Unit 2: Chemical Bonding and Molecular Structure
Table of Contents
Valence Bond (VB) Approach
VBT describes a covalent bond as the overlap of half-filled atomic orbitals from two different atoms. The electrons in the overlapping orbitals must have opposite spins.
Shapes of Molecules (VSEPR and Hybridization)
VBT uses the concepts of Hybridization and VSEPR Theory to explain molecular shapes.
- VSEPR (Valence Shell Electron Pair Repulsion): Electron pairs (both bonding and lone pairs) around a central atom repel each other and arrange themselves to be as far apart as possible, minimizing repulsion. The order of repulsion is: LP-LP > LP-BP > BP-BP.
- Hybridization: The mixing of atomic orbitals (s, p, d) of slightly different energies to form a new set of identical "hybrid orbitals" that are degenerate and better suited for bonding.
| Hybridization | Electron Pairs | Electron Geometry | Example(s) | Shape |
|---|---|---|---|---|
| sp | 2 | Linear | BeCl2 | Linear (180°) |
| sp² | 3 | Trigonal Planar | BF3, SO2 (AX₂E) | Trigonal Planar (120°), Bent |
| sp³ | 4 | Tetrahedral | CH4, NH3 (AX₃E), H2O (AX₂E₂) | Tetrahedral (109.5°), Trigonal Pyramidal, Bent |
| dsp² | 4 | Square Planar | [Ni(CN)4]2- | Square Planar (90°) |
| sp³d | 5 | Trigonal Bipyramidal | PCl5, SF4 (AX₄E), ClF3 (AX₃E₂) | Trigonal Bipyramidal (90°, 120°), See-Saw, T-Shape |
| sp³d² | 6 | Octahedral | SF6, BrF5 (AX₅E), XeF4 (AX₄E₂) | Octahedral (90°), Square Pyramidal, Square Planar |
- In Trigonal Bipyramidal (sp³d) geometry, lone pairs always occupy the equatorial positions first to minimize repulsion.
- In Octahedral (sp³d²) geometry, the first lone pair can go anywhere. The second lone pair goes trans (180°) to the first one.
Concept of Resonance
Resonance is used when a single Lewis structure cannot adequately describe the bonding in a molecule. The actual structure is an average or "hybrid" of two or more resonance structures (or canonical forms), which differ only in the placement of π-electrons and lone pairs.
Example: Ozone (O3). The two O-O bonds are identical in length, which is explained by resonance between two contributing structures.
Molecular Orbital (MO) Approach
MOT is a more advanced model where all atomic orbitals (AOs) combine to form an equal number of molecular orbitals (MOs) that are delocalized over the *entire* molecule.
Rules for the LCAO Method
MOs are formed by the Linear Combination of Atomic Orbitals (LCAO). For effective combination, the AOs must have:
- Similar Energy: (e.g., 2s combines with 2s, 2p with 2p).
- Same Symmetry: (e.g., a pz orbital can combine with another pz (to make σ) but not with a px (which makes π)).
- Sufficient Overlap: The atoms must be close enough.
Bonding, Antibonding, and Nonbonding MOs
- Bonding MO (BMO): Formed by constructive interference (ψA + ψB). It is lower in energy than the original AOs and is stabilizing. (e.g., σss, πpp).
- Antibonding MO (ABMO): Formed by destructive interference (ψA - ψB). It is higher in energy than the AOs and is destabilizing. It has a node between the nuclei. (e.g., σ*ss, π*pp).
- Nonbonding MO: An atomic orbital that does not interact with any other orbital (due to symmetry or energy mismatch). Its energy in the MO diagram is unchanged.
MO Treatment of Homonuclear Diatomic Molecules
Electrons are filled into MOs using Aufbau, Pauli, and Hund's rules. The order of filling depends on s-p mixing.
| For B2, C2, N2 (s-p mixing) | For O2, F2, Ne2 (no s-p mixing) |
|---|---|
| σ1s < σ*1s < σ2s < σ*2s < π2p_{x,y} < σ2pz < π*2p_{x,y} < σ*2pz | σ1s < σ*1s < σ2s < σ*2s < σ2pz < π2p_{x,y} < π*2p_{x,y} < σ*2pz |
- Bond Order (BO): BO = (1) / (2) × (Bonding e- - Antibonding e-)
- N2: (10 bonding e-, 4 antibonding e-) → BO = 3, Diamagnetic (no unpaired e-).
- O2: (10 bonding e-, 6 antibonding e-) → BO = 2, Paramagnetic (2 unpaired e- in π* orbitals).
- B2: (6 bonding e-, 4 antibonding e-) → BO = 1, Paramagnetic (2 unpaired e- in π orbitals).
MO Treatment of Heteronuclear Diatomic Molecules
For molecules like CO and NO, the MO diagram is asymmetric because the AOs of the more electronegative atom (O) are lower in energy.
- CO (14 e-): Similar to N2. BO = 3. Diamagnetic.
- NO (15 e-): Has one extra electron in a π* orbital. BO = 2.5. Paramagnetic.
- NO+ (14 e-): Isoelectronic with N2 and CO. BO = 3. Diamagnetic.
Comparison of VB and MO Approaches
| Valence Bond Theory (VBT) | Molecular Orbital Theory (MOT) |
|---|---|
| Considers bonds as localized between two atoms. | Considers electrons as delocalized over the entire molecule. |
| A bond is formed by the overlap of atomic orbitals. | A bond is formed by the combination of atomic orbitals (LCAO). |
| Concept of hybridization is central to explaining geometry. | Geometry is inherent in the symmetry of the MOs formed. |
| Simple to apply and visualize. | More complex, but provides a more accurate picture. |
| Fails to explain the paramagnetism of O2. | Correctly predicts the paramagnetism of O2 (and B2). |
| Does not easily explain fractional bond orders or excited states. | Easily explains bond order, magnetic properties, and electronic spectra. |