Topic 3: Intermolecular Forces and Properties

The forces between particles, not just within them, decide whether a substance is a gas, liquid, or solid and how it behaves. This unit connects molecular-scale attractions to measurable properties such as boiling point, vapor pressure, and gas behavior. It also covers how we describe gases mathematically, prepare and quantify solutions, separate mixtures, and probe matter with light.

Types of Intermolecular Forces (IMFs)

Intermolecular forces are attractions between separate particles, and they are much weaker than the covalent or ionic bonds inside particles. London dispersion forces (LDFs) arise from temporary, instantaneous dipoles created by shifting electron clouds; they exist between all particles and grow stronger with more electrons and larger, more polarizable surfaces. Dipole-dipole forces act between permanent dipoles in polar molecules, lining up partial positive ends with partial negative ends. Hydrogen bonding is an especially strong dipole interaction that occurs when H is bonded directly to N, O, or F and is attracted to a lone pair on N, O, or F of a neighboring molecule. Ion-dipole forces, important in solutions, occur between an ion and a polar molecule such as water.

IMF Trends and Physical Properties

Stronger IMFs raise boiling point, melting point, viscosity, and surface tension while lowering vapor pressure and volatility. To compare two substances, identify the strongest IMF each can use, then if they are tied compare molar mass and shape. Worked example: rank CH4, NH3, and Ne by boiling point. CH4 is nonpolar (LDF only, 16 g/mol), Ne is a single atom (LDF only, 20 g/mol), and NH3 has hydrogen bonding. So NH3 boils highest, then Ne (slightly more electrons than CH4), then CH4. Remember that long straight-chain molecules have more contact area and stronger LDFs than compact branched isomers of equal mass.

Solids, Liquids, and Gases

In gases, particles are far apart with negligible attraction, move rapidly, and fill any container, giving high compressibility. Liquids have particles in contact but free to move past one another, producing a definite volume but no fixed shape. Solids hold particles in fixed positions, giving definite shape and volume and very low compressibility. Solid types differ in their attractions: ionic solids (strong electrostatic, high melting, brittle), covalent network solids like diamond (very high melting), metallic solids (mobile electrons, conductive, malleable), and molecular solids held by IMFs (low melting, soft). The dominant attraction explains the observed macroscopic property.

The Ideal Gas Law and Partial Pressures

The ideal gas law is PV = nRT, where P is pressure, V is volume, n is moles, T is absolute temperature in kelvin, and R is 0.08206 L atm per mol K. Worked example: find the volume of 2.0 mol of gas at 1.5 atm and 300 K. V = nRT/P = (2.0)(0.08206)(300)/1.5 = 32.8 L. For a mixture, Dalton's law states the total pressure equals the sum of partial pressures: P_total = P1 + P2 + ... Each gas's partial pressure equals its mole fraction times the total: P_i = X_i times P_total. Always convert temperature to kelvin and keep units consistent with the chosen value of R.

Kinetic Molecular Theory and Maxwell-Boltzmann

Kinetic molecular theory (KMT) models an ideal gas as many tiny particles in constant random motion, with negligible volume, no IMFs, and perfectly elastic collisions. The average kinetic energy of the particles depends only on absolute temperature, so at the same temperature all gases have the same average KE regardless of mass. Lighter molecules therefore move faster on average. A Maxwell-Boltzmann diagram plots the distribution of molecular speeds: raising temperature shifts the peak to higher speed and flattens and broadens the curve, while heavier gases at the same temperature show a taller, narrower distribution centered at lower speed. The area under any curve stays constant because the number of molecules is fixed.

Deviations from Ideal Behavior

Real gases deviate from PV = nRT most at high pressure and low temperature. At high pressure the actual particle volume is no longer negligible compared to the container, so the measured volume is larger than ideal predicts. At low temperature, particles move slowly enough that IMFs pull them together, reducing the measured pressure below ideal. Gases with strong IMFs or large size (for example NH3 or larger hydrocarbons) deviate more than small nonpolar gases like helium. The closer the conditions are to high temperature and low pressure, the more accurately a gas obeys the ideal gas law.

Solutions and Molarity

A solution is a homogeneous mixture of a solute dissolved in a solvent. Dissolving requires that solute-solvent attractions be comparable to the attractions broken apart, summarized as like dissolves like: polar and ionic solutes dissolve in polar solvents, nonpolar in nonpolar. Concentration is most often expressed as molarity, M = moles of solute per liter of solution. Worked example: dissolving 0.50 mol NaCl in enough water to make 250 mL gives M = 0.50 / 0.250 = 2.0 M. Dilution follows M1V1 = M2V2; for instance, diluting 50 mL of 2.0 M stock to 200 mL gives (2.0)(50)/200 = 0.50 M.

Separation Techniques

Mixtures are separated by exploiting differences in physical properties. Filtration separates an insoluble solid from a liquid using a barrier with small pores. Distillation separates liquids by differences in boiling point, vaporizing the more volatile component first and condensing it. Chromatography separates components by differences in their attraction to a stationary phase versus a mobile phase; components that interact more strongly with the mobile phase travel farther. Paper chromatography reports a retention factor, Rf = distance traveled by the spot divided by distance traveled by the solvent front, a value between 0 and 1 used to identify components.

Spectroscopy and the Beer-Lambert Law

Spectroscopy uses the interaction of light with matter to identify and quantify substances. Different regions of light probe different motions: microwave excites rotation, infrared excites bond vibration (useful for functional groups), and ultraviolet-visible excites electron transitions (useful for colored species). The Beer-Lambert law relates absorbance to concentration: A = (epsilon)(b)(c), where epsilon is molar absorptivity, b is path length, and c is concentration. Because A is directly proportional to c, a calibration curve of absorbance versus known concentrations lets you read off the concentration of an unknown. Worked idea: if doubling a solution's concentration doubles its measured absorbance, the sample is obeying Beer's law.