Slovakia Maturita Chemistry cheat sheet
Slovakia Maturita Chemistry · Keep the key ideas within reach.
General chemistry: structure of matter, atomic structure, periodic table, bonding
Atomic number Z = number of protons; mass number A = protons + neutrons; isotopes differ in neutron count.
Electron configuration follows Aufbau principle: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p; Hund's rule: maximize unpaired electrons.
Periodic trends: atomic radius decreases across a period, increases down a group; electronegativity increases across, decreases down.
Ionic bond forms between metal and nonmetal via electron transfer; covalent bond involves sharing; metallic bond involves delocalized electrons.
VSEPR theory: electron pairs repel; molecular geometry depends on number of bonding and lone pairs (e.g., CH4 tetrahedral, H2O bent).
Polarity: a molecule is polar if it has polar bonds and asymmetric geometry (e.g., CO2 nonpolar, H2O polar).
Intermolecular forces: London dispersion < dipole-dipole < hydrogen bonding; hydrogen bonds occur when H is bonded to N, O, or F.
Oxidation number rules: free element = 0; monatomic ion = charge; oxygen usually -2, hydrogen +1; sum in neutral compound = 0.
Chemical calculations, stoichiometry and quantitative relationships
Mole concept: n = m/M; Avogadro constant NA = 6.022×10^23 mol^-1; molar volume at STP (0°C, 101.325 kPa) = 22.4 dm^3/mol.
Stoichiometry: use balanced equation coefficients to convert moles of reactant to moles of product.
Limiting reactant: reactant that produces the least amount of product; excess reactant remains after reaction.
Percent yield = (actual yield / theoretical yield) × 100%.
Concentration: c = n/V (mol/dm^3); dilution: c1V1 = c2V2.
Mass percent: (mass of component / total mass) × 100%.
Empirical formula from percent composition: divide by atomic masses, then by smallest ratio; molecular formula = empirical × n.
Gas laws: ideal gas equation pV = nRT; R = 8.314 J/(mol·K); at constant T and p, volumes react in simple whole-number ratios (Gay-Lussac).
Inorganic chemistry: elements, inorganic compounds and their reactions
Alkali metals (Group 1) react vigorously with water to form hydroxide and hydrogen gas: 2M + 2H2O → 2MOH + H2.
Halogens (Group 17) are diatomic, oxidizing agents; reactivity decreases down the group (F2 > Cl2 > Br2 > I2).
Common acids: HCl (hydrochloric), H2SO4 (sulfuric), HNO3 (nitric); common bases: NaOH, KOH, Ca(OH)2.
Oxides: basic (metal oxides, e.g., Na2O), acidic (nonmetal oxides, e.g., CO2, SO3), amphoteric (e.g., Al2O3, ZnO).
Reactivity series: K > Na > Ca > Mg > Al > Zn > Fe > Sn > Pb > H > Cu > Ag > Au; metals above H displace H2 from acids.
Displacement reaction: more reactive metal displaces less reactive from its salt solution (e.g., Zn + CuSO4 → ZnSO4 + Cu).
Ammonia (NH3) is a weak base; it forms NH4+ in acid; used in fertilizer production (Haber process).
Carbonates react with acids to produce CO2 gas: CaCO3 + 2HCl → CaCl2 + CO2 + H2O.
Organic chemistry: hydrocarbons, derivatives, nomenclature and reactions
IUPAC naming: longest carbon chain, lowest locants for substituents, suffix denotes principal functional group (-ol, -al, -one, -oic acid).
Alkanes: saturated, general formula CnH2n+2; undergo substitution (e.g., halogenation with Cl2/UV).
Alkenes: CnH2n, contain C=C; undergo addition (H2, HX, H2O) and polymerisation; test with bromine water (decolourised).
Alkynes: CnH2n-2, contain C≡C; undergo addition; terminal alkynes are weakly acidic.
Aromatic hydrocarbons (arenes) undergo electrophilic substitution (nitration, halogenation) rather than addition.
Alcohols: contain -OH; primary oxidise to aldehydes then carboxylic acids; secondary to ketones; tertiary resistant.
Carboxylic acids: contain -COOH; react with alcohols to form esters (Fischer esterification, acid catalyst); with bases to form salts.
Aldehydes and ketones: carbonyl group C=O; aldehydes reduce Tollens' and Fehling's reagents; ketones do not.
Biochemistry: biomolecules, metabolism and natural substances
Carbohydrates: monosaccharides (glucose, fructose) are reducing sugars; sucrose is non-reducing; starch and cellulose are polysaccharides of glucose.
Proteins: polymers of amino acids linked by peptide bonds; primary, secondary (α-helix, β-sheet), tertiary, quaternary structure.
Enzymes are biological catalysts; activity depends on pH and temperature; denaturation occurs at high temperature or extreme pH.
Lipids: fats and oils are triglycerides (glycerol + fatty acids); saturated fats are solid, unsaturated are liquid; phospholipids form membranes.
Nucleic acids: DNA and RNA are polymers of nucleotides; DNA contains deoxyribose and bases A, T, G, C; RNA contains ribose and uracil instead of thymine.
Glycolysis converts glucose to pyruvate, producing net 2 ATP and 2 NADH; occurs in cytoplasm.
Photosynthesis: 6CO2 + 6H2O → C6H12O6 + 6O2; light-dependent reactions produce ATP and NADPH; Calvin cycle fixes CO2.
Vitamins: water-soluble (C, B-complex) and fat-soluble (A, D, E, K); deficiency causes specific diseases (e.g., scurvy from vitamin C deficiency).
Chemical thermodynamics, kinetics, equilibrium and electrochemistry
First law of thermodynamics: ΔU = q + w; enthalpy change ΔH = qp at constant pressure; exothermic ΔH < 0, endothermic ΔH > 0.
Hess's law: total enthalpy change is independent of pathway; ΔH for reaction = sum of ΔHf(products) - sum of ΔHf(reactants).
Gibbs free energy: ΔG = ΔH - TΔS; spontaneous if ΔG < 0.
Reaction rate increases with concentration, temperature, surface area, and catalyst; catalyst lowers activation energy.
Equilibrium: dynamic state where forward and reverse rates are equal; equilibrium constant K = [products]^coefficients / [reactants]^coefficients.
Le Chatelier's principle: system shifts to counteract change in concentration, pressure, or temperature.
Electrochemistry: oxidation at anode, reduction at cathode; in galvanic cell, anode is negative, cathode positive; E°cell = E°cathode - E°anode.
Electrolysis: external voltage drives non-spontaneous reaction; products depend on electrode potentials and concentration.
Acids, bases, pH and analytical chemistry
Brønsted-Lowry acid is a proton donor; base is a proton acceptor; conjugate acid-base pairs differ by one H+.
pH = -log[H3O+]; pOH = -log[OH-]; pH + pOH = 14 at 25°C; acidic pH < 7, basic pH > 7.
Strong acids (HCl, HNO3, H2SO4) dissociate completely; weak acids (CH3COOH) partially; Ka expresses acid strength.
Neutralisation: acid + base → salt + water; titration uses burette and indicator to find equivalence point.
Indicators: phenolphthalein (colourless to pink, pH 8.2-10), methyl orange (red to yellow, pH 3.1-4.4).
Buffer solution resists pH change; composed of weak acid and its conjugate base (e.g., CH3COOH/CH3COO-).
Qualitative analysis: flame tests (Na yellow, K lilac, Ca brick-red, Cu blue-green); precipitation reactions identify ions (e.g., Cl- with AgNO3 gives white AgCl).
Redox titration (e.g., KMnO4) uses colour change of the titrant to indicate endpoint; calculations use mole ratios.
Laboratory work, safety and laboratory techniques
Safety: wear goggles and lab coat; never pipette by mouth; know location of eyewash, fire extinguisher, and safety shower.
Heating test tubes: point mouth away from people; use test-tube holder; do not heat closed systems.
Acid spill: neutralise with sodium bicarbonate (NaHCO3) and rinse with plenty of water; mercury spill requires special cleanup.
Filtration separates solid from liquid; distillation separates liquids by boiling point; crystallisation purifies solids.
Chromatography separates mixtures based on differential adsorption or solubility (paper, thin-layer, column).
Measuring volume: use graduated cylinder for approximate volumes; pipette and burette for precise volumes; read meniscus at eye level.
Drying agents: anhydrous CaCl2, silica gel; desiccator stores moisture-sensitive substances.
Waste disposal: never pour chemicals down the sink unless instructed; collect halogenated solvents and heavy-metal solutions separately.
Chemistry in society: environment, industry, materials and everyday life
Greenhouse gases: CO2, CH4, N2O, water vapour; contribute to global warming by absorbing infrared radiation.
Air pollutants: SO2 and NOx cause acid rain; CO binds to haemoglobin; particulate matter affects respiratory health.
Water treatment: coagulation, filtration, disinfection (chlorination or ozonation); eutrophication from excess nitrates and phosphates.
Haber process: N2 + 3H2 ⇌ 2NH3 (iron catalyst, high pressure, moderate temperature) for fertilisers.
Contact process: production of sulfuric acid from SO2 oxidation to SO3 (vanadium(V) oxide catalyst).
Polymers: addition polymers (polyethylene, PVC, polystyrene) and condensation polymers (PET, nylon); recycling codes identify plastics.
Fossil fuels: coal, oil, natural gas; combustion produces CO2; alternative energy sources include solar, wind, hydro, and hydrogen fuel cells.
Hard water contains Ca2+ and Mg2+ ions; causes scale; softened by ion exchange or precipitation.