Properties of Substances and Mixtures (AP Exams) — Complete Study Guide

Intermolecular & Interparticle Forces (IMFs)

1. London Dispersion Forces (LDFs)

• The weakest type of IMF exists between all atoms and molecules.

• Caused by temporary (instantaneous) dipoles when electron clouds shift momentarily.

• Strength increases with:

  • Larger molar mass (more electrons = more polarizable).

  • Greater surface area (more contact = stronger attractions).

• Important in nonpolar substances (e.g., noble gases, N₂, O₂, hydrocarbons).

• Example: I₂(s) has stronger LDFs than F₂(g) due to more electrons and larger molecular size.

2. Dipole-Induced Dipole Interactions

• Occur when a polar molecule induces a dipole in a nonpolar molecule by distorting its electron cloud.

• Strength depends on:

  • Polarity of the polar molecule.

  • Polarizability of the nonpolar molecule.

• Example: O₂ (nonpolar) dissolving in H₂O (polar)—temporary attraction forms.

3. Dipole–Dipole Interactions

• Occur between two polar molecules with permanent dipoles.

• Molecules align positive end to negative end.

• Stronger than dispersion forces but weaker than hydrogen bonding.

• Greater dipole moment = stronger interactions.

• Example: HCl molecules attract through dipole–dipole forces.

4. Ion–Dipole Interactions

• Occur between an ion and a polar molecule.

• Very strong IMF; key in dissolving ionic compounds in water.

• Example: Na⁺ and Cl⁻ ions surrounded by water molecules in an aqueous solution.

• Cations attract oxygen (δ–) and anions attract hydrogen (δ+).

5. Molecular Dipole Moment

• A measure of net polarity in a molecule; vector sum of all bond dipoles.

• Molecules can have polar bonds but no dipole moment if symmetrical (e.g., CO₂, CCl₄).

• Measured in Debye units (D).

• Polar molecules exhibit stronger IMFs, higher boiling/melting points, and greater solubility in polar solvents.

6. Hydrogen Bonding

• Special case of strong dipole–dipole interaction.

• Occurs when H is bonded to N, O, or F and attracted to a lone pair on another N, O, or F atom.

• Responsible for:

  • High boiling points of water and alcohol.

  • Ice’s lower density (hydrogen bonds form an open lattice).

  • DNA base pairing (A–T, G–C hydrogen bonds).

• Example: H₂O, NH₃, HF.

7. Interactions in Large Biomolecules

• Biological macromolecules (proteins, DNA) rely on multiple weak IMFs:

  • Hydrogen bonding stabilizes secondary structures (α-helices, β-sheets).

  • Van der Waals (dispersion) forces help protein folding.

  • Ion–dipole and dipole–dipole interactions maintain tertiary and quaternary structures.

• Weak forces collectively create strong structural stability.

Properties of Solids, Liquids, & Gases

1. Solids & Liquids

• Determined by IMF strength:

  • Stronger IMFs → higher melting/boiling points.

  • Weak IMFs → more fluid and volatile substances.

• Solids

  • Properties:

    • Fixed shape and volume

    • Strong particle attractions

• Liquids

  • Properties:

    • Fixed volume

    • Variable shape

    • Moderate IMFs.

2. Ionic Solids

• Composed of positive and negative ions in a crystal lattice.

• Strong ionic bonds 

  • Properties:

    • Have high melting point

    • Hard

    • Brittle

• Conductivity:

  • Solid: nonconductive (ions fixed in lattice).

  • Molten/dissolved: conductive (ions free to move).

• Example: NaCl, MgO.

3. Covalent Network Solids

• Atoms held by covalent bonds in a continuous network.

  • Properties:

    • Very hard

    • Have high melting points

    • Poor conductors

• Examples:

  • Diamond (C) — tetrahedral bonding.

  • Quartz (SiO₂) — 3D silicon–oxygen framework.

  • Graphite (C) — sheets of covalent layers (conducts electricity due to delocalized electrons).

4. Molecular Solids

• Made of discrete molecules held by IMFs (not covalent or ionic bonds).

  • Properties:

    • Low melting points

    • Soft

    • Poor electrical conductors

• Examples: CO₂ (dry ice), H₂O (ice), C₁₂H₂₂O₁₁ (sucrose).

5. Metallic Solids

• Consist of metal cations surrounded by a sea of delocalized electrons.

  • Properties:

    • Malleable

    • Ductile

    • Electrically/Thermally Conductive

• Strength varies with metal ion charge and electron density.

• Examples: Fe, Cu, Ag, Na.

6. Noncovalent Interactions in Large Molecules

• Includes:

  • Dispersion

  • Dipole–dipole

  • Hydrogen bonding

  • Ion–dipole

• Dictate biological recognition (enzyme–substrate binding, DNA pairing).

• Weaker individually but strong in large numbers.

7. The Structure of Solids

• Described by crystal lattice patterns:

  • Simple cubic

  • Body-centered cubic (BCC)

  • Face-centered cubic (FCC)

• Unit cell: smallest repeating unit.

• Affects density, melting point, and stability.

8. The Liquid Phase

• Particles close together but can move freely.

• Surface tension and viscosity depend on IMF strength.

• Evaporation rate inversely proportional to IMF strength.

9. The Gas Phase

• Particles far apart; negligible attraction.

• Compressible and expandable.

• Follows gas laws under ideal conditions (low pressure, high temperature).

Gas Laws

1. The Ideal Gas Law

• PV = nRT

  • P = pressure (atmospheres, atm)

  • V = volume (liters, L)

  • n = moles (moles, mol)

  • R = gas constant = 0.0821 L·atm/mol·K

  • T = temperature (kelvin, K).

• Applies best to gases with weak IMFs and low density.

• Derivatives:

  • Boyle’s Law: P ∝ 1/V (constant T, n).

  • Charles’s Law: V ∝ T (constant P, n).

  • Avogadro’s Law: V ∝ n (constant P, T).

2. Partial Pressure & Total Pressure

• Dalton’s Law: Pₜₒₜ = P₁ + P₂ + P₃ + …

• Meaning: 

  • The total pressure of a mixture of gases is equal to the sum of the partial pressures of all the individual gases in the mixture.

  • Each gas exerts pressure proportional to its mole fraction.

  • Used for gas collection over water and mixture analysis.

3. Graphical Representations of Gas Laws

• Boyle’s Law

• Pressure (P) vs. Volume (V)

  • Shape: A curve (hyperbola) that slopes downward from left to right.

  • Meaning:

    • When volume decreases, pressure increases.

    • When volume increases, pressure decreases.

  • Explanation: The gas particles have less space to move when volume is smaller, so they hit the walls of the container more often creating higher pressure.

  • Example: If you compress a syringe (decrease volume), you feel more resistance (increased pressure).

• Pressure (P) vs. 1/Volume (1/V)

  • Shape: A straight line passing through the origin.

  • Meaning:

    • As 1/V increases, P increases proportionally.

  • Explanation: This linear relationship confirms the inverse proportionality between pressure and volume.

• Charles’s Law

• Volume (V) vs. Temperature (T)

  • Shape: A straight line that slopes upward from left to right.

  • Meaning:

    • When temperature increases, volume increases.

    • When temperature decreases, volume decreases.

  • Explanation: Heating a gas causes its particles to move faster and push outward, expanding the volume if the pressure remains constant.

  • Example: A balloon expands when heated because the gas molecules inside move faster and occupy more space.

• Avogadro’s Law

• Volume (V) vs. Moles (n)

  • Shape: A straight line that slopes upward from left to right.

  • Meaning:

    • When the number of moles increases, the volume increases.

    • When the number of moles decreases, the volume decreases.

  • Explanation: Adding more gas particles (moles) increases the number of collisions with the container walls, so the gas expands to maintain constant pressure and temperature.

  • Example: Filling a balloon with more air makes it larger because more gas particles occupy more space.

4. Kinetic Molecular Theory (KMT)

• Gas particles are in constant random motion.

• Volume of individual molecules ≈ negligible.

• Collisions are elastic (no energy loss).

• Average kinetic energy ∝ temperature (Kelvin).

• Explains pressure as particle collisions with container walls.

5. The Maxwell–Boltzmann Distribution

• Describes range of molecular speeds in a gas sample.

• At higher temperatures:

  • Distribution broadens.

  • Average kinetic energy increases.

• Lighter molecules move faster on average.

6. Non-Ideal Behavior of Gases

• Deviate from ideal gas law at high pressures or low temperatures.

• Attractive forces lower pressure; finite volume reduces free space.

• Corrected by van der Waals equation:
  [(P + a(n/V)^2)(V - nb) = nRT]
  where a and b correct for attractions and volume, respectively.

Solutions & Mixtures

1. Calculations About Solutions

• Molarity (M): M = moles solute / liters solution.

• Molality (m): m = moles solute / kg solvent.

• Percent composition, mole fraction, ppm/ppb used in concentration analysis.

2. Homogeneous & Heterogeneous Mixtures

• Homogeneous: uniform composition (solutions).

• Heterogeneous: non-uniform, distinct phases (suspensions, emulsions).

• Example: air = homogeneous; sand in water = heterogeneous.

3. Molarity

• Used for stoichiometric calculations.

• Dilution formula: M₁V₁ = M₂V₂.

• Important for titrations and reaction concentration control.

4. Particulate Representations of Solutions

• Diagrams depict:

  • Solvent–solute interactions (ion–dipole, hydrogen bonding).

  • Ion dissociation in ionic solutes.

  • Molecular solutes staying intact but evenly distributed.

5. Separation by Chromatography

• Separates components based on affinity for mobile and stationary phases.

• More polar substances interact strongly with polar stationary phase → move slower.

• Used for pigment separation, purity tests, forensic analysis.

6. Separation by Distillation

• Relies on differences in boiling points.

• Lower boiling point component vaporizes first.

• Used to purify liquids or separate miscible liquids.

7. Solubility of Ionic & Molecular Compounds

• “Like dissolves like”: polar dissolves polar; nonpolar dissolves nonpolar.

• Ionic compounds dissolve via ion–dipole interactions.

• Molecular solutes dissolve via dipole–dipole or hydrogen bonding.

• Solubility depends on temperature and IMF balance.

Spectroscopy

1. The Electromagnetic Spectrum

• Ordered by wavelength (λ) and frequency (ν):

  • Radio → Microwave → Infrared → Visible → UV → X-ray → Gamma.

• Energy related by: E = hν = hc/λ.

• Shorter wavelength → higher energy.

2. Transitions Associated with Radiation

• Microwave: molecular rotations.

• Infrared (IR): bond vibrations.

• Visible/UV: electronic transitions.

• X-rays: inner electron transitions or ejection (photoelectron spectroscopy).

3. Properties of Photons

• Photon energy depends on frequency: E = hν.

• Photoelectric effect: light of sufficient energy ejects electrons from metal surface.

• Demonstrates quantized nature of light.

4. Beer–Lambert Law

• Relates absorbance (A) to concentration (c) and path length (l):
  A = εlc, where ε = molar absorptivity.

• Used in UV–Vis spectroscopy to determine solution concentration.

• Higher concentration → higher absorbance (linear relationship).

✅ AP Exam Focus Tips

• Be able to compare IMF strength and predict relative boiling/melting points.

• Explain why certain substances dissolve or remain insoluble using molecular reasoning.

• Interpret phase diagrams and gas law graphs.

• Perform stoichiometric and molarity calculations accurately.

• Describe spectroscopic evidence for molecular structure or concentration.

• Relate microscopic behavior (particle level) to macroscopic properties (observable traits).