SecondaryCode 5070
Chemistry
49 topics and 227 objectives, each one something you should be able to do.
Written by Kindred in its own plain words. Your syllabus is still the official list, so check it for anything exact.
What this subject asks you to be able to do
1.1Solids, liquids and gases5 objectives
- give the properties that set solids, liquids and gases apart
- outline the structures of solids, liquids and gases by describing how far apart the particles sit, how they are arranged and how they move
- outline and account for changes of state (melting, boiling, evaporating, freezing and condensing) using kinetic particle theory
- interpret heating and cooling curves and account for them using kinetic particle theory
- outline and account for, using kinetic particle theory, how temperature and pressure affect the volume of a gas
1.2Diffusion2 objectives
- Describe and account for diffusion using kinetic particle theory.
- Describe and account for the influence of relative molecular mass upon the rate at which gases diffuse.
2.1Elements, compounds and mixtures1 objective
- Outline how elements, compounds and mixtures differ from one another
2.2Atomic structure and the Periodic Table6 objectives
- outline how an atom is built, with a central nucleus holding neutrons and protons and electrons arranged in shells around it
- give the relative charge and relative mass of each of a proton, a neutron and an electron
- define proton number (atomic number) to mean how many protons sit in an atom's nucleus
- define mass number, also known as nucleon number, to mean the combined count of protons and neutrons within an atom's nucleus
- work out the electronic configuration for both elements and ions whose proton number is 1 to 20, e.g. 2,8,3
- State that: (a) each noble gas in Group VIII has a full outer electron shell (b) for Groups I to VII, the group number equals the number of electrons in the outer shell (c) the period number equals the count of occupied electron shells
2.3Isotopes4 objectives
- Define isotopes as atoms of one element that share a proton number but differ in their number of neutrons
- State that, because they have equal numbers of electrons and therefore an identical electronic configuration, isotopes of one element share the same chemical properties
- Interpret and use the symbols that represent atoms, e.g. 126C, as well as those that represent ions, e.g. 3517Cl –
- Work out the relative atomic mass of a given element, using the relative masses of its isotopes together with their abundances
2.4Ion and ionic bonds5 objectives
- Describe how positive ions (cations) and negative ions (anions) are formed
- Describe how ionic compounds are built as a giant lattice structure, with positive and negative ions alternating in a regular arrangement
- State that the strong electrostatic attraction holding oppositely charged ions together is called an ionic bond
- Describe how ionic bonds form between metal ions and non-metal ions, including showing this bonding with dot-and-cross diagrams
- Describe these properties of ionic compounds and use their structure and bonding to explain each one: (a) high melting and boiling points (b) good electrical conductivity when molten or aqueous, but poor conductivity as a solid
2.5Simple molecules and covalent bonds3 objectives
- State that a covalent bond comes about when a shared pair of electrons joins two atoms, giving them noble gas electronic configurations
- Describe how covalent bonds are formed in simple molecules, including H2, Cl2, H2O, CH4, NH3, HCl, CH3OH, C2H4, O2, CO2 and N2, and show the electronic configurations of these molecules, and of others like them, with dot-and-cross diagrams
- Describe these properties of simple molecular compounds and use their structure and bonding to explain each one: (a) low melting and boiling points, explained by weak intermolecular forces (the particular kinds of intermolecular force need not be named) (b) poor electrical conductivity
2.6Giant covalent structures3 objectives
- Describe the giant covalent structure that each of these has: graphite, diamond and silicon(IV) oxide, SiO2
- Relate what graphite and diamond are used for to the structure and bonding of each, limited to: (a) graphite's use as a lubricant and in electrodes (b) diamond's use in cutting tools
- Describe how the properties of diamond and silicon(IV) oxide are similar, linking this to their structures
2.7Metallic bonding2 objectives
- Describe metallic bonding: in a giant metallic lattice, positive ions are held to a ‘sea’ of delocalised electrons by electrostatic attraction
- Explain these properties of metals through their structure and bonding: (a) good electrical conductivity (b) malleability and ductility
3.1Formulae7 objectives
- Give the formula of any element or compound that the subject content mentions by name
- Define a compound's molecular formula as the kinds of atom in one of its molecules and how many there are of each
- Define a compound's empirical formula as the formula that gives its different atoms or ions in their simplest whole number ratio
- Deduce a simple compound's formula from a model or diagram of it, using the relative numbers of its atoms or ions
- Deduce an ionic compound's formula, working from the charges its ions carry
- Construct word, symbol and ionic equations showing how reactants give rise to products, including state symbols
- Deduce, from the relevant information you are given, the symbol equation for a chemical reaction, showing state symbols
3.2Relative masses of atoms and molecules2 objectives
- Describe relative atomic mass, Ar, as an element's average isotope mass, measured in units of 1/12th of the mass of one 12C atom
- Define relative molecular mass, Mr, as the total of the relative atomic masses. For ionic compounds, the quantity to use is called relative formula mass, Mr.
3.3The mole and the Avogadro constant8 objectives
- State that amount of substance is measured in moles, symbol mol, that each mole holds 6.02×10^23 particles, e.g. atoms, ions, molecules, and that this figure is called the Avogadro constant
- Use the relationship mass (g) / molar mass (g/mol) = amount of substance (mol) to work out: (a) the amount of substance (b) the mass (c) the molar mass (d) the relative atomic mass, or the relative molecular or formula mass (e) the number of particles, by means of the Avogadro constant
- apply the molar gas volume, set at 24 dm^3 for r.t.p. (room temperature and pressure), in any calculation that concerns gases
- state that concentration may be expressed in g/dm^3 or mol/dm^3
- work out stoichiometric reacting masses, limiting reactants, gas volumes at r.t.p., solution volumes and solution concentrations given in g/dm^3 and mol/dm^3, including converting between cm^3 and dm^3
- Use data gathered in a titration experiment to work out a solution's concentration
- work out empirical formulae and molecular formulae, provided suitable data are given
- find the percentage yield, the percentage composition by mass and the percentage purity, where suitable data are supplied
4.1Electrolysis11 objectives
- Define electrolysis as a process in which an electric current brings about decomposition of an ionic compound that is molten or present in aqueous solution
- Identify, in a simple electrolytic cell: (a) the positive electrode, called the anode (b) the negative electrode, called the cathode (c) the molten or aqueous substance being electrolysed, called the electrolyte
- Describe how charge is transferred during electrolysis, covering: (a) electrons moving through the external circuit (b) electrons being lost or gained at the electrodes (c) ions moving through the electrolyte
- identify what is produced at each electrode and describe what is seen when electrolysing: (a) molten lead(II) bromide (b) concentrated aqueous sodium chloride (c) dilute sulfuric acid, with inert platinum or carbon/graphite electrodes
- Identify what is produced at each electrode, and describe what is seen, when aqueous copper(II) sulfate is electrolysed with inert carbon/graphite electrodes and also with copper electrodes
- state that the cathode produces metals or hydrogen, whereas the anode produces non-metals other than hydrogen
- Predict, at each electrode, the identity of the products when a binary compound is electrolysed while molten
- Predict which products form at each electrode when a halide compound is electrolysed as a dilute or a concentrated aqueous solution
- Construct ionic half-equations that describe the reaction at the cathode, to show reduction, and the reaction at the anode, to show oxidation
- state that metal items are electroplated in order to make them look better and resist corrosion
- Describe the procedure by which metals undergo electroplating
4.2Hydrogen–oxygen fuel cells2 objectives
- State that a hydrogen–oxygen fuel cell makes electricity from hydrogen and oxygen, and that water is the only chemical product formed
- Describe, for powering vehicles, the advantages and disadvantages of hydrogen–oxygen fuel cells set against those of gasoline/petrol engines
5.1Exothermic and endothermic reactions7 objectives
- State that an exothermic reaction gives thermal energy out to its surroundings, so that the temperature of those surroundings rises
- State that an endothermic reaction absorbs thermal energy from its surroundings, which causes the temperature of those surroundings to fall
- State that the thermal energy transferred during a reaction is known as its enthalpy change, ΔH, and that ΔH has a negative value for exothermic reactions and a positive value for endothermic ones
- Define activation energy, Ea, to be the smallest amount of energy that particles need when they collide in order to react
- Draw, label and interpret, from the information supplied, reaction pathway diagrams for exothermic reactions and for endothermic reactions, showing: (a) reactants (b) products (c) the reaction's enthalpy change, ΔH (d) activation energy, Ea
- State that bond breaking counts as an endothermic process and bond making as an exothermic one, and account for how a reaction's enthalpy change arises from bond breaking and bond making
- Work out a reaction's enthalpy change from bond energies
6.1Physical and chemical changes1 objective
- Identify physical and chemical changes and outline how the two kinds differ from each other
6.2Rate of reaction5 objectives
- Describe the collision theory with reference to (a) how many particles are present in each unit of volume, (b) how often particles collide with one another, (c) the kinetic energy that particles carry, (d) the activation energy, Ea
- State that a catalyst raises the rate of a reaction, lowers the activation energy, Ea, of that reaction, and remains unchanged once the reaction has finished
- Describe, and account for, how the rate of reactions is influenced by: (a) altering the concentration of solutions (b) altering the pressure of gases (c) altering the surface area of solids (d) altering the temperature (e) introducing or withdrawing a catalyst, including enzymes, by means of collision theory
- Describe and evaluate practical ways of investigating a reaction's rate, including tracking how the mass of a product or reactant changes and tracking the gas that forms
- Interpret the data that experiments on reaction rates produce, including graphs
6.3Reversible reactions and equilibrium11 objectives
- State that certain chemical reactions are reversible, and that the ⇌ symbol is used to show this
- Describe how a reversible reaction's direction can be reversed by varying conditions, in these cases: (a) applying heat to hydrated compounds (b) introducing water to anhydrous compounds, including copper(II) sulfate and cobalt(II) chloride
- State that a reversible reaction taking place in a closed system has reached equilibrium once: (a) the forward and reverse reactions proceed at the same rate (b) reactant and product concentrations are no longer changing
- Predict and explain, using the information supplied, how the position of equilibrium of a reversible reaction responds to: (a) a change in temperature (b) a change in pressure (c) a change in concentration (d) the use of a catalyst
- State the symbol equation, for ammonia manufacture by the Haber process, N2(g) + 3H2(g) ⇌ 2NH3(g)
- Name the origins of the hydrogen gas (from methane) and the nitrogen gas (from air) used in the Haber process
- Give the usual conditions of the Haber process: 450°C, 20000kPa/200 atm, and an iron catalyst
- Give the symbol equation for turning sulfur dioxide into sulfur trioxide in the Contact process: 2SO2(g) + O2(g) ⇌ 2SO3(g)
- State where the Contact process obtains its sulfur dioxide (by burning sulfur, or by roasting sulfide ores) and its oxygen (from air)
- State the usual conditions the Contact process uses to turn sulfur dioxide into sulfur trioxide: 450°C, 200kPa/2 atm and a vanadium(V) oxide catalyst
- Account for, using reaction rate and equilibrium position, why the usual conditions named are chosen in both the Haber process and the Contact process, including safety factors and economic issues
6.4Redox10 objectives
- Apply a Roman numeral to show an element's oxidation number inside a compound
- Define a redox reaction as one in which reduction and oxidation take place at the same time.
- Define oxidation by reference to: (a) taking in oxygen (b) giving up electrons (c) a rise in oxidation number
- Define reduction by reference to: (a) losing oxygen (b) taking on electrons (c) a fall in oxidation number
- Identify redox reactions as ones in which oxygen is gained and lost, or electrons are gained and lost
- Identify redox reactions from shifts in oxidation number, applying these rules: (a) an element in the uncombined form has an oxidation number of zero (b) for a monatomic ion, the oxidation number matches the charge on the ion (c) the oxidation numbers within a compound add up to zero (d) the oxidation numbers within an ion total the charge on that ion
- Identify redox reactions from the colour shifts observed when acidified aqueous potassium manganate(VII) or aqueous potassium iodide is used.
- Define an oxidising agent: it oxidises a different substance and, in doing so, is reduced
- Define a reducing agent: it reduces a different substance and, in doing so, is oxidised
- Identify, within redox reactions, the oxidation, oxidising agents, reduction and reducing agents.
7.1The characteristic properties of acids and bases12 objectives
- State that acids in aqueous solution contain H+ ions, and alkalis in aqueous solution contain OH– ions
- Define an acid as a proton donor and a base as a proton acceptor
- State that a base is a metal's oxide or hydroxide, and that alkalis are bases which are soluble
- Describe, as the characteristic properties of acids, the way they react with: (a) metals (b) bases (c) carbonates
- Describe, as the characteristic properties of bases, the way they react with: (a) acids (b) ammonium salts
- State that when an acid reacts with a base, neutralisation takes place
- Outline the neutralisation of an acid by an alkali, forming water: H+(aq) + OH−(aq) → H2O(l)
- Outline what acids and alkalis do to (a) litmus, (b) thymolphthalein, (c) methyl orange
- Define strong and weak acids: in aqueous solution, a strong acid is completely dissociated, whereas a weak acid is partially dissociated
- Name some strong acids, including hydrochloric acid, nitric acid and sulfuric acid, and construct symbol equations showing that each one dissociates completely, e.g. HCl(aq) → H+(aq) + Cl–(aq)
- Name some weak acids, including carboxylic acids, and construct symbol equations showing that each one is partially dissociated, e.g. for ethanoic acid, CH3COOH(aq) ⇌ H+(aq) + CH3COO–(aq)
- Describe how universal indicator paper, through its colour and the pH it shows, lets you compare hydrogen ion concentration, neutrality, relative acidity and relative alkalinity
7.2Oxides2 objectives
- Describe an amphoteric oxide as one that gives a salt and water when it reacts with an acid, and also when it reacts with a base
- Classify each oxide, according to the metallic or non-metallic character of its element, as basic (including CuO and CaO), acidic (including SO2 and CO2) or amphoteric (limited to Al2O3 and ZnO)
7.3Preparation of salts5 objectives
- Describe how a soluble salt is prepared, separated and purified when an acid reacts with: (a) an alkali, using titration (b) an excess of metal (c) an excess of insoluble base (d) an excess of insoluble carbonate
- Describe how an insoluble salt is made by precipitation
- Describe the general rule on the solubility of salts: (a) all sodium, potassium and ammonium salts are soluble (b) all nitrates are soluble (c) all chlorides except lead and silver chloride are soluble (d) all sulfates except barium, calcium and lead sulfate are soluble (e) all carbonates except sodium, potassium and ammonium carbonate are insoluble (f) all hydroxides except sodium, potassium, ammonium and calcium hydroxide are insoluble, and calcium hydroxide is partially soluble
- Define hydrated and anhydrous substances: a hydrated substance has water chemically combined in it, while an anhydrous substance contains no water
- Define water of crystallisation: the molecules of water held within hydrated crystals, including CuSO4•5H2O and CoCl2•6H2O
8.1Arrangement of elements6 objectives
- Outline how the Periodic Table is laid out, with elements arranged into periods and groups and ordered by increasing proton number/atomic number
- Outline how an element's character shifts from metallic to non-metallic as you move along a period
- Describe how the charge on the ions that an element forms depends on its group number
- Account for elements in the same group sharing similar chemical properties, using their electronic configurations
- Explain how knowing an element's place in the Periodic Table helps to predict its properties
- Identify the trends within groups, using information supplied about their elements
8.2Group I properties2 objectives
- Describe lithium, sodium and potassium, the alkali metals of Group I, as relatively soft metals that show general trends going down the group, limited to: (a) melting point decreasing (b) density increasing (c) reactivity increasing
- Predict, from information supplied about them, the properties of the remaining Group I elements
8.3Group VII properties4 objectives
- Describe chlorine, bromine and iodine, the halogens of Group VII, as diatomic non-metals that show general trends going down the group, limited to: density increasing, reactivity decreasing
- State what the halogens look like at r.t.p.: bromine is a red-brown liquid, chlorine is a pale yellow-green gas and iodine is a grey-black solid
- Describe and account for the reactions in which a halogen displaces other halide ions
- Predict, from information supplied about them, the properties of the remaining Group VII elements
8.4Transition elements1 objective
- Describe the transition elements as dense metals with high melting points which often act as catalysts, both as elements and in compounds, form coloured compounds and have variable oxidation numbers
8.5Noble gases1 objective
- Outline how the noble gases of Group VIII are unreactive and monatomic, then account for this using their electronic configuration
9.1Properties of metals2 objectives
- Compare the broad physical properties of metals with those of non-metals, including: melting points and boiling points, malleability and ductility, electrical conductivity, thermal conductivity
- Describe the chemical properties that metals generally show, limited to how they react with dilute acids, cold water and steam, and oxygen
9.2Uses of metals1 objective
- describe what metals are used for, based on their physical properties, including: aluminium used to build aircraft owing to its low density; aluminium used in overhead electrical cables owing to its low density and good electrical conductivity; aluminium used for food containers owing to its resistance to corrosion; copper used in electrical wiring owing to its good electrical conductivity and ductility
9.3Alloys and their properties4 objectives
- Describe an alloy as a metal mixed with other elements, including: brass, which is copper mixed with zinc, and stainless steel, which is iron mixed with other elements such as chromium, nickel and carbon
- Explain, from its structure, why an alloy may be harder and stronger than pure metals: its atoms or ions are of different sizes, so the layers can no longer slide past one another
- outline the applications of alloys based on their physical properties, including stainless steel in cutlery, chosen for its hardness and its ability to withstand rusting
- Identify which structure diagrams show an alloy
9.4Reactivity series5 objectives
- Give the sequence in which potassium, sodium, calcium, magnesium, aluminium, carbon, zinc, iron, hydrogen, copper, silver, gold appear in the reactivity series.
- Describe the relative reactivities of metals, based on the tendency of each to form positive ions, from any displacement reactions they undergo with the aqueous ions of magnesium, zinc, iron, copper and silver
- Describe the reactions taking place, if any, when potassium, sodium and calcium meet cold water, magnesium meets steam, and magnesium, zinc, iron, copper, silver and gold meet dilute hydrochloric acid; then explain these outcomes using where the metals sit in the reactivity series.
- Explain why aluminium appears to be unreactive, with reference to its oxide layer
- Deduce a reactivity order from a supplied set of experimental results.
9.5Corrosion of metals5 objectives
- Give the circumstances that must be present for iron and steel to rust and produce hydrated iron(III) oxide
- Describe the manner in which barrier methods stop rusting by keeping oxygen or water away
- Give a few widespread barrier methods, including painting, greasing and coating with plastic
- Account for how sacrificial protection works, relating it to the reactivity series and also to electron loss
- Describe how galvanising with zinc works both as a barrier method and as sacrificial protection
9.6Extraction of metals3 objectives
- Outline how easily a metal can be obtained from its ores, linked to where that metal sits in the reactivity series
- Describe how iron is extracted from hematite in a blast furnace, including a symbol equation for every step, limited to: burning carbon (coke) to give heat and form carbon dioxide; reducing carbon dioxide to form carbon monoxide; reducing iron(III) oxide with carbon monoxide; breaking down calcium carbonate/limestone by heat (thermal decomposition) to give calcium oxide; and the formation of slag
- Describe how purified bauxite/aluminium oxide is used to extract aluminium, including: what cryolite does, why the carbon anodes have to be replaced regularly, and the reactions at each electrode, including ionic half-equations
10.1Water7 objectives
- Describe the chemical tests which reveal whether water is present, making use of anhydrous cobalt(II) chloride and anhydrous copper(II) sulfate
- Describe the procedure for checking how pure water is by measuring its melting point and its boiling point
- Account for the use of distilled water in practical chemistry in preference to tap water, given that it contains fewer chemical impurities
- State that water drawn from natural sources can have substances in it, including: sewage, plastics, harmful microbes, metal compounds, dissolved oxygen, nitrates from fertilisers, and phosphates from fertilisers and detergents
- State that, among these substances, some are beneficial, including: some metal compounds, which supply minerals essential for life, and dissolved oxygen, which aquatic life depends on
- State that a number of these substances can do harm, including: certain metal compounds, which are toxic; certain plastics, which harm aquatic life; sewage, which carries harmful microbes that cause disease; and nitrates and phosphates, which lead to deoxygenation of the water and harm to aquatic life
- Describe the treatment applied to the domestic water supply, with reference to: sedimentation and filtration to take out solids, carbon used to take out tastes and odours, chlorination to destroy microbes
10.2Fertilisers2 objectives
- State that ammonium salts are a type of fertiliser, as are nitrates
- Describe how NPK fertilisers are used to supply nitrogen, phosphorus and potassium, which boosts plant growth
10.3Air quality and climate8 objectives
- Give the makeup of uncontaminated, moisture-free air: about 78% nitrogen, N2, 21% oxygen, O2, with what remains being a blend of noble gases and carbon dioxide, CO2
- State where each of these air pollutants comes from: (a) carbon dioxide, released when carbon-containing fuels undergo complete combustion (b) carbon monoxide and particulates, released when carbon-containing fuels undergo incomplete combustion (c) methane, released as vegetation undergoes decomposition and in waste gases from animals' digestion (d) oxides of nitrogen, released by car engines (e) sulfur dioxide, released by the combustion of fossil fuels that contain sulfur compounds
- State how these air pollutants cause harm: (a) carbon dioxide: climate change, because higher levels of the gas increase global warming (b) carbon monoxide: it is a toxic gas (c) particulates: a greater risk of cancer and respiratory problems (d) methane: climate change, because higher levels of the gas increase global warming (e) oxides of nitrogen: they cause acid rain, respiratory problems and photochemical smog (f) sulfur dioxide: it causes acid rain
- Describe the mechanism by which carbon dioxide and methane, acting as greenhouse gases, produce global warming, limited to: (a) absorbing, reflecting and emitting thermal energy (b) lowering the rate at which thermal energy is lost to space
- State and explain ways of lessening the harm done by these environmental problems, limited to: (a) climate change: farming fewer livestock, planting trees, using less fossil fuel, and making more use of hydrogen and of renewable energy, e.g. solar, wind (b) acid rain: fitting catalytic converters to vehicles, and cutting sulfur dioxide emissions by switching to low-sulfur fuels and by flue gas desulfurisation, which uses calcium oxide
- Explain how car engines produce oxides of nitrogen, and describe how a catalytic converter removes them, e.g. 2CO + 2NO → 2CO2 + N2
- Describe photosynthesis as a reaction in which carbon dioxide and water combine to make glucose and oxygen, with chlorophyll present and drawing on energy from light
- Give the word equation and the symbol equation that represent photosynthesis
11.1Formulae, functional groups and terminology8 objectives
- State that a structural formula is a clear, unambiguous depiction of how a molecule's atoms are laid out, including CH2=CH2, CH3CH2OH, CH3COOCH3
- Draw and interpret a molecule's displayed formula so that it shows every atom and every bond.
- Write and interpret the general formula that the compounds of one homologous series share, limited to: (a) CnH2n+2 for alkanes (b) CnH2n for alkenes (c) CnH2n+1OH for alcohols (d) CnH2n+1COOH for carboxylic acids
- Define structural isomers as compounds that share a molecular formula but whose structural formulae differ, including C4H10, with the isomers CH3CH2CH2CH3 and CH3CH(CH3)CH3, and C4H8, with the isomers CH3CH2CH=CH2 and CH3CH=CHCH3
- Identify a functional group as an atom, or a group of atoms, that gives a homologous series its chemical properties
- Describe the general characteristics shared by the members of a homologous series: they (a) have the same functional group (b) share one general formula (c) differ from the next member by a –CH2– unit (d) show a trend in their physical properties (e) have similar chemical properties
- State that in a saturated compound's molecules, every carbon–carbon bond is a single bond
- State that the molecules of an unsaturated compound contain at least one carbon–carbon bond that is not a single bond
11.2Naming organic compounds3 objectives
- Give the names of these unbranched compounds, and sketch their structural and displayed formulae: (a) alkanes (b) alkenes, including but-1-ene and but-2-ene (c) alcohols, including propan-1-ol, propan-2-ol, butan-1-ol and butan-2-ol (d) carboxylic acids (e) products of the reactions set out in 11.4–11.7 that contain as many as four carbon atoms per molecule
- State which family a compound belongs to when it is given by a chemical name ending in -ane, -ene, -ol or -oic acid, or by a molecular, structural or displayed formula
- Name and draw the displayed formula of each unbranched ester that can be produced from unbranched alcohols and carboxylic acids, with no more than four carbon atoms in each
11.3Fuels7 objectives
- Name coal, natural gas and petroleum as the three fossil fuels
- Name methane as the compound that makes up most of natural gas
- State that a hydrocarbon is a compound made of only carbon and hydrogen
- State that a mixture of hydrocarbons makes up petroleum
- Describe the way petroleum is divided into useful fractions using fractional distillation
- Describe how the properties of petroleum fractions change from the bottom of the fractionating column to its top, limited to: (a) chain length decreases (b) volatility gets higher (c) boiling points get lower (d) viscosity gets lower
- Name what each fraction is used for: (a) the refinery gas fraction, gas for heating and cooking (b) the gasoline/petrol fraction, fuel for cars (c) the naphtha fraction, chemical feedstock (d) the kerosene/paraffin fraction, jet fuel (e) the diesel oil/gas oil fraction, fuel for diesel engines (f) the fuel oil fraction, fuel for ships and for heating homes (g) the lubricating oil fraction, lubricants, waxes and polishes (h) the bitumen fraction, road making
11.4Alkanes4 objectives
- State that alkanes are saturated hydrocarbons whose bonds are all single covalent bonds
- Describe alkanes' properties as generally unreactive, except with regard to combustion and substitution by chlorine.
- State that a substitution reaction involves swapping one atom or group of atoms out for a different atom or group of atoms
- Describe how alkanes undergo substitution with chlorine in a photochemical reaction, where ultraviolet light supplies the activation energy, Ea, and, limited to monosubstitution, draw the structural or displayed formulae of what is produced
11.5Alkenes6 objectives
- State that alkenes have a double covalent bond between carbon atoms and so count as unsaturated hydrocarbons
- Describe how alkenes and hydrogen are manufactured by cracking larger alkane molecules with a catalyst and a high temperature
- Describe why larger alkane molecules are cracked.
- Describe the test that distinguishes saturated from unsaturated hydrocarbons through how each reacts with aqueous bromine.
- State that only one product comes out of an addition reaction.
- Describe the properties of alkenes shown by their addition reactions with: (a) aqueous bromine or bromine (b) hydrogen with a nickel catalyst present (c) steam with an acid catalyst present, and for each product draw its structural or displayed formula
11.6Alcohols3 objectives
- Describe how ethanol is manufactured by: (a) fermenting aqueous glucose at 25–35°C, with yeast present and oxygen absent (b) catalytic addition, in which steam is added to ethene at 300°C and 6000kPa/60 atm with an acid catalyst present, including a comparison of the two methods' advantages and disadvantages
- Describe how alcohols undergo combustion
- State how ethanol is used: (a) as a solvent (b) as a fuel
11.7Carboxylic acids3 objectives
- Describe how carboxylic acids react with: (a) metals (b) bases (c) carbonates, including the name and formula of each salt formed
- Describe how ethanoic acid is made when ethanol is oxidised: (a) using acidified aqueous potassium manganate(VII) (b) through bacterial oxidation in the process of making vinegar
- Describe what takes place when a carboxylic acid is combined with an alcohol in the presence of an acid catalyst, producing an ester
11.8Polymers12 objectives
- define a polymer as a big molecule made by joining many smaller molecules known as monomers
- Identify, for both addition and condensation polymers, the repeat units and/or linkages present.
- Deduce an addition polymer's structure or repeat unit when given its alkene, and deduce the alkene when given the polymer
- Deduce a condensation polymer's structure or repeat unit when given its monomers, and deduce the monomers when given the polymer, limited to: (a) polyamides, formed when a dicarboxylic acid reacts with a diamine (b) polyesters, formed when a dicarboxylic acid reacts with a diol
- outline how addition polymerisation and condensation polymerisation differ
- State that plastics consist of polymers
- Describe the ways in which the properties of plastics affect how they are disposed of
- Describe the problems that plastics cause for the environment, limited to: (a) disposal of them at land fill sites (b) the way they accumulate in oceans (c) the toxic gases formed when they are burned
- Outline and sketch the structures of: (a) nylon, which is a polyamide (b) PET, which is a polyester
- State that PET can be broken back down into its monomers, which can then be polymerised again
- Describe proteins as natural polyamides, built from amino acid monomers that have the general structure: H O H H O N H C C R, in which R stands for one of various kinds of side-chain
- Describe proteins and sketch their structure in the form: N H C O N H C O N H C O
12.1Experimental design3 objectives
- Name suitable apparatus for measuring time, temperature, mass and volume, including: (a) balances (b) burettes (c) gas syringes (d) measuring cylinders (e) stop-watches (f) thermometers (g) volumetric pipettes
- Suggest the benefits and drawbacks of experimental methods and equipment
- Describe each of these terms: (a) solvent, the substance a solute dissolves in (b) solute, the substance that dissolves in a solvent (c) solution, a mixture in which one or more solutes are dissolved in a solvent (d) saturated solution, a solution with the maximum concentration of solute that the solvent can dissolve at a particular temperature (e) residue, the substance left behind after evaporation, distillation, filtration or a similar process (f) filtrate, the liquid or solution that comes out after passing through a filter
12.2Acid–base titrations2 objectives
- Outline an acid–base titration, covering the employment of (a) a burette, (b) a volumetric pipette and (c) an appropriate indicator
- Describe how an indicator is used to pinpoint when a titration reaches its end-point
12.3Chromatography4 objectives
- Describe the method of paper chromatography for splitting up mixtures of soluble substances, making use of a suitable solvent.
- Describe how locating agents are used when separating mixtures that contain colourless substances; specific locating agents do not need to be known
- Interpret simple chromatograms in order to recognise: (a) unidentified substances by matching them with reference substances (b) whether substances are pure or impure
- Give and apply the equation that defines Rf: Rf = distance moved by the substance / distance moved by the solvent
12.4Separation and purification3 objectives
- Describe and explain how separation and purification are carried out using: (a) filtration (b) crystallisation (c) simple distillation (d) fractional distillation (e) a suitable solvent
- Suggest which separation and purification techniques fit, using details given about the substances concerned
- Identify substances and assess how pure they are from data on melting points and boiling points
12.5Identification of ions and gases4 objectives
- Describe the tests used to identify these anions: (a) carbonate, CO3 2–: react it with dilute acid, then test the gas given off for carbon dioxide (b) chloride, Cl –, bromide, Br –, and iodide, I –: acidify with dilute nitric acid and then add aqueous silver nitrate to it (c) nitrate, NO3 –: reduce it using aqueous sodium hydroxide and aluminium foil, then test the gas given off for ammonia (d) sulfate, SO4 2–: acidify with dilute nitric acid and then add aqueous barium nitrate to it (e) sulfite, SO3 2–: react it with acidified aqueous potassium manganate(VII)
- Describe how aqueous sodium hydroxide and aqueous ammonia can be used to recognise the aqueous cations listed here: (a) aluminium, Al 3+ (b) ammonium, NH4 + (c) calcium, Ca2+ (d) chromium(III), Cr3+ (e) copper(II), Cu2+ (f) iron(II), Fe2+ (g) iron(III), Fe3+ (h) zinc, Zn2+
- Describe the tests used to recognise the following gases: (a) ammonia, NH3, with moist red litmus paper (b) carbon dioxide, CO2, with limewater (c) chlorine, Cl2, with moist litmus paper (d) hydrogen, H2, with a lit splint (e) oxygen, O2, with a glowing splint (f) sulfur dioxide, SO2, with acidified aqueous potassium manganate(VII)
- Describe how a flame test is used to identify these cations: (a) lithium, Li+ (b) sodium, Na+ (c) potassium, K+ (d) calcium, Ca2+ (e) barium, Ba2+ (f) copper(II), Cu2+
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