AP CHEM MIDTERM

Unit 1 — Atomic Structure & Properties

1.1 — The Mole & Molar Relationships

• Practice: 5F

• Key Ideas:

  • Use dimensional analysis to convert between mass ↔ moles ↔ particles.

  • Avogadro’s number links moles to particles.

  • The relationship for moles: n = \frac{m}{M} is essential.

  • Molar mass (g/mol) = average mass in amu of one particle.

• Be ready to:

  • Convert grams → moles → atoms/molecules.

  • Identify limiting reagents using mole ratios.

  • Interpret particle diagrams in terms of moles.

1.3 — Empirical Formulas

• Practice: 5F, 6G

• Key Ideas:

  • Empirical formula = simplest whole-number ratio.

  • Law of definite proportions: mass ratios are constant in pure compounds.

  • Convert % composition → grams → moles → ratio.

• Be ready to:

  • Determine empirical formula from mass %.

  • Compare molecular vs empirical formulas.

  • Use combustion analysis data.

1.7 — Periodic Trends (Focus: Atomic radius)

• Key Ideas:

  • Atomic radius increases down groups, decreases across periods.

  • Trends driven by:

  • Effective nuclear charge ($Z_{eff}$)

  • Shielding effects

  • Distance from the nucleus

• Be ready to:

  • Rank elements by atomic radius.

  • Explain trends using electron structure.

1.8 — Valence Electrons & Reactivity

• Practice: none listed, but conceptual

• Key Ideas:

  • Reactivity is influenced by valence electrons and nuclear attraction.

  • Elements in the same group tend to form similar compounds.

  • Typical ionic charges follow periodic table positions.

• Be ready to:

  • Predict ion charges.

  • Identify whether elements form ionic vs covalent bonds.

  • Explain why alkali metals are highly reactive.

Unit 3 — Intermolecular Forces & Gases

3.4 — Ideal Gas Law

• Practice: 5C

• Key Ideas:

  • Ideal gas law equation: PV = nRT

  • Relates macroscopic properties of gases.

  • Assumes:

  • No intermolecular forces (IMFs)

  • Negligible particle volume

• Be ready to:

  • Solve for pressure (P), volume (V), number of moles (n), or temperature (T).

  • Compare gases under identical conditions.

  • Interpret gas diagrams.

3.5 — Kinetic Molecular Theory

• Practice: 4A

• Key Ideas:

  • Kinetic energy is proportional to temperature in Kelvin.

  • All gases at the same temperature have the same average kinetic energy.

  • Lighter gases have higher speeds.

• Be ready to:

  • Compare speeds of different gases.

  • Explain Maxwell–Boltzmann distributions.

3.6 — Real Gases & Deviations

• Practice: 4B, 4D, 6E

• Key Ideas:

  • Real gases deviate from ideal behavior due to:

  • Intermolecular forces (attractive forces lower pressure)

  • Particle volume (high pressure leads to higher measured pressure)

  • Deviations are strongest under:

  • Low temperatures (near condensation)

  • High pressures

• Be ready to:

  • Argue whether a gas behaves ideally.

  • Compare gases with different net strengths of intermolecular forces.

  • Explain model diagrams showing deviations from ideal behavior.

3.7 — Solutions & Molarity

• Practice: 2C

• Key Ideas:

  • Solutions are defined as homogeneous mixtures.

  • Molarity is calculated as: M = \frac{\text{mol solute}}{\text{L solution}}

• Be ready to:

  • Calculate molarity, volume, or number of moles.

  • Identify homogeneous versus heterogeneous mixtures.

3.9 — Separation Techniques

• Practice: 5D

• Key Ideas:

  • Chromatography separates components by their polarity and strength of intermolecular interactions.

  • Distillation takes advantage of boiling point differences (stronger intermolecular forces result in lower vapor pressures).

• Be ready to:

  • Predict which component will travel farther in chromatography.

  • Determine which component distills first based on boiling points.