Molecular Shape, Polarity, and Intermolecular Forces in Organic Chemistry

Shape and Reactivity of Molecules

Importance of Molecular Shape

The shape of a molecule plays a crucial role in determining its chemical properties and reactivity. Understanding molecular geometry helps predict how molecules interact, react, and function in organic chemistry.

• Molecular shape affects physical and chemical properties such as boiling point, solubility, and reactivity.

• Shape is determined by the arrangement of atoms and electron pairs around a central atom.

Valence Shell Electron Pair Repulsion (VSEPR) Theory

Basic Principles of VSEPR Theory

VSEPR theory is used to predict the geometry of molecules based on the repulsion between electron pairs (bonding and lone pairs) around a central atom.

• Valence Shell Electron Pair Repulsion (VSEPR): Electron pairs (bonding or lone pairs) arrange themselves as far apart as possible to minimize repulsion.

• Regions of electron density include:

  • Single bond = 1 region

  • Double bond = 1 region

  • Triple bond = 1 region

  • Lone pair = 1 region

Examples of Molecular Geometry

• CO2 (Carbon Dioxide):

  • Lewis Structure: O = C = O

  • Bonds around carbon are 180° apart

  • Molecular shape: linear

• BF3 (Boron Trifluoride):

  • 120° bond angles around boron

  • Molecular shape: trigonal planar

• CH4 (Methane):

  • 4 regions of electron density

  • Molecular shape: tetrahedral, bond angles 109.5°

  • 3D representation: wedge (in front), dash (behind), lines (in plane)

• NH3 (Ammonia):

  • 3 bonds + 1 lone pair = 4 regions of electron density

  • Electron region arrangement: tetrahedral

  • Molecular shape: trigonal pyramidal (lone pair not seen in shape)

Electronic Arrangement/Geometry

• 2 regions: linear

• 3 regions: trigonal planar

• 4 regions: tetrahedral

• Only nuclei (atoms) are visible in molecular shape; electrons are too small to be seen directly.

Summary Table: Molecular Geometry and Electron Domains

This table summarizes the relationship between the number of bonds, lone pairs, electronic geometry, and molecular shape for common molecules.

Bonds	Lone Pairs	Electronic Geometry	Molecular Shape	Example
2	0	Linear	Linear	CO2
3	0	Trigonal Planar	Trigonal Planar	BF3
2	1	Trigonal Planar	Bent	SO2
4	0	Tetrahedral	Tetrahedral	CH4
3	1	Tetrahedral	Trigonal Pyramidal	NH3
2	2	Tetrahedral	Bent	H2O

Molecular Polarity

Definition and Determinants

Molecular polarity is determined by the shape of the molecule and the distribution of charge due to differences in electronegativity between atoms.

• Polarity results from an unequal distribution of charge within a molecule.

• Depends on both bond polarity and molecular geometry.

Examples of Polarity

• HCl:

  • Polar bond (Cl more electronegative than H)

  • Polar molecule

• CO2:

  • Polar bonds (O more electronegative than C)

  • Linear shape causes dipole moments to cancel

  • Nonpolar molecule

• H2O:

  • Polar bonds (O more electronegative than H)

  • Bent shape means dipoles do not cancel

  • Polar molecule

• NH3:

  • Polar bonds (N more electronegative than H)

  • Trigonal pyramidal shape, dipoles add up

  • Polar molecule

• CCl4:

  • Tetrahedral shape with identical substituents

  • All bond dipoles cancel

  • Nonpolar molecule

Intermolecular Forces

Types of Intermolecular Forces

Intermolecular forces are interactions between atoms or molecules that affect physical properties such as boiling and melting points.

• London Dispersion Forces:

  • Momentary forces of attraction between all molecules due to temporary uneven electron distribution.

  • Present in all molecules; only significant in nonpolar molecules.

  • Strength increases with molar mass.

  • Also called induced dipole or dispersion forces.

• Dipole-Dipole Attraction:

  • Occurs between polar molecules with permanent dipoles.

  • Results from unequal distribution of electrons.

  • Interaction between two dipoles.

Summary Table: Intermolecular Forces

Type of Force	Occurs Between	Relative Strength	Example
London Dispersion	All molecules/atoms	Weakest	Noble gases, nonpolar molecules
Dipole-Dipole	Polar molecules	Intermediate	HCl, SO2
Hydrogen Bonding	H bonded to N, O, or F	Strongest (of van der Waals)	H2O, NH3

Additional info: Hydrogen bonding is a special case of dipole-dipole interaction, particularly strong when hydrogen is bonded to highly electronegative atoms (N, O, F).