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Chemistry Unit 4: States of matter, gases and solutions

Intermolecular forces decide whether a substance is a solid, liquid or gas. The gas laws relate pressure, volume, temperature and amount.

Covers: Intermolecular forces, phase changes, gas laws, solutions and molarity, colligative properties

Study guide

Matter exists as solids, liquids and gases depending on the balance between particle motion and intermolecular forces (IMFs), the attractions between molecules. London dispersion forces exist in all molecules and grow with molecular size. Dipole-dipole forces act between polar molecules. Hydrogen bonds, the strongest IMF, form when H is bonded to N, O or F. Stronger IMFs mean higher boiling points, higher viscosity and lower vapor pressure.

Phase changes (melting, freezing, vaporization, condensation, sublimation, deposition) happen at constant temperature while energy is absorbed or released. A heating curve shows flat sections during phase changes, and a phase diagram shows which phase exists at each temperature and pressure.

Gases are described by pressure, volume, temperature (always in Kelvin, K = °C + 273) and moles. Boyle's law says P₁V₁ = P₂V₂ at constant temperature. Charles's law says V₁/T₁ = V₂/T₂ at constant pressure. Gay-Lussac's law relates pressure and temperature, and the combined gas law joins them all. The ideal gas law, PV = nRT, ties everything together, with R = 0.0821 L·atm/(mol·K). At STP, one mole of gas occupies about 22.4 L. Dalton's law says partial pressures add to the total pressure.

A solution is a solute dissolved in a solvent, and "like dissolves like." Molarity is moles of solute per liter of solution, and dilution follows M₁V₁ = M₂V₂. Colligative properties depend only on the number of solute particles: dissolving solute lowers the freezing point and raises the boiling point, which is why salt melts ice.

Ion-dipole forces form when an ion dissolves in a polar solvent: water's partial negative oxygen faces cations, and its partial positive hydrogens face anions. These forces are stronger for ions with a higher charge and a smaller radius, and they are why ionic compounds dissolve in water. During a phase change the temperature stays constant, so the particles' average kinetic energy doesn't change. Only their potential energy changes as intermolecular forces are broken or formed. The heat involved is q = m·ΔH_fus for melting and q = m·ΔH_vap for boiling (in J/g, or use moles with kJ/mol). For water, ΔH_fus is 334 J/g.

Kinetic molecular theory models an ideal gas as tiny particles in constant, random, straight-line motion. The particles take up no volume, don't attract each other, and collide elastically, and their average kinetic energy is proportional to the Kelvin temperature. Real gases behave least ideally at high pressure and low temperature. A saturated solution holds the maximum amount of solute at that temperature, an unsaturated one holds less, and a supersaturated one temporarily holds more. A solubility curve shows that maximum at each temperature, and most solids dissolve better in hotter water, while gases dissolve better in colder water. Parts per million = (grams of solute ÷ grams of solution) × 10⁶.

Beer's law says absorbance is directly proportional to concentration (A = εbc). To find an unknown concentration, measure standards of known concentration, plot absorbance against concentration to make a calibration line, and read the unknown's concentration from its absorbance.

Key ideas

• IMF strength: London dispersion < dipole-dipole < hydrogen bonding.
• Stronger IMFs mean higher boiling points.
• Gas laws need temperature in Kelvin.
• PV = nRT ties together all gas variables.
• Molarity = moles of solute per liter of solution.

Formulas and rules

Common mistakes

• Using °C in gas laws instead of K.
• Dividing by liters of solvent instead of liters of solution for molarity.

Words to know

Ideal gas law

PV = nRT

Kelvin conversion

K = °C + 273 (25°C = 298 K).

Hydrogen bonding

The strongest intermolecular force; H bonded to N, O or F.

Molarity

Moles of solute ÷ liters of solution.

Boyle's law

P₁V₁ = P₂V₂ at constant temperature.

Intermolecular forces

Attractions between molecules.

London dispersion forces

Weak attractions in all molecules; stronger with size.

Dipole-dipole forces

Attractions between polar molecules.

Vapor pressure

Pressure of a vapor above its liquid; lower with stronger IMFs.

Phase diagram

Shows which phase exists at each temperature and pressure.

Charles's law

V₁/T₁ = V₂/T₂ at constant pressure.

STP molar volume

1 mol of gas ≈ 22.4 L at STP.

Dalton's law

Partial pressures add to the total pressure.

Dilution equation

M₁V₁ = M₂V₂

Colligative properties

Depend only on the number of solute particles (freezing-point depression, boiling-point elevation).

Sample practice questions

27 more questions like these, from level 1 to 5, are in the practice set.