Državna Matura Chemistry flashcards

32 cards. Read the term and the definition here. The deck at the top of this page is optional.

Electron configuration rule

Aufbau principle: electrons fill orbitals in order of increasing (n+l), and for equal (n+l) the lower n fills first. Hund's rule and Pauli exclusion apply within subshells.

VSEPR lone pair effect

Lone pairs occupy more space than bonding pairs, so they repel bonds more strongly. This compresses bond angles below the ideal, e.g. water's 104.5° versus tetrahedral 109.5°.

Hydrogen bond

A strong dipole–dipole attraction between H bonded to N, O or F and a lone pair on a neighbouring N, O or F. It raises boiling points and explains water's high boiling point.

Oxidation number

A bookkeeping charge assigned to an atom. Free elements are 0, monatomic ions equal their charge, oxygen is usually −2 and hydrogen +1, and the sum equals the species charge.

Oxidising agent

The species that accepts electrons and is itself reduced, causing another species to be oxidised. Its oxidation number decreases during the reaction.

Precipitation reaction

A double displacement in which two soluble salts exchange ions and one product is insoluble. Solubility rules predict the precipitate, e.g. AgCl from Ag⁺ and Cl⁻.

Molar volume at STP

One mole of any ideal gas occupies 22.4 dm³ at 0 °C and 101.3 kPa, or about 24.5 dm³ at 25 °C and 101.3 kPa. Used to convert gas volume to moles.

Limiting reagent

The reactant that is completely consumed first and therefore determines the maximum amount of product. Compare mole ratios from the balanced equation to identify it.

Hess's law

The enthalpy change of a reaction depends only on initial and final states, not the pathway. Reversing a step changes the sign of its ΔH; multiplying coefficients multiplies ΔH.

Standard enthalpy of formation

The enthalpy change when one mole of a compound forms from its elements in their standard states at 298 K and 100 kPa. Elements in standard states have ΔH_f = 0.

Amphoteric oxide

An oxide that reacts with both acids and bases to form salts. Aluminium oxide and zinc oxide are examples; they react with HCl and with NaOH.

Halogen displacement

A halogen higher in group 17 displaces halide ions of a halogen lower down, because oxidising strength decreases down the group. Chlorine liberates bromine from bromide solution.

Transition metal properties

Typical features include variable oxidation states, coloured compounds, catalytic activity and paramagnetism due to unpaired d electrons. Zn is not a typical transition metal in this sense.

Flame test colours

Characteristic emission colours identify metal ions: sodium yellow, potassium lilac, calcium brick-red, copper blue-green, lithium crimson. A clean nichrome wire is used.

Markovnikov's rule

In addition of HX to an unsymmetrical alkene, hydrogen adds to the carbon with more hydrogens, so the halogen ends up on the more substituted carbon.

Esterification

A carboxylic acid reacts with an alcohol in the presence of an acid catalyst to form an ester and water. The reaction is reversible and the ester is often fragrant.

Tollens' reagent test

Ammoniacal silver nitrate oxidises aldehydes to carboxylate while silver ions reduce to a silver mirror. Ketones give no reaction, distinguishing the two carbonyl types.

Peptide bond

An amide linkage formed between the carboxyl group of one amino acid and the amino group of another, releasing water. Repeating peptide bonds form the protein backbone.

Equilibrium constant K_c

The ratio of product concentrations to reactant concentrations, each raised to its stoichiometric coefficient, at equilibrium. Pure solids and liquids are excluded. It depends only on temperature.

Le Chatelier's principle

When a system at equilibrium is disturbed, it shifts in the direction that opposes the change. Concentration, pressure and temperature changes shift position; only temperature changes K.

Reaction quotient Q

The same expression as K but using current, not equilibrium, concentrations. If Q < K the reaction proceeds forward; if Q > K it proceeds in reverse.

pH definition

pH = −log[H⁺], so [H⁺] = 10^(−pH). At 25 °C pH + pOH = 14 and neutral solutions have pH 7 because K_w = 1.0 × 10⁻¹⁴.

Buffer solution

A weak acid and its conjugate base (or weak base and conjugate acid) that resists pH change on adding small amounts of acid or base. pH follows the Henderson–Hasselbalch equation.

Equivalence point in titration

The point where the acid and base have reacted in exact stoichiometric amounts. Its pH depends on the salt formed, not always 7.

Standard cell potential

E°_cell = E°_cathode − E°_anode, measured against the standard hydrogen electrode (0.00 V). A positive value indicates a spontaneous galvanic reaction.

Electrolysis of brine

Passing current through concentrated sodium chloride solution gives hydrogen at the cathode, chlorine at the anode and sodium hydroxide remaining in solution.

Faraday's law of electrolysis

The amount of substance deposited is proportional to the charge passed: n(e⁻) = It/F, with F = 96 485 C mol⁻¹. Mass = M·I·t/(z·F).

Titration procedure

Rinse and fill the burette with titrant, record the initial reading, add titrant with swirling until the indicator changes, then repeat for concordant titres.

Dilution formula

c₁V₁ = c₂V₂, because the moles of solute stay constant when solvent is added. Concentrations must use the same units and volumes the same units.

Percentage yield

(actual yield/theoretical yield) × 100. Theoretical yield comes from stoichiometry with the limiting reagent; losses and side reactions make actual yield lower.

Significant figures in results

The answer carries no more significant figures than the least precise measurement used. Burette readings are recorded to 0.05 cm³ and molar masses to the given data precision.

Ideal gas law

pV = nRT with R = 8.314 J mol⁻¹ K⁻¹, p in Pa, V in m³ and T in kelvin. Rearranged, M = mRT/(pV) gives molar mass from gas data.