AP Chemistry Full Study Guide Notes

Unit 1: Atomic Structure and Properties

• Subatomic Particles:

  • Protons, neutrons, electrons are the fundamental constituents of atoms.

  • Isotopes are variants of an element with different numbers of neutrons.

  • Mass number is the total number of protons and neutrons in an atom's nucleus.

• Electron Configuration:

  • Aufbau principle: Electrons fill the lowest energy levels first.

  • Hund's rule: Electrons individually occupy each orbital within a subshell before doubling up.

  • Pauli exclusion principle: No two electrons in an atom can have the same set of quantum numbers.

• Photoelectron Spectroscopy (PES):

  • A technique used to identify elements based on electron binding energies.

  • Each element exhibits a unique PES spectrum, acting as a fingerprint.

• Periodic Trends:

  • Atomic radius: Increases down a group (due to added energy levels) and decreases across a period (due to increasing nuclear charge).

  • Ionization energy: Increases across a period (due to increasing nuclear charge) and decreases down a group (due to increased shielding and distance from the nucleus).

  • Electronegativity: Increases across a period (due to increasing nuclear charge) and decreases down a group (due to increased shielding and distance from the nucleus).

Unit 2: Molecular and Ionic Compound Structure and Properties

• Types of Bonds:

  • Ionic Bond:

    • Transfer of electrons between a metal and a nonmetal.

    • Results in the formation of ions and electrostatic attraction.

  • Covalent Bond:

    • Sharing of electrons between two nonmetals.

    • Can be polar or nonpolar depending on electronegativity differences.

  • Metallic Bond:

    • A "sea of electrons" shared among metal atoms.

    • Accounts for the conductivity and malleability of metals.

• Bond Polarity & Electronegativity:

  • Electronegativity is the ability of an atom to attract electrons in a chemical bond.

  • Polar bonds occur when there is an unequal sharing of electrons due to electronegativity differences.

• Lewis Structures:

  • Representations of molecules showing valence electrons and bonds.

  • Octet rule: Atoms tend to gain, lose, or share electrons to achieve a full outer shell of eight electrons.

  • Resonance structures: Multiple valid Lewis structures can be drawn for a molecule.

  • Formal charge: Helps determine the most stable Lewis structure.

• VSEPR Theory:

  • Valence Shell Electron Pair Repulsion theory.

  • Predicts molecular geometry based on minimizing electron pair repulsion around a central atom.

  • Determines bond angles based on the arrangement of atoms.

• Intermolecular Forces (IMFs):

  • London Dispersion Forces:

    • Present in all molecules.

    • Temporary, induced dipoles due to electron movement.

  • Dipole-Dipole Forces:

    • Occur between polar molecules.

    • Positive end of one molecule attracts the negative end of another.

  • Hydrogen Bonding:

    • Strongest type of IMF.

    • Occurs when hydrogen is bonded to highly electronegative atoms (N, O, F).

• Properties:

  • Melting/boiling points: Affected by the strength of IMFs.

  • Conductivity: Ability to conduct electricity, depends on the presence of mobile charge carriers.

Unit 3: Intermolecular Forces and Properties

• States of Matter:

  • Gas: Particles widely separated, weak IMFs.

  • Liquid: Particles closer together, stronger IMFs than gases.

  • Solid: Particles tightly packed, strong IMFs.

• IMFs and Phase Changes:

  • Phase changes involve breaking or forming IMFs.

  • Energy is required to overcome IMFs during melting, boiling, and sublimation.

• Gas Laws:

  • Ideal Gas Law: $PV = nRT$

    • P = pressure, V = volume, n = number of moles, R = ideal gas constant, T = temperature.

  • Boyle's Law: P1V1 = P2V2

  • Charles's Law: $\frac{V<em>1}{T</em>1} = \frac{V<em>2}{T</em>2}$

  • Avogadro's Law: $\frac{V<em>1}{n</em>1} = \frac{V<em>2}{n</em>2}$

  • Combined Gas Law: $\frac{P<em>1V</em>1}{T<em>1} = \frac{P</em>2V<em>2}{T</em>2}$

• Deviations from Ideal Behavior:

  • Real gases deviate from ideal behavior at high pressures and low temperatures.

  • Van der Waals equation accounts for these deviations.

• Solution Properties:

  • Solubility: Ability of a substance to dissolve in a solvent.

  • Concentration (Molarity): $M = \frac{moles \space of \space solute}{Liters \space of \space solution}$

  • Electrolytes: Substances that conduct electricity when dissolved in water.

• Beer-Lambert Law:

  • $A = \epsilon b c$

    • A = absorbance, \epsilon = molar absorptivity, b = path length, c = concentration.

Unit 4: Chemical Reactions

• Types of Reactions:

  • Synthesis: Two or more reactants combine to form a single product.

  • Decomposition: A single reactant breaks down into two or more products.

  • Combustion: A substance reacts with oxygen to produce heat and light.

  • Single Replacement: One element replaces another in a compound.

  • Double Replacement: Two compounds exchange ions.

  • Acid-Base: A reaction between an acid and a base.

  • Redox: A reaction involving the transfer of electrons.

• Balancing Equations:

  • Ensuring that the number of atoms of each element is the same on both sides of the equation.

• Net Ionic Equations:

  • Equations that show only the species that participate in the reaction.

• Stoichiometry:

  • Mole conversions: Using molar mass to convert between mass and moles.

  • Limiting reagent: The reactant that is completely consumed in a reaction.

  • Percent yield: $\% \space yield = \frac{actual \space yield}{theoretical \space yield} \times 100$

• Titrations:

  • Equivalence point: The point at which the acid and base have completely reacted.

  • Indicators: Substances that change color at or near the equivalence point.

Unit 5: Kinetics

• Reaction Rates:

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

• Rate Laws:

  • $rate = k[A]^n[B]^m$

    • k = rate constant, [A] and [B] are concentrations of reactants, n and m are reaction orders.

• Integrated Rate Laws:

  • 0 order: $[A]<em>t = -kt + [A]</em>0$

  • 1st order: $ln[A]<em>t = -kt + ln[A]</em>0$

  • 2nd order: $\frac{1}{[A]<em>t} = kt + \frac{1}{[A]</em>0}$

• Mechanisms:

  • Elementary steps: Individual steps in a reaction mechanism.

  • Rate-determining step: The slowest step in the mechanism, which determines the overall rate of the reaction.

• Catalysts:

  • Lower activation energy, speeding up the reaction without being consumed.

Unit 6: Thermodynamics

• Heat and Enthalpy (H):

  • Endothermic (+H): Absorbs heat from the surroundings.

  • Exothermic (-H): Releases heat to the surroundings.

  • $q = mc\Delta T$

    • q = heat, m = mass, c = specific heat capacity, \Delta T = change in temperature.

  • Calorimetry: Measurement of heat flow.

• Hess's Law:

  • The enthalpy change for a reaction is independent of the pathway.

• Bond Enthalpies:

  • $\Delta H = \sum bonds \space broken - \sum bonds \space formed$

• Entropy (S) and Gibbs Free Energy (G):

  • $G = H - TS$

    • G = Gibbs free energy, H = enthalpy, T = temperature, S = entropy.

  • Spontaneity: \Delta G < 0 for a spontaneous reaction.

Unit 7: Equilibrium

• Equilibrium Expressions:

  • $K_{eq} = \frac{[products]}{[reactants]}$

• Reaction Quotient (Q):

  • Predicting shifts in equilibrium based on comparing Q and K.

• Le Châtelier's Principle:

  • Change in concentration: Equilibrium shifts to relieve the stress.

  • Change in pressure: Equilibrium shifts to decrease or increase the number of gas molecules.

  • Change in temperature: Equilibrium shifts to absorb or release heat.

• ICE Tables:

  • Used to calculate equilibrium concentrations.

• Kp and Kc Conversions:

  • $K<em>p = K</em>c(RT)^{\Delta n}$

    • $\Delta n$ = change in moles of gas.

Unit 8: Acids and Bases

• Strong vs Weak Acids/Bases:

  • Strong acids/bases: Completely dissociate in water.

  • Weak acids/bases: Partially dissociate in water.

• pH and pOH:

  • $pH = -log[H^+]$

  • $pOH = -log[OH^-]$

  • $pH + pOH = 14$

• Ka, Kb, and Kw:

  • $K_a$ = acid dissociation constant.

  • $K_b$ = base dissociation constant.

  • $K_w$ = ion product constant for water.

• Titration Curves:

  • Strong acid/strong base: Equivalence point at pH = 7.

  • Weak acid/strong base: Equivalence point at pH > 7.

• Buffers:

  • Resist changes in pH.

  • Henderson-Hasselbalch equation: $pH = pK_a + log(\frac{[A^-]}{[HA]})$

Unit 9: Applications of Thermodynamics

• Electrochemistry:

  • Galvanic cells: Produce electricity through spontaneous reactions.

  • Electrolytic cells: Use electricity to drive non-spontaneous reactions.

• Anode (oxidation): Where oxidation occurs.

• Cathode (reduction): Where reduction occurs.

• Cell notation: A shorthand notation for electrochemical cells.

• Standard cell potentials (E):

• Nernst Equation:
  Relates \E to concentration.

• $\Delta G$ and \E:

  • $\Delta G = -nFE$

• Thermodynamic Favorability:
  Coupled reactions, spontaneous processes

Test Strategies

• Practice with FRQs and MCQs from past exams (College Board).

• Use process of elimination on MCQs.

• Show units and work on FRQs.

Suggested Resources

• AP Chemistry Crash Course (book).

• Khan Academy, Fiveable, Bozeman Science.

• Practice Tests: Princeton Review, Barron's, College Board released exams.