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Chemistry 2 Unit 1: Intermolecular forces and solutions

Intermolecular forces explain boiling points, solubility and other physical properties, and solutes change a solvent's properties.

Covers: Types of intermolecular forces, properties of liquids and solids, phase diagrams, solutions and solubility, concentration units, colligative properties

Study guide

Intermolecular forces (IMFs) are attractions between molecules, much weaker than covalent bonds within them. London dispersion forces exist in all molecules and come from temporary dipoles; they grow stronger with more electrons and larger, more polarizable molecules. Dipole-dipole forces attract polar molecules. Hydrogen bonding is an especially strong dipole-dipole force when H is bonded to N, O or F and attracted to a lone pair on another N, O or F. Ion-dipole forces occur between ions and polar molecules, as when salt dissolves in water. Stronger IMFs lead to higher boiling points, higher melting points, higher viscosity, higher surface tension and lower vapor pressure.

Solids can be ionic, molecular, covalent network (like diamond) or metallic, each with characteristic properties. A phase diagram shows the phase of a substance at each temperature and pressure, including the triple point and critical point. Heating curves show that temperature stays constant during phase changes, using heat of fusion and vaporization. A liquid boils when its vapor pressure equals the external pressure.

In solutions, 'like dissolves like': polar and ionic solutes dissolve in polar solvents, and nonpolar solutes dissolve in nonpolar solvents. Gas solubility increases with pressure (Henry's law) and decreases with temperature, while most solids become more soluble at higher temperatures. Concentration can be expressed as molarity (mol solute / L solution), molality (mol solute / kg solvent), mole fraction and mass percent. Colligative properties depend on the number of solute particles, not their identity: vapor pressure lowering, boiling point elevation (ΔTb = i·Kb·m), freezing point depression (ΔTf = i·Kf·m) and osmotic pressure (Π = iMRT). The van 't Hoff factor i is the number of particles a solute forms; NaCl gives about 2. Chromatography and distillation separate mixtures based on IMFs.

For an ideal gas, PV = nRT (R = 0.08206 L·atm/(mol·K)), with temperature in kelvins. At constant volume, pressure is proportional to temperature. In a gas mixture, each gas's partial pressure equals its mole fraction times the total pressure (Dalton's law). Real gases deviate from ideal behavior at high pressure and low temperature, when molecular volume and attractions matter. A Maxwell–Boltzmann distribution shows how many particles have each speed or kinetic energy. At a higher temperature the curve flattens and shifts toward higher speeds, while the area under it stays the same. Lighter gases have broader distributions with higher average speeds at the same temperature.

By Coulomb's law, the attraction between charges grows with the size of the charges and weakens as the distance between them grows. So smaller ions and more highly charged ions, like Mg²⁺ compared with Na⁺, attract water molecules more strongly and form stronger ionic lattices. In ion–dipole interactions, water's partially positive hydrogens point toward anions and its partially negative oxygen points toward cations. Among molecules with similar molar mass, long, unbranched shapes have more surface contact than compact, branched ones. So they have stronger London dispersion forces and higher boiling points, as n-pentane boils higher than neopentane.

In paper or thin-layer chromatography, a compound's Rf is the distance it travels divided by the distance the solvent front travels. Compounds more attracted to the mobile phase than the stationary phase travel farther. To make a solution of known molarity, dissolve a measured mass of solute in some solvent in a volumetric flask, then add solvent to the calibration mark and mix. Absorbance follows the Beer–Lambert law, A = εbc. At a fixed wavelength and path length, absorbance is directly proportional to concentration, so a calibration line of absorbance versus concentration can be used to find an unknown concentration.

Key ideas

• London dispersion forces exist in all molecules and grow with size.
• Hydrogen bonding requires H bonded to N, O or F.
• Stronger IMFs mean higher boiling points and lower vapor pressure.
• Like dissolves like.
• Colligative properties depend on the number of solute particles.

Formulas and rules

Common mistakes

• Thinking hydrogen bonds are bonds within a molecule; they're attractions between molecules.
• Forgetting the van 't Hoff factor for ionic solutes.

Words to know

Intermolecular forces

Attractions between molecules.

London dispersion forces

Weak attractions from temporary dipoles, present in all molecules.

Polarizability

How easily a molecule's electron cloud is distorted.

Dipole-dipole force

Attraction between polar molecules.

Hydrogen bonding

Strong attraction involving H bonded to N, O or F.

Ion-dipole force

Attraction between an ion and a polar molecule.

Vapor pressure

Pressure of a vapor in equilibrium with its liquid.

Viscosity

A liquid's resistance to flow.

Surface tension

Energy needed to increase a liquid's surface area.

Covalent network solid

A solid held together by a network of covalent bonds.

Phase diagram

A graph of phase versus temperature and pressure.

Triple point

Conditions where solid, liquid and gas coexist.

Critical point

Conditions beyond which liquid and gas are indistinguishable.

Heat of vaporization

Energy needed to vaporize a substance.

Henry's law

Gas solubility is proportional to its partial pressure.

Molality

Moles of solute per kilogram of solvent.

Colligative property

A property that depends on the number of solute particles.

Van 't Hoff factor

The number of particles a solute forms in solution.

Osmotic pressure

Pressure needed to stop osmosis.

Sample practice questions

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