Explain this observation using molecular orbitals – [Free] B102
explain this observation using molecular orbitals
Question:
Explain the observation that oxygen (O₂) is paramagnetic using Molecular Orbital Theory.
Answer:
🔍 The Observation
Experimental evidence shows that oxygen (O₂) is paramagnetic, meaning it is attracted to a magnetic field. This is unexpected when looking only at the Lewis structure of O₂, which shows all electrons paired in a double bond.
📘 Molecular Orbital Theory Basics
In Molecular Orbital (MO) Theory, atomic orbitals combine to form molecular orbitals that are classified as:
- Bonding orbitals: lower energy, stabilize the molecule
- Antibonding orbitals: higher energy, destabilize the molecule
Electron filling follows the Aufbau principle, Pauli exclusion principle, and Hund’s rule.
🧪 Constructing the MO Diagram for O₂
Each oxygen atom has 8 electrons, so O₂ has a total of 16 electrons.
Filling order of molecular orbitals:
- Bonding orbitals: 8 electrons
- Antibonding orbitals: 4 electrons
⚡ Explaining Paramagnetism
The π*(2px) and π*(2py) orbitals each contain one unpaired electron. According to Hund’s Rule, these orbitals are singly occupied to minimize electron repulsion.
Unpaired electrons result in a magnetic moment, which explains the paramagnetic behavior of O₂.
🔢 Bond Order Calculation
Bond order = ½ (Number of bonding electrons − Number of antibonding electrons)
This confirms a double bond in O₂, consistent with experimental observations.
✅ Conclusion
Molecular Orbital Theory successfully explains the paramagnetism of O₂ by revealing the presence of two unpaired electrons in its MO diagram. These electrons are not shown in traditional Lewis structures. Additionally, MO Theory validates the bond order of 2, confirming a double bond between the oxygen atoms.