Oxygen is paramagnetic because it possesses unpaired electrons in its molecular orbitals, leading it to be attracted to an external magnetic field. In contrast, nitrogen is diamagnetic because all its electrons are paired, causing it to be weakly repelled by a magnetic field. This fundamental difference in magnetic behavior arises from the distinct arrangements of electrons within their molecular orbitals, as explained by Molecular Orbital (MO) Theory.
Understanding Molecular Orbital Theory
MO theory describes how atomic orbitals combine to form molecular orbitals, which span the entire molecule. Electrons then fill these molecular orbitals, following principles similar to filling atomic orbitals:
- Aufbau Principle: Electrons fill lower-energy orbitals first.
- Pauli Exclusion Principle: Each orbital can hold a maximum of two electrons with opposite spins.
- Hund's Rule: When degenerate (same energy) orbitals are available, electrons fill them singly with parallel spins before pairing up.
The specific order and energy levels of these molecular orbitals vary slightly depending on the atoms involved, particularly for diatomic molecules like nitrogen (N₂) and oxygen (O₂).
Nitrogen (N₂): A Diamagnetic Molecule
Nitrogen (N) has 7 electrons, so a nitrogen molecule (N₂) has a total of 14 electrons. According to MO theory, the molecular orbital energy order for N₂ (and other diatomic molecules up to nitrogen) is typically:
- σ2s (bonding)
- **σ2s*** (antibonding)
- π2p (bonding) - two degenerate orbitals
- σ2p (bonding)
- **π2p*** (antibonding) - two degenerate orbitals
- **σ2p*** (antibonding)
When the 14 electrons of N₂ are filled into these orbitals:
- (σ2s)²
- (σ2s*)²
- (π2p)⁴ (filling both degenerate π2p orbitals with two electrons each)
- (σ2p)²
All 14 electrons in N₂ occupy these molecular orbitals in pairs. Since there are no unpaired electrons in any of its molecular orbitals, nitrogen exhibits diamagnetism.
Oxygen (O₂): A Paramagnetic Molecule
Oxygen (O) has 8 electrons, so an oxygen molecule (O₂) has a total of 16 electrons. A crucial difference in the MO energy ordering for O₂ (and other diatomic molecules starting from oxygen) is that the σ2p bonding orbital is lower in energy than the π2p bonding orbitals due to less significant s-p orbital mixing. The revised order for O₂ is:
- σ2s (bonding)
- **σ2s*** (antibonding)
- σ2p (bonding)
- π2p (bonding) - two degenerate orbitals
- **π2p*** (antibonding) - two degenerate orbitals
- **σ2p*** (antibonding)
When the 16 electrons of O₂ are filled into these orbitals:
- (σ2s)²
- (σ2s*)²
- (σ2p)²
- (π2p)⁴ (filling both degenerate π2p orbitals with two electrons each)
- *(π2p)²* (filling the two degenerate π2p antibonding orbitals)
According to Hund's Rule, when filling the two degenerate π2p antibonding orbitals, the last two electrons enter these orbitals singly with parallel spins before pairing up. This results in two unpaired electrons in the π2p antibonding orbitals. These unpaired electrons give oxygen its paramagnetic properties.
Summary of Magnetic Properties
| Property | Nitrogen (N₂) | Oxygen (O₂) |
|---|---|---|
| Electron Count | 14 total electrons | 16 total electrons |
| MO Order | π2p bonding < σ2p bonding | σ2p bonding < π2p bonding |
| Unpaired Electrons | None | Two (in π2p* antibonding orbitals) |
| Magnetic Behavior | Diamagnetic (weakly repelled by B-field) | Paramagnetic (attracted to B-field) |
| Origin | All electrons paired | Presence of unpaired electrons |
What are Paramagnetism and Diamagnetism?
- Paramagnetism: A form of magnetism where certain materials are weakly attracted by an externally applied magnetic field. This attraction is due to the presence of one or more unpaired electrons in the atoms or molecules of the material. The unpaired electrons possess a magnetic moment that aligns with the external field.
- Diamagnetism: A property of all materials, which creates an induced magnetic field in an opposite direction to an externally applied magnetic field, causing a repulsive effect. Diamagnetism occurs in substances where all electrons are paired, meaning their individual magnetic moments cancel each other out.
This clear distinction in molecular orbital filling, specifically the presence of unpaired electrons in O₂ versus their absence in N₂, is the definitive reason for their differing magnetic properties.
