Bond Enthalpy: The Energy Cost of Breaking a Bond
Think of a chemical bond as a spring holding two atoms together. To pull the atoms apart, you have to do work — you have to put energy in. That energy, per mole of bonds broken, is the bond enthalpy. The stronger the bond, the more energy you need to supply, and the larger the bond enthalpy.
The intuition is simple: breaking bonds costs energy; forming bonds releases energy. A reaction is exothermic if the energy released in forming new bonds is greater than the energy consumed in breaking old ones.
The Precise Definition
Bond enthalpy (symbol: ΔHbond or B.E.) is defined under a specific set of conditions:
The average enthalpy change when one mole of a particular covalent bond is broken, with all species in the gaseous state.
Three key points are packed into that sentence:
- One mole of bonds — we measure the energy for Avogadro's number of bonds, not just one.
- Gaseous state — the atoms or molecules must be in the gas phase. This removes complications from intermolecular forces (like hydrogen bonding in liquid water) and lets us isolate the bond energy itself.
- Average — in a molecule like water (H2O), the two O–H bonds are not identical in energy. The first O–H bond in H2O requires about 502 kJ/mol to break, but the second (in the remaining OH radical) requires about 427 kJ/mol. So we report the average bond enthalpy for O–H: roughly 464 kJ/mol.
Bond enthalpy is always positive — it is the energy absorbed to break a bond. Bond formation releases the same amount of energy (negative enthalpy change).
How It's Used: Estimating Reaction Enthalpies
You can estimate the enthalpy change of a reaction (ΔHrxn) using bond enthalpies:
ΔHrxn=∑(bond enthalpies of bonds broken)−∑(bond enthalpies of bonds formed)
The logic: you put energy in to break bonds (positive), and you get energy out when bonds form (negative). So:
ΔHrxn=Energy in−Energy out
This method gives an estimate, not an exact value. Bond enthalpies are averages taken from many different molecules, so they don't account for the exact molecular environment. For precise work, use standard enthalpies of formation.
Example: Combustion of Methane
Consider: CH4(g)+2O2(g)→CO2(g)+2H2O(g)
Bonds broken (energy in):
- 4 C–H bonds: 4×413=1652 kJ/mol
- 2 O=O bonds: 2×498=996 kJ/mol
- Total in: 1652+996=2648 kJ/mol
Bonds formed (energy out):
- 2 C=O bonds: 2×799=1598 kJ/mol
- 4 O–H bonds: 4×464=1856 kJ/mol
- Total out: 1598+1856=3454 kJ/mol
ΔHrxn=2648−3454=−806 kJ/mol
The negative sign tells you the reaction is exothermic — more energy is released in forming bonds than was consumed in breaking them. …