• UNIT 1

  • Moles, atoms, atomic mass (EZ)

  • 6.02x10^23

  • Electron configs

  • Coulomb's law; Greater magnitude of charge and distance

  • Valence electrons are farther 

  • Each peak in a photoelectron spectrum is a sublevel

  • Peaks on the left require more energy

  • Anion –, cation +

  • UNIT 2

  • Ionic bonds, electrostatic forces, + and -, brittle, high melting point, conduct electricity when dissolved

  • Covalent bonds, two nonmetals, lower melting points, don’t conduct electricity

  • Polar; share unequally

  • nonpolar; share equally

  • Ionic compounds have a repeating lattice structure, cations and anions alternate

  • Metallic; a sea of electrons

  • Lewis diagram

  • VSEPR

  • Linear: 180 degrees

  • UNIT 3

  • Intermolecular forces

  • LDFs: usually weakest except in large molecules, found in ALL molecules, more electrons = more polarizable=stronger LDFs

  • Dipole-dipole forces (POLAR molecules), moderate strength

  • Hydrogen bonding, usually strongest, FON 

  • Solids, crystalline

  • Liquids have more freedom of motion

  • Gases have independent motion

  • PV=nRT 

  • Gas acts most ideal at high temp and low pressure, also very small molecules

  • High temp, higher average kinetic energy, greater velocity

  • Molarity mol/L

  • “Like dissolved like”

  • Polar molecules dissolve in polar solvents(like water) because of dipole-dipole or h-bond forces between the solute and solvent

  • Nonpolar molecules dissolve in nonpolar solvents because of LDFs

  • Planck’s constant

  • The higher the absorbance, the higher the concentration

  • UNIT 4

  • Omit spectator ions

  • Net ionic equation

  • Coefficients form mole ratio

  • In precipitation reactions, two solutions are mixed, and a solid is formed

  • Redox reactions, OIL RIG

  • Acid-base reactions, acid and bade form conjugate acid/base. Acid proton donor, base proton acceptor

  • UNIT 5

  • Coefficients determine rate

  • Double concentration and rate quadruples: 2nd Order

  • Double concentration and rate doubles: 1st Order

  • Double concentration and rate don't change: 0th Order

  • 0 order = [A]sub t - [A]sub 0 = -kt

  • 1st order = ln[A] sub t - ln [A] sub 0 = -kt

  • 2nd order = 1/[A]sub t - 1/[A]sub 0 = kt

  • Reaction mechanism

  • Slow step determines the rate of reaction

  • Molecules have to collide with enough energy and the correct orientation to react

  • Activation energy, is the energy required to start

  • Net loss of heat to the surroundings; exothermic

  • Speed up reaction: higher temp, higher concentration, larger surface area, add catalyst

  • Catalyst lowers activation energy

  • UNIT 6

  • Endothermic absorbs heat, exothermic releases heat

  • q=mCdeltaT

  • Heat change= change in enthalpy

  • Estimate using bond enthalpies

  • Enthalpy of formation products-reactants

  • Hess’s law adds delta H together

  • UNIT 7

  • Equilibrium DOES NOT STOP reaction

  • Forward and reverse reaction rates are equal

  • Ration quotient = Q

  • No liquids or solids

  • Q = k at equilibrium

  • k>1 lots of product, equilibrium is at the right

  • K<1 lot of reactant, the equilibrium lies to the left

  • Given initial concentration/pressure, use an ICE chart

  • Le Chatliers, the reaction will try to compensate 

  • UNIT8

  • pH=-log[H+]

  • Strong acids and strong bases ionize completely

  • Weak acids and bases are equilibrium problems (ice chart)

  • ACid-base situations, try to find concentration

  • The endpoint is a complete reaction

  • Halfway point pH=pKa

  • Buffers:weak acid and its conjugate base

  • Hendresson-hasselbpach equation 

  • pH= pKa + log([A-]/[HA])

  • UNIT 9

  • Entropy (S) disorder

  • Solids -> gases increasing entropy

  • Higher temp = more entropy

  • Gibbs free energy, thermodynamic favorability

  • ΔG=ΔH-T(ΔS)

  • Negative G is favored

  • ΔG = -RTlnK

  • Every galvanic cell has two half-reactions, redox

  • RED CAT and AN OX (reduction at the cathode, oxidation at the anode)

  • Electrons move through the wire FAT CAT (from the anode to the cathode)

  • The salt bridge allows ions to flow freely through the cell

  • Anions to the anode, cations to the cathode

  • Voltage  calculated using standard reduction potentials

  • As the call runs, the voltage slowly drops until it reaches 0 (dead battery, equilibrium)

  • Standard conditions, 25 degrees C, 1 atm, 1 M

  • Nernst equation for nonstandard conditions (E = E° - (RT/nF)lnQ

  • Galvanic cells are thermodynamically favored

  • ΔG = -nFE

  • External energy powers the electrolysis process

  • I = q/t (Current in amps = coulombs/time in seconds)