Myanmar Matriculation Chemistry cheat sheet

Mole relations: n = m/M; n = V/22.4 dm³ for gases at STP (0 °C, 1 atm); n = C × V(dm³) for solutions, so C = n/V.

Avogadro constant N_A = 6.022 × 10²³ mol⁻¹; number of particles = n × N_A. Molar volume at STP = 22.4 dm³ mol⁻¹ (22 400 cm³ mol⁻¹).

Gas laws: Boyle P₁V₁ = P₂V₂ (constant T); Charles V₁/T₁ = V₂/T₂ (constant P); combined PV/T = constant; ideal gas PV = nRT with R = 8.314 J K⁻¹ mol⁻¹ = 0.0821 dm³ atm K⁻¹ mol⁻¹.

Percentage yield = (actual mass ÷ theoretical mass) × 100; percentage purity = (mass of pure substance ÷ mass of sample) × 100.

Empirical formula from percentage composition: divide each % by atomic mass, then divide all by the smallest quotient; molecular formula = (empirical formula)ₙ where n = M_r ÷ empirical formula mass.

Limiting reactant: convert all reactant amounts to moles, divide each by its stoichiometric coefficient; the smallest quotient identifies the limiting reactant and fixes the maximum product.

Concentration conversions: molarity = (mass concentration in g dm⁻³) ÷ M_r; dilution obeys C₁V₁ = C₂V₂ (moles of solute unchanged).

Titration calculation route: moles of standard in burette reading → mole ratio from balanced equation → moles of analyte in pipetted volume → scale to required volume → mass or concentration.

Atomic structure: proton number Z fixes the element; nucleon number A = Z + neutrons; isotopes share Z but differ in neutron number, so chemical properties are nearly identical while relative atomic mass is a weighted mean of isotope masses.

Electronic configuration follows the Aufbau order 1s 2s 2p 3s 3p 4s 3d 4p; the octet rule governs ion formation - metals lose electrons to reach the preceding noble-gas configuration, non-metals gain electrons.

Periodic trends across a period (left to right): nuclear charge increases, atomic radius decreases, ionisation energy increases, electronegativity increases, metallic character decreases. Down a group the reverse occurs because shells are added.

Bonding decision rule: metal + non-metal gives ionic (transfer of electrons, high melting point, conducts when molten or aqueous); non-metal + non-metal gives covalent (shared pairs, usually low melting point); metal + metal gives metallic bonding.

Intermolecular forces ranked by strength: hydrogen bonding (H bonded to N, O or F) > permanent dipole–dipole > instantaneous dipole–induced dipole (van der Waals). Stronger forces raise boiling point.

Oxidation number rules: free element = 0; monatomic ion = its charge; oxygen usually −2 (peroxide −1); hydrogen usually +1 (metal hydride −1); sum of oxidation numbers in a neutral compound = 0, in a polyatomic ion = ion charge.

Redox identification: oxidation is loss of electrons (oxidation number rises); reduction is gain of electrons (oxidation number falls); the oxidising agent is itself reduced, the reducing agent is itself oxidised.

Electrolysis: at the cathode (negative) cations are reduced; at the anode (positive) anions are oxidised. Mass deposited follows Faraday's law, m = (Q × M)/(z × F) with F = 96 500 C mol⁻¹ and z the electron count.

Reaction-rate factors: increasing concentration or pressure, raising temperature, increasing surface area of a solid, and adding a catalyst all raise collision frequency or the fraction of successful collisions; a catalyst lowers activation energy without changing ΔH.

Acid–base definitions: Brønsted–Lowry acid = proton donor, base = proton acceptor; strong acids (HCl, HNO₃, H₂SO₄) ionise completely, weak acids (CH₃COOH, H₂CO₃) ionise partially and set up an equilibrium.

pH = −log₁₀[H⁺]; [H⁺] = 10^(−pH); at 25 °C pH + pOH = 14 and [H⁺][OH⁻] = K_w = 1 × 10⁻¹⁴. For a strong monoprotic acid, pH = −log₁₀(acid concentration).

Neutralisation: acid + base → salt + water; acid + carbonate → salt + water + CO₂; acid + metal → salt + hydrogen. Use the mole ratio from the balanced equation, not the volumes alone.

Salt preparation choice: soluble salts by titration (sodium, potassium, ammonium) or by excess insoluble base/carbonate plus filtration; insoluble salts by precipitation (double decomposition) followed by filtration and washing.

Solubility rules: all nitrates and all sodium, potassium and ammonium salts are soluble; chlorides are soluble except AgCl and PbCl₂; sulfates are soluble except BaSO₄, PbSO₄ and CaSO₄ (slightly soluble); carbonates are insoluble except those of Na, K and NH₄.

Tests: CO₂ turns limewater milky; O₂ relights a glowing splint; H₂ gives a 'pop' with a lighted splint; NH₃ turns damp red litmus blue; Cl₂ bleaches damp litmus; SO₂ decolourises acidified potassium dichromate(VI).

Flame tests: lithium crimson, sodium yellow, potassium lilac, calcium brick-red, copper blue-green, barium apple-green. Cation confirmation uses NaOH: Cu²⁺ blue precipitate, Fe²⁺ green, Fe³⁺ reddish-brown, Zn²⁺ and Al³⁺ white precipitates soluble in excess NaOH.

Organic homologous series: alkanes CₙH₂ₙ₊₂ (substitution, saturated); alkenes CₙH₂ₙ (addition, decolourise bromine water); alcohols CₙH₂ₙ₊₁OH (oxidation to aldehyde then acid); carboxylic acids (esterify with alcohols, release CO₂ from carbonates).

Esterification: carboxylic acid + alcohol ⇌ ester + water, with concentrated H₂SO₄ as catalyst; the reaction is reversible and the ester is identified by its sweet smell. Addition polymerisation of ethene gives poly(ethene); condensation polymerisation releases a small molecule such as water.

Enthalpy conventions: ΔH is negative for exothermic reactions (heat released, products lower in energy) and positive for endothermic reactions; bond breaking absorbs energy, bond forming releases it.

Hess's law: the total enthalpy change is independent of route, so ΔH = ΣΔH_f(products) − ΣΔH_f(reactants) or, using bond energies, ΔH = Σ(bonds broken) − Σ(bonds formed).

Calorimetry: q = mcΔT, with m in grams, c = 4.2 J g⁻¹ K⁻¹ for water; ΔH (kJ mol⁻¹) = −q ÷ (1000 × moles reacting). Heat loss to surroundings makes measured values smaller in magnitude than accepted ones.

Le Chatelier's principle: a system at equilibrium shifts to oppose any change. Raising pressure shifts toward fewer gas moles; raising temperature shifts in the endothermic direction; a catalyst changes only the rate of approach, not the position.

Equilibrium constant K_c = [products]^coefficients ÷ [reactants]^coefficients for the balanced equation; K_c depends only on temperature. A large K_c means products predominate.

Rates versus equilibrium: rate measures how fast concentrations change; equilibrium describes the final ratio when forward and reverse rates are equal and concentrations stay constant.

Electrochemical cells: in a galvanic cell oxidation occurs at the anode (negative terminal) and reduction at the cathode (positive terminal); electrons flow through the external circuit from anode to cathode.

E° values decide spontaneity: the half-cell with the more positive (less negative) E° is reduced; E°cell = E°cathode − E°anode, and a positive E°cell means the reaction is feasible.

Qualitative analysis of anions: dilute acid gives CO₂ from carbonate, SO₂ from sulfite; BaCl₂ followed by acid gives a white precipitate for sulfate; AgNO₃ gives cream AgBr for bromide and yellow AgI for iodide.