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AP Chemistry Speed Review Notes

Introduction

  • This is a speed review of the major topics in AP Chemistry.
  • It's not a replacement for a full course but a good review.
  • Ultimate Review Packet at ultimatereviewpacket.com offers study guides, answers, review videos, tips, and a full-length exam for $24.99 (with a 40% discount for teachers paying for the whole class).

Unit 1: Atoms

  • The mole is used to count large numbers of atoms and molecules.
  • One mole of an element = its atomic mass in grams.
  • One mole of a compound = the sum of atomic masses in grams.
    • 1 mole of iron (Fe) ≈ 55.85 grams.
    • 1 mole of water (H₂O) ≈ 18.02 grams.
  • Both contain 6.022 \times 10^{23} particles.
  • Electron configurations:
    • Neon (Ne): 1s^2 2s^2 2p^6
  • Atoms are most stable with eight electrons (octet) in their valence shell.
  • Coulomb's Law: describes the electrostatic force between charged particles.
    • Greater charge magnitude = stronger attraction.
    • Shorter distance = stronger attraction.
    • Valence electrons are held less tightly because they are farther from the nucleus.
  • Photoelectron Spectroscopy (PES):
    • Each peak represents a sublevel.
    • Peak height indicates the number of electrons.
    • Sublevels on the left require more energy to remove electrons.
    • Sublevels on the right require less energy.
    • Example: Calcium (Ca) PES diagram.
  • Periodic Table Trends:
    • Atomic radius increases down and to the left.
    • First ionization energy increases up and to the right.
    • Anions (negative ions) are larger than their neutral atoms.
    • Cations (positive ions) are smaller than their neutral atoms.

Unit 2: Chemical Compounds

  • Ionic bonds: electrostatic forces between metals and nonmetals.
  • Covalent bonds: sharing of electrons between nonmetals.
    • Polar: unequal sharing.
    • Nonpolar: equal sharing.
  • Covalent bonds form molecules.
  • Ionic compounds form 3D lattices.
  • Metallic bonding: electrons move freely among positive metal ions (sea of electrons).
  • Lewis Dot Diagrams: visualize molecule shapes.
    • Draw all valence electrons.
    • Arrange to give each atom eight valence electrons (octet rule), forming double or triple bonds if needed.
  • Molecular Geometry:
    • Tetrahedral: 109.5° bond angle.
    • Linear: 180° bond angle.
    • Trigonal planar: 120° bond angle.

Unit 3: Intermolecular Forces

  • Dispersion Forces: weak interactions, stronger in larger molecules with more electrons (more polarizable).
  • Dipole-Dipole Forces: between polar molecules, stronger than dispersion forces.
  • Hydrogen Bonding: strong force between molecules with O-H, N-H, or F-H bonds.
  • States of Matter:
    • Solids: crystalline, tightly packed molecules, fixed shape and volume.
    • Liquids: more space, molecules can slide, flow.
    • Gases: independent molecules, expand/compress easily.
  • Ideal Gas Law: PV = nRT (Pressure, Volume, moles, Ideal gas constant, Temperature).
  • Real gases approximate ideal conditions at small molecule size, weak attractions, high temperature, or low pressure.
  • Maxwell-Boltzmann Distribution: shows molecular speeds at different temperatures.
    • Higher temperature: more molecules move faster.
    • Lower temperature: most molecules move slower.
  • Molarity: concentration = moles of solute / liters of solution.
  • "Like dissolves like": polar solutes dissolve in polar solvents, nonpolar in nonpolar.
  • Light Interaction: \lambda \nu = c (wavelength x frequency = speed of light).
    • E = h \nu (energy of a photon = Planck's constant x frequency).
  • Spectrophotometry: measures solution concentration based on absorbance.
    • Higher absorbance = higher concentration.

Unit 4: Chemical Reactions

  • Net Ionic Equations: omit spectator ions (e.g., Na^+, K^+, NO_3^-.
  • Balancing Equations: same number of atoms of each element on both sides using coefficients.
  • Coefficients form a mole ratio.
  • Stoichiometry: convert to moles, use mole ratio, convert to final unit.
  • Types of Reactions:
    • Precipitation: two solutions mix, forming a solid.
    • Oxidation-Reduction (Redox): electron loss (oxidation) and gain (reduction).
    • Acid-Base: acid reacts with base to form conjugate acid and base.
      • Acid = proton (H+) donor.
      • Base = proton acceptor.
      • Acid has one more H+ than its conjugate base.

Unit 5: Kinetics

  • Relative rates: coefficients describe relative rates (e.g., N2 + 3H2 \rightarrow 2NH_3, ammonia appears twice as fast as nitrogen disappears).
  • Rate Law: Rate = k[A]^m[B]^n (k = rate constant, m and n are orders).
    • Double concentration, rate quadruples: second order.
    • Double concentration, rate doubles: first order.
    • Double concentration, rate unchanged: zero order.
  • Integrated Rate Laws: calculate amount left after time (relate concentration to time).
    • Includes rate constant k, time t, initial concentration [A]0, and concentration at time t, [A]t
  • Reaction Mechanisms: multiple steps, slow step determines overall rate.
  • Molecules must collide with enough energy and correct orientation to react.
    • Transition state: high-energy peak.
    • Activation energy: energy to start reaction.
  • Exothermic reaction: net heat loss to surroundings.
  • Speeding up Reactions: increase temperature, decrease particle size, increase reactant concentration, add a catalyst (lowers activation energy).

Unit 6: Thermodynamics

  • Endothermic: absorbs heat.
  • Exothermic: releases heat.
  • Heat Transfer: q = mc \Delta T (heat, mass, specific heat capacity, change in temperature).
  • Enthalpy Change (\Delta H): heat change for a reaction (kJ/mol).
    • Estimated by: Bonds broken - bonds formed.
    • Or: Enthalpies of formation of products - reactants.
  • Hess's Law: if reactions add up to a new reaction, their \Delta H values add up to the new \Delta H.

Unit 7: Equilibrium

  • At equilibrium, forward and reverse rates are equal.
  • Reaction Quotient (Q): products / reactants (raised to coefficient powers), omitting liquids and solids.
  • At equilibrium, Q = K (equilibrium constant).
  • If Q \neq K, the reaction shifts to reach equilibrium.
  • Large K: lots of product.
  • Small K: lots of reactant.
  • ICE Box: Initial, Change, Equilibrium concentrations.
  • Le Chatelier's Principle: adding a component shifts to the opposite side; removing shifts to replenish.
  • Changing temperature changes K value.

Unit 8: Acids and Bases

  • Key Equations:
    • pH = -log[H^+].
    • pOH = -log[OH^-].
    • At 25°C: pH + pOH = 14.
    • At 25°C: [H^+][OH^-] = 1 \times 10^{-14} = K_w
  • Strong acids/bases ionize completely.
    • E.g., 0.50 M nitric acid: [H^+] = 0.50 M, so pH = -log(0.50).
  • Weak acids/bases dissociate less than 1% (equilibrium problem).
    • Use ICE box and equilibrium constant expression (Ka for weak acid, Kb for weak base).
  • Acid-Base Titrations: find concentration of acid or base.
    • Endpoint: indicator changes color.
    • Titration Curve: pH vs. volume of base added.
      • Inflection point: equivalence point.
      • pH at equivalence point indicates the strength of acid/base.
      • Halfway to equivalence point, pH = pKa, so Ka = 10^{-pH}.
  • Buffers: weak acid and conjugate base mixtures that resist pH changes.
    • Henderson-Hasselbalch Equation: estimate buffer pH.

Unit 9: Applications of Thermodynamics

  • Entropy (S): disorder.
    • Solids < Liquids < Solutions < Gases.
    • Higher temperature = more entropy.
    • \Delta S: positive if disorder increases.
  • Gibbs Free Energy ($\Delta G$): thermodynamic favorability.
    • \Delta G = \Delta H - T\Delta S
    • Negative \Delta G: thermodynamically favored.
    • Positive \Delta G: not favored.
    • \Delta G = -RTlnK
  • Electrochemistry: galvanic cells have oxidation and reduction half-reactions.
    • Cathode: reduction.
    • Anode: oxidation.
    • Electrons flow from anode to cathode.
    • Salt bridge: ions flow (anions to anode, cations to cathode).
  • Galvanic cell voltage can be calculated using standard reduction potentials.
  • Voltage drops to zero at equilibrium.
  • Standard conditions: 25°C, 1 M solutions.
  • Nernst Equation: calculates voltage under non-standard conditions.
  • Galvanic cells are thermodynamically favored: \Delta G = -nFE (n = electrons transferred, F = Faraday's constant, E = voltage).
  • Electrolysis: external electricity powers a reaction.
    • I = q/t (current = charge / time).
    • Use coulombs to calculate the amount of metal plated out.