Bagrut Chemistry cheat sheet

Quantum numbers: n = shell (1,2,3…), ℓ = 0…n−1 (s,p,d,f), mℓ = −ℓ…+ℓ, ms = ±½; no two electrons in an atom share all four.

Orbital filling follows Aufbau (lowest n+ℓ first; for equal n+ℓ, lower n first), Pauli exclusion, and Hund's rule (maximize unpaired spins within a subshell).

Electron configuration exceptions to memorize: Cr = [Ar]3d⁵4s¹, Cu = [Ar]3d¹⁰4s¹; 4s fills before 3d but 3d electrons are removed first in cations.

Effective nuclear charge Z_eff = Z − S; across a period Z_eff rises, so atomic radius decreases and first ionization energy generally increases.

Down a group: n increases, radius increases, ionization energy decreases, metallic character increases, electron affinity magnitude generally decreases.

Successive ionization energies jump sharply when a new inner shell is breached - the size of the jump identifies the group number.

Periodic anomalies: IE(Be) > IE(B) and IE(N) > IE(O) because of filled 2s² and half-filled 2p³ stability.

Electron affinity is the energy released when one electron is added to a gaseous atom; noble gases have ~zero or negative (unfavorable) values.

Electronegativity difference ΔEN: <0.4 nonpolar covalent, 0.4–1.7 polar covalent, >1.7 predominantly ionic (Pauling scale, qualitative only).

VSEPR: electron domains repel and spread; lone pairs repel more than bonding pairs, compressing bond angles (CH₄ 109.5° > NH₃ 107° > H₂O 104.5°).

Hybridization from steric number: 2 → sp (linear), 3 → sp² (trigonal planar), 4 → sp³ (tetrahedral); lone pairs occupy hybrid orbitals.

Polar bond ≠ polar molecule: CO₂ and CCl₄ are nonpolar because bond dipoles cancel by symmetry; H₂O and NH₃ are polar.

Intermolecular force strength: ion–dipole > hydrogen bond > dipole–dipole > London dispersion; dispersion grows with polarizability (electron count, size).

Hydrogen bonding requires H bonded to N, O, or F; it explains water's high boiling point, surface tension, and ice's lower density.

Bond order from MO theory: (bonding − antibonding)/2; O₂ has bond order 2 and is paramagnetic because of two unpaired π* electrons.

Lattice energy (Born–Landé) increases with higher ionic charges and smaller ionic radii; it drives melting point and solubility trends.

1 mol = 6.022×10²³ particles; molar mass in g/mol equals the sum of atomic masses in the formula.

Mole conversions: n = m/M, n = N/N_A, and for gases at STP (0 °C, 1 bar or 1 atm) use the stated molar volume (22.7 L at 1 bar; 22.4 L at 1 atm).

Mole ratio from a balanced equation converts between substances; identify the limiting reagent by dividing moles by stoichiometric coefficient and taking the smallest.

Percent yield = (actual/theoretical)×100; theoretical yield derives from the limiting reagent only.

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

Empirical formula from mass %: divide each % by atomic mass, then divide by the smallest ratio; multiply to integers. Molecular formula = empirical × (M_molecular/M_empirical).

Solution concentration: c = n/V (mol/L); dilution obeys c₁V₁ = c₂V₂; mass % = (solute mass / solution mass) × 100.

For a reaction in solution, moles = c × V(L); convert volumes to liters before multiplying.

Ideal gas law PV = nRT with R = 0.08206 L·atm/(mol·K) or 8.314 J/(mol·K); T must be in kelvin.

Combined gas law P₁V₁/T₁ = P₂V₂/T₂ for fixed moles; Boyle (T const), Charles (P const), Gay-Lussac (V const) are special cases.

Dalton's law: P_total = ΣP_i; partial pressure P_i = x_i·P_total where x_i is mole fraction.

Kinetic molecular theory: average kinetic energy ∝ T (KE = 3/2 kT per molecule); root-mean-square speed ∝ √(T/M), so lighter gases effuse faster (Graham's law).

Real gases deviate from ideal behavior at high pressure and low temperature; van der Waals corrects for molecular volume (b) and attraction (a).

Phase changes: melting/freezing, vaporization/condensation, sublimation/deposition; during a phase change at constant pressure, temperature stays constant while latent heat is absorbed or released.

Heating curve: q = mcΔT for a single phase, q = nΔH_fus or nΔH_vap during a phase change; sum the segments for total heat.

Phase diagram: triple point (all three phases coexist), critical point (liquid–gas boundary ends); the slope of the solid–liquid line is negative for water.

At equilibrium, forward and reverse rates are equal and concentrations are constant (not necessarily equal); the system is dynamic at the molecular level.

Equilibrium constant K = [products]^coeff / [reactants]^coeff (activities of pure solids and liquids = 1); K depends only on temperature.

Reaction quotient Q compared to K: Q < K shifts right (forward), Q > K shifts left (reverse), Q = K at equilibrium.

Le Chatelier: adding reactant or removing product shifts right; increasing pressure shifts toward the side with fewer gas moles; increasing temperature shifts in the endothermic direction.

A catalyst lowers activation energy and speeds both directions equally - it does not change K or equilibrium position.

For K_p and K_c: K_p = K_c(RT)^Δn, where Δn = moles of gaseous products − gaseous reactants.

Manipulating K: reversing a reaction gives 1/K; multiplying coefficients by n gives K^n; adding reactions multiplies their K values.

A very small K (<10⁻³) means reactants dominate at equilibrium; a very large K (>10³) means products dominate.

Brønsted–Lowry: acid donates a proton, base accepts one; conjugate acid–base pairs differ by one H⁺.

pH = −log[H₃O⁺]; pOH = −log[OH⁻]; pH + pOH = 14 at 25 °C; [H₃O⁺][OH⁻] = Kw = 1.0×10⁻¹⁴.

Strong acids (HCl, HBr, HI, HNO₃, H₂SO₄, HClO₄) and strong bases (group 1 hydroxides, Ca(OH)₂, Ba(OH)₂) dissociate completely; weak ones use Ka or Kb.

For a weak acid: Ka·Kb = Kw for a conjugate pair; pKa = −log Ka; larger Ka/pKa smaller means stronger acid.

Henderson–Hasselbalch: pH = pKa + log([A⁻]/[HA]); a buffer is most effective when pH ≈ pKa (ratio 0.1–10) and resists pH change on adding small amounts of acid or base.

Titration: equivalence point is where moles acid = moles base stoichiometrically; at half-equivalence for a weak acid, pH = pKa.

Salt hydrolysis: strong acid + strong base → neutral; weak acid + strong base → basic (anion hydrolyzes); strong acid + weak base → acidic (cation hydrolyzes).

Solubility product Ksp = [cation]^m[anion]^n for a sparingly soluble salt; a precipitate forms when the ion product Q exceeds Ksp.

First law: ΔU = q + w; at constant pressure, enthalpy change ΔH = q_p and ΔH = ΔU + Δ(PV).

Exothermic: ΔH < 0 (releases heat); endothermic: ΔH > 0 (absorbs heat).

Hess's law: ΔH for a reaction is independent of pathway; sum ΔH of steps (reversing a step flips its sign, scaling multiplies ΔH).

Standard enthalpy of formation ΔH°f: ΔH°rxn = ΣΔH°f(products) − ΣΔH°f(reactants); ΔH°f of an element in its standard state = 0.

Calorimetry: q = mcΔT for solutions (c_water ≈ 4.18 J/(g·°C)); for a bomb calorimeter q = C_cal·ΔT, where C_cal is heat capacity.

Bond enthalpy: ΔH ≈ Σ(bonds broken) − Σ(bonds formed); breaking bonds is endothermic, forming bonds is exothermic.

Gibbs free energy: ΔG = ΔH − TΔS; spontaneous when ΔG < 0, and ΔG° = −RT ln K.

Entropy ΔS > 0 favors more microstates: increasing moles of gas, temperature, or disorder (solid → liquid → gas).

Oxidation is loss of electrons (oxidation number rises); reduction is gain of electrons (oxidation number falls); OIL RIG.

Oxidation number rules: free element = 0; monatomic ion = charge; O usually −2 (peroxide −1); H usually +1 (hydride −1); sum in a neutral compound = 0.

Balance redox in acidic solution with H₂O and H⁺; in basic solution add OH⁻ to neutralize H⁺ (or balance as if acidic then add OH⁻ to both sides).

Galvanic (voltaic) cell: oxidation at anode (negative), reduction at cathode (positive); electrons flow anode → cathode through the external circuit.

Cell potential E°cell = E°cathode − E°anode (both as reduction potentials); a positive E°cell means spontaneous (ΔG° = −nFE°cell).

Electrolytic cell: external power drives a nonspontaneous reaction; the anode is positive and the cathode negative (opposite of galvanic).

Faraday's law: moles of electrons = It/F with F ≈ 96,485 C/mol; mass deposited = (M·It)/(nF) where n is electrons per ion.

Nernst equation: E = E° − (RT/nF)ln Q, or E = E° − (0.0592/n)log Q at 25 °C.

Rate = −Δ[reactant]/Δt = Δ[product]/Δt, divided by stoichiometric coefficients; units are mol·L⁻¹·s⁻¹.

Rate law rate = k[A]^m[B]^n is determined experimentally; orders m and n are not taken from the balanced equation.

Arrhenius: k = A·e^(−Ea/RT), so ln k = ln A − Ea/RT; a plot of ln k vs 1/T has slope −Ea/R.

Half-life: first order t₁/₂ = 0.693/k (independent of concentration); zero order t₁/₂ = [A]₀/(2k); second order t₁/₂ = 1/(k[A]₀).

Collision theory: reaction requires collisions with sufficient energy (≥ Ea) and correct orientation; a catalyst provides an alternative lower-Ea pathway.

A catalyst increases rate without being consumed and does not alter ΔH or the equilibrium constant.

Rate-determining step is the slowest elementary step; the overall rate law matches that step's reactants (and any prior fast equilibria).

Increasing temperature raises both collision frequency and the fraction of molecules with E ≥ Ea; the exponential energy factor dominates.

Homologous series and general formulas: alkanes CnH2n+2, alkenes/cycloalkanes CnH2n, alkynes CnH2n−2, alcohols CnH2n+1OH, carboxylic acids CnH2nO₂.

Functional groups: alcohol –OH, ether –O–, aldehyde –CHO, ketone C=O (internal), carboxylic acid –COOH, ester –COO–, amine –NH₂, amide –CONH₂.

IUPAC naming: longest parent chain, number to give the lowest locants to the principal functional group, then to substituents alphabetically.

Isomerism: structural (chain, position, functional group) vs stereoisomerism (geometric cis/trans, optical enantiomers with a chiral carbon bearing four different groups).

Alkane substitution: free-radical halogenation (initiation, propagation, termination); alkenes undergo addition (H₂, HX, X₂, H₂O) at the C=C double bond.

Markovnikov's rule: in HX addition to an alkene, H adds to the carbon with more hydrogens; the halogen goes to the more substituted carbon.

Oxidation ladder: primary alcohol → aldehyde → carboxylic acid; secondary alcohol → ketone; tertiary alcohol resists oxidation.

Esterification: carboxylic acid + alcohol ⇌ ester + water (acid catalyst, reversible); hydrolysis reverses it; amines + acids form salts, and amides form from acid + amine.

Atomic number Z = protons (defines the element); mass number A = protons + neutrons; isotopes differ in neutrons, ions differ in electrons.

Atomic mass on the periodic table is the abundance-weighted average of isotope masses: (Σ fractional abundance × isotope mass).

Bohr model: E_n = −13.6 eV/n² for hydrogen; photon energy of a transition = |E_final − E_initial| = hc/λ.

Emission spectra are discrete lines from excited atoms relaxing; absorption spectra show dark lines at the same wavelengths.

Half-life t₁/₂: after n half-lives, remaining fraction = (1/2)^n; radioactive decay is first order.

Alpha decay emits ⁴₂He (A decreases by 4, Z by 2); beta-minus decay converts a neutron to a proton (Z +1, A unchanged); gamma emission changes neither A nor Z.

The periodic table is ordered by increasing Z, not atomic mass; Mendeleev's mass ordering was corrected by Moseley's atomic-number ordering.

Formal charge = valence electrons − (nonbonding electrons) − ½(bonding electrons); choose the Lewis structure with formal charges closest to zero.

Bond length decreases and bond energy increases with bond order: C–C > C=C > C≡C in strength, and C≡C < C=C < C–C in length.

Resonance structures differ only in electron placement; the real molecule is a hybrid with bond orders between the extremes (e.g., benzene C–C all 1.5).

Metallic bonding: delocalized valence electrons in a lattice explain conductivity, malleability, and luster; stronger with more delocalized electrons.

Network covalent solids (diamond, SiO₂, graphite) have very high melting points; molecular solids melt low because only intermolecular forces break.

Ionic compounds conduct only when molten or dissolved (mobile ions); solid ionic lattices are insulators.

Dipole moment μ = Q·r; a molecule with polar bonds can have μ = 0 if bond dipoles cancel by symmetry.

If the problem gives amounts of two reactants, compute the limiting reagent before anything else - the other is in excess.

For gas volumes in a reaction, the volume ratio equals the mole ratio at the same T and P (Avogadro's law).

When a reaction goes to completion and one product is a gas, the volume of gas collected = moles × molar volume at the stated conditions.

Percent purity = (mass of pure substance / mass of sample) × 100; use the pure mass in stoichiometry.

Combustion analysis: moles C = moles CO₂, moles H = 2 × moles H₂O; oxygen in the sample is found by mass difference.

Molarity after mixing: c_final = (n₁ + n₂)/(V₁ + V₂) when solutes are the same; for precipitation, use the limiting reagent's ion stoichiometry.

Titration relation: c_acid × V_acid × (moles base/moles acid) = c_base × V_base for a balanced neutralization.

Standard conditions: STP for gases is 0 °C and 1 bar (22.71 L/mol) or 0 °C and 1 atm (22.4 L/mol) - check which convention the problem uses.

Density of an ideal gas: ρ = PM/(RT); molar mass from gas density M = ρRT/P.

Graham's law: rate₁/rate₂ = √(M₂/M₁) for effusion of two gases at the same T.

Vapor pressure rises with temperature; a liquid boils when its vapor pressure equals the external pressure.

At the critical temperature, the liquid and gas phases become indistinguishable; above it, no amount of pressure liquefies the gas.

Heat of vaporization is much larger than heat of fusion for the same substance because vaporization fully separates molecules.

Colligative properties (vapor-pressure lowering, boiling-point elevation ΔTb = i·Kb·m, freezing-point depression ΔTf = i·Kf·m, osmosis) depend on particle count, not identity.

van't Hoff factor i equals the number of ions per formula unit for strong electrolytes (i ≈ 2 for NaCl, 3 for CaCl₂).

ICE table: Initial, Change (using the stoichiometric ratio x), Equilibrium; substitute into K and solve (small-x approximations allowed when K is very small).

For a weak acid, if [HA]/Ka > 100, the approximation [H⁺] ≈ √(Ka·c) is valid.

Degree of dissociation α = (moles dissociated)/(initial moles); for weak acids α ≈ √(Ka/c).

Le Chatelier with inert gas: adding inert gas at constant volume changes nothing; at constant pressure it dilutes the mixture and shifts toward more gas moles.

Equilibrium constant is unchanged by concentration, pressure, or catalyst changes; only temperature alters K.

For the reaction aA ⇌ bB, K = [B]^b/[A]^a; if the reaction is written twice, K is squared.

Solubility and Ksp: for a salt MₐX_b, s = (Ksp/(a^a·b^b))^(1/(a+b)) when no common ion is present.

Strong acid pH: [H⁺] = c_acid for monoprotic strong acids; for H₂SO₄, treat the first proton as complete and the second as weak.

Strong base pH: [OH⁻] = c_base × (number of OH⁻ per formula unit); then pH = 14 − pOH.

Weak acid: Ka = [H⁺][A⁻]/[HA]; if [H⁺] ≈ √(Ka·c), then pH = ½(pKa − log c).

Weak base: Kb = [BH⁺][OH⁻]/[B]; pOH ≈ ½(pKb − log c) when the approximation holds.

Buffer capacity is highest when [A⁻] = [HA]; adding strong acid converts A⁻ to HA, adding strong base converts HA to A⁻.

Titration curve of a weak acid with strong base: initial pH > 1, buffered region around pKa, steep jump at equivalence, basic pH at equivalence.

Amphiprotic salts (e.g., NaHCO₃) can act as acid or base; pH ≈ ½(pKa₁ + pKa₂).

Strong acid + weak base titration: equivalence pH < 7 because the conjugate acid hydrolyzes.

q_system + q_surroundings = 0 for an isolated system; heat lost by the reaction is gained by the calorimeter.

Bond breaking requires energy (ΔH > 0 for that step); bond formation releases energy (ΔH < 0 for that step).

ΔH°rxn can be estimated from average bond enthalpies: ΔH ≈ Σ(bonds broken) − Σ(bonds formed).

Phase-change enthalpies: q = n·ΔH_fus at melting point, q = n·ΔH_vap at boiling point; no temperature change during the phase change.

Standard state: 1 bar, 298 K, elements in their most stable form; ΔH°f of O₂(g), N₂(g), C(graphite) = 0.

Calorimetry sign: if the water temperature rises, the reaction is exothermic (ΔH < 0) and q_water is positive.

ΔG° = ΔH° − TΔS°; a reaction is spontaneous at all T if ΔH < 0 and ΔS > 0, and nonspontaneous at all T if ΔH > 0 and ΔS < 0.

ΔS°rxn = ΣS°(products) − ΣS°(reactants); S° values are positive and gases have much larger S° than liquids or solids.

Assign oxidation numbers before balancing; the total increase in oxidation number equals the total decrease.

In a galvanic cell, the more negative E° electrode is the anode (oxidation); the more positive is the cathode (reduction).

Standard hydrogen electrode (SHE) is the reference with E° = 0.00 V; all standard potentials are reduction potentials.

A species can disproportionate only if its potential lies between the potentials of its oxidized and reduced forms.

Electrolysis of water: cathode produces H₂ (reduction), anode produces O₂ (oxidation); the ratio is 2:1 by volume.

Electroplating: the object to be plated is the cathode; the plating metal is the anode; the electrolyte contains the plating metal's ions.

Corrosion (rusting) is an electrochemical process: iron is oxidized at an anodic site; oxygen and water are reduced at the cathodic site.

A salt bridge or porous membrane maintains electrical neutrality by allowing spectator ions to migrate.

Zero order: [A] vs t is linear with slope −k; first order: ln[A] vs t is linear with slope −k; second order: 1/[A] vs t is linear with slope +k.

For first-order reactions, concentration falls by a fixed fraction per half-life; the half-life is independent of starting concentration.

The rate constant k has units that depend on overall order: zero order mol·L⁻¹·s⁻¹, first order s⁻¹, second order L·mol⁻¹·s⁻¹.

Rate law from mechanism: use the slow step's reactants; if an intermediate appears, substitute it using the fast equilibrium expression.

A catalyst lowers Ea for both forward and reverse reactions by the same amount, so ΔH and K are unchanged.

Increasing reactant concentration raises rate only if that reactant appears in the rate law; zero-order reactants do not affect rate.

Enzyme-catalyzed reactions show saturation: rate becomes independent of substrate concentration at high [S] (Michaelis–Menten behavior).

Temperature coefficient rule of thumb: rate roughly doubles per 10 °C rise near room temperature, consistent with the Arrhenius equation.

Degree of unsaturation (index of hydrogen deficiency) = (2C + 2 + N − H − X)/2; each ring or π bond counts as one.

Alkenes: addition of Br₂ decolorizes bromine water; addition of H₂O (acid-catalyzed) gives Markovnikov alcohol.

Alcohols: oxidation with dichromate (acidified) turns orange Cr₂O₇²⁻ to green Cr³⁺; primary → aldehyde/acid, secondary → ketone.

Esterification is reversible and acid-catalyzed; hydrolysis of an ester under base (saponification) is irreversible and gives the carboxylate salt.

Amines are basic (lone pair on N); amides are much less basic because the lone pair is delocalized into the carbonyl.

Aldehydes are reducing (Tollens' silver mirror, Fehling's); ketones are not; this distinguishes the two carbonyl types.

Chirality requires a carbon with four different substituents; a racemic mixture has equal enantiomers and zero net optical rotation.

Cis/trans isomerism requires restricted rotation (C=C or a ring) and two different groups on each of the two carbons.

Atomic radius: decreases left→right across a period (higher Z_eff); increases top→bottom down a group (more shells).

Ionization energy: increases across a period, decreases down a group; anomalies at group 2→13 and group 15→16.

Electron affinity: becomes more negative (more favorable) across a period; noble gases are the exception.

Metallic character decreases across a period and increases down a group; oxides become more acidic across a period.

Cation radius < parent atom; anion radius > parent atom; isoelectronic series: larger nuclear charge → smaller radius (Na⁺ < Mg²⁺ < Al³⁺).

Fluorine has the highest electronegativity (3.98), then O (3.44), then N and Cl (~3.0); noble gases are not assigned values in the standard scale.

Group 1 hydroxides become more basic and more soluble down the group; group 2 sulfates become less soluble down the group.

Rank IMFs for similar molar mass: hydrogen bonding > dipole–dipole > London dispersion; ion–dipole is stronger than any of these.

London dispersion increases with the number of electrons and surface area; straight-chain isomers have higher boiling points than branched isomers.

Hydrogen bonding explains why HF, H₂O, and NH₃ boil far above their heavier homologues.

Solubility rule: like dissolves like - polar solutes dissolve in polar solvents; nonpolar solutes in nonpolar solvents.

Viscosity increases with stronger intermolecular forces and longer molecules; it decreases with temperature.

Surface tension arises from unbalanced forces at the liquid surface; it increases with stronger IMFs and decreases with temperature.

A molecule with a hydrogen-bond donor (O–H, N–H) and acceptor (lone pair) can hydrogen-bond; ethers are acceptors but not donors.

Particles to moles: n = N/6.022×10²³; moles to mass: m = n·M.

Mass of one molecule = molar mass / N_A; volume of one gas molecule at STP = molar volume / N_A.

Percent yield calculation order: balance equation → limiting reagent → theoretical yield → actual/theoretical × 100.

Mass of element in a compound = (atoms of element × atomic mass / molar mass) × sample mass.

For solutions, moles of solute = molarity × volume in liters; for gases at STP, moles = volume / 22.4 L (or 22.7 L).

Hydrated salts: molar mass includes water of crystallization; mass of anhydrous salt = total mass × (M_anhydrous/M_hydrate).

When mixing solutions that react, determine the limiting reagent from moles, not from concentrations.

At constant T, P₁V₁ = P₂V₂; at constant P, V₁/T₁ = V₂/T₂; at constant V, P₁/T₁ = P₂/T₂.

Ideal gas law can be rearranged for molar mass: M = mRT/(PV).

Partial pressure of a gas collected over water: P_gas = P_total − P_water vapor (look up vapor pressure at that T).

Mole fraction x_i = n_i/n_total; the sum of all mole fractions equals 1.

Real gas behavior is most ideal at low pressure and high temperature (far from condensation).

A heating curve's flat plateaus correspond to phase changes; the sloped segments correspond to heating within a single phase.

Critical pressure is the minimum pressure needed to liquefy a gas at its critical temperature.

Adding a solid or pure liquid does not shift equilibrium because its activity is constant.

Adding an inert gas at constant volume does not change partial pressures of reactants or products; no shift.

Increasing temperature favors the endothermic direction; if ΔH > 0, K increases with T.

Decreasing volume (increasing pressure) favors the side with fewer moles of gas; no shift if Δn = 0.

A catalyst does not change equilibrium concentrations; it only shortens the time to reach equilibrium.

Removing a product continuously (e.g., by precipitation or gas evolution) drives the reaction to the right.

For the same reaction written in reverse, K_reverse = 1/K_forward.

A buffer is made from a weak acid and its conjugate base (or weak base and conjugate acid) in comparable amounts.

Adding strong acid to a buffer: HA forms, A⁻ decreases; pH changes only slightly because the ratio [A⁻]/[HA] changes little.

Buffer pH after adding strong acid: use the new moles of HA and A⁻ in Henderson–Hasselbalch.

Salt of weak acid + strong base is basic: Kb = Kw/Ka of the conjugate acid; pH > 7.

Salt of weak base + strong acid is acidic: Ka = Kw/Kb of the conjugate base; pH < 7.

Salt of weak acid + weak base: pH depends on the relative Ka and Kb of the conjugate pair.

Amphoteric hydroxides (e.g., Al(OH)₃) dissolve in both strong acid and strong base.

Write the target reaction; identify given reactions that contain the target's reactants and products.

Reverse a given reaction to flip the sign of its ΔH; multiply a reaction to scale its ΔH by the same factor.

Add the manipulated reactions; species appearing on both sides cancel.

ΔH°rxn from formation enthalpies uses products minus reactants, each multiplied by its stoichiometric coefficient.

ΔH°rxn from combustion enthalpies uses reactants minus products (combustion values are negative).

Enthalpy of reaction from bond enthalpies is approximate because average bond values ignore molecular environment.

Calorimetry with a coffee-cup calorimeter assumes constant pressure, so q = ΔH; a bomb calorimeter measures q_v = ΔU.

Write separate half-reactions; balance atoms, then charge with electrons; multiply to equalize electrons.

In a galvanic cell, electrons flow from the anode to the cathode through the wire; conventional current flows the opposite way.

Anions migrate toward the anode and cations toward the cathode in the salt bridge or electrolyte.

E°cell > 0 means the reaction is spontaneous as written; E°cell < 0 means it is nonspontaneous.

ΔG° = −nFE°cell; a positive E°cell corresponds to a negative ΔG° (spontaneous).

Electrolysis stoichiometry: charge Q = It (coulombs); moles of electrons = Q/F; moles of metal = moles e⁻/n.

Overpotential and competing oxidations (e.g., water oxidation) can change the expected products in aqueous electrolysis.

Elementary step rate law: rate = k[reactants]^coefficients for that step.

If the slow step contains an intermediate, solve the fast equilibrium for the intermediate concentration and substitute.

The molecularity of an elementary step (unimolecular, bimolecular, termolecular) equals its number of reactant species.

A large Ea means strong temperature dependence; a small Ea means rate changes little with temperature.

Catalysts can be homogeneous (same phase) or heterogeneous (different phase, often solid surface).

Inhibitors slow reactions by raising Ea or blocking active sites; they do not change equilibrium.

Reaction rate decreases as reactants are consumed because collisions become less frequent.

Parent chain choice: longest chain containing the principal functional group; if tied, choose the one with more substituents.

Numbering priority: principal functional group > multiple bonds > substituents; lowest locant set wins.

Prefixes: meth-, eth-, prop-, but-, pent-, hex- for 1–6 carbons; substituents use -yl (methyl, ethyl, chloro, bromo).

Suffixes: -ane (alkane), -ene (alkene), -yne (alkyne), -ol (alcohol), -al (aldehyde), -one (ketone), -oic acid (carboxylic acid).

Isomer count for small alkanes: C₄H₁₀ has 2, C₅H₁₂ has 3, C₆H₁₄ has 5 structural isomers.

Geometric isomers require a C=C with two different groups on each carbon; identical groups on one carbon rule out cis/trans.

Enantiomers rotate plane-polarized light in opposite directions; diastereomers have different physical properties.