Chapter 11: Chemical Bonding II – Molecular Shapes, VSEPR & MO Theory (Study Notes)

Chapter 11: Chemical Bonding II – Molecular Shapes, VSEPR & MO Theory

Valence Shell Electron Pair Repulsion Theory (VSEPR Theory)

The VSEPR theory is used to predict the shapes of molecules based on the repulsion between electron groups around a central atom. Electron groups (bonding pairs and lone pairs) arrange themselves to minimize repulsion, resulting in specific molecular geometries.

• Electron Groups: Regions of electron density (bonds or lone pairs) around a central atom.

• Bonding Pairs: Shared pairs of electrons between atoms.

• Lone Pairs: Non-bonding pairs of electrons localized on a single atom.

• Key Principle: Electron groups will orient themselves as far apart as possible.

Example: How many electron groups are on the nitrogen atom in NH3 based on its Lewis Dot Structure?

Additional info: For NH3, there are three bonding pairs and one lone pair, totaling four electron groups.

Electron Geometry

Electron geometry describes the arrangement of all electron groups (bonding and lone pairs) around the central atom. It is determined by the number of electron groups.

Electron Groups	Geometry	Example
2	Linear	CO2
3	Trigonal Planar	BF3
4	Tetrahedral	CH4
5	Trigonal Bipyramidal	PCl5
6	Octahedral	SF6

Example: Determine the electron geometry for H2S (hydrogen sulfide). H2S has four electron groups (two bonding pairs, two lone pairs) → Tetrahedral electron geometry.

Molecular Geometry

Molecular geometry refers to the arrangement of only the atoms (not lone pairs) in a molecule. It is derived from the electron geometry but considers only the positions of the atoms.

Electron Groups	Lone Pairs	Molecular Geometry	Example
2	0	Linear	CO2
3	0	Trigonal Planar	BF3
3	1	Bent	SO2
4	0	Tetrahedral	CH4
4	1	Trigonal Pyramidal	NH3
4	2	Bent	H2O
5	0	Trigonal Bipyramidal	PCl5
5	1	Seesaw	SF4
5	2	T-shaped	ClF3
5	3	Linear	XeF2
6	0	Octahedral	SF6
6	1	Square Pyramidal	BrF5
6	2	Square Planar	XeF4

Example: Determine the molecular geometry for SOCl2. SOCl2 has four electron groups (one lone pair), so its molecular geometry is trigonal pyramidal.

Equatorial and Axial Positions

In molecules with trigonal bipyramidal geometry (five electron groups), positions are classified as equatorial (in the plane) and axial (above and below the plane). Lone pairs preferentially occupy equatorial positions to minimize repulsion.

• Equatorial Positions: Three positions in the plane of the molecule.

• Axial Positions: Two positions perpendicular to the plane.

• Application: Predicting the most stable arrangement of atoms and lone pairs in molecules like SF4 and PCl5.

Example: Based on equatorial and axial positions, predict the most likely structure of PF5.

Molecular Polarity

Molecular polarity depends on both the polarity of individual bonds and the overall shape of the molecule. A molecule is polar if it has a net dipole moment due to asymmetrical distribution of electron density.

• Nonpolar Molecule: All bonds are nonpolar or the molecule is symmetrical (dipoles cancel).

• Polar Molecule: At least one bond is polar and the molecule has an asymmetrical shape.

Electron Groups	Perfect Shape	Nonpolar Example	Polar Example
2	Linear	CO2	NO
3	Trigonal Planar	BF3	SO2
4	Tetrahedral	CH4	NH3
5	Trigonal Bipyramidal	PCl5	SF4
6	Octahedral	SF6	BrF5

Example: Is CCl4 polar or nonpolar? CCl4 is nonpolar due to its symmetrical tetrahedral shape.

Bond Angles

Bond angles are determined by the number of electron groups and the presence of lone pairs. Ideal bond angles are observed in perfect geometries, but lone pairs compress bond angles due to greater repulsion.

Electron Groups	Geometry	Ideal Bond Angle
2	Linear	180°
3	Trigonal Planar	120°
4	Tetrahedral	109.5°
5	Trigonal Bipyramidal	90°, 120°
6	Octahedral	90°

Example: What is the H–C–H bond angle in CH4? 109.5° (tetrahedral geometry).

Hybridization

Hybridization describes the mixing of atomic orbitals to form new hybrid orbitals suitable for bonding. The type of hybridization depends on the number of electron groups around the central atom.

Electron Groups	Hybridization	Geometry
2	sp	Linear
3	sp2	Trigonal Planar
4	sp3	Tetrahedral
5	sp3d	Trigonal Bipyramidal
6	sp3d2	Octahedral

Example: What is the hybridization of the sulfur atom in SF6? sp3d2 hybridization.

Molecular Orbital Theory

Molecular Orbital (MO) Theory explains bonding by combining atomic orbitals to form molecular orbitals that are delocalized over the entire molecule. Electrons fill these orbitals according to the Pauli Exclusion Principle and Hund's Rule.

• Bonding Molecular Orbitals: Lower energy, promote bond formation.

• Antibonding Molecular Orbitals: Higher energy, oppose bond formation.

• Bond Order: Indicates the stability and strength of a bond.

Bond Order Formula:

$\text{Bond Order} = \frac{1}{2}(\text{Number of bonding electrons} - \text{Number of antibonding electrons})$

Example: Determine the bond order of the NO- ion.

MO Theory: Homonuclear and Heteronuclear Diatomic Molecules

MO diagrams differ for homonuclear (same element) and heteronuclear (different elements) diatomic molecules. The energy ordering of molecular orbitals depends on the atomic number and electronegativity of the atoms involved.

• Homonuclear Diatomics: O2, N2, F2, etc.

• Heteronuclear Diatomics: CO, NO, HF, etc.

Example: Construct the MO diagram for CO and determine its bond order.

Practice Problems

• Determine electron and molecular geometries for various molecules (e.g., CH4, SOCl2, SF4).

• Predict molecular polarity for compounds such as CCl4, BF3, and SCl2.

• Calculate bond angles and hybridization for given molecules.

• Draw and interpret molecular orbital diagrams for diatomic molecules.

• Calculate bond order and relate it to molecular stability.

Additional info: These notes cover the essential concepts and applications of VSEPR theory, electron and molecular geometry, molecular polarity, bond angles, hybridization, and molecular orbital theory, as required for a General Chemistry college course.