Reaction Profiles and Bond Energies

This lesson covers reaction profile diagrams in detail, the concept of bond energy, and how to calculate the overall energy change of a reaction using bond energies as required by the Edexcel GCSE Combined Science specification (1SC0). Bond energy calculations are a Higher tier topic but understanding reaction profiles is needed by all students.

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Reaction Profile Diagrams

A reaction profile (also called an energy level diagram) is a graph that shows how the energy of the reacting system changes during the course of a reaction.

What to Include on a Reaction Profile

Label	What it represents
Reactants	The starting energy level (labelled on the left)
Products	The final energy level (labelled on the right)
Activation energy (Eₐ)	The energy difference between the reactants and the peak of the curve
Overall energy change (ΔH)	The energy difference between reactants and products
Progress of reaction	The x-axis label
Energy	The y-axis label

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Exothermic Reaction Profile

graph LR
    subgraph "Exothermic Profile"
        direction TB
        R["Reactants<br/>———————<br/>(high energy)"]
        T["Peak (Eₐ)"]
        P["Products<br/>———————<br/>(low energy)"]
    end

Feature	Value
ΔH	Negative (energy released)
Products	Lower than reactants
Eₐ	Measured from reactants up to the peak

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Endothermic Reaction Profile

graph LR
    subgraph "Endothermic Profile"
        direction TB
        R2["Reactants<br/>———————<br/>(low energy)"]
        T2["Peak (Eₐ)"]
        P2["Products<br/>———————<br/>(high energy)"]
    end

Feature	Value
ΔH	Positive (energy absorbed)
Products	Higher than reactants
Eₐ	Measured from reactants up to the peak

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Effect of a Catalyst on the Reaction Profile

A catalyst provides an alternative pathway with a lower activation energy. On a reaction profile, this is shown by a curve with a lower peak:

Feature	Without catalyst	With catalyst
Eₐ	Higher	Lower
ΔH	Unchanged	Unchanged
Reactant level	Unchanged	Unchanged
Product level	Unchanged	Unchanged

The catalyst only changes the height of the activation energy barrier — everything else stays the same.

Exam Tip: When drawing a catalysed reaction on a profile, draw a second, lower curve between the same reactant and product levels. Label both Eₐ values clearly.

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Bond Energy

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

Key Principle

• Breaking bonds requires energy (endothermic process).
• Making bonds releases energy (exothermic process).

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

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Bond Energy Values

You will be given bond energy values in the exam. Here are some common ones:

Bond	Bond energy (kJ/mol)
H–H	436
O=O	498
O–H	464
C–H	413
C=O	805
C–C	347
C=C	614
C–O	358
N≡N	945
N–H	391
Cl–Cl	243
H–Cl	432

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Calculating Overall Energy Change from Bond Energies (Higher)

The Formula

$$\text{Overall energy change} = \text{Energy to break bonds (reactants)} - \text{Energy released making bonds (products)}$$Overall energy change=Energy to break bonds (reactants)−Energy released making bonds (products)

Or equivalently:

$$\Delta H = \Sigma \text{(bonds broken)} - \Sigma \text{(bonds made)}$$ΔH=Σ(bonds broken)−Σ(bonds made)

Result	Meaning
Positive ΔH	More energy needed to break bonds than released making them → endothermic
Negative ΔH	More energy released making bonds than needed to break them → exothermic

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Worked Example 1: Combustion of Hydrogen

Equation: 2H₂ + O₂ → 2H₂O

Step 1: Identify all bonds broken (in reactants)

Bond	Number	Energy per bond (kJ/mol)	Total (kJ)
H–H	2	436	872
O=O	1	498	498
		Total energy in	1370

Step 2: Identify all bonds made (in products)

Bond	Number	Energy per bond (kJ/mol)	Total (kJ)
O–H	4 (two in each H₂O, × 2 molecules)	464	1856
		Total energy out	1856

Step 3: Calculate ΔH

$$\Delta H = 1370 - 1856 = -486 \text{ kJ/mol}$$ΔH=1370−1856=−486 kJ/mol

The answer is negative, so the reaction is exothermic. This makes sense — burning hydrogen is a combustion reaction that releases heat.

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Worked Example 2: Combustion of Methane

Equation: CH₄ + 2O₂ → CO₂ + 2H₂O

Step 1: Bonds broken

Bond	Number	Energy (kJ/mol)	Total (kJ)
C–H	4	413	1652
O=O	2	498	996
		Total	2648

Step 2: Bonds made

Bond	Number	Energy (kJ/mol)	Total (kJ)
C=O	2	805	1610
O–H	4	464	1856
		Total	3466

Step 3: Calculate

$$\Delta H = 2648 - 3466 = -818 \text{ kJ/mol}$$ΔH=2648−3466=−818 kJ/mol

Again negative — the combustion of methane is exothermic.

Exam Tip: Draw out the structural formulas of all reactants and products to count bonds accurately. For example, CO₂ is O=C=O (two C=O bonds) and H₂O is H–O–H (two O–H bonds). Mistakes in counting bonds are the most common error.

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Worked Example 3: Breaking Down Nitrogen (Endothermic)

Equation: N₂ + 3H₂ → 2NH₃

Bonds broken

Bond	Number	Energy (kJ/mol)	Total (kJ)
N≡N	1	945	945
H–H	3	436	1308
		Total	2253

Bonds made

Bond	Number	Energy (kJ/mol)	Total (kJ)
N–H	6 (3 in each NH₃ × 2)	391	2346
		Total	2346

Calculate

$$\Delta H = 2253 - 2346 = -93 \text{ kJ/mol}$$ΔH=2253−2346=−93 kJ/mol

This is actually exothermic (negative ΔH), which is correct — the Haber process is exothermic in the forward direction.

Exam Tip: Bond energy calculations give approximate values because the values used are averages across many different molecules. The actual energy change may differ slightly from the calculated value.

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Common Mistakes