Molecular Geometry and Intermolecular Forces: VSEPR Theory and Types of Intermolecular Attractions

Molecular Geometry

Introduction to Molecular Structure

Molecular geometry refers to the three-dimensional arrangement of atoms within a molecule. Understanding molecular geometry is essential for predicting molecular properties, reactivity, and physical behavior. The Valence Shell Electron Pair Repulsion (VSEPR) theory is the primary model used to predict the shapes of molecules based on the repulsions between electron pairs around a central atom.

• Structure: The 3D shape of a molecule, determined by the positions of atoms bonded to the central atom.

• VSEPR Theory: Electron pairs (bonding and lone pairs) arrange themselves as far apart as possible to minimize repulsion and maximize stability.

Steps for Predicting Molecular Geometry Using VSEPR

1. Write the Lewis structure of the molecule.

2. Count the total number of atoms and lone pairs attached to the central atom.

3. Arrange atoms and lone pairs around the central atom to minimize repulsion (maximize distance).

4. Determine the electron pair geometry (includes both bonding and lone pairs).

5. Determine the molecular geometry (positions of atoms only, ignoring lone pairs).

Electron Pair Geometry vs. Molecular Geometry

• If the central atom has no lone pairs, the electron pair geometry and molecular geometry are the same.

• If the central atom has lone pairs, the electron pair geometry and molecular geometry differ. The molecular geometry is found by "removing" the lone pairs from the electron pair geometry.

Common Electron Pair and Molecular Geometries

Electron Pair Geometry	Molecular Geometry	Example	Bond Angle
Linear	Linear	CO2	180°
Trigonal Planar	Trigonal Planar	BF3	120°
Trigonal Planar	Bent	SO2	~120°
Tetrahedral	Tetrahedral	CH4	109.5°
Tetrahedral	Trigonal Pyramidal	NH3	107.3°
Tetrahedral	Bent	H2O	104.5°
Trigonal Bipyramidal	Trigonal Bipyramidal	PCl5	90°, 120°
Trigonal Bipyramidal	Seesaw	SF4	~90°, ~120°
Trigonal Bipyramidal	T-shaped	ClF3	~90°
Trigonal Bipyramidal	Linear	XeF2	180°
Octahedral	Octahedral	SF6	90°
Octahedral	Square Pyramidal	BrF5	~90°
Octahedral	Square Planar	XeF4	90°

Additional info: The above table is a standard summary of VSEPR geometries, expanded for clarity.

Placement of Lone Pairs in Trigonal Bipyramidal Geometry

• Lone pairs prefer equatorial positions over axial positions to minimize 90° interactions and repulsion.

• If a lone pair occupies an axial position, it has three 90° interactions; in an equatorial position, only two 90° interactions.

• Rule: Lone pairs always go to equatorial positions in trigonal bipyramidal geometry for maximum stability.

Examples: Predicting Geometries

Molecule	Electron Pair Geometry	Molecular Geometry
BBr3	Trigonal Planar	Trigonal Planar
CF4	Tetrahedral	Tetrahedral
NF3	Tetrahedral	Trigonal Pyramidal
PF5	Trigonal Bipyramidal	Trigonal Bipyramidal
SeF4	Trigonal Bipyramidal	Seesaw
SF6	Octahedral	Octahedral

Examples: Molecular Geometry of Organic Molecules

Molecule	Electron Count	Molecular Geometry
C2H4 (Ethene)	12 e-	Each C: Trigonal Planar
CH3OH (Methanol)	14 e-	C: Tetrahedral, O: Bent
NH2CONH2 (Urea)	24 e-	C: Trigonal Planar, N: Trigonal Pyramidal

Intermolecular Forces

Introduction to Intermolecular Forces

Intermolecular forces are the attractive forces between molecules, also known as Van der Waals forces. These forces are much weaker than the covalent or ionic bonds that hold atoms together within a molecule, but they play a crucial role in determining the physical properties of substances, such as boiling and melting points.

Comparison: Intermolecular Forces vs. Chemical Bonds

• Strength: Intermolecular forces (1–40 kJ/mol) are much weaker than chemical bonds (150–800 kJ/mol).

• Distance Dependence: The strength of intermolecular forces decreases rapidly as the distance between molecules increases.

Types of Intermolecular Forces

1. Dipole–Dipole Forces: Attractive forces between polar molecules due to the interaction of permanent dipoles. The strength increases with greater molecular polarity and smaller atomic size.

2. Hydrogen Bonding: A special, strong type of dipole–dipole interaction that occurs when hydrogen is bonded to highly electronegative atoms (N, O, or F). Hydrogen bonds are responsible for the high boiling point of water ($H_2O$) compared to hydrogen sulfide ($H_2S$).

3. London Dispersion Forces (Van der Waals/Dispersion Forces): Weak, temporary attractive forces that arise from momentary dipoles induced by the movement of electrons. The strength increases with the number of electrons and the polarizability of the molecule.

Relative Strengths of Intermolecular Forces

• Hydrogen Bonding > Dipole–Dipole Forces > London Dispersion Forces

• All molecules exhibit London forces, but only polar molecules exhibit dipole–dipole forces, and only molecules with N–H, O–H, or F–H bonds exhibit hydrogen bonding.

Examples and Applications

• Water ($H_2O$): Exhibits hydrogen bonding, leading to a high boiling point (100°C).

• Hydrogen Sulfide ($H_2S$): Lacks hydrogen bonding, resulting in a much lower boiling point (–61°C).

• Halogens: The strength of London forces increases down the group: $F_2 < Cl_2 < Br_2 < I_2$.

Additional info: The above content expands on the original notes with standard textbook context and examples for clarity and completeness.