JCE Chemistry cheat sheet

A pure substance has a fixed composition: an element contains one type of atom, a compound contains two or more elements chemically combined in fixed proportions.

A mixture contains two or more substances not chemically combined; components keep their own properties and can be separated by physical methods (filtration, distillation, chromatography).

Filtration separates an insoluble solid from a liquid; evaporation or crystallisation recovers a dissolved solid from solution.

Simple distillation separates a solvent from a solution; fractional distillation separates miscible liquids with different boiling points.

Chromatography separates dissolved substances by differing solubilities in a mobile phase and attractions to a stationary phase; the Rf value = distance moved by spot / distance moved by solvent front.

Changes of state are physical: melting (s→l), boiling (l→g), condensing (g→l), freezing (l→s) and sublimation (s→g); mass is conserved.

Kinetic particle theory: particles in solids vibrate about fixed positions, in liquids move randomly but stay close, in gases move rapidly and far apart.

Heating a substance raises the average kinetic energy of its particles; during a change of state, energy is used to overcome intermolecular forces, so temperature stays constant.

Proton number (atomic number, Z) = number of protons = number of electrons in a neutral atom; it determines the element.

Nucleon number (mass number, A) = protons + neutrons; neutrons = A − Z.

Isotopes are atoms of the same element with different numbers of neutrons; they have identical chemical properties but different masses.

Electron shell capacity: 2, 8, 8 (up to calcium); the outer-shell electrons determine chemical properties and group number.

Ions form when atoms lose or gain electrons: metals lose electrons to form positive ions (cations), non-metals gain electrons to form negative ions (anions).

Relative atomic mass, Ar = (sum of isotope mass × abundance) / total abundance, using percentage or fractional abundances.

The periodic table is arranged in order of increasing proton number; elements in the same group have the same number of outer electrons.

Metals are on the left and centre; non-metals on the right; the noble gases (Group 0) have full outer shells and are unreactive.

Ionic bonding: transfer of electrons from a metal to a non-metal, producing oppositely charged ions held in a giant ionic lattice by strong electrostatic attraction.

Ionic compounds have high melting points, conduct electricity when molten or dissolved (mobile ions), and are usually soluble in water.

Covalent bonding: sharing of electron pairs between non-metal atoms; a single bond shares one pair, a double bond shares two pairs.

Simple molecular substances have weak intermolecular forces, so low melting/boiling points; they do not conduct electricity because they have no free ions or electrons.

Giant covalent structures (diamond, silicon dioxide) have very high melting points; graphite conducts because each carbon has one delocalised electron.

Metallic bonding: positive ions in a lattice surrounded by a sea of delocalised electrons, giving conductivity, malleability and ductility.

Alloys are harder than pure metals because atoms of different sizes disrupt the layers, preventing them from sliding over each other.

Use the pattern: metal + non-metal → ionic; non-metal + non-metal → covalent; metal + metal → metallic.

Moles: n = m / M, where m is mass in grams and M is molar mass in g/mol.

Gas volume at room temperature and pressure: 1 mol occupies 24 dm³ (24 000 cm³).

Concentration: c = n / V, with V in dm³; convert cm³ to dm³ by dividing by 1000.

Avogadro constant = 6.02 × 10²³ particles per mole.

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

Balanced equations conserve atoms and charge; use state symbols (s), (l), (g), (aq) where required.

Empirical formula: divide moles of each element by the smallest value; molecular formula = empirical formula × whole-number multiplier.

Water of crystallisation or reacting masses: convert mass to moles, use the mole ratio from the balanced equation, then convert back to mass.

Acids release H⁺ ions in aqueous solution; bases accept H⁺; alkalis are soluble bases that release OH⁻ ions.

pH scale: pH < 7 acidic, pH = 7 neutral, pH > 7 alkaline; each unit represents a tenfold change in H⁺ concentration.

Neutralisation: acid + base → salt + water; acid + metal → salt + hydrogen; acid + carbonate → salt + water + carbon dioxide.

Salt name from the acid: hydrochloric → chloride, sulfuric → sulfate, nitric → nitrate, ethanoic → ethanoate, phosphoric → phosphate.

Preparation of a soluble salt: react acid with excess insoluble base or carbonate, filter off excess, then crystallise the filtrate.

Titration: use a pipette for a measured volume of alkali, a burette for the acid, and a suitable indicator (e.g. phenolphthalein or methyl orange).

Strong acids (HCl, HNO₃, H₂SO₄) ionise completely; weak acids (ethanoic, citric) ionise only partially.

Tests: H⁺ with litmus or pH paper; CO₂ turns limewater milky; hydrogen gives a squeaky pop with a lit splint.

Rate of reaction = amount of reactant used or product formed / time; measured by mass loss, gas volume, or time to a colour change.

Collision theory: particles must collide with energy at least equal to the activation energy and with correct orientation for a reaction to occur.

Increasing temperature increases rate by giving more particles energy above the activation energy and increasing collision frequency.

Increasing concentration or pressure increases rate by increasing the frequency of collisions.

Increasing surface area (smaller particles) increases rate by exposing more particles to collisions.

A catalyst increases rate by providing an alternative pathway with lower activation energy; it is not used up and does not change the yield.

Reversible reactions reach dynamic equilibrium when forward and reverse rates are equal and concentrations remain constant.

Le Chatelier: increasing temperature shifts equilibrium in the endothermic direction; increasing pressure shifts it toward fewer gas molecules.

Exothermic reactions release energy to the surroundings (temperature rises); ΔH is negative.

Endothermic reactions absorb energy from the surroundings (temperature falls); ΔH is positive.

Bond breaking requires energy (endothermic); bond forming releases energy (exothermic). ΔH = energy in (bonds broken) − energy out (bonds formed).

Activation energy is the minimum energy needed to start a reaction; it is shown on an energy profile diagram.

Electrolysis uses a direct current to decompose an ionic compound; the cathode is negative and attracts cations (reduction), the anode is positive and attracts anions (oxidation).

At the cathode, the less reactive metal or hydrogen is produced; at the anode, halides give halogen, otherwise oxygen from water.

In aqueous solutions, hydrogen is produced at the cathode unless the metal is less reactive than hydrogen (e.g. Cu, Ag).

A simple cell produces voltage from two different metals in an electrolyte; a larger difference in reactivity gives a larger voltage.

Group 1 alkali metals: soft, low density, react vigorously with water giving hydrogen and an alkaline hydroxide; reactivity increases down the group.

Group 7 halogens: diatomic non-metals; reactivity decreases down the group; a more reactive halogen displaces a less reactive one from its salt.

Group 0 noble gases: full outer shells, monatomic, very unreactive; boiling points increase down the group.

Across a period, metallic character decreases and non-metallic character increases; atomic radius decreases.

Down a group, atomic radius increases, ionisation energy decreases, and reactivity of metals increases while that of non-metals decreases.

Transition metals: hard, dense, high melting points, form coloured compounds, and act as catalysts (e.g. iron in the Haber process).

Halogen displacement example: chlorine + potassium bromide → potassium chloride + bromine, because chlorine is more reactive.

Group 1 oxides are basic; Group 7 elements form acidic compounds with hydrogen (e.g. HCl).

The reactivity series (most to least): potassium, sodium, calcium, magnesium, aluminium, (carbon), zinc, iron, (hydrogen), copper, silver, gold.

Metals above carbon in the series are extracted by electrolysis of their molten compounds; metals below carbon are reduced with carbon.

Iron is extracted in a blast furnace: haematite (Fe₂O₃) is reduced by carbon monoxide; limestone removes impurities as slag.

Aluminium is extracted by electrolysis of aluminium oxide dissolved in molten cryolite; cryolite lowers the melting point and saves energy.

Reaction with acid: metal + acid → salt + hydrogen; only metals above hydrogen react.

Displacement: a more reactive metal displaces a less reactive metal from its salt solution.

Rusting requires both oxygen and water; prevention includes painting, oiling, galvanising and sacrificial protection with zinc or magnesium.

Alloys such as steel (iron + carbon) and brass (copper + zinc) are harder and more corrosion-resistant than the pure metals.

Alkanes are saturated hydrocarbons with the general formula CₙH₂ₙ₊₂; they undergo complete combustion to CO₂ and H₂O.

Alkenes are unsaturated with the general formula CₙH₂ₙ; they contain a C=C double bond and decolourise bromine water.

Fractional distillation of crude oil separates fractions by boiling point; smaller molecules have lower boiling points and are more volatile.

Cracking breaks large alkanes into smaller, more useful alkanes and alkenes using heat and a catalyst.

Homologous series: same general formula, similar chemical properties, gradual change in physical properties, differ by CH₂.

Isomers have the same molecular formula but different structural formulae; they have different physical properties.

Polymers form by addition polymerisation of alkenes (e.g. ethene → poly(ethene)); the repeat unit has the same atoms as the monomer.

Combustion of hydrocarbons in limited oxygen produces carbon monoxide and soot; CO is toxic and binds to haemoglobin.

Accuracy: use a pipette for fixed volumes and a burette for variable volumes; read the bottom of the meniscus at eye level.

Precision: use apparatus with smaller scale divisions (e.g. a 0.1 cm³ burette rather than a 1 cm³ measuring cylinder).

Reliability: repeat measurements and calculate a mean, discarding anomalous results.

Control variables: keep temperature, concentration and volume constant when investigating the effect of one variable.

Hazard symbols: corrosive, toxic, flammable, oxidising, irritant, harmful, environmental hazard; wear goggles and use a fume cupboard for toxic gases.

Tests: flame tests identify metal ions (lithium crimson, sodium yellow, potassium lilac, calcium orange-red, copper blue-green).

Gas tests: hydrogen - squeaky pop; oxygen - relights a glowing splint; carbon dioxide - turns limewater milky; chlorine - bleaches damp litmus.

When diluting acid, add acid to water slowly, never water to acid, to avoid violent boiling and spitting.

Accuracy: use a pipette for fixed volumes and a burette for variable volumes; read the bottom of the meniscus at eye level.

Precision: use apparatus with smaller scale divisions (e.g. a 0.1 cm³ burette rather than a 1 cm³ measuring cylinder).

Reliability: repeat measurements and calculate a mean, discarding anomalous results.

Control variables: keep temperature, concentration and volume constant when investigating the effect of one variable.

Hazard symbols: corrosive, toxic, flammable, oxidising, irritant, harmful, environmental hazard; wear goggles and use a fume cupboard for toxic gases.

Tests: flame tests identify metal ions (lithium crimson, sodium yellow, potassium lilac, calcium orange-red, copper blue-green).

Gas tests: hydrogen - squeaky pop; oxygen - relights a glowing splint; carbon dioxide - turns limewater milky; chlorine - bleaches damp litmus.

When diluting acid, add acid to water slowly, never water to acid, to avoid violent boiling and spitting.