462

Step 6:

ΔH = 2648 − 3462

ΔH = −814 kJ/mol

Conclusion: ΔH is negative, so combustion of methane is exothermic. This is expected — burning methane releases a large amount of energy, which is why natural gas is used as a fuel.

Exam tip: The calculated value (−814 kJ/mol) differs slightly from the true value (−890 kJ/mol) because bond energies are average values measured across many different compounds. The actual bond energy in a specific molecule may differ slightly from the average. You do not need to worry about this discrepancy in the exam — just use the values given.

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