DPST1031 Chap.5 Valence Bond, Hybrid and Molecular Orbitals
Valence Bond, Hybrid and Molecular Orbitals
Lewis structures say where electron pairs are; this chapter explains why the bonds form and how strong they are. Valence bond theory treats a bond as overlap between orbitals on neighbouring atoms, distinguishes head-on σ bonds from sideways π bonds, and uses sp, sp² and sp³ hybrid orbitals to account for linear, trigonal planar and tetrahedral geometry.
Molecular orbital theory then combines atomic orbitals across the whole molecule into bonding and antibonding orbitals, giving bond orders, bond length trends and the paramagnetism of oxygen that a Lewis structure cannot show. The chapter ends by comparing the strengths and weaknesses of the two models, a named Mastery outcome.
What this chapter covers
- 01
Bond formation and the potential energy curve
- 02
Sigma and pi bonds from orbital overlap
- 03
Hybrid orbitals: sp, sp² and sp³
- 04
Hybridisation in organic molecules
- 05
Bonding and antibonding molecular orbitals
- 06
MO diagrams for second-row diatomics
- 07
Bond order, bond length and stability
- 08
Strengths and weaknesses of each model
Worked example · free
Bond order and magnetism of O₂
- +1O₂ has 12 valence electrons: σ2s², σ*2s², σ2p², π2p⁴, π*2p², using the oxygen ordering with σ2p below π2p.
- +1Bonding electrons 8, antibonding electrons 4, so bond order = (8 − 4)/2 = 2.
- +1The two π*2p electrons occupy separate orbitals with parallel spins, as Hund's rule requires.
- +1Two unpaired electrons make O₂ paramagnetic, which a fully paired Lewis structure cannot show.
Key terms
- Sigma bond
- A bond formed by head-on orbital overlap with electron density along the internuclear axis.
- Pi bond
- A bond formed by sideways overlap of parallel p orbitals above and below the bond axis.
- Hybridisation
- The mixing of atomic orbitals on one atom to form equivalent hybrid orbitals.
- Molecular orbital
- An orbital spread over a whole molecule, formed by combining atomic orbitals.
- Antibonding orbital
- A higher-energy molecular orbital with a node between the nuclei that weakens bonding.
- Diamagnetic species
- A species in which every electron is paired, so it is weakly repelled by a magnetic field.
Valence Bond, Hybrid and Molecular Orbitals FAQ
How do I find the hybridisation of an atom?
Count the electron domains around it using VSEPR, counting lone pairs and each bond once. Two domains means sp, three means sp² and four means sp³. The number of unhybridised p orbitals left over tells you how many π bonds that atom can form.
Why is O₂ paramagnetic when its Lewis structure has all electrons paired?
The Lewis structure forces every electron into a pair. In the molecular orbital diagram, oxygen's last two electrons occupy two separate π*2p orbitals of equal energy, and Hund's rule puts them there unpaired. That prediction matches the observation that liquid oxygen is attracted to a magnet.
Why does the orbital order change between N₂ and O₂?
For boron, carbon and nitrogen the 2s and 2p atomic orbitals are close in energy and mix, which pushes the σ2p molecular orbital above the π2p pair. In oxygen and fluorine the gap is larger, so σ2p falls below π2p. Use the correct ordering before filling electrons.
What does a bond order of zero mean?
It means the bonding and antibonding electrons cancel, so there is no net bond and the species is not expected to exist as a stable molecule. He₂, with two bonding and two antibonding electrons, is the standard example used to show this prediction.
Why can a molecule not rotate freely about a double bond?
The π part of a double bond comes from sideways overlap of parallel p orbitals above and below the bond axis. Twisting one end would turn those p orbitals out of line and break the overlap, which costs a lot of energy, so rotation about C=C is restricted while rotation about C−C is free.
Exam move
Link this chapter back to VSEPR first: for every molecule you drew in Chapter 4, write the hybridisation of the central atom and the number of σ and π bonds. Then learn the two diatomic orderings and practise filling diagrams for H₂ through F₂ and their common ions, writing the bond order and magnetism each time. Rank sets of species by bond length and bond energy from bond order alone.
Finally, prepare a short comparison of the two models with one property each explains well and one it cannot, because mastery questions ask for exactly that judgement.
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