GCSE Chemistry (AQA) equations and key facts

Atomic structure and the periodic table

Atoms, elements, compounds and mixtures

• Filtration: separates an insoluble solid from a liquid.
• Crystallisation: gets a soluble solid out of its solution.
• Simple distillation: collects the solvent from a solution, e.g. pure water from salt solution.
• Fractional distillation: separates liquids that mix and have different boiling points.
• Paper chromatography: separates dissolved substances, e.g. the dyes in an ink.
• State symbols: (s) solid, (l) liquid, (g) gas, (aq) dissolved in water.
• To balance an equation, change the large numbers in front of the formulae, never the small numbers inside them.

Atomic models, subatomic particles, isotopes and electronic structure

• Relative charge: proton +1, neutron 0, electron −1.
• Relative mass: proton 1, neutron 1, electron very small.
• number of neutrons = mass number − atomic number
• In an atom, electrons = protons. A positive ion has lost electrons; a negative ion has gained electrons.
• Radius of an atom is about 0.1 nm (1 × 10⁻¹⁰ m); the radius of the nucleus is less than 1/10 000 of this.
• relative atomic mass = sum of (mass number × percentage of each isotope) ÷ 100
• Electronic structure example: sodium (11 electrons) is 2,8,1.

The periodic table: development, Group 0, Group 1 and Group 7

• Group 0 (noble gases): boiling points increase down the group.
• Group 1: one outer electron, forms 1+ ions. Down the group the outer electron is further from the nucleus and is lost more easily.
• Group 1 metal + water → metal hydroxide + hydrogen, e.g. 2Li + 2H₂O → 2LiOH + H₂
• Group 7 (halogens): seven outer electrons, molecules of two atoms (Cl₂), form 1− ions with metals.
• Group 7: melting and boiling points increase down the group; reactivity decreases because an electron is gained less easily.
• A more reactive halogen displaces a less reactive halogen from a solution of its salt.

Properties of transition metals

• Transition metals: high melting point, high density, hard and strong. Group 1 metals: low melting point, low density, soft.
• Group 1 metals react vigorously with cold water; transition metals react slowly or not at all.
• The Roman numeral in a name gives the charge on the metal ion: iron(III) is Fe³⁺, copper(II) is Cu²⁺.
• Transition metal compounds are often coloured, e.g. many copper(II) compounds are blue. Group 1 compounds are white.
• Iron is the catalyst in the Haber process; manganese(IV) oxide is a catalyst for the decomposition of hydrogen peroxide.
• A catalyst makes a reaction faster and is not used up, so its mass is the same at the end.

Bonding, structure, and the properties of matter

Ionic, covalent and metallic bonding

• Ion charges: Group 1 forms 1+, Group 2 forms 2+, Group 6 forms 2−, Group 7 forms 1−.
• In the formula of an ionic compound the total positive charge equals the total negative charge.
• One shared pair of electrons = one covalent bond. H₂ has a single bond, O₂ a double bond, N₂ a triple bond.
• After bonding, most atoms have a full outer shell (8 electrons; 2 for hydrogen).
• Dot and cross diagrams show where the electrons come from but not the 3D shape of the molecule.
• Metal + non-metal: ionic. Non-metal + non-metal: covalent. Metal only: metallic.

How bonding and structure are related to properties

• Ionic compounds: high melting points; conduct when molten or dissolved in water, but not when solid, because the ions are then fixed in place.
• Larger ion charges give stronger attraction and a higher melting point.
• Small molecules: low melting and boiling points; do not conduct because the molecules have no overall charge.
• Bigger molecules have stronger intermolecular forces, so polymers are solids at room temperature.
• Giant covalent structures: very high melting points; all atoms joined by strong covalent bonds.
• Metals: conduct electricity and heat through delocalised electrons; layers of atoms slide, so metals can be bent and shaped.
• Alloys are harder than pure metals: atoms of different sizes distort the layers so they slide less easily.

Structure and bonding of carbon; nanoparticles

• Diamond: 4 bonds per carbon atom, no delocalised electrons.
• Graphite: 3 bonds per carbon atom, 1 delocalised electron per atom, weak forces between layers.
• Graphene: one layer of graphite, one atom thick; strong and conducts electricity.
• Buckminsterfullerene, C₆₀, is a hollow sphere. Carbon nanotubes are cylinders with a very high length to diameter ratio and high tensile strength.
• Sizes: nanoparticles 1 to 100 nm; fine particles 100 to 2500 nm; coarse particles (dust) 2500 to 10 000 nm.
• When the side of a cube gets 10 times smaller, its surface area to volume ratio gets 10 times larger.
• Nanoparticles are used in sun creams, cosmetics, medicine, electronics and as catalysts; the risk is that their effects inside the body are not fully known.

Quantitative chemistry

Conservation of mass and relative formula mass

• total mass of reactants = total mass of products
• M_r = sum of the A_r values of all the atoms in the formula
• A number outside a bracket multiplies everything inside it: Ca(OH)₂ has 1 Ca, 2 O and 2 H.
• Metal heated in air: mass of solid increases (oxygen is added).
• Metal carbonate heated or added to acid: mass decreases (carbon dioxide escapes).
• In a sealed container the total mass does not change.
• uncertainty = range ÷ 2, written as mean ± uncertainty

Moles, reacting masses and limiting reactants (HT)

• number of moles = mass ÷ M_r (mass in g)
• mass = number of moles × M_r
• number of particles = number of moles × 6.02 × 10²³
• The large numbers in a balanced equation are the mole ratio, e.g. 2Mg + O₂ → 2MgO means 2 mol Mg gives 2 mol MgO.
• To find the limiting reactant, divide the moles of each reactant by its number in the equation; the smaller result is the limiting one.
• To balance an equation from masses, change each mass to moles and find the simplest whole number ratio.
• volume of gas at room temperature and pressure = number of moles × 24 dm³

Yield and atom economy

• percentage yield = (mass of product actually made ÷ maximum theoretical mass of product) × 100
• mass actually made = (percentage yield ÷ 100) × theoretical mass
• atom economy = (M_r of desired product from the equation ÷ sum of M_r of all reactants from the equation) × 100
• Include the large numbers from the equation when adding up the M_r values.
• A reaction with only one product has an atom economy of 100%.
• Percentage yield can never be more than 100%.

Concentrations of solutions and volumes of gases

• 1 dm³ = 1000 cm³, so volume in dm³ = volume in cm³ ÷ 1000
• mass of solute (g) = concentration (g/dm³) × volume (dm³)
• moles of solute = concentration (mol/dm³) × volume (dm³)
• concentration in g/dm³ = concentration in mol/dm³ × M_r
• volume of gas (dm³) = moles × 24, at room temperature and pressure
• At the same temperature and pressure, gas volumes react in the same ratio as the numbers in the balanced equation.

Chemical changes

Reactivity of metals

• Order, most reactive first: potassium, sodium, lithium, calcium, magnesium, (carbon), zinc, iron, (hydrogen), copper, gold
• Oxidation = gain of oxygen; reduction = loss of oxygen
• Oxidation is loss of electrons, reduction is gain of electrons (OIL RIG)
• A more reactive metal displaces a less reactive metal from its compound
• Metals less reactive than carbon are extracted by heating the oxide with carbon
• Gold is so unreactive that it is found in the Earth as the metal itself
• Ionic equations leave out the ions that do not change (spectator ions)

Reactions of acids

• acid + metal → salt + hydrogen
• acid + alkali or base (metal oxide or hydroxide) → salt + water
• acid + metal carbonate → salt + water + carbon dioxide
• Neutralisation: H⁺(aq) + OH⁻(aq) → H₂O(l)
• Soluble salt: add excess insoluble base to warm acid, filter, heat the filtrate gently, leave to crystallise
• Titration: measure the alkali with a pipette, add acid from a burette until the indicator changes colour
• When pH falls by 1, the H⁺ concentration becomes 10 times greater

Electrolysis

• Cathode (negative): positive ions gain electrons (reduction), e.g. Pb²⁺ + 2e⁻ → Pb
• Anode (positive): negative ions lose electrons (oxidation), e.g. 2Br⁻ → Br₂ + 2e⁻
• Solution, cathode: hydrogen forms if the metal is more reactive than hydrogen; if not, the metal forms (copper, silver)
• Solution, anode: a halogen forms if halide ions are present; if not, oxygen forms
• Oxygen at the anode: 4OH⁻ → O₂ + 2H₂O + 4e⁻
• Aluminium oxide is dissolved in molten cryolite to lower the melting point and save energy
• Aluminium: Al³⁺ + 3e⁻ → Al at the cathode; oxygen at the carbon anodes, which burn away to carbon dioxide and are replaced

Energy changes

Exothermic and endothermic reactions

• Exothermic: temperature rises, products are at a lower energy than the reactants, energy change is negative
• Endothermic: temperature falls, products are at a higher energy than the reactants, energy change is positive
• Activation energy = energy at the top of the curve − energy of the reactants
• Overall energy change = energy of the products − energy of the reactants
• Bond breaking is endothermic; bond making is exothermic
• Energy change = energy needed to break bonds − energy released making bonds
• Uses: hand warmers and self-heating cans (exothermic); sports injury cold packs (endothermic)

Chemical cells and fuel cells

• Cell = two different metals + an electrolyte
• Voltage depends on the two metals and on the electrolyte; a bigger gap in reactivity gives a bigger voltage
• Non-rechargeable cells stop when a reactant is used up; rechargeable cells are recharged by an external current
• Fuel cell overall: 2H₂ + O₂ → 2H₂O
• Negative electrode: H₂ → 2H⁺ + 2e⁻
• Positive electrode: O₂ + 4H⁺ + 4e⁻ → 2H₂O
• Fuel cells: only water is made and no recharging is needed, but hydrogen is flammable and hard to store

The rate and extent of chemical change

Calculating and measuring rates of reaction

• mean rate = quantity of reactant used ÷ time taken
• mean rate = quantity of product formed ÷ time taken
• Units: g/s, cm³/s or mol/s
• Rate at a given time = gradient of the tangent = change in y ÷ change in x
• Steeper curve = faster reaction; flat curve = reaction finished
• Disappearing cross: a shorter time means a faster rate

Factors affecting rate, collision theory and catalysts

• Higher concentration or pressure: more particles in the same volume, so more frequent collisions
• Higher temperature: more frequent collisions and more of them have the activation energy
• Smaller pieces: larger surface area to volume ratio, so more frequent collisions
• Catalyst: lower activation energy; it is not used up and is not written in the equation
• On a reaction profile a catalyst lowers the peak only; the reactant and product energies stay the same
• Enzymes are biological catalysts
• Cube of side a: surface area = 6 × a × a, volume = a × a × a

Reversible reactions and dynamic equilibrium

• At equilibrium: forward rate = reverse rate, and the amounts stay constant (they are not necessarily equal)
• Add more of a reactant: the position moves towards the products
• Remove a product: the position moves towards the products
• Higher temperature favours the endothermic direction; lower temperature favours the exothermic direction
• Higher pressure favours the side with fewer gas molecules; lower pressure favours the side with more
• Equal numbers of gas molecules on both sides: pressure has no effect on the position
• Hydrated copper(II) sulfate (blue) ⇌ anhydrous copper(II) sulfate (white) + water

Organic chemistry

Crude oil, alkanes, fractional distillation and cracking

• Alkane general formula: C_nH_2n+2
• Methane CH₄, ethane C₂H₆, propane C₃H₈, butane C₄H₁₀
• Larger molecules: higher boiling point, higher viscosity, lower flammability
• Complete combustion: hydrocarbon + oxygen → carbon dioxide + water
• Catalytic cracking: hydrocarbon vapour passed over a hot catalyst; steam cracking: vapour mixed with steam at a very high temperature
• Alkenes turn orange bromine water colourless; alkanes do not
• Fractions from the top down: petroleum gases, petrol, kerosene (aircraft fuel), diesel oil, heavy fuel oil

Alkenes, alcohols and carboxylic acids

• Alkenes: general formula C_nH_2n; the first four are ethene, propene, butene and pentene.
• Alkene + hydrogen (catalyst) → alkane; alkene + steam (catalyst) → alcohol; alkene + halogen → a compound with two halogen atoms, e.g. C₂H₄ + Br₂ → C₂H₄Br₂.
• Alkenes turn orange bromine water colourless; alkanes do not. Alkenes burn in air with a smoky flame because combustion is incomplete.
• Alcohols (–OH): methanol, ethanol, propanol, butanol. Carboxylic acids (–COOH): methanoic, ethanoic, propanoic, butanoic acid.
• Fermentation: glucose → ethanol + carbon dioxide (C₆H₁₂O₆ → 2C₂H₅OH + 2CO₂), with yeast, at about 30 °C, with no air.
• Carboxylic acid + alcohol ⇌ ester + water, e.g. ethanoic acid + ethanol ⇌ ethyl ethanoate + water (acid catalyst).

Synthetic and naturally occurring polymers

• To draw a repeating unit: change C=C to C–C, keep all the other atoms, and draw a bond out of each end through the brackets, with n after the bracket.
• Naming: the polymer of ethene is poly(ethene), of propene is poly(propene), of chloroethene is poly(chloroethene).
• Addition polymerisation gives one product; condensation polymerisation gives the polymer and a small molecule such as water.
• Polyester: a diol (two –OH groups) + a dicarboxylic acid (two –COOH groups) → polyester + water.
• Amino acids have an amine group, –NH₂, and a carboxylic acid group, –COOH. They join by condensation to make polypeptides; proteins are made from different amino acids in one chain.
• DNA is two polymer chains, made from four different nucleotides, twisted into a double helix.
• Starch and cellulose are polymers of glucose; proteins are polymers of amino acids.

Chemical analysis

Purity, formulations and chromatography

• Pure: a sharp melting point that matches the data value. Impure: melts over a range and at a lower temperature; impurities also raise the boiling point.
• Stationary phase = the paper (it does not move). Mobile phase = the solvent (it moves).
• A substance that is more attracted to the solvent travels further; one that is more attracted to the paper travels less far.
• R_f = distance moved by the substance ÷ distance moved by the solvent. Measure both from the start line.
• R_f has no unit and is always less than 1.
• A pure substance gives one spot in every solvent; a mixture can give two or more spots.
• Spots at the same height on the same paper (the same R_f in the same solvent) are probably the same substance.

Tests for gases and ions; instrumental methods

• Gases: hydrogen burns with a squeaky pop (lighted splint); oxygen relights a glowing splint; carbon dioxide turns limewater milky; chlorine bleaches damp litmus paper white.
• Flame colours: lithium crimson, sodium yellow, potassium lilac, calcium orange-red, copper green.
• With sodium hydroxide solution: Cu²⁺ blue precipitate, Fe²⁺ green, Fe³⁺ brown; Al³⁺, Ca²⁺ and Mg²⁺ white, and only the aluminium precipitate dissolves in excess.
• Carbonate: add dilute acid; carbon dioxide is given off, which turns limewater milky.
• Halides: add dilute nitric acid, then silver nitrate solution. Chloride gives a white precipitate, bromide cream, iodide yellow.
• Sulfate: add dilute hydrochloric acid, then barium chloride solution; a white precipitate forms.
• Flame emission spectroscopy: the line spectrum identifies the metal ions, even in a mixture, and the intensity of the light gives their concentration.

Chemistry of the atmosphere

Composition and evolution of the Earth's atmosphere

• Air today: about 80% nitrogen, about 20% oxygen, and small proportions of argon (and other noble gases), carbon dioxide and water vapour.
• Volcanoes released carbon dioxide, water vapour and nitrogen, with small amounts of methane and ammonia.
• Photosynthesis: carbon dioxide + water → glucose + oxygen (6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂), using light.
• Algae first produced oxygen about 2.7 billion years ago; as oxygen built up, animals could evolve.
• Carbon dioxide dissolved in the oceans and formed carbonate precipitates, which became sediments and then rocks such as limestone.
• Coal formed from plant remains; crude oil and natural gas formed from plankton buried in mud. All of these hold carbon.
• The theories are uncertain because the Earth is about 4.6 billion years old and there is little evidence from so long ago.

Greenhouse gases, climate change and atmospheric pollutants

• The main greenhouse gases are carbon dioxide, methane and water vapour.
• More carbon dioxide: burning fossil fuels and cutting down forests. More methane: cattle, rice fields and landfill sites.
• Effects of climate change: rising sea levels and flooding, more extreme weather, changes in rainfall, and changes to where crops and wildlife can live.
• Carbon footprint: the total amount of carbon dioxide and other greenhouse gases given out over the full life cycle of a product, service or event.
• Incomplete combustion (too little oxygen) gives carbon monoxide, a toxic gas with no colour or smell, and soot (carbon particles), which causes global dimming and breathing problems.
• Sulfur dioxide comes from sulfur impurities in the fuel; oxides of nitrogen form when nitrogen and oxygen from the air react in a hot engine. Both cause acid rain and breathing problems.

Using resources

Using the Earth's resources and obtaining potable water

• Fresh water to potable water: choose a suitable source, pass it through filter beds to remove solids, then sterilise it with chlorine, ozone or ultraviolet light.
• Desalination: distillation, or reverse osmosis using membranes. Both need large amounts of energy, so they are expensive.
• Sewage treatment: screening and grit removal, then sedimentation to give sewage sludge and effluent.
• Sludge is digested by bacteria without air (anaerobic); effluent is treated by bacteria with air (aerobic).
• Phytomining: plants take up metal compounds, are harvested and burned, and the ash contains the metal compounds.
• Bioleaching: bacteria produce a solution called a leachate that contains metal compounds.
• Copper is then obtained from these compounds by displacement with scrap iron or by electrolysis.

Life cycle assessment and recycling

• LCA stages: raw materials, manufacturing and packaging, use and operation, disposal; include transport and distribution at each stage.
• To compare two products fairly, work out the impact per use, not just the impact of making one item.
• A shortened or selective LCA can leave out stages, so it can be misused to support advertising claims.
• Reuse: the object is used again as it is, e.g. a glass bottle is washed and refilled.
• Recycle: the material is processed into something new, e.g. glass is crushed and melted, metals are melted and recast.
• Recycling metals uses less energy than extracting them from ores, saves limited ore, and reduces mining and waste.
• Adding scrap steel to iron from the blast furnace reduces the amount of iron ore that must be extracted.

Using materials

• Rusting needs iron + oxygen + water. Barriers: paint, grease or oil, electroplating. Aluminium is protected by its own thin oxide layer.
• Sacrificial protection: a more reactive metal such as zinc corrodes in place of the iron. Galvanising coats iron with zinc, which works as a barrier and sacrificially.
• Bronze = copper + tin. Brass = copper + zinc. Steels = iron + carbon (and other metals).
• High carbon steel is strong but brittle; low carbon steel is softer and easily shaped; stainless steel (with chromium and nickel) is hard and resists corrosion. Aluminium alloys have low density.
• Gold: 24 carat is 100% gold, so percentage of gold = (carat ÷ 24) × 100. 18 carat is 75% gold.
• Soda-lime glass: heat sand, sodium carbonate and limestone. Borosilicate glass: sand and boron trioxide; it melts at a higher temperature. Clay ceramics: shape wet clay, then heat it in a furnace.
• Thermosoftening polymers melt when heated (no cross-links between chains); thermosetting polymers do not melt (strong cross-links). LD and HD poly(ethene) are both made from ethene under different conditions.

The Haber process and NPK fertilisers

• nitrogen + hydrogen ⇌ ammonia (N₂ + 3H₂ ⇌ 2NH₃); the forward reaction is exothermic.
• Raw materials: nitrogen from the air, hydrogen from natural gas.
• Conditions: about 450 °C, about 200 atmospheres, iron catalyst.
• Ammonia makes ammonium salts and nitric acid, e.g. ammonia + nitric acid → ammonium nitrate (NH₃ + HNO₃ → NH₄NO₃).
• Potassium chloride, potassium sulfate and phosphate rock are mined. Phosphate rock is insoluble, so it is treated with acid to make soluble salts.
• Phosphate rock + nitric acid → phosphoric acid and calcium nitrate; + sulfuric acid → single superphosphate; + phosphoric acid → triple superphosphate.
• In the laboratory: dilute solutions, titration, then crystallisation, in small batches. In industry: ammonia gas and concentrated acid in a continuous process, on a large scale.