By reference to Molecular orbital Theory and giving the appropriate MO diagrams explain why the bond in N2 is very strong short (bond dissociation entalpy 945 kjmol-1,bond length 109.8 pm) while that in O2 is weaker and longer (bond dissociation entalpy 497 kJ mol-1,bond length 120.8 pm)
To explain the difference in bond strength and length between N2 and O2 using the molecular orbital (MO) theory, we need to consider the electronic structure and bonding in these molecules.
In MO theory, we use a combination of atomic orbitals (AOs) to form molecular orbitals (MOs), which are regions of space where electrons are likely to be found. The number of MOs formed is equal to the number of AOs used.
Let's start with N2:
1. N2 molecule: N≡N
- Each nitrogen atom has five valence electrons: 2s²2p³.
- These valence electrons will occupy MOs formed from the combination of atomic orbitals.
2. Formation of molecular orbitals in N2:
- Two 2s orbitals combine to form two σ MOs:
2s + 2s → σ2s, σ*2s
- Two 2p orbitals combine to form two σ MOs and two π MOs:
2pz + 2pz → σ2pz, σ*2pz
2px + 2px → π2px, π*2px
2py + 2py → π2py, π*2py
3. Filling of molecular orbitals in N2:
- The electrons fill the molecular orbitals in order of increasing energy, following the aufbau principle and Pauli exclusion principle.
- The molecular orbitals are filled with a maximum of two electrons with opposite spins.
4. Bonding and antibonding molecular orbitals in N2:
- The σ2s and σ2pz orbitals are lower in energy and are bonding orbitals, stabilizing the molecule.
- The π2px, π2py, π*2px, and π*2py orbitals are higher in energy and are antibonding orbitals, destabilizing the molecule.
- The antibonding π* orbitals are not involved in the bonding and do not contribute significantly to the stability of the molecule.
Now let's move to O2:
1. O2 molecule: O=O
- Each oxygen atom has six valence electrons: 2s²2p⁴.
2. Formation of molecular orbitals in O2:
- Two 2s orbitals combine to form two σ MOs:
2s + 2s → σ2s, σ*2s
- Two 2p orbitals combine to form two σ MOs and two π MOs:
2pz + 2pz → σ2pz, σ*2pz
2px + 2px → π2px, π*2px
2py + 2py → π2py, π*2py
3. Filling of molecular orbitals in O2:
- The electrons fill the molecular orbitals following the aufbau principle and Pauli exclusion principle.
4. Bonding and antibonding molecular orbitals in O2:
- The σ2s, σ2pz, and π2pz orbitals are lower in energy and are bonding orbitals, stabilizing the molecule.
- The π2px, π2py, π*2px, and π*2py orbitals are higher in energy and are antibonding orbitals, destabilizing the molecule.
- The antibonding π* orbitals contribute to the destabilization of the molecule.
Now let's analyze the differences between N2 and O2 based on their MO diagrams:
1. Bond strength:
- In N2, all the bonding molecular orbitals are filled, resulting in a strong bond.
- In O2, there are two antibonding π* orbitals with electron occupation, which weakens the overall bond strength.
2. Bond length:
- In N2, the absence of electrons in the antibonding π* orbitals keeps the bond length shorter.
- In O2, the presence of electrons in the antibonding π* orbitals causes electron-electron repulsion, resulting in a longer bond length.
Therefore, the combination of a stronger bond and shorter bond length in N2 compared to O2 can be attributed to the electronic structure and occupation of molecular orbitals as explained by MO theory.