Elements, compounds and mixtures

• A compound has different properties from the elements it contains and a fixed composition; it can only be separated by a chemical reaction.
• A mixture keeps the properties of its components and can be separated by physical methods (filtration, distillation, chromatography).

Atomic structure

An atom has a tiny central nucleus containing protons and neutrons, surrounded by electrons in shells.

• Proton number (atomic number, Z): the number of protons in the nucleus of an atom.
• Mass number (nucleon number, A): the total number of protons and neutrons in the nucleus of an atom.
• Number of neutrons = mass number − proton number. Atoms are neutral, so number of electrons = number of protons.
• Notation: ²³₁₁Na has 11 protons, 11 electrons and 12 neutrons.

Electronic configuration (first 20 elements)

• Shells fill in order: first shell holds up to 2, second up to 8, third up to 8 (for elements 1–20).
• Examples: C 2,4; O 2,6; Na 2,8,1; Cl 2,8,7; Ca 2,8,8,2.
• The number of occupied shells = the period number.
• The number of outer-shell electrons = the group number (Groups I–VII).
• Noble gases (Group VIII) have a full outer shell, which makes them unreactive. Helium has 2 outer electrons; the others have 8.

Isotopes

Isotopes are different atoms of the same element that have the same number of protons but different numbers of neutrons.

• Example: ³⁵₁₇Cl (18 neutrons) and ³⁷₁₇Cl (20 neutrons).
• Isotopes of an element have the same chemical properties.
• Supplement This is because they have the same number of electrons in their outer shell (the same electronic configuration), and chemical reactions involve electrons.

Supplement Calculating relative atomic mass

Relative atomic mass (A_r) is the weighted average mass of the isotopes of an element, using their relative abundances.

Ions and ionic bonding

• An ion is a charged particle formed when an atom gains or loses electrons.
• Cations are positive ions, formed when atoms (usually metals) lose electrons. Anions are negative ions, formed when atoms (usually non-metals) gain electrons.
• An ionic bond is a strong electrostatic attraction between oppositely charged ions.
• Ionic bonds form between metals and non-metals. Electrons are transferred from the metal to the non-metal so both gain a full outer shell (noble gas configuration).

• Core: dot-and-cross diagrams for Group I metals with Group VII non-metals (e.g. NaCl, KBr).
• Supplement Dot-and-cross diagrams for other metal/non-metal combinations, e.g. MgO (Mg²⁺ and O²⁻, two electrons transferred) and CaCl₂ (one Ca²⁺ and two Cl⁻).
• Supplement An ionic compound has a giant lattice structure: a regular arrangement of alternating positive and negative ions, e.g. sodium chloride.

Properties of ionic compounds

• High melting and boiling points.
• Good electrical conductivity when molten or in aqueous solution; poor when solid.
• Supplement Explanation: many strong electrostatic attractions between oppositely charged ions in the giant lattice need a lot of energy to overcome. When solid the ions are held in fixed positions; when molten or dissolved the ions are free to move and carry charge.

Simple molecules and covalent bonds

• A covalent bond forms when a pair of electrons is shared between two atoms, giving each a noble gas electronic configuration. Covalent bonding occurs between non-metal atoms.
• Core dot-and-cross diagrams: H₂, Cl₂, H₂O, CH₄, NH₃, HCl.
• Supplement Also: CH₃OH, C₂H₄ (one C=C double bond), O₂ (double bond), CO₂ (two C=O double bonds), N₂ (triple bond).

Properties of simple molecular substances

• Low melting and boiling points (many are gases or liquids at room temperature).
• Poor electrical conductivity: there are no ions or free electrons to carry charge.
• Supplement Explanation: the covalent bonds within molecules are strong, but the intermolecular forces (attractive forces between molecules) are weak, so little energy is needed to separate the molecules. Larger molecules have stronger intermolecular forces and higher boiling points.

Giant covalent structures

• Both have very high melting points because many strong covalent bonds must be broken.
• Supplement Silicon(IV) oxide, SiO₂, has a giant covalent structure similar to diamond: each Si is bonded to 4 O atoms and each O to 2 Si atoms. It is hard, has a very high melting point and does not conduct electricity.

Supplement Metallic bonding

• Metallic bonding is the electrostatic attraction between a lattice of positive metal ions and a “sea” of delocalised electrons.
• Good electrical conductivity: the delocalised electrons are free to move through the structure and carry charge.
• Malleable and ductile: the layers of positive ions can slide over each other without breaking the metallic bonding, because the delocalised electrons still hold them together.

Formulae

• The molecular formula is the number and type of different atoms in one molecule, e.g. C₂H₆ for ethane.
• Supplement The empirical formula is the simplest whole-number ratio of the different atoms or ions in a compound, e.g. CH₃ for ethane.
• You should be able to deduce a formula from a diagram or model by counting each type of atom.
• Ionic formulae come from ion charges: the total positive charge must equal the total negative charge.

Examples: Al³⁺ + O²⁻ gives Al₂O₃; Ca²⁺ + NO₃⁻ gives Ca(NO₃)₂. Roman numerals give the charge: iron(III) chloride is FeCl₃.

Equations

• Supplement Ionic equations show only the ions that change; spectator ions are left out. Neutralisation: H⁺(aq) + OH⁻(aq) → H₂O(l). Precipitation: Ag⁺(aq) + Cl⁻(aq) → AgCl(s).
• You may be asked to deduce a balanced equation from information given in words, including state symbols.

Relative masses

• 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 carbon-12.
• Relative molecular mass, M_r: the sum of the relative atomic masses of all the atoms in a molecule. The same term (relative formula mass) is used for ionic compounds.
• Example: M_r of CaCO₃ = 40 + 12 + (3 × 16) = 100. Relative masses have no units.
• Calculate reacting masses in simple proportions, e.g. if 100 g CaCO₃ gives 56 g CaO, then 25 g gives 14 g. Mass is conserved in reactions.

Concentration

• Concentration can be measured in g/dm³ or mol/dm³.
• 1 dm³ = 1000 cm³; divide cm³ by 1000 to convert to dm³.

Supplement The mole and the Avogadro constant

• The mole (mol) is the unit of amount of substance. One mole contains 6.02 × 10²³ particles (the Avogadro constant).
• The mass of one mole of a substance in grams equals its A_r or M_r (the molar mass, in g/mol).

Reacting masses and gas volumes

Limiting reactant

The limiting reactant is the one that is completely used up; it controls the amount of product. The other reactant is in excess.

Titration calculations

Empirical and molecular formulae

Percentage yield, composition and purity

• Percentage yield = (actual yield ÷ theoretical yield) × 100
• Percentage composition by mass of an element = (total A_r of that element in the formula ÷ M_r) × 100
• Percentage purity = (mass of pure substance ÷ mass of impure sample) × 100

Electrolysis: key terms

Electrolysis is the decomposition of an ionic compound, when molten or in aqueous solution, by the passage of an electric current.

Supplement Transfer of charge during electrolysis

Predicting products

• Molten binary compound (two elements, inert electrodes): the metal forms at the cathode and the non-metal at the anode.
• In general, metals or hydrogen form at the cathode and non-metals (other than hydrogen) form at the anode.
• Aqueous solutions also contain H⁺ and OH⁻ ions from water.
• At the cathode: hydrogen forms if the metal is more reactive than hydrogen; the metal forms if it is less reactive (e.g. copper, silver).
• At the anode: with a concentrated halide solution the halogen forms; with a dilute halide solution, or a sulfate, oxygen forms.

Required electrolytes

Gas tests

• Hydrogen: a lighted splint gives a “pop”.
• Oxygen: relights a glowing splint.
• Chlorine: bleaches damp litmus paper.

Supplement Ionic half-equations

Reduction happens at the cathode (electrons gained); oxidation happens at the anode (electrons lost). Balance atoms and charges.

Electroplating

Electroplating uses electrolysis to coat one metal with a thin layer of another metal.

• Uses: to improve appearance (e.g. silver or gold plating on jewellery) and to resist corrosion (e.g. chromium on steel taps, tin on steel cans).
• Supplement Copper is purified the same way: impure copper as the anode, pure copper as the cathode, copper(II) sulfate solution as electrolyte. Impurities fall below the anode.

Hydrogen–oxygen fuel cells

• A fuel cell uses hydrogen and oxygen to produce electricity, with water as the only chemical product.
• Overall reaction: 2H₂(g) + O₂(g) → 2H₂O(l)

Supplement Fuel cells compared with petrol engines

Exothermic and endothermic reactions

Energy is conserved in every reaction. What changes is where it is stored: some energy is held in the chemical bonds (the system), and some is thermal energy in everything around the reaction (the surroundings — the solvent, test tube, air and thermometer).

• Exothermic reaction: a reaction that transfers thermal energy to the surroundings, so the temperature of the surroundings increases.
• Endothermic reaction: a reaction that takes in thermal energy from the surroundings, so the temperature of the surroundings decreases.

Measuring the temperature change in the lab

• Carry out the reaction in an insulated container (a polystyrene cup with a lid) to reduce thermal energy loss to the air.
• Record the starting temperature, mix, stir, and record the highest (or lowest) temperature reached.
• A rise in temperature shows an exothermic reaction; a fall shows an endothermic reaction.

Reaction pathway diagrams

A reaction pathway diagram (energy level diagram) plots the energy of the chemicals on the vertical axis against the progress of the reaction on the horizontal axis.

Supplement Enthalpy change and activation energy

• Enthalpy change, ΔH: the transfer of thermal energy during a reaction. Units: kJ/mol.
• For an exothermic reaction ΔH is negative (the system loses energy).
• For an endothermic reaction ΔH is positive (the system gains energy).
• Activation energy, E_a: the minimum energy that colliding particles must have in order to react.

Drawing and labelling a pathway diagram (Supplement)

Supplement Bond breaking and bond making

• Bond breaking is endothermic — energy must be supplied to pull atoms apart.
• Bond making is exothermic — energy is released when new bonds form.
• If more energy is released making bonds than is taken in breaking bonds, the reaction is exothermic overall.
• If more energy is taken in breaking bonds than is released making bonds, the reaction is endothermic overall.

Calculating ΔH from bond energies

ΔH = total energy needed to break bonds in reactants − total energy released making bonds in products.

Physical and chemical changes

• Physical change: no new substance is made and the change is usually easy to reverse, e.g. melting ice, dissolving salt, boiling water.
• Chemical change: one or more new substances are formed; often accompanied by an energy change, a colour change, fizzing or a precipitate, and usually hard to reverse, e.g. burning magnesium, rusting.

Rate of reaction

Rate measures how quickly reactants are used up or products are formed. The five factors on the syllabus:

• Catalyst: a substance that increases the rate of a reaction and is unchanged in mass and chemical composition at the end of the reaction.
• Enzymes are proteins that act as biological catalysts. They work best at an optimum temperature and pH; high temperatures denature them.
• Fine powders (e.g. flour, coal dust) suspended in air can burn so fast that they cause explosions — a hazard in mills and mines.

Supplement Collision theory in detail

• Particles react only when they collide with energy equal to or greater than the activation energy. These are successful collisions.
• Rate depends on the frequency of successful collisions (number per second), not just the number of collisions.
• Temperature has two effects: collisions become more frequent, and a larger proportion of particles have energy ≥ E_a. The second effect is the more important one.
• A catalyst lowers the activation energy, so a larger proportion of collisions are successful at the same temperature.

Practical methods for following rate

Interpreting rate graphs

• On a graph of volume of gas (or loss of mass) against time, the gradient shows the rate.
• The curve is steepest at the start (highest concentration of reactants) and levels off when one reactant is used up.
• A faster reaction gives a steeper initial curve. If the amounts of reactants are the same, it levels off at the same final volume, just sooner.
• Halving the amount of the limiting reactant halves the final volume of gas.
• Supplement Calculate average rate = change in volume ÷ time taken; find the rate at a point by drawing a tangent and working out its gradient.

Reversible reactions

• Some reactions can go in both directions; this is shown by the symbol ⇌.
• Heating blue hydrated copper(II) sulfate drives off water and leaves white anhydrous copper(II) sulfate. Adding water turns it blue again:
  CuSO₄·5H₂O(s) ⇌ CuSO₄(s) + 5H₂O(l)
• Pink hydrated cobalt(II) chloride becomes blue when heated (anhydrous) and turns pink again when water is added:
  CoCl₂·6H₂O(s) ⇌ CoCl₂(s) + 6H₂O(l)
• These colour changes are used as chemical tests for water (anhydrous CuSO₄ white → blue; anhydrous CoCl₂ blue → pink). To test whether water is pure, check that it boils at exactly 100 °C and freezes at 0 °C.

Supplement Equilibrium

• A reversible reaction in a closed system reaches equilibrium when the rate of the forward reaction equals the rate of the reverse reaction and the concentrations of reactants and products no longer change.
• Equilibrium is dynamic: both reactions keep happening.

Supplement The Haber process (ammonia)

N₂(g) + 3H₂(g) ⇌ 2NH₃(g)   ΔH negative (forward reaction exothermic)

Supplement The Contact process (sulfuric acid)

2SO₂(g) + O₂(g) ⇌ 2SO₃(g)   ΔH negative

Supplement Explaining the chosen conditions

• Temperature: a low temperature would give a higher equilibrium yield (forward reaction is exothermic) but the rate would be too slow. 450 °C is a compromise between rate and yield.
• Pressure (Haber): high pressure moves equilibrium to the right (4 moles of gas → 2) and increases rate, but very high pressures are expensive and dangerous; 20 000 kPa balances yield against cost and safety.
• Pressure (Contact): the equilibrium already lies well to the right, so only a low pressure is needed; higher pressure is not worth the extra cost.
• Catalyst: increases the rate so equilibrium is reached faster; it does not change the yield.

Redox

• Oxidation is gain of oxygen; reduction is loss of oxygen.
• Redox reactions involve simultaneous oxidation and reduction, e.g. in CuO + H₂ → Cu + H₂O, copper(II) oxide is reduced and hydrogen is oxidised.
• Roman numerals show the oxidation number of an element in a compound, e.g. iron(II) sulfate contains Fe²⁺, iron(III) chloride contains Fe³⁺.

Supplement Electrons and oxidation numbers

• Oxidation is loss of electrons and an increase in oxidation number. Reduction is gain of electrons and a decrease in oxidation number.
• Rules: an uncombined element has oxidation number 0; a simple ion has oxidation number equal to its charge; in a compound the oxidation numbers add up to 0; in an ion they add up to the charge. Oxygen is usually −2 and hydrogen usually +1.
• An oxidising agent oxidises another substance and is itself reduced. A reducing agent reduces another substance and is itself oxidised.
• Example: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)
  Zn → Zn²⁺ + 2e⁻ (oxidation; Zn is the reducing agent)
  Cu²⁺ + 2e⁻ → Cu (reduction; Cu²⁺ is the oxidising agent)
• Tests using colour changes: acidified potassium manganate(VII) is an oxidising agent and turns from purple to colourless when it reacts with a reducing agent. Aqueous potassium iodide is a reducing agent and turns from colourless to brown (iodine forms) when it reacts with an oxidising agent.

Properties of acids

• Acids contain hydrogen ions, H⁺(aq), when dissolved in water.
• Acids have a pH below 7, turn blue litmus red, turn methyl orange red and keep thymolphthalein colourless.

Reactions of acids (learn the general word equations)

Hydrochloric acid gives chlorides, sulfuric acid gives sulfates, nitric acid gives nitrates.

Properties of bases and alkalis

• Bases are metal oxides and hydroxides; they neutralise acids to form a salt and water.
• Alkalis are soluble bases. They contain hydroxide ions, OH⁻(aq), when dissolved in water.
• Alkalis have a pH above 7, turn red litmus blue, turn methyl orange yellow and turn thymolphthalein blue.
• Bases react with ammonium salts when warmed, releasing ammonia gas:
  NH₄Cl(s) + NaOH(aq) → NaCl(aq) + H₂O(l) + NH₃(g)
  (Test for NH₄⁺: ammonia turns damp red litmus paper blue.)

The pH scale and neutralisation

• pH runs from 1 to 14. Below 7 is acidic, 7 is neutral, above 7 is alkaline. Measure it with universal indicator (compare to a colour chart) or a pH meter.
• Neutralisation is the reaction between an acid and an alkali: H⁺(aq) + OH⁻(aq) → H₂O(l)

Supplement Proton transfer and strength of acids

• An acid is a proton (H⁺) donor; a base is a proton acceptor.
• A strong acid is completely dissociated in aqueous solution, e.g. hydrochloric acid: HCl(aq) → H⁺(aq) + Cl⁻(aq)
• A weak acid is only partially dissociated in aqueous solution, e.g. ethanoic acid: CH₃COOH(aq) ⇌ H⁺(aq) + CH₃COO⁻(aq)
• At the same concentration, a strong acid has a higher H⁺ concentration, so a lower pH and reacts faster (e.g. with magnesium).
• The lower the pH, the higher the H⁺ concentration. Each drop of 1 pH unit means a tenfold increase in H⁺ concentration.

Oxides

• Acidic oxides are usually oxides of non-metals, e.g. CO₂, SO₂. They react with bases to form a salt and water.
• Basic oxides are usually oxides of metals, e.g. CuO, CaO. They react with acids to form a salt and water.
• Supplement Amphoteric oxides react with both acids and bases to produce a salt and water, e.g. aluminium oxide, Al₂O₃, and zinc oxide, ZnO.

Preparing salts

A salt is formed when the hydrogen of an acid is replaced by a metal (or ammonium) ion. First decide whether the salt is soluble.

Method A: soluble salt from an insoluble solid (excess metal, base or carbonate)

Example: making copper(II) sulfate crystals from excess copper(II) oxide and dilute sulfuric acid.

Method B: soluble salt by titration (acid + alkali)

• Used for sodium, potassium and ammonium salts, where both reactants are solutions so an excess cannot be filtered off.
• Pipette a known volume of alkali into a conical flask, add a few drops of indicator (e.g. methyl orange) and add acid from a burette until the indicator just changes colour. Record the volume.
• Repeat with the same volumes but no indicator, then evaporate and crystallise as in Method A.

Supplement Method C: insoluble salt by precipitation

• Mix two solutions that between them contain the ions of the insoluble salt, e.g. barium chloride and sodium sulfate:
  BaCl₂(aq) + Na₂SO₄(aq) → BaSO₄(s) + 2NaCl(aq)
  Ionic equation: Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s)
• Filter to collect the precipitate, wash it with distilled water to remove soluble impurities, then dry it in a warm oven.

Water of crystallisation

• Hydrated substance: a substance that is chemically combined with water, e.g. CuSO₄·5H₂O.
• Anhydrous substance: a substance containing no water.
• Water of crystallisation: the water molecules present in the crystals of a hydrated substance.

Arrangement of elements

• Elements are arranged in order of increasing proton number (atomic number).
• Periods are horizontal rows; groups are vertical columns.
• Across a period, elements change from metallic to non-metallic character (e.g. period 3: Na, Mg, Al are metals; Si is a metalloid; P, S, Cl, Ar are non-metals).
• The group number equals the number of outer-shell electrons (Group I has 1, Group VII has 7). Group VIII (Group 0) noble gases have a full outer shell.
• The period number equals the number of occupied electron shells, e.g. chlorine (2,8,7) is in period 3, Group VII.
• Elements in the same group have similar chemical properties because they have the same number of outer-shell electrons.
• The Periodic Table can be used to predict the charge on an ion: Group I forms 1+, Group II 2+, Group III 3+, Group VI 2−, Group VII 1−.

Group I: the alkali metals

• Lithium, sodium and potassium are soft metals (can be cut with a knife) with relatively low melting points and low densities.
• Freshly cut surfaces are shiny but tarnish quickly in air, so they are stored under oil.
• They react with cold water to form a metal hydroxide (alkaline solution) and hydrogen:
  2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)

Trends down Group I

• Reactivity increases.
• Melting point decreases.
• Density generally increases (potassium is an exception, less dense than sodium).
• Use the trends to predict: rubidium and caesium are even softer, melt at lower temperatures and react explosively with water.

Group VII: the halogens

• Halogens exist as diatomic molecules: Cl₂, Br₂, I₂.

Trends down Group VII

• Colour becomes darker.
• Density and melting/boiling points increase (gas → liquid → solid).
• Reactivity decreases — the opposite of Group I.
• Predictions: fluorine is a very pale yellow gas and the most reactive halogen; astatine would be a dark solid and the least reactive.

Displacement reactions

A more reactive halogen displaces a less reactive halogen from an aqueous solution of its halide.

Cl₂(aq) + 2KBr(aq) → 2KCl(aq) + Br₂(aq)

• Supplement Ionic equation: Cl₂(aq) + 2Br⁻(aq) → 2Cl⁻(aq) + Br₂(aq)
• Supplement This is a redox reaction: chlorine gains electrons (reduced, Cl₂ + 2e⁻ → 2Cl⁻) and bromide ions lose electrons (oxidised, 2Br⁻ → Br₂ + 2e⁻). Chlorine is the oxidising agent.

Transition elements

Transition elements (e.g. iron, copper, nickel, chromium) are metals with:

Noble gases (Group VIII)

• Helium, neon, argon, krypton, xenon and radon are unreactive, monatomic gases.
• Explanation: their atoms have a full outer shell of electrons (a stable electronic configuration; helium has 2, the others have 8), so they do not need to lose, gain or share electrons.
• Uses linked to unreactivity: argon in light bulbs and as an inert atmosphere in welding; helium in balloons (also low density).

Properties of metals

Physical properties (compared with non-metals)

General chemical properties

• With dilute acids: metal + acid → salt + hydrogen, e.g. Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g)
• With cold water: reactive metals give a metal hydroxide + hydrogen, e.g. 2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)
• With steam: less reactive metals give a metal oxide + hydrogen, e.g. Mg(s) + H₂O(g) → MgO(s) + H₂(g)
• With oxygen: metal + oxygen → metal oxide, e.g. 2Cu(s) + O₂(g) → 2CuO(s). Metal oxides are basic.

Uses of metals

Always link the use to a property.

Alloys

• An alloy is a mixture of a metal with other elements (metals or non-metals such as carbon).
• Brass is a mixture of copper and zinc.
• Stainless steel is a mixture of iron with other elements such as chromium, nickel and carbon.
• Alloys can be harder and stronger than the pure metals, which is why they are more useful.
• Stainless steel is used in cutlery because it is hard and resistant to rusting.
• Supplement In a pure metal the positive ions are all the same size and arranged in regular layers that slide easily over one another. In an alloy, atoms of different sizes distort the layers, so the layers cannot slide as easily. This makes the alloy harder and stronger.
• You should be able to identify an alloy from a diagram: a lattice with a few differently sized circles among the regular metal ions.

Reactivity series

Most reactive at the top (potassium), least reactive at the bottom (gold). Carbon and hydrogen are included as reference points.

Displacement reactions

• A more reactive metal displaces a less reactive metal from a solution of its salt or from its oxide.
• Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s): blue solution fades, pink-brown copper forms, temperature rises.
• Thermite: 2Al(s) + Fe₂O₃(s) → Al₂O₃(s) + 2Fe(l)
• You can use displacement results (and reactions with water and acid) to put unfamiliar metals into order of reactivity.
• Supplement Reactivity is the tendency of a metal to form positive ions. A more reactive metal loses electrons more readily. Ionic equation: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s). Zinc is oxidised (loses electrons); copper ions are reduced (gain electrons).
• Supplement Aluminium seems unreactive (e.g. no reaction with water, slow start with acid) because it is covered by a thin, tough, unreactive layer of aluminium oxide that prevents contact with water or acid.

Corrosion of metals

• Rusting is the corrosion of iron and steel. It needs both water and oxygen. Rust is hydrated iron(III) oxide.
• Classic experiment: iron nails in three tubes: (1) air and water → rusts; (2) boiled water with an oil layer (no oxygen) → no rust; (3) dry air with anhydrous calcium chloride (no water) → no rust.
• Salt water speeds up rusting.

Preventing rusting

• Barrier methods keep out water and oxygen: painting, greasing or oiling, coating with plastic.
• Galvanising: coating iron with a layer of zinc.
• Supplement Barrier methods work only while the coating is unbroken.
• Supplement Sacrificial protection: a more reactive metal (zinc or magnesium) is attached to the iron. The more reactive metal loses electrons (is oxidised) in preference to iron, so iron does not rust. Used on ship hulls and pipelines. Zn(s) → Zn²⁺(aq) + 2e⁻
• Supplement Galvanising protects in two ways: the zinc acts as a barrier, and if scratched the zinc still corrodes sacrificially instead of the iron.

Extraction of metals

• The ease of extracting a metal depends on its position in the reactivity series.
• Metals above carbon (K, Na, Ca, Mg, Al) are extracted by electrolysis of the molten compound.
• Metals below carbon (Zn, Fe, Cu) are extracted by reduction with carbon (or carbon monoxide).
• Very unreactive metals (silver, gold) are found uncombined (native).

Iron from hematite in the blast furnace

Raw materials: hematite (iron ore, mainly Fe₂O₃), coke (carbon), limestone (CaCO₃) and hot air.

Supplement You must be able to write all five symbol equations above. Word equations are enough for Core.

Aluminium from bauxite

• The main ore of aluminium is bauxite, which is purified to aluminium oxide (alumina), Al₂O₃.
• Aluminium is too reactive to be reduced by carbon, so it is extracted by electrolysis.
• Supplement Aluminium oxide is dissolved in molten cryolite. This lowers the operating temperature (saving energy and cost) and makes a better conductor than molten alumina alone.
• Supplement Cathode (−): Al³⁺ + 3e⁻ → Al (molten aluminium collects at the bottom).
• Supplement Anode (+): 2O²⁻ → O₂ + 4e⁻. The carbon (graphite) anodes react with the oxygen to form carbon dioxide, C + O₂ → CO₂, so they burn away and must be replaced regularly.

Water

Chemical tests for water

• These tests show that water is present; they do not show it is pure.
• Physical test for purity: pure water melts at exactly 0 °C and boils at exactly 100 °C (at atmospheric pressure).
• Distilled water is used in practical chemistry instead of tap water because it contains fewer chemical impurities that could interfere with results.

Substances found in water from natural sources and from pollution

Supplement You need to state which substances are beneficial and which are harmful, and why, as shown in the table. Deoxygenation happens because nitrates and phosphates cause rapid growth of algae; when the algae die, bacteria decompose them and use up the dissolved oxygen, so fish die (eutrophication).

Treatment of the domestic water supply

Fertilisers

• Common fertilisers are ammonium salts (e.g. ammonium nitrate, ammonium sulfate, ammonium phosphate) and nitrates (e.g. potassium nitrate).
• NPK fertilisers provide the three elements needed for improved plant growth: nitrogen (N), phosphorus (P) and potassium (K).
• Nitrogen is needed for making proteins, so plants grow larger and greener; phosphorus helps root growth; potassium helps flowering and fruiting.
• Ammonium salts are made by neutralising ammonia with an acid, e.g. NH₃(aq) + HNO₃(aq) → NH₄NO₃(aq).

Air quality and climate

Composition of clean, dry air

Sources and adverse effects of air pollutants

Acid rain damages buildings made of limestone and marble, corrodes metal structures, kills trees and makes lakes too acidic for fish.

Supplement The greenhouse effect

• The Sun's radiation passes through the atmosphere and warms the Earth's surface.
• The Earth's surface emits thermal (infrared) energy.
• Greenhouse gases (carbon dioxide, methane) absorb some of this thermal energy and re-emit it in all directions, including back towards the Earth, while some is reflected back into space.
• Higher concentrations of greenhouse gases mean more thermal energy is kept in the atmosphere, so the average temperature rises (global warming) and climate changes.

Supplement Strategies to reduce environmental effects

Climate change

• Planting trees (they remove CO₂ by photosynthesis).
• Reducing livestock farming (less methane).
• Decreasing use of fossil fuels.
• Increasing use of hydrogen as a fuel and of renewable energy such as wind and solar.

Acid rain

• Using catalytic converters in vehicles.
• Using low-sulfur fuels.
• Flue gas desulfurisation using calcium oxide in power stations: CaO(s) + SO₂(g) → CaSO₃(s)

Catalytic converters

Catalytic converters remove oxides of nitrogen and carbon monoxide (and unburnt hydrocarbons) from exhaust gases using a platinum or rhodium catalyst:

• 2CO(g) + 2NO(g) → 2CO₂(g) + N₂(g)
• 2NO₂(g) + 4CO(g) → N₂(g) + 4CO₂(g)
• 2CO(g) + O₂(g) → 2CO₂(g)

Photosynthesis

• Word equation: carbon dioxide + water → glucose + oxygen, in the presence of light and chlorophyll.
• Supplement Symbol equation: 6CO₂(g) + 6H₂O(l) → C₆H₁₂O₆(aq) + 6O₂(g)
• Photosynthesis removes carbon dioxide from the atmosphere, which is why deforestation adds to climate change.

Formulae and terminology

• A homologous series is a family of similar compounds with similar chemical properties due to the presence of the same functional group.
• Supplement Members of a homologous series have the same general formula, differ from one member to the next by a CH₂ unit, show a trend in physical properties (e.g. boiling point rises with chain length) and have similar chemical properties.
• Saturated: only single carbon–carbon bonds. Unsaturated: contains one or more C=C double (or triple) bonds.
• Supplement Structural isomers are compounds with the same molecular formula but different structural formulae, e.g. butane CH₃CH₂CH₂CH₃ and methylpropane CH₃CH(CH₃)CH₃; or but-1-ene and but-2-ene.

Naming organic compounds

• Core: name and draw displayed formulae of methane, ethane, ethene, ethanol, ethanoic acid and the products of their reactions.
• Supplement Name and draw structural and displayed formulae of unbranched alkanes, alkenes, alcohols and carboxylic acids with up to four carbon atoms, and esters made from alcohols and acids each with up to four carbon atoms.
• Supplement Ester names: alcohol part first (ending -yl), acid part second (ending -oate). Methanol + propanoic acid → methyl propanoate, CH₃CH₂COOCH₃.

Fuels and petroleum

• Fossil fuels: coal, natural gas and petroleum. The main constituent of natural gas is methane.
• Petroleum is a mixture of hydrocarbons, separated into useful fractions by fractional distillation.
• The crude oil is vaporised; vapours rise up a column that is hot at the bottom and cooler at the top. Each fraction condenses where the temperature falls below its boiling point range.

Moving from the bottom to the top of the column: chain length decreases, volatility increases, boiling point decreases and viscosity decreases.

Alkanes

• Saturated hydrocarbons; generally unreactive apart from combustion.
• Complete combustion: CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l)
• Substitution with chlorine in the presence of ultraviolet light (a photochemical reaction): a hydrogen atom is replaced by a chlorine atom.
• Supplement CH₄(g) + Cl₂(g) → CH₃Cl(g) + HCl(g) (UV light). Further substitution can give CH₂Cl₂, CHCl₃ and CCl₄.

Alkenes

• Made by cracking larger alkane molecules, using a high temperature and a catalyst. Cracking produces alkenes and hydrogen as well as smaller alkanes, e.g. C₁₀H₂₂ → C₈H₁₈ + C₂H₄.
• Reasons for cracking: supply of long-chain fractions exceeds demand, while demand for shorter chains (petrol) and alkenes (for polymers) is greater than supply.
• Test for unsaturation: add aqueous bromine. Alkenes turn it from orange to colourless; alkanes leave it orange.
• Supplement Alkenes undergo addition reactions in which only one product forms:

Alcohols

• Ethanol is made by fermentation of aqueous glucose at 25–35 °C with yeast, in the absence of oxygen: C₆H₁₂O₆(aq) → 2C₂H₅OH(aq) + 2CO₂(g)
• Or by catalytic addition of steam to ethene (see table above).
• Combustion: C₂H₅OH(l) + 3O₂(g) → 2CO₂(g) + 3H₂O(l)
• Uses of ethanol: as a solvent and as a fuel.

Carboxylic acids

• Ethanoic acid is a weak acid and shows typical acid reactions, forming ethanoate salts:
• With metals: 2CH₃COOH + Mg → (CH₃COO)₂Mg + H₂ (magnesium ethanoate)
• With bases: CH₃COOH + NaOH → CH₃COONa + H₂O (sodium ethanoate)
• With carbonates: 2CH₃COOH + Na₂CO₃ → 2CH₃COONa + H₂O + CO₂
• Supplement Ethanoic acid forms by oxidation of ethanol: (a) by bacterial oxidation during fermentation (wine turning to vinegar), or (b) with acidified aqueous potassium manganate(VII), which turns from purple to colourless.
• Supplement Carboxylic acids react with alcohols, with a concentrated sulfuric acid catalyst and heat, to make esters and water: CH₃COOH + C₂H₅OH ⇌ CH₃COOC₂H₅ + H₂O (ethyl ethanoate).

Polymers

• Polymers are large molecules built up from many small units called monomers. Plastics are made from polymers.
• Addition polymerisation: many alkene monomers join; the C=C opens. Ethene forms poly(ethene).
• Disposal problems of plastics (they are non-biodegradable): they take up space and persist in landfill; they accumulate in oceans, harming marine life; burning them forms toxic gases.
• Supplement Identify the repeat unit and linkages in a polymer; deduce the polymer from a monomer and vice versa. Draw the repeat unit in brackets with continuation bonds through them and "n" outside, e.g. propene gives poly(propene) with repeat unit –[CH₂–CH(CH₃)]–.
• Supplement Addition polymerisation gives only the polymer; condensation polymerisation joins monomers with two functional groups and also releases a small molecule such as water.
• Supplement Nylon is a polyamide (amide linkage –CONH–) from a dicarboxylic acid and a diamine. PET is a polyester (ester linkage –COO–) from a dicarboxylic acid and a diol. PET can be converted back into monomers and re-polymerised, which aids recycling.
• Supplement Proteins are natural polyamides formed by condensation polymerisation of amino acid monomers, joined by amide linkages.

Experimental design

Measuring apparatus

• Be ready to suggest advantages and disadvantages of methods and apparatus, e.g. a gas syringe is more accurate than collecting over water, but collecting over water cannot be used for gases that dissolve (ammonia, hydrogen chloride).
• Read volumes at the bottom of the meniscus, at eye level.

Key terms

Acid–base titrations

• End-point colours: methyl orange goes from yellow (alkali) to orange at the end-point (red in excess acid); thymolphthalein goes from blue to colourless as acid is added to alkali.
• Titre = final reading − initial reading. To make a pure salt, repeat without indicator using the known volumes, then crystallise.

Chromatography

• Paper chromatography separates mixtures of soluble coloured substances using a suitable solvent.
• Method: draw a start line in pencil (graphite does not dissolve); spot the samples on it; stand the paper in solvent with the solvent level below the start line (so the spots do not dissolve into the bulk solvent); cover and allow the solvent to rise; remove and mark the solvent front.
• A pure substance gives one spot; an impure substance (mixture) gives two or more spots.
• Identify unknowns by comparing their spots with those of known substances run on the same paper: spots at the same height match.
• Supplement Colourless substances are made visible with a locating agent, sprayed on the dried chromatogram.
• Supplement R_f = distance travelled by substance ÷ distance travelled by solvent. R_f has no units and is always less than 1.

Separation and purification

• Choose a technique from the information given, e.g. solubility and boiling point data.
• Assessing purity: a pure substance has a sharp melting point and boiling point. Impurities lower the melting point and raise the boiling point, and make them occur over a range of temperatures.
• Purity matters in food and medicines, where impurities could be harmful.

Identification of ions and gases

Tests for anions

Tests for aqueous cations

Aluminium and zinc give identical results with sodium hydroxide; use aqueous ammonia to tell them apart (only zinc hydroxide dissolves in excess ammonia).

Flame tests for metal ions

Use a clean nichrome or platinum wire dipped in concentrated hydrochloric acid, then in the sample, and hold it in a blue (roaring) Bunsen flame.

Tests for gases

Water is tested with anhydrous copper(II) sulfate (white to blue) or cobalt(II) chloride (blue to pink); see Topic 10.