MATSEC SEC Chemistry cheat sheet

Mole concept: n = m/M; for gases at s.t.p. (0 °C, 1 atm) one mole occupies 22.4 dm³; molar volume at room temperature (25 °C) is about 24 dm³.

Avogadro constant N_A = 6.02 × 10²³ mol⁻¹; number of particles = n × N_A; concentration c = n/V with V in dm³ (mol dm⁻³).

Empirical formula: divide each percentage or mass by atomic mass, then divide by the smallest ratio; molecular formula = (empirical formula) × n where n = M_r/empirical mass.

Percentage yield = (actual mass/theoretical mass) × 100; percentage atom economy = (M_r of desired product/Σ M_r of all reactants) × 100.

Strong acids (HCl, HNO₃, H₂SO₄) ionise completely; weak acids (ethanoic, carbonic) ionise partially and show an equilibrium arrow.

Neutralisation titration: moles acid = moles base when using the stoichiometric ratio; n(H⁺) = c_acid × V_acid × basicity; indicator choice: phenolphthalein for strong–weak, methyl orange for weak base–strong acid.

pH scale: pH = −log₁₀[H⁺]; pH < 7 acidic, 7 neutral at 25 °C, > 7 alkaline; each unit of pH represents a tenfold change in [H⁺].

Oxidation is loss of electrons/increase in oxidation number; reduction is gain of electrons/decrease in oxidation number; oxidising agent is itself reduced.

Tests: CO₂ turns limewater milky; O₂ relights a glowing splint; H₂ gives a 'pop' with a lighted splint; Cl₂ bleaches damp litmus; water turns anhydrous copper(II) sulfate blue and cobalt chloride paper pink.

Ionic bonding: metal loses electrons to non-metal; giant lattice of ions; high melting point; conducts when molten or aqueous, not when solid.

Covalent bonding: shared electron pairs; simple molecular substances have weak intermolecular forces so low melting points and do not conduct.

Giant covalent structures: diamond (each C bonded to 4 others, very hard, non-conductor), graphite (layers of hexagonal sheets, delocalised electrons, lubricant and conductor).

Metallic bonding: lattice of positive ions in a sea of delocalised electrons; explains conductivity, malleability and ductility.

Electrolysis of molten lead(II) bromide: Pb²⁺ ions go to the cathode (reduction), Br⁻ ions to the anode (oxidation); electrode equations Pb²⁺ + 2e⁻ → Pb and 2Br⁻ → Br₂ + 2e⁻.

Group 1 metals: reactivity increases down the group (Li → Na → K) because the single outer electron is lost more easily as atomic radius increases and shielding rises.

Group 7 halogens: reactivity decreases down the group; displacement reactions show a more reactive halogen displacing a less reactive halide from its salt solution.

Group 0 noble gases: full outer shell, monatomic, unreactive; boiling point increases down the group as intermolecular forces strengthen.

Transition metals: variable oxidation states, coloured compounds, catalytic activity; e.g. Fe²⁺/Fe³⁺ interconversion and Cu²⁺ blue solutions.

Rate increases with higher concentration, higher temperature, larger surface area and added catalyst; temperature raises both collision frequency and the proportion of collisions exceeding activation energy.

Catalyst provides an alternative pathway of lower activation energy; it is not consumed overall and does not change the position of equilibrium.

Exothermic reaction: ΔH negative, products lower in energy than reactants; endothermic: ΔH positive, products higher in energy.

Bond enthalpy calculation: ΔH = Σ bonds broken (reactants) − Σ bonds formed (products).

Le Chatelier's principle: a system at equilibrium shifts to oppose a change in concentration, pressure or temperature; catalysts do not shift the position.

Haber process (N₂ + 3H₂ ⇌ 2NH₃, exothermic): moderate temperature ~450 °C, high pressure ~200 atm, iron catalyst; compromise conditions balance yield and rate.

Contact process (2SO₂ + O₂ ⇌ 2SO₃, exothermic): vanadium(V) oxide catalyst, ~450 °C, ~1–2 atm; SO₃ absorbed in concentrated sulfuric acid.

Collision theory: reaction requires particles to collide with sufficient energy (≥ activation energy) and correct orientation.

Homologous series: same general formula and functional group, similar chemical properties, gradual change in physical properties; alkanes CₙH₂ₙ₊₂, alkenes CₙH₂ₙ, alcohols CₙH₂ₙ₊₁OH.

Alkanes: saturated, undergo substitution with halogens in UV light (free-radical mechanism); burn to CO₂ and H₂O; incomplete combustion gives CO.

Alkenes: unsaturated, decolourise bromine water; addition reactions with H₂ (Ni catalyst), Br₂, HBr and steam (H₃PO₄ catalyst) to give an alcohol.

Alcohols: ethanol by fermentation of glucose (yeast, ~30 °C, anaerobic) or by hydration of ethene (steam, H₃PO₄, ~300 °C); oxidised by acidified potassium dichromate(VI) to ethanoic acid.

Carboxylic acids: weak acids, react with carbonates to give CO₂, with alcohols to form esters (concentrated H₂SO₄ catalyst); esters have characteristic sweet smells.

Polymers: addition polymerisation of ethene gives poly(ethene); condensation polymerisation of a diol and a dicarboxylic acid gives a polyester plus water.

Isomers have the same molecular formula but different structural formulae; e.g. butane and 2-methylpropane, or ethanol and methoxyethane.

Flame tests: Li crimson, Na yellow-orange, K lilac, Ca brick-red; cations identified by NaOH precipitate colours (Cu²⁺ blue, Fe²⁺ green, Fe³⁺ brown, Zn²⁺ and Al³⁺ white).

Anions: carbonate gives CO₂ with dilute acid; sulfate gives white precipitate with BaCl₂ then dilute HCl; halides give cream (Br⁻) or yellow (I⁻) precipitate with AgNO₃; nitrate gives brown gas with aluminium foil and NaOH.

Reactivity series (decreasing): K, Na, Ca, Mg, Al, Zn, Fe, Pb, (H), Cu, Ag, Au; metals above hydrogen displace H₂ from dilute acids.

Extraction: metals above carbon in the series are extracted by electrolysis (e.g. aluminium from molten alumina in cryolite); metals below carbon are reduced with carbon (e.g. iron in the blast furnace).

Blast furnace: haematite (Fe₂O₃) + coke + limestone; coke burns to CO, CO reduces iron oxide; limestone removes silica impurity as slag (calcium silicate).

Rusting requires both oxygen and water; prevented by painting, oiling, galvanising or sacrificial protection with a more reactive metal such as zinc or magnesium.

Alloys: brass (Cu/Zn), bronze (Cu/Sn), steel (Fe/C); alloying disrupts layers of atoms so the material is harder than the pure metal.

Crude oil is separated by fractional distillation on the basis of boiling point; cracking converts long-chain alkanes into shorter, more useful alkanes and alkenes.

Greenhouse gases (CO₂, CH₄, water vapour) absorb infrared radiation; consequences include rising sea levels and climate change.

Acid rain: SO₂ from burning fossil fuels and NOₓ from vehicle engines dissolve to form sulfurous/sulfuric and nitric acids, damaging buildings, soils and aquatic life.

Carbon cycle and combustion: complete combustion gives CO₂ and H₂O; incomplete combustion gives CO and soot, both harmful.

Titration technique: rinse burette with the solution it will contain, remove air bubble, read at eye level at the bottom of the meniscus; use a white tile to see the end-point colour change.

Concordant titres are within about 0.10 cm³ of each other; average only concordant values, then use the balanced equation to compute concentration.

Crystallisation: heat to concentrate the solution, allow to cool slowly to form crystals, filter, wash with a little cold solvent and dry between filter papers.

Distillation separates a solvent from a solution or two liquids with sufficiently different boiling points; a thermometer bulb is placed at the side-arm for the vapour temperature.

Chromatography: a mobile phase carries components over a stationary phase; R_f = distance moved by spot / distance moved by solvent front; each spot must be below the solvent front line.

Gas tests and purity: a pure substance melts and boils at a sharp, fixed temperature; impurities lower the melting point and widen the melting range.

Significant figures and units: give answers to the same number of significant figures as the least precise measurement; convert cm³ to dm³ by dividing by 1000.

Percentage composition by mass = (total A_r of element in formula / M_r of compound) × 100; use this for water of crystallisation and fertiliser calculations.

Atom notation: mass number A = protons + neutrons; atomic number Z = protons; isotopes have same Z but different A, identical chemical properties but different physical properties.

Electron configuration fills shells 2, 8, 8 (up to Ca); period number = number of occupied shells; group number = number of outer-shell electrons.

Ion formation: metals lose electrons to reach a noble-gas configuration (e.g. Mg → Mg²⁺ + 2e⁻); non-metals gain electrons (e.g. Cl + e⁻ → Cl⁻).

Relative atomic mass is the weighted mean of isotope masses relative to 1/12 the mass of a carbon-12 atom; calculate from percentage abundances.

Rutherford's alpha-particle experiment showed the atom has a small, dense, positively charged nucleus with mostly empty space; Bohr proposed quantised electron shells.

Ionic radius versus atomic radius: cations are smaller than their atoms, anions larger; this affects lattice strength and melting point trends.

Period 3 trends: Na, Mg, Al are metals; Si is a metalloid; P, S, Cl are non-metals; oxides change from basic (Na₂O, MgO) through amphoteric (Al₂O₃) to acidic (SO₂, SO₃, Cl₂O₇).

Formulae and valency: use the charges of common ions (e.g. Al³⁺ and O²⁻ give Al₂O₃; Ca²⁺ and NO₃⁻ give Ca(NO₃)₂) to write balanced formulae.