Chemistry Formula Sheet
42 formulas across 5 chapters — with variables explained and exam tips where needed.
Ch 1Basic Concepts of Chemistry(8 formulas)
Number of moles
moln = m/M
m = mass (g), M = molar mass (g/mol)
Number of particles
N = n × Nₐ
Nₐ = 6.022×10²³ mol⁻¹ (Avogadro's number)
Molarity
mol/LM = n_solute / V_solution(L)
Mole fraction
χ_A = n_A / (n_A + n_B + ...)
Parts per million
ppm = (mass of solute / mass of solution) × 10⁶
Percentage yield
% yield = (actual yield / theoretical yield) × 100
Theoretical yield is based on limiting reagent
Limiting reagent identification
Divide moles of each reactant by its stoichiometric coefficient — smallest ratio → limiting reagent
Empirical formula steps
% → g (assume 100g) → mol (÷M) → simple ratio → empirical formula
Multiply ratios to get whole numbers if needed
Ch 5States of Matter(8 formulas)
Ideal gas law
PV = nRT
R = 0.0821 L·atm/mol·K = 8.314 J/mol·K
Combined gas law
P₁V₁/T₁ = P₂V₂/T₂
If one variable is constant, simplify accordingly
Boyle's law
PV = constant (at constant T, n)
P₁V₁ = P₂V₂
Charles's law
V/T = constant (at constant P, n)
V₁/T₁ = V₂/T₂ — use Kelvin
van der Waals equation
(P + an²/V²)(V − nb) = nRT
a = intermolecular attraction, b = volume excluded per mole
Corrects for real gas behaviour
Graham's law of effusion
r₁/r₂ = √(M₂/M₁)
r = rate of effusion, M = molar mass
Lighter gas effuses faster
Average kinetic energy
KE_avg = 3/2 RT (per mole) or 3/2 kT (per molecule)
RMS speed
m/su_rms = √(3RT/M)
Ch 6Thermodynamics(8 formulas)
Relation between ΔH and ΔU
ΔH = ΔU + Δn_g RT
Δn_g = moles of gaseous products − moles of gaseous reactants
For reactions with no gas: ΔH ≈ ΔU
Hess's law
ΔH_reaction = Σ ΔH_products − Σ ΔH_reactants (using standard enthalpies)
Enthalpy is a state function — path doesn't matter
Bond enthalpy method
ΔH_rxn = Σ BE(bonds broken) − Σ BE(bonds formed)
Energy is absorbed to break bonds and released to form bonds
Gibbs free energy
ΔG = ΔH − TΔS
T = temperature in K, ΔS = entropy change
Spontaneity condition
ΔG < 0 → spontaneous | ΔG = 0 → equilibrium | ΔG > 0 → non-spontaneous
Standard Gibbs energy and cell EMF
ΔG° = −nFE°
n = moles of electrons, F = 96485 C/mol
ΔG and equilibrium constant
ΔG° = −RT lnK
K > 1 → ΔG° < 0 → products favoured
First law of thermodynamics
ΔU = Q + W (IUPAC sign convention)
Note: W = −PΔV; work done ON system is positive
Ch 7Equilibrium(10 formulas)
Equilibrium constant Kc
Kc = [products]^stoich / [reactants]^stoich (at equilibrium)
Equilibrium constant Kp
Kp = Kc(RT)^Δn_g
R = 0.0821 L·atm/mol·K, T in K
Ionic product of water
Kw = [H⁺][OH⁻] = 1×10⁻¹⁴ at 25°C
Acid dissociation constant
Ka = [H⁺][A⁻] / [HA]
Base dissociation constant
Kb = [BH⁺][OH⁻] / [B]
Relation Ka × Kb
Ka × Kb = Kw (conjugate acid-base pair)
Stronger acid has weaker conjugate base
pH definition
pH = −log[H⁺] | pOH = −log[OH⁻] | pH + pOH = 14
Henderson-Hasselbalch equation
pH = pKa + log([A⁻]/[HA])
Used to find pH of buffer solutions
Degree of dissociation (weak acid)
α = √(Ka/C)
C = initial concentration of weak acid
Valid when α << 1
pH of weak acid
pH = ½(pKa − log C)
Easier form of Henderson equation for pure weak acid
Ch 8Redox Reactions(8 formulas)
Oxidation number: free element
ON = 0 (e.g. Fe, Cl₂, S₈)
Oxidation number: monoatomic ion
ON = charge of ion (e.g. Na⁺ → +1, Cl⁻ → −1)
Oxidation number: oxygen
ON(O) = −2 usually; −1 in peroxides (H₂O₂); +2 in OF₂
Oxidation number: hydrogen
ON(H) = +1 usually; −1 in metal hydrides (NaH)
n-factor (acid-base)
n-factor = basicity of acid or acidity of base
n-factor (redox)
n-factor = change in oxidation number per formula unit
In ionic equations: n-factor = total electrons transferred per formula unit
Equivalents (for calculations)
Equivalents = moles × n-factor | N = M × n-factor
Half-reaction method (acidic)
Add H₂O to balance O, then H⁺ to balance H, then e⁻ to balance charge