8.4 Bond Energy and Bond Dissociation Energy

This topic explains bond energy and bond dissociation energy — the energies involved in breaking and forming chemical bonds.

Definition

Bond Dissociation Energy

The amount of energy required to break one mole of a particular bond to form neutral atoms is called the bond dissociation energy.

Simple explanation: Bond dissociation energy is the energy you need to put in to break a bond.

Bond Energy

The amount of energy released when neutral atoms form one mole of a bond is called bond energy.

Simple explanation: Bond energy is the energy you get out when a bond forms.

Main Idea

• Old bonds are broken → energy is absorbed.
• New bonds are formed → energy is released.

The difference between the bond dissociation energy and the bond energy determines whether the reaction absorbs or releases energy.

Simple explanation: Breaking bonds costs energy; making bonds pays energy back. The net result decides whether the reaction is exothermic or endothermic.

Core Concept

Key Relationship

The enthalpy change in a chemical reaction is the sum of energies absorbed and released in bond breaking and bond forming.

Δ H^° = Sum of bond dissociation energies of reactants - Sum of bond energies of products

Simple explanation: ΔH = (energy to break bonds in reactants) − (energy released when bonds form in products).

Interpreting the Result

Detailed Explanation

Bond Dissociation Energy — In Detail

Definition: The amount of energy required to break one mole of a particular bond to form neutral atoms.

Example: Breaking the H–H bond in H₂:

H₂(g) → 2H(g) Δ H = +436 kJ/mol

• 436 kJ/mol is the bond dissociation energy of H–H.
• The positive sign means energy is absorbed.

Simple explanation: To break one mole of H–H bonds, you must put in 436 kJ of energy.

Bond Energy — In Detail

Definition: The amount of energy released when neutral atoms form one mole of a bond.

Example: Forming the H–H bond:

2H(g) → H₂(g) Δ H = -436 kJ/mol

• 436 kJ/mol is the bond energy of H–H.
• The negative sign means energy is released.

Simple explanation: When two H atoms join to form H₂, 436 kJ of energy is released.

Bond Energy vs Bond Dissociation Energy

Simple explanation: They are the same magnitude but opposite in sign — one is energy in, the other is energy out.

Calculating Enthalpy Change from Bond Energies

Steps (from the book):

Step 1: Write the balanced chemical equation.

Step 2: Show all the reactants and products in the gaseous state.

Step 3: Substitute the relevant bond energy values in the formula and solve:

Δ H^° = Sum of bond dissociation energies of reactants - Sum of bond energies of products

Example: Calculate the enthalpy of the following reaction from the given bond energy data.

Reaction:

H₂(g) + F₂(g) → 2HF(g)

• Bond energy of H–H = **436 kJ/mol**
• Bond energy of F–F = **155 kJ/mol**
• Bond energy of H–F = **567 kJ/mol**

Solution:

Step 1: Balanced equation:

H₂(g) + F₂(g) → 2HF(g)

• 1 mole H–H = 436 kJ
• 1 mole F–F = 155 kJ
• Total energy absorbed = 436 + 155 = 591 kJ

• 2 moles H–F = 2 × 567 = **1134 kJ**
• Total energy released = 1134 kJ

Step 4: Apply the formula:

Δ H^° = 591 - 1134 = -543 kJ

Final Answer: ΔH° = −543 kJ → the reaction is exothermic.

Simple explanation: More energy is released forming H–F bonds than is absorbed breaking H–H and F–F bonds, so the reaction releases energy.

Note from the Book

> "Note that the enthalpy of reaction calculated using bond energy data are often different from values determined experimentally."

Simple explanation: Bond energy calculations give approximate values; experimental values may differ slightly because bond energies are averages.

Concept Assessment Exercise 8.2 (from the book)

Example: Calculate the enthalpy of the following reaction from the given bond energy data.

Reaction:

H₂(g) + Cl₂(g) → 2HCl(g)

• H–H = **436 kJ/mol**
• Cl–Cl = **243 kJ/mol**
• H–Cl = **432 kJ/mol**

Solution:

• H–H = 436 kJ
• Cl–Cl = 243 kJ
• Total = **679 kJ**

Bonds formed (products): - 2 × H–Cl = 2 × 432 = 864 kJ

**ΔH° = 679 − 864 = −185 kJ**

Final Answer: ΔH° = −185 kJ → exothermic.

Intermediate Level

Why Bond Energies Are Averages

• The exact energy to break a bond depends on the molecular environment.
• For example, the C–H bond in CH₄ is slightly different from the C–H bond in C₂H₆.
• So we use average bond energies for calculations.

Simple explanation: Bond energy tables give average values; actual values may vary slightly.

Using Bond Energies to Calculate ΔH

General steps:

• Write the balanced equation.
• Draw the structures to identify which bonds are broken and formed.
• Sum the bond energies of all bonds broken (reactants).
• Sum the bond energies of all bonds formed (products).
• Apply:

Example: H₂ + F₂ → 2HF

Δ H = 591 - 1134 = -543 kJ

Exothermic (negative ΔH).

Why the Sign Matters

• Negative ΔH → more energy released than absorbed → exothermic.
• Positive ΔH → more energy absorbed than released → endothermic.

Common Bond Energies

Important Points

• - Bond dissociation energy: energy required to break one mole of a bond to form neutral atoms.
• - Bond energy: energy released when neutral atoms form one mole of a bond.
• - Breaking bonds: energy absorbed.
• - Forming bonds: energy released.
• - Formula:
• Δ H^° = Sum of bond dissociation energies of reactants - Sum of bond energies of products
• - Negative ΔH → exothermic; positive ΔH → endothermic.
• - Bond energy values are averages — calculated values may differ slightly from experimental values.
• - **Example (H₂ + F₂ → 2HF):** ΔH° = **−543 kJ** (exothermic).
• - **Example (H₂ + Cl₂ → 2HCl):** ΔH° = **−185 kJ** (exothermic).

Remember

• - Bond dissociation energy = energy to break a bond.
• - Bond energy = energy released when a bond forms.
• - Breaking bonds → energy absorbed.
• - Forming bonds → energy released.
• - ΔH = bonds broken − bonds formed.
• - Negative ΔH → exothermic.
• - Positive ΔH → endothermic.
• - **Example:** H₂ + F₂ → 2HF, ΔH = −543 kJ.

Common Mistakes

• 1. Confusing bond energy and bond dissociation energy. Bond dissociation = breaking (absorbed); bond energy = forming (released).
• 2. Forgetting the sign convention. Breaking bonds is positive (energy in); forming bonds is negative (energy out).
• 3. Forgetting to multiply by coefficients. For 2HF, multiply H–F bond energy by 2.
• 4. Thinking bond energies are exact values. They are average values.
• 5. Subtracting in the wrong order. Use: ΔH = (bonds broken) − (bonds formed).
• 6. Forgetting to show gaseous states. Bond energies apply to gaseous species.

Quick Revision

• Bond dissociation energy: the amount of energy required to break one mole of a particular bond to form neutral atoms.
• Bond energy: the amount of energy released when neutral atoms form one mole of a bond.
• Enthalpy change formula:
• Breaking bonds: energy absorbed (positive).
• Forming bonds: energy released (negative).
• Example (H₂ + F₂ → 2HF):
• Bonds broken: H–H (436) + F–F (155) = 591 kJ
• Bonds formed: 2 × H–F (567) = 1134 kJ
• ΔH° = 591 − 1134 = −543 kJ (exothermic)
• Example (H₂ + Cl₂ → 2HCl):
• Bonds broken: H–H (436) + Cl–Cl (243) = 679 kJ
• Bonds formed: 2 × H–Cl (432) = 864 kJ
• ΔH° = 679 − 864 = −185 kJ (exothermic)
• Bond energy values are averages — calculated values may differ slightly from experimental values.

60-Second Revision

• Bond dissociation energy = break a bond → absorbed.
• Bond energy = form a bond → released.
• ΔH = bonds broken − bonds formed.
• Negative ΔH → exothermic.
• Positive ΔH → endothermic.
• **Example:** H₂ + F₂ → 2HF, ΔH = −543 kJ.
• Warning: breaking bonds is positive (energy in); forming bonds is negative (energy out).