Every Chemistry concept in the middle school Science Bowl study set (grades 6-8), 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.
Atoms are the smallest units of an element that retain its chemical properties. Atoms consist of a nucleus (protons and neutrons) surrounded by electrons. Atomic number = number of protons (identifies the element). Mass number = protons + neutrons. An element is a pure substance made of only one type of atom. There are 118 known elements on the periodic table.
Solid: definite shape and volume; particles vibrate in place. Liquid: definite volume, no definite shape; particles slide past each other. Gas: no definite shape or volume; particles move freely and rapidly. Plasma: high-energy ionized gas (stars, lightning). Changes of state: melting (solid→liquid), freezing, evaporation, condensation, sublimation (solid→gas), deposition.
Physical change: alters the form or appearance of matter but not its chemical composition (cutting, melting, dissolving). Chemical change: produces new substances with different properties (combustion, rusting, cooking). Evidence of chemical change: color change, gas production, precipitate formation, temperature change, light emission. Chemical changes are generally harder to reverse.
Element: pure substance made of one type of atom (oxygen, gold). Compound: pure substance formed when two or more elements chemically combine in fixed ratios (H₂O, NaCl); properties differ from its elements. Mixture: two or more substances combined physically; can be separated (salt water, air). Homogeneous mixture (solution): uniform composition; heterogeneous: non-uniform.
Acids taste sour, turn blue litmus red, and have pH below 7 (examples: vinegar, lemon juice, stomach acid). Bases taste bitter, feel slippery, turn red litmus blue, and have pH above 7 (examples: baking soda, soap, ammonia). Neutral substances have pH = 7 (pure water). Neutralization: acid + base → salt + water.
Organizes 118 elements by atomic number. Periods (rows): elements across a row have the same number of electron shells. Groups/families (columns): elements share similar properties. Metals (left side): shiny, conduct heat/electricity. Nonmetals (right side): poor conductors. Metalloids (along the staircase): intermediate properties. Noble gases (Group 18): very unreactive.
Chemical reactions rearrange atoms to form new substances. Reactants (left side) → Products (right side). Law of Conservation of Mass: mass is conserved; atoms are neither created nor destroyed, so equations must be balanced. Coefficients show the ratio of molecules. Types: synthesis, decomposition, single replacement, double replacement, combustion.
Density = mass / volume (D = m/V). Units: g/cm³ or g/mL. Water density = 1 g/cm³. Objects denser than the liquid they are placed in sink; less dense objects float. Density is an intensive property — it does not change with sample size. Temperature affects density: most substances become less dense as they are heated (water is an exception near 4°C).
Atoms bond to achieve stable electron configurations (usually 8 valence electrons — the octet rule). Ionic bonds: metal transfers electrons to nonmetal; forms ions (charged atoms). The opposite charges attract. Covalent bonds: two nonmetals share electrons. Metallic bonds: metal atoms share a sea of electrons. Bond type determines the properties of the compound formed.
Dalton (1803): atoms are indivisible solid spheres; different elements have different atoms. Thomson (1897): discovered electron; 'plum pudding' model — electrons embedded in positive sphere. Rutherford (1911): gold foil experiment — most of atom is empty space; dense positive nucleus. Bohr (1913): electrons orbit nucleus in fixed energy levels; electrons absorb/emit photons when changing levels. Modern quantum mechanical model: electrons exist in probability clouds (orbitals), not fixed paths.
Ions are atoms or groups of atoms with a net electrical charge. Cations: positive ions formed when an atom loses electrons (metals; e.g., Na → Na⁺). Anions: negative ions formed when an atom gains electrons (nonmetals; e.g., Cl → Cl⁻). Ionic compounds form when cations and anions attract each other. Common polyatomic ions: OH⁻ (hydroxide), SO₄²⁻ (sulfate), NO₃⁻ (nitrate), NH₄⁺ (ammonium), CO₃²⁻ (carbonate). Ions in solution conduct electricity (electrolytes).
Chemical formula shows types and numbers of atoms in a compound (H₂O: 2 hydrogen + 1 oxygen). Subscripts show number of atoms; coefficients show number of molecules. Balancing equations: must have equal numbers of each atom on both sides (Law of Conservation of Mass). Steps: write unbalanced equation, count atoms, add coefficients to balance (never change subscripts). Diatomic elements: H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂ — always appear as pairs in their elemental form.
Solution: homogeneous mixture of solute (substance being dissolved) and solvent (substance doing the dissolving). Solubility: maximum amount of solute that dissolves in a given amount of solvent at a specific temperature. 'Like dissolves like': polar solvents (water) dissolve polar and ionic solutes; nonpolar solvents dissolve nonpolar solutes. Temperature: increases solubility of most solids; decreases solubility of gases. Saturated: holds maximum solute; unsaturated: can dissolve more; supersaturated: holds more than maximum (unstable).
Oxidation: loss of electrons (OIL — Oxidation Is Loss). Reduction: gain of electrons (RIG — Reduction Is Gain). Redox reactions involve simultaneous oxidation and reduction. Oxidizing agent: gains electrons and is reduced. Reducing agent: loses electrons and is oxidized. Examples: rusting (iron oxidized by oxygen), combustion (fuel oxidized), photosynthesis/respiration (electron transfers in cells). Electrochemical cells use redox reactions to produce electricity (batteries).
Synthesis: two or more substances combine to form one product (A + B → AB; e.g., 2H₂ + O₂ → 2H₂O). Decomposition: one compound breaks down into simpler substances (AB → A + B; e.g., H₂O₂ → H₂O + O₂). Single replacement: one element replaces another in a compound (A + BC → AC + B). Double replacement: ions exchange between two compounds in solution. Combustion: substance reacts with oxygen, releasing heat and light; produces CO₂ and H₂O (complete combustion).
Filtration: separates insoluble solid from liquid using filter paper (e.g., sand from water). Evaporation: removes liquid solvent, leaves solid solute (e.g., salt from saltwater). Distillation: separates liquids by boiling point differences; liquid boils, vapor rises, condenses in cooler tube. Chromatography: separates mixtures based on how far components travel through a medium with a solvent (different Rf values). Magnetism: separates magnetic materials (iron filings from sand). Centrifugation: dense particles settle faster under centrifugal force.
Metals (left and center of periodic table): shiny luster, good conductors of heat and electricity, malleable (can be hammered into sheets), ductile (drawn into wire), high melting points (except mercury), tend to lose electrons and form cations. Nonmetals (upper right): dull, poor conductors (except graphite), brittle in solid form, lower melting points, tend to gain electrons and form anions. Metalloids (boron, silicon, germanium, arsenic, antimony, tellurium): properties intermediate between metals and nonmetals; used as semiconductors.
Exothermic reactions: release energy (usually as heat) to surroundings; products have lower energy than reactants; ΔH < 0 (examples: combustion, rusting, hand warmers). Endothermic reactions: absorb energy from surroundings; products have higher energy than reactants; ΔH > 0 (examples: photosynthesis, cold packs, cooking an egg). Activation energy: minimum energy needed to start a reaction; catalysts lower activation energy. Bond breaking requires energy; bond forming releases energy.
Carbon is the basis of all organic (living) chemistry. Carbon has 4 valence electrons → forms 4 covalent bonds → can bond with itself to form long chains, rings, and branched structures. Hydrocarbons: compounds containing only C and H. Methane (CH₄, natural gas), ethane (C₂H₆), propane (C₃H₈), butane (C₄H₁₀). Carbon dioxide (CO₂) and carbonate ions (CO₃²⁻) are inorganic carbon compounds. Organic compounds include fats, sugars, proteins, and DNA.
Water (H₂O): polar molecule (slight negative charge near O, slight positive near H). Hydrogen bonds between water molecules give it unique properties: high specific heat capacity (resists temperature change), high surface tension (insects can walk on water), cohesion (water sticks to water) and adhesion (water sticks to other surfaces). Water expands when it freezes (ice less dense than liquid → ice floats → aquatic life can survive under ice). Universal solvent for polar and ionic compounds.
Rate of reaction: how quickly reactants are converted to products. Factors that increase reaction rate: higher temperature (particles have more kinetic energy, collide more often and with more energy), higher concentration (more particles per volume → more collisions), greater surface area (smaller particles expose more surface for collisions), addition of a catalyst (lowers activation energy without being consumed). Collision theory: for reaction to occur, particles must collide with sufficient energy and correct orientation.
Some atomic nuclei are unstable and emit particles or energy (radioactive decay). Alpha particles (α): 2 protons + 2 neutrons (helium nucleus); stopped by paper; least penetrating. Beta particles (β): fast-moving electron; stopped by thin metal. Gamma rays (γ): high-energy electromagnetic radiation; most penetrating; requires lead or thick concrete shielding. Half-life: time for half of radioactive atoms to decay. Carbon-14 dating: used to date organic materials (t½ = 5,730 years). Uses: medicine (PET scan), power generation, food irradiation.
Atoms of the same element with the same number of protons but different numbers of neutrons. They have the same atomic number but different mass numbers. Examples: carbon-12 (6p, 6n) and carbon-14 (6p, 8n); hydrogen (1p, 0n), deuterium (1p, 1n), tritium (1p, 2n). Most elements exist as a mixture of isotopes; atomic mass on the periodic table is a weighted average. Radioactive isotopes (radioisotopes) have unstable nuclei. Stable isotopes are used in medical imaging and archaeological dating.
Gas particles move rapidly and randomly in all directions; fill any container; spread to occupy available volume (diffusion). Gas pressure: force per unit area exerted by gas particles hitting container walls; measured in Pascals, atm, or mmHg (1 atm = 101,325 Pa = 760 mmHg). Boyle's Law: pressure and volume are inversely proportional at constant temperature (P₁V₁ = P₂V₂). Charles' Law: volume and temperature (Kelvin) are directly proportional at constant pressure (V₁/T₁ = V₂/T₂). STP: 0°C and 1 atm.
Electrolytes: substances that dissolve in water to produce ions that conduct electricity (salts, acids, bases). Electrolysis: using electrical current to drive a non-spontaneous chemical reaction. Electroplating: using electrolysis to coat an object with a thin layer of metal (e.g., chrome plating). Batteries (galvanic/voltaic cells): chemical reactions produce electricity spontaneously. In a battery: anode (negative terminal, where oxidation occurs) and cathode (positive terminal, where reduction occurs).
pH measures how acidic or basic a solution is on a scale from 0 to 14. pH 7 = neutral (pure water). pH < 7 = acidic (more H⁺ ions). pH > 7 = basic/alkaline (more OH⁻ ions). Each whole number represents a 10× change in concentration. Strong acids have very low pH (hydrochloric acid ~1); strong bases have very high pH (sodium hydroxide ~14). Indicators change color at different pH values (litmus: red in acid, blue in base; universal indicator shows full range of colors).
Law of Conservation of Mass: mass is neither created nor destroyed in chemical reactions; total mass of reactants = total mass of products (Lavoisier, 1789). Mole: a counting unit for chemistry; 1 mole = 6.022 × 10²³ particles (Avogadro's number). Molar mass: mass of one mole of a substance in grams (equals atomic/formula mass in grams). Example: water (H₂O) molar mass = 18 g/mol. Allows converting between mass, moles, and number of particles.
Combustion is a rapid chemical reaction between a fuel and oxygen that releases heat and light. Complete combustion: fuel reacts with plentiful oxygen → carbon dioxide + water vapor (e.g., methane: CH₄ + 2O₂ → CO₂ + 2H₂O). Incomplete combustion (limited oxygen): produces carbon monoxide (CO, toxic) and soot (unburned carbon). Fire triangle: fuel + oxygen + heat — remove any one to extinguish fire. Common fuels: hydrocarbons (natural gas, gasoline, wood). Burning is an exothermic oxidation reaction.
Alloy: a mixture of two or more metals (or a metal and a nonmetal) with properties superior to the pure metals. Steel: iron + carbon (stronger and harder than iron). Stainless steel: iron + carbon + chromium + nickel (corrosion-resistant). Bronze: copper + tin. Brass: copper + zinc. Solder: tin + lead (low melting point for joining metals). Properties of materials: conductivity, strength, hardness, flexibility/ductility, melting point, corrosion resistance. Ceramics: hard, brittle, good insulators. Polymers: flexible, low density, electrical insulators.
Flame tests: technique to identify metal ions based on the color of light they emit when heated. When atoms are heated, electrons jump to higher energy levels; as they fall back, they emit photons of specific wavelengths (colors). Characteristic colors: lithium (red), sodium (yellow-orange), potassium (lilac/violet), copper (blue-green), barium (pale green), calcium (brick red), strontium (crimson). Used in fireworks to produce colors. More precise identification uses a spectroscope to analyze emission spectra.
Polymers are large molecules made of many repeating units (monomers) linked together. Natural polymers: DNA (nucleotide monomers), proteins (amino acid monomers), cellulose and starch (glucose monomers), rubber (isoprene monomers). Synthetic polymers: polyethylene (plastic bags, from ethylene monomers), nylon (clothing, from diamine + diacid), PVC, Teflon, polystyrene. Addition polymers: monomers add together directly. Condensation polymers: monomers join releasing a small molecule (usually water). Plastics are synthetic polymers and are largely non-biodegradable.
Corrosion: deterioration of metals due to chemical reactions with their environment. Rusting: specific to iron; iron reacts with oxygen and water → iron oxide (rust; Fe₂O₃·nH₂O). Conditions needed: iron, oxygen, and water (electrolyte accelerates corrosion). Rusting is an electrochemical process — iron loses electrons (oxidizes). Prevention methods: painting or coating (barrier), galvanizing (coating with zinc — zinc oxidizes sacrificially instead of iron), stainless steel (chromium forms protective oxide layer), cathodic protection (attaches more reactive metal).
Binary ionic compounds: name cation, then anion with -ide suffix (NaCl = sodium chloride; CaCl₂ = calcium chloride; MgO = magnesium oxide). When metal has multiple possible charges: use Roman numerals in parentheses (FeCl₂ = iron(II) chloride; FeCl₃ = iron(III) chloride). Binary covalent compounds (nonmetal + nonmetal): use Greek prefixes for number of atoms (CO₂ = carbon dioxide; N₂O₄ = dinitrogen tetroxide). Acids ending in -ic and -ous. Common bases end in -hydroxide.
Natural materials: wood, cotton, wool, silk, leather, stone, clay, rubber — come from living organisms or the environment. Synthetic (man-made) materials: nylon, polyester, plastic, Kevlar, fiberglass, synthetic rubber — made from chemical processes, often from petroleum. Advantages of synthetic: can be engineered for specific properties (strength, flexibility, durability). Disadvantages: often non-biodegradable, petroleum-based. Composite materials: combine materials (fiberglass = glass fibers in plastic resin; concrete = cement + aggregate; carbon fiber composites).
Greenhouse gases: CO₂, CH₄, N₂O, water vapor, ozone — absorb infrared radiation and warm the planet. Combustion of fossil fuels is the main source of CO₂ increase. Ozone: stratospheric ozone (good — blocks UV) vs. ground-level ozone (bad — component of smog). Acid rain: SO₂ and NOₓ from burning fossil fuels react with water → H₂SO₄ and HNO₃; damages ecosystems and infrastructure. Water treatment: coagulation, filtration, disinfection (chlorination or UV). Plastic pollution: non-degradable polymers accumulate in ocean (microplastics).
Percentage composition: mass of each element / molar mass of compound × 100%. Example: H₂O — H: (2×1)/18 × 100% = 11.1%; O: 16/18 × 100% = 88.9%. Empirical formula: simplest whole-number ratio of atoms in a compound. To find: convert percentage to grams, divide by molar mass to get moles, divide by smallest mole value, round to whole numbers. Molecular formula = empirical formula × n, where n = molecular mass / empirical formula mass. Example: empirical CH₂O, molecular C₆H₁₂O₆ (n=6).