NY Regents Chemistry cheat sheet

Matter is classified as a substance (element or compound, fixed composition) or a mixture (homogeneous or heterogeneous, variable composition); physical changes do not alter composition, chemical changes do.

Significant figures: nonzero digits and trapped zeros count; leading zeros never count; trailing zeros count only with a decimal point. Addition/subtraction round to the fewest decimal places; multiplication/division to the fewest significant figures.

Density = mass/volume; it is an intensive property, so it does not change with sample size and can be used as a conversion factor between mass and volume.

Common metric conversions: 1 kg = 1000 g, 1 L = 1000 mL, 1 cm^3 = 1 mL, 1 kJ = 1000 J, 1 mol = 6.02 x 10^23 particles (Avogadro's number).

Percent error = |accepted - measured| / accepted x 100; percent composition by mass = (mass of component / total mass) x 100.

Temperature scales: K = degrees C + 273; a temperature change of 1 K equals a change of 1 degree C, so temperature differences convert without the 273 offset.

Heat flow q = mc(delta T); a positive q means heat is absorbed (or the surroundings lost it), and specific heat capacity is an intensive property in J/(g degrees C).

Separations rely on physical property differences: filtration (particle size), distillation (boiling point), chromatography (solubility/attraction to phases), and evaporation (volatility).

Atomic number Z = number of protons; in a neutral atom electrons = protons; mass number A = protons + neutrons, so neutrons = A - Z.

Isotopes share Z but differ in A; average atomic mass = sum of (isotopic mass x fractional abundance), which is why it is not a whole number.

Electron configuration fills in order of increasing energy: 1s 2s 2p 3s 3p 4s 3d 4p; the valence shell is the outermost shell (highest principal quantum number).

Excited-state configurations violate the normal filling order or show an electron promoted to a higher sublevel; they are temporary and release energy as photons when the electron falls back.

A neutral atom's electron configuration can be written as the nearest noble gas core plus remaining electrons; ions gain or lose electrons without changing the nucleus.

Periodic trends: atomic radius decreases left to right and increases down a group; first ionization energy and electronegativity increase left to right and decrease down.

Metals lose electrons to form cations, nonmetals gain electrons to form anions, and Group 18 (noble gases) have full valence shells and low reactivity.

Groups: alkali metals (1), alkaline earth metals (2), halogens (17), noble gases (18); elements in the same group share valence electron count and similar chemistry.

Ionic bonds form between metals and nonmetals by electron transfer; the resulting lattice of cations and anions gives high melting points and conductivity when molten or dissolved.

Covalent bonds form between nonmetals by electron sharing; bond order and bond energy rise from single to double to triple bonds while bond length decreases.

Electronegativity difference guides classification: roughly 0-0.4 nonpolar covalent, 0.4-1.7 polar covalent, above 1.7 largely ionic (a rule of thumb, not a strict cutoff).

A polar bond has unequal sharing; a molecule is polar only if its bond dipoles do not cancel, so symmetric shapes such as linear CO2 and tetrahedral CH4 are nonpolar.

VSEPR electron-domain counts give shapes: 2 linear, 3 trigonal planar, 4 tetrahedral, with lone pairs reducing bond angles (for example about 105 degrees in water).

Intermolecular forces ranked by strength: hydrogen bonding (H bonded to N, O, or F) > dipole-dipole > London dispersion; all three are weaker than covalent or ionic bonds.

Hydrogen bonding raises boiling point and surface tension, which is why water boils far above H2S despite its lower molar mass.

Network solids (diamond, quartz, ionic crystals) are held together by continuous covalent or ionic bonding rather than discrete molecules, giving very high melting points.

Balance equations by changing coefficients only; the number of atoms of each element must be equal on both sides, and charges must balance in ionic equations.

Reaction types: synthesis (A + B -> AB), decomposition (AB -> A + B), single replacement (A + BC -> AC + B), double replacement (AB + CD -> AD + CB), combustion of a hydrocarbon yielding CO2 and H2O.

Mole conversions: moles = mass / molar mass; moles = particles / 6.02 x 10^23; at STP, moles = volume of gas / 22.4 L.

Stoichiometry path: balance the equation, convert given quantity to moles, apply the mole ratio from coefficients, convert to the requested unit.

The limiting reactant is consumed first and determines the maximum product; the excess reactant remains, and theoretical yield is computed from the limiting reactant.

Percent yield = (actual yield / theoretical yield) x 100; percent error compares a measured value with an accepted value.

Solubility rules: most nitrate, sodium, potassium, and ammonium salts are soluble; most chlorides are soluble except AgCl, PbCl2, and Hg2Cl2; most carbonates and hydroxides are insoluble.

Net ionic equations remove spectator ions; a precipitate, gas, or weak electrolyte (such as water) must appear on the product side for a double replacement to occur.

Collision theory: reaction rate rises with concentration, surface area, temperature, and catalysts; temperature increases both collision frequency and the fraction of collisions with energy at least the activation energy.

A catalyst lowers activation energy by providing an alternate pathway and is not consumed; it does not change the enthalpy change or the equilibrium position.

Potential energy diagrams show activation energy as the peak minus the reactant energy; exothermic reactions release energy (products lower), endothermic absorb it (products higher).

Le Chatelier's principle: adding reactant shifts right, adding product shifts left, increasing pressure shifts toward fewer moles of gas, and heating shifts an endothermic reaction right.

Equilibrium is dynamic: forward and reverse rates are equal and concentrations are constant, but the reaction has not stopped; K expresses the ratio of product to reactant concentrations at equilibrium.

Enthalpy: delta H is negative for exothermic and positive for endothermic; bond breaking absorbs energy and bond formation releases energy.

Hess's law: the enthalpy change of an overall reaction is the sum of the enthalpy changes of its steps, allowing delta H to be found from known reactions.

Entropy (disorder) tends to increase; a spontaneous process at constant temperature and pressure is favored by negative delta H and positive delta S, consistent with delta G = delta H - T delta S.

Hydrocarbons: alkanes are saturated with single bonds (C_nH_2n+2), alkenes have one double bond (C_nH_2n), alkynes have one triple bond (C_nH_2n-2); each carbon forms four bonds.

Homologous series differ by one CH2 unit; isomers share a molecular formula but differ in structure, giving different physical and chemical properties.

Functional groups determine properties: alcohols (-OH), organic acids (-COOH), aldehydes (-CHO), ketones (C=O), esters (-COO-), amines (-NH2), and halocarbons (-X).

IUPAC naming: find the longest carbon chain, number to give the lowest locants to substituents or multiple bonds, and end with -ane, -ene, or -yne.

Organic reactions: substitution (alkane with halogen), addition (alkene or alkyne adding H2, Br2, or H2O), polymerization (monomers joining), and esterification (acid plus alcohol).

Alpha decay emits a helium nucleus (mass number -4, atomic number -2); beta decay converts a neutron to a proton (mass number unchanged, atomic number +1); gamma emission changes neither.

Nuclear reactions conserve mass number and atomic number, and mass defect converts to energy by E = mc^2; nuclear changes release far more energy than chemical changes.

Half-life is the time for half the radioactive nuclei to decay; after n half-lives the remaining fraction is (1/2)^n, and artificial transmutation requires a bombarding particle.