Specification (2011)

Specification (2011)2017-11-09T22:17:26+00:00
Section 1: Principles of chemistry
a) States of matter
1.01understand the arrangement, movement and energy of the particles in each of the three states of matter: solid, liquid and gas
1.02understand how the interconversions of solids, liquids and gases are achieved and recall the names used for these interconversions
1.03explain the changes in arrangement, movement and energy of particles during these interconversions.
b) Atoms
1.04describe and explain experiments to investigate the small size of particles and their movement including:
i) dilution of coloured solutions
ii) diffusion experiments
1.05understand the terms atom and molecule
1.06understand the differences between elements, compounds and mixtures
1.07describe experimental techniques for the separation of mixtures, including simple distillation, fractional distillation, filtration, crystallisation and paper chromatography
1.08explain how information from chromatograms can be used to identify the composition of a mixture.
c) Atomic structure
1.09understand that atoms consist of a central nucleus, composed of protons and neutrons, surrounded by electrons, orbiting in shells
1.10recall the relative mass and relative charge of a proton, neutron and electron
1.11understand the terms atomic number, mass number, isotopes and relative atomic mass (Ar)
1.12calculate the relative atomic mass of an element from the relative abundances of its isotopes
1.13understand that the Periodic Table is an arrangement of elements in order of atomic number
1.14deduce the electronic configurations of the first 20 elements from their positions in the Periodic Table
1.15deduce the number of outer electrons in a main group element from its position in the Periodic Table.
d) Relative formula masses and molar volumes of gases
1.16calculate relative formula masses (Mr) from relative atomic masses (Ar)
1.17understand the use of the term mole to represent the amount of substance
1.18(Triple only.) understand the term mole as the Avogadro number of particles (atoms, molecules, formulae, ions or electrons) in a substance
1.19carry out mole calculations using relative atomic mass (Ar) and relative formula mass (Mr)
1.20(Triple only.) understand the term molar volume of a gas and use its values (24 dm3 and 24,000 cm3) at room temperature and pressure (rtp) in calculations.
e) Chemical formulae and chemical equations
1.21write word equations and balanced chemical equations to represent the reactions studied in this specification
1.22use the state symbols (s), (l), (g) and (aq) in chemical equations to represent solids, liquids, gases and aqueous solutions respectively
1.23understand how the formulae of simple compounds can be obtained experimentally, including metal oxides, water and salts containing water of crystallisation
1.24calculate empirical and molecular formulae from experimental data
1.25calculate reacting masses using experimental data and chemical equations
1.26(Triple only.) calculate percentage yield
1.27carry out mole calculations using volumes and molar concentrations.
f) Ionic compounds
1.28describe the formation of ions by the gain or loss of electrons
1.29understand oxidation as the loss of electrons and reduction as the gain of electrons
1.30recall the charges of common ions in this specification
1.31deduce the charge of an ion from the electronic configuration of the atom from which the ion is formed
1.32explain, using dot and cross diagrams, the formation of ionic compounds by electron transfer, limited to combinations of elements from Groups 1, 2, 3 and 5, 6, 7
1.33understand ionic bonding as a strong electrostatic attraction between oppositely charged ions
1.34understand that ionic compounds have high melting and boiling points because of strong electrostatic forces between oppositely charged ions
1.35(Triple only.) understand the relationship between ionic charge and the melting point and boiling point of an ionic compound
1.36(Triple only.) describe an ionic crystal as a giant three-dimensional lattice structure held together by the attraction between oppositely charged ions
1.37(Triple only.) draw a diagram to represent the positions of the ions in a crystal of sodium chloride.
g) Covalent substances
1.38describe the formation of a covalent bond by the sharing of a pair of electrons between two atoms
1.39understand covalent bonding as a strong attraction between the bonding pair of electrons and the nuclei of the atoms involved in the bond
1.40explain, using dot and cross diagrams, the formation of covalent compounds by electron sharing for the following substances:
i) hydrogen
ii) chlorine
iii) hydrogen chloride
iv) water
v) methane
vi) ammonia
vii) oxygen
viii) nitrogen
ix) carbon dioxide
x) ethane
xi) ethene
1.41understand that substances with simple molecular structures are gases or liquids, or solids with low melting points
1.42explain why substances with simple molecular structures have low melting and boiling points in terms of the relatively weak forces between the molecules
1.43explain the high melting and boiling points of substances with giant covalent structures in terms of the breaking of many strong covalent bonds
1.44(Triple only.) draw diagrams representing the positions of the atoms in diamond and graphite
1.45(Triple only.) explain how the uses of diamond and graphite depend on their structures, limited to graphite as a lubricant and diamond in cutting.
h) Metallic crystals
1.46understand that a metal can be described as a giant structure of positive ions surrounded by a sea of delocalised electrons
1.47explain the electrical conductivity and malleability of a metal in terms of its structure and bonding.
i) Electrolysis
1.48understand that an electric current is a flow of electrons or ions
1.49understand why covalent compounds do not conduct electricity
1.50understand why ionic compounds conduct electricity only when molten or in solution
1.51describe experiments to distinguish between electrolytes and nonelectrolytes
1.52understand that electrolysis involves the formation of new substances when ionic compounds conduct electricity
1.53describe experiments to investigate electrolysis, using inert electrodes, of molten salts such as lead(II) bromide and predict the products
1.54(Triple only.) describe experiments to investigate electrolysis, using inert electrodes, of aqueous solutions such as sodium chloride, copper(II) sulfate and dilute sulfuric acid and predict the products
1.55write ionic half-equations representing the reactions at the electrodes during electrolysis
1.56(Triple only.) recall that one faraday represents one mole of electrons
1.57(Triple only.) calculate the amounts of the products of the electrolysis of molten salts and aqueous solutions.
Section 2: Chemistry of the elements
a) The Periodic Table
2.01understand the terms group and period
2.02recall the positions of metals and non-metals in the Periodic Table
2.03explain the classification of elements as metals or non-metals on the basis of their electrical conductivity and the acid-base character of their oxides
2.04understand why elements in the same group of the Periodic Table have similar chemical properties
2.05understand that the noble gases (Group 0) are a family of inert gases and explain their lack of reactivity in terms of their electronic configurations.
b) Group 1 elements: lithium, sodium and potassium
2.06describe the reactions of these elements with water and understand that the reactions provide a basis for their recognition as a family of elements
2.07describe the relative reactivities of the elements in Group 1
2.08(Triple only.) explain the relative reactivities of the elements in Group 1 in terms of distance between the outer electrons and the nucleus.
c) Group 7 elements: chlorine, bromine and iodine
2.09recall the colours and physical states of the elements at room temperature
2.10make predictions about the properties of other halogens in this group
2.11understand the difference between hydrogen chloride gas and hydrochloric acid
2.12explain, in terms of dissociation, why hydrogen chloride is acidic in water but not in methylbenzene
2.13describe the relative reactivities of the elements in Group 7
2.14describe experiments to demonstrate that a more reactive halogen will displace a less reactive halogen from a solution of one of its salts
2.15understand these displacement reactions as redox reactions.
d) Oxygen and oxides
2.16recall the gases present in air and their approximate percentage by volume
2.17explain how experiments involving the reactions of elements such as copper, iron and phosphorus with air can be used to investigate the percentage by volume of oxygen in air
2.18describe the laboratory preparation of oxygen from hydrogen peroxide, using manganese(IV) oxide as a catalyst
2.19describe the reactions of magnesium, carbon and sulfur with oxygen in air, and the acid-base character of the oxides produced
2.20describe the laboratory preparation of carbon dioxide from calcium carbonate and dilute hydrochloric acid
2.21describe the formation of carbon dioxide from the thermal decomposition of metal carbonates such as copper(II) carbonate
2.22describe the properties of carbon dioxide, limited to its solubility and density
2.23explain the use of carbon dioxide in carbonating drinks and in fire extinguishers, in terms of its solubility and density
2.24understand that carbon dioxide is a greenhouse gas and may contribute to climate change.
e) Hydrogen and water
2.25describe the reactions of dilute hydrochloric and dilute sulfuric acids with magnesium, aluminium, zinc and iron
2.26describe the combustion of hydrogen
2.27describe the use of anhydrous copper(II) sulfate in the chemical test for water
2.28describe a physical test to show whether water is pure.
f) Reactivity series
2.29understand that metals can be arranged in a reactivity series based on the reactions of the metals and their compounds: potassium, sodium, lithium, calcium, magnesium, aluminium, zinc, iron, copper, silver and gold
2.30describe how reactions with water and dilute acids can be used to deduce the following order of reactivity: potassium, sodium, lithium, calcium, magnesium, zinc, iron and copper
2.31deduce the position of a metal within the reactivity series using displacement reactions between metals and their oxides, and between metals and their salts in aqueous solutions
2.32understand oxidation and reduction as the addition and removal of oxygen respectively
2.33understand the terms redox, oxidising agent, reducing agent
2.34describe the conditions under which iron rusts
2.35describe how the rusting of iron may be prevented by grease, oil, paint, plastic and galvanising
2.36understand the sacrificial protection of iron in terms of the reactivity series.
g) Tests for ions and gases
2.37describe tests for the cations:
i) Li+, Na+, K+, Ca2+ using flame tests
ii) NH4+, using sodium hydroxide solution and identifying the ammonia evolved
iii) Cu2+, Fe2+ and Fe3+, using sodium hydroxide solution
2.38describe tests for the anions:
i) Cl-, Br- and I-, using dilute nitric acid and silver nitrate solution
ii) SO42-, using dilute hydrochloric acid and barium chloride solution
iii) CO32-, using dilute hydrochloric acid and identifying the carbon dioxide evolved
2.39describe tests for the gases:
i) hydrogen
ii) oxygen
iii) carbon dioxide
iv) ammonia
v) chlorine
Section 3: Organic chemistry
a) Introduction
3.01explain the terms homologous series, hydrocarbon, saturated, unsaturated, general formula and isomerism.
b) Alkanes
3.02recall that alkanes have the general formula CnH2n+2
3.03draw displayed formulae for alkanes with up to five carbon atoms in a molecule, and name the straight-chain isomers
3.04recall the products of the complete and incomplete combustion of alkanes
3.05describe the substitution reaction of methane with bromine to form bromomethane in the presence of UV light.
c) Alkenes
3.06recall that alkenes have the general formula CnH2n
3.07draw displayed formulae for alkenes with up to four carbon atoms in a molecule, and name the straight-chain isomers (knowledge of cis- and transisomers is not required)
3.08describe the addition reaction of alkenes with bromine, including the decolourising of bromine water as a test for alkenes.
d) Ethanol
3.09(Triple only.) describe the manufacture of ethanol by passing ethene and steam over a phosphoric acid catalyst at a temperature of about 300°C and a pressure of about 60-70 atm
3.10(Triple only.) describe the manufacture of ethanol by the fermentation of sugars, for example glucose, at a temperature of about 30°C
3.11(Triple only.) evaluate the factors relevant to the choice of method used in the manufacture of ethanol, for example the relative availability of sugar cane and crude oil
3.12(Triple only.) describe the dehydration of ethanol to ethene, using aluminium oxide.
Section 4: Physical chemistry
a) Acids, alkalis and salts
4.01describe the use of the indicators litmus, phenolphthalein and methyl orange to distinguish between acidic and alkaline solutions
4.02understand how the pH scale, from 0-14, can be used to classify solutions as strongly acidic, weakly acidic, neutral, weakly alkaline or strongly alkaline
4.03describe the use of universal indicator to measure the approximate pH value of a solution
4.04define acids as sources of hydrogen ions, H+, and alkalis as sources of hydroxide ions, OH¯
4.05predict the products of reactions between dilute hydrochloric, nitric and sulfuric acids; and metals, metal oxides and metal carbonates (excluding the reactions between nitric acid and metals)
4.06understand the general rules for predicting the solubility of salts in water:
i) all common sodium, potassium and ammonium salts are soluble
ii) all nitrates are soluble
iii) common chlorides are soluble, except silver chloride
iv) common sulfates are soluble, except those of barium and calcium
v) common carbonates are insoluble, except those of sodium, potassium and ammonium
4.07describe experiments to prepare soluble salts from acids
4.08describe experiments to prepare insoluble salts using precipitation reactions
4.09describe experiments to carry out acid-alkali titrations.
b) Energetics
4.10understand that chemical reactions in which heat energy is given out are described as exothermic and those in which heat energy is taken in are endothermic
4.11describe simple calorimetry experiments for reactions such as combustion, displacement, dissolving and neutralisation in which heat energy changes can be calculated from measured temperature changes
4.12(Triple only.) calculate molar enthalpy change from heat energy change
4.13understand the use of ΔH to represent enthalpy change for exothermic and endothermic reactions
4.14represent exothermic and endothermic reactions on a simple energy level diagram
4.15understand that the breaking of bonds is endothermic and that the making of bonds is exothermic
4.16(Triple only.) use average bond energies to calculate the enthalpy change during a simple chemical reaction.
c) Rates of reaction
4.17describe experiments to investigate the effects of changes in surface area of a solid, concentration of solutions, temperature and the use of a catalyst on the rate of a reaction
4.18describe the effects of changes in surface area of a solid, concentration of solutions, pressure of gases, temperature and the use of a catalyst on the rate of a reaction
4.19understand the term activation energy and represent it on a reaction profile
4.20explain the effects of changes in surface area of a solid, concentration of solutions, pressure of gases and temperature on the rate of a reaction in terms of particle collision theory
4.21explain that a catalyst speeds up a reaction by providing an alternative pathway with lower activation energy.
d) Equilibria
4.22understand that some reactions are reversible and are indicated by a special symbol ⇋ in equations
4.23describe reversible reactions such as the dehydration of hydrated copper(II) sulfate and the effect of heat on ammonium chloride
4.24understand the concept of dynamic equilibrium
4.25predict the effects of changing the pressure and temperature on the equilibrium position in reversible reactions.
Section 5: Chemistry in industry
a) Extraction and uses of metals
5.01explain how the methods of extraction of the metals in this section are related to their positions in the reactivity series
5.02describe and explain the extraction of aluminium from purified aluminium oxide by electrolysis, including:
i) the use of molten cryolite as a solvent and to decrease the required operating temperature
ii) the need to replace the positive electrodes
iii) the cost of the electricity as a major factor
5.03write ionic half-equations for the reactions at the electrodes in aluminium extraction
5.04describe and explain the main reactions involved in the extraction of iron from iron ore (haematite), using coke, limestone and air in a blast furnace
5.05explain the uses of aluminium and iron, in terms of their properties.
b) Crude oil
5.06understand that crude oil is a mixture of hydrocarbons
5.07describe and explain how the industrial process of fractional distillation separates crude oil into fractions
5.08recall the names and uses of the main fractions obtained from crude oil: refinery gases, gasoline, kerosene, diesel, fuel oil and bitumen
5.09describe the trend in boiling point and viscosity of the main fractions
5.10understand that incomplete combustion of fuels may produce carbon monoxide and explain that carbon monoxide is poisonous because it reduces the capacity of the blood to carry oxygen
5.11understand that, in car engines, the temperature reached is high enough to allow nitrogen and oxygen from air to react, forming nitrogen oxides
5.12understand that nitrogen oxides and sulfur dioxide are pollutant gases which contribute to acid rain, and describe the problems caused by acid rain
5.13understand that fractional distillation of crude oil produces more long-chain hydrocarbons than can be used directly and fewer short-chain hydrocarbons than required and explain why this makes cracking necessary
5.14describe how long-chain alkanes are converted to alkenes and shorter-chain alkanes by catalytic cracking, using silica or alumina as the catalyst and a temperature in the range of 600-700°C.
c) Synthetic polymers
5.15understand that an addition polymer is formed by joining up many small molecules called monomers
5.16draw the repeat unit of addition polymers, including poly(ethene), poly(propene) and poly(chloroethene)
5.17deduce the structure of a mon/tagomer from the repeat unit of an addition polymer
5.18describe some uses for polymers, including poly(ethene), poly(propene) and poly(chloroethene)
5.19explain that addition polymers are hard to dispose of as their inertness means that they do not easily biodegrade
5.20(Triple only.) understand that some polymers, such as nylon, form by a different process called condensation polymerisation
5.21(Triple only.) understand that condensation polymerisation produces a small molecule, such as water, as well as the polymer.
d) The industrial manufacture of chemicals
5.22understand that nitrogen from air, and hydrogen from natural gas or the cracking of hydrocarbons, are used in the manufacture of ammonia
5.23describe the manufacture of ammonia by the Haber process, including the essential conditions:
i) a temperature of about 45°C
ii) a pressure of about 200 atmospheres
iii) an iron catalyst
5.24understand how the cooling of the reaction mixture liquefies the ammonia produced and allows the unused hydrogen and nitrogen to be recirculated
5.25describe the use of ammonia in the manufacture of nitric acid and fertilisers
5.26(Triple only.) recall the raw materials used in the manufacture of sulfuric acid
5.27(Triple only.) describe the manufacture of sulfuric acid by the contact process, including the essential conditions:
(Triple only.) i) a temperature of about 450°C
(Triple only.) ii) a pressure of about 2 atmospheres
(Triple only.) iii) a vanadium(V) oxide catalyst
5.28(Triple only.) describe the use of sulfuric acid in the manufacture of detergents, fertilisers and paints
5.29(Triple only.) describe the manufacture of sodium hydroxide and chlorine by the electrolysis of concentrated sodium chloride solution (brine) in a diaphragm cell
5.30(Triple only.) write ionic half-equations for the reactions at the electrodes in the diaphragm cell
5.31(Triple only.) describe important uses of sodium hydroxide, including the manufacture of bleach, paper and soap; and of chlorine, including sterilising water supplies and in the manufacture of bleach and hydrochloric acid.
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Section 1: Principles of chemistry

      a) States of matter

      b) Atoms

      c) Atomic structure

     d) Relative formula masses and molar volumes of gases

     e) Chemical formulae and chemical equations

     f) Ionic compounds

     g) Covalent substances

     h) Metallic crystals

     i) Electrolysis

 Section 2: Chemistry of the elements

     a) The Periodic Table

     b) Group 1 elements: lithium, sodium and potassium

     c) Group 7 elements: chlorine, bromine and iodine

     d) Oxygen and oxides

     e) Hydrogen and water

     f) Reactivity series

     g) Tests for ions and gases

Section 3: Organic chemistry

     a) Introduction

     b) Alkanes

     c) Alkenes

     d) Ethanol

Section 4: Physical chemistry

     a) Acids, alkalis and salts

     b) Energetics

     c) Rates of reaction

     d) Equilibria

Section 5: Chemistry in industry

     a) Extraction and uses of metals

     b) Crude oil

     c) Synthetic polymers

     d) The industrial manufacture of chemicals