Every Chemistry concept in the high school Science Bowl study set (grades 9-12), with a short definition for each. Read through them here, then drill them as flashcards or practise real Chemistry questions.
Protons (+) and neutrons (neutral) in nucleus; electrons (−) in orbitals. Atomic number = protons; mass number = protons + neutrons. Isotopes: same element, different neutron count. Ions: gained electrons = anion (−); lost electrons = cation (+). Electron configuration follows Aufbau principle.
Atomic radius decreases across a period (more protons pull electrons in) and increases down a group (more electron shells). Electronegativity and ionization energy increase across a period, decrease down a group. Electron affinity follows similar trends. Noble gases have the highest ionization energies per period.
Ionic: metal + nonmetal, electron transfer, forms crystal lattice. Covalent: nonmetal + nonmetal, electron sharing. Polar covalent: unequal sharing. Metallic: delocalized electron sea. Bond strength order (single < double < triple). Octet rule: atoms gain/lose/share electrons to achieve 8 valence electrons.
Boyle's Law: P₁V₁ = P₂V₂ (constant T). Charles' Law: V₁/T₁ = V₂/T₂ (constant P). Gay-Lussac's: P₁/T₁ = P₂/T₂ (constant V). Ideal Gas Law: PV = nRT (R = 8.314 J/mol·K). Avogadro's: at same T and P, equal volumes contain equal numbers of moles.
Arrhenius: acid produces H⁺, base produces OH⁻. Brønsted-Lowry: acid donates proton, base accepts proton. Lewis: acid accepts electron pair. pH = −log[H⁺]; strong acids fully dissociate (HCl, HNO₃, H₂SO₄, HBr, HI, HClO₄). Conjugate acid-base pairs differ by one H⁺.
Enthalpy (ΔH): heat at constant pressure; exothermic ΔH < 0, endothermic ΔH > 0. Entropy (ΔS): disorder. Gibbs free energy: ΔG = ΔH − TΔS; spontaneous if ΔG < 0. Hess's Law: enthalpy is path-independent (state function). Standard enthalpy of formation: ΔHf° for 1 mol from elements in standard state.
Rate depends on concentration, temperature, surface area, and catalysts. Activation energy (Eₐ): minimum energy for reaction. Arrhenius equation: k = Ae^(−Eₐ/RT). Rate law: rate = k[A]^m[B]^n (determined experimentally). Half-life (t½): time for half of reactant to be consumed; constant for first-order reactions.
OIL RIG: Oxidation Is Loss, Reduction Is Gain of electrons. Galvanic (voltaic) cell: spontaneous redox; anode (oxidation, −) and cathode (reduction, +). Electrolytic cell: non-spontaneous, driven by external current. EMF (cell potential) measured in volts. Nernst equation relates EMF to concentration.
Dynamic state where forward = reverse reaction rate. Le Chatelier's Principle: system shifts to oppose any stress. Kc = [products]/[reactants] (excludes solids and pure liquids). Kp relates to partial pressures of gases. Q vs K: Q < K → shifts right; Q > K → shifts left.
Molarity (M) = moles solute / liters solution. Molality (m) = moles / kg solvent. Dilution: M₁V₁ = M₂V₂. Solubility rules: all nitrates soluble; most chlorides soluble (except Ag⁺, Pb²⁺, Hg₂²⁺). Like dissolves like: polar solvents dissolve polar/ionic solutes; nonpolar solvents dissolve nonpolar solutes.
Mole concept: 1 mol = 6.022×10²³ particles (Avogadro's number). Molar mass: sum of atomic masses in g/mol. Balanced equations give mole ratios. Limiting reagent: reactant completely consumed first; determines maximum product yield. Percent yield = (actual yield / theoretical yield) × 100%.
Radioactive decay types: alpha (α, ²⁴He nucleus emitted, penetration lowest), beta-minus (β⁻, neutron → proton + electron), beta-plus (positron emission), gamma (γ, high-energy photon, no mass/charge change). Half-life: constant for each isotope. Fission splits heavy nuclei; fusion combines light nuclei.
Hydrocarbons: alkanes (single bonds, CₙH₂ₙ₊₂), alkenes (double bond, CₙH₂ₙ), alkynes (triple bond, CₙH₂ₙ₋₂), benzene (aromatic). Functional groups: hydroxyl (−OH, alcohols), carbonyl (C=O), carboxyl (−COOH, acids), amino (−NH₂), ester (−COO−). IUPAC naming by longest carbon chain.
London dispersion forces (van der Waals): weakest; in all molecules; increase with molecular mass and surface area. Dipole-dipole: between polar molecules. Hydrogen bonds: strongest intermolecular; N–H, O–H, or F–H with electronegative atom on adjacent molecule. IMF strength determines boiling point and vapor pressure.
VSEPR theory: electron pairs repel to maximize separation. Shapes by electron pairs: 2 = linear, 3 = trigonal planar, 4 = tetrahedral, 5 = trigonal bipyramidal, 6 = octahedral. Lone pairs compress bond angles. Water (2 lone pairs): bent (~104.5°). NH₃ (1 lone pair): trigonal pyramidal (~107°).
Four quantum numbers describe each electron: n (principal, energy level, 1,2,3…), l (angular momentum, subshell: 0=s, 1=p, 2=d, 3=f), mₗ (magnetic, orbital orientation), mₛ (spin, +½ or −½). Pauli exclusion: no two electrons can share all four quantum numbers. Aufbau and Hund's rules fill orbitals.
Synthesis (combination): A + B → AB. Decomposition: AB → A + B. Single displacement: A + BC → AC + B. Double displacement (metathesis): AB + CD → AD + CB. Combustion: fuel + O₂ → CO₂ + H₂O. Acid-base neutralization: acid + base → salt + water.
Properties that depend only on the number of solute particles, not their identity. Adding a non-volatile solute: raises boiling point (ΔTb = Kb·m·i), lowers freezing point (ΔTf = Kf·m·i), lowers vapor pressure (Raoult's Law), and increases osmotic pressure (π = MRT·i). i = van 't Hoff factor.
Graph of pressure vs temperature showing solid, liquid, and gas regions. Triple point: all three phases coexist. Critical point: above this T and P, liquid and gas are indistinguishable (supercritical fluid). Sublimation: solid → gas directly (CO₂ at 1 atm). Normal boiling/melting points are at 1 atm.
Resist changes in pH when small amounts of acid or base are added. Composed of a weak acid and its conjugate base (or weak base and its conjugate acid). Henderson-Hasselbalch equation: pH = pKa + log([A⁻]/[HA]). Most effective buffering at pH = pKa ± 1. Blood is buffered by H₂CO₃/HCO₃⁻ at pH ≈ 7.4.
Assign oxidation numbers: free elements = 0; monatomic ions = charge; O usually −2 (except peroxides: −1); H usually +1 (except metal hydrides: −1); sum = charge of species. In redox reactions, the oxidizing agent is reduced (gains electrons); the reducing agent is oxidized (loses electrons).
Metal ion (Lewis acid) surrounded by ligands (Lewis bases donating electron pairs). Coordination number: number of ligand bonds to metal. Ligand types: monodentate (one bond, e.g., NH₃, Cl⁻), bidentate (two bonds, e.g., en), polydentate (chelate). Complex ion charge = metal oxidation state + ligand charges.
Structural isomers: same molecular formula, different connectivity (e.g., butane vs isobutane). Geometric (cis-trans) isomers: same connectivity, different spatial arrangement around a double bond or ring. Optical isomers (enantiomers): non-superimposable mirror images; differ at chiral carbon; rotate plane-polarized light oppositely.
Emission spectra: atoms emit characteristic photons when electrons fall to lower energy levels (Bohr model). Absorption spectra: atoms absorb specific wavelengths, leaving dark lines. Beer-Lambert Law: A = εlc (absorbance = molar absorptivity × path length × concentration). NMR, IR, and mass spectrometry identify organic compounds.
Polymers: large molecules made of repeating monomer units. Addition polymers (e.g., polyethylene): monomers with double bonds add without losing atoms. Condensation polymers (e.g., nylon, polyester): monomers join by losing small molecules (H₂O). Proteins are natural polymers (amino acid monomers joined by peptide bonds).
Elementary steps that show how reactants become products. Rate-determining step: slowest step; controls overall reaction rate. Intermediates: formed in one step, consumed in a later step. Transition state (activated complex): highest-energy point in an elementary step. SN1 and SN2 are common substitution mechanisms in organic chemistry.
Titration: technique to determine unknown concentration by reacting with a standard solution. Equivalence point: moles of acid = moles of base. Endpoint: indicator changes color. Common indicators: phenolphthalein (colorless → pink at pH ~8.2–10), methyl orange (red → yellow at pH ~3.1–4.4). Back titration: excess reagent is titrated to find original amount. Neutralization: strong acid + strong base → salt + water; pH at equivalence = 7.
Equilibrium constant for dissolution of sparingly soluble ionic compound. For AgCl: AgCl(s) ⇌ Ag⁺ + Cl⁻; Ksp = [Ag⁺][Cl⁻]. Common ion effect: adding a shared ion reduces solubility (Le Chatelier). Precipitation occurs when Q > Ksp. Selective precipitation separates ions by exploiting different Ksp values. Molar solubility s relates to Ksp (e.g., for AB type: Ksp = s²).
Electrons fill orbitals by Aufbau principle (lowest energy first), Hund's rule (one electron per orbital before pairing), and Pauli exclusion (no two electrons with same four quantum numbers). Orbital filling order: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d... Exceptions: Cr ([Ar]3d⁵4s¹) and Cu ([Ar]3d¹⁰4s¹) for half-filled/full-d stability. Valence electrons: outermost shell electrons; determine chemical behavior.
Dot diagrams showing valence electrons. Steps: count total valence electrons, place least electronegative atom in center, form single bonds, complete octets of outer atoms, place remaining electrons on central atom, convert lone pairs to double/triple bonds if needed. Formal charge = (valence e⁻) − (lone pair e⁻) − ½(bonding e⁻). Resonance structures: multiple valid Lewis structures; actual molecule is a hybrid (e.g., benzene, O₃, NO₃⁻).
Weak acid equilibrium: HA ⇌ H⁺ + A⁻; Ka = [H⁺][A⁻]/[HA]. Weak base: B + H₂O ⇌ BH⁺ + OH⁻; Kb. Relationship: Ka × Kb = Kw = 10⁻¹⁴. pKa = −log(Ka); lower pKa = stronger acid. For weak acid solutions: [H⁺] ≈ √(Ka × C). Buffer: pH = pKa + log([A⁻]/[HA]) (Henderson-Hasselbalch). Amphoteric species (like water, amino acids) can act as both acid and base.
Mixing of atomic orbitals to form new hybrid orbitals with equivalent energy. sp: linear geometry (180°); two sp orbitals + two unhybridized p (e.g., BeCl₂, alkynes). sp²: trigonal planar (120°); three sp² + one p (e.g., BF₃, alkenes, benzene). sp³: tetrahedral (109.5°); four sp³ (e.g., CH₄, NH₃, H₂O). sp³d: trigonal bipyramidal. sp³d²: octahedral. Lone pairs occupy hybrid orbitals and compress bond angles.
Half-reaction method: split into oxidation and reduction half-reactions; balance atoms (add H₂O for O, H⁺ for H in acidic; OH⁻ and H₂O in basic); balance charge by adding electrons; multiply half-reactions so electrons cancel; add half-reactions together. Oxidation number method: identify changes, multiply to equalize electron gain/loss, balance remaining atoms. Always confirm atoms and charge balance.
ΔG = ΔH − TΔS (Gibbs free energy). Spontaneous if ΔG < 0. Four cases: ΔH < 0 and ΔS > 0 → always spontaneous; ΔH > 0 and ΔS < 0 → never spontaneous; ΔH < 0 and ΔS < 0 → spontaneous at low T; ΔH > 0 and ΔS > 0 → spontaneous at high T. Standard ΔG° = −RT ln K; relates equilibrium constant to thermodynamics. ΔG° < 0 means K > 1 (products favored).
Colloids: mixtures with particles 1–1000 nm dispersed in a medium (between solution and suspension). Types: aerosol (liquid/solid in gas — fog, smoke), foam (gas in liquid), emulsion (liquid in liquid — milk, mayonnaise), sol (solid in liquid), gel (liquid in solid). Tyndall effect: colloids scatter light (distinguishes from true solutions). Coagulation: colloid particles aggregate and settle. Surfactants stabilize emulsions.
Binary ionic: cation name first, then anion with -ide suffix (NaCl = sodium chloride). Transition metals: use Roman numerals for charge (FeCl₃ = iron(III) chloride). Polyatomic ions: memorize common ones (SO₄²⁻ = sulfate, NO₃⁻ = nitrate, PO₄³⁻ = phosphate, CO₃²⁻ = carbonate, NH₄⁺ = ammonium). Binary covalent: use Greek prefixes (mono-, di-, tri-…); less electronegative element named first (N₂O₄ = dinitrogen tetroxide). Acids: HCl = hydrochloric; H₂SO₄ = sulfuric.
Twenty standard amino acids; structure: central α-carbon bonded to amino group (−NH₂), carboxyl group (−COOH), H, and variable R-group (side chain). Classified by R-group: nonpolar, polar uncharged, acidic (Asp, Glu), basic (Lys, Arg, His). Peptide bond formation: condensation reaction between −COOH and −NH₂; water is released. Isoelectric point (pI): pH at which net charge = 0. Essential amino acids must be obtained from diet.
Reaction rate v = Vmax[S] / (Km + [S]). Km (Michaelis constant): substrate concentration at half-maximal velocity; low Km = high enzyme affinity. Vmax: maximum rate when all enzyme active sites are saturated. Lineweaver-Burk (double reciprocal) plot linearizes the data. Competitive inhibition: raises apparent Km, same Vmax (overcome by high [S]). Noncompetitive inhibition: lowers Vmax, same Km (cannot be overcome).
Binding energy = mass defect × c². Iron-56 has the highest binding energy per nucleon — fusing lighter nuclei or fissioning heavier nuclei toward iron releases energy. Radioactive decay law: N = N₀e^(−λt); λ = ln2/t½. Activity A = λN (measured in Becquerels or Curies). Mass spectrometry separates ions by mass-to-charge ratio. Radiocarbon dating (C-14, t½ = 5730 yr) used for organic material up to ~50,000 years.
Technique separating mixtures based on differential migration through a stationary phase with a mobile phase. Paper chromatography: Rf = distance traveled by compound / distance traveled by solvent; separates pigments by polarity. Thin-layer chromatography (TLC): similar, uses silica on glass/plastic. Gas chromatography (GC): volatile compounds separated by boiling point and polarity. HPLC: high-pressure liquid chromatography for complex mixtures. Column chromatography: gravity-fed separation.