Državna Matura Chemistry cheat sheet

Electron configuration fills by increasing (n+l); for equal (n+l) the lower n fills first; exceptions Cr [Ar]3d⁵4s¹ and Cu [Ar]3d¹⁰4s¹.

Quantum numbers: n ≥ 1, l = 0…n−1, m_l = −l…+l, m_s = ±½; no two electrons in an atom share all four.

Periodic trends across a period: atomic radius decreases, ionisation energy and electronegativity increase (noble gases excepted for IE).

Bond polarity follows ΔEN: <0.4 nonpolar covalent, 0.4–1.7 polar covalent, >1.7 predominantly ionic (Pauling scale).

VSEPR: count bonding pairs plus lone pairs around the central atom; lone pairs repel more than bond pairs, compressing bond angles.

Hybridisation: 2 electron domains sp (180°), 3 domains sp² (120°), 4 domains sp³ (109.5°); lone pairs reduce angles below ideal.

Intermolecular forces ranked: ion–dipole > hydrogen bond > dipole–dipole > London dispersion; H-bonds need H bonded to N, O or F.

Metallic bonding: lattice of cations in a delocalised electron sea explains conductivity, malleability and lustre.

Balance by conservation of atoms and charge; in ionic equations the sum of charges must be equal on both sides.

Common reaction classes: synthesis, decomposition, single displacement, double displacement (precipitation), combustion, neutralisation, redox.

Oxidation is loss of electrons (oxidation number increases); reduction is gain; the oxidising agent is itself reduced.

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

Activity series predicts single displacement: a metal displaces ions of metals below it (e.g. Zn + Cu²⁺ → Zn²⁺ + Cu).

Solubility: most nitrates and alkali-metal salts are soluble; most chlorides soluble except AgCl, PbCl₂; most sulfates soluble except BaSO₄, PbSO₄, CaSO₄ (sparingly).

Gas evolution signs: CO₂ from carbonates plus acid, H₂ from active metal plus acid, O₂ from peroxide decomposition, NH₃ from ammonium salts plus base.

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

n = m/M; at STP (0 °C, 101.3 kPa) 1 mol of ideal gas occupies 22.4 dm³; at 25 °C, 24.5 dm³.

Mole ratio from balanced equation converts moles of one species to another; limiting reagent is the one giving the least product.

Percentage yield = (actual/theoretical) × 100; percentage composition = (mass of element/M) × 100.

Empirical formula from mass or percentage composition: divide by atomic masses, then divide by the smallest ratio.

q = m·c·ΔT (c_water = 4.18 J g⁻¹ K⁻¹); ΔH = −q/n for the reaction as written, sign from the thermometer direction.

Hess's law: ΔH of an overall reaction is independent of pathway; reverse a step to change the sign, multiply ΔH when multiplying coefficients.

ΔH_rxn = ΣΔH_f(products) − ΣΔH_f(reactants); ΔH_f of an element in its standard state is zero.

Bond enthalpy estimate: ΔH ≈ Σ(bonds broken) − Σ(bonds formed); exothermic when more energy is released than absorbed.

Groups: alkali metals (1), alkaline earths (2), halogens (17), noble gases (18); periods numbered by highest principal quantum number.

Alkali metals react with water giving hydroxide and H₂; reactivity increases down the group as ionisation energy falls.

Halogen reactivity decreases down the group; a halogen displaces halide ions of elements below it (Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂).

Transition metals show variable oxidation states, coloured complexes, catalytic activity and paramagnetism from unpaired d electrons.

Common oxides: Na₂O and MgO basic, Al₂O₃ amphoteric, SO₂ and CO₂ acidic, NO and CO neutral.

Nitrogen: N₂ inert (triple bond); NH₃ basic; HNO₃ strong oxidising acid; nitrates decompose on heating.

Sulfur: SO₂ reducing and bleaching, SO₃ forms H₂SO₄; H₂S is a weak acid and reducing agent giving black PbS with Pb²⁺.

Halide identification: AgNO₃ gives white AgCl (soluble in NH₃), pale yellow AgBr (partly soluble), yellow AgI (insoluble).

Homologous series differ by CH₂; general formulas: alkanes CₙH₂ₙ₊₂, alkenes CₙH₂ₙ, alkynes CₙH₂ₙ₋₂, alcohols CₙH₂ₙ₊₁OH.

IUPAC naming: longest chain, lowest locants, suffix by priority (−oic acid > −al > −one > −ol > −ene > −ane).

Alkanes undergo radical substitution with halogens (UV light); alkenes undergo electrophilic addition (H₂, HX, H₂O, Br₂) and decolourise bromine water.

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.

Alcohols: primary oxidise to aldehydes then carboxylic acids; secondary to ketones; tertiary resist oxidation; esterification with acids gives esters plus water.

Functional group tests: Fehling's/Tollens' for aldehydes, NaHCO₃ fizz for carboxylic acids, Br₂ water for unsaturation, 2,4-DNP for carbonyls.

Carbohydrates: monosaccharides (glucose, fructose) reduce Tollens'; sucrose is a non-reducing disaccharide; starch gives a blue colour with iodine.

Proteins are polyamides of α-amino acids linked by peptide bonds; denaturation disrupts secondary and tertiary structure, not primary sequence.

At equilibrium forward and reverse rates are equal and concentrations stay constant, but they are not necessarily equal.

K_c = [products]^coefficients / [reactants]^coefficients; pure solids and liquids are omitted.

Reaction quotient Q compared with K: Q < K shifts right, Q > K shifts left, Q = K at equilibrium.

Le Chatelier: adding reactant shifts right; increasing pressure shifts toward fewer gas moles; raising temperature shifts in the endothermic direction.

A catalyst speeds attainment of equilibrium but changes neither K nor the equilibrium position.

For gases K_p = K_c(RT)^Δn, where Δn is the change in moles of gas (products − reactants).

Only temperature changes K; concentration and pressure changes shift position but leave K unchanged.

Haber process: N₂ + 3H₂ ⇌ 2NH₃ is exothermic, so high yield needs low temperature and high pressure, with a catalyst for rate.

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

pH = −log[H⁺]; [H⁺] = 10^(−pH); pH + pOH = 14 at 25 °C; K_w = 1.0 × 10⁻¹⁴.

Strong acids (HCl, HNO₃, H₂SO₄) ionise completely; weak acids (CH₃COOH, H₂CO₃) partially ionise with K_a = [H⁺][A⁻]/[HA].

For a weak acid, [H⁺] ≈ √(K_a·c); for a weak base, [OH⁻] ≈ √(K_b·c).

Henderson–Hasselbalch: pH = pK_a + log([base]/[acid]); a buffer resists pH change on adding small amounts of acid or base.

At the equivalence point the pH depends on the salt: strong acid–strong base pH 7, weak acid–strong base >7, strong acid–weak base <7.

Titration calculation: c₁V₁/n₁ = c₂V₂/n₂, with n from the balanced equation.

Indicators change colour over roughly pK_a ± 1; phenolphthalein suits weak acid–strong base, methyl orange strong acid–weak base titrations.

Oxidation at the anode, reduction at the cathode; in a galvanic cell the anode is negative, in an electrolytic cell it is positive.

Standard cell potential E°_cell = E°_cathode − E°_anode; a positive value means the reaction is spontaneous.

In cell notation the anode is written on the left; the salt bridge maintains electrical neutrality and completes the circuit.

Standard hydrogen electrode (SHE) is assigned E° = 0.00 V and is the reference for all standard potentials.

Electrolysis of molten salts or aqueous solutions: at the cathode the species with the higher reduction potential is discharged, water competing with metal ions.

Electrolysis of brine gives H₂ at the cathode, Cl₂ at the anode and NaOH in solution.

Faraday's laws: moles of electrons = I·t/F (F = 96 485 C mol⁻¹); mass deposited = (M·I·t)/(z·F).

Corrosion of iron is an electrochemical process; sacrificial anodes (Zn, Mg) protect steel because they are more easily oxidised.

Volumetric flask: dissolve, transfer with rinsing, fill to the mark with the meniscus at eye level; do not heat it.

Burette is rinsed with the titrant and read at the bottom of the meniscus; the burette tip must be free of air bubbles before titrating.

Pipette delivers a fixed volume; rinse with the solution it will deliver; drain by gravity, touching the tip to the flask wall.

Filtration separates insoluble solid from liquid; distillation separates liquids by boiling point; extraction uses differing solubility.

Drying agents (anhydrous CaCl₂, Na₂SO₄) remove water; desiccators store hygroscopic solids.

Safety: fume hood for volatile or toxic reagents, goggles and gloves, never pipette by mouth, add acid to water not water to acid.

Qualitative tests: flame tests (Na yellow, K lilac, Ca brick-red, Cu blue-green); BaCl₂ for sulfate (white precipitate).

Significant figures in measurements: record the volume to the precision of the instrument, e.g. burette to 0.05 cm³, and calculate uncertainty from repeated titres.

Dimensional analysis: set up each step so units cancel, leaving the requested unit.

Density ρ = m/V; convert cm³ to dm³ by dividing by 1000 before using molar volume.

Concentration c = n/V (mol dm⁻³); dilution c₁V₁ = c₂V₂; mass concentration = c·M.

Percentage error = |measured − accepted|/accepted × 100; mean of repeated titres should exclude outliers.

Avogadro's number N_A = 6.022 × 10²³ mol⁻¹; number of particles = n·N_A.

Molar mass from formula: sum atomic masses; molar mass of a gas from ideal gas law M = mRT/(pV).

Graph interpretation: slope of a linear plot gives the rate or constant; extrapolate to the y-intercept only where the relationship is linear.

Report answers with correct significant figures and clear units; a result without units is incomplete.