Short Answer

O₂ is paramagnetic because it contains two unpaired electrons in its molecular orbitals. These unpaired electrons create magnetic moments that cause the molecule to be attracted toward a magnetic field. This behavior cannot be explained by Lewis structures but is clearly shown by molecular orbital (MO) theory.

According to MO theory, the electrons in O₂ fill bonding and antibonding orbitals in a specific order. The π* antibonding orbitals contain two unpaired electrons, and this is the direct reason for the paramagnetic nature of oxygen gas.

Detailed Explanation :

Why O₂ Is Paramagnetic

Oxygen (O₂) is one of the most important diatomic molecules in nature, and its magnetic behavior provides a clear demonstration of the power of molecular orbital (MO) theory. While Lewis structures incorrectly suggest that all the electrons in O₂ are paired, experiments show that oxygen is attracted to a magnetic field. This property is known as paramagnetism. MO theory explains this behavior by showing that O₂ actually contains unpaired electrons in its antibonding orbitals.

To understand why O₂ is paramagnetic, we need to look at how electrons fill molecular orbitals, how bonding and antibonding orbitals are arranged, and how unpaired electrons influence magnetic behavior.

1. Electron Configuration of Individual Oxygen Atoms

Each oxygen atom has:

• Atomic number = 8
• Electronic configuration = 1s² 2s² 2p⁴

When two oxygen atoms combine, their atomic orbitals mix to form molecular orbitals. The electrons from both atoms must then fill these molecular orbitals in order of increasing energy.

2. Molecular Orbital Arrangement in O₂

The MO energy level diagram for O₂ includes the following orbitals (from low to high energy):

1. σ2s (bonding)
2. σ2s* (antibonding)
3. σ2p (bonding)
4. π2p (bonding)
5. π2p* (antibonding)
6. σ2p* (antibonding)

O₂ has a total of 16 electrons (8 from each oxygen atom). These electrons fill the orbitals in order, following the Pauli exclusion principle and Hund’s rule.

3. Occupation of Orbitals and Presence of Unpaired Electrons

When electrons reach the π2p* antibonding orbitals, they fill as follows:

• Two π* orbitals exist
• Electrons enter singly into each orbital before pairing, according to Hund’s rule

This results in:

• Two unpaired electrons in the π* antibonding orbitals

These unpaired electrons are the reason O₂ is paramagnetic.

If all electrons were paired, O₂ would be diamagnetic. But experimental evidence strongly confirms the presence of unpaired electrons, proving MO theory correct.

4. How Unpaired Electrons Cause Paramagnetism

Unpaired electrons behave like tiny magnets because they have:

• Intrinsic spin
• Magnetic moment

When a magnetic field is applied:

• These unpaired electrons align with the field
• The molecule is pulled toward the magnetic source

This attraction is the characteristic sign of paramagnetism.

Thus, O₂ shows:

• Strong attraction to a magnetic field
• Increased paramagnetic strength compared to molecules with only one unpaired electron

5. Bond Order in O₂ and Its Relation to Paramagnetism

The bond order of O₂ is calculated from MO theory:

This bond order indicates:

• A double bond in O₂
• Moderate stability

However, bond order does not tell us about magnetism.
Magnetic behavior depends solely on the presence of unpaired electrons.

Therefore:

• O₂ has bond order 2
• But it is paramagnetic because of unpaired e⁻ in π* orbitals

6. Comparison with Other Oxygen Species

MO theory also explains magnetism in oxygen ions:

O₂⁺ (superoxide ion)

• One electron removed
• Still has an unpaired electron
  → Paramagnetic

O₂²⁻ (peroxide ion)

• Two extra electrons
• All electrons paired
  → Diamagnetic

This proves magnetic properties follow the electron arrangement in MO diagrams.

7. Why Lewis Structures Fail to Explain O₂ Magnetism

Lewis structure of O₂ shows:

• A double bond between the atoms
• All electrons paired

If this were true, O₂ would be diamagnetic.
But experiments show the opposite.

MO theory succeeds because it:

• Accounts for orbital mixing
• Provides correct electron distribution
• Shows unpaired electrons clearly

Thus, MO theory gives the real picture of O₂’s magnetic behavior.

Conclusion

O₂ is paramagnetic because it contains two unpaired electrons in the π* antibonding molecular orbitals. These unpaired electrons create magnetic moments that cause oxygen molecules to be attracted to an external magnetic field. While Lewis structures fail to explain this behavior, molecular orbital theory clearly shows the correct electron arrangement and successfully describes the magnetic nature of oxygen. Therefore, O₂’s paramagnetism is a direct consequence of its molecular orbital configuration.