CLEP Chemistry cheat sheet

Atomic number Z = protons; mass number A = protons + neutrons; isotopes differ in neutrons only.

Electron configuration follows Aufbau order: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p.

Pauli exclusion: no two electrons share all four quantum numbers; each orbital holds at most 2 electrons with opposite spin.

Hund's rule: within a subshell, electrons singly occupy degenerate orbitals with parallel spins before pairing.

Periodic trends: atomic radius decreases left-to-right and increases down a group; ionization energy and electronegativity follow the reverse pattern.

Formal charge = (valence electrons) − (nonbonding electrons) − ½(bonding electrons).

Bond polarity follows electronegativity difference: <0.4 nonpolar covalent, 0.4–1.7 polar covalent, >1.7 largely ionic.

VSEPR electron-domain count sets geometry: 2 linear, 3 trigonal planar, 4 tetrahedral, 5 trigonal bipyramidal, 6 octahedral.

Ideal gas law: PV = nRT with R = 0.08206 L·atm/(mol·K) or 8.314 J/(mol·K).

Combined gas law: P1V1/T1 = P2V2/T2 at constant moles; temperatures must be in kelvin.

Dalton's law: total pressure equals the sum of partial pressures; partial pressure = mole fraction × total pressure.

Kinetic molecular theory: average kinetic energy is proportional to absolute temperature and independent of gas identity.

Graham's law: rate ratio for effusion = sqrt(M2/M1), the lighter gas effusing faster.

Real gases deviate from ideality at high pressure and low temperature, where molecular volume and attractions matter.

Phase changes: melting and vaporization absorb heat; freezing and condensation release heat; sublimation goes solid to gas directly.

Heating-curve plateaus occur at phase transitions because added heat changes potential energy, not temperature.

Synthesis: A + B → AB; decomposition: AB → A + B; single replacement: A + BC → AC + B.

Double replacement: AB + CD → AD + CB, driven by formation of a precipitate, gas, or weak electrolyte.

Combustion of a hydrocarbon with sufficient O2 yields CO2 and H2O; insufficient O2 yields CO or soot.

Acid–base neutralization: acid + base → salt + water; Brønsted acids donate protons, bases accept protons.

Redox is identified by changes in oxidation number; oxidation is an increase, reduction is a decrease.

Precipitation is predicted by solubility rules: most nitrates and alkali-metal salts are soluble; most carbonates and phosphates are not.

Activity series ranks metals by oxidation tendency; a metal displaces only ions of metals below it.

Disproportionation is a redox reaction in which one element is both oxidized and reduced.

Balance equations by changing coefficients only, never subscripts; total atoms of each element must match on both sides.

Mole bridge: mass = moles × molar mass; particles = moles × 6.022×10^23; gas volume at STP = moles × 22.4 L.

Mole ratio from balanced coefficients converts moles of one species to moles of another.

Limiting reactant is the one producing the least product; excess reactant remains after reaction.

Percent yield = (actual yield / theoretical yield) × 100; theoretical yield comes from the limiting reactant.

Percent composition = (mass of element in formula / molar mass of compound) × 100.

Empirical formula from percent composition: divide each mass by atomic mass, then divide by the smallest ratio.

Molarity M = moles of solute / liters of solution; dilution obeys M1V1 = M2V2.

Equilibrium constant K expresses products over reactants, each raised to its stoichiometric coefficient; pure solids and liquids are omitted.

For aA + bB ⇌ cC + dD, Kc = [C]^c[D]^d / ([A]^a[B]^b).

Reaction quotient Q compared with K predicts direction: Q < K shifts right, Q > K shifts left, Q = K at equilibrium.

Le Châtelier's principle: a system at equilibrium shifts to oppose a change in concentration, pressure, or temperature.

Raising temperature favors the endothermic direction; lowering temperature favors the exothermic direction.

For gases, Kp = Kc(RT)^Δn, where Δn is change in moles of gas.

Water ion product Kw = [H+][OH−] = 1.0×10^−14 at 25 °C.

pH = −log[H+]; pOH = −log[OH−]; pH + pOH = 14 at 25 °C.

For a weak acid, Ka = [H+][A−]/[HA]; smaller Ka means weaker acid.

Reaction rate is the change in concentration of a reactant or product per unit time, often divided by its coefficient.

Rate law: rate = k[A]^m[B]^n; exponents are determined experimentally, not from balanced coefficients.

Reaction order is the sum of exponents; zero order means rate is independent of that reactant's concentration.

Rate constant k depends on temperature and activation energy, not on concentration.

Arrhenius equation: k = Ae^(−Ea/RT); higher Ea or lower T gives a smaller k.

Catalysts lower activation energy and speed both forward and reverse reactions without being consumed.

Collision theory requires sufficient energy and proper orientation for a successful reaction.

Half-life of a first-order reaction is constant: t½ = 0.693/k.

First law: ΔE = q + w; for gas work at constant pressure, w = −PΔV.

Enthalpy change of reaction = sum of bond energies broken − sum of bond energies formed.

Hess's law: ΔH for a reaction is the same regardless of pathway; reverse a step to flip the sign of ΔH.

Exothermic reactions have ΔH < 0; endothermic reactions have ΔH > 0.

Standard enthalpy of formation of an element in its standard state is zero.

ΔH°rxn = ΣΔH°f(products) − ΣΔH°f(reactants).

Gibbs free energy: ΔG = ΔH − TΔS; negative ΔG means spontaneous, positive means nonspontaneous.

Entropy increases with greater disorder: gas > liquid > solid, more particles, and higher temperature.

Calorimetry: q = mcΔT, where c is specific heat capacity of the substance.

Alkali metals (Group 1) are highly reactive, form +1 ions, and react vigorously with water to give H2 and hydroxide.

Halogens (Group 17) form −1 ions and diatomic molecules; reactivity decreases down the group.

Noble gases (Group 18) have full valence shells and are largely unreactive.

Transition metals show variable oxidation states and form colored complexes; many are catalysts.

Common strong acids: HCl, HBr, HI, HNO3, H2SO4, HClO4; common strong bases: Group 1 hydroxides and Ba(OH)2.

Ammonia is a weak base; its conjugate acid is NH4+; it forms complexes with transition-metal ions.

Carbon dioxide is an acidic oxide; it forms carbonic acid in water and reacts with base to give carbonate.

Hydrogen bonding in water accounts for its high boiling point, high specific heat, and surface tension.

Significant figures: multiplication and division keep the least number of sig figs; addition and subtraction keep the least decimal places.

Precision describes reproducibility; accuracy describes closeness to the true value.

Percent error = |measured − accepted| / accepted × 100.

Titration equivalence point is where moles of acid equal moles of base per stoichiometry; an indicator marks the endpoint.

Filtration separates a solid from a liquid; distillation separates liquids by boiling point; chromatography separates by differential attraction.

Spectrophotometry uses absorbance proportional to concentration via Beer's law, A = εbc.

A volumetric flask is used to prepare a solution of exact concentration; a buret delivers variable, measured volumes.

Proper technique includes reading a meniscus at eye level and rinsing electrodes with deionized water.

Random errors scatter in both directions; systematic errors shift results consistently in one direction.