Chemistry · Ch 4 — Chemical Bonding and Molecular Structure
Bond Order
Bond Order
Bond Order
The Lewis model gives us a simple, direct way to measure the strength of a covalent bond. This measure is called the bond order.
In the Lewis picture, the bond order is simply the number of bonds between two atoms in a molecule. Each bond corresponds to one shared pair of electrons.
Consider the simplest case: a hydrogen molecule, . The two hydrogen atoms share a single pair of electrons, forming one bond. The bond order is 1.
Now look at oxygen, . The two oxygen atoms share two pairs of electrons — a double bond. The bond order is 2.
For nitrogen, , three pairs of electrons are shared — a triple bond. The bond order is 3.
This pattern extends to heteronuclear diatomic molecules. In carbon monoxide, CO, the carbon and oxygen atoms share three pairs of electrons (a triple bond), so the bond order is also 3.
The bond order is always a positive integer in the Lewis model for molecules that obey the octet rule. Fractional bond orders appear in resonance structures (like ozone or the benzene ring), but that is a separate topic.
Bond Order and Bond Enthalpy
There is a direct relationship between bond order and the energy required to break a bond. The bond enthalpy (or bond dissociation energy) is the energy needed to break one mole of bonds in the gaseous state.
For , with a bond order of 3, the bond enthalpy is 946 kJ mol⁻¹. This is one of the highest values known for any diatomic molecule. A triple bond is exceptionally strong.
The general rule is straightforward: as bond order increases, bond enthalpy increases. More bonds between two atoms mean more electron density holding the nuclei together, so more energy is required to pull them apart.
Bond Order and Bond Length
Bond length is the distance between the nuclei of two bonded atoms. It follows the opposite trend to bond enthalpy.
A single bond (bond order 1) is the longest. A double bond (bond order 2) is shorter. A triple bond (bond order 3) is the shortest of the three.
The general rule: as bond order increases, bond length decreases. More bonds pull the nuclei closer together.
Do not memorise these trends in isolation. Understand the physical reason: a higher bond order means more shared electrons, which creates a stronger electrostatic attraction between the two positively charged nuclei and the shared negative charge. This stronger attraction pulls the nuclei closer (shorter bond) and requires more energy to separate them (higher bond enthalpy).
Isoelectronic Molecules and Ions
A powerful observation emerges when we compare species that have the same number of electrons — these are called isoelectronic species.
Isoelectronic molecules and ions have identical bond orders.
Consider the following examples: …