CHEM1011 · Chemistry 1a
Molecular Shape: VSEPR, Valence Bond and Molecular Orbital Theory
Week 3 turns Lewis structures into three-dimensional shapes with VSEPR (electron-domain geometry, molecular shape and bond angles up to six domains) and into orbital models with valence bond theory (sp/sp²/sp³ hybrids, σ/π overlap). Molecular-orbital diagrams for homonuclear diatomics — filling MOs, computing bond order and explaining O₂'s paramagnetism — are Mastery content and a signature final-exam item.
What this chapter covers
- 01VSEPR: electron domains (bonding + lone pairs; a multiple bond = one domain) arrange to minimise repulsion
- 02Domain geometries and ideal angles: 2 linear 180°, 3 trigonal planar 120°, 4 tetrahedral 109.5°, 5 trigonal bipyramidal 90°/120°, 6 octahedral 90°
- 03Lone pairs repel more than bonding pairs, compressing angles below ideal (NH₃ ~107°, H₂O ~104.5°)
- 04Valence bond theory: bonds as orbital overlap; σ (head-on) and π (side-on) bonds
- 05Hybridisation up to sp³: sp (linear, 2 domains), sp² (trigonal planar, 3), sp³ (tetrahedral, 4)
- 06Molecular-orbital theory by LCAO: bonding and antibonding MOs; fill by Aufbau/Hund/Pauli
- 07Bond order = ½(bonding electrons − antibonding electrons); the second-row ordering swap at O₂
- 08O₂ paramagnetism and the MO-vs-VBT comparison — each model's strengths and limits (Mastery)
Bond order and magnetism of O₂ from an MO diagram
- +1Count valence electrons: each oxygen has the 2s²2p⁴ valence set = 6, so O₂ has 12 valence electrons to place (core 1s electrons are omitted, they don't affect bonding).
- +1Fill the second-row MOs in the O₂/F₂ ordering (σ2p below π2p): σ2s² σ*2s² σ2p² π2p⁴ π*2p² — that accounts for all 12 electrons.
- +1Bond order = ½(bonding − antibonding) = ½[(2 + 2 + 4) − (2 + 2)] = ½(8 − 4) = 2, i.e. a double bond.
- +1The last two electrons go singly into the two degenerate π*2p orbitals with parallel spins (Hund's rule), so O₂ has 2 unpaired electrons and is paramagnetic.
Key terms
- VSEPR theory
- Valence-shell electron-pair repulsion: electron domains around a central atom arrange to minimise repulsion, setting the electron-domain geometry; lone pairs then determine the molecular shape and compress bond angles.
- Electron domain
- A region of electron density around the central atom — a lone pair, or a single, double or triple bond (each multiple bond counts as one domain). The domain count fixes the geometry.
- Hybridisation
- Mixing atomic orbitals into equivalent hybrids that match the observed shape: sp (linear, 2 domains), sp² (trigonal planar, 3), sp³ (tetrahedral, 4). The σ framework comes from hybrids; π bonds use unhybridised p orbitals.
- Molecular orbital (MO)
- An orbital spread over the whole molecule, formed by the linear combination of atomic orbitals (LCAO). Two atomic orbitals give one lower-energy bonding MO and one higher-energy antibonding MO (σ*/π*).
- Bond order
- Bond order = ½(bonding electrons − antibonding electrons). A value greater than zero implies a bound molecule; higher bond order means a shorter, stronger bond (He₂ has bond order 0 and does not exist).
- Paramagnetism (MO prediction)
- A species with unpaired electrons in its MO diagram is paramagnetic; O₂'s two unpaired π*2p electrons are correctly predicted by MO theory, a success over the Lewis/VBT picture.
Molecular Shape: VSEPR, Valence Bond and Molecular Orbital Theory FAQ
How is electron-domain geometry different from molecular shape?
Electron-domain geometry counts all domains — bonding and lone pairs — to set the arrangement (e.g. four domains = tetrahedral). Molecular shape describes only where the atoms are, so lone pairs change it: four domains with one lone pair gives trigonal pyramidal (NH₃), with two lone pairs gives bent (H₂O). Bond angles also shrink below the ideal because lone pairs repel more strongly than bonding pairs.
Why does MO theory predict O₂ is paramagnetic when the Lewis structure looks fine?
The Lewis structure O=O shows all electrons paired, predicting diamagnetism — but liquid O₂ is drawn to a magnet, so it must have unpaired electrons. MO theory places the last two electrons singly in the two degenerate π*2p orbitals (Hund's rule), giving two unpaired electrons and correctly predicting paramagnetism. This is the classic example of MO theory succeeding where the localised Lewis/VBT picture fails.
When would I use valence bond theory versus molecular orbital theory?
Use VBT (hybrid orbitals, localised σ/π bonds) when you want an intuitive picture of shape and connectivity — it maps cleanly onto VSEPR. Use MO theory (delocalised bonding/antibonding orbitals) when you need bond order for odd-electron or ionised species, magnetism, or spectra. Neither is 'the truth'; each model has strengths and limits, and the exam may ask you to say why chemists switch between them.
Can Sia help me with VSEPR and MO diagrams?
Yes. Sia can take a Lewis structure through the VSEPR steps to a shape and bond angle, assign the hybridisation, and build a second-row MO diagram, fill it, compute the bond order and call the magnetism. It explains each step and checks your working; it does not do graded assessment, and UNSW academic-integrity rules apply.
Exam move
Split this topic into a shapes half and an orbitals half and drill each. For VSEPR, memorise the domain-count table (2 linear through 6 octahedral, with ideal angles) and practise the lone-pair variants that give bent, trigonal pyramidal, seesaw, T-shaped, square pyramidal and square planar shapes, always noting the small angle compression from lone-pair repulsion. Tie each shape to its hybridisation (sp/sp²/sp³) so the VBT question falls out of the VSEPR answer. For the Mastery MO layer, learn the second-row filling order cold — including the σ2p/π2p swap at O₂ — and treat bond order = ½(bonding − antibonding) plus a Hund's-rule magnetism call as a single routine. Keep O₂ paramagnetism as your worked showcase of MO over VBT. Rehearse comparing X₂ with X₂⁺ and X₂⁻ so you can predict how bond order and stability change on ionisation. When an MO diagram confuses you, ask Sia to rebuild it orbital by orbital.
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