Bonding in Some Diatomic Molecules (Applications of Molecular Orbital Theory)

Bonding in Some Diatomic Molecules:

Hydrogen Molecule (H₂):

The electronic configuration of H is 1s¹. Thus the formation of the H₂ molecule involves the linear combination of 1s orbital of one H-atom with 1s-orbital of another H-atom forming two molecular orbitals. The molecular orbital formed by the addition overlap is known as bonding molecular orbital and is represented as σ1s while the other formed due to subtraction overlap is known as antibonding molecular orbital and is represented as σ*1s. Two 1s electrons one from each H-atom, pass into the σ1s orbital (bonding molecular orbital) because it is of lower energy while σ*1s orbital (antibonding molecular orbital) with higher energy remains vacant. Thus, the electronic configuration of the H₂ molecule is [σ(1s)]².

Thus, the two H-atoms in a molecule of H₂ are joined by a single covalent bond. It has no unpaired electron, so it is diamagnetic.

Hydrogen Molecule Ion (H₂⁺):

It is formed by the combination of H-atom containing one electron and hydrogen ion (H⁺) containing no electron. Thus, this molecule has only one electron which is present in the lowermost molecular orbital i.e. σ1s. So electronic configuration of H₂⁺ is [σ(1s)]¹.

Since bond order is positive, so the molecule is stable, though stability is not very high due to low value. The presence of one unpaired electron shows that it is paramagnetic.

Hydrogen Molecule Ion (H₂^–):

It is formed by the combination of H-atom containing molecular orbitals of H₂⁺ ions having one electron (in 1s orbital) and hydrogen ion (H^–) containing two electrons (in 1s orbital). Thus, the molecule has three electrons, two are present in bonding molecular orbital (σ1s being of lower energy) and one in antibonding molecular orbital (σ*1s). Thus the electronic configuration of H₂^– ion is [σ(1s)]² [σ*(1s)]¹.

Like H₂⁺ ion, its bond order is also +1/2 but this ion, H₂^–, is less stable than H₂⁺ ion due to the presence of one electron in the antibonding orbital which results in repulsion and decreases the stability. Due to the presence of one unpaired electron, it is paramagnetic.

Helium Molecule (He₂):

The electronic configuration of He is 1s² indicating that in the He₂ molecule there will be 4 electrons. Two of these electrons are present in σ1s and two in σ*1s molecular orbital. Therefore, the electronic configuration of the He₂ molecule will be [σ(1s)]² [σ*(1s)]².

Since the bond order of the He₂ molecule is zero, therefore, this molecule is unstable and does not exist.

Nitrogen Molecule (N₂):

The electronic configuration of N is 1s², 2s² 2p_x¹ 2p_y¹ 2p_z¹ which shows that the N₂ molecule has 14 electrons which are to be accommodated in orbitals in the increasing order of their energy. The molecular orbital configuration of the N₂ molecule is [σ(1s)]² [σ*(1s)]² [σ(2s)]² [σ*(2s)]² [π(2p_x)]² [π(2p_y)]² [σ(2p_z)]².

The structure [σ(1s)]² [σ*(1s)]² indicated a closed K shell structure which does not enter into bonding and is usually written as KK. The electronic configuration of the N₂ molecule may also be written as KK [σ(2s)]² [σ*(2s)]² [π(2p_x)]² [π(2p_y)]² [σ(2p_z)]².

The two N-atoms in a molecule of N₂ are joined by a triple covalent bond (N≡N). It has no unpaired electron and, hence, is a diamagnetic molecule.

Oxygen Molecule (O₂):

The electronic configuration of O is 1s², 2s² 2p_x² 2p_y¹ 2p_z¹ indicating that it has 8 electrons so that the molecule of O₂ has 16 electrons which are to be accommodated in various molecular orbitals in the increasing order of their energy. The molecular orbital configuration of O₂ is [σ(1s)]² [σ*(1s)]² [σ(2s)]² [σ*(2s)]² [σ(2p_z)]² [π(2p_x)]² [π(2p_y)]² [π*(2p_x)]¹ [π*(2p_y)]¹.

Since 1s electrons are not involved in bonding and so the molecular orbital configuration may also be written as KK [σ(2s)]² [σ*(2s)]² [σ(2p_z)]² [π(2p_x)]² [π(2p_y)]² [π*(2p_x)]¹ [π*(2p_y)]¹.

Thus, the two O-atom in a molecule of O₂ is joined by a double bond (O=O). Since it contains two unpaired electrons (in π*2p_x and π*2p_y orbitals) so it is paramagnetic in nature.