6.2 Bond Energy

This topic explains bond energy — what it is, how it relates to bond breaking and bond formation, and how it is used to calculate enthalpy changes.

Definition

Bond Dissociation Energy: The energy required when one mole of a particular bond breaks apart to form neutral atoms.

Bond Energy: The amount of energy released when one mole of neutral atoms forms a bond.

Main Idea

When a chemical reaction occurs, old bonds are broken and new bonds are formed. Breaking bonds always requires energy, and forming a bond always releases energy. The difference between the bond dissociation energy and the bond energy determines whether the reaction absorbs or releases energy. For any chemical reaction, the enthalpy change is the sum of the bond dissociation energies of the reactants minus the sum of the bond energies of the products.

Core Concept

Δ H = ∑ H_{(bonds broken)} - ∑ H_{(bonds formed)}

• Bond breaking: Energy absorbed (endothermic).
• Bond formation: Energy released (exothermic).
• Net energy change: Determines whether the reaction is exothermic or endothermic.

Detailed Explanation

What is Bond Energy?

• When a chemical reaction occurs, old bonds are broken and new bonds are formed.
• Breaking bonds always requires energy, and forming a bond always releases energy.
• The energy required when one mole of a particular bond breaks apart to form neutral atoms is called the bond dissociation energy.
• The amount of energy released when one mole of neutral atoms forms a bond is called bond energy.
• The difference between the bond dissociation energy and the bond energy determines whether the reaction absorbs or releases energy.

Enthalpy Change and Bond Energy

For any chemical reaction, the enthalpy change is the sum of the bond dissociation energies of the reactants minus the sum of the bond energies of the products.

Δ H = ∑ H_{(bonds broken)} - ∑ H_{(bonds formed)}

Example 6.3: Bond Energy of H−Cl

We can theoretically calculate the bond energy of a bond from known bond energy data. For example, we can calculate the bond energy of H−Cl as follows:

• Bond dissociation energy of H₂ = 436 kJ/mol
• Bond dissociation energy of Cl₂ = 243 kJ/mol

Required: - Bond energy of HCl

• Total energy absorbed in bond breaking = Total energy released in bond formation
• Bond dissociation energy of H−H + Bond dissociation energy of Cl−Cl = 2 × Bond energy of H−Cl

436 + 242 = 2(Bond energy of H−Cl)

679 = 2(Bond energy of H−Cl)

Bond energy of H−Cl = 339 kJ/mol

• Theoretically calculated bond energy of HCl = 339 kJ/mol
• Experimentally determined bond energy of HCl = 431.5 kJ/mol

• This is because the theoretical values are averages and do not take into account situations where the bond strength differs from the average.
• For example, unequal distribution of electron pairs due to differences in electronegativity values.
• This energy difference also appears as the enthalpy of reaction.

Example: Reaction Between H₂ and O₂

For the reaction:

2H_2(g) + O_2(g) → 2H₂O_(l)

Δ H = 2 × B.D.E of H₂ + B.D.E of O₂ - 4 × B.E of O−H

= 2(436 kJ) + 493.6 kJ - 4(460 kJ)

= 872 kJ + 493.6 kJ - 1840 kJ

= -474.4 kJ

Thus, the reaction between hydrogen and oxygen to form water is exothermic.

Calculation of Bond Energy for the Enthalpy of Reaction

You can calculate the bond energy for the enthalpy of a reaction. For example, you can calculate the bond energy per mole of methane gas by breaking the C−H bonds as follows:

CH_4(g) → 4H_(g) + C_(s) Δ H^° = +1662 kJ

In this reaction, 4 moles of C−H bonds are broken and 1662 kJ of energy is absorbed.

• The bond dissociation energy per mole of CH₄ is +1662 kJ.
• As 4 mol of C−H bonds are broken:

Intermediate Level

Bond Energy and Enthalpy Change

Δ H = ∑ H_{(bonds broken)} - ∑ H_{(bonds formed)}

Example 6.3: Bond Energy of H−Cl

• B.D.E of H₂ = 436 kJ/mol
• B.D.E of Cl₂ = 243 kJ/mol

Solution: 436 + 242 = 2(B.E of H−Cl) 679 = 2(B.E of H−Cl) B.E of H−Cl = 339 kJ/mol

Theoretical value: 339 kJ/mol Experimental value: 431.5 kJ/mol

Example: H₂ + O₂ Reaction

2H₂ + O₂ → 2H₂O \] \[ Δ H = 2(436) + 493.6 - 4(460) \] \[ = 872 + 493.6 - 1840 \] \[ = -474.4 kJ

Bond Energy of C−H

CH₄ → 4H + C Δ H^° = +1662 kJ \] \[ Average C−H bond energy = 16624 = +415.5 kJ

Important Points

• 1. Bond breaking: Energy absorbed.
• 2. Bond formation: Energy released.
• 3. ΔH = Σ(bonds broken) − Σ(bonds formed).
• 4. Example: H−Cl bond energy = 339 kJ/mol (theoretical), 431.5 kJ/mol (experimental).
• 5. Example: H₂ + O₂ → 2H₂O, ΔH = −474.4 kJ.
• 6. Average C−H bond energy = 415.5 kJ/mol.

Remember

• - Bond breaking: Energy absorbed
• - Bond formation: Energy released
• - ΔH = Σ(bonds broken) − Σ(bonds formed)
• - H−Cl: 339 kJ/mol (theoretical)
• - **H₂ + O₂:** ΔH = −474.4 kJ
• - C−H: 415.5 kJ/mol

Common Mistakes

• 1. Confusing bond dissociation energy with bond energy: Dissociation = breaking; Bond energy = forming.
• 2. Forgetting the formula: ΔH = Σ(bonds broken) − Σ(bonds formed).
• 3. Mixing up theoretical and experimental values: Theoretical values are averages.
• 4. Ignoring the sign of ΔH: Negative = exothermic; Positive = endothermic.

Quick Revision

• Bond breaking: Energy absorbed.
• Bond formation: Energy released.
• ΔH = Σ(bonds broken) − Σ(bonds formed).
• Example: H−Cl bond energy = 339 kJ/mol (theoretical), 431.5 kJ/mol (experimental).
• Example: H₂ + O₂ → 2H₂O, ΔH = −474.4 kJ.
• Average C−H bond energy = 415.5 kJ/mol.

60-Second Revision

• Bond breaking: Energy absorbed
• Bond formation: Energy released
• ΔH = Σ(bonds broken) − Σ(bonds formed)
• H−Cl: 339 kJ/mol (theoretical)
• **H₂ + O₂:** ΔH = −474.4 kJ
• C−H: 415.5 kJ/mol

QUALITY VALIDATION

• Topic: 6.2 Bond Energy
• Chapter: 6 (Energetics)

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