AQA GCSE Chemistry Foundation: grade by grade

Every skill, from the first marks to the top grade. For each grade you also need the skills for the grades below it. Tick them off in the app's Notes section.

Grade 3

Grade 3 checklist (PDF)

Tell elements from compounds using formulae An element's formula has one symbol (Fe, O₂); a compound's formula has two or more different symbols (CO₂, NaCl). Atoms, elements and compounds
Recall symbols of the first 20 elements Also the Group 1 and Group 7 elements and others in the course, with the correct capital and lower-case letters. Atoms, elements and compounds
Define a mixture Two or more elements or compounds that are not chemically combined together. Mixtures
State that models change with new evidence A scientific model may be changed or replaced when new experimental evidence doesn't fit it. Development of the model of the atom
Recall the charges of subatomic particles Proton +1, neutron 0, electron −1. Relative charges of subatomic particles
State where each particle is found Protons and neutrons are in the nucleus; electrons are in shells around it. Relative charges of subatomic particles
Recall relative masses of subatomic particles Proton 1, neutron 1, electron very small. Size and mass of atoms
Recall how many electrons each shell holds For the first 20 elements: 2 in the first shell, then up to 8 in the second and 8 in the third. Electronic structure
Identify groups and periods Groups are the vertical columns; periods are the horizontal rows. The periodic table
State how early tables were ordered By atomic weight, because protons, neutrons and electrons had not been discovered. Development of the periodic table
Locate metals and non-metals in the table Metals are on the left and towards the bottom; non-metals are on the right and towards the top. Metals and non-metals
Name Group 0 and some of its elements The noble gases: helium, neon, argon, krypton, xenon and radon. Group 0
Explain the name 'alkali metals' They react with water to form hydroxides that dissolve to give alkaline solutions. Group 1
Recall the halogens and their states At room temperature chlorine is a green gas, bromine a red-brown liquid and iodine a grey-black solid. Group 7
Locate the transition metals in the table They are the central block, between Group 2 and Group 3. Comparison with Group 1 elements
Name the three types of strong bond Ionic, covalent and metallic. Chemical bonds
State that metals lose and non-metals gain electrons Metal atoms lose outer electrons to form positive ions; non-metal atoms gain electrons to form negative ions. Ionic bonding
Describe an ionic compound as a giant structure It is a giant structure (lattice) of ions arranged in a regular pattern. Ionic compounds
Describe a covalent bond as a shared pair A covalent bond is a pair of electrons shared between two atoms. Covalent bonds are strong. Covalent bonding
State that metals have giant structures The atoms in a metal are arranged in a regular pattern in a giant structure. Metallic bonding
Name the changes of state Melting and freezing happen at the melting point; boiling and condensing happen at the boiling point. The three states of matter
Recall the four state symbols (s) solid, (l) liquid, (g) gas, (aq) aqueous: dissolved in water. State symbols
State that ionic compounds have high melting points Ionic compounds such as sodium chloride are solids at room temperature with high melting and boiling points. Properties of ionic compounds
State that small molecules have low boiling points Substances made of small molecules are usually gases or liquids at room temperature. Properties of small molecules
State that polymers have very large molecules Polymer molecules are long chains made of many repeating units. Polymers
Name examples of giant covalent structures Diamond and graphite (both forms of carbon) and silicon dioxide (silica). Giant covalent structures
State that pure metals can be bent and shaped The atoms in a pure metal are arranged in layers. Properties of metals and alloys
State that metals conduct electricity and heat Metals are good conductors of electricity and of thermal energy. Metals as conductors
State that diamond is a form of carbon Diamond is made only of carbon atoms, in a giant covalent structure. Diamond
Describe graphite as layers of carbon atoms Graphite is made of layers of carbon atoms arranged in hexagonal rings. Graphite
Describe graphene as one layer of graphite Graphene is a single layer of carbon atoms in hexagonal rings, one atom thick. Graphene and fullerenes
State the law of conservation of mass No atoms are lost or made in a reaction, so the mass of the products equals the mass of the reactants. Conservation of mass and balanced equations
Count the atoms in a formula Multiply the small numbers by the big number in front, e.g. 2H₂O contains 4 hydrogen atoms and 2 oxygen atoms. Conservation of mass and balanced equations
Read relative atomic masses from the periodic table A_r is the larger of the two numbers in each box, e.g. O = 16, Na = 23. Relative formula mass
State that gases can enter or leave In an open container, gases can escape into the air or react from the air. Mass changes when a reactant or product is a gas
Calculate a mean from repeat readings Add up the results, leaving out any anomalous ones, and divide by how many you added. Chemical measurements and uncertainty
Recall the unit of concentration Mass of solute per volume of solution, usually grams per dm³ (g/dm³). Concentration of solutions
Name the oxide formed by a metal A metal reacting with oxygen forms the metal oxide, e.g. magnesium + oxygen → magnesium oxide. Metal oxides
Recall the order of the reactivity series Potassium, sodium, lithium, calcium, magnesium, (carbon), zinc, iron, (hydrogen), copper, from most to least reactive. The reactivity series
Explain why gold is found as the metal Gold is so unreactive that it is found in the Earth as the metal itself, not as a compound. Extraction of metals and reduction
Name the products of metal + acid A salt and hydrogen, e.g. magnesium + hydrochloric acid → magnesium chloride + hydrogen. Reactions of acids with metals
Name the salt type from each acid Hydrochloric acid gives chlorides, nitric acid gives nitrates and sulfuric acid gives sulfates. Neutralisation and salt production
Name reactants needed to make a salt For example, copper(II) oxide and sulfuric acid make copper(II) sulfate. Soluble salts
Use pH to classify solutions Below 7 is acidic, 7 is neutral and above 7 is alkaline. The pH scale and neutralisation
Name the electrodes and their charges The cathode is the negative electrode and the anode is the positive electrode. The process of electrolysis
State the products from molten lead bromide Lead is produced at the cathode and bromine at the anode. Electrolysis of molten ionic compounds
Define exothermic and endothermic reactions Exothermic reactions transfer energy to the surroundings so their temperature rises; endothermic reactions take in energy so their temperature falls. Energy transfer in exothermic and endothermic reactions
Give examples of exothermic reactions Combustion, many oxidation reactions and neutralisation are all exothermic. Energy transfer in exothermic and endothermic reactions
State what activation energy is The minimum amount of energy that particles must have to react when they collide. Reaction profiles
State what a cell does A cell contains chemicals that react to produce electricity. Cells and batteries
Name what a hydrogen fuel cell uses Hydrogen (the fuel) and oxygen or air are supplied to the cell from outside. Fuel cells
Say what the rate of reaction measures Rate is how quickly a reactant is used up or a product is formed. Calculating rates of reactions
Name the five factors that affect rate Concentration of solutions, pressure of gases, surface area of solids, temperature and the presence of a catalyst. Factors which affect the rate of reaction
State that particles must collide to react A reaction can only happen when reactant particles collide with each other. Collision theory and activation energy
State what a catalyst does It increases the rate of a reaction but is not used up. Catalysts
Recognise the reversible reaction symbol ⇌ The ⇌ symbol shows that a reaction can go in both directions. Reversible reactions
Recall what exothermic and endothermic mean Exothermic reactions transfer energy to the surroundings; endothermic reactions take in energy from the surroundings. Energy changes and reversible reactions
State what crude oil is A finite resource found in rocks: the remains of an ancient biomass, mainly plankton, that was buried in mud. Crude oil, hydrocarbons and alkanes
Define a hydrocarbon A compound whose molecules are made of hydrogen and carbon atoms only. Crude oil, hydrocarbons and alkanes
Name fuels produced from crude oil Petrol, diesel oil, kerosene, heavy fuel oil and liquefied petroleum gases (LPG). Fractional distillation and petrochemicals
Name materials made by the petrochemical industry Solvents, lubricants, polymers and detergents. Fractional distillation and petrochemicals
Name the products of complete combustion A hydrocarbon burns completely in oxygen to form carbon dioxide and water. Properties of hydrocarbons
Define cracking Breaking down large hydrocarbon molecules into smaller, more useful molecules. Cracking and alkenes
State what an alkene is A hydrocarbon with a carbon–carbon double bond, C=C. Structure and formulae of alkenes
Name the functional group of the alkenes C=C, the carbon–carbon double bond. Reactions of alkenes
Name the functional group of the alcohols The –OH group. Alcohols
Give the main uses of alcohols As fuels and solvents; ethanol is the main alcohol in alcoholic drinks. Alcohols
Name the functional group of carboxylic acids The –COOH group. Carboxylic acids
Define monomer and polymer Monomers are small molecules that join together in large numbers to make a very large molecule, a polymer. Addition polymerisation
Name natural polymers that are important for life DNA, proteins, starch and cellulose. DNA and other naturally occurring polymers
Define a pure substance in chemistry A single element or a single compound, not mixed with any other substance. Pure substances
Define a formulation A mixture that has been designed as a useful product. Formulations
Give examples of formulations Fuels, cleaning agents, paints, medicines, alloys, fertilisers and foods. Formulations
Tell mixtures from pure substances on chromatograms A mixture gives two or more spots; a pure substance gives a single spot. Chromatography
Describe the test for hydrogen Hold a burning splint at the open end of a test tube of the gas. Test for hydrogen
State the positive result for hydrogen Hydrogen burns rapidly with a pop sound (a 'squeaky pop'). Test for hydrogen
Describe the test for oxygen Put a glowing splint into a test tube of the gas. Test for oxygen
State the positive result for oxygen The glowing splint relights (bursts back into flame). Test for oxygen
Describe the test for carbon dioxide Bubble the gas through limewater, or shake the gas with limewater. Test for carbon dioxide
State the positive result for carbon dioxide The limewater turns milky (cloudy). Test for carbon dioxide
Describe the test for chlorine Put damp litmus paper into the gas. Test for chlorine
State the positive result for chlorine The litmus paper is bleached and turns white. Test for chlorine
Recall the five flame test colours Lithium crimson, sodium yellow, potassium lilac, calcium orange-red, copper green. Flame tests
Recall the coloured hydroxide precipitates Copper(II) blue, iron(II) green, iron(III) brown. Metal hydroxides
Describe the test for carbonate ions Add dilute acid and bubble any gas given off through limewater. Carbonates
State the positive result for carbonates Fizzing, and the gas turns limewater milky (cloudy). Carbonates
Recall the silver halide precipitate colours Chloride white, bromide cream, iodide yellow. Halides
Describe the test for sulfate ions Add dilute hydrochloric acid, then barium chloride solution, to a solution of the sample. Sulfates
State the positive result for sulfates A white precipitate of barium sulfate forms. Sulfates
State what instrumental methods are Methods that use instruments (machines) to detect and identify elements and compounds. Instrumental methods
Name the two main gases in air Nitrogen makes up about four-fifths of the air and oxygen about one-fifth. Proportions of gases in the atmosphere
Give the percentages of nitrogen and oxygen Air is about 80% nitrogen and about 20% oxygen. Proportions of gases in the atmosphere
Say where the early atmosphere came from Intense volcanic activity during the Earth's first billion years released the gases that formed it. The Earth's early atmosphere
Name the process that produced the oxygen Photosynthesis by algae and plants produced the oxygen that is now in the atmosphere. How oxygen increased
Name a process that removed carbon dioxide Photosynthesis by algae and plants took carbon dioxide out of the atmosphere. How carbon dioxide decreased
Name three greenhouse gases Water vapour, carbon dioxide and methane are greenhouse gases. Greenhouse gases
Give two activities that release carbon dioxide Burning fossil fuels and deforestation both increase the carbon dioxide in the atmosphere. Human activities that increase greenhouse gases
State a major cause of climate change An increase in the average global temperature is a major cause of climate change. Global climate change
Give one way to reduce carbon dioxide emissions For example, generate electricity from renewable resources instead of burning fossil fuels. The carbon footprint and its reduction
Name the products of complete combustion A fuel containing carbon and hydrogen forms carbon dioxide and water when it burns in plenty of oxygen. Atmospheric pollutants from fuels
State that carbon monoxide is toxic Carbon monoxide is a toxic (poisonous) gas. Properties and effects of atmospheric pollutants
State what humans use Earth's resources for To provide warmth, shelter, food and transport. Using resources and sustainable development
State what potable water means Potable water is water that is safe to drink. Potable water
Say why waste water must be treated Towns and industry produce large amounts of waste water that must be treated before it is released into the environment. Waste water treatment
Name the four stages of an LCA Extracting and processing raw materials, manufacturing and packaging, use and operation, and disposal. Life cycle assessment
State the three ways to reduce resource use Reduce use, reuse and recycle. Ways of reducing the use of resources
Define corrosion and give an example Corrosion is the destruction of materials by chemical reactions with substances in the environment; rusting is an example. Corrosion and its prevention
State what an alloy is A mixture of a metal with at least one other element, usually another metal (or carbon, in steel). Alloys as useful materials
Name the raw materials for soda-lime glass Sand, sodium carbonate and limestone, heated together. Ceramics, polymers and composites
State what the Haber process makes Ammonia, which is used to make nitrogen-based fertilisers. The Haber process
State why farmers use fertilisers Compounds of nitrogen, phosphorus and potassium improve agricultural productivity (bigger crop yields). Production and uses of NPK fertilisers

Grade 4

Grade 4 checklist (PDF)

Name compounds from their formulae A metal with a non-metal ends in -ide (NaCl, sodium chloride); -ate means oxygen is also present (CaCO₃, calcium carbonate). Atoms, elements and compounds
Count the atoms in a formula In Ca(OH)₂ the 2 multiplies everything in the brackets: 1 Ca, 2 O and 2 H, so 5 atoms. Atoms, elements and compounds
Match each separation method to its use Filtration, crystallisation, simple distillation, fractional distillation or chromatography, depending on what is in the mixture. Mixtures
Describe filtration and crystallisation Filter off an insoluble solid; evaporate some water from a solution, then cool it so crystals of the dissolved solid form. Mixtures
Describe the plum pudding model The atom is a ball of positive charge with negative electrons embedded in it. Development of the model of the atom
Put the atomic models in order Indivisible spheres → plum pudding → nuclear model → Bohr's shells → protons → neutrons (Chadwick). Development of the model of the atom
Define atomic number The number of protons in an atom; all atoms of the same element have the same number. Relative charges of subatomic particles
Define mass number The total number of protons and neutrons in an atom. Size and mass of atoms
Find protons, neutrons and electrons in atoms Protons = electrons = atomic number; neutrons = mass number − atomic number. Size and mass of atoms
Explain what relative atomic mass means An average value for the atoms of an element that takes account of the abundance of each isotope. Relative atomic mass
Write electronic structures of the first 20 elements Fill the innermost shells first, e.g. sodium (11 electrons) is 2,8,1. Electronic structure
Draw electron shell diagrams Draw circles for the shells around the nucleus and dots or crosses for the electrons. Electronic structure
State how the elements are ordered In order of atomic (proton) number. The periodic table
Describe problems with early tables They were incomplete, and strict atomic weight order put some elements in groups with the wrong properties. Development of the periodic table
Define metals by the ions they form Elements that react to form positive ions are metals; elements that do not are non-metals. Metals and non-metals
Compare physical properties of metals and non-metals Metals are usually shiny, high melting, malleable conductors; non-metals are usually dull, brittle, low melting and poor conductors. Metals and non-metals
State that noble gases are unreactive They are unreactive and do not easily form molecules; they exist as single atoms. Group 0
Describe their reactions with water They float and fizz; sodium melts into a ball; potassium's hydrogen burns with a lilac flame. Group 1
State the reactivity trend down Group 1 Reactivity increases going down the group: lithium, then sodium, then potassium. Group 1
Describe the trends down Group 7 Relative molecular mass, melting point and boiling point increase; reactivity decreases. Group 7
Compare melting points and densities Transition metals have much higher melting points and densities than Group 1 metals. Comparison with Group 1 elements
Compare strength and hardness Transition metals are strong and hard; Group 1 metals are soft enough to cut with a knife. Comparison with Group 1 elements
State three typical transition metal properties Ions with different charges, coloured compounds, and use as catalysts. Typical properties
Choose the bond type from the elements Metal + non-metal gives ionic bonding; non-metals only give covalent bonding; metals (elements or alloys) have metallic bonding. Chemical bonds
Name the particles in each type of bond Ionic: oppositely charged ions; covalent: atoms sharing pairs of electrons; metallic: positive ions and delocalised electrons. Chemical bonds
Work out the charge from the group number Group 1 → 1+, Group 2 → 2+, Group 6 → 2−, Group 7 → 1−. Ionic bonding
Name the force holding the ions together Strong electrostatic forces of attraction between oppositely charged ions. Ionic compounds
Recognise small molecules from their formulae Formulae of a few non-metal atoms, e.g. H₂, Cl₂, HCl, H₂O, NH₃, CH₄. Covalent bonding
Describe what delocalised electrons are The outer shell electrons of metal atoms are delocalised: free to move through the whole structure. Metallic bonding
Describe particles in solids, liquids and gases Regular and vibrating in solids, close but moving around in liquids, far apart and moving quickly in gases. The three states of matter
Know that aqueous means dissolved in water Any solution in water, such as sodium chloride solution or dilute hydrochloric acid, is (aq). State symbols
State when ionic compounds conduct electricity When melted (molten) or dissolved in water, but not when solid. Properties of ionic compounds
State that small molecules don't conduct electricity The molecules have no overall electric charge. Properties of small molecules
State that polymers are solids at room temperature The forces between their large molecules are strong enough to keep them solid. Polymers
State that they have very high melting points They are solids at room temperature with very high melting and boiling points. Giant covalent structures
Define an alloy A mixture of a metal with other elements, usually other metals. Properties of metals and alloys
Name the particles that carry the charge Delocalised electrons carry electrical charge through the metal. Metals as conductors
Describe the bonding in diamond Each carbon atom forms four covalent bonds with other carbon atoms. Diamond
State the number of bonds each carbon forms Each carbon atom forms three covalent bonds with three other carbon atoms. Graphite
Describe fullerenes as hollow carbon molecules They are based on hexagonal rings of carbon atoms, sometimes with rings of five or seven atoms. Graphene and fullerenes
Name buckminsterfullerene and give its formula The first fullerene discovered: C₆₀, a hollow, spherical molecule. Graphene and fullerenes
State the size range of nanoparticles Nanoparticles are 1 to 100 nm across and contain a few hundred atoms. Sizes of particles and their properties
List the main uses of nanoparticles Medicine, electronics, cosmetics and sun creams, deodorants, and catalysts. Uses of nanoparticles
Calculate a missing mass in a reaction Total mass of reactants = total mass of products, so subtract the known masses from the total. Conservation of mass and balanced equations
Calculate M_r of a simple formula Add the A_r of every atom, e.g. CO₂ = 12 + (2 × 16) = 44. Relative formula mass
Explain a mass decrease when gas forms The gas produced escapes into the air, so its mass is no longer measured. Mass changes when a reactant or product is a gas
Find the range of a set of results Range = highest value − lowest value. Chemical measurements and uncertainty
Convert cm³ to dm³ Divide by 1000, e.g. 250 cm³ = 0.250 dm³. Concentration of solutions
State what is meant by yield The amount of product obtained from a reaction. Percentage yield
State what atom economy measures The amount of the starting materials that ends up as the useful (desired) product. Atom economy
Define oxidation and reduction using oxygen Oxidation is the gain of oxygen and reduction is the loss of oxygen. Metal oxides
Describe metal reactions with water and acid Say what you would see, e.g. fizzing, and name the products (metal hydroxide or salt, plus hydrogen). The reactivity series
State which metals carbon can extract Metals below carbon in the reactivity series, such as zinc, iron and copper, can be extracted by heating their oxides with carbon. Extraction of metals and reduction
Name the salt from the acid used Hydrochloric acid gives chlorides and sulfuric acid gives sulfates. Reactions of acids with metals
Describe the test for hydrogen A lit splint held at the mouth of the tube burns the hydrogen with a squeaky pop. Reactions of acids with metals
Predict products of acid with base or alkali acid + metal oxide or metal hydroxide → salt + water. Neutralisation and salt production
Predict products of acid with a carbonate acid + metal carbonate → salt + water + carbon dioxide. Neutralisation and salt production
Describe filtering and crystallising Filter off the excess solid, heat to evaporate some water, then leave the solution to cool and crystallise. Soluble salts
Measure pH with universal indicator Add the indicator and match its colour to a pH colour chart. The pH scale and neutralisation
Name the equipment used in a titration Burette, pipette with pipette filler, conical flask, indicator and white tile. Titrations
Define an electrolyte A molten or dissolved ionic compound, which conducts electricity because its ions are free to move. The process of electrolysis
Predict products for any molten binary compound The metal forms at the cathode and the non-metal forms at the anode. Electrolysis of molten ionic compounds
State when electrolysis is used for extraction When the metal is too reactive to be extracted by carbon, or when it reacts with carbon. Using electrolysis to extract metals
State the products from sodium chloride solution Hydrogen forms at the cathode and chlorine at the anode, leaving sodium hydroxide solution. Electrolysis of aqueous solutions
Give examples of endothermic reactions Thermal decomposition and the reaction of citric acid with sodium hydrogencarbonate are endothermic. Energy transfer in exothermic and endothermic reactions
Classify a reaction from temperature data A rise in temperature means the reaction is exothermic; a fall means it is endothermic. Energy transfer in exothermic and endothermic reactions
Label the parts of a reaction profile Identify the reactants, products, activation energy and overall energy change on a given profile. Reaction profiles
Identify exothermic or endothermic from a profile Products lower than reactants means exothermic; products higher than reactants means endothermic. Reaction profiles
Describe how to make a simple cell Put two different metals in contact with an electrolyte and connect them with wires, e.g. to a voltmeter. Cells and batteries
Explain how a battery gives a bigger voltage A battery is two or more cells connected in series, so their voltages add together. Cells and batteries
State the product of a hydrogen fuel cell The only product is water: 2H₂ + O₂ → 2H₂O. Fuel cells
Calculate a mean rate of reaction Divide the quantity of reactant used or product formed by the time taken, e.g. 30 cm³ ÷ 20 s = 1.5 cm³/s. Calculating rates of reactions
Give the correct unit for rate Use g/s for a mass or cm³/s for a volume of gas, matching the units of the quantity and the time. Calculating rates of reactions
State how each factor changes the rate Increasing the concentration, pressure, surface area or temperature, or adding a catalyst, increases the rate. Factors which affect the rate of reaction
Define activation energy The minimum amount of energy that particles must have to react. Collision theory and activation energy
Know that enzymes are biological catalysts Enzymes are the catalysts that speed up reactions in living things. Catalysts
Describe what a reversible reaction is The products of the reaction can react to produce the original reactants. Reversible reactions
Know the reverse has the opposite energy change If the forward reaction is exothermic, the reverse reaction is endothermic, and the other way round. Energy changes and reversible reactions
Define a closed system Apparatus that prevents the reactants and products from escaping. Equilibrium
Name the first four alkanes and their formulae Methane CH₄, ethane C₂H₆, propane C₃H₈ and butane C₄H₁₀. Crude oil, hydrocarbons and alkanes
Define a fraction A mixture of hydrocarbons with a similar number of carbon atoms, and so similar boiling points. Fractional distillation and petrochemicals
Recall how properties change with molecule size Bigger molecules: higher boiling point, higher viscosity and lower flammability. Properties of hydrocarbons
Name the two types of product Cracking produces smaller alkanes and another type of hydrocarbon called alkenes. Cracking and alkenes
Describe the bromine water test for alkenes Shake with bromine water: it turns from orange to colourless if an alkene is present. Cracking and alkenes
Name the first four alkenes Ethene C₂H₄, propene C₃H₆, butene C₄H₈ and pentene C₅H₁₀. Structure and formulae of alkenes
Explain why alkenes burn with smoky flames In air they tend to undergo incomplete combustion, which produces soot (carbon). Reactions of alkenes
Name and draw the first four alcohols Methanol, ethanol, propanol and butanol; ethanol is CH₃CH₂OH. Alcohols
State the conditions for fermentation Sugar solution, yeast, a warm temperature (about 30 °C) and no air. Alcohols
Name the first four carboxylic acids Methanoic, ethanoic, propanoic and butanoic acid; ethanoic acid is CH₃COOH. Carboxylic acids
Name the polymer made from an alkene Write 'poly' and put the monomer's name in brackets: propene gives poly(propene). Addition polymerisation
Name the monomers of each natural polymer DNA: nucleotides. Proteins: amino acids. Starch and cellulose: sugars (glucose). DNA and other naturally occurring polymers
State what DNA does DNA encodes the genetic instructions for the development and functioning of living organisms and viruses. DNA and other naturally occurring polymers
Contrast everyday and chemical meanings of pure In everyday language 'pure' means nothing has been added (e.g. pure milk), but the product can still be a mixture. Pure substances
Use melting points to spot pure samples A pure substance melts at a specific temperature; a mixture melts over a range of temperatures. Pure substances
Identify formulations from given information Look for a product made by mixing measured amounts of components, each with a purpose. Formulations
Match spots to known reference substances Spots of the same substance travel the same distance on the same paper with the same solvent. Chromatography
Describe the paper chromatography method Pencil start line, small spots, solvent below the line, lid on, mark the solvent front. Chromatography
Tell the hydrogen and oxygen tests apart Hydrogen: burning splint gives a pop. Oxygen: glowing splint relights. Test for hydrogen
Tell the oxygen and hydrogen tests apart Oxygen: glowing splint relights. Hydrogen: burning splint gives a pop. Test for oxygen
Name the chemical in limewater Limewater is an aqueous solution of calcium hydroxide. Test for carbon dioxide
Give a safety precaution for chlorine Chlorine is toxic, so use small amounts in a fume cupboard. Test for chlorine
Identify a metal ion from a flame colour Match the colour seen to the metal ion that gives it. Flame tests
Describe how to carry out a flame test Dip a clean wire loop in the sample and hold it in a hot, blue Bunsen flame. Flame tests
Name the ions giving white precipitates Aluminium, calcium and magnesium ions all give a white precipitate. Metal hydroxides
Describe the test for halide ions Add dilute nitric acid, then silver nitrate solution, to a solution of the sample. Halides
State the three advantages of instrumental methods They are accurate, sensitive and rapid. Instrumental methods
Describe how flame emission spectroscopy works The sample is put into a flame and the light given out passes through a spectroscope, giving a line spectrum. Flame emission spectroscopy
Name the gases present in small proportions Carbon dioxide, water vapour and noble gases such as argon are present in small proportions. Proportions of gases in the atmosphere
State how long the proportions have been similar The proportions of the gases have been much the same as today for about 200 million years. Proportions of gases in the atmosphere
Name the main gas in the early atmosphere It was mainly carbon dioxide with little or no oxygen, like the atmospheres of Mars and Venus today. The Earth's early atmosphere
Describe how the oceans formed Water vapour released by volcanoes condensed to form the oceans. The Earth's early atmosphere
Write the word equation for photosynthesis carbon dioxide + water → glucose + oxygen, with light as the energy source. How oxygen increased
Name rocks and fuels that lock up carbon Sedimentary rocks such as limestone, and the fossil fuels coal, crude oil and natural gas. How carbon dioxide decreased
State why greenhouse gases are needed They keep the Earth's temperature high enough to support life. Greenhouse gases
Give two activities that release methane Farming cattle and other livestock, growing rice in paddy fields, and waste decaying in landfill sites. Human activities that increase greenhouse gases
Give potential effects of climate change For example, rising sea levels, more frequent and severe storms, and changes in rainfall. Global climate change
Define carbon footprint The total amount of carbon dioxide and other greenhouse gases emitted over the full life cycle of a product, service or event. The carbon footprint and its reduction
Name the pollutants released by burning fuels Carbon dioxide, water vapour, carbon monoxide, sulfur dioxide, oxides of nitrogen and particulates. Atmospheric pollutants from fuels
Name the gases that cause acid rain Sulfur dioxide and oxides of nitrogen cause acid rain. Properties and effects of atmospheric pollutants
Give effects of acid rain It damages buildings and statues, harms trees and plants, and makes lakes acidic, harming aquatic life. Properties and effects of atmospheric pollutants
Define finite and renewable resources A finite resource is used faster than it forms, so it will run out; a renewable resource is replaced as fast as it is used. Using resources and sustainable development
Give natural products replaced by synthetic ones For example, wool and cotton supplemented by synthetic fibres such as polyester, or natural rubber by synthetic rubber. Using resources and sustainable development
Explain why potable water is not pure It contains dissolved substances, whereas pure water contains only water molecules. Potable water
List the steps in treating fresh water Choose a suitable source of fresh water, pass it through filter beds, then sterilise it. Potable water
State what sewage treatment must remove Sewage and agricultural waste water need organic matter and harmful microbes removed. Waste water treatment
State what an LCA is used for To assess the environmental impact of a product over its whole life, including transport at each stage. Life cycle assessment
Name materials made from limited raw materials Metals, glass, building materials, clay ceramics and most plastics. Ways of reducing the use of resources
State the conditions needed for rusting Iron needs both air (oxygen) and water to rust. Corrosion and its prevention
Name barrier methods of preventing corrosion Greasing, painting and electroplating keep air and water away from the metal. Corrosion and its prevention
Name the metals in bronze and brass Bronze is copper and tin; brass is copper and zinc. Alloys as useful materials
Recall a use for each named alloy For example, bronze for statues, brass for musical instruments, stainless steel for cutlery, aluminium alloys for aircraft. Alloys as useful materials
Describe how clay ceramics are made Wet clay is shaped and then heated in a furnace, e.g. to make pottery and bricks. Ceramics, polymers and composites
Name the two parts of a composite A matrix (binder) surrounds and binds together fibres or fragments of the reinforcement. Ceramics, polymers and composites
Give the sources of nitrogen and hydrogen Nitrogen comes from the air; hydrogen comes from natural gas (methane). The Haber process
Recall the conditions used An iron catalyst, about 450 °C and about 200 atmospheres. The Haber process
Explain what an NPK fertiliser is A formulation of salts containing appropriate percentages of nitrogen, phosphorus and potassium. Production and uses of NPK fertilisers
Name the sources of the three elements Nitrogen from ammonia; potassium from mined potassium chloride and potassium sulfate; phosphorus from mined phosphate rock. Production and uses of NPK fertilisers

Grade 5

Grade 5 checklist (PDF)

Write word equations for reactions Reactants on the left, products on the right, joined by an arrow, e.g. magnesium + oxygen → magnesium oxide. Atoms, elements and compounds
Describe simple distillation The solvent boils off, is cooled and condensed in a condenser, and is collected, leaving the dissolved solid behind. Mixtures
Describe fractional distillation Liquids with different boiling points are separated using a fractionating column; the lowest boiling point liquid is collected first. Mixtures
Describe the nuclear model A tiny, positively charged nucleus that contains almost all the mass, with electrons outside it. Development of the model of the atom
Explain why atoms have no overall charge The number of electrons equals the number of protons, so the charges cancel. Relative charges of subatomic particles
Identify an element from its proton number Different elements have different numbers of protons, so the atomic number tells you the element. Relative charges of subatomic particles
Define isotopes Atoms of the same element with different numbers of neutrons: same atomic number, different mass number. Size and mass of atoms
Recall the sizes of atoms and nuclei Atom radius about 0.1 nm (1 × 10⁻¹⁰ m); nucleus radius less than 1/10 000 of that (about 1 × 10⁻¹⁴ m). Size and mass of atoms
Calculate A_r for two isotopes Multiply each mass number by its percentage abundance, add the results, then divide by 100. Relative atomic mass
Explain why A_r is rarely whole It is an average of isotopes with different mass numbers, weighted by how common each one is. Relative atomic mass
Link electronic structure to group and period Outer electrons = group number (Groups 1 to 7); number of occupied shells = period. Electronic structure
Identify an element from its electronic structure Add up the electrons to get the atomic number, e.g. 2,8,5 is 15, which is phosphorus. Electronic structure
Explain why group members react similarly They have the same number of electrons in their outer shell. The periodic table
Link position to electronic structure Group number = outer electrons (Groups 1 to 7); period = number of occupied shells. The periodic table
Describe what Mendeleev did He left gaps for undiscovered elements and changed the order of some elements to fit their properties. Development of the periodic table
Compare the oxides of metals and non-metals Metal oxides are basic; non-metal oxides are usually acidic. Metals and non-metals
Explain their unreactivity using electrons Their atoms have stable arrangements of electrons: a full outer shell of 8 (2 for helium). Group 0
Describe the boiling point trend Boiling points increase going down the group, as relative atomic mass increases. Group 0
Write equations for their reactions With water, oxygen and chlorine, e.g. 2K + 2H₂O → 2KOH + H₂. Group 1
Describe the compounds halogens form Ionic compounds with metals (halide ions, 1−); covalent molecules with non-metals, e.g. HCl. Group 7
Predict whether a displacement reaction happens A more reactive halogen displaces a less reactive one from a solution of its salt. Group 7
Compare reactivity with oxygen, water and halogens Group 1 metals react vigorously; transition metals react slowly, or not at all, at room temperature. Comparison with Group 1 elements
Give examples of coloured compounds e.g. copper(II) sulfate solution is blue; iron(III) compounds are orange-brown. Typical properties
Give examples of transition metal catalysts Iron in the Haber process; nickel for adding hydrogen to alkenes; manganese(IV) oxide for decomposing hydrogen peroxide. Typical properties
Decide the bond type from a formula Use the periodic table to decide whether each element in, for example, KBr, SO₂ or Zn is a metal or a non-metal. Chemical bonds
Write the electronic structure of an ion For example, Na (2,8,1) becomes Na⁺ (2,8) and O (2,6) becomes O²⁻ (2,8). Ionic bonding
Draw dot and cross diagrams for ionic compounds Show each ion in square brackets with its charge, using dots and crosses for electrons from different atoms. Ionic bonding
State that the forces act in all directions Each ion is attracted to all the oppositely charged ions around it, throughout the lattice. Ionic compounds
Recognise an ionic structure from a diagram Look for a regular lattice of two types of ion, often labelled with charges, and no separate molecules. Ionic compounds
Draw dot and cross diagrams for simple molecules Hydrogen, chlorine, hydrogen chloride, water, ammonia and methane, which have single bonds only. Covalent bonding
Describe the metallic bond Strong electrostatic attraction between positive metal ions and the shared delocalised electrons. Metallic bonding
Predict states from melting and boiling points Below the melting point: solid; between the melting and boiling points: liquid; above the boiling point: gas. The three states of matter
Add state symbols to a simple equation For example, Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g). State symbols
Explain high melting points using electrostatic forces A lot of energy is needed to overcome the many strong electrostatic forces between oppositely charged ions. Properties of ionic compounds
Explain low boiling points using intermolecular forces Only weak intermolecular forces must be overcome, which needs little energy. Properties of small molecules
Describe the bonds in a polymer chain The atoms in each chain are linked to each other by strong covalent bonds. Polymers
Recognise a polymer from a diagram Look for a repeating unit in brackets with bonds through the brackets and n after them. Polymers
Explain their very high melting points Many strong covalent bonds must be broken, which needs a lot of energy. Giant covalent structures
Explain why most metals have high melting points Giant structures with strong metallic bonding need a lot of energy to overcome. Properties of metals and alloys
Explain why pure metals are soft The layers of atoms can slide over each other easily. Properties of metals and alloys
Explain electrical conduction in metals The delocalised electrons are free to move through the structure, carrying charge. Metals as conductors
Explain why diamond is very hard Many strong covalent bonds hold the atoms rigidly in a 3D network. Diamond
Explain diamond's very high melting point A lot of energy is needed to break the many strong covalent bonds. Diamond
Explain why graphite is soft and slippery There are no covalent bonds between the layers, so the layers can slide over each other. Graphite
Give uses of fullerenes and nanotubes Fullerenes: drug delivery, lubricants, catalysts; nanotubes: nanotechnology, electronics and strengthening materials. Graphene and fullerenes
Convert between nanometres and metres 1 nm = 1 × 10⁻⁹ m, so 50 nm = 5 × 10⁻⁸ m. Sizes of particles and their properties
Classify particles as nano, fine or coarse Nano 1–100 nm; fine (PM_2.5) 100–2500 nm; coarse (PM₁₀) 2500–10 000 nm. Sizes of particles and their properties
Give advantages of using nanoparticles Their high surface area to volume ratio makes them effective, so smaller quantities are needed. Uses of nanoparticles
Balance a symbol equation Change only the big numbers in front of formulae until every element has the same number of atoms on both sides. Conservation of mass and balanced equations
Calculate M_r for formulae with brackets Multiply everything inside the bracket, e.g. Ca(OH)₂ = 40 + 2 × (16 + 1) = 74. Relative formula mass
Explain why a heated metal gains mass Oxygen atoms from the air combine with the metal, so the metal oxide has a greater mass than the metal. Mass changes when a reactant or product is a gas
Identify an anomalous result A result that clearly does not fit the pattern of the others; leave it out of the mean. Chemical measurements and uncertainty
Calculate concentration in g/dm³ Concentration = mass of solute (g) ÷ volume of solution (dm³). Concentration of solutions
Calculate percentage yield % yield = mass of product actually made ÷ maximum theoretical mass × 100. Percentage yield
Give reasons why yield is below 100% The reaction is reversible, some product is lost when it is separated, or some reactants react in unexpected ways. Percentage yield
Calculate atom economy from an equation Atom economy = M_r of desired product ÷ sum of M_r of all reactants × 100. Atom economy
Identify what is oxidised or reduced In an equation, the substance that gains oxygen is oxidised and the substance that loses oxygen is reduced. Metal oxides
Deduce an order of reactivity from results The more vigorous the reaction (more bubbles, bigger temperature rise), the more reactive the metal. The reactivity series
Predict whether a displacement reaction happens A metal displaces another metal from its compound only if it is more reactive. The reactivity series
Write equations for reduction with carbon For example, 2ZnO + C → 2Zn + CO₂. Extraction of metals and reduction
Tell the difference between bases and alkalis A base neutralises an acid; an alkali is a base that dissolves in water, such as a soluble metal hydroxide. Neutralisation and salt production
Explain why excess solid is added Excess solid makes sure all the acid reacts, so the salt solution contains no leftover acid. Soluble salts
Tell when the solid is in excess Solid stays undissolved at the bottom, and for a carbonate the fizzing stops. Soluble salts
Name the ions in acids and alkalis Acids produce H⁺ ions in water and alkalis produce OH⁻ ions in water. The pH scale and neutralisation
Compare universal indicator with a pH probe A pH probe gives a more precise numerical reading; universal indicator gives an approximate value from a colour. The pH scale and neutralisation
Describe the pH change during neutralisation As alkali is added to acid, the pH rises to 7 at neutralisation and above 7 when alkali is in excess. The pH scale and neutralisation
Describe how to carry out a titration Add acid from a burette to a measured volume of alkali with indicator until the colour just changes. Titrations
Choose concordant results and find the mean Use titres within 0.10 cm³ of each other and leave out the rough titre. Titrations
Predict which electrode each ion moves to Positive ions move to the cathode and negative ions move to the anode. The process of electrolysis
Describe what is seen at each electrode For lead bromide, silvery lead forms at the cathode and brown bromine vapour at the anode. Electrolysis of molten ionic compounds
Name the products from aluminium oxide Aluminium forms at the cathode and oxygen forms at the anode. Using electrolysis to extract metals
Explain why electrolysis is expensive Large amounts of energy are needed to melt the compound and to produce the electric current. Using electrolysis to extract metals
Predict the product at the cathode Hydrogen forms if the metal is more reactive than hydrogen; otherwise the metal forms. Electrolysis of aqueous solutions
Predict the product at the anode The halogen forms if halide ions are present; otherwise oxygen forms. Electrolysis of aqueous solutions
Explain energy conservation in a reaction Energy is not created or destroyed, so in an exothermic reaction the products have less energy than the reactants by the amount transferred. Energy transfer in exothermic and endothermic reactions
Read energy values from a profile Activation energy = peak − reactants; overall energy change = products − reactants. Reaction profiles
Explain why non-rechargeable cells stop working The chemical reactions stop when one of the reactants has been used up. Cells and batteries
Describe how fuel cells produce a potential difference The fuel is oxidised electrochemically inside the cell, which produces a potential difference. Fuel cells
Interpret graphs of product formed against time The steeper the curve, the faster the rate; a horizontal line means the reaction has stopped. Calculating rates of reactions
Describe the gas-volume method for rates Measure the volume of gas given off at regular time intervals with a gas syringe or an upturned measuring cylinder. Factors which affect the rate of reaction
Describe the disappearing-cross (turbidity) method Time how long it takes for a cross under the flask to disappear as a cloudy precipitate forms. Factors which affect the rate of reaction
Explain concentration and pressure effects There are more particles in the same volume, so collisions are more frequent. Collision theory and activation energy
Explain the effect of surface area More particles of the solid are exposed to the other reactant, so collisions are more frequent. Collision theory and activation energy
Identify a catalyst from experimental data It speeds up the reaction, has the same mass at the end and is not in the equation. Catalysts
Describe heating and cooling ammonium chloride Heating decomposes it into ammonia and hydrogen chloride; cooling makes these gases re-form the white solid. Reversible reactions
Write equations for reversible reactions Use ⇌ in word and balanced symbol equations, e.g. NH₄Cl ⇌ NH₃ + HCl. Reversible reactions
State that equal energy is transferred each way The energy taken in one way is exactly the same as the energy given out the other way. Energy changes and reversible reactions
Describe the copper sulfate reversible reaction Heating blue hydrated copper sulfate gives white anhydrous copper sulfate and water; adding water reverses it. Energy changes and reversible reactions
State when equilibrium is reached When the forward and reverse reactions happen at exactly the same rate. Equilibrium
Know that amounts stay constant at equilibrium The concentrations of reactants and products stop changing, but they are not necessarily equal. Equilibrium
Use the general formula of the alkanes Recognise an alkane from its formula, and write the formula for any number of carbon atoms: 10 carbon atoms gives C₁₀H₂₂. Crude oil, hydrocarbons and alkanes
Draw the displayed formula of an alkane Show every atom and every bond: each carbon atom forms four single bonds and each hydrogen atom forms one. Crude oil, hydrocarbons and alkanes
Describe the temperature gradient in the column The fractionating column is hottest at the bottom and coolest at the top. Fractional distillation and petrochemicals
Explain what feedstock means A raw material used to make other chemicals; fractions are feedstock for the petrochemical industry. Fractional distillation and petrochemicals
Choose a suitable fuel using property data For example, a fuel for a gas stove must have a boiling point below room temperature and be very flammable. Properties of hydrocarbons
State what happens to the fuel in combustion The carbon and hydrogen in the fuel are oxidised, and energy is released. Properties of hydrocarbons
Describe the conditions for catalytic and steam cracking Catalytic: vapour passed over a hot catalyst. Steam: vapour mixed with steam and heated to a very high temperature. Cracking and alkenes
Give uses of the products of cracking Smaller alkanes are used as fuels such as petrol; alkenes are used to make polymers and many other chemicals. Cracking and alkenes
Use the general formula of the alkenes C_nH_2n: the number of hydrogen atoms is double the number of carbon atoms. Structure and formulae of alkenes
Explain what unsaturated means Alkenes have two fewer hydrogen atoms than the alkane with the same number of carbon atoms, because they contain a C=C. Structure and formulae of alkenes
Name the products of addition reactions Hydrogen gives an alkane, steam gives an alcohol, and a halogen gives a compound with two halogen atoms. Reactions of alkenes
Give the conditions for adding hydrogen and steam Hydrogen: a nickel catalyst. Steam: high temperature, high pressure and a catalyst (phosphoric acid). Reactions of alkenes
Describe the reactions of alcohols Sodium: fizzing (hydrogen). Water: dissolves, neutral solution. Oxidising agent: a carboxylic acid forms. Alcohols
Describe the reaction with carbonates Fizzing as carbon dioxide is given off; a salt and water also form, e.g. sodium ethanoate. Carboxylic acids
Describe how an ester is made Carboxylic acid + alcohol, with an acid catalyst, gives an ester and water: ethanoic acid + ethanol → ethyl ethanoate + water. Carboxylic acids
Recognise monomers that form addition polymers The monomer must contain a C=C double bond. Addition polymerisation
Describe the structure of DNA Two polymer chains, made from four different nucleotides, in the form of a double helix. DNA and other naturally occurring polymers
Describe how impurities change melting and boiling Impurities lower the melting point and widen the melting range; dissolved impurities raise the boiling point. Pure substances
Interpret heating curves to judge purity A flat section at a fixed temperature shows a pure substance changing state; a sloping section shows a mixture. Pure substances
Explain why quantities are carefully measured The right amount of each component gives the product the properties it needs. Formulations
Suggest the purpose of a component Use the information given, e.g. a pigment gives colour and a sweetener improves taste. Formulations
Calculate an Rf value Rf = distance moved by substance ÷ distance moved by solvent, both measured from the start line. Chromatography
Name reactions that produce hydrogen Metals with acids, very reactive metals with water, and electrolysis of many solutions (at the negative electrode). Test for hydrogen
Name reactions that produce oxygen Decomposition of hydrogen peroxide, electrolysis (at the positive electrode) and photosynthesis. Test for oxygen
Describe how to test gas from a reaction Use a delivery tube to bubble the gas from the reaction through limewater. Test for carbon dioxide
Name reactions that produce carbon dioxide Acids with metal carbonates, and complete combustion of fuels that contain carbon. Test for carbon dioxide
Name a reaction that produces chlorine Electrolysis of a chloride solution, such as sodium chloride solution: chlorine forms at the positive electrode. Test for chlorine
Explain why the wire loop must be clean Traces of other metal ions would add their own colour and give a false result. Flame tests
Use excess sodium hydroxide to identify aluminium Only aluminium hydroxide dissolves in excess sodium hydroxide solution. Metal hydroxides
Distinguish calcium ions from magnesium ions Both precipitates stay in excess, so do a flame test: calcium gives orange-red. Metal hydroxides
Name the products of acid and carbonate A salt, water and carbon dioxide. Carbonates
Identify a compound from two ion tests Combine the metal ion test result with the halide test result, e.g. lilac flame + yellow precipitate → potassium iodide. Halides
Identify a compound using two ion tests e.g. a blue precipitate with sodium hydroxide and a white precipitate with barium chloride → copper(II) sulfate. Sulfates
Explain accurate, sensitive and rapid Accurate: close to the true value. Sensitive: detects very small amounts. Rapid: gives results quickly. Instrumental methods
Identify a metal ion from reference spectra Each metal ion gives its own pattern of lines; the ion is present if all its lines appear. Flame emission spectroscopy
Calculate the volume of a gas in air Multiply the total volume by the percentage divided by 100, e.g. 20 ÷ 100 × 500 cm³ = 100 cm³ of oxygen. Proportions of gases in the atmosphere
Name other gases in the early atmosphere Nitrogen from volcanoes gradually built up, and there may have been small proportions of methane and ammonia. The Earth's early atmosphere
State when algae first produced oxygen Algae first produced oxygen about 2.7 billion years ago, and oxygen appeared in the atmosphere soon after. How oxygen increased
Describe how limestone formed Shells and skeletons of sea creatures, made of calcium carbonate, settled, were buried and were compressed into rock. How carbon dioxide decreased
Describe how coal formed Remains of plants, mainly trees, were buried and compressed over millions of years. How carbon dioxide decreased
Describe what happens to the Sun's radiation Short wavelength radiation from the Sun passes through the atmosphere and is absorbed by the Earth's surface, warming it. Greenhouse gases
Explain how deforestation increases carbon dioxide Burning the trees releases carbon dioxide, and fewer trees are left to remove it by photosynthesis. Human activities that increase greenhouse gases
Describe four different potential effects Give four clearly different effects, each with a short consequence, e.g. sea levels rise, causing flooding. Global climate change
Describe actions to reduce emissions Alternative energy, energy conservation, carbon capture and storage, carbon taxes and licences, carbon off-setting and carbon neutrality. The carbon footprint and its reduction
Describe how carbon monoxide and soot form Incomplete combustion in a limited supply of oxygen produces carbon monoxide and carbon particles (soot). Atmospheric pollutants from fuels
Describe how sulfur dioxide forms Sulfur in the fuel reacts with oxygen when the fuel burns. Atmospheric pollutants from fuels
Explain why carbon monoxide is hard to detect It is colourless and odourless, so you cannot see or smell it. Properties and effects of atmospheric pollutants
Define sustainable development Development that meets the needs of current generations without compromising the ability of future generations to meet their own needs. Using resources and sustainable development
Give reasons for each treatment step Filter beds remove solid particles; sterilising with chlorine, ozone or ultraviolet light kills microbes. Potable water
Name two methods of desalination Distillation, and processes that use membranes such as reverse osmosis. Potable water
List the stages of sewage treatment in order Screening and grit removal, sedimentation, anaerobic digestion of the sludge, and aerobic biological treatment of the effluent. Waste water treatment
Compare plastic and paper bags using an LCA Plastic bags come from crude oil (finite) but can be reused many times; paper bags come from trees but are reused less. Life cycle assessment
Describe how glass and metals are recycled Glass is crushed and melted to make different glass products; metals are melted and recast or reformed. Ways of reducing the use of resources
Explain why recycling saves resources and energy Less of a limited raw material is used, and usually less energy than extracting and processing new material. Ways of reducing the use of resources
Describe the rusting experiment with three tubes Nails in water open to air, in boiled water under oil, and in dry air; only the first rusts. Corrosion and its prevention
Compare high carbon, low carbon and stainless steel High carbon steel is strong but brittle; low carbon steel is softer and easily shaped; stainless steel is hard and resists corrosion. Alloys as useful materials
Give examples of composites and their parts Reinforced concrete is steel rods in concrete; fibreglass is glass fibres in a polymer resin. Ceramics, polymers and composites
Describe how ammonia is separated and gases recycled On cooling, ammonia liquefies and is removed; unreacted nitrogen and hydrogen are recycled. The Haber process
Name ammonium salts made from ammonia With nitric acid, ammonium nitrate; with sulfuric acid, ammonium sulfate; with phosphoric acid, ammonium phosphate. Production and uses of NPK fertilisers

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