AP Chemistry Review: Atoms

Introduction to AP Chemistry Review

• The video provides a quick review of major AP Chemistry topics.
• It is intended a starting point for review, not a full course replacement.
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Unit 1: Atoms

• The mole is used to count large numbers of atoms and molecules.
• One mole of an element mass equals its atomic mass in grams.
• One mole of a compound mass equals the sum of its constituents' atomic masses in grams.
• About 55.85 grams makes one mole of iron.
• Approximately 18.02 grams constitutes one mole of water (H2O).
• Both contain 6.022 x 10^23 particles.
• The electron configuration of neon is 1s2, 2s2, 2p6.
• Atoms achieve maximum stability with 8 valence electrons (octet rule).
• Coulomb's Law: Greater charge = stronger attractive force.
• Closer proximity of charged particles = stronger attractive force.
• Valence electrons are held more loosely because they are farther from the nucleus.
• Photoelectron spectroscopy: Each peak signifies a sublevel.
• Higher peaks mean more electrons.
• Removing electrons from sublevels on the left requires more energy.
• Removing electrons from sublevels on the right requires less energy.
• The diagram represents the element calcium.
• Atomic radius increases down and to the left on the periodic table.
• First ionization energy is highest for atoms in the top right corner of the periodic table.
• Anions (negatively charged atoms) are larger because they gain electrons.
• Cations (positively charged atoms) are smaller because they lose electrons.

Unit 2: Chemical Compounds

• Ionic bonds form between metals and nonmetals via electrostatic forces.
• Covalent bonds form between nonmetals via electron sharing.
• Covalent bonds can be polar (unequal sharing) or nonpolar (equal sharing).
• Covalent bonds create molecules.
• Ionic compounds exist as 3D lattices of alternating cations and anions.
• Metals and metal alloys use metallic bonding characterized by mobile electrons.
• Lewis dot diagrams show molecule shapes and valence electrons.
• Atoms arrange to achieve eight valence electrons (octet).
• Double or triple bonds are formed if needed.
• Molecules have shapes like tetrahedral (109.5 degree angle).
• Linear shapes have 180 degree bond angles.
• Trigonal planar shapes have 120 degree bond angles.

Unit 3: Intermolecular Forces

• Dispersion forces are weak, increasing with molecule size/electron count.
• More electrons means a molecule is more polarizable.
• Dispersion forces dominate in nonpolar molecules.
• Dipole-dipole forces act between polar molecules.
• Dipole-dipole forces are generally stronger than dispersion forces.
• Hydrogen bonding is strong between molecules with O-H, N-H, or F-H bonds.
• Solids are crystalline, molecules tightly packed, with fixed shape/volume.
• Liquids have more space between molecules, allowing them to flow.
• Gases have independent molecules, enabling easy expansion and compression.
• PV=nRT (Ideal Gas Law): pressure, volume, moles, temperature relationships.
• Ideal Gas Law assumes small molecules, weak attractions, high temperature, or low pressure.
• Higher temperatures raise the average kinetic energy of molecules.
• Maxwell-Boltzmann distribution shows faster molecule movement at higher temperatures.
• Molarity: solution concentration in moles of solute per liter of solution.
• "Like dissolves like": polar solutes in polar solvents, nonpolar in nonpolar.
• Wavelength (lambda) x frequency (nu) equals the speed of light.
• Planck’s constant multiplied by frequency yields energy of a single photon.
• Spectrophotometry measures solution concentration; higher absorbance = higher concentration.
• Spectrophotometry constructs graphs to find the concentration of an unknown solution.

Unit 4: Chemical Reactions

• Net ionic equations exclude spectator ions (e.g., sodium, potassium cations, nitrate anions).
• Balanced equations have the same number of atoms for each element on both sides.
• Coefficients in balanced equations show mole ratios.
• Balanced equations represent a recipe for the reaction.
• Stoichiometric calculations always start by converting quantities to moles.
• Then coefficients from balanced equations are used to form a mole ratio.
• Finally, convert to the desired units.
• Precipitation reactions occur when mixing two solutions to form a solid.
• Oxidation-reduction (redox) reactions involve electron loss (oxidation) and gain (reduction).
• Acid-base reactions involve an acid reacting with a base to form a conjugate acid and base.
• Acids donate protons (H+), while bases accept them.
• An acid has one more H+ than its conjugate base.

Unit 5: Kinetics

• Balanced equations show relative rates: e.g., if NH3 coefficient is twice that of N2, ammonia appearance rate is twice nitrogen disappearance rate.
• Rate law: Rate = k [A]^m [B]^n (k is the rate constant, exponents are orders).
• If doubling concentration quadruples the rate, it’s second-order.
• If doubling concentration doubles the rate, it’s first-order.
• If doubling concentration doesn’t change the rate, it’s zero-order.
• Integrated rate law equations calculate the amount remaining after time: Rate constant k, time elapsed t, initial concentration A subzero, and elapsed concentration A sub T.
• Most reactions proceed via multiple steps, together called a reaction mechanism.
• The slowest step dictates the overall reaction rate.
• A slow step rate law becomes the rate law for the whole reaction.
• Molecules must collide with enough energy and proper orientation to react.
• This forms a high-energy transition state at the peak of the reaction energy graph.
• Activation energy is the energy required to initiate a reaction.
• Exothermic reactions release heat into the environment.
• Reactions speed up with increased temperature, decreased particle size, increased reactant concentration.
• Catalysts provide an alternate mechanism, which lowers activation energy.

Unit 6: Thermodynamics

• Endothermic reactions absorb heat, exothermic reactions release it.
• Q = M C delta T is the heat transfer equation.
• Q = heat in Joules, M = mass in grams, C = specific heat capacity, delta T = temperature change.
• Change in enthalpy (delta H) is the reaction heat change in kilojoules per mole.
• Estimate delta H using bond enthalpies (broken bonds - formed bonds).
• Or use enthalpy of formation (products - reactants).
• Hess’s Law: summing delta H values of individual reactions to get the delta H of a new reaction.

Unit 7: Equilibrium

• Equilibrium: forward and reverse reaction rates are equal, concentrations remain constant.
• Reaction quotient (Q) = [Products]/[Reactants], raised to the power of the coefficients, omit liquids and solids.
• At equilibrium, Q = K (equilibrium constant).
• If Q ≠ K, the reaction proceeds until it reaches equilibrium.
• Large K values means lots of product, little reactant.
• Small K values means lots of reactant, little product.
• An ICE (Initial, Change, Equilibrium) box finds final concentrations.
• Le Chatelier’s Principle: Adding a component shifts the reaction to the other side; removing a product shifts the reaction to replenish it at the expense of the reactants.
• ONLY temperature changes alter the equilibrium constant (K).

Unit 8: Acids and Bases

• pH = -log[H+]
• pOH = -log[OH-]
• pH + pOH = 14 (at 25°C)
• [H+][OH-] = 1 x 10^-14 = Kw (at 25°C)
• Strong acids/bases ionize fully: e.g., 0.50 M nitric acid has [H+] of 0.50 M, pH = -log(0.50).
• Weak acids/bases dissociate partially: use ICE box to solve equilibrium problems, using Ka for acids and Kb for bases.
• Acid-base titrations find the concentration of an acid or base via an indicator color change at the endpoint.
• A titration curve plots pH against base volume.
• The inflection point is the equivalence point.
• The pH halfway to the equivalence point indicates the Ka of the acid.
• Buffers resist pH changes with a mix of weak acid and its conjugate base.
• The Henderson-Hasselbalch equation calculates buffer pH.

Unit 9: Applications of Thermodynamics

• Entropy (S) is the disorder in matter: solids have least, gases have most.
• Higher temperatures increases entropy.
• Predict entropy changes by looking at a reaction
• Gibbs Free Energy (delta G) indicates thermodynamic favorability: Delta G equals delta H minus temperature in Kelvins times delta S.
• Negative delta G values are thermodynamically favored.
• Positive delta G values are not favored.
• Delta G = -RTlnK, showing the relationship between thermodynamic favorability and the equilibrium constant.
• Galvanic cells have two half-reactions: oxidation and reduction.
• Reduction occurs at the cathode; oxidation occurs at the anode.
• Electrons flow from anode to cathode.
• The salt bridge allows ion flow, maintaining charge balance.
• Anions move toward the anode, cations toward the cathode.
• Cell voltage is calculated using standard reduction potentials.
• Voltage drops as the galvanic cell operates, reaching zero at equilibrium.
• Standard conditions: 25°C, 1 M concentration.
• The Nernst Equation calculates voltage under non-standard conditions.
• Galvanic cells are thermodynamically favored.
• Delta G = -nFE (n = number of electrons transferred, E = overall cell voltage).
• Electrolysis: Uses an external electricity source to drive a reaction.
• I = Q/t (Electrical current, amp = charge (Q: coulombs) /time (t: seconds)), finds how much of an element is plated out.
• Number of coulombs is used to calculate the amount of metal plated out.
• Entropy increases when solids transition to liquids or gases.