UAE TLS Chemistry Test cheat sheet

Atomic number Z = number of protons; mass number A = protons + neutrons; isotopes differ only in neutron count.

Electron configuration fills orbitals by Aufbau order, obeying Pauli exclusion (max 2 electrons per orbital) and Hund's rule (parallel spins in degenerate orbitals).

Maximum electrons per shell level n is 2n²; subshell capacities: s=2, p=6, d=10, f=14.

Ionisation energy increases across a period and decreases down a group; successive ionisation energies rise sharply when a new shell is breached.

Atomic radius decreases across a period (higher effective nuclear charge) and increases down a group (extra shells).

Bohr energy levels: E = -13.6/n² eV for hydrogen; photon energy of a transition equals the level difference, E = hf = hc/λ.

Mass spectrometry gives isotope masses and relative abundances; relative atomic mass = Σ(isotope mass × fractional abundance).

Mole of any substance contains Avogadro's number, 6.022 × 10²³ particles.

Ionic bonds form between metals and non-metals by electron transfer; lattice energy magnitude rises with higher ionic charge and smaller ionic radius.

Covalent bonds share electron pairs; bond order from Lewis structures predicts shorter, stronger bonds at higher order.

VSEPR: electron domains repel to maximise separation; lone pairs repel more than bonding pairs, compressing bond angles.

Molecular shapes by domain count: 2 linear, 3 trigonal planar, 4 tetrahedral, 5 trigonal bipyramidal, 6 octahedral; lone pairs give bent, trigonal pyramidal, seesaw, or square planar variants.

Polarity requires polar bonds plus an asymmetric geometry; symmetric arrangements (CO₂, CCl₄) cancel bond dipoles.

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

Metallic bonding: delocalised electrons explain conductivity, malleability, and ductility of metals.

Sigma bonds are end-on overlaps; pi bonds are side-on p-orbital overlaps and prevent free rotation.

Moles n = mass / molar mass; for gases at STP (273 K, 1 atm) molar volume is 22.4 dm³ mol⁻¹.

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

Balanced equations give mole ratios; convert all quantities to moles before applying the ratio.

Limiting reagent is the reactant giving the smallest product yield from its mole ratio; excess reagent remains unreacted.

Percentage yield = (actual yield / theoretical yield) × 100; atom economy = (molar mass of desired product / total molar masses of reactants) × 100.

Concentration c = n / V; in mol dm⁻³ with V in dm³; 1 dm³ = 1000 cm³.

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

Gas law for non-standard conditions: PV = nRT with R = 8.314 J mol⁻¹ K⁻¹ and T in kelvin.

Kinetic theory assumes particles in constant random motion; pressure arises from collisions with container walls.

Average kinetic energy of gas particles is proportional to absolute temperature (KE ∝ T in kelvin).

Ideal gas equation PV = nRT; deviations occur at high pressure and low temperature where particle volume and attractions matter.

Boyle's law: P ∝ 1/V at constant T; Charles's law: V ∝ T at constant P; Gay-Lussac: P ∝ T at constant V.

Dalton's law: total pressure equals the sum of partial pressures of non-reacting gases.

Diffusion and effusion rates follow Graham's law: rate ∝ 1/√molar mass.

Phase changes: melting and boiling absorb energy (endothermic); freezing and condensation release energy (exothermic).

Boiling point rises with stronger intermolecular forces and higher external pressure; vapour pressure reaches external pressure at the boiling point.

Enthalpy change ΔH is negative for exothermic and positive for endothermic reactions, measured at constant pressure.

Calorimetry: q = mcΔT; for a reaction in solution, ΔH = -q / n (per mole of limiting reactant).

Hess's law: total ΔH is independent of route, so ΔH can be summed over steps.

Standard enthalpy of formation ΔHf° is zero for elements in their standard states; ΔH°reaction = ΣΔHf°(products) - ΣΔHf°(reactants).

Bond enthalpy method: ΔH = Σ(bonds broken) - Σ(bonds formed).

Gibbs free energy ΔG = ΔH - TΔS; a reaction is spontaneous when ΔG < 0.

Entropy S increases with more gas moles, higher temperature, and greater disorder; ΔS is positive for such changes.

Second law: total entropy of the universe increases for any spontaneous process.

Rate is the change in concentration per unit time; rate laws take the form rate = k[A]^m[B]^n with orders determined experimentally.

Collision theory: reaction requires sufficient energy (activation energy) and correct orientation of colliding particles.

Rate increases with higher concentration, higher temperature, larger surface area, and presence of a catalyst.

A catalyst lowers activation energy and speeds both forward and reverse reactions without changing ΔH or equilibrium position.

Arrhenius equation k = A·e^(-Ea/RT); a plot of ln k against 1/T gives slope -Ea/R.

Le Chatelier's principle: a system at equilibrium shifts to oppose imposed changes in concentration, pressure, or temperature.

Equilibrium constant Kc = [products]^coefficients / [reactants]^coefficients; pure solids and liquids are excluded.

For gases Kp = Kc(RT)^Δn, where Δn is the change in moles of gas.

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

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

Strong acids and bases dissociate completely; weak acids and bases establish equilibrium with Ka or Kb.

For a weak acid, Ka = [H⁺]²/[HA] approximately, so [H⁺] = √(Ka × [HA]).

Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25 °C.

Neutralisation: acid + base → salt + water; titration equivalence uses mole ratio from the balanced equation.

Salt hydrolysis: salts of strong acid and weak base give acidic solutions; strong base and weak acid give basic solutions.

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

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

Oxidation numbers: free elements 0; monatomic ions equal charge; oxygen usually -2; hydrogen usually +1; sum equals overall charge.

In a galvanic cell, oxidation occurs at the anode and reduction at the cathode; electrons flow anode to cathode.

Standard cell potential E°cell = E°cathode - E°anode; positive E°cell indicates a spontaneous reaction.

In electrolysis, the anode is positive and the cathode is negative; electrical energy drives a non-spontaneous reaction.

Faraday's law: moles of electrons = charge / 96500 C mol⁻¹; charge Q = I × t.

Oxidising agent is reduced; reducing agent is oxidised.

Standard hydrogen electrode has E° = 0.00 V and serves as the reference for standard electrode potentials.

Alkanes are saturated hydrocarbons, CₙH₂ₙ₊₂; they undergo substitution with halogens under UV light.

Alkenes contain C=C, formula CₙH₂ₙ; they undergo addition reactions with H₂, Br₂, HX, and water.

Alcohols contain -OH; primary alcohols oxidise to aldehydes then carboxylic acids; secondary to ketones; tertiary resist oxidation.

Aldehydes (CHO) reduce Tollens' reagent to a silver mirror; ketones do not.

Carboxylic acids (COOH) donate H⁺ and react with alcohols in acid-catalysed esterification, forming esters plus water.

Esters (COO) hydrolyse to alcohol and carboxylic acid; boiling with aqueous alkali gives the salt (saponification).

IUPAC naming: longest carbon chain gives the parent, numbered to give the lowest locants to functional groups.

Functional group tests: bromine water decolourises with alkenes; sodium carbonate gives CO₂ with carboxylic acids.

Flame tests: lithium crimson, sodium yellow, potassium lilac, calcium brick-red, copper blue-green.

Qualitative cation tests: NaOH gives coloured precipitates (Cu²⁺ blue, Fe²⁺ green, Fe³⁺ brown); NH₃(aq) distinguishes them.

Anion tests: carbonate gives CO₂ with acid; sulfate gives white BaSO₄ precipitate with Ba²⁺; chloride gives white AgCl precipitate with Ag⁺ that darkens in light.

Titration technique: rinse burette with the titrant, read the meniscus at eye level, use a white tile, and add indicator dropwise near the endpoint.

Filtration separates insoluble solid from liquid; crystallisation purifies a soluble solid; distillation separates miscible liquids by boiling point.

Chromatography separates mixtures by differential affinity; Rf = distance travelled by solute / distance travelled by solvent front.

Percentage error = (uncertainty / measured value) × 100; record to the precision of the instrument (e.g., burette ±0.05 cm³).

Safety: use fume cupboard for toxic vapours, wear goggles and gloves, and never add water to concentrated acid (add acid to water).