HESI A2 Chemistry cheat sheet

Matter is anything with mass and volume; a pure substance has fixed composition (element or compound), while a mixture's components keep their identities and can be separated physically.

Homogeneous mixtures are uniform throughout (salt water, air); heterogeneous mixtures show visibly distinct phases (sand in water, oil and water).

Physical changes alter state or form without changing composition (melting, dissolving, crushing); chemical changes produce new substances (burning, rusting, reacting with acid, color change or gas evolution).

Density = mass / volume; common units g/cm3 for solids, g/mL for liquids, and g/L for gases. Water's density near 4 C is 1.00 g/mL, so objects with density above 1.00 g/mL sink in water.

SI base units: meter (length), kilogram (mass), second (time), kelvin (temperature), mole (amount), ampere (current), candela (luminous intensity).

Kelvin conversion: K = C + 273.15 (use 273 for most exam calculations). Absolute zero is 0 K = -273.15 C.

Significant figures: nonzero digits and zeros between them are significant; leading zeros are not; trailing zeros count only when a decimal point is present. Add/subtract to the least decimal place; multiply/divide to the least number of significant figures.

Metric prefixes: kilo- 10^3, centi- 10^-2, milli- 10^-3, micro- 10^-6, nano- 10^-9. Dimensional analysis cancels units by multiplying by conversion factors equal to 1.

Classification flow: can it be separated by physical means? Yes = mixture (uniform = homogeneous, not uniform = heterogeneous). No = pure substance; can it be decomposed chemically? Yes = compound, no = element.

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

Isotopes share the same Z but differ in neutron count; average atomic mass is the weighted average of isotope masses by natural abundance.

Electron configuration order (Aufbau): 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p. Each orbital holds two electrons; s holds 2, p holds 6, d holds 10, f holds 14.

Maximum electrons in shell n is 2n^2. Valence electrons are those in the highest principal energy level; main-group elements have valence counts matching their group number (Groups 1-2, 13-18).

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

Groups: alkali metals (1, highly reactive, form +1 ions), alkaline earth metals (2, form +2 ions), halogens (17, form -1 ions), noble gases (18, full valence shell, unreactive).

Metals lose electrons to form cations, are malleable, ductile, and conductive; nonmetals gain or share electrons and are poor conductors; metalloids have intermediate properties.

Quantum numbers: n (size/level), l (shape: 0=s, 1=p, 2=d, 3=f), ml (orientation, -l to +l), ms (spin, +1/2 or -1/2). Pauli exclusion principle: no two electrons in an atom share all four quantum numbers.

Moles and particles: 1 mol = 6.022 x 10^23 particles (Avogadro's number). Molar mass in g/mol equals the atomic or formula mass in amu.

Ionic bonds transfer electrons between metal and nonmetal, forming a cation and anion held by electrostatic attraction; covalent bonds share electron pairs, usually between nonmetals.

Ionic compound formulas are electrically neutral: total positive charge equals total negative charge. Use the crisscross method to set subscripts, then reduce to the lowest whole-number ratio.

Binary ionic names: cation name first, then anion root with -ide (NaCl = sodium chloride, CaO = calcium oxide). Transition metals need a Roman numeral unless the charge is fixed (FeCl3 = iron(III) chloride).

Polyatomic ions to memorize: nitrate NO3-, sulfate SO4^2-, carbonate CO3^2-, phosphate PO4^3-, hydroxide OH-, ammonium NH4+, acetate C2H3O2-.

Covalent prefixes: mono-1, di-2, tri-3, tetra-4, penta-5, hexa-6. Examples: CO = carbon monoxide, CO2 = carbon dioxide, N2O4 = dinitrogen tetroxide.

Acids: binary acids use hydro- + root + -ic acid (HCl = hydrochloric acid); oxyacids with -ate become -ic acid and -ite become -ous acid (HNO3 = nitric acid, HNO2 = nitrous acid).

VSEPR shapes from bonding pairs plus lone pairs on the central atom: 2 = linear, 3 = trigonal planar, 4 = tetrahedral (109.5 degrees), 3 with one lone pair = trigonal pyramidal, 2 with two lone pairs = bent.

Polarity: a bond is polar when electronegativity difference is roughly 0.4-1.7; above about 1.7 is largely ionic. A molecule is polar only if polar bonds are arranged asymmetrically.

Intermolecular forces from weakest to strongest: London dispersion, dipole-dipole, hydrogen bonding. Hydrogen bonds occur when H is bonded to N, O, or F.

Balance equations by changing coefficients only, never subscripts. Conservation of mass requires equal atoms of each element on both sides.

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

Mole ratio from a balanced equation converts moles of one substance to moles of another. General flow: mass -> moles (divide by molar mass) -> mole ratio -> moles -> mass (multiply by molar mass).

Limiting reactant is fully consumed and sets the maximum product; the excess reactant is left over. Compare the mole ratios of each reactant to the coefficients to identify the limiting one.

Percent yield = (actual yield / theoretical yield) x 100. Theoretical yield comes from stoichiometry; actual yield is measured.

Molarity M = moles of solute / liters of solution. Dilution: M1 x V1 = M2 x V2, because moles of solute stay constant.

Solubility rules: most Group 1 and ammonium salts are soluble; most nitrates and acetates are soluble; most chlorides are soluble except AgCl, PbCl2, and Hg2Cl2; most carbonates, phosphates, and hydroxides are insoluble except with Group 1 or ammonium.

pH = -log[H+]; pOH = -log[OH-]; pH + pOH = 14 at 25 C. pH below 7 is acidic, above 7 is basic. Strong acids and bases dissociate completely.

Colligative property examples: adding solute lowers freezing point and raises boiling point relative to the pure solvent; osmosis is water moving across a membrane toward higher solute concentration.

Ideal gas law: PV = nRT, with R = 0.0821 L atm / (mol K) or 8.314 J / (mol K). Temperature must be in kelvin.

Combined gas law: P1V1/T1 = P2V2/T2. At constant temperature this reduces to Boyle's law (P1V1 = P2V2); at constant pressure to Charles's law (V1/T1 = V2/T2).

STP (0 C, 1 atm) gives 22.4 L per mole of any ideal gas. Dalton's law of partial pressures: total pressure is the sum of the partial pressures.

Exothermic reactions release heat (delta H negative); endothermic reactions absorb heat (delta H positive). Breaking bonds absorbs energy; forming bonds releases energy.

Specific heat formula: q = m x c x delta T. Water's specific heat is 4.184 J/(g C), meaning water resists temperature change more than most substances.

Calorimetry: heat gained by the surroundings equals heat lost by the system in an insulated container, so q of reaction = -(q of water).

Kinetics: reaction rate rises with higher temperature, higher concentration, higher surface area, and added catalyst. A catalyst lowers activation energy without being consumed.

Equilibrium (Le Chatelier): adding reactant shifts right, adding product shifts left; increasing pressure shifts toward fewer gas moles; a catalyst changes only the rate, not the position.

Percent by mass = (mass of solute / mass of solution) x 100. Percent by volume = (volume of solute / volume of solution) x 100.

Mole fraction of a component = moles of that component / total moles of all components; all mole fractions sum to 1.

Molality m = moles of solute / kilograms of solvent; unlike molarity, molality does not change with temperature because it uses mass.

ppm = (mass of solute / mass of solution) x 10^6. Useful for very dilute solutions such as water contaminants.

Electrolytes dissociate into ions in water and conduct electricity (salts, strong acids, strong bases); nonelectrolytes dissolve without ions (sugar, ethanol).

A saturated solution holds the maximum solute at that temperature; solubility of most solids rises with temperature, while gas solubility falls as temperature rises.

Henry's law: gas solubility is proportional to the partial pressure of that gas above the liquid. Higher pressure means more dissolved gas.

Lab safety rules: add acid to water, not water to acid; wear goggles; never pipette by mouth; read the meniscus at eye level; heat test tubes pointing away from people.

Precision describes how closely repeated measurements agree; accuracy describes how close a measurement is to the true value. Random error affects precision; systematic error affects accuracy.

Mole map: mass <-> moles via molar mass; particles <-> moles via 6.022 x 10^23; gas volume at STP <-> moles via 22.4 L/mol; solution volume <-> moles via molarity.

Percent composition of an element = (total mass of that element in the formula / molar mass of the compound) x 100.

Empirical formula from percent composition: assume 100 g, convert grams to moles, divide all by the smallest mole value, then round to whole numbers.

Molecular formula = (empirical formula) x n, where n = molar mass / empirical formula mass.

Dimensional analysis: set up conversion factors so unwanted units cancel diagonally. Example: 25.0 g NaCl x (1 mol / 58.44 g) = 0.428 mol.

Molar mass of a compound sums the atomic masses of each atom in the formula. Example: H2O = 2(1.008) + 16.00 = 18.02 g/mol.

When converting units, state the given quantity with units, multiply by ratios equal to 1, and check that the final unit matches what was asked.

Common conversion: 1 L = 1000 mL; 1 kg = 1000 g; 1 cm3 = 1 mL; 1 atm = 760 mmHg = 101.3 kPa.

The periodic table is ordered by increasing atomic number. Rows are periods (energy levels); columns are groups (similar valence electron configurations).

Blocks: s-block (Groups 1-2 plus He), p-block (Groups 13-18), d-block (transition metals, Groups 3-12), f-block (lanthanides and actinides).

Metallic character increases down a group and decreases across a period; nonmetallic character does the opposite.

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

Half-life: after n half-lives, the remaining fraction is (1/2)^n. Example: after 3 half-lives, 1/8 of the original sample remains.

Fission splits a heavy nucleus and releases energy; fusion combines light nuclei and powers stars. Both release energy because the products are more tightly bound.

Ion charge prediction: Group 1 forms +1, Group 2 forms +2, Group 16 forms -2, Group 17 forms -1, because atoms gain or lose electrons to reach a noble-gas configuration.

Electron affinity and ionization energy generally increase across a period, with noble gases having the highest ionization energies because of full valence shells.

Evidence of a chemical reaction: color change, gas bubbles, precipitate formation, temperature change, or light emission.

A precipitate is an insoluble solid that forms when two solutions are mixed. Predict it by combining the cation of one reactant with the anion of the other and applying solubility rules.

Net ionic equations show only the species that change; spectator ions appear on both sides and are canceled.

Aqueous strong acids and bases are written as separate ions in ionic equations; weak acids and molecular compounds stay intact.

Combustion of a hydrocarbon always yields CO2 and H2O when oxygen is sufficient. Balance carbon, then hydrogen, then oxygen last.

Oxidation is loss of electrons (oxidation number increases); reduction is gain of electrons (oxidation number decreases). In redox reactions both occur together.

Oxidation number rules: elements in their standard state are 0; monatomic ions equal their charge; oxygen is usually -2; hydrogen is usually +1; the sum equals the overall charge.

Activity series: a metal higher on the series displaces a metal lower on it from a compound, which determines whether a single replacement reaction occurs.

Read a graduated cylinder at the bottom of the meniscus at eye level. Estimate one digit beyond the smallest marked increment.

Percent error = |measured value - accepted value| / accepted value x 100.

Random error scatters results in both directions; systematic error shifts all results in one direction, such as a mis-calibrated balance.

Precision is expressed by how close repeated trials are to each other; report it with standard deviation or a range.

Temperature conversions: F = (C x 9/5) + 32; C = (F - 32) x 5/9; K = C + 273.15.

Scientific notation: a number is written as a x 10^n where 1 <= a < 10. Multiplication adds exponents; division subtracts them.

Rounding rules: if the digit dropped is 5 or more, round up; if less than 5, leave the digit unchanged.

An exact number, such as a counted object or a defined conversion factor, has unlimited significant figures and does not limit the answer.

Arrhenius acid releases H+ in water; Arrhenius base releases OH-. Bronsted-Lowry acid donates a proton; base accepts one.

Strong acids include HCl, HBr, HI, HNO3, H2SO4, and HClO4. Strong bases include Group 1 hydroxides and Ca(OH)2, Ba(OH)2.

pH = -log[H+]; [H+] = 10^(-pH). Every pH unit represents a tenfold change in hydrogen ion concentration.

Neutral water at 25 C has [H+] = [OH-] = 1 x 10^-7 M, so pH = 7. Kw = [H+][OH-] = 1 x 10^-14.

Titration: at the equivalence point, moles of H+ equal moles of OH-. Use M_acid x V_acid x n_acid = M_base x V_base x n_base.

Neutralization reaction: acid + base -> salt + water. Example: HCl + NaOH -> NaCl + H2O.

A buffer resists pH change and consists of a weak acid with its conjugate base (or a weak base with its conjugate acid).

For solution stoichiometry, convert volume and molarity to moles first: moles = M x V (liters), then use the balanced-equation mole ratio.

Gas stoichiometry at STP: convert moles to liters using 22.4 L/mol after applying the mole ratio.

At non-STP conditions, use the ideal gas law PV = nRT to convert between moles and volume.

Limiting reactant problems in solution: calculate moles of each reactant from M x V, divide each by its coefficient, and the smallest quotient is limiting.

Mass of product from a solution: moles of limiting reactant -> mole ratio -> moles of product -> multiply by molar mass.

Percent by mass from solution stoichiometry: divide the mass of the dissolved solute by the total mass of the solution and multiply by 100.

Dilution before reaction changes molarity but not moles of solute; recalculate M after dilution using M1V1 = M2V2 before applying the mole ratio.

Dalton's atomic theory: matter is made of indivisible atoms; atoms of one element are identical; compounds form in fixed whole-number ratios.

Thomson's cathode-ray experiment discovered the electron and led to the plum-pudding model of a positive sphere with embedded electrons.

Rutherford's gold-foil experiment showed that most alpha particles pass through but some deflect sharply, implying a small dense positive nucleus.

Bohr model: electrons occupy fixed energy levels and emit or absorb photons of specific energy when they move between levels.

The quantum mechanical model describes electrons as probability clouds (orbitals) rather than fixed orbits.

The photoelectric effect and line spectra supported quantized energy, not the continuous classical model.

Atomic emission spectra are unique to each element, which is the basis for identifying elements by spectroscopy.

London dispersion forces exist in all molecules and grow stronger with more electrons and larger surface area.

Dipole-dipole forces occur between polar molecules; hydrogen bonding is a strong dipole interaction when H is bonded to N, O, or F.

Stronger intermolecular forces raise boiling point, melting point, surface tension, and viscosity.

Ion-dipole forces occur between ions and polar molecules and help explain why ionic compounds dissolve in water.

Like dissolves like: polar solvents dissolve polar and ionic solutes; nonpolar solvents dissolve nonpolar solutes.

Vapor pressure rises with weaker intermolecular forces and with higher temperature.

Hydrogen bonding in water explains its unusually high boiling point, high specific heat, and high surface tension.

Collision theory: reactions occur when particles collide with sufficient energy and proper orientation.

Activation energy is the minimum energy needed to start a reaction; a catalyst lowers it without changing the overall energy change.

Increasing temperature raises the fraction of collisions with enough energy, so rate increases.

Increasing concentration or surface area increases collision frequency and rate.

At equilibrium, forward and reverse reaction rates are equal and concentrations remain constant, but both reactions continue.

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

Equilibrium constant K expresses the ratio of product to reactant concentrations at equilibrium; a large K favors products.

Enthalpy change delta H is negative for exothermic reactions and positive for endothermic reactions.

Bond breaking absorbs energy; bond formation releases energy. Net delta H is the difference between energy absorbed and released.

q = m x c x delta T calculates heat absorbed or released by a substance.

In a calorimeter, heat lost by the reaction equals heat gained by the water, so q_reaction = -(m x c x delta T) of the water.

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

Hess's law: the total enthalpy change is the same regardless of the pathway, so reactions can be added to find an unknown delta H.

Energy diagrams show reactants, products, activation energy, and delta H; exothermic reactions have products lower in energy than reactants.

Write formulas from names by identifying the cation and anion, then balancing total charges.

When a polyatomic ion appears more than once in a formula, enclose it in parentheses with a subscript, as in Ca(OH)2.

Binary molecular compounds use prefixes, but the first element never uses mono-; the second element ends in -ide.

Roman numerals in a name indicate the charge of a variable-charge metal, as in copper(II) sulfate = CuSO4.

Hydrates include water molecules in the formula and are named with a prefix and hydrate, as in CuSO4 x 5H2O = copper(II) sulfate pentahydrate.

Common acids: HCl hydrochloric, H2SO4 sulfuric, HNO3 nitric, H3PO4 phosphoric, CH3COOH acetic.

Check a formula's neutrality by summing positive and negative charges to zero.

Moles = mass / molar mass, and mass = moles x molar mass.

Molarity M = moles / liters; moles = M x liters; liters = moles / M.

To convert molarity to percent by mass, find the mass of solute using molar mass, then divide by the total mass of solution.

Molality uses kilograms of solvent, not liters of solution, so it is temperature independent.

To prepare a solution, dissolve the calculated moles of solute in enough solvent to reach the final volume.

Dilution: M1V1 = M2V2; the moles of solute are unchanged, only the volume increases.

Mixing two solutions of the same solute: total moles = M1V1 + M2V2 and final molarity = total moles / total volume.

Elements in the same group have the same number of valence electrons and similar bonding behavior.

Metals bond with nonmetals by transferring electrons to form ionic compounds; nonmetals bond with nonmetals by sharing electrons to form covalent compounds.

Electronegativity difference predicts bond type: 0 to about 0.4 nonpolar covalent, 0.4 to 1.7 polar covalent, above 1.7 ionic.

Metallic bonding involves delocalized electrons and explains conductivity, malleability, and luster.

Noble gases are stable because their valence shells are full, so they rarely form compounds.

Ion size: cations are smaller than their neutral atom because electrons are lost; anions are larger because electrons are gained.

The octet rule states that atoms gain, lose, or share electrons to achieve eight valence electrons, with hydrogen and helium stabilizing at two.

Filtration separates an insoluble solid from a liquid by particle size.

Distillation separates liquids by differences in boiling point.

Evaporation removes a volatile solvent to recover a dissolved solid.

Chromatography separates components by differing attraction to a stationary phase and a mobile phase.

Decanting pours off a liquid while leaving a settled solid behind.

Centrifugation spins a mixture so denser particles settle faster.

Physical separation methods do not change the chemical identities of the substances involved.