Intermolecular Forces: Types, Properties, and Applications

Intermolecular Forces

Overview and Applications

Intermolecular forces (IMFs) are the forces of attraction or repulsion between neighboring molecules, distinct from the stronger intramolecular forces (covalent or ionic bonds) within molecules. IMFs play a crucial role in determining the physical properties of substances, such as boiling and melting points, solubility, and molecular organization.

• Superhydrophobic, Hydrophobic, Hydrophilic Surfaces and Coatings: IMFs influence how surfaces interact with water and other substances, affecting applications in materials science and biology.

• Chemical Separations: Techniques like chromatography, azeotrope formation, and distillation rely on differences in IMFs between substances.

• Protein and DNA Structure: IMFs, especially hydrogen bonding, are essential for the stability and replication of biological macromolecules.

• Drug/Protein Interactions: Molecular shape and IMFs determine how drugs bind to proteins.

Consequences of Intermolecular Forces

Physical Properties and Non-Ideal Gas Behavior

IMFs affect the behavior of gases, liquids, and solids, especially under high pressure and low temperature conditions.

• Non-Ideal Gas Behavior: At high pressures and low temperatures, IMFs cause gases to deviate from ideal behavior, as described by the van der Waals equation.

• Condensation: Sufficiently strong IMFs can cause a gas to condense into a liquid.

• Density: IMFs keep molecules in close proximity, resulting in relatively high density for liquids at given pressure and temperature.

Van der Waals Forces

Definition and Molecular Properties

Van der Waals forces encompass all types of intermolecular attractions, including dipole-dipole, dipole-induced dipole, and London dispersion forces. These forces are responsible for the non-ideal behavior of gases and many physical properties of substances.

• Dipole Moment (μ): A measure of the separation of positive and negative charges in a molecule. It is produced when bond dipoles within a molecule do not cancel out.

• Polarizability (α): Indicates how easily the electron cloud of a molecule can be distorted. Larger atoms and molecules are generally more polarizable.

Equation:

• Dipole moment: $\mu = q \times r$

• Polarizability: $\alpha$ (unit: volume)

Dipole Moment and Polarizability

For Molecules

These properties determine the strength and type of IMFs present in a substance.

• Dipole Moment: Calculated by summing bond dipoles (magnitude and direction) to determine the net dipole for a molecule.

• Polarizability: Increases with atomic/molecular size; the unit is volume. Perturbation can be caused by ions, molecular dipoles, or electric fields.

Types of Intermolecular Forces

Dipole-Dipole Interactions

These occur between polar molecules with permanent dipoles. Molecules align so that the positive end of one is near the negative end of another, resulting in relatively strong attractions.

• Permanent Dipole: Molecules with a net dipole moment.

• Boiling Point Example: CH4 (nonpolar, μ = 0) has a much lower boiling point than CH3F (polar, μ = 1.65 D).

Dipole-Induced Dipole Interactions

These occur when a polar molecule induces a dipole in a neighboring nonpolar molecule, increasing the strength of attraction.

• Ion-Induced Dipole: An ion can also induce a dipole in a nonpolar molecule, leading to similar interactions.

London Dispersion Forces

London dispersion forces are present in all molecules, but are the only IMFs in nonpolar substances. They arise from temporary (instantaneous) dipoles due to random electron movement.

• Instantaneous Dipole: A temporary uneven distribution of electrons creates a dipole, which can induce a dipole in a neighboring molecule.

• Strength: Generally weak, but increases with molecular size and polarizability.

• Boiling Point Trend: More polarizable molecules (larger atoms) have higher boiling points.

Example:

• He: -268.9°C, Ne: -246.05°C, Ar: -185.8°C

Properties of Selected Nonpolar Compounds

Melting and Boiling Points of Halogens

The physical state and boiling/melting points of halogens are determined by the strength of London dispersion forces, which increase with the number of electrons.

Molecule	Total Number of Electrons	Melting Point (°C)	Boiling Point (°C)	Physical State at Room Temperature
F2	18	-220	-188	gas
Cl2	34	-102	-34	gas
Br2	70	-7	59	liquid
I2	106	114	184	solid

Trend: More electrons (larger atoms) result in greater polarizability and stronger London dispersion forces.

Boiling Point Trends for Halocarbons

Halogen Substitution Effects

Boiling points of halocarbons increase with the size and polarizability of the halogen atom. The trend is explained by the increasing strength of London dispersion forces and, in some cases, dipole-dipole interactions.

Boiling Points of Hydrides (Groups 14, 15, 16, 17)

Hydrogen Bonding Effects

Hydrides of N, O, and F (NH3, H2O, HF) have unusually high boiling points compared to other group members due to strong hydrogen bonding.

Hydrogen Bonding

Definition and Examples

Hydrogen bonding is a special type of dipole-dipole interaction that occurs when a hydrogen atom is covalently bonded to a highly electronegative atom (F, O, N, sometimes Cl or Br). The hydrogen atom interacts with a lone pair on another electronegative atom in a neighboring molecule.

• Intermolecular Hydrogen Bonding: The bond forms between two separate molecules.

• Example: In gaseous HF, molecules can arrange into cyclic (HF)n structures due to hydrogen bonding.

Equation:

• Hydrogen bond: $\text{D-H} \cdots \text{A}$ (D = donor atom, H = hydrogen, A = acceptor atom)

Summary Table: Types of Intermolecular Forces

Type of IMF	Occurs Between	Relative Strength	Example
London Dispersion	All molecules (esp. nonpolar)	Weak	He, CH4
Dipole-Dipole	Polar molecules	Moderate	CH3F, HCl
Hydrogen Bonding	H bonded to F, O, N (sometimes Cl, Br)	Strong	H2O, NH3, HF
Ion-Dipole	Ion and polar molecule	Strongest	Na+ in H2O

Additional info: The notes also reference the van der Waals equation for non-ideal gases, which is:

$\left( P + a \frac{n^2}{V^2} \right) (V - nb) = nRT$

where a and b are constants accounting for intermolecular forces and molecular volume, respectively.