CHEMICAL EQUILIBRIUM:
  For aA + bB ⇌ cC + dD:
    Kc = [C]^c[D]^d / ([A]^a[B]^b)
    Kp = Kc(RT)^Δn where Δn = moles of gas products - moles of gas reactants
  
  Le Chatelier's Principle:
    Add reactant → shifts right | Remove product → shifts right
    Increase pressure → shifts to side with fewer moles of gas
    Increase temperature → shifts in endothermic direction (absorbs heat)
  
  Degree of dissociation (α):
    For A ⇌ nB: if initial moles = 1, at equilibrium: (1-α) A, nα B
    Total moles at equilibrium = 1 - α + nα = 1 + (n-1)α

IONIC EQUILIBRIUM:
  pH = -log[H⁺] | pOH = -log[OH⁻] | pH + pOH = 14 at 25°C
  
  Strong acids: HCl, H₂SO₄, HNO₃, HBr, HI, HClO₄
  Strong bases: NaOH, KOH, Ba(OH)₂, Ca(OH)₂
  
  Weak acid: Ka = [H⁺][A⁻]/[HA]
    [H⁺] = √(Ka × C) for weak acid of concentration C
    pH = ½(pKa - log C) = ½pKa + ½pC
  
  Buffer solution:
    Henderson-Hasselbalch: pH = pKa + log([Salt]/[Acid])
    Best buffer at pH = pKa (equal concentrations of acid and salt)
  
  Solubility product: Ksp = [Cation]^m [Anion]^n
    Precipitation occurs when ionic product > Ksp
    Common ion effect: decrease in solubility when common ion added

GALVANIC CELL:
  Cell notation: anode | anode solution || cathode solution | cathode
    Anode: oxidation (negative electrode, loses electrons)
    Cathode: reduction (positive electrode, gains electrons)
  
  EMF of cell: E_cell = E_cathode - E_anode (reduction potentials)
  
  Nernst equation: E = E° - (0.0592/n) log Q (at 25°C)
    At equilibrium: E = 0, so E° = (0.0592/n) log K
    Relation: ΔG° = -nFE° | ΔG° = -RT ln K

ELECTROLYSIS:
  Faraday's 1st law: mass deposited ∝ charge (Q = It)
  Faraday's 2nd law: mass ∝ equivalent weight
  
  Mass deposited: m = (M × I × t) / (n × F)
    where M = molar mass, n = n-factor, F = 96500 C/mol
  
  Products of electrolysis:
    At cathode: cation with highest reduction potential reduced
    At anode: anion with lowest oxidation potential oxidized (or O₂ from H₂O)