Bulgaria DZI Chemistry cheat sheet

Atomic number Z = number of protons; in a neutral atom it also equals the electron count; mass number A = protons + neutrons, so neutrons = A − Z.

Isotopes share Z but differ in neutron number; relative atomic mass is the abundance-weighted mean of isotope masses.

Electron configuration follows the Aufbau order 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p, with Pauli exclusion (max 2 electrons per orbital) and Hund's rule (parallel spins fill degenerate orbitals first).

An orbital holds at most 2 electrons; subshell capacities are s = 2, p = 6, d = 10, f = 14; a shell n holds at most 2n² electrons.

Across a period, nuclear charge rises and atomic radius falls; down a group, added shells increase radius.

Ionisation energy increases across a period and decreases down a group; each successive ionisation energy is larger than the previous one.

Metallic character decreases left to right and increases down a group; nonmetals occupy the upper right of the periodic table.

Groups: alkali metals (1), alkaline earths (2), halogens (17), noble gases (18); the s-, p-, d- and f-blocks follow the subshell being filled.

Anions form by gaining electrons (radius larger than the atom); cations form by losing electrons (radius smaller than the atom).

Ionic bonding: electron transfer between a metal of low ionisation energy and a nonmetal of high electron affinity; lattice energy rises with higher ionic charge and smaller ionic radius.

Covalent bonding: shared electron pairs; bond order and bond energy rise from single to double to triple bonds while bond length falls.

Electronegativity difference rules of thumb: below about 0.4 nonpolar covalent, 0.4–1.7 polar covalent, above 1.7 predominantly ionic.

VSEPR: electron pairs repel and spread out; lone pairs repel more strongly than bonding pairs, reducing bond angles (CH4 109.5°, NH3 107°, H2O 104.5°).

Molecular polarity requires polar bonds that do not cancel by symmetry; CO2 is nonpolar, H2O and NH3 are polar.

Intermolecular forces in increasing strength: London dispersion < dipole–dipole < hydrogen bonding; hydrogen bonds occur with H attached to N, O or F.

Giant covalent solids (diamond, SiO2) have high melting points; molecular solids are soft with low melting points; metals conduct via delocalised electrons.

Allotropes are different structural forms of one element (diamond/graphite, O2/O3, white/red phosphorus).

Amount of substance n = m / M (mol), and for gases at standard conditions n = V / 22.4 (V in litres, 22.4 L/mol).

Avogadro's number NA = 6.02 × 10²³ particles per mole; number of particles = n × NA.

Molar concentration c = n / V (mol/L); dilution obeys c1V1 = c2V2.

Balance equations by conserving each element and charge; use the mole ratio from coefficients to convert between substances.

Limiting reagent: compute n/coefficient for each reactant; the smallest quotient is the limiting one and controls the yield.

Percentage yield = (actual mass / theoretical mass) × 100; mass of element in a compound = (Ar × atoms / Mr) × mass of sample.

Reaction types: combination, decomposition, single displacement (activity series), double displacement (driven by precipitate, gas or water formation), combustion.

Gas volume ratio equals mole ratio at the same temperature and pressure (Gay-Lussac/Avogadro).

Empirical formula from percentage composition: divide each percentage by the atomic mass, then divide by the smallest quotient.

Electrolytes dissociate into ions in water (salts, strong acids, strong bases); nonelectrolytes do not conduct because they yield no ions.

Strong electrolytes dissociate essentially completely, weak electrolytes partially; degree of dissociation α = dissociated amount / total amount.

Arrhenius acids release H3O+ (H+) in water, bases release OH−; the Brønsted–Lowry view is proton donor/acceptor.

pH = −log[H+]; pOH = −log[OH−]; at 25 °C pH + pOH = 14 and Kw = [H+][OH−] = 1 × 10⁻¹⁴.

Neutralisation: acid + base → salt + water; n(H+) = n(OH−) at the equivalence point.

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

Solubility product Ksp applies to sparingly soluble salts; a precipitate forms when the ion product exceeds Ksp.

Colligative properties depend on particle count: boiling-point elevation and freezing-point depression are larger for electrolytes that dissociate into more ions.

Electrical conductivity of a solution increases with ion concentration and ion mobility.

Hydrogen: lightest element, forms H2; reacts with halogens and oxygen; a strong reducing agent in the activity series.

Halogens F2 > Cl2 > Br2 > I2 in oxidising strength; displacement reactions follow this order (Cl2 displaces Br− and I−).

Hydrogen halides dissolve to give strong acids HF < HCl < HBr < HI in acid strength; HF is weak and attacks glass (SiO2).

Oxygen and sulfur: O2 supports combustion; SO2 is a reducing agent, SO3 + H2O → H2SO4; concentrated H2SO4 is a dehydrating and oxidising agent.

Nitrogen: N2 is inert (triple bond); NH3 is a weak base; HNO3 is a strong oxidising acid giving NO or NO2 with metals.

Phosphorus: white phosphorus (P4) is reactive and toxic, red phosphorus is stable; P2O5 + H2O → H3PO4.

Carbon: CO is a neutral toxic reducing agent; CO2 is acidic (CO2 + H2O ⇌ H2CO3) and turns limewater milky (CaCO3 precipitate).

Silicon: SiO2 is an acidic oxide reacting with HF and molten alkali; silicates form the basis of glass.

Noble gases have filled valence shells, very high ionisation energies and essentially no chemical reactivity under ordinary conditions.

Activity (reactivity) series: K > Na > Ca > Mg > Al > Zn > Fe > Sn > Pb > (H) > Cu > Ag > Au; metals above H displace H2 from dilute acids.

Metal + acid → salt + H2 (for metals above hydrogen); very active metals (Na, K, Ca) also react with cold water to give hydroxide + H2.

Thermal decomposition of carbonates: CaCO3 → CaO + CO2 on heating; hydroxides of active metals decompose similarly.

Amphoteric hydroxides and oxides: Al(OH)3, Zn(OH)2, Al2O3, ZnO dissolve in both acid and alkali.

Aluminium is extracted by electrolysis of alumina in molten cryolite; iron in a blast furnace from Fe2O3 with CO as reducing agent.

Rusting requires both oxygen and water; galvanising (Zn coating) protects iron by sacrificial oxidation of zinc.

Copper: Cu + dilute acid gives no H2, but Cu reacts with concentrated HNO3 (→ NO2) and dilute HNO3 (→ NO).

Transition metals show variable oxidation states (Fe²⁺/Fe³⁺, Cu⁺/Cu²⁺), coloured ions and catalytic activity.

Qualitative tests: flame colours (Na yellow, K lilac, Ca brick-red, Cu blue-green); Fe³⁺ + SCN− gives blood-red colour; BaSO4 white precipitate confirms sulfate.

Alkanes CnH2n+2: saturated, tetrahedral carbon, free-radical substitution with halogens under UV light; low reactivity otherwise.

Alkenes CnH2n: one C=C, planar sp² carbons, electrophilic addition of H2, HX, X2 and H2O; decolourise bromine water.

Alkynes CnH2n−2: one C≡C, linear sp carbons, addition reactions and acidic terminal hydrogen.

Markovnikov's rule: in HX addition to an unsymmetrical alkene, H adds to the carbon bearing more hydrogens.

Aromatic hydrocarbons: benzene C6H6 is planar with delocalised π electrons and undergoes substitution rather than addition.

Isomerism: chain, position and geometric (cis/trans) isomerism occur in hydrocarbons; cis/trans requires restricted rotation about C=C.

Combustion: complete gives CO2 + H2O; incomplete gives CO or soot; alkenes and alkynes burn with a smokier flame than alkanes.

Polymerisation: ethene → polyethene, propene → polypropene by addition polymerisation across the double bond.

Naming: longest chain gives the parent, lowest locants for substituents, suffix -ane/-ene/-yne.

Alcohols R–OH: hydrogen bonding raises boiling points; oxidised to aldehydes (primary, mild) or carboxylic acids (strong oxidant); tertiary alcohols resist oxidation.

Phenols: –OH on the aromatic ring, weakly acidic, give violet colour with FeCl3, undergo electrophilic substitution on the ring.

Aldehydes R–CHO: reduce Tollens' reagent (silver mirror) and Fehling's solution; ketones do not.

Carboxylic acids R–COOH: weak acids, react with carbonates to release CO2, form esters with alcohols under acid catalysis (esterification).

Esters R–COO–R': pleasant odours, hydrolysed by acid or alkali (saponification) to acid/salt plus alcohol.

Fats and oils are triesters of glycerol; soaps are alkali-metal salts of long-chain fatty acids.

Amines R–NH2: basic (lone pair on N), more basic than ammonia for alkylamines in water; react with acids to give salts.

Amino acids contain –NH2 and –COOH, are amphoteric and join by peptide bonds to form proteins.

Carbohydrates: glucose is a reducing sugar (aldehyde group); sucrose is nonreducing; starch gives a blue colour with iodine.

Enthalpy change ΔH = H(products) − H(reactants); negative for exothermic, positive for endothermic reactions.

Hess's law: ΔH of a reaction is path-independent; sum the steps' enthalpies to obtain the overall value.

Bond breaking absorbs energy, bond formation releases energy; ΔH ≈ energy in broken bonds − energy in formed bonds.

Rate of reaction increases with concentration, surface area, temperature and catalysts; rate = change in amount / time.

Collision theory: only collisions with sufficient energy (≥ activation energy) and correct orientation are effective.

A catalyst lowers activation energy without changing ΔH or the equilibrium position.

Le Chatelier's principle: raising temperature shifts an endothermic direction; raising pressure shifts toward fewer gas moles; catalyst changes only the rate of approach.

Equilibrium constant K = [products]^coefficients / [reactants]^coefficients (for gases Kp uses partial pressures); K depends only on temperature.

For a reaction at equilibrium, Q < K means the forward reaction is favoured; Q > K means the reverse reaction is favoured.