4.1Atomic structure and the periodic table
4.1.1A simple model of the atom, symbols, relative atomic mass, electronic charge and isotopes
4.1.1.1 Atoms, elements and compounds
03.2 — Identifying an element from its relative atomic mass using the periodic table
4.1.1.2 Mixtures
03.2 — Suggesting a method to separate an insoluble product from a reaction mixture (WS 2.3 — selecting a separation technique)
03.3 — Suggesting an impurity removed by rinsing a solid product with water (WS 2.3 — identifying an impurity source)
03.4 — Suggesting why a solid product was warmed after purification (WS 2.4 — practical technique reasoning)
4.1.1.3 The development of the model of the atom
04.1 — Identifying the diagram showing the plum pudding model
04.2 — Identifying the diagram showing the model from the alpha particle scattering experiment
04.3 — Identifying the diagram showing the model resulting from Bohr's work
04.6 — Explaining how Chadwick's work led to a new understanding of atomic structure
04.4 — Explaining how alpha particle scattering evidence changed the plum pudding model
05.1 — Describing three differences between the nuclear model and the plum pudding model of the atom
05.2 — Describing the change Bohr made to the nuclear model of the atom
01.1 — Describing the plum pudding model of the atom
01.2 — Ordering electrons, neutrons and protons by date of discovery
03.1 — Naming two historical models of the atom shown in a diagram
03.2 — Comparing an early model of the atom with the model used today
05.1 — Describing the plum pudding model of the atom
05.2 — Giving two conclusions from the alpha particle scattering experiment
05.3 — Identifying the change Bohr made to the model of the atom
4.1.1.4 Relative electrical charges of subatomic particles
02.1 — Giving the numbers of protons, neutrons and electrons in an aluminium atom from its notation
4.1.1.5 Size and mass of atoms
02.2 — Giving the numbers of protons, neutrons and electrons in an iron atom from its atomic/mass number
04.4 — Defining the mass number of an atom
04.1 — Completing a table of subatomic particle properties
04.2 — Defining mass number
04.3 — Explaining why two isotopes of the same element have different mass numbers
02.3 — Giving the numbers of electrons and neutrons in a named gallium isotope
05.4 — Calculating how many times larger an atom's radius is than its nucleus's radius
4.1.1.6 Relative atomic mass
04.5 — Estimating the relative atomic mass of an element from isotope abundance data
01.3 — Identifying the modern name for 'atomic weight'
02.4 — Calculating the relative atomic mass of a metal from isotope abundance data
02.1 — Giving the meaning of 'isotopes' in terms of subatomic particles
02.2 — Calculating the relative atomic mass of gallium from isotope abundance data
05.5 — Calculating the relative atomic mass of neon from three isotopes' abundance data
01.5 — Calculating the relative atomic mass of an element from isotope abundance data
03.3 — Defining the term 'isotopes', referring to subatomic particles
01.6 — Explaining what is meant by 'isotopes' in terms of subatomic particles
01.7 — Estimating the relative atomic mass of an element from isotope abundance data
4.1.1.7 Electronic structure
01.4 — Completing the electronic structure diagram for an aluminium atom [Figure — print if needed]
01.3 — Completing the electronic structure diagram for a sodium atom [Figure — print if needed]
4.1.2The periodic table
4.1.2.1 The periodic table
01.4 — Completing a sentence on how elements are ordered in the modern periodic table
01.3 — Predicting the number of outer shell electrons in an atom of a given element, with a reason
02.1 — Giving a similarity and a difference between the electronic structures of two Group 1 elements
02.2 — Explaining why aluminium is positioned in Group 3 of the periodic table
4.1.2.2 Development of the periodic table
01.1 — Explaining why Mendeleev reversed the order of two elements in his periodic table
01.2 — Explaining why Mendeleev's periodic table became more widely accepted than earlier versions
05.3 — Explaining why a suggested reason for Mendeleev reversing element order cannot be correct
05.4 — Giving the correct reason why Mendeleev reversed the order of some element pairs
02.5 — Giving two reasons why gallium's discovery helped Mendeleev's periodic table become accepted
01.4 — Suggesting why a newly identified element was not accepted by other scientists until several years later
01.4 — Explaining why Mendeleev did not place two elements in atomic-weight order
4.1.2.3 Metals and non-metals
01.1 — Identifying which section of the periodic table is likely to contain a described element
08.3 — Predicting and explaining whether an element reacts with metals, based on its position in the periodic table
01.1 — Describing where non-metallic elements are found in the periodic table
4.1.2.4 Group 0
05.4 — Identifying the correct statement about the noble gases
08.1 — Explaining why argon does not form compounds, in terms of a stable outer electron shell
02.4 — Plotting density data for Group 0 elements on a graph [Figure — print if needed] (MS 4c — plotting data on a graph)
02.5 — Estimating a Group 0 element's density from a graph
01.5 — Explaining why argon is unreactive
4.1.2.5 Group 1
08.2 — Giving a reason why water contact with sodium metal would be hazardous
01.6 — Describing the observations when sodium reacts with chlorine
04.1 — Giving two observations when potassium is added to water
04.2 — Completing and balancing the equation for potassium reacting with water
04.3 — Explaining why reactivity changes going down Group 1
05.1 — Predicting an observation showing that rubidium is more reactive than potassium
05.2 — Explaining why rubidium is more reactive than potassium
05.3 — Completing and balancing the equation for rubidium reacting with water
02.2 — Giving two observations when potassium reacts with water
02.3 — Identifying and explaining the colour of universal indicator in potassium hydroxide solution
4.1.2.6 Group 7
07.1 — Naming the products of chlorine solution reacting with potassium iodide solution
07.2 — Explaining why chlorine is more reactive than iodine
01.5 — Predicting the formula and physical state of an astatine molecule
08.2 — Explaining the trend in boiling points of the halogens shown in a data table
08.4 — Giving a reason why a halogen-reaction experiment should be done in a fume cupboard
08.5 — Explaining why halogen reactivity decreases going down the group
02.6 — Identifying which combination of solutions would produce a halogen displacement reaction
02.7 — Identifying the correct trend in relative molecular mass and boiling point going down Group 7
04.1 — Predicting the boiling point of chlorine from Group 7 trend data
04.2 — Identifying the state of fluorine at a given temperature from Group 7 data
04.3 — Identifying the number of outer shell electrons in an astatine atom
04.4 — Predicting the formula of a compound formed between hydrogen and astatine
04.5 — Explaining the trend in reactivity of the halogens going down Group 7
04.6 — Writing a balanced equation for a halogen displacement reaction
01.2 — Predicting the boiling point of bromine from Group 7 trend data
01.3 — Identifying the colour of the final solution in a halogen displacement reaction
01.4 — Identifying the ionic equation for chlorine reacting with potassium iodide
4.1.3Properties of transition metals (chemistry only)
4.1.3.1 Comparison with Group 1 elements
02.3 — Giving two property differences between a transition metal and a Group 1 metal
01.3 — Giving two physical property differences between Group 1 and transition elements
03.4 — Giving two property differences between a transition metal and a Group 1 metal
02.3 — Comparing the chemical and physical properties of transition elements and Group 1 elements using data
4.1.3.2 Typical properties
05.4 — Identifying the metal oxide most likely to catalyse a given reaction
02.1 — Identifying two typical properties of transition metals
01.2 — Identifying which section of the periodic table is likely to contain an element forming ions of two different charges
07.4 — Identifying the ratio of ions in an iron oxide with mixed oxidation states
4.2Bonding, structure, and the properties of matter
4.2.1Chemical bonds, ionic, covalent and metallic
4.2.1.1 Chemical bonds
01.6 — Naming the type of bonding between metals and non-metals
09.2 — Explaining how a covalent bond holds two atoms together
4.2.1.2 Ionic bonding
02.4 — Identifying the most likely formula of a gallium ion
04.4 — Drawing a dot and cross diagram for sodium and oxygen atoms reacting to form sodium oxide [Figure — print if needed]
01.7 — Describing what happens when a magnesium atom reacts with an oxygen atom, in terms of electrons
05.7 — Identifying which element has atoms with the same electronic structure as the chloride ion
05.1 — Describing what happens to calcium and chlorine atoms when calcium chloride forms
02.2 — Describing what happens when a calcium atom reacts with a sulfur atom, in terms of electrons and ions
03.1 — Describing the electron transfer and ion formulae when potassium reacts with sulfur
4.2.1.3 Ionic compounds
02.1 — Determining the formula of iron pyrites from a structure diagram
03.2 — Giving a limitation of the ball and stick model for representing an ionic structure
4.2.1.4 Covalent bonding
05.1 — Completing a dot and cross diagram for ammonia, outer shell electrons only [Figure — print if needed]
05.2 — Giving a limitation of using a dot and cross diagram to represent ammonia
01.2 — Comparing the structure and bonding of a covalent molecule, an ionic compound and a giant covalent structure
01.3 — Completing a dot and cross diagram for propanone using single bond lines [Figure — print if needed]
01.4 — Determining the molecular formula of propanone from a diagram
07.5 — Completing a dot and cross diagram to show outer shell electrons in Si2H6 [Figure — print if needed]
04.4 — Drawing a dot and cross diagram for hydrogen chloride, outer shell electrons only [Figure — print if needed]
08.2 — Predicting the formula of a compound formed between phosphorus and hydrogen
07.2 — Completing a dot and cross diagram for hydrazine, outer shell electrons only [Figure — print if needed]
07.3 — Giving an advantage of a ball and stick model over a dot and cross diagram
03.3 — Completing a dot and cross diagram for hydrogen sulfide, outer shell electrons only [Figure — print if needed]
4.2.2How bonding and structure are related to the properties of substances
4.2.2.1 The three states of matter
07.4 — Giving two limitations of a simple particle model for a gas
08.1 — Identifying the states of bromine at two given temperatures
08.3 — Explaining why it is incorrect to describe a single molecule as having a boiling point
4.2.2.2 State symbols
08.1 — Identifying what a given state symbol represents
4.2.2.3 Properties of ionic compounds
04.6 — Explaining why sodium oxide has a high melting point
05.2 — Giving a reason why solid calcium chloride cannot be electrolysed
06.1 — Explaining why aluminium oxide has a very high melting point
01.5 — Explaining why potassium iodide solution conducts electricity
03.6 — Matching properties of an ionic compound to their explanations [Figure — print if needed]
4.2.2.4 Properties of small molecules
07.3 — Explaining why hydrogen chloride is a gas at room temperature, in terms of structure and bonding
08.4 — Explaining why titanium chloride would not be expected to be a liquid at room temperature
05.3 — Explaining why ammonia has a low boiling point, in terms of structure and bonding
01.1 — Identifying two substances that have intermolecular forces between their particles
01.5 — Identifying why propanone has a low boiling point
04.5 — Explaining why methane is a gas and poly(ethene) is a solid at room temperature
08.1 — Explaining why propane has a low boiling point
02.4 — Explaining why substances made of small molecules do not conduct electricity
01.1 — Explaining why iodine has a higher boiling point than chlorine
03.5 — Matching properties of a covalent compound to their explanations [Figure — print if needed]
4.2.2.5 Polymers
02.5 — Explaining why poly(ethene) has a higher melting point than methane
4.2.2.7 Properties of metals and alloys
06.2 — Explaining why alloys are harder than pure metals
06.3 — Suggesting why alloys do not conduct electricity as well as pure metals, in terms of structure and bonding
07.2 — Explaining why alloying iron with other metals makes it harder than pure iron
4.2.2.8 Metals as conductors
09.3 — Sketching a comparison line on a graph for a metal container repeat of the investigation, with explanation [Figure — print if needed] (WS 3.1 — sketching a comparative result)
06.1 — Identifying a substance from melting point, boiling point and conductivity data
01.5 — Describing how metals conduct electricity, in terms of electrons
06.1 — Evaluating the use of three metals for electrical wiring using given data (WS 1.4 — evaluating materials using data)
06.2 — Describing how metals conduct electricity
07.1 — Describing how iron conducts thermal energy
01.2 — Identifying which property of gallium is not typical of most metals
4.2.3Structure and bonding of carbon
4.2.3.1 Diamond
03.1 — Describing the structure and bonding of diamond
03.2 — Explaining why diamond has a very high melting point
4.2.3.2 Graphite
06.2 — Explaining why graphite conducts electricity in terms of its structure and bonding
01.6 — Explaining why graphite is a good electrical conductor and is soft and slippery, in terms of structure and bonding
02.3 — Describing the structure and bonding of graphite
4.2.3.3 Graphene and fullerenes
03.1 — Explaining why carbon nanotubes conduct electricity
03.2 — Evaluating materials for use as badminton racket frames using given property data (WS 1.4 — evaluating applications using given data)
01.1 — Identifying the shape of a Buckminsterfullerene molecule
01.2 — Giving one use of a fullerene
03.3 — Naming a type of molecule shown in a diagram (fullerene)
03.4 — Suggesting why a fullerene molecule is suitable for moving drugs around the body
4.2.4Bulk and surface properties of matter including nanoparticles (chemistry only)
4.2.4.1 Sizes of particles and their properties
03.3 — Calculating the surface area of a zinc oxide nanoparticle, giving the answer in standard form
05.5 — Calculating the simplest surface area to volume ratio of a cubic nanoparticle
4.2.4.2 Uses of nanoparticles
03.4 — Suggesting why nanoparticles cost less than fine particles to use in suncream
10.1 — Suggesting two reasons why nanoparticles rather than fine particles are used for self-cleaning window coatings
05.6 — Explaining why nanoparticles rather than fine particles are used in a scratch-resistant coating
4.3Quantitative chemistry
4.3.1Chemical measurements, conservation of mass and the quantitative interpretation of chemical equations
4.3.1.1 Conservation of mass and balanced chemical equations
03.1 — Describing a method used to investigate the law of conservation of mass
03.5 — Showing that a given equation obeys the law of conservation of mass
4.3.1.2 Relative formula mass
03.1 — Calculating the relative atomic mass of an element from a given sum of relative formula masses
07.5 — Calculating the percentage by mass of iron in an iron oxide
03.4 — Calculating the relative formula mass of aluminium sulfate
4.3.1.4 Chemical measurements
04.5 — Calculating a mean result including its uncertainty
4.3.2Use of amount of substance in relation to masses of pure substances
4.3.2.1 Moles (HT only)
03.5 — Calculating the number of molecules that can be made from one mole of atoms, using the Avogadro constant
4.3.2.2 Amounts of substances in equations (HT only)
07.5 — Calculating the mass of oxygen produced from a given mass of aluminium oxide electrolysed
07.4 — Calculating the minimum mass of magnesium needed to reduce a given mass of silicon dioxide
09.2 — Identifying the correct expression for the mass of carbon needed to produce a given amount of iron
05.2 — Calculating a mass of reactant needed to produce a given mass of salt, from a balanced equation
4.3.2.3 Using moles to balance equations (HT only)
08.1 — Suggesting why a step is needed in a copper oxide reduction experiment method (WS 2.2 — explaining a practical procedure step)
08.2 — Explaining why excess hydrogen gas must be burned off in the experiment (WS 2.4 — health and safety in practical work)
08.3 — Calculating the mass of copper and mass of water produced from experimental mass data
08.4 — Determining which of two possible equations matches the experimental results
08.6 — Calculating masses from an iron-and-chlorine reaction experiment to determine the product formula
07.3 — Identifying the correct equation for iron reacting with chlorine from given mole data
02.1 — Determining the empirical formula of calcium sulfide from a structure diagram
4.3.2.4 Limiting reactants (HT only)
08.7 — Explaining why titanium chloride is the limiting reactant, showing working
06.2 — Showing which reactant is the limiting reactant in a given reaction mixture
07.3 — Explaining why the volume of gas collected stops increasing at a given point
4.3.2.5 Concentration of solutions
09.5 — Calculating the mass of sodium hydroxide in a given volume of solution of known concentration
05.2 — Identifying which combination of changes increases the concentration of an acid
4.3.3Yield and atom economy of chemical reactions (chemistry only)
4.3.3.1 Percentage yield
08.8 — Calculating the actual mass of titanium produced, given percentage yield
01.5 — Calculating the actual mass of salt produced from a given percentage yield
09.2 — Calculating the percentage yield of zinc iodide from given masses
05.3 — Calculating the actual mass of product from a given percentage yield
4.3.3.2 Atom economy
02.5 — Calculating the percentage atom economy for nickel production
03.5 — Calculating the percentage atom economy for a precipitation reaction
03.6 — Giving a reason why high atom economy reactions are preferred in industry
03.3 — Calculating the percentage atom economy for extracting tin
03.3 — Calculating the percentage atom economy for cadmium production
03.6 — Explaining why high atom economy reactions are important in industry
05.4 — Calculating an atom economy from a given equation
05.5 — Comparing the atom economies of two reactions and giving a reason for the difference
4.3.4Using concentrations of solutions in mol/dm³ (chemistry only) (HT only)
4.3.4 Using concentrations of solutions in mol/dm³
09.3 — Calculating the concentration of sulfuric acid in mol/dm³ from concordant titration results (MS 2b — using concordant results as a mean)
09.5 — Calculating the concentration of potassium hydroxide solution in mol/dm³ and g/dm³
09.4 — Calculating the mass of citric acid required for a solution of given concentration in mol/dm³
09.7 — Calculating the concentration of sodium hydroxide solution in mol/dm³ from titration data
09.4 — Calculating the mass of a solid acid needed for a solution of given concentration in mol/dm³
09.5 — Calculating the concentration of a solution in mol/dm³ from titration data
08.3 — Calculating the concentration of hydrochloric acid in mol/dm³ from titration data
05.5 — Calculating the concentration of an acid in mol/dm³ from titration data
09.4 — Calculating the volume of acid needed to neutralise a given amount of alkali, using concentration in mol/dm³ — RP2 context (MS 3c)
08.4 — Calculating the concentration of sodium hydroxide solution in mol/dm³ from titration data — RP2 context (MS 3c)
08.5 — Calculating the concentration of sodium hydroxide solution in mol/dm³ from titration data — RP2 context (MS 3c)
08.6 — Calculating the mass of sodium hydroxide in a given volume of solution of known concentration in mol/dm³
4.3.5Use of amount of substance in relation to volumes of gases (chemistry only) (HT only)
4.3.5 Use of amount of substance in relation to volumes of gases
06.5 — Suggesting a reason for a difference in gas volumes collected despite equal moles produced
06.6 — Calculating the amount in moles of chlorine gas collected, using molar gas volume at RTP
07.7 — Calculating the volume of a given mass of chlorine gas at room temperature and pressure
07.6 — Calculating the volume of hydrogen gas needed for a car journey, from energy and molar volume data
05.6 — Calculating the volume of chlorine gas needed to react with a given mass of iron
08.2 — Calculating the volume of oxygen required to react with a given volume of hydrogen sulfide
07.6 — Calculating the total volume of gases present after a reaction involving an excess reagent
07.5 — Identifying the volume of oxygen produced from a given volume of hydrogen, using a balanced equation
09.6 — Calculating the relative atomic mass of a metal from percentage atom economy and gas volume data (4.1.1.6 and 4.3.3.2 noted)
10.2 — Calculating the volume of chlorine gas needed to react completely with a given mass of titanium dioxide (4.3.2.2 — reacting-mass mole ratio noted)
07.6 — Calculating the volume of gas produced at room temperature and pressure from a given mass of reactant
07.5 — Calculating the number of moles of gas from a given volume at room temperature and pressure
4.4Chemical changes
4.4.1Reactivity of metals
4.4.1.1 Metal oxides
09.1 — Identifying which substance is reduced in an equation, in terms of oxygen
03.1 — Identifying which substance is reduced in a reaction, in terms of oxygen
4.4.1.2 The reactivity series
05.1 — Plotting reactivity-series temperature-change data as a bar chart [Figure — print if needed] (WS 3.1 / MS 2c — presenting data as a bar chart)
05.3 — Describing a method to place an unknown metal in a reactivity series, giving valid results
02.2 — Explaining how given observations show that silver is less reactive than copper
02.3 — Planning an investigation to identify three unknown metals by their acid reactions, giving valid results (WS 2.2 — planning a procedure with valid results)
06.3 — Planning an investigation to compare the reactivity of an unknown metal with zinc (WS 2.2 — planning a comparative investigation)
08.4 — Suggesting two observations when barium reacts with hydrochloric acid
08.5 — Writing a balanced symbol equation for barium reacting with hydrochloric acid
09.4 — Identifying which metal has the greatest tendency to form positive ions, using given data
09.5 — Identifying which metal is aluminium from reactivity data, with an explanation
01.6 — Giving reasons why two named metals are not used to produce a salt with sulfuric acid
06.1 — Ordering five metals by reactivity from given data, with justification
09.5 — Explaining why calcium reacts more vigorously than magnesium with hydrochloric acid
4.4.1.3 Extraction of metals and reduction
02.4 — Explaining why carbon can be used to extract nickel from nickel oxide
08.1 — Suggesting a hazard associated with the chlorination stage of titanium extraction
08.3 — Suggesting why a stage of titanium extraction is carried out in argon rather than air
03.4 — Evaluating three possible methods to extract tungsten using comparison data (WS 1.4 — evaluating extraction methods using data)
07.1 — Explaining why silicon can be extracted by reduction with carbon
07.2 — Explaining why carbon is used rather than aluminium to reduce silicon dioxide — extraction/reactivity content, not electrolysis
07.3 — Giving a reason why products are difficult to separate when magnesium reduces silicon dioxide
03.2 — Explaining why carbon can be used to extract tin from tin oxide
4.4.1.4 Oxidation and reduction in terms of electrons (HT only)
08.5 — Explaining why sodium being oxidised to sodium ions makes this an oxidation reaction
08.6 — Completing the half equation for the oxidation of sodium
07.2 — Identifying what happens at the negative electrode during aluminium production
06.3 — Completing an ionic equation, with state symbols, for a metal displacement reaction
06.4 — Explaining why a displacement reaction is both an oxidation and a reduction
04.5 — Explaining why oxygen is described as being reduced in a reaction with sodium
04.5 — Completing the ionic equation for a metal displacement reaction
04.6 — Explaining why a metal is described as being oxidised in a displacement reaction
09.3 — Identifying why a given ionic equation represents a redox reaction
04.4 — Completing the ionic equation, with state symbols, for zinc reacting with copper sulfate solution
06.1 — Explaining why zinc losing electrons is an oxidation reaction
4.4.2Reactions of acids
4.4.2.2 Neutralisation of acids and salt production
01.2 — Giving the formula of calcium nitrate from its ion charges
02.2 — Completing a word equation for calcium hydroxide reacting with an acid
05.1 — Naming the salt produced by neutralising hydrochloric acid with potassium hydroxide
07.1 — Identifying an apparatus error and its effect in a gas-collection investigation (WS 2.4 — correct apparatus use)
07.2 — Suggesting a cause for an anomalous result in a gas-collection investigation (WS 3.7 — identifying a cause of anomaly)
07.4 — Suggesting further work to be more certain of a minimum reactant mass (WS 2.7 — improving certainty of a result)
07.6 — Suggesting an apparatus improvement for more accurate results, with a reason (WS 3.7 — improving accuracy)
07.7 — Evaluating two students' claims about a source of experimental error (WS 3.7 — evaluating a claim about error)
4.4.2.3 Soluble salts
01.1 — Identifying another substance type that reacts with an acid to form a soluble salt
01.3 — Describing a method to make pure, dry magnesium sulfate crystals from an oxide and dilute acid — RP1: preparation of a pure, dry sample of a soluble salt from an insoluble oxide or carbonate
05.3 — Identifying which insoluble solids can be used to make a copper salt with dilute hydrochloric acid
05.4 — Identifying steps in a method for preparing a soluble salt — RP1: preparation of a pure, dry sample of a soluble salt from an insoluble oxide or carbonate
05.5 — Describing how the filtrate should be evaporated gently — RP1: preparation of a pure, dry sample of a soluble salt from an insoluble oxide or carbonate
04.1 — Giving an observation showing a solid reactant is in excess
04.2 — Explaining why excess solid is used rather than excess acid in salt preparation
04.3 — Naming another compound that could be reacted with acid to produce the same salt
04.4 — Describing how to obtain crystals of a salt from its solution
03.1 — Planning a method to make pure, dry crystals of a salt from a carbonate and dilute acid — RP1: preparation of a pure, dry sample of a soluble salt from an insoluble oxide or carbonate
03.2 — Naming two other substances that could be reacted with dilute acid to make the same salt
01.1 — Completing the word equation for copper carbonate reacting with sulfuric acid
01.2 — Giving an observation showing a reactant is in excess in a salt-preparation method
01.3 — Giving a reason for filtering the reaction mixture — RP1: preparation of a pure, dry sample of a soluble salt from an insoluble oxide or carbonate
01.4 — Naming equipment used to warm a filtrate gently — RP1: preparation of a pure, dry sample of a soluble salt from an insoluble oxide or carbonate
09.1 — Planning a method to obtain crystals of zinc iodide from zinc and iodine (WS 2.2 — planning a preparative method; not RP1, as this uses direct reaction rather than an insoluble oxide/carbonate with acid)
04.1 — Suggesting an improvement to a salt-preparation step to ensure all solid dissolves — RP1: preparation of a pure, dry sample of a soluble salt from an insoluble oxide or carbonate (WS 2.7 — improving a method)
04.2 — Suggesting an improvement to remove remaining solid impurity from a solution — RP1: preparation of a pure, dry sample of a soluble salt from an insoluble oxide or carbonate
04.3 — Suggesting a safety precaution during a salt-preparation step — RP1: preparation of a pure, dry sample of a soluble salt from an insoluble oxide or carbonate (WS 2.4 — health and safety in practical work)
05.1 — Describing how to make copper chloride crystals from copper carbonate and dilute acid — RP1: preparation of a pure, dry sample of a soluble salt from an insoluble oxide or carbonate
4.4.2.4 The pH scale and neutralisation
02.1 — Identifying the ion produced by all acids in aqueous solution
05.2 — Writing the ionic equation for the neutralisation of hydrochloric acid with potassium hydroxide
08.4 — Explaining why the electrical conductivity of a mixture was zero at the point of neutralisation, referring to ions (4.2.2.3 — ionic conduction noted)
08.5 — Giving a reason why conductivity increased after further alkali was added (4.2.2.3 — ionic conduction noted)
05.4 — Giving the formula of the ion that makes a solution alkaline
08.1 — Identifying the ion that all alkalis contain in solution
08.2 — Writing the ionic equation, with state symbols, for a neutralisation reaction (4.2.2.2 — state symbols)
4.4.2.5 Titrations (chemistry only)
09.4 — Giving a reason for using a burette to measure the sodium hydroxide solution
02.3 — Naming a piece of titration apparatus shown in a diagram (WS 2.3 — identifying apparatus)
02.4 — Reading a volume from a titration apparatus diagram (WS 2.6 — reading a measurement from apparatus)
02.5 — Describing how titrations could be used to compare the concentration of two acid samples — RP2: determination of the reacting volumes of a strong acid and a strong alkali by titration
09.5 — Describing how to complete a titration
09.6 — Giving two reasons why a burette is used to measure the acid in a titration
09.3 — Suggesting two improvements to increase the accuracy of a titration result (WS 3.7 — improving accuracy)
08.2 — Explaining why a mean of concordant titration volumes was used in a calculation (MS 2b — using concordant results as a mean)
05.3 — Naming a titration indicator and describing its colour change — RP2: determination of the reacting volumes of solutions of a strong acid and a strong alkali by titration (AT 8)
09.3 — Describing titration setup steps needed before adding acid from a filled burette — RP2: determination of the reacting volumes of solutions of a strong acid and a strong alkali by titration (AT 8)
08.2 — Suggesting two improvements to a titration plan, with reasons — RP2: determination of the reacting volumes of solutions of a strong acid and a strong alkali by titration (WS 2.7, AT 8)
08.3 — Reading a burette scale from a diagram — RP2 context (AT 8)
08.3 — Describing how to complete a titration, naming a suitable indicator and its colour change — RP2: determination of the reacting volumes of solutions of a strong acid and a strong alkali by titration (AT 8)
08.4 — Calculating a mean titre from concordant results — RP2: determination of the reacting volumes of solutions of a strong acid and a strong alkali by titration (MS 2b)
4.4.2.6 Strong and weak acids (HT only)
09.1 — Explaining why an acid can be described as both strong and dilute
09.2 — Calculating the pH of a hydrochloric acid solution of given concentration
09.1 — Classifying a solution by concentration and acid strength
09.2 — Identifying which solution would have the lowest pH
08.1 — Explaining why the pH of an acid depends on its strength and its concentration
05.1 — Explaining what is meant by 'weak acid' in terms of ionisation
05.6 — Calculating the hydrogen ion concentration of a solution from its pH, using a given example
09.1 — Explaining what is meant by a weak acid
09.2 — Explaining what happens to the pH of an acid as it is diluted
08.5 — Identifying the correct statement comparing a strong and a weak acid solution
08.6 — Explaining why one acid solution has a lower pH than another, using given data
08.1 — Explaining what is meant by a strong acid
4.4.3Electrolysis
4.4.3.1 The process of electrolysis
06.1 — Explaining a feature of the electrolysis setup used to investigate different substances
04.1 — Explaining a colour change observed at the positive electrode during electrolysis
4.4.3.2 Electrolysis of molten ionic compounds
06.2 — Suggesting why the products of electrolysis must be kept apart
06.3 — Identifying which type of particle passes through a mesh during electrolysis of molten sodium chloride
4.4.3.3 Using electrolysis to extract metals
07.1 — Explaining why a mixture is used as the electrolyte in aluminium extraction
07.3 — Completing the balanced half-equation for the positive electrode in aluminium extraction
07.4 — Explaining why the positive electrode must be continually replaced in aluminium extraction
06.2 — Naming the substance added to aluminium oxide to form the electrolyte mixture
06.3 — Giving a reason why a mixture is used as the electrolyte in aluminium extraction
06.4 — Writing the half equation for oxygen production at the positive electrode in aluminium extraction (4.4.3.5 — half equation noted)
06.5 — Explaining why the positive electrode must be continually replaced in aluminium extraction
4.4.3.4 Electrolysis of aqueous solutions
06.3 — Identifying an apparatus error in an electrolysis gas-collection investigation and how to fix it
06.4 — Describing the trends shown in gas-volume-vs-time electrolysis results
07.6 — Explaining why sodium chloride solution cannot be used to produce sodium metal by electrolysis
04.2 — Naming the substance in solution that provides hydroxide ions
04.3 — Describing how a solid forms at the negative electrode during electrolysis
04.4 — Naming the products at each electrode when potassium iodide solution is electrolysed
07.3 — Completing a table of electrode products for two aqueous salt solutions electrolysed
07.4 — Suggesting how to find the total mass of copper produced when some falls from the electrode (WS 2.6 — devising an accurate combined measurement)
07.5 — Explaining how results support a proportionality conclusion for mass and time (WS 3.5 — interpreting data to support a conclusion)
07.6 — Explaining how results support a proportionality conclusion for mass and current (WS 3.5 — interpreting data to support a conclusion)
07.7 — Suggesting why the blue colour of copper nitrate solution fades during electrolysis
07.8 — Calculating the number of copper atoms produced during electrolysis from current and time (4.3.2.1 — moles noted)
06.4 — Identifying two additional ions present in aqueous sodium chloride solution
06.5 — Naming the alkaline solution produced during electrolysis of sodium chloride solution
06.6 — Explaining how the alkaline solution is produced, referring to the electrode processes
07.2 — Suggesting why sodium is not a product when cryolite containing sodium and aluminium ions is electrolysed
07.3 — Explaining how oxygen was produced from water during electrolysis of an aqueous solution
07.4 — Suggesting an apparatus change to compare gas volumes more accurately, with a reason (WS 3.7 — improving accuracy)
05.3 — Naming the product formed at the negative electrode when aqueous calcium chloride solution is electrolysed
05.4 — Identifying the correct half equation for the positive electrode in electrolysis of aqueous calcium chloride solution (4.4.3.5 — half equation noted)
05.5 — Interpreting apparatus and results from an electrolysis investigation of aqueous copper chromate solution
01.6 — Identifying the products of electrolysing potassium iodide solution
4.4.3.5 Representation of reactions at electrodes as half equations (HT only)
07.2 — Completing and balancing the half equation for bromine production at the positive electrode
06.1 — Identifying the correct half equation for the production of sodium
07.1 — Completing the half equation for the reaction at the negative electrode
4.5Energy changes
4.5.1Exothermic and endothermic reactions
4.5.1.1 Energy transfer during exothermic and endothermic reactions
05.2 — Giving a reason why a stated conclusion about a reaction being endothermic is incorrect
09.1 — Suggesting why a polystyrene cup was used instead of a glass beaker in a temperature-change investigation — RP4 context (WS 2.3 — selecting appropriate apparatus)
09.2 — Plotting data, drawing lines of best fit and extending them to find an intersection point [Figure — print if needed] (WS 3.1 / MS 4c — plotting data and drawing lines of best fit)
09.3 — Determining a volume from a graph intersection point
09.4 — Determining an overall temperature change from a graph
09.1 — Suggesting a reason for an anomalous point in temperature-change data (WS 3.7 — identifying a source of anomaly)
09.2 — Explaining the shape of a temperature-change graph in terms of energy transfers
02.1 — Planning a method to investigate the effect of mass of sodium carbonate on highest temperature reached (WS 2.2 — planning a procedure)
02.2 — Determining the gradient of a line of best fit, with units (MS 4d — determining the slope of a linear graph)
02.3 — Determining the initial temperature of a reaction mixture from a graph intercept
02.4 — Identifying the sketch graph matching results when a reactant is added in excess
02.1 — Identifying the independent and dependent variables in a temperature-change investigation
02.2 — Plotting data and drawing a line of best fit [Figure — print if needed] (WS 3.2 — presenting data; MS 4c)
02.3 — Determining the initial temperature of the water by extending a line of best fit
02.4 — Explaining how results show a dissolving process is endothermic
02.5 — Explaining why a mean result was reported with an uncertainty range (WS 3.4 — representing uncertainty)
02.6 — Identifying the type of error shown by a set of results (WS 3.7 — identifying error type)
04.1 — Drawing two crossing lines of best fit on a graph [Figure — print if needed]
04.2 — Explaining results shown on a graph, excluding anomalous points
04.3 — Explaining why using a polystyrene cup gives more accurate results than a glass beaker
04.6 — Suggesting a reason for random error in an experiment where no measuring mistakes were made (WS 3.7 — identifying a source of random error)
4.5.1.2 Reaction profiles
05.4 — Drawing a fully labelled reaction profile for a metal displacement reaction [Figure — print if needed]
05.7 — Labelling activation energy and overall energy change on a reaction profile [Figure — print if needed]
06.1 — Defining activation energy
07.1 — Identifying two errors in a student-drawn reaction profile
08.3 — Completing a reaction profile: the curve, activation energy and overall energy change [Figure — print if needed]
02.5 — Identifying what two labels represent on a reaction profile
02.6 — Explaining how a reaction profile shows a reaction is exothermic
04.3 — Completing a reaction profile: labelling the activation energy and overall energy change [Figure — print if needed]
08.2 — Identifying the correct reaction profile and labels for an exothermic reaction
09.1 — Completing a reaction profile with labelled arrows for energy released and activation energy [Figure — print if needed]
4.5.1.3 The energy change of reactions (HT only)
07.4 — Calculating a bond energy value from an overall energy change and other bond energies
05.5 — Calculating the overall energy change of a reaction from given bond energies
05.6 — Explaining why a reaction is exothermic, using a calculated energy value
08.4 — Calculating a bond energy value from an overall energy change and other bond energies
04.1 — Identifying the correct expression for calculating an overall energy change from bond energies
04.2 — Explaining why a reaction releases energy to the surroundings
08.3 — Calculating a bond energy value from an overall energy change and other bond energies
07.4 — Calculating a bond energy value from an overall energy change and other bond energies
09.3 — Calculating a bond energy value from an overall energy change and other bond energies
09.4 — Evaluating a statement linking the number of electron shells to bond strength, in terms of energy (4.1.1.7 — electronic structure noted)
4.5.2Chemical cells and fuel cells (chemistry only)
4.5.2.1 Cells and batteries
03.1 — Identifying which electrode/electrolyte combination gives a non-zero cell voltage
03.2 — Explaining why alkaline batteries eventually stop working
03.3 — Explaining why alkaline batteries cannot be recharged
06.1 — Identifying two variables that should be controlled in a cell-voltage investigation
06.2 — Justifying an order of reactivity for six metals using cell voltage data
06.3 — Identifying which pair of metals would produce the greatest cell voltage
07.1 — Explaining the difference between the processes in electrolysis and in a chemical cell
04.7 — Completing and labelling a diagram of a simple cell [Figure — print if needed]
06.2 — Planning an experiment to investigate how cell voltage varies with electrolyte concentration (WS 2.2 — planning a procedure)
06.2 — Identifying the least reactive metal from simple-cell voltage data, with a reason
06.3 — Predicting and explaining the voltage of a simple cell made from two named metals
4.5.2.2 Fuel cells
03.4 — Completing the balanced equation for the overall hydrogen fuel cell reaction
03.5 — Evaluating hydrogen fuel cells against rechargeable lithium-ion batteries using comparison data
06.4 — Explaining a newspaper description of hydrogen-fuel-cell trains as the new 'steam trains'
07.2 — Giving two advantages of hydrogen fuel cells over rechargeable cells for powering cars
07.3 — Writing a half equation for a hydrogen fuel cell electrode reaction
07.5 — Suggesting a way to reduce the volume needed to store hydrogen gas for a car
06.3 — Describing how a hydrogen fuel cell produces a potential difference
07.1 — Evaluating the use of hydrazine compared with hydrogen in fuel cells using given data
06.4 — Writing the word equation for the overall reaction in a hydrogen fuel cell
06.5 — Writing the two half equations for the electrode reactions in a hydrogen fuel cell
4.8Chemical analysis
4.8.1Purity, formulations and chromatography
4.8.1.1 Pure substances
04.4 — Describing how a distillation technique works, referring to the processes at two labelled points (AT 4 — distillation/separation technique)
04.5 — Reading a thermometer value during a distillation process (AT 1 — accurate temperature reading)
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