TransferCredit.org Guide to AP® Chemistry

This is the heaviest unit, about a fifth of the multiple-choice section. It links particles to what you can measure: intermolecular forces explain boiling points, solubility and the kind of solid; gas laws and kinetic molecular theory explain pressure, volume and temperature; and absorbed light measures concentration. Expect calculations with PV = nRT, partial pressures, molarity, c = λν, E = hν and A = εbc.

Key terms, with examples

intermolecular forces

Coulombic attractions between separate particles (molecules, atoms or ions). Much weaker than covalent bonds, they set boiling points, vapor pressures and solubility.

Example: Boiling water pulls H₂O molecules away from each other; the O–H bonds inside every molecule stay intact.

vapor pressure and boiling point

Vapor pressure is the pressure of a liquid's vapor at equilibrium. Stronger intermolecular forces mean lower vapor pressure and a higher boiling point, reached when vapor pressure equals outside pressure.

Example: At 25 °C, diethyl ether's vapor pressure is far higher than water's, so ether evaporates faster and boils at only 35 °C.

crystalline vs amorphous solid

A crystalline solid has particles in a regular, repeating 3-D arrangement; an amorphous solid has no long-range order. In both, particles vibrate but don't move past each other.

Example: Quartz, crystalline SiO₂, melts at a sharp temperature; window glass, which is amorphous, softens gradually over a range.

ideal gas law

PV = nRT links the pressure, volume, moles and kelvin temperature of an ideal gas. With P in atm and V in liters, R = 0.08206 L·atm/(mol·K).

Example: 2.00 mol at 273.15 K and 1.00 atm: V = nRT/P = 44.8 L, or 22.4 L per mole at STP.

kinetic molecular theory (KMT)

A model of gas behavior: particles move constantly and randomly and collide elastically; in an ideal gas their own volume is negligible and they don't attract each other.

Example: Pumping more air into a bike tire raises its pressure because more particles strike the inner walls each second.

solution (homogeneous mixture)

A mixture whose properties are the same throughout; it can be a solid, liquid or gas. In a heterogeneous mixture, properties depend on where you sample.

Example: Air, vinegar and brass are solutions; sand stirred into water is a heterogeneous mixture.

chromatography

Separates components by their different attractions to a moving mobile phase and a fixed stationary phase. A component attracted more to the stationary phase travels less far.

Example: On a polar silica TLC plate, a nonpolar compound travels farther than a polar one, revealing their relative polarities.

regions of the electromagnetic spectrum

Absorbed radiation changes a molecule according to its energy: microwave radiation changes rotation, infrared changes vibration, and ultraviolet/visible light causes electronic transitions.

Example: A microwave oven speeds up water's rotation; an IR spectrum shows C=O bonds vibrating; UV light excites electrons in sunscreen molecules.

Beer-Lambert law

Absorbance equals molar absorptivity × path length × concentration: A = εbc. With ε and b fixed, absorbance is proportional to concentration.

Example: A = 0.360, ε = 1.20 × 10⁴ M⁻¹cm⁻¹, b = 1.00 cm: c = 0.360/(1.20 × 10⁴) = 3.00 × 10⁻⁵ M.

London dispersion forces

Attractions between temporary, fluctuating dipoles. They exist between all molecules and are often the strongest net force between large ones. Not a synonym for "van der Waals forces".

Example: Iodine, I₂, is a solid at room temperature while fluorine, F₂, is a gas: I₂'s big electron cloud gives stronger dispersion forces.

ionic solid properties

Strong attractions between ions give low vapor pressures and high melting and boiling points. Ionic solids are brittle and conduct electricity only when the ions can move: melted or dissolved.

Example: A hammer blow shatters a salt crystal: sliding one layer lines up like charges, which repel.

particle model of solids, liquids and gases

Solid and liquid particles touch, giving similar molar volumes; liquid particles also slide past each other. Gas particles are far apart, with no definite shape or volume.

Example: One mole of liquid water fills about 18 mL; as steam at 100 °C and 1 atm it fills about 30 L.

What this area covers

How the exam asks it

Many questions name two or more substances and ask which has the higher boiling point, vapor pressure or solubility, and why: name the strongest intermolecular force in each, then compare. Gas questions mix calculations (PV = nRT, partial pressures, mole fractions) with particle reasoning from kinetic molecular theory and Maxwell-Boltzmann curves described in words. Lab questions ask how a separation or a spectrophotometer reading works and what a mistake does to the result. The exam doesn't test phase diagrams, colligative properties, or molality and percent-by-mass calculations.

Common traps

Trap 1. Boiling and melting separate whole molecules; they don't break the covalent bonds inside them. Compare intermolecular forces, not bond strengths.

Trap 2. Gas laws and KMT need kelvins. Warming a gas from 20 °C to 40 °C raises its kelvin temperature by under 7%, not 100%.

Trap 3. Hydrogen bonding needs H bonded directly to N, O or F. In CH₃OCH₃ and CH₂O every H is on carbon, so these molecules can't hydrogen-bond with each other.

Memory trick

"Rotate, vibrate, elevate": microwave, infrared, ultraviolet/visible, in order of rising photon energy.

Real gases misbehave when cold and crowded: low temperature and high pressure.

Acids and bases is the second-biggest unit (11–15%). Expect pH calculations for strong and weak acids and bases, mixtures of acids and bases, titration curves and buffers, plus questions on how molecular structure and pH decide which form of an acid is present.

Key terms, with examples

pH and pOH

Log scales of acidity: pH = −log[H₃O⁺] and pOH = −log[OH⁻]. Each unit is a tenfold change in concentration; a lower pH means more H₃O⁺.

Example: [H₃O⁺] = 1.0 × 10⁻³ M gives pH 3.00; a pH of 4.70 means [H₃O⁺] = 10^(−4.70) = 2.0 × 10⁻⁵ M.

strong acid

An acid that ionizes completely in water, so [H₃O⁺] equals the acid's starting concentration. The six to know: HCl, HBr, HI, HClO₄, H₂SO₄ and HNO₃.

Example: 0.020 M HNO₃ gives [H₃O⁺] = 0.020 M, so pH = −log(0.020) = 1.70.

weak acid and Ka

An acid that ionizes only slightly, so most molecules stay as HA. Ka = [H₃O⁺][A⁻]/[HA]; when ionization is small, [H₃O⁺] ≈ √(Ka × c).

Example: 0.10 M acetic acid, Ka = 1.8 × 10⁻⁵: [H₃O⁺] ≈ √(1.8 × 10⁻⁶) = 1.3 × 10⁻³ M, so pH = 2.87.

strong acid–strong base mixture

H₃O⁺ + OH⁻ → 2 H₂O goes to completion. Subtract moles, divide the excess reagent's moles by the total volume, then find the pH.

Example: Mixing 5.0 mmol HCl with 4.0 mmol NaOH leaves 1.0 mmol H₃O⁺ in 90.0 mL total, so pH = 1.95.

titration curve

A graph of pH versus volume of titrant added. The steep jump marks the equivalence point, where the moles of titrant equal the moles of analyte at the start.

Example: 25.00 mL of HCl needs 20.00 mL of 0.1000 M NaOH: 0.002000 mol ÷ 0.02500 L = 0.08000 M HCl.

acid strength and molecular structure

An acid is stronger when its conjugate base is more stable: electronegative atoms, inductive effects and resonance spread the negative charge. Carboxylic acids are weak; ammonia is a weak base.

Example: CCl₃COOH is far stronger than CH₃COOH: three electronegative chlorine atoms pull electron density away and stabilize the anion.

protonation state and pH

Compare pH with pKa: if pH < pKa, the acid form HA is the major species; if pH > pKa, the base form A⁻ is; at pH = pKa they're equal.

Example: Formic acid (pKa 3.75) at pH 7.4: the pH is far above the pKa, so it is almost all formate ion, HCOO⁻.

buffer

A solution with large amounts of both members of a conjugate pair. Its weak acid neutralizes added OH⁻ and its conjugate base neutralizes added H₃O⁺, so the pH barely changes.

Example: In an acetic acid/acetate buffer, added acid is used up: CH₃COO⁻ + H₃O⁺ → CH₃COOH + H₂O.

pH and solubility

A salt dissolves more at low pH when its anion is a weak base (such as F⁻ or CO₃²⁻) or OH⁻: H₃O⁺ removes the anion, shifting the dissolving equilibrium right.

Example: CaF₂ dissolves more in acid: H₃O⁺ + F⁻ → HF + H₂O lowers [F⁻], so more calcium fluoride dissolves.

autoionization of water (Kw)

Water transfers protons to itself: 2 H₂O(l) ⇄ H₃O⁺(aq) + OH⁻(aq). Kw = [H₃O⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25°C, so pH + pOH = 14.00.

Example: At 25°C, [OH⁻] = 1.0 × 10⁻⁴ M means [H₃O⁺] = 1.0 × 10⁻¹⁰ M and pH = 10.00.

strong base

A base that dissociates completely in water: the group 1 and group 2 hydroxides. [OH⁻] equals the concentration for NaOH but twice it for Ba(OH)₂.

Example: 0.0050 M Ba(OH)₂ gives [OH⁻] = 0.010 M: pOH = 2.00, so pH = 12.00.

pKa and pKb

pKa = −log Ka and pKb = −log Kb. A smaller pKa means a larger Ka and a stronger acid; a smaller pKb means a stronger base.

Example: An acid with Ka = 1.0 × 10⁻³ has pKa 3.00, so it is stronger than an acid with pKa 5.00.

What this area covers

How the exam asks it

Start every pH question by naming the major species. Strong acids and bases ionize completely, so [H₃O⁺] or [OH⁻] comes straight from the concentration; for a weak acid or base, set up Ka or Kb. For a mixture, react the strong acid or base completely first, then see what is left: excess strong reagent, a buffer, or a lone conjugate base or acid (a weak acid mixed with a weak base only reaches an equilibrium). Titration questions describe a curve in words; read pKa at the half-equivalence point and pick indicators by pKa. Ka and Kb values are given. The exam doesn't ask you to calculate a buffer's pH change after acid or base is added, or a salt's solubility at a given pH.

Common traps

Trap 1. pH 7 is neutral only at 25°C, and an equivalence point is at pH 7 only for a strong acid with a strong base. A weak acid titrated with a strong base ends above 7.

Trap 2. In the Henderson-Hasselbalch equation the base goes on top: log([A⁻]/[HA]). More base than acid puts the pH above the pKa, not below it.

Trap 3. React first, then calculate: after a strong base is added to a weak acid, use the moles left after the complete reaction, not the starting amounts.

Memory trick

Half-equivalence, half-way: the pH equals the pKa.

A buffer's two members each have a job: the acid eats added OH⁻, the base eats added H₃O⁺.