Bond enthalpy (also called bond energy or bond dissociation enthalpy) is a fundamental concept in IB Chemistry Topic 5. It helps you estimate the enthalpy change of reactions by examining the strengths of the chemical bonds broken and formed. Understanding bond enthalpy allows you to analyze stability, compare molecules, and solve energetic calculations quickly and confidently.
What Is Bond Enthalpy?
Bond enthalpy is the energy required to break one mole of a specific bond in the gaseous state.
Key points:
- It refers to breaking a bond → always endothermic
- Units: kJ mol⁻¹
- Applied only to gaseous molecules
- Indicates bond strength
A larger bond enthalpy means a stronger bond.
Example:
- O=O (double bond) has a higher bond enthalpy than O–O (single bond)
- C–H bonds are relatively strong
- C–C single bonds are weaker than C=C double bonds
Bond enthalpy provides insight into molecular stability.
Why Bond Enthalpy Is Always Endothermic
Breaking a bond requires energy because the bonded atoms are more stable together.
Energy must be supplied to overcome the attraction between nuclei and shared electrons.
Therefore:
Bond breaking = endothermic (positive ΔH)
Bond forming = exothermic (negative ΔH)
This is the basis of bond enthalpy calculations.
Average Bond Enthalpy
In many molecules, a particular bond does not have the exact same strength in every environment.
For example, the O–H bonds in water differ slightly from those in ethanol.
To simplify, IB Chemistry uses average bond enthalpies, which are:
- Averaged over similar molecules
- Sufficiently accurate for estimation
- Found in the IB data booklet
This is why answers from bond enthalpy calculations are approximate.
Using Bond Enthalpy to Calculate Enthalpy Change
The key formula used in IB is:
ΔH = Σ(bonds broken) – Σ(bonds formed)
Step-by-step:
1. Break bonds in the reactants
- This requires energy
- Add all bond enthalpy values together
2. Form new bonds in the products
- This releases energy
- Subtract these values
3. Use data booklet values
- Only use average bond enthalpies provided
Worked Example (IB-Style)
Calculate ΔH for the reaction:
H₂ + Cl₂ → 2HCl
Bond enthalpies (from data booklet):
- H–H = 436 kJ mol⁻¹
- Cl–Cl = 242 kJ mol⁻¹
- H–Cl = 431 kJ mol⁻¹
1. Bonds broken
H–H + Cl–Cl
= 436 + 242
= 678 kJ
2. Bonds formed
Two H–Cl bonds
= 2 × 431
= 862 kJ
3. Apply formula
ΔH = 678 – 862
= –184 kJ mol⁻¹
This negative value shows the reaction is exothermic.
When Bond Enthalpies Are Useful
Bond enthalpy calculations are useful for:
- Estimating enthalpy changes
- Checking whether a reaction is endothermic or exothermic
- Predicting relative stability of molecules
- Understanding combustion and formation reactions
- Comparing bond strengths (e.g., C–C vs C=C)
IB exam questions often include reaction mechanisms, where knowing which bonds break and form helps you understand energy profiles.
Factors That Influence Bond Enthalpy
Bond enthalpy depends on:
1. Bond Order
Higher bond order → stronger bond → higher bond enthalpy
Example:
C≡C > C=C > C–C
2. Bond Length
Shorter bonds are stronger
Example:
H–F > H–Cl > H–Br
3. Electronegativity Differences
More polarized bonds often have higher bond enthalpies.
4. Molecular Environment
Resonance and molecular geometry affect bond strength
(e.g., delocalization lowers bond enthalpy).
Limitations of Bond Enthalpy Calculations
- Not exact—values are averages
- Only applies to gas-phase reactions
- Cannot be used for reactions involving ionic bonds or lattice energies
- Less reliable when resonance structures are present (e.g., benzene)
Despite limitations, they remain extremely helpful for estimation and conceptual understanding.
FAQs
Why must all substances be in the gas phase for bond enthalpy?
Bond enthalpy is defined for gaseous molecules because it isolates the energy required to break a specific bond without additional interactions from solvents or crystal lattices.
Why are calculations approximate?
Because average bond enthalpies are used, not exact values for the specific molecule.
Why do triple bonds have higher bond enthalpy?
Triple bonds involve more overlapping electron density, making them shorter and stronger than double or single bonds.
Conclusion
Bond enthalpy measures the energy needed to break covalent bonds in gaseous molecules. It allows you to estimate enthalpy changes by comparing the energy required to break bonds with the energy released when new bonds form. Though approximate, this method is invaluable for understanding reaction energetics, stability, and molecular structure in IB Chemistry.
