Estonian Basic-School Chemistry Exam cheat sheet

Physical properties (melting point, boiling point, density, solubility, conductivity) are observed without changing the substance's identity; chemical properties describe how it reacts.

Pure substances have fixed melting and boiling points; impurities lower the melting point and widen the boiling range.

Density = mass ÷ volume (ρ = m/V); water's density is 1.0 g/cm³ at 4 °C, so objects denser than water sink.

Solubility rules: most nitrates and sodium, potassium, ammonium salts are soluble; most chlorides soluble except AgCl and PbCl₂; most sulfates soluble except BaSO₄, PbSO₄, CaSO₄ (slightly).

Metals are shiny, malleable, ductile and conduct heat and electricity; non-metals are usually brittle, dull and poor conductors (graphite is an exception).

A mixture can be separated by physical methods (filtration, evaporation, distillation, magnetism) because its components keep their properties.

Hardness, brittleness and elasticity are mechanical properties; a substance's use in engineering follows from these plus corrosion resistance.

Alloy = homogeneous mixture of a metal with other elements; alloying usually increases hardness and lowers melting point compared with the pure metal.

Atom: protons (Z, positive, in nucleus) + neutrons (nucleus, no charge) + electrons (negative, shells). Atomic number Z = proton count = electron count in a neutral atom; mass number A = Z + N.

Isotopes: same Z, different N; they share chemical properties but differ in mass (e.g. ¹²C and ¹⁴C).

Ion: atom that has gained electrons (anion, negative) or lost electrons (cation, positive); charge = protons − electrons.

Period number = number of electron shells; group number (main groups) = number of outer-shell electrons for groups 1–2 and 13–18.

Group 1 alkali metals: one outer electron, very reactive, reactivity increases down the group; group 17 halogens: seven outer electrons, reactivity decreases down the group.

Noble gases (group 18) have full outer shells, so they are chemically inert under ordinary conditions.

Ionic bond: electron transfer between metal and non-metal, held by electrostatic attraction; covalent bond: shared electron pair between non-metals.

Metallic bond: positive metal ions in a lattice surrounded by delocalised electrons, which explains conductivity and malleability.

Law of conservation of mass: atoms are neither created nor destroyed, so every equation must be balanced - the number of each element's atoms is equal on both sides.

Balancing procedure: write correct formulas, count atoms per element, adjust coefficients only (never subscripts), then re-check and add state symbols.

Reaction types: synthesis (A + B → AB), decomposition (AB → A + B), single displacement (A + BC → AC + B), double displacement (AB + CD → AD + CB), combustion.

Combustion of a hydrocarbon: CₓHᵧ + O₂ → CO₂ + H₂O; incomplete combustion also gives CO and soot.

Exothermic reactions release energy (combustion, neutralisation, rusting); endothermic reactions absorb it (thermal decomposition of CaCO₃, dissolving NH₄NO₃).

Activation energy is the minimum energy needed to start a reaction; a catalyst lowers it and speeds up the reaction without being consumed.

Reaction rate increases with higher temperature, higher concentration, larger surface area of solid, and presence of a catalyst.

Signs of a chemical change: gas evolution, precipitate formation, colour change, heat or light release - not just dissolving or melting.

Oxide classes: basic oxides (metal + oxygen, e.g. Na₂O, CaO) react with acids; acidic oxides (non-metal + oxygen, e.g. CO₂, SO₂) react with bases; amphoteric oxides (Al₂O₃, ZnO) react with both.

Metal + oxygen → basic oxide; non-metal + oxygen → acidic oxide; acidic oxide + water → acid; basic oxide + water → base (only for soluble bases).

Acid + metal → salt + hydrogen (metals above hydrogen in the activity series, e.g. Zn, Mg, Fe; not Cu).

Acid + base → salt + water (neutralisation); acid + carbonate → salt + water + CO₂; acid + metal oxide → salt + water.

Activity series (decreasing): K, Na, Ca, Mg, Al, Zn, Fe, Sn, Pb, H, Cu, Ag, Au - a more active metal displaces a less active one from its salt solution.

Base + acidic oxide → salt + water; base + salt → new base + new salt if a precipitate or gas forms.

Salt preparation routes: metal + acid, oxide + acid, hydroxide + acid, carbonate + acid, or precipitation from two soluble salts.

Common strong acids: HCl, H₂SO₄, HNO₃; common strong bases: NaOH, KOH, Ca(OH)₂; weak: CH₃COOH, NH₃·H₂O.

Solution = solute dissolved in solvent; in aqueous solutions water is the solvent. Concentration c = n/V (mol/dm³) or mass fraction w = m(solute)/m(solution) × 100%.

Dilution rule: n stays constant, so c₁V₁ = c₂V₂; adding water decreases concentration but does not change the amount of solute.

Solubility is the maximum mass of solute in 100 g solvent at a given temperature; for most solids it rises with temperature, for gases it falls.

pH scale 0–14 at 25 °C: pH < 7 acidic, pH = 7 neutral, pH > 7 basic; each unit is a tenfold change in H⁺ concentration.

Acids release H⁺ (protons) in water and turn litmus red; bases release OH⁻ and turn litmus blue; indicators: litmus, methyl orange (red in acid), phenolphthalein (pink in base).

Neutralisation: acid + base → salt + water; H⁺ + OH⁻ → H₂O. The enthalpy change is exothermic.

Titration: known concentration in the burette neutralises a measured volume of analyte with indicator; at the equivalence point n(acid) : n(base) follows the balanced equation.

Strong acid + strong base gives pH 7 at equivalence; weak acid + strong base gives pH > 7, so choose the indicator accordingly.

Hydrocarbons contain only C and H; alkanes CₙH₂ₙ₊₂ (single bonds, saturated), alkenes CₙH₂ₙ (one C=C, unsaturated), alkynes CₙH₂ₙ₋₂ (one C≡C).

Homologous series: each member differs from the next by CH₂; physical properties change gradually (boiling point rises with chain length).

IUPAC naming: longest carbon chain gives the root (-ane, -ene, -ol, -oic acid); number from the end nearest the functional group; state substituent positions.

Alcohols contain –OH: methanol, ethanol (C₂H₅OH), glycerol; ethanol is produced by fermentation of sugars and is oxidised to ethanoic acid.

Carboxylic acids contain –COOH: methanoic and ethanoic acid; they are weak acids and react with metals, bases and carbonates like inorganic acids.

Polymers: monomers join by addition (ethene → polyethene) or condensation (amino acids → proteins; glucose → starch); plastics are long-chain synthetic polymers.

Fossil fuels (coal, oil, natural gas) are hydrocarbon mixtures; fractional distillation separates them by boiling point into petrol, kerosene, diesel and heavy residues.

Environmental issues: CO₂ from combustion drives the greenhouse effect; SO₂ and NOₓ cause acid rain; unburnt hydrocarbons and CO are toxic; plastic waste resists degradation.

Amount of substance n = m/M, where m is mass in grams and M is molar mass in g/mol; also n = N/N_A (N_A = 6.02 × 10²³ mol⁻¹) and n = V/V_m (V_m = 22.4 dm³/mol at STP).

Molar mass M is the sum of relative atomic masses from the periodic table, e.g. M(H₂O) = 2×1 + 16 = 18 g/mol.

Concentration c = n/V, where V is in dm³ (1 dm³ = 1 L = 1000 cm³); rearrange to n = cV and V = n/c.

Mass fraction w = m(solute)/m(solution); to convert to concentration use c = (w × ρ)/M with ρ in g/dm³.

Stoichiometry steps: balance the equation → convert known quantity to moles → apply mole ratio from coefficients → convert back to the requested unit.

Mole ratio example: 2H₂ + O₂ → 2H₂O means 2 mol H₂ reacts with 1 mol O₂ producing 2 mol H₂O.

Limiting reactant: the reagent that runs out first; calculate the product amount from it, not from the excess reagent.

Percentage yield = (actual mass ÷ theoretical mass) × 100%; percentage composition of an element = (n × A_r ÷ M) × 100%.

Industrial processes: Haber–Bosch (N₂ + 3H₂ ⇌ 2NH₃, catalyst Fe, high pressure, moderate temperature) for fertilisers; contact process (SO₂ → SO₃, catalyst V₂O₅) for sulfuric acid.

Limestone (CaCO₃) is thermally decomposed to quicklime (CaO) and CO₂; CaO + water gives slaked lime Ca(OH)₂, used to neutralise acidic soil and flue gases.

Water treatment: coagulation, sedimentation, filtration and chlorination/ozonation for disinfection; hard water contains Ca²⁺ and Mg²⁺ and forms scale and poor soap lather.

Air composition by volume: about 78% N₂, 21% O₂, 0.9% Ar, 0.04% CO₂; pollutants include CO, SO₂, NOₓ, particulates and ground-level ozone.

Waste hierarchy: prevent, reuse, recycle, recover energy, dispose; separate collection of glass, paper, metal and biowaste improves recycling.

Green chemistry: atom economy, energy efficiency, use of renewable feedstock and non-toxic solvents; catalysts reduce energy demand.

Materials: glass (SiO₂-based), ceramics, cement and concrete, steel (Fe + C), and bioplastics are chosen by properties such as strength, corrosion resistance and recyclability.

Fertilisers supply N, P, K; excess nitrate and phosphate runoff causes eutrophication - algal blooms and oxygen depletion in water bodies.

Before any experiment: read the procedure, identify hazards from labels (corrosive, toxic, flammable, oxidising) and locate the eyewash, fire blanket and extinguisher.

Safety rules: goggles on, tie hair back, never taste or sniff chemicals directly, add acid to water (not water to acid), heat test tubes pointing away from people.

Glassware: measuring cylinder for approximate volumes, pipette for exact fixed volume, burette for titration, volumetric flask for a precise solution, beaker for mixing.

Separation techniques: filtration (insoluble solid from liquid), evaporation/crystallisation (soluble solid), simple distillation (solvent from solution), chromatography (dyes and mixtures).

Heating: use a water bath for flammable liquids, a Bunsen burner with a blue flame for high temperature; clamp and support apparatus securely.

Observation vs conclusion: record what you directly see (colour, gas, precipitate, temperature change) and only then interpret it in terms of a reaction.

Test for gases: H₂ - squeaky pop; O₂ - relights a glowing splint; CO₂ - turns limewater milky; NH₃ - turns damp red litmus blue.

Error types: random errors reduce by repeating and averaging; systematic errors (e.g. uncalibrated balance) shift all results the same way; report results with units and sensible significant figures.