Intermolecular Forces and the Physical Properties of Liquids and Solids

Intermolecular Forces and States of Matter

Overview of Intermolecular Forces

The physical state of a substance—whether it is a gas, liquid, or solid—is determined by the magnitude of intermolecular forces (IMFs) acting between its particles. These attractive forces are responsible for holding molecules together and influence many physical properties, such as boiling and melting points.

• Intermolecular forces are attractive forces between particles in a substance.

• The balance between kinetic energy (which tends to separate particles) and intermolecular attractions (which hold particles together) determines the state of matter.

Intra- Versus Intermolecular Forces

Definitions and Examples

It is important to distinguish between intramolecular and intermolecular forces:

• Intramolecular forces are the forces within a molecule, such as ionic bonding and covalent bonding. These are much stronger than intermolecular forces.

• Example: Breaking the two O–H covalent bonds in 1 mol of water requires 927 kJ.

• Intermolecular forces are forces between molecules, including ion-dipole, dipole-dipole, and London dispersion forces.

• Example: Overcoming the intermolecular attractions between water molecules to convert 1 mol of liquid water to vapor at 100°C requires only 41 kJ.

Types of Intermolecular Forces

Van der Waals Forces

Attractive forces between atoms or molecules in a pure substance are collectively called van der Waals forces. The main types are:

• Dipole-dipole interactions: Attractive forces between polar molecules. The strength depends on the magnitude of the dipole moment.

• Ion-dipole forces: Occur between an ion and the partial charge on the end of a polar molecule. These are especially important in solutions of ionic substances in polar solvents (e.g., NaCl in water).

• London dispersion forces: Result from Coulombic attractions between instantaneous dipoles of nonpolar molecules. These are present in all molecules but are the only IMFs in nonpolar substances.

Dipole-Dipole Forces

• Exist between oppositely charged ends of polar molecules.

• Effective only when polar molecules are very close together.

• The magnitude of the attractive forces depends on the magnitude of the dipole moment.

• Example: Solid and liquid arrangements of polar molecules show different strengths of dipole-dipole interactions.

Ion-Dipole Forces

• Occur between an ion and the partial charge on a polar molecule.

• Strength depends on the charge and size of the ion and the magnitude of the dipole.

• Example: Na+ (weak interaction) vs. Mg2+ (strong interaction) with water molecules.

• Important for dissolving ionic compounds in polar solvents.

London Dispersion Forces

• Result from temporary fluctuations in electron distribution, creating instantaneous dipoles.

• All molecules experience dispersion forces, but they are the only IMFs in nonpolar molecules.

• Strength increases with molecular size and shape.

• Example: Heavier noble gases and halogens have higher boiling points due to stronger dispersion forces.

Hydrogen Bonding

Definition and Occurrence

Hydrogen bonding is a special, very strong type of dipole-dipole interaction that occurs only in molecules containing hydrogen bonded to a small, highly electronegative atom (N, O, or F).

• Hydrogen bonds are responsible for the high boiling points of water, ammonia, and hydrogen fluoride.

• Hydrogen bonding is expected in molecules with –OH or –NH groups.

• Example: Water (H2O), methanol (CH3OH), and ammonia (NH3) all exhibit hydrogen bonding.

Effects of Hydrogen Bonding

• Hydrogen bonding leads to anomalously high boiling points for compounds containing N, O, or F compared to other group members.

• Having more than one –OH group increases the number of hydrogen bonds and raises the boiling point (e.g., 1-propanol vs. 1,2,3-propanetriol).

• Hydrogen bonding plays a crucial role in biological molecules, such as DNA base pairing (Thymine–Adenine: 2 hydrogen bonds; Cytosine–Guanine: 3 hydrogen bonds).

Comparing Intermolecular Forces

Relative Strengths and Characteristics

Force	Strength	Characteristics
Ion-dipole	Moderate (10–50 kJ/mol)	Occurs between ions and polar solvents
Dipole-dipole	Weak (3–4 kJ/mol)	Occurs between polar molecules
London dispersion	Weak (1–10 kJ/mol)	Occurs between all molecules; strength depends on size, polarizability
Hydrogen bond	Moderate (10–40 kJ/mol)	Occurs between molecules with O–H, N–H, and F–H bonds

Intermolecular Forces and Physical Properties

Boiling and Melting Points

The magnitude of intermolecular forces directly affects boiling and melting points. Stronger IMFs require more energy to separate particles, resulting in higher boiling and melting points.

• Boiling points of compounds with similar molecular masses increase with dipole moment (see Table 12.1).

• Halogens and noble gases show increasing boiling points with increasing molar mass due to stronger London dispersion forces (see Table 12.2).

• Alkanes: Boiling point increases with the number of carbon atoms due to increased London dispersion forces.

• Molecular shape affects boiling point: More compact molecules have less surface area for interactions and lower boiling points (e.g., pentane vs. 2,2-dimethylpropane).

Factors Affecting London Dispersion Forces

• Relative atomic weight of the atoms involved

• Number of atoms in the molecule

• Shape of the molecule

Application: Identifying Intermolecular Forces

Strategy for Determining IMFs

To identify the types of intermolecular forces present in a molecule:

• Draw the Lewis dot structure and apply VSEPR theory to determine molecular polarity.

• Nonpolar molecules exhibit only dispersion forces.

• Polar molecules exhibit dipole-dipole interactions and dispersion forces.

• Polar molecules with N–H, F–H, or O–H bonds exhibit dipole-dipole interactions (including hydrogen bonding) and dispersion forces.

Example:

• CCl4: Nonpolar, only dispersion forces.

• CH3COOH: Polar, contains O–H bond, exhibits hydrogen bonding, dipole-dipole, and dispersion forces.

• CH3COCH3: Polar, but no N–H, F–H, or O–H bonds, so only dipole-dipole and dispersion forces.

• H2S: Polar, but no N–H, F–H, or O–H bonds, so only dipole-dipole and dispersion forces.

Summary Table: Dipole Moments and Boiling Points

Compound	Structural Formula	Dipole Moment (D)	Boiling Point (°C)
Propane	CH3CH2CH3	0.1	–42
Dimethyl ether	CH3OCH3	1.3	–25
Methyl chloride	CH3Cl	1.9	–24
Acetaldehyde	CH3CHO	2.7	21
Acetonitrile	CH3CN	2.9	82

Additional info:

• Hydrogen bonding is responsible for the unique structure of ice and snowflakes, as well as the double helix structure of DNA.

• London dispersion forces are the only IMFs present in noble gases and nonpolar molecules.

• Understanding IMFs is essential for predicting solubility, boiling/melting points, and other physical properties of substances.