The very premise of a covalent bond is electron sharing. As two atoms share one electron each to form one covalent bond, they may likely share more than once and form more bonds. This information on the number of connections between two atoms is revealed from the Bond Order.
So, the Bond Order measures the number of bonds between the two atoms in a molecule. The number can be integers like 1, 2, or 3 for single, double, or triple bonds or non-integers like 0.5, 1.3, 1.5, etc.
Integer Bond Order
If a molecule has a bond order of 1, then the two atoms are joined by a single covalent bond.
For example, the bond order of the H-H bond in the H2 molecule is 1. The bond order of 2 and 3 means that the bond joining the two atoms is two and three- for example, O2 (O=O) and N2 (N ≡N) molecules.
Non-Integer Bond Order
In molecules having multiple bonds (above bond order 1), the second and third bond is formed by pie electrons. The delocalization of pie electrons such that they do not remain confined between two atoms but are distributed to more than two atoms (a phenomenon known as resonance) alters the bond order. So, in such instances, instead of having a integral bond order, the bond order is fractional.
For example, the delocalization of pie electrons in Benzene results in bonds losing their identity as single and double bonds. The single bonds are not entirely single, and the double bond is only partially double. So, there is no fixed bond order. The net bond order is 1.5 for one-and-a-half bonds.
Another example is nitrate anion (NO3-), where nitrogen forms single and double covalent bonds with oxygen atoms.
There are two N-O single bonds and one N=O double bond. So, the total bond order is 4 (1+1+2).
However, the pie electrons on the oxygen atoms and the double bond are not fixed but delocalized. The delocalization occurs over the three N-O bond groups (also known as the covalent bond pairs). So, the net bond order of the molecule or for each bond is 4/3 or 1.33.
Such molecules with partial pie bonds resulting from resonance or electron delocalization have fractional bond order between 1 and 2.
The bond order of dihydrogen cation H2+ is another interesting case. The number of electrons forming a covalent bond in dihydrogen cation H2+ is one, indicating a one-electron covalent bond. Since the total contribution for bond order, 1 is two electrons, the bond order of dihydrogen cation H2+ is 0.5, half of one.
How Bond Order Corresponds to the Bond Strength and Bond Length
Bond Order helps in understanding the relative stability of the bonds by comparing it to bonds with order 1. It means that a bond with a bond order of 2 is twice as stronger as a bond with a bond order of 1. Or a bond order 0.5 has half the strength of a single covalent bond.
So, the high bond order would naturally indicate higher stability, translating to more input of energy (bond energy) that would be required to break such stable bonds. Therefore, higher bond order would also indicate bonds having higher bond energy.
Such stable molecules with higher bond order and high bond energy would also mean that they are tightly bound, reducing the bond length.
The triple bonds with bond order two have the shortest bond length and highest bond energy than double and single bonds.
Molecule | Bond Order | Bond Energy (kJ/mole) | Bond Length (pm) |
|---|---|---|---|
C-C | 1 | 348 | 153.3 |
C=C | 2 | 614 | 133.9 |
C≡C | 3 | 839 | 120.3 |
N≡N (N2) | 3 | 946 | 110 |
C≡O | 3 | 1077 | 112.8 |
I-I (I2) | 1 | 151 | 267 |
In ethyne molecule H-C≡C-H, the bond order of the Carbon-Carbon bond is 3, whereas the Bond Order of the C-H bond is 1.
Nitrogen (N2), a homonuclear diatomic molecule, has a bond order of 3 (N≡N), a bond length of 110 pm, and a bond enthalpy of 946 kJ/mol, the highest for any homonuclear diatomic molecule.
I2 is also a diatomic molecule, but it is large-sized. It has a bond length of 267 pm and a bond order 1. Moreover, the two Iodine atoms have three lone pairs causing repulsion, destabilizing the covalent bond strength reflecting in the I2 molecule to have a bond enthalpy of only 151 kJ/mol.
Carbon monoxide (CO) is another example of a heteronuclear diatomic molecule having a bond order of 3 (C≡O) and a high bond enthalpy for any heteronuclear diatomic molecule of 1077 kJ/mol.
