O Level / IGCSE Chemistry equations and key facts

States of matter

Solids, liquids and gases

• Solid: particles close together in a regular arrangement, vibrating about fixed positions
• Liquid: particles close together in an irregular arrangement, moving past one another
• Gas: particles far apart, arranged randomly, moving quickly in all directions
• Flat part of a heating or cooling curve = a change of state, with the temperature constant
• Below the melting point: solid. Between melting point and boiling point: liquid. Above the boiling point: gas
• Gas volume: higher temperature (at constant pressure) → larger volume; higher pressure (at constant temperature) → smaller volume
• Gas heated at constant volume: particles hit the walls harder and more often, so the pressure rises

Diffusion

• Diffusion: net movement of particles from higher concentration to lower concentration by random motion
• Speed of diffusion: gases fast, liquids slow, solids almost none
• Smaller M_r → faster diffusion; gases with the same M_r diffuse at the same rate
• Higher temperature → particles have more kinetic energy and move faster → faster diffusion
• Ammonia (M_r 17) and hydrogen chloride (M_r 36.5) in a tube: the white ring forms nearer the hydrogen chloride end
• NH₃(g) + HCl(g) → NH₄Cl(s), a white solid

Atoms, elements and compounds

Elements, compounds and mixtures

• Element: one type of atom only; cannot be broken down chemically
• Compound: two or more elements chemically bonded in fixed proportions, so it has a formula
• Mixture: two or more substances not chemically bonded, in any proportions
• A compound has different properties from the elements in it; substances in a mixture keep their own properties
• Mixtures are separated by physical methods; compounds are split only by chemical reactions
• Air is a mixture (mainly nitrogen and oxygen); water, carbon dioxide and sodium chloride are compounds
• A pure substance melts and boils at one fixed temperature; a mixture melts and boils over a range

Atomic structure and the Periodic Table

• Proton: relative charge +1, relative mass 1. Neutron: charge 0, mass 1. Electron: charge −1, mass 1/1840
• Proton number (atomic number) = number of protons in the nucleus
• Nucleon number (mass number) = number of protons + number of neutrons
• Number of neutrons = nucleon number − proton number
• In an atom, electrons = protons. Charge on an ion = number of protons − number of electrons
• Number of outer shell electrons = group number (Groups I to VII); number of occupied shells = period number
• Group VIII (noble gases) have a full outer shell

Isotopes

• Isotopes: atoms of the same element with the same number of protons but different numbers of neutrons
• Same number of electrons and same electronic configuration → same chemical properties
• In a symbol: top number = nucleon number, bottom number = proton number
• Neutrons = top number − bottom number
• Electrons in an ion = proton number − charge, e.g. Cl⁻ has 17 + 1 = 18 and Na⁺ has 11 − 1 = 10
• Relative atomic mass = (mass of isotope 1 × % abundance + mass of isotope 2 × % abundance) ÷ 100

Ions and ionic bonds

• Cation: positive ion, formed when an atom loses electrons. Anion: negative ion, formed when an atom gains electrons
• Ion charges: Group I 1+, Group II 2+, Group III 3+, Group V 3−, Group VI 2−, Group VII 1−
• Ionic bond: strong electrostatic attraction between oppositely charged ions
• Structure: giant lattice with a regular arrangement of alternating positive and negative ions
• High melting and boiling points: a lot of energy is needed to overcome the strong attractions between the ions
• Conducts electricity when molten or aqueous (ions free to move); does not conduct when solid (ions held in fixed positions)

Simple molecules and covalent bonds

• Covalent bond: a pair of electrons shared between two atoms, giving noble gas electronic configurations
• Single bonds only: H₂, Cl₂, HCl, H₂O (2 bonds), NH₃ (3 bonds), CH₄ (4 bonds), CH₃OH
• Double bonds: O₂ (one O=O), CO₂ (two C=O), C₂H₄ (one C=C)
• Triple bond: N₂ (three shared pairs)
• Number of bonds an atom forms = number of electrons it needs to fill its outer shell (H 1, O 2, N 3, C 4)
• Low melting and boiling points: only weak intermolecular forces have to be overcome
• Poor electrical conductivity: no ions and no free electrons to carry charge

Giant covalent structures

• Diamond: each carbon atom covalently bonded to 4 others; very hard; no free electrons, so it does not conduct
• Graphite: each carbon atom covalently bonded to 3 others in layers of hexagons; weak forces between the layers
• Graphite has one delocalised electron per carbon atom, so it conducts electricity
• Silicon(IV) oxide, SiO₂: each silicon atom bonded to 4 oxygen atoms; each oxygen atom bonded to 2 silicon atoms
• Diamond and silicon(IV) oxide are both hard, have very high melting points and do not conduct electricity
• All three have very high melting points: many strong covalent bonds must be broken
• Uses: diamond in cutting tools (hard); graphite as a lubricant (layers slide) and as an electrode (conducts)

Metallic bonding

• Metallic bonding: electrostatic attraction between positive ions in a giant metallic lattice and a sea of delocalised electrons
• The delocalised electrons are the outer shell electrons of the metal atoms (Na gives 1, Mg gives 2, Al gives 3)
• Good electrical conductivity: delocalised electrons are free to move through the lattice
• Malleable (can be hammered into sheets) and ductile (can be drawn into wires): layers of ions slide over each other
• The positive ions stay in the lattice and only vibrate; they do not carry the current in a solid metal
• A piece of metal is neutral overall: the charge of the ions equals the charge of the delocalised electrons

Stoichiometry

Formulae

• Molecular formula: the number and type of different atoms in one molecule
• Empirical formula: the simplest whole number ratio of the different atoms or ions, e.g. C₂H₆ has the empirical formula CH₃
• Ionic formula: total positive charge = total negative charge, e.g. Al³⁺ and O²⁻ give Al₂O₃
• Ions to remember: NH₄⁺, OH⁻, NO₃⁻, SO₄²⁻, CO₃²⁻
• State symbols: (s) solid, (l) liquid, (g) gas, (aq) aqueous, meaning dissolved in water
• Ionic equation: leave out the spectator ions, e.g. Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s)
• Neutralisation: H⁺(aq) + OH⁻(aq) → H₂O(l)

Relative masses of atoms and molecules

• Relative atomic mass, A_r: the average mass of the isotopes of an element compared to 1/12th of the mass of an atom of ¹²C
• Relative molecular mass, M_r: the sum of the relative atomic masses of all the atoms in the formula
• Relative formula mass: the same sum, used for ionic compounds
• A number outside a bracket multiplies everything inside it: (NH₄)₂ means 2 N and 8 H
• Water of crystallisation is added on: ·5H₂O adds 5 × 18 = 90
• A_r and M_r have no units

The mole and the Avogadro constant

• One mole contains 6.02 × 10²³ particles; number of particles = amount (mol) × 6.02 × 10²³
• Amount (mol) = mass (g) ÷ molar mass (g/mol)
• Volume of a gas at r.t.p. (dm³) = amount (mol) × 24
• Concentration (mol/dm³) = amount (mol) ÷ volume (dm³); to change cm³ to dm³, divide by 1000
• Concentration in g/dm³ = concentration in mol/dm³ × molar mass
• Percentage yield = actual mass of product ÷ theoretical mass of product × 100
• Percentage purity = mass of pure substance ÷ mass of impure sample × 100

Electrochemistry

Electrolysis

• Electrolysis = decomposition of an ionic compound, molten or in aqueous solution, by passing an electric current
• Cathode (negative): metals or hydrogen form, e.g. Pb²⁺ + 2e⁻ → Pb. Anode (positive): non-metals form, e.g. 2Br⁻ → Br₂ + 2e⁻
• Molten lead(II) bromide: grey molten lead at the cathode, red-brown bromine gas at the anode
• Concentrated aqueous sodium chloride: hydrogen at the cathode, chlorine at the anode. A dilute halide solution gives oxygen at the anode instead
• Dilute sulfuric acid: hydrogen at the cathode, oxygen at the anode, in a volume ratio of 2 : 1
• Aqueous copper(II) sulfate: with graphite electrodes, pink-brown copper on the cathode, oxygen bubbles at the anode and the blue colour fades. With copper electrodes, the anode loses mass, the cathode gains mass and the blue colour stays
• Electroplating: object = cathode, plating metal = anode, electrolyte contains ions of the plating metal. Fuel cell: 2H₂ + O₂ → 2H₂O, water is the only chemical product

Chemical energetics

Exothermic and endothermic reactions

• Exothermic: thermal energy given out, temperature of surroundings rises, ΔH is negative
• Endothermic: thermal energy taken in, temperature of surroundings falls, ΔH is positive
• Activation energy (E_a) = the minimum energy that colliding particles must have to react
• On a pathway diagram, E_a is the rise from the reactants level to the top of the curve; ΔH is the difference between the reactants level and the products level
• Bond breaking is endothermic; bond making is exothermic
• ΔH = energy taken in to break bonds in the reactants − energy given out when bonds form in the products (kJ/mol)

Chemical reactions

Physical and chemical changes

• Physical change: no new substance; examples are melting, boiling, freezing, condensing, sublimation, dissolving
• Chemical change: at least one new substance forms; examples are burning, rusting, electrolysis, fermentation, thermal decomposition
• Evidence of a chemical change: gas given off, precipitate formed, lasting colour change, temperature change
• Physical changes are usually easy to reverse; many chemical changes are hard to reverse
• Mass is conserved: a solid that gains mass on heating in air has combined with oxygen; one that loses mass has given off a gas

Rate of reaction

• average rate = change in amount ÷ time taken, e.g. in cm³/s or g/s
• Higher concentration (or higher gas pressure): more particles per unit volume, so collisions are more frequent
• Larger surface area (powder instead of lumps): more particles of the solid are exposed, so collisions are more frequent
• Higher temperature: particles have more kinetic energy, so they collide more often and more of the collisions have at least the activation energy
• Catalyst: increases the rate, lowers the activation energy and is unchanged at the end. Enzymes are biological catalysts
• The final amount of product depends only on the amounts of reactants, not on how fast the reaction goes

Reversible reactions and equilibrium

• Equilibrium: rate of forward reaction = rate of reverse reaction; concentrations no longer change
• Raising the temperature shifts the equilibrium in the endothermic direction; lowering it shifts in the exothermic direction
• Raising the pressure shifts the equilibrium to the side with fewer gas molecules
• Adding a reactant or removing a product shifts the equilibrium to the right. A catalyst does not change the position; equilibrium is just reached faster
• Haber process: N₂(g) + 3H₂(g) ⇌ 2NH₃(g); 450 °C, 200 atm (20 000 kPa), iron catalyst; nitrogen from air, hydrogen from methane
• Contact process: 2SO₂(g) + O₂(g) ⇌ 2SO₃(g); 450 °C, 2 atm (200 kPa), vanadium(V) oxide catalyst; SO₂ from burning sulfur or roasting sulfide ores, oxygen from air
• CuSO₄·5H₂O (blue) ⇌ CuSO₄ (white) + 5H₂O; CoCl₂·6H₂O (pink) ⇌ CoCl₂ (blue) + 6H₂O

Redox

• Oxidation Is Loss of electrons, Reduction Is Gain of electrons (OIL RIG)
• Oxidation number of an uncombined element is 0; of a simple ion it equals the charge, e.g. Fe³⁺ is +3
• Oxidation numbers add up to 0 in a compound and to the charge in an ion. In most compounds O is −2 and H is +1
• Oxidising agent: gains electrons, is reduced. Reducing agent: loses electrons, is oxidised
• Acidified aqueous potassium manganate(VII) turns from purple to colourless when it oxidises a reducing agent
• Aqueous potassium iodide turns from colourless to brown when an oxidising agent oxidises iodide ions to iodine

Acids, bases and salts

The characteristic properties of acids and bases

• acid + metal → salt + hydrogen (unreactive metals such as copper do not react)
• acid + base → salt + water; acid + carbonate → salt + water + carbon dioxide
• base + ammonium salt → salt + water + ammonia (on warming)
• Neutralisation: H⁺(aq) + OH⁻(aq) → H₂O(l)
• Litmus: red in acid, blue in alkali. Methyl orange: red in acid, yellow in alkali. Thymolphthalein: colourless in acid, blue in alkali
• Strong acids (hydrochloric, nitric, sulfuric): HCl(aq) → H⁺(aq) + Cl⁻(aq). Weak acids (carboxylic acids): CH₃COOH(aq) ⇌ H⁺(aq) + CH₃COO⁻(aq)
• Universal indicator: pH below 7 acidic (red for strong acid), pH 7 neutral (green), pH above 7 alkaline (purple for strong alkali)

Oxides

• Basic oxides (metals), e.g. CuO, CaO: react with acids → salt + water
• Acidic oxides (non-metals), e.g. SO₂, CO₂: react with alkalis → salt + water
• Amphoteric oxides, Al₂O₃ and ZnO: react with both acids and alkalis → salt + water
• CuO + H₂SO₄ → CuSO₄ + H₂O
• CO₂ + 2NaOH → Na₂CO₃ + H₂O
• Across a period, oxides change from basic (metals, left) to acidic (non-metals, right)

Preparation of salts

• All sodium, potassium and ammonium salts are soluble; all nitrates are soluble
• Chlorides are soluble except silver and lead; sulfates are soluble except barium, calcium and lead
• Carbonates are insoluble except sodium, potassium and ammonium; hydroxides are insoluble except sodium, potassium, ammonium and calcium (partially soluble)
• Acid + excess metal, insoluble base or insoluble carbonate: warm, add solid until no more reacts, filter, evaporate to the crystallisation point, cool, dry the crystals
• Acid + alkali: titrate with an indicator, repeat with the same volumes and no indicator, then crystallise
• Insoluble salt: mix two solutions, filter, wash the residue with distilled water, dry
• Hydrated = chemically combined with water; anhydrous = containing no water. Water of crystallisation is the water in hydrated crystals, e.g. CuSO₄·5H₂O, CoCl₂·6H₂O

The Periodic Table

Arrangement of elements

• Elements are arranged in order of increasing proton number (atomic number)
• Period number = number of occupied electron shells; group number = number of outer-shell electrons
• Same group → same number of outer electrons → similar chemical properties
• Ion charges: Group I 1+, Group II 2+, Group III 3+, Group V 3−, Group VI 2−, Group VII 1−; Group VIII elements do not usually form ions
• Across a period: metallic character decreases, non-metallic character increases
• Properties change steadily down a group, so a missing value can be estimated from the trend

Group I properties

• Group I metals: relatively soft, low density, low melting point, one outer electron, form 1+ ions
• Down the group: melting point decreases, density increases (in general), reactivity increases
• metal + water → metal hydroxide + hydrogen, e.g. 2Na + 2H₂O → 2NaOH + H₂
• Lithium fizzes steadily; sodium melts into a ball and darts about on the surface; potassium catches fire with a lilac flame
• The solution left is strongly alkaline (pH about 13 to 14)
• Elements lower down (rubidium, caesium) are softer, melt at lower temperatures and react more violently

Group VII properties

• At r.t.p.: chlorine is a pale yellow-green gas, bromine is a red-brown liquid, iodine is a grey-black solid
• Down the group: density increases, melting and boiling points increase, reactivity decreases
• Reactivity order: chlorine > bromine > iodine
• A halogen displaces any halogen below it from a solution of its halide; it cannot displace one above it
• Cl₂ + 2KBr → 2KCl + Br₂ (solution turns orange)
• Cl₂ + 2KI → 2KCl + I₂ and Br₂ + 2KI → 2KBr + I₂ (solution turns brown)
• Reactivity falls down the group because the larger atom attracts an extra electron less strongly

Transition elements

• Transition elements have high densities and high melting points.
• They have variable oxidation numbers, e.g. iron is +2 in FeCl₂ and +3 in FeCl₃.
• They form coloured compounds, e.g. hydrated copper(II) sulfate is blue and iron(II) compounds are pale green.
• They often act as catalysts, as elements and in compounds, e.g. iron in the Haber process and vanadium(V) oxide in the Contact process.
• Oxidation numbers in a compound add up to zero; oxygen is −2 and a Group I metal is +1.
• Ions of Group I and Group II metals are colourless in solution.

Noble gases

• Noble gases are in Group VIII and are unreactive, monatomic gases.
• Monatomic means made of single atoms: He, Ne, Ar (never He₂ or Ar₂).
• Electronic configurations: helium 2; neon 2,8; argon 2,8,8.
• A full outer shell is stable, so the atom does not lose, gain or share electrons.
• Other atoms form ions with a noble gas configuration, e.g. Na⁺ is 2,8 like neon.
• Uses depend on being unreactive: argon fills lamps; helium fills balloons (low density and does not burn).

Metals

Properties of metals

• Metals: good conductors of heat and electricity, malleable, ductile, usually high melting and boiling points.
• Non-metals: poor conductors, brittle if solid, usually low melting and boiling points. Graphite is the non-metal that conducts electricity.
• metal + dilute acid → salt + hydrogen, e.g. Mg + 2HCl → MgCl₂ + H₂
• metal + cold water → metal hydroxide + hydrogen, e.g. 2Na + 2H₂O → 2NaOH + H₂
• metal + steam → metal oxide + hydrogen, e.g. Mg + H₂O → MgO + H₂
• metal + oxygen → metal oxide, e.g. 2Mg + O₂ → 2MgO (white solid) and 2Cu + O₂ → 2CuO (black solid)

Uses of metals

• Aluminium in aircraft: low density.
• Aluminium in overhead electrical cables: low density and good electrical conductivity.
• Aluminium in food containers: resistance to corrosion (protective oxide layer).
• Copper in electrical wiring: good electrical conductivity and ductility.
• Low density means a small mass for a given volume (density = mass ÷ volume).
• Ductile means it can be drawn into wires; malleable means it can be hammered into sheets.

Alloys and their properties

• Alloy: a mixture of a metal with other elements.
• Brass = copper + zinc.
• Stainless steel = iron + other elements such as chromium, nickel and carbon.
• Alloys are harder and stronger because the different sized atoms stop the layers sliding over each other.
• Stainless steel is used for cutlery because it is hard and resists rusting.
• In a diagram: a pure metal has identical atoms in regular rows; an alloy has atoms of different sizes and the rows are distorted.

Reactivity series

• Order, most reactive first: potassium, sodium, calcium, magnesium, aluminium, carbon, zinc, iron, hydrogen, copper, silver, gold.
• Potassium, sodium and calcium react with cold water: metal + water → metal hydroxide + hydrogen.
• Magnesium reacts with steam: Mg + H₂O → MgO + H₂.
• Magnesium, zinc and iron react with dilute hydrochloric acid to give a salt and hydrogen; copper, silver and gold do not react.
• A more reactive metal displaces a less reactive one from its salt solution: Zn + CuSO₄ → ZnSO₄ + Cu.
• Ionic equation: Zn + Cu²⁺ → Zn²⁺ + Cu (zinc atoms lose electrons, copper ions gain them).
• Aluminium appears unreactive because of its protective oxide layer.

Corrosion of metals

• Rusting needs both oxygen and water.
• Rust is hydrated iron(III) oxide.
• Barrier methods (painting, greasing, coating with plastic) keep oxygen and water away from the iron.
• Sacrificial protection: a metal above iron in the reactivity series corrodes instead of the iron.
• The sacrificial metal loses electrons, e.g. Zn → Zn²⁺ + 2e⁻.
• Galvanising = coating iron with zinc; the zinc still protects the iron when the coating is scratched.

Extraction of metals

• Metals above carbon are extracted by electrolysis; metals below carbon by heating the oxide with carbon.
• Coke burns to give heat: C + O₂ → CO₂.
• Carbon dioxide is reduced by more coke: CO₂ + C → 2CO.
• Carbon monoxide reduces the ore: Fe₂O₃ + 3CO → 2Fe + 3CO₂.
• Limestone: CaCO₃ → CaO + CO₂, then CaO + SiO₂ → CaSiO₃ (slag).
• Aluminium, cathode: Al³⁺ + 3e⁻ → Al. Anode: 2O²⁻ → O₂ + 4e⁻.
• The hot carbon anodes burn away in the oxygen (C + O₂ → CO₂), so they are replaced regularly.

Chemistry of the environment

Water

• Anhydrous cobalt(II) chloride turns from blue to pink when water is added.
• Anhydrous copper(II) sulfate turns from white to blue when water is added.
• Pure water melts at 0 °C and boils at 100 °C; impurities lower the melting point and raise the boiling point.
• Beneficial: dissolved oxygen (for aquatic life) and some metal compounds (essential minerals).
• Harmful: toxic metal compounds, plastics, sewage (microbes cause disease), nitrates and phosphates (deoxygenation of water).
• Nitrates come from fertilisers; phosphates come from fertilisers and detergents.
• Treatment, in order: sedimentation and filtration (remove solids), carbon (removes tastes and odours), chlorination (kills microbes).

Fertilisers

• NPK = nitrogen, phosphorus, potassium.
• Ammonium salts and nitrates are used as fertilisers, e.g. NH₄NO₃, (NH₄)₂SO₄, KNO₃.
• Percentage by mass of an element = (number of atoms × A_r) ÷ M_r × 100.
• Ammonia + acid → ammonium salt, e.g. 2NH₃ + H₂SO₄ → (NH₄)₂SO₄.
• Ammonium salt + alkali → ammonia gas, so nitrogen is lost if lime is added together with an ammonium fertiliser.

Air quality and climate

• Clean, dry air: about 78% N₂, 21% O₂, the rest noble gases and CO₂.
• Complete combustion of carbon-containing fuels gives carbon dioxide; incomplete combustion gives carbon monoxide (toxic) and particulates (respiratory problems, cancer).
• Methane comes from decomposing vegetation and from digestion in animals; sulfur dioxide from burning fossil fuels that contain sulfur.
• Oxides of nitrogen come from car engines; they cause acid rain, photochemical smog and respiratory problems.
• Catalytic converter: 2CO + 2NO → 2CO₂ + N₂.
• Flue gas desulfurisation: calcium oxide removes sulfur dioxide, CaO + SO₂ → CaSO₃.
• Photosynthesis (needs light and chlorophyll): carbon dioxide + water → glucose + oxygen; 6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂.

Organic chemistry

Formulae, functional groups and terminology

• Alkanes: C_nH_2n+2. Alkenes: C_nH_2n.
• Alcohols: C_nH_2n+1OH. Carboxylic acids: C_nH_2n+1COOH.
• Functional groups: C=C in alkenes, –OH in alcohols, –COOH in carboxylic acids.
• The next member of a series has one more CH₂, so its M_r is 14 higher.
• Structural isomers of C₄H₁₀: CH₃CH₂CH₂CH₃ and CH₃CH(CH₃)CH₃.
• Structural isomers of C₄H₈: CH₃CH₂CH=CH₂ and CH₃CH=CHCH₃.
• Saturated: all carbon–carbon bonds are single. Unsaturated: one or more carbon–carbon bonds are not single, e.g. C=C.

Naming organic compounds

• Stems: meth- 1 carbon, eth- 2, prop- 3, but- 4, pent- 5
• Alkane, -ane: C_nH_2n+2, e.g. butane CH₃CH₂CH₂CH₃
• Alkene, -ene: C_nH_2n, contains C=C, e.g. but-1-ene CH₂=CHCH₂CH₃ and but-2-ene CH₃CH=CHCH₃
• Alcohol, -ol: contains –OH, e.g. propan-1-ol CH₃CH₂CH₂OH and propan-2-ol CH₃CH(OH)CH₃
• Carboxylic acid, -oic acid: contains –COOH, e.g. propanoic acid CH₃CH₂COOH (3 carbons in total)
• Ester name = alkyl part from the alcohol + -oate part from the acid, e.g. ethyl ethanoate CH₃COOCH₂CH₃
• In an ester formula written RCOOR', the part before COO (with that carbon) is from the acid; the part after it is from the alcohol

Fuels

• Fossil fuels: coal, natural gas, petroleum. Methane, CH₄, is the main constituent of natural gas
• Hydrocarbon: a compound containing hydrogen and carbon only
• Fractional distillation separates petroleum because the hydrocarbons have different boiling points
• Going up the column: shorter chains, lower boiling points, higher volatility, lower viscosity
• Order from top to bottom: refinery gas, gasoline (petrol), naphtha, kerosene (paraffin), diesel oil (gas oil), fuel oil, lubricating oil, bitumen
• Uses: refinery gas for heating and cooking; gasoline for cars; naphtha as a chemical feedstock; kerosene for jet fuel
• Uses: diesel oil for diesel engines; fuel oil for ships and home heating; lubricating oil for lubricants, waxes and polishes; bitumen for roads

Alkanes

• Alkanes: saturated hydrocarbons, single covalent bonds only, general formula C_nH_2n+2
• Generally unreactive, except for combustion and substitution by chlorine
• Complete combustion gives carbon dioxide and water: CH₄ + 2O₂ → CO₂ + 2H₂O
• Substitution: one atom or group of atoms is replaced by another atom or group of atoms
• CH₄ + Cl₂ → CH₃Cl + HCl (in ultraviolet light); CH₃Cl is chloromethane
• Ultraviolet light provides the activation energy, E_a; it is not a catalyst
• Substitution always gives two products; alkanes do not decolourise aqueous bromine

Alkenes

• Alkenes: unsaturated hydrocarbons with a C=C double bond, general formula C_nH_2n
• Cracking: large alkane → smaller alkane + alkene (or alkene + hydrogen), using a high temperature and a catalyst
• Test: an alkene turns aqueous bromine from orange to colourless; with an alkane it stays orange
• Addition reaction: only one product is formed
• With bromine: C₂H₄ + Br₂ → CH₂BrCH₂Br
• With hydrogen, nickel catalyst: C₂H₄ + H₂ → C₂H₆ (an alkane)
• With steam, acid catalyst: C₂H₄ + H₂O → C₂H₅OH (an alcohol)

Alcohols

• Fermentation: C₆H₁₂O₆ → 2C₂H₅OH + 2CO₂
• Fermentation conditions: yeast, 25–35 °C, aqueous glucose, no oxygen
• From ethene: C₂H₄ + H₂O → C₂H₅OH
• Ethene method conditions: 300 °C, 6000 kPa (60 atm), acid catalyst
• Fermentation: renewable, low energy, but slow, batch process, impure product
• Ethene method: fast, continuous, pure product, but non-renewable raw material and high energy use
• Complete combustion: C₂H₅OH + 3O₂ → 2CO₂ + 3H₂O. Uses of ethanol: solvent and fuel

Carboxylic acids

• Acid + metal → salt + hydrogen: 2CH₃COOH + Mg → (CH₃COO)₂Mg + H₂
• Acid + base → salt + water: CH₃COOH + NaOH → CH₃COONa + H₂O
• Acid + carbonate → salt + water + carbon dioxide: 2CH₃COOH + Na₂CO₃ → 2CH₃COONa + H₂O + CO₂
• Salt names: ethanoic acid gives ethanoates, e.g. sodium ethanoate, CH₃COONa
• Ethanol is oxidised to ethanoic acid by acidified aqueous potassium manganate(VII), which turns from purple to colourless
• Bacterial oxidation of ethanol in air also gives ethanoic acid (vinegar production)
• Acid + alcohol ⇌ ester + water, acid catalyst: CH₃COOH + C₂H₅OH ⇌ CH₃COOC₂H₅ + H₂O

Polymers

• Polymer: a large molecule built up from many smaller molecules called monomers
• Addition polymer: alkene monomers, C=C opens, one product only; repeat unit has the same atoms as the monomer with C–C in place of C=C
• Example: ethene, CH₂=CH₂, gives poly(ethene) with repeat unit –CH₂–CH₂–
• Condensation polymer: two different groups react and a small molecule such as water is formed as well
• Polyamide (nylon, proteins): amide linkage –CONH–, from –COOH and –NH₂ groups
• Polyester (PET): ester linkage –COO–, from –COOH and –OH groups; PET can be converted back into monomers and re-polymerised
• Plastics problems: fill landfill sites, build up in oceans, and give toxic gases when burnt

Experimental techniques and chemical analysis

Experimental design

• Volumetric pipette: one fixed volume (e.g. 25.0 cm³), very accurate. Burette: any volume, read to 0.05 cm³. Measuring cylinder: quick but less accurate
• Gas syringe: measures the volume of a gas. Balance: mass. Stop-watch: time. Thermometer: temperature
• Solvent: a substance that dissolves a solute. Solute: a substance that is dissolved in a solvent. Solution: solute(s) dissolved in a solvent
• Saturated solution: contains the maximum concentration of solute dissolved in the solvent at a specified temperature
• Residue: what remains after filtration, evaporation or distillation. Filtrate: the liquid or solution that passes through a filter
• Methyl orange: red in acid, yellow in alkali. Thymolphthalein: colourless in acid, blue in alkali
• End-point: the point where one drop from the burette makes the indicator just change colour

Chromatography

• R_f = distance travelled by substance ÷ distance travelled by solvent
• Both distances are measured from the start line. R_f has no unit and is never greater than 1
• One spot: pure substance. Two or more spots: mixture (impure)
• Spots at the same height (same R_f) in the same solvent are the same substance
• The start line is drawn in pencil because pencil does not dissolve in the solvent
• The start line must be above the solvent level, or the spots dissolve into the solvent in the beaker
• A substance that stays on the start line is insoluble in that solvent. A locating agent makes colourless spots visible

Separation and purification

• Filtration: separates an insoluble solid from a liquid (residue on the paper, filtrate passes through)
• Crystallisation: obtains a soluble solid from its solution. Heat to the crystallisation point, cool, filter, then dry the crystals
• Simple distillation: obtains the solvent from a solution, e.g. pure water from sea water
• Fractional distillation: separates miscible liquids with different boiling points. The liquid with the lowest boiling point distils first
• Using a suitable solvent: choose a solvent that dissolves only one of the solids, then filter
• Pure substance: sharp melting point and sharp boiling point
• Impurities lower the melting point, raise the boiling point, and make both happen over a range of temperature

Identification of ions and gases

• Carbonate: add dilute acid, effervescence, gas turns limewater milky. Sulfite: turns acidified aqueous potassium manganate(VII) from purple to colourless
• Halides: add dilute nitric acid, then aqueous silver nitrate. Chloride white, bromide cream, iodide yellow precipitate
• Sulfate: dilute nitric acid, then aqueous barium nitrate, white precipitate. Nitrate: warm with aluminium foil and aqueous sodium hydroxide, ammonia given off
• With aqueous sodium hydroxide: Cu²⁺ light blue, Fe²⁺ green, Fe³⁺ red-brown, Ca²⁺ white precipitate, all insoluble in excess. Al³⁺ and Zn²⁺ white, Cr³⁺ green, all soluble in excess. NH₄⁺ gives ammonia on warming
• With aqueous ammonia: Zn²⁺ white precipitate, soluble in excess. Al³⁺ white, insoluble. Cu²⁺ light blue, soluble in excess to a dark blue solution. Cr³⁺ green, Fe²⁺ green, Fe³⁺ red-brown, all insoluble. Ca²⁺ no precipitate or a very slight white one
• Gases: ammonia turns damp red litmus blue. Carbon dioxide turns limewater milky. Chlorine bleaches damp litmus. Hydrogen pops with a lighted splint. Oxygen relights a glowing splint. Sulfur dioxide turns acidified potassium manganate(VII) from purple to colourless
• Flame tests: Li⁺ red, Na⁺ yellow, K⁺ lilac, Ca²⁺ orange-red, Ba²⁺ light green, Cu²⁺ blue-green