AEPA Chemistry NT306 cheat sheet

Scientific method sequence: observation, question, hypothesis, controlled experiment, data analysis, conclusion, peer review, replication.

Independent variable is manipulated by the researcher; dependent variable is measured; controlled variables are held constant.

A hypothesis must be falsifiable: a specific, testable prediction that could be shown false by data.

Accuracy is closeness to the true value; precision is repeatability of measurements; random error affects precision, systematic error affects accuracy.

Significant figures: nonzero digits count, interior zeros count, leading zeros do not, trailing zeros count only with a decimal point.

Percent error = |accepted - experimental| / accepted x 100%.

A theory is a well-tested explanatory framework supported by many lines of evidence; a law describes a consistent relationship, often mathematically.

Correlation does not establish causation; controlled experiments or strong mechanistic evidence are required for causal claims.

SI base units: meter, kilogram, second, ampere, kelvin, mole, candela.

Pure substances are elements or compounds; mixtures are homogeneous (uniform, one phase) or heterogeneous (nonuniform phases).

Physical changes alter state or form without changing composition; chemical changes produce new substances with different properties.

Intensive properties (density, boiling point, color) do not depend on amount; extensive properties (mass, volume) do.

Density = mass / volume; a substance floats in a fluid of greater density and sinks in one of lesser density.

Separation methods: filtration (insoluble solid), distillation (boiling point difference), chromatography (differential affinity), evaporation.

Diatomic elements: H2, N2, O2, F2, Cl2, Br2, I2.

Naming ionic compounds: cation name first, anion with -ide suffix; use Roman numerals for variable-charge metals such as Fe2+ and Fe3+.

Percent composition by mass = (mass of element in formula / molar mass of compound) x 100%.

Molar mass is the sum of atomic masses in grams per mole; Avogadro's number is 6.022 x 10^23 particles per mole.

Reaction types: synthesis (A + B -> AB), decomposition, single replacement, double replacement, combustion.

Balancing rules: change coefficients only, never subscripts; conserve atoms of each element on both sides.

Solubility rules: most nitrates and alkali-metal salts are soluble; most chlorides soluble except AgCl, PbCl2, Hg2Cl2; most carbonates and phosphates insoluble except with alkali metals or ammonium.

Net ionic equations remove spectator ions; only species that change appear.

Acid-base neutralization: acid + base -> salt + water; strong acids and bases dissociate completely.

pH = -log[H+]; pOH = -log[OH-]; pH + pOH = 14 at 25 degrees C.

Oxidation is loss of electrons (increase in oxidation number); reduction is gain of electrons (decrease in oxidation number).

Limiting reactant is consumed first and determines maximum product; excess reactant remains after reaction.

Percent yield = (actual yield / theoretical yield) x 100%.

First law: energy is conserved; delta E = q + w, where w = -P delta V for pressure-volume work.

Enthalpy change: delta H negative for exothermic, positive for endothermic; Hess's law sums step enthalpies to get the overall reaction.

Bond breaking absorbs energy; bond formation releases energy; net delta H reflects the balance.

Calorimetry: q = mc delta T; for a calorimeter use the heat capacity C and q = C delta T.

Gibbs free energy: delta G = delta H - T delta S; spontaneous when delta G is negative.

Entropy increases with greater disorder: solid to liquid to gas, more moles of gas, higher temperature.

Reaction rate increases with higher concentration, higher temperature, greater surface area, and presence of a catalyst.

A catalyst lowers activation energy and speeds both forward and reverse reactions without changing delta H or equilibrium position.

Equilibrium: rates of forward and reverse reactions are equal; Le Chatelier's principle predicts shifts when concentration, pressure, or temperature changes.

Atomic number Z equals protons; mass number A equals protons plus neutrons; isotopes differ in neutron count.

Electron configuration follows Aufbau order, Pauli exclusion (two electrons per orbital, opposite spins), and Hund's rule (maximize unpaired spins).

Orbital shapes: s spherical, p dumbbell (three orientations), d five orientations; each orbital holds two electrons.

Periodic trends: atomic radius decreases across a period and increases down a group; ionization energy and electronegativity generally increase across a period and decrease down a group.

Valence electrons determine group chemistry; main-group elements follow the octet rule, with exceptions such as H (duet) and expanded octets for period 3 and beyond.

Ionic bonds form by electron transfer between metals and nonmetals; covalent bonds share electrons between nonmetals; metallic bonds have delocalized electrons.

VSEPR: electron pairs repel and spread out; lone pairs occupy more space, reducing bond angles below ideal values.

Polarity depends on electronegativity differences and molecular geometry; symmetric geometries can cancel bond dipoles and yield nonpolar molecules.

Intermolecular forces in increasing strength: London dispersion < dipole-dipole < hydrogen bonding; these govern boiling point and solubility trends.