ORT Chemistry cheat sheet

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

Quantum numbers: n (1,2,3…), l = 0…n−1, m_l = −l…+l, m_s = ±1/2; no two electrons in an atom share all four.

Orbital filling order follows the Aufbau principle: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p.

Hund's rule: degenerate orbitals are singly occupied with parallel spins before pairing; explains paramagnetism of unpaired electrons.

Across a period, effective nuclear charge rises, atomic radius falls, ionization energy and electronegativity rise.

Down a group, radius and metallic character increase; ionization energy and electronegativity decrease.

Anomalies: Cr is [Ar]3d5 4s1 and Cu is [Ar]3d10 4s1 because half-filled and filled d subshells are more stable.

Group valence-electron counts: group 1 → ns1, group 2 → ns2, groups 13–18 → ns2 np1–6; d-block fills (n−1)d.

Ionic bonding: large ΔEN (>~1.7) with electron transfer; lattice energy rises with higher charges and smaller ion radii (Born–Landé).

Covalent bond order = (bonding electrons − antibonding electrons)/2; higher bond order means shorter, stronger bond.

VSEPR: electron pairs repel; 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, 180°), 3 → sp2 (trigonal planar, 120°), 4 → sp3 (tetrahedral, 109.5°).

Polarity requires polar bonds plus an asymmetric arrangement; CO2 is nonpolar, H2O and NH3 are polar.

Molecular orbital theory: O2 has two unpaired electrons in π* orbitals, explaining its paramagnetism; N2 has bond order 3.

Hydrogen bonding (N–H, O–H, F–H with lone pair) raises boiling point and solubility in water; London dispersion grows with molar mass.

Formal charge = valence electrons − nonbonding electrons − ½ bonding electrons; prefer structures with charges minimized.

s-Block: group 1 metals react with water giving MOH + H2; reactivity increases down the group; group 2 hydroxides become more soluble down the group.

Alkali metal peroxides (Na2O2) and superoxides (KO2) form with larger cations; Li forms mainly the normal oxide Li2O.

p-Block trends: nonmetallic character increases across a period; halogens are oxidizing agents, F2 strongest, I2 weakest.

Halide reducing power increases down the group (I− > Br− > Cl− > F−); HF is a weak acid, HCl–HI are strong.

Nitrogen: NH3 is a weak base, HNO3 a strong oxidizing acid; NO2 is a brown gas, N2O4 colorless (dimerization equilibrium).

d-Block: variable oxidation states (Mn: +2 to +7); Cr2O7^2− (orange) and CrO4^2− (yellow) interconvert with pH.

KMnO4 in acid is reduced to Mn2+ (colorless/faint pink); in neutral or alkaline media to MnO2 (brown).

Transition-metal ions are colored and often paramagnetic due to d–d transitions and unpaired d electrons; Zn2+ is colorless and diamagnetic.

Mole n = m/M = N/N_A = V(gas)/22.4 L at STP; molar volume 22.4 L/mol at 0 °C and 101.3 kPa.

Balance equations by atom conservation; use mole ratios from coefficients for limiting-reagent problems.

Limiting reagent: divide moles of each reactant by its coefficient; the smallest quotient is limiting.

Percent yield = (actual/theoretical) × 100%; percent composition = (mass of element/molar mass) × 100%.

Ideal gas law PV = nRT; at constant T and P, volumes react in the same ratio as moles (Gay-Lussac).

Enthalpy: ΔH = Σ bond energies broken − Σ bond energies formed; Hess's law allows summing stepwise ΔH values.

Reaction rate rises with concentration, temperature, surface area, and catalysts; a catalyst lowers activation energy without changing ΔH.

Le Chatelier: adding reactant or removing product shifts forward; raising pressure shifts toward fewer gas moles; catalysts change neither K nor position.

Molarity c = n/V(L); dilution M1V1 = M2V2; molality uses kg solvent and is temperature-independent.

Strong electrolytes dissociate fully (HCl, NaOH, NaCl); weak electrolytes partially (CH3COOH, NH3·H2O) with Ka or Kb.

Water autoionization Kw = [H+][OH−] = 1.0×10−14 at 25 °C; pH + pOH = 14.

pH = −log[H+] for strong acids; for weak acids [H+] ≈ √(Ka·c); buffer pH = pKa + log([base]/[acid]) (Henderson–Hasselbalch).

Solubility product Ksp = [cation]^m[anion]^n; precipitation occurs when the ion product Q exceeds Ksp.

Common-ion effect suppresses dissociation of a weak electrolyte or lowers salt solubility.

Titration equivalence: strong acid–strong base pH 7; weak acid–strong base equivalence pH > 7; choose indicator whose range brackets the equivalence pH.

Colligative properties (ΔTb = Kb·m, ΔTf = Kf·m, π = cRT) depend on particle count; electrolytes use the van't Hoff factor i.

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

Oxidation number rules: free element 0; O usually −2 (peroxides −1); H usually +1 (hydrides −1); sum equals charge.

Balance redox in acid with H2O and H+; in base add OH− to neutralize H+ after balancing.

Standard cell potential E°cell = E°cathode − E°anode; a positive E°cell means a spontaneous reaction (ΔG° = −nFE°cell).

In a galvanic cell, oxidation occurs at the anode (negative) and reduction at the cathode (positive); electrons flow anode → cathode.

In electrolysis, the anode is positive and oxidation occurs there; the cathode is negative and reduction occurs there.

Nernst equation: E = E° − (0.0592/n)·log Q at 25 °C; changing concentration shifts potential.

Electrolysis product priority: at the cathode the more easily reduced species (Cu2+ before H2O); at the anode the more easily oxidized (Cl− before H2O, except with inert vs active electrodes).

Alkanes C_nH_2n+2: substitution with halogens under UV light via free radicals; no addition reactions.

Alkenes C_nH_2n and alkynes C_nH_2n−2: electrophilic addition of H2, X2, HX, H2O; Markovnikov places H on the carbon with more hydrogens.

Benzene: aromatic, planar, delocalized π system; undergoes electrophilic substitution (nitration, halogenation with FeX3, Friedel–Crafts) rather than addition.

Alcohols: oxidation of primary gives aldehyde then acid; secondary gives ketone; tertiary resists oxidation.

Aldehydes reduce Tollens' reagent (silver mirror) and Fehling's solution; ketones do not. Carboxylic acids + alcohols ⇌ esters (acid catalyst).

Phenols are weakly acidic (resonance-stabilized phenoxide) and give violet color with FeCl3; amines are basic, aliphatic stronger than aromatic.

Esterification and hydrolysis are reversible; saponification of fats with NaOH gives glycerol and soap.

Isomerism: chain, position, functional group, cis–trans, and optical (chiral carbon with four different groups).

Free-radical substitution: initiation (homolysis by UV/heat), propagation, termination; order of halogen reactivity Cl2 > Br2.

SN1: two steps via carbocation, rate depends only on substrate; favored by tertiary carbon and polar protic solvent; racemization.

SN2: one step, backside attack, rate depends on substrate and nucleophile; favored by methyl/primary carbon and polar aprotic solvent; inversion.

E1/E2 elimination competes with substitution; strong bulky base and heat favor E2; Zaitsev's rule favors the more substituted alkene.

Electrophilic aromatic substitution: activating groups (–OH, –NH2, –R) are ortho/para directors; deactivating groups (–NO2, –COOH, –SO3H) are meta directors.

Nucleophilic addition to carbonyls: aldehydes more reactive than ketones; Grignard reagents give alcohols; HCN gives cyanohydrins.

Retrosynthesis: disconnect the target at a bond formed by a known reaction (ester → acid + alcohol; ether → alkoxide + alkyl halide via Williamson).

Oxidation/reduction reagents: KMnO4 or K2Cr2O7 oxidize alcohols; LiAlH4 reduces acids/esters to alcohols; NaBH4 reduces aldehydes/ketones only.

Addition polymerization (polyethylene, PVC, polystyrene) opens double bonds; condensation polymerization (nylon, PET) releases small molecules such as H2O.

Degree of polymerization n = molar mass of polymer / molar mass of repeat unit; crystallinity raises density and strength.

Thermoplastics soften on heating and are recyclable; thermosets cross-link irreversibly (bakelite, epoxy).

Proteins are polyamides of α-amino acids linked by peptide bonds; primary, secondary (α-helix, β-sheet), tertiary, quaternary structure.

Carbohydrates: monosaccharides (glucose, fructose) are reducing sugars; sucrose is nonreducing; starch gives blue with iodine.

Lipids: fats are triacylglycerols; saturated fats are solid, unsaturated liquid; saponification yields soap and glycerol.

Nucleic acids: DNA deoxyribose double helix with A–T, G–C pairs; RNA ribose single strand with U replacing T.

Enzymes are protein catalysts with active sites; activity has an optimum pH and temperature and can be denatured by heat or heavy metals.

Never pipette by mouth; use a pipette bulb. Add acid to water, not water to acid, to control heat.

Flame tests: Li crimson, Na yellow, K lilac (view through cobalt glass), Ca brick-red, Ba yellow-green, Cu blue-green.

Cation group separation: group I precipitated as chlorides (Ag+, Pb2+, Hg2^2+), group II as sulfides in acid, group III as sulfides/hydroxides in base.

Anion tests: CO3^2− effervesces with acid (CO2 turns limewater milky); SO4^2− gives white BaSO4 insoluble in acid; Cl− gives white AgCl soluble in ammonia.

NH4+ detected by warming with NaOH, releasing NH3 that turns moist red litmus blue.

Heating and weighing: cool a crucible in a desiccator before weighing; use a watch glass to cover beakers.

Titration technique: rinse burette with titrant, read at eye level at the bottom of the meniscus, add dropwise near the endpoint.

Waste disposal: halogenated solvents, heavy-metal solutions and concentrated acids go into labeled waste containers, never the sink.