Liquids, Solids, and Intermolecular Forces (Ch. 11): Study Notes

Liquids, Solids, and Intermolecular Forces (Ch. 11)

11.2 Solids, Liquids, and Gases: A Molecular Comparison

The three states of matter—solid, liquid, and gas—differ in the arrangement and movement of their particles, which is determined by the balance between kinetic energy and intermolecular forces.

• Solids: High density, definite shape and volume, strong intermolecular forces.

• Liquids: High density, indefinite shape, definite volume, moderate intermolecular forces.

• Gases: Low density, indefinite shape and volume, weak intermolecular forces.

State	Density	Shape	Volume	Strength of IMF
Gas	Low	Indefinite	Indefinite	Weak
Liquid	High	Indefinite	Definite	Moderate
Solid	High	Definite	Definite	Strong

Key Point: The physical state of a substance is a balance between the kinetic energy (which moves particles apart) and intermolecular forces (which hold particles together).

11.3 Intermolecular Forces

Intermolecular forces (IMF) are attractive forces between molecules. They are much weaker than covalent bonds and are responsible for many physical properties of substances, such as boiling and melting points.

• Boiling point is an indicator of IMF strength: higher boiling point means stronger IMF.

There are four main types of intermolecular forces, collectively called van der Waals forces:

1. London dispersion forces (LDF): Present in all molecules (polar and nonpolar). Caused by instantaneous dipoles due to electron movement. Strength increases with polarizability (ease of electron cloud distortion) and molecular size/shape.

2. Dipole-dipole forces: Occur between polar molecules. The positive end of one molecule is attracted to the negative end of another. Strength increases with the magnitude of the molecular dipole.

3. Hydrogen bonding: A special, strong type of dipole-dipole force. Occurs when hydrogen is covalently bonded to N, O, or F and interacts with N, O, or F on another molecule.

4. Ion-dipole forces: Exist between an ion and a polar molecule. Important for dissolving ionic compounds in polar solvents (e.g., NaCl in water).

Relative Strengths: Ion-dipole > Hydrogen bond > Dipole-dipole > London dispersion

London Dispersion Forces (LDF)

• Present in all molecules and atoms.

• Result from instantaneous and induced dipoles.

• Strength increases with:

  • Number of electrons (higher molar mass = stronger LDF)

  • Molecular surface area (more area = stronger LDF)

• Example: Rank Br2, F2, Cl2 by LDF strength: F2 < Cl2 < Br2 (increasing molar mass)

Application: Geckos stick to surfaces due to LDF between their feet and the wall.

Dipole-Dipole Forces

• Occur between polar molecules (molecules with a permanent dipole).

• The larger the dipole moment, the stronger the force.

• Example: Methanal (CH2O) has a higher boiling point than ethene (C2H4) due to stronger dipole-dipole forces.

Hydrogen Bonding

• Occurs when H is bonded to N, O, or F and interacts with N, O, or F on another molecule.

• Much stronger than regular dipole-dipole forces.

• Donor: N, O, or F covalently bonded to H.

• Acceptor: Electronegative atom (N, O, F) not covalently bonded to H.

• Biological Importance: Hydrogen bonding is crucial in protein structure and DNA base-pair binding.

• Example: Ethanol (C2H5OH) has a higher boiling point than dimethyl ether (CH3OCH3) due to hydrogen bonding.

Ion-Dipole Forces

• Occur between an ion and a polar molecule.

• Responsible for the dissolution of ionic compounds in water.

Summary Table: Types of Intermolecular Forces

Type	Occurs Between	Relative Strength	Example
London Dispersion	All molecules/atoms	Weakest	Noble gases, nonpolar molecules
Dipole-Dipole	Polar molecules	Intermediate	CH2O
Hydrogen Bond	H bonded to N, O, or F	Strong	H2O, NH3
Ion-Dipole	Ion + polar molecule	Strongest	Na+ in H2O

11.5 Intermolecular Forces in Action

Intermolecular forces influence many observable properties of liquids.

• Surface tension: Energy required to increase the surface area of a liquid. Stronger IMFs = higher surface tension.

• Viscosity: Resistance to flow. Increases with stronger IMFs and decreases with higher temperature.

• Capillary action: Ability of a liquid to flow in narrow spaces against gravity. Occurs when adhesive forces (liquid to surface) are greater than cohesive forces (liquid to itself).

Substance	Formula	Viscosity (mPa·s)
Hexane	CH3(CH2)4CH3	0.300
Heptane	CH3(CH2)5CH3	0.387
Octane	CH3(CH2)6CH3	0.508
Nonane	CH3(CH2)7CH3	0.665
Decane	CH3(CH2)8CH3	0.838

Example: Water forms a concave meniscus in glass (adhesion > cohesion), while mercury forms a convex meniscus (cohesion > adhesion).

5.2 Pressure: The Result of Molecular Collisions

Pressure is the amount of force applied per unit area.

• Formula: $p = \frac{F}{A}$

• Gases exert pressure on surfaces due to molecular collisions.

• The atmosphere exerts pressure on everything at Earth's surface.

11.6 – 11.8 Phase Changes

Phase Transitions

• A phase transition occurs when a substance changes from one state to another (e.g., solid to liquid).

• Always involves a change in energy.

• The energy required for a phase change is independent of the path taken.

Vaporization and Vapour Pressure

• Vaporization: Liquid to gas transition; molecules at the surface gain enough energy to escape IMF.

• Condensation: Gas to liquid transition.

• Vapour pressure: Pressure of gas above a liquid at equilibrium.

• Vapour pressure increases with temperature.

• Normal boiling point: Temperature at which vapour pressure equals atmospheric pressure.

Sublimation and Fusion

• Sublimation: Solid to gas transition (e.g., dry ice).

• Deposition: Gas to solid transition.

• Fusion (melting): Solid to liquid transition.

• Freezing: Liquid to solid transition.

Heating Curves

• During a phase change, temperature remains constant while energy is used to break IMFs.

• It takes less energy to melt ice (solid to liquid) than to vaporize water (liquid to gas) because IMFs are stronger in the liquid-gas transition.

11.9 Phase Diagrams

Phase diagrams show the state of a substance at various pressures and temperatures, and the boundaries where phase transitions occur.

• Triple point: Unique pressure and temperature where solid, liquid, and gas coexist in equilibrium.

• Critical point: Beyond this, the substance becomes a supercritical fluid (properties of both liquid and gas).

• Supercritical fluid: State with properties of both a liquid and a gas.

Example: The triple point of water is at $T_3 = 0.01\ ^\circ\mathrm{C}$ and $p_3 = 0.006\ \mathrm{atm}$.

11.10 – 11.12 Crystalline Solids

Solids can be classified based on their structure and bonding.

• Amorphous solids: Lack long-range order (e.g., glass).

• Crystalline solids: Have regular, repeating patterns.

Type	Components	Bonding	Examples
Ionic	Cations & anions	Electrostatic	NaCl
Metallic	Metal atoms	Metallic bonds	Fe, Cu
Network covalent	Atoms	Covalent bonds	Diamond, SiO2
Molecular	Molecules	IMFs	Ice, CO2

• Unit cell: The simplest repeating unit in a crystal lattice.

• Lattice points: Positions in the crystal occupied by atoms, ions, or molecules.

• Coordination number: Number of nearest neighbors to a particle in the crystal.

Types of cubic unit cells:

• Simple cubic: Coordination number 6, 1 atom per unit cell.

• Body-centered cubic (bcc): Coordination number 8, 2 atoms per unit cell.

• Face-centered cubic (fcc): Coordination number 12, 4 atoms per unit cell.

10.8 Band Theory of Solids

Band theory extends molecular orbital theory to solids, explaining properties such as conductivity and color.

• In solids, atomic orbitals combine to form bands of energy levels.

• Valence band: Highest occupied band.

• Conduction band: Lowest unoccupied band.

• The size of the band gap determines if a material is a conductor, semiconductor, or insulator.

• Semiconductors can absorb photons to excite electrons from the valence to the conduction band, generating current (basis for LEDs).