Unit 3: Intermolecular Forces (IMF) and Properties (Simple Studies AP Chem Study Guide)

Unit 3: Intermolecular Forces (IMF) and Properties

Vocabulary

• Amorphous Solids

  • Particles lack a repeating lattice form (pseudo solids).

  • No set melting point; gradual melting occurs over a temperature range.

• Barometer

  • Instrument used to measure atmospheric pressure.

• Boiling

  • Formation of bubbles as liquid particles vaporize inside.

  • Phase Change: Liquid to gas.

• Condensation

  • Change of a substance from gas to liquid state.

• Deposition

  • Direct change from gas to solid without becoming a liquid.

• Equilibrium

  • Balance reached between evaporation and condensation of trapped molecules.

• Evaporation

  • Surface molecules gain energy, overcoming intermolecular forces (IMF).

  • Change occurs only from the liquid's surface.

  • Phase Change: Liquid to gas.

• Ionic Solids

  • Crystals formed from cations and anions, making ionic bonds difficult to overcome.

  • Characteristics: hard, brittle, non-conductive, typically soluble in water.

• Intermolecular Attractions

  • Forces that determine the attraction between molecules.

• Kinetic Molecular Theory

  • States that all particles of matter are in constant motion.

• Melting Point

  • Temperature where solid turns to liquid; equals freezing point.

  • Higher IMF correlates with higher melting points.

• Metallic Solids

  • Electrons are free-moving in overlapping electron clouds.

  • Properties: high conductivity of heat/electricity, high melting points, malleable, ductile.

• Molecular Solids

  • Formed by covalently bonded molecules held by electrostatic forces.

  • Often nonpolar and insoluble in water, but soluble in nonpolar solvents.

Phase Changes

• Network Covalent

  • Atoms are covalently bonded in a continuous network.

  • Characterized by extreme hardness, high melting points, non-solubility in water, and non-conductivity.

• Phase Change

  • Rearrangement of molecules while maintaining constant temperature.

• Pressure

  • Measure of the force exerted by gas particles upon collision with container walls.

• Temperature

  • Represents the average kinetic energy of particles.

• Temperature Change

  • Variation in molecular motion, dependent upon heat capacity.

• Sublimation

  • Occurs when solid converts to vapor without passing liquid phase due to vapor pressure matching external pressure.

• Standard Atmosphere (atm)

  • Defined as 760 mm Hg.

• Standard Temperature

  • Defined as 0°C or 273 K.

• Standard Pressure

  • Defined as 1 atm or 101.325 kPa.

• Surface Tension

  • Tendency of liquids to minimize surface area due to molecular interactions below the surface.

• Volatility

  • Measure of a substance's evaporation rate, affected by temperature and IMF.

Types of Intermolecular Forces

• London Dispersion Forces

  • Weak attractions caused by transient polarity.

  • Stronger in heavier molecules.

• Dipole-Dipole Attractions

  • Exist between polar molecules, attracting oppositely charged ends.

• Hydrogen Bonds

  • Strong dipole interactions when hydrogen bonds to F, O, or N.

  • Hydrogen’s electrons are pulled away, leaving a positively charged nucleus that attracts lone pairs in adjacent molecules.

Dissolving Process

1. Break apart solute (ΔH1: Positive)

2. Break apart solvent (ΔH2: Positive)

3. Solute-solvent interaction (ΔH3: Usually negative)

   • ΔH solution = ΔH1 + ΔH2 + ΔH3

   • If ΔH solution is negative and substantial, solute dissolves.

   • If ΔH solution is positive and small, solute does not dissolve.

   • If ΔH solution is small (±), entropy increase favors dissolution.

• ΔH1 is always opposite of lattice energy and always positive.

• ΔH2 and ΔH3 are known as enthalpy of hydration.

States of Matter

• Gases

  • Take shape and volume of their container; can expand to fill.

  • Exhibit very low density; compressible.

  • Undergo diffusion and effusion.

• Ideal Gases

  • Hypothetical; fit all assumptions of kinetic molecular theory.

  • Have significant empty space and particle collisions are elastic (no energy loss).

  • Gas particles are in constant random motion without attractions.

  • Gas temperature depends on average kinetic energy.

• Real Gases

  • Have volume and attract each other; behavior is more ideal at low pressures and high temperatures.

Characteristics of States

• Solids

  • Fixed shape and volume; very strong IMF.

  • Non-fluid, high density, not compressible, rigid structure.

• Liquids

  • Take shape of the container, not volume; weaker IMF than solids.

  • Fluids; not compressible and diffuse slower than gases but faster than solids.

• Vapor Pressure

  • Expressed in atm; affected by temperature, container size, and molecular bonds.

  • Indicates extent of evaporation; increases with temperature.

• Boiling Point

  • Temperature at which vapor pressure equals external pressure; depends on atmospheric pressure and IMF.

Gas Laws

• Ideal Gas Law

  • PV = nRT.

• Dalton’s Law of Partial Pressures

  • Total pressure equals the sum of individual pressures of mixed gases.

Reference

This note is from https://www.simplestudies.org/groups/ap-chemistry