Armenia Chemistry Unified Exam cheat sheet

Atomic number Z = number of protons; mass number A = protons + neutrons; isotopes have same Z, different A.

Electron configuration follows Aufbau principle: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p.

Hund's rule: within a subshell, electrons occupy degenerate orbitals singly with parallel spins before pairing.

Pauli exclusion principle: no two electrons in an atom can have the same four quantum numbers.

Periodic law: properties of elements are periodic functions of their atomic numbers.

Groups (vertical columns) have similar valence electron configurations; periods (horizontal rows) show trends.

Effective nuclear charge increases across a period, causing atomic radius to decrease and ionization energy to increase.

Quantum numbers: n (principal, size), l (azimuthal, shape), ml (magnetic, orientation), ms (spin).

Ionic bond: electrostatic attraction between oppositely charged ions formed by electron transfer; typically metal + nonmetal.

Covalent bond: sharing of electron pairs between atoms; bond order = number of shared electron pairs.

Electronegativity difference > 1.7 generally indicates ionic character; 0.4–1.7 polar covalent; < 0.4 nonpolar covalent.

VSEPR theory: electron domains repel to maximize separation; lone pairs repel more than bonding pairs, reducing bond angles.

Hybridization: steric number 2 → sp; 3 → sp²; 4 → sp³; 5 → sp³d; 6 → sp³d².

Intermolecular forces strength: ion-dipole > hydrogen bonding > dipole-dipole > London dispersion.

Metallic bonding: delocalized valence electrons in a lattice of cations; explains conductivity and malleability.

Crystal lattice energy increases with higher ionic charges and smaller ionic radii (Coulomb's law).

Atomic radius increases down a group and decreases across a period (left to right).

Ionization energy generally increases across a period and decreases down a group; exceptions occur at filled/half-filled subshells.

Electron affinity becomes more negative across a period (except noble gases) and less negative down a group.

Metallic character increases down a group and decreases across a period.

Group 1 alkali metals react vigorously with water: 2M + 2H₂O → 2MOH + H₂; reactivity increases down the group.

Group 17 halogens: F₂ > Cl₂ > Br₂ > I₂ in oxidizing power; displacement reactions follow this order.

Group 18 noble gases have full valence shells (He: 1s²; others ns²np⁶) and low reactivity.

Oxides of nonmetals are acidic; oxides of metals are basic; amphoteric oxides (Al₂O₃, ZnO) react with both acids and bases.

Transition metals have partially filled d subshells (or form ions with them) and show variable oxidation states.

Ligand: Lewis base that donates an electron pair to a central metal ion to form a coordinate covalent bond.

Coordination number = number of donor atoms bonded to the central metal ion.

Common geometries: coordination number 2 linear; 4 tetrahedral or square planar; 6 octahedral.

Crystal field splitting: d orbitals split into lower and higher energy sets; Δ (splitting) depends on ligand and metal.

Spectrochemical series (weak to strong field): I⁻ < Br⁻ < Cl⁻ < F⁻ < OH⁻ < H₂O < NH₃ < en < NO₂⁻ < CN⁻ < CO.

High-spin complexes have small Δ and maximum unpaired electrons; low-spin complexes have large Δ and minimum unpaired electrons.

Color in transition metal complexes arises from d–d electron transitions; complementary color is observed.

1 mole = 6.022 × 10²³ particles (Avogadro's number).

Molar mass (g/mol) = mass of one mole of substance; numerically equal to atomic/molecular mass in amu.

n = m / M; n = N / N_A; n = V / V_m (at STP, V_m = 22.4 L/mol for ideal gas).

Balanced chemical equation gives mole ratios (stoichiometric coefficients) for reactants and products.

Limiting reactant: reactant that produces the least amount of product; excess reactant remains after reaction.

Percent yield = (actual yield / theoretical yield) × 100%.

Empirical formula = simplest whole-number ratio of atoms; molecular formula = (empirical formula)ₙ where n = molar mass / empirical formula mass.

Mass percent of element = (mass of element in 1 mol compound / molar mass of compound) × 100%.

Rate of reaction = change in concentration of reactant or product per unit time; usually positive and expressed as −Δ[reactant]/Δt or Δ[product]/Δt.

Rate law: rate = k[A]^m[B]^n; m and n are reaction orders determined experimentally, not from coefficients.

Zero-order reaction: rate = k; [A] = [A]₀ − kt; half-life t₁/₂ = [A]₀/(2k).

First-order reaction: rate = k[A]; ln[A] = ln[A]₀ − kt; t₁/₂ = 0.693/k (independent of initial concentration).

Second-order reaction: rate = k[A]²; 1/[A] = 1/[A]₀ + kt; t₁/₂ = 1/(k[A]₀).

Arrhenius equation: k = A e^(−Ea/RT); ln(k₂/k₁) = (Ea/R)(1/T₁ − 1/T₂).

Catalyst increases reaction rate by lowering activation energy; it does not change ΔH or equilibrium position.

Collision theory: reaction occurs when molecules collide with sufficient energy (≥ Ea) and proper orientation.

First law of thermodynamics: ΔU = q + w; for work done on the system w = −PΔV (constant pressure).

Enthalpy change ΔH = q_p (heat exchanged at constant pressure); ΔH < 0 exothermic, ΔH > 0 endothermic.

Hess's law: ΔH for a reaction is the sum of ΔH values for any series of steps into which the reaction can be divided.

Standard enthalpy of reaction: ΔH°_rxn = ΣnΔH°_f(products) − ΣmΔH°_f(reactants).

Entropy S is a measure of disorder; ΔS > 0 for increasing disorder (e.g., solid → liquid → gas).

Gibbs free energy: ΔG = ΔH − TΔS; ΔG < 0 spontaneous (favorable), ΔG > 0 nonspontaneous, ΔG = 0 at equilibrium.

Equilibrium constant K: for aA + bB ⇌ cC + dD, K = [C]^c[D]^d / ([A]^a[B]^b); pure solids and liquids are omitted.

Le Chatelier's principle: a system at equilibrium shifts to counteract a change in concentration, pressure, or temperature.

For gases, K_p = K_c(RT)^Δn, where Δn = moles gaseous products − moles gaseous reactants.

Molarity (M) = moles of solute / liters of solution; molality (m) = moles of solute / kg of solvent.

Mass percent = (mass solute / mass solution) × 100%; ppm = (mass solute / mass solution) × 10⁶.

Dilution: M₁V₁ = M₂V₂ (moles of solute constant).

Raoult's law: vapor pressure of solution P = X_solvent P°_solvent; nonvolatile solute lowers vapor pressure.

Colligative properties depend on number of solute particles: boiling point elevation ΔT_b = K_b·m·i; freezing point depression ΔT_f = K_f·m·i.

Osmotic pressure π = iMRT (van't Hoff factor i accounts for dissociation).

Solubility of most solids increases with temperature; solubility of gases decreases with increasing temperature (Henry's law: S_gas = k_H P_gas).

Like dissolves like: polar solutes dissolve in polar solvents; nonpolar in nonpolar.

Arrhenius acid releases H⁺ (H₃O⁺) in water; Arrhenius base releases OH⁻ in water.

Brønsted–Lowry acid is a proton donor; base is a proton acceptor; conjugate acid-base pairs differ by one H⁺.

Strong acids/bases ionize completely; weak acids/bases ionize partially (equilibrium).

pH = −log[H⁺]; pOH = −log[OH⁻]; pH + pOH = 14 at 25 °C; [H⁺][OH⁻] = 1.0 × 10⁻¹⁴.

For weak acid HA: K_a = [H⁺][A⁻]/[HA]; for weak base B: K_b = [BH⁺][OH⁻]/[B]; K_a·K_b = K_w.

Salt hydrolysis: salt of strong acid + strong base → neutral; strong acid + weak base → acidic; weak acid + strong base → basic.

Buffer solution resists pH change; Henderson–Hasselbalch: pH = pK_a + log([A⁻]/[HA]).

Electrolytes conduct electricity in solution or molten state; nonelectrolytes do not.

Oxidation is loss of electrons (increase in oxidation number); reduction is gain of electrons (decrease in oxidation number).

Oxidizing agent is reduced; reducing agent is oxidized.

Standard cell potential: E°_cell = E°_cathode − E°_anode (reduction potentials).

Spontaneous redox reaction: E°_cell > 0; ΔG° = −nFE°_cell.

In a galvanic cell, oxidation occurs at the anode (negative electrode) and reduction at the cathode (positive electrode).

In electrolysis, the anode is positive and the cathode is negative; external power drives nonspontaneous reactions.

Faraday's laws: mass deposited ∝ charge passed; m = (M·I·t)/(n·F), F = 96485 C/mol e⁻.

Nernst equation: E = E° − (0.0592/n) log Q at 25 °C.

Alkanes: CₙH₂ₙ₊₂, single bonds, sp³ carbons, substitution reactions (e.g., halogenation with light).

Alkenes: CₙH₂ₙ, one C=C double bond, sp² carbons, addition reactions (e.g., H₂, Br₂, HX, H₂O).

Alkynes: CₙH₂ₙ₋₂, one C≡C triple bond, sp carbons, addition reactions; terminal alkynes are weakly acidic.

Aromatic hydrocarbons (arenes) contain a benzene ring; undergo substitution rather than addition.

Alcohols (R–OH): hydroxyl group; primary/secondary/tertiary; oxidation yields aldehydes/ketones or carboxylic acids.

Aldehydes (R–CHO) and ketones (R–CO–R') contain carbonyl group; aldehydes are easily oxidized (Tollens' and Fehling's tests).

Carboxylic acids (R–COOH) are weak acids; react with alcohols to form esters (Fischer esterification, acid catalyst).

Amines (R–NH₂) are basic; amides (R–CO–NH₂) are neutral; amino acids contain both –NH₂ and –COOH.

Always add concentrated acid to water (not water to acid) to control heat release.

Read burette volumes at the bottom of the meniscus at eye level; record to nearest 0.01 mL.

Use appropriate glassware: volumetric flask for precise dilution; graduated cylinder for approximate volumes; pipette for accurate delivery.

Titration endpoint is indicated by a color change of indicator; use phenolphthalein for strong base–weak acid, methyl orange for strong acid–weak base.

Wear goggles and gloves; know location of eyewash, safety shower, and fire extinguisher; never taste or directly smell chemicals.

Dilution calculation: M₁V₁ = M₂V₂; percent error = |experimental − accepted| / accepted × 100%.

Significant figures: addition/subtraction → least decimal places; multiplication/division → least significant figures.

Density = mass / volume; use it to convert between mass and volume in stoichiometric calculations.