AP Chemistry Exam Preparation Notes

Unit 0 - Foundations

• Separation of Matter

  • Compounds: separate into elements via chemical changes.

  • Mixtures: separate by physical changes.

• Conservation

  • Mass and energy conserved in both chemical and physical changes.

• Volume Measurement

  • Read liquid volume between graduated markings.

  • Always read one decimal place beyond calibration.

• Measurement Devices Precision

  • Ranked from least to most precise:

  • Beaker < Graduated Cylinder < Volumetric Flasks < Burette

  • Volumetric flasks measure one specific volume.

  • Burette reads top to bottom.

Unit 1 - Moles, Atomic Structure, Periodicity

• Percent Error Formula

  • ext{Percent Error} = rac{| ext{experimental} - ext{theoretical}|}{ ext{theoretical}}

• Empirical Formula Procedure

  • Rhyme: % to mass, mass to mole, divide by small, times until whole (simplest ratio).

• Molecular Formula

  • A whole number multiple of the empirical formula.

• Bond Types

  • Ionic Bonds: Formed when a metal transfers electrons to a nonmetal; opposite charges attract.

  • Covalent Bonds: Formed by sharing electrons between nonmetals.

  • Polarity: Increased electronegativity difference increases bond polarity.

• Key Facts

  • Density: ext{Density} = rac{ ext{mass}}{ ext{volume}}

  • Elements in the same group share similar properties (Periodic Law).

  • Metals left of zig-zag line, nonmetals right.

• Ionization Energy

  • Higher energy level = lower ionization energy; easier to remove.

  • Across a period: increased Zeff = smaller radius = higher ionization energy.

• PES Graphs

  • Height of peaks indicates number of electrons; larger binding energy means electrons are closer to nucleus.

• Electron Configuration for Cations

  • Remove valence electrons first; cations are smaller than their atoms, anions are larger due to increased electron-electron repulsions.

• Isotopes

  • Same number of protons, different number of neutrons.

• Mass Spectroscopy

  • Measures atomic masses of isotopes.

Unit 2 - Bonding, Molecular Geometry, and Polarity

• Carbon Bonds

  • Carbon forms a total of 4 bonds.

• Bond Angles

  • 4 domains: 109.5°; 3 domains: 120°; 2 domains: 180°.

• Hybrid Orbitals

  • 4 domains: sp³; 3 domains: sp²; 2 domains: sp.

• Polarity of Molecules

  • Asymmetrical = polar; symmetrical = nonpolar.

• Types of Bonds

  • Single bond: sigma; double bond: sigma + pi; triple bond: sigma + 2 pi.

• Lattice Energy

  • Energy needed to break ionic bonds; increases with charge and decreases with larger ionic radii.

Unit 3A - Properties of Liquids and Solids - IMFs

• Intermolecular Forces (IMFs)

  • Weakest to strongest: London Dispersion < Dipole-Dipole < Hydrogen Bonding < Ion-Dipole.

• Polarizability

  • Larger molecules = stronger London Dispersion forces.

• Vapor Pressure and Volatility

  • Increase in IMFs decreases vapor pressure and volatility.

• Conductivity

  • Molecular solids: low melting/boiling points, non-conductive.

  • Ionic solids: high melting/boiling points; non-conductive as solids, conductive in liquid/aqueous states.

  • Covalent network solids (SiO₂, diamonds) have very high melting/boiling points.

• Metallic Bonds

  • Always conductive; hardness varies.

Unit 3B - Gases, Solutions, and Chromatography

• Gas Properties

  • Homogeneous mixtures due to constant random motion.

  • Compressible due to large particle spacing.

• Gas Pressure

  • Caused by collisions with container walls; more collisions = higher pressure.

• Gas Laws

  • PV = nRT (P = pressure, V = volume, n = moles, R = gas constant; T = temperature in Kelvin).

  • Doubling volume halves pressure.

  • Temperature and volume are directly related.

• Molar Mass

  • ext{Molar Mass} = rac{dRT}{P} (d = density in g/L).

• Kinetic Energy

  • ext{Temperature} = ext{Average Kinetic Energy}.

• Collecting Gas by Water Displacement

  • P{total} = P{dry ext{ }gas} + P_{water ext{ }vapor}.

• Ideal vs. Real Gases

  • Real gases deviate from ideal behavior under high pressure or low temperature.

Gas Properties

• Homogeneous mixtures due to constant random motion.

• Compressible due to large particle spacing.

Gas Pressure

• Caused by collisions with container walls; more collisions = higher pressure.

Gas Laws

• PV = nRT (P = pressure, V = volume, n = moles, R = gas constant; T = temperature in Kelvin).

• Doubling volume halves pressure.

• Temperature and volume are directly related.

Molar Mass

• ext{Molar Mass} = rac{dRT}{P} (d = density in g/L).

Kinetic Energy

• ext{Temperature} = ext{Average Kinetic Energy}.

Collecting Gas by Water Displacement

• P{total} = P{dry ext{ }gas} + P_{water ext{ }vapor}.

Ideal vs. Real Gases

• Real gases deviate from ideal behavior under high pressure or low temperature.

Unit 4 - Thermodynamics and Thermochemistry

Laws of Thermodynamics

1. Energy cannot be created or destroyed, only transformed.

2. In a closed system, total entropy tends to increase.

Enthalpy (H)

• Change in enthalpy: \Delta H = H{products} - H{reactants}.

• Exothermic reactions: \Delta H < 0 (releases heat).

• Endothermic reactions: \Delta H > 0 (absorbs heat).

Calorimetry

• Measures heat transfer in chemical reactions.

• q = mC\Delta T (where m = mass, C = specific heat, \Delta T = change in temperature).

Unit 5 - Kinetics and Chemical Equilibrium

Reaction Rates

• Factors affecting reaction rates: concentration, temperature, surface area, catalyst presence.

Rate Laws

• Generally in the form of: rate = k[A]^m[B]^n (where k = rate constant, A & B are reactants, m & n are reaction orders).

Equilibrium

• At equilibrium, rates of forward and reverse reactions are equal; concentrations remain constant.

• Equilibrium constant expression: K = \frac{[products]}{[reactants]} (for a-b reactions, aA + bB \rightleftharpoons cC + dD).

• Le Chatelier's Principle: if a system at equilibrium experiences a change in concentration, temperature, or pressure, the equilibrium shifts to counteract that change.

Unit 6 - Acids, Bases, and pH

Acid-Base Definitions

• Arrhenius: Acids produce H+; bases produce OH-.

• Bronsted-Lowry: Acids donate protons; bases accept protons.

pH Scale

• pH = -log[H+]; scale ranges from 0 (acidic) to 14 (basic).

Neutralization Reactions

• Acid + base = salt + water; typically exothermic.

Indicators

• Substances that change color based on pH, such as litmus or phenolphthalein.

Buffer Solutions

• Solutions that resist changes in pH upon the addition of acids or bases; often contain a weak acid and its conjugate base.