Explain this observation using molecular orbitals – [Free] B102
explain this observation using molecular orbitals
Question:
Explain the observation that oxygen (Oβ) is paramagnetic using Molecular Orbital Theory.
Answer:
π The Observation
Experimental evidence shows that oxygen (Oβ) is paramagnetic, meaning it is attracted to a magnetic field. This is unexpected when looking only at the Lewis structure of Oβ, which shows all electrons paired in a double bond.
π 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 Hundβs rule.
π§ͺ Constructing the MO Diagram for Oβ
Each oxygen atom has 8 electrons, so Oβ has a total of 16 electrons.
Filling order of molecular orbitals:
- Bonding orbitals: 8 electrons
- Antibonding orbitals: 4 electrons
β‘ Explaining Paramagnetism
The Ο*(2px) and Ο*(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 Oβ.
π’ Bond Order Calculation
Bond order = Β½ (Number of bonding electrons β Number of antibonding electrons)
This confirms a double bond in Oβ, consistent with experimental observations.
β Conclusion
Molecular Orbital Theory successfully explains the paramagnetism of Oβ 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.