In general chemistry courses, students learn that covalent bonds can come as either single, double or triple bonds, which are identifies by their bond order. Both bond length and bond energy changes as the bond order increases and as the number of electrons shared between two atoms in a molecule increases, the bond order of a bond increases, the strength of the bond increases and the distance between nuclei decreases (Table \(\PageIndex{1}\)).
Table \(\PageIndex{1}\): General Correlation between Bond Strength, length and order in Covalent bonds
Bond
Bond Order
Bond Enthalpy (kJ/mol)
Bond Length (Å)
\(\ce{C-C}\)
1
348
1.54
\(\ce{C=C}\)
2
614
1.34
\(\ce{C#C}\)
3
839
1.20
\(\ce{N-N}\)
1
163
1.47
\(\ce{N=N}\)
2
418
1.24
\(\ce{N#N}\)
3
941
1.10
The above trend can be observed in the first row diatomics in Figure \(\PageIndex{1}\). The bond order can be determined directly form the molecular orbital electron configurations.
\[\text{bond order} = \dfrac{\text{number of electrons in bonding MOs}- \text{number of electrons in antbonding MOs}}{2} \label{BO}\]
For diatomics, the occupations can correlate to bond length, bond energies (Figure \(\PageIndex{1}\)) and behavior in applied magnetic fields.
Figure \(\PageIndex{1}\): Plot of bond length (left) and bond energy (right) for first row diatomics.
Molecular Oxygen is Paramagnetic
We now turn to a molecular orbital description of the bonding in \(\ce{O2}\). It so happens that the molecular orbital description of this molecule provided an explanation for a long-standing puzzle that could not be explained using other bonding models. None of the other bonding models (e.g., Valence Bond theory or Lewis bonding) can predict the presence of two unpaired electrons in \(\ce{O_2}\). Chemists had long wondered why, unlike most other substances, liquid \(\ce{O_2}\) is attracted into a magnetic field. As shown in Figure \(\PageIndex{2}\), it actually remains suspended between the poles of a magnet until the liquid boils away. The only way to explain this behavior was for \(\ce{O_2}\) to have unpaired electrons, making it paramagnetic. This result was one of the earliest triumphs of molecular orbital theory over the other bonding approaches.
Figure\(\PageIndex{2}\): Liquid O2 Suspended between the Poles of a Magnet.Because the O2 molecule has two unpaired electrons, it is paramagnetic. Consequently, it is attracted into a magnetic field, which allows it to remain suspended between the poles of a powerful magnet until it evaporates.