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Unit 3 · Properties of Substances & Mixtures Unit Hub Flashcards Cheat Sheet Essentials Visual Review MC Practice SAQ Practice

AP Chemistry Unit 3 Visual Review

A topic-by-topic visual walkthrough of Properties of Substances & Mixtures — intermolecular forces, gases, solutions, and spectroscopy.

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TOPIC 3.1 Intermolecular Forces (IMFs) RANKED WEAKEST → STRONGEST (all much weaker than bonds) London dispersion Temporary dipoles from shifting electrons. In ALL molecules; stronger with more electrons (bigger). Only force in nonpolar molecules. Dipole–dipole Between POLAR molecules — the δ+ of one attracts the δ− of another. Requires a permanent dipole (polar molecule). Hydrogen bonding A strong dipole–dipole when H is bonded to N, O, or F. Explains water's high boiling point. "N-O-F" is the memory trick. Ion–dipole Between an ION and a polar molecule. The strongest IMF. Drives dissolving of salts in water. e.g., Na⁺ surrounded by H₂O. Why IMFs matter IMFs are attractions BETWEEN molecules (not the bonds within them). Stronger IMFs mean molecules stick together more, which raises boiling point, melting point, viscosity, and surface tension. Comparing boiling points? Compare IMF strength. e.g., H₂O boils higher than H₂S because water H-bonds; larger nonpolar molecules boil higher due to more LDFs. Order of strength (general): ion–dipole > hydrogen bonding > dipole–dipole > London dispersion. Stronger IMFs → higher boiling/melting points. Hydrogen bonding needs H bonded to N, O, or F. The Review Hub · AP Chemistry Unit 3 TOPIC 3.2 Properties of Solids FOUR TYPES OF SOLID — PROPERTIES FOLLOW THE PARTICLES & FORCES Molecular solids Molecules held by IMFs (ice, dry ice, sugar). LOW melting points, soft, don't conduct. Ionic solids Ion lattice, Coulombic forces (NaCl). High mp, brittle, conduct when molten/aqueous. Covalent network solids Atoms bonded in a giant network (diamond, quartz, SiO₂). VERY high mp, extremely hard, don't conduct. Metallic solids Cations in a sea of electrons (Cu, Fe). Malleable, lustrous, excellent conductors. The pattern: properties reveal the bonding The AP exam often gives you properties and asks you to identify the solid type — or vice versa. Key clues: • Conducts electricity as a SOLID → metallic • Conducts only when melted/dissolved → ionic • Very high mp + insulator + super hard → covalent network • Low mp + soft → molecular solid Stronger particle-to-particle forces always mean higher melting points and greater hardness. Four solid types — a solid's melting point and conductivity reveal what holds its particles together. The Review Hub · AP Chemistry Unit 3 TOPIC 3.3 Solids, Liquids & Gases Solid fixed shape & volume particles vibrate in place Liquid fixed volume, takes shape particles slide past each other Gas fills the container particles move fast & far apart Phase changes & energy Adding energy overcomes IMFs and moves toward gas: solid → (melting) → liquid → (vaporization) → gas. Removing energy reverses it: gas → (condensation) → liquid → (freezing) → solid. Sublimation = solid → gas directly. During a phase change, TEMPERATURE stays constant — the added energy breaks IMFs instead of speeding up particles. Stronger IMFs require more energy to melt/boil → higher melting and boiling points. State depends on the balance of particle energy vs. IMFs; phase changes happen at constant temperature. The Review Hub · AP Chemistry Unit 3 TOPIC 3.4 The Ideal Gas Law PV = nRT What each variable is P = pressure (atm) · V = volume (L) n = moles of gas · T = temperature (must be in KELVIN) R = 0.08206 L·atm/(mol·K), the gas constant Always convert °C → K (K = °C + 273) before plugging in. The relationships inside it At fixed T & n: P and V are INVERSE (Boyle's law). At fixed P & n: V and T are DIRECT (Charles's law). Combined gas law: P₁V₁/T₁ = P₂V₂/T₂ At STP (0 °C, 1 atm), 1 mol of ideal gas = 22.4 L. Worked example What volume does 2.0 mol of gas occupy at 1.0 atm and 300 K? V = nRT ÷ P = (2.0)(0.0821)(300) ÷ 1.0 = 49 L To find any variable, rearrange PV = nRT and keep units consistent (L, atm, mol, K). The law is "ideal" — it assumes no IMFs and zero molecular volume (see Topic 3.6). PV = nRT links pressure, volume, moles, and temperature — always use Kelvin. The Review Hub · AP Chemistry Unit 3 TOPIC 3.5 Kinetic Molecular Theory (KMT) MAXWELL–BOLTZMANN DISTRIBUTION lower T higher T # of molecules molecular speed (kinetic energy) → KMT's assumptions • Gas particles are in constant, random motion • Their volume is negligible vs. the container • No IMFs between them (no attraction/repulsion) • Collisions are perfectly elastic (no energy lost) Temperature = average kinetic energy Kelvin temperature is directly proportional to the AVERAGE KE of the particles. Higher T → faster average speed, and a broader, flatter distribution curve. At the same temperature, lighter gas molecules move faster than heavier ones (same average KE). The area under both curves is equal (same number of molecules). KMT explains WHY the ideal gas law works. Pressure comes from particle–wall collisions. KMT models gases as tiny, fast, non-interacting particles — Kelvin temperature ∝ average KE. The Review Hub · AP Chemistry Unit 3 TOPIC 3.6 Deviation from the Ideal Gas Law Real gases aren't quite ideal The ideal gas law assumes gas particles have NO volume and NO attractions. Real particles have both, so real gases deviate from PV = nRT most under conditions that squeeze particles close together. Deviation is largest at HIGH pressure and LOW temperature. High pressure Particles are forced close together, so their own VOLUME becomes significant compared to the container — actual volume is a bit larger than ideal. Empty space assumption breaks down. Low temperature Slow-moving particles let IMFs pull them together, reducing collisions with the walls — so actual pressure is a bit lower than ideal. "No attractions" assumption breaks down. Which gases behave most ideally? Small, nonpolar gases with weak IMFs — like helium and hydrogen — behave most ideally. Large or polar gases (with strong IMFs) deviate the most. Real gases deviate most at high pressure and low temperature — where volume & IMFs matter. The Review Hub · AP Chemistry Unit 3 TOPIC 3.7 Solutions & Mixtures The parts of a solution A solution is a HOMOGENEOUS mixture. Solute — the substance being dissolved (smaller amount). Solvent — what does the dissolving (larger amount). e.g., salt water: salt = solute, water = solvent. Molarity — the key concentration unit Molarity (M) = moles solute ÷ liters solution Units: mol/L. A "1.0 M" solution has 1 mole of solute per liter of the FINAL solution (not per liter of solvent). Worked example What is the molarity of a solution made by dissolving 58.5 g of NaCl in enough water to make 0.500 L? 1. Moles NaCl = 58.5 g ÷ 58.5 g/mol = 1.00 mol 2. Molarity = 1.00 mol ÷ 0.500 L = 2.00 M For ionic solutes, remember they dissociate: 2.00 M NaCl gives 2.00 M Na⁺ AND 2.00 M Cl⁻ in solution. CaCl₂ at 1.0 M → 1.0 M Ca²⁺ and 2.0 M Cl⁻ (one formula unit releases three ions). Ion concentrations matter for conductivity, colligative properties, and reaction stoichiometry. Molarity = moles ÷ liters; ionic solutes dissociate, so count each ion's concentration. The Review Hub · AP Chemistry Unit 3 TOPIC 3.8 Representations of Solutions PARTICULATE VIEW — NaCl DISSOLVING ++++ −−−− ions separate & are surrounded by water (not shown) Dilution M₁V₁ = M₂V₂ Adding solvent lowers concentration but keeps the MOLES of solute constant — so M₁V₁ (before) = M₂V₂ (after). Example: dilute 10 mL of 6.0 M HCl to 60 mL. M₂ = (6.0 × 10) ÷ 60 = 1.0 M HCl A 6× volume increase gives a 6× lower concentration. Reading particulate diagrams The AP exam uses particle-level pictures to test whether you understand solutions at the molecular scale: • A soluble ionic compound is shown as SEPARATED ions (dissociated), each surrounded by solvent. • A molecular (covalent) solute like glucose stays as WHOLE molecules — it dissolves but does not dissociate. More particles drawn = higher concentration; correct ion ratios must match the formula. Dilution keeps moles constant: M₁V₁ = M₂V₂; ionic solutes appear as separated ions. The Review Hub · AP Chemistry Unit 3 TOPIC 3.9 Separation of Solutions & Mixtures Filtration Separates an insoluble SOLID from a liquid, by particle size. e.g., sand out of water. Distillation Separates liquids by BOILING POINT — heat, vaporize, then condense. e.g., salt from seawater. Chromatography Separates components by their different attractions to mobile vs. stationary phase. e.g., separating ink dyes. How chromatography works A MOBILE phase (solvent) carries the mixture across a STATIONARY phase (paper or gel). Each component is attracted differently to the two phases based on its intermolecular forces and polarity. • Components that "like" the mobile phase travel FAR (move fast with the solvent). • Components that "like" the stationary phase stay CLOSE to the start (move slowly). The result is separated spots or bands — the basis for the "like dissolves like" idea in Topic 3.10. All these methods are PHYSICAL — they don't change the chemical identity of the components. Mixtures separate by physical properties — size (filtration), boiling point (distillation), or attraction (chromatography). The Review Hub · AP Chemistry Unit 3 TOPIC 3.10 Solubility "Like dissolves like" A solute dissolves in a solvent when their INTERMOLECULAR FORCES are compatible. Polar solvents dissolve polar & ionic solutes (water dissolves salt); nonpolar solvents dissolve nonpolar solutes (oil dissolves grease). Saturation levels Unsaturated — more solute can still dissolve. Saturated — the maximum has dissolved (equilibrium). Supersaturated — an unstable excess; will crash out. A saturated solution is at dynamic equilibrium. What changes solubility Temperature: most SOLIDS dissolve more in hot water; GASES dissolve LESS as temperature rises. Pressure: raises gas solubility (Henry's law — think soda). Pressure has almost no effect on solids and liquids. "Insoluble" ionic compounds still dissolve a tiny bit — that's the topic of solubility equilibria (Ksp) in Unit 7. Like dissolves like — polar dissolves polar/ionic; nonpolar dissolves nonpolar. The Review Hub · AP Chemistry Unit 3 TOPIC 3.11 Spectroscopy & the Electromagnetic Spectrum THE EM SPECTRUM — low energy (left) → high energy (right) radio microwave infrared visible UV X-ray gamma long wavelength, low frequency short λ, high frequency Different regions probe different things Microwave → molecular rotation Infrared (IR) → bond vibrations (identifies bonds/groups) UV–visible → electron transitions (color, concentration) Higher-energy light drives bigger energy changes. How spectroscopy works A molecule absorbs light whose energy exactly matches the energy gap of a transition (rotation, vibration, or electron jump). By measuring WHICH wavelengths are absorbed, we learn about the molecule's structure. Wavelength (λ) and frequency (ν) are inversely related: shorter wavelength = higher frequency = higher energy. Molecules absorb light matching an energy gap; different EM regions probe rotation, vibration, and electrons. The Review Hub · AP Chemistry Unit 3 TOPIC 3.12 Properties of Photons Light comes in packets — photons A photon's energy depends only on its frequency: E = hν = hc ÷ λ h = Planck's constant (6.626 × 10⁻³⁴ J·s), c = speed of light The relationships • Energy is DIRECTLY proportional to frequency (ν) • Energy is INVERSELY proportional to wavelength (λ) • c = λν (wavelength × frequency = speed of light) So: high frequency ↔ short wavelength ↔ high energy. Worked example & why it matters Find the energy of a photon with frequency 5.0 × 10¹⁴ Hz. E = hν = (6.626 × 10⁻³⁴ J·s)(5.0 × 10¹⁴ /s) = 3.3 × 10⁻¹⁹ J per photon A photon is absorbed only if its energy EXACTLY matches the gap between two allowed energy levels. This is why elements emit and absorb specific colors — the basis of atomic emission spectra and PES (Topic 1.6). UV and X-rays (high energy) can ionize or damage molecules; radio and microwaves (low energy) cannot. To convert per-photon energy to per-mole, multiply by Avogadro's number. E = hν = hc/λ — higher frequency (shorter wavelength) means a higher-energy photon. The Review Hub · AP Chemistry Unit 3 TOPIC 3.13 The Beer-Lambert Law ABSORBANCE vs. CONCENTRATION slope = ε·b absorbance (A) concentration (c) → The law A = ε b c A = absorbance · ε = molar absorptivity (constant) b = path length · c = concentration Absorbance has no units; it's a ratio. Absorbance ∝ concentration With ε and b fixed, A is DIRECTLY proportional to c — a straight line through the origin. Why it's useful Measure a solution's absorbance in a spectrophotometer, then use a calibration line (A vs. c of known standards) to find the UNKNOWN concentration of a colored solution — a common AP lab technique. A = εbc — absorbance is directly proportional to concentration, so it can measure an unknown. The Review Hub · AP Chemistry Unit 3
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This is great for review the night before the exam — fast, visual, and covers everything you need to remember about Unit 3's intermolecular forces content.