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Textbook pagesRead textbookTried exam questions in textbookTried homework questionsAsked for help Principles of Chemistry – States of Matter (C1 work)1.1understand the arrangement, movement and energy of the particles in solid, liquid and gas11.2understand how to change between the states the names used for these interconversions21.3explain the changes in arrangement, movement and energy of particles during these interconversions.2-3, 31Atoms1.4describe and explain experiments to investigate the small size of particles and their movement including:i dilution of coloured solutionsii diffusion experiments3-41.5understand the terms atom and molecule6-71.6understand the differences between elements, compounds and mixtures301.7describe experimental techniques for the separation of mixtures, including simple distillation, fractional distillation, filtration, crystallisation and paper chromatography89-911.8explain how information from chromatograms can be used to identify the composition of a mixture.911.9understand that atoms consist of a central nucleus, composed of protons and neutrons, surrounded by electrons, orbiting in shells71.10recall the relative mass and relative charge of a proton, neutron and electron71.11understand the terms atomic number, mass number, isotopes and relative atomic mass (Ar)7-81.12calculate the relative atomic mass of an element from the relative abundances of its isotopes7-81.13understand that the Periodic Table is an arrangement of elements in order of atomic number101.14deduce the electronic configurations of the first 20 elements from their positions in the Periodic Table10-111.15deduce the number of outer electrons in a main group element from its position in the Periodic Table.10Periodic Table 2.1understand the terms group and period992.2recall the positions of metals and non-metals in the Periodic Table1002.3explain the classification of elements as metals or non-metals on the basis of their electrical conductivity and the acid-base character of their oxides1002.4understand why elements in the same group of the Periodic Table have similar chemical properties92.5understand that the noble gases (Group 0) are a family of inert gases and explain their lack of reactivity in terms of their electronic configurations.101Group 12.6describe the reactions of these elements with water and understand that the reactions provide a basis for their recognition as a family of elements102-1032.7describe the relative reactivities of the elements in Group 11042.8explain the relative reactivities of the elements in Group 1 in terms of distance between the outer electrons and the nucleus104Group 72.9recall the colours and physical states of the elements at room temperature1052.10make predictions about the properties of other halogens in this group1052.11understand the difference between hydrogen chloride gas and hydrochloric acid782.12explain, in terms of dissociation, why hydrogen chloride is acidic in water but not in methylbenzene782.13describe the relative reactivities of the elements in Group 71062.14describe experiments to demonstrate that a more reactive halogen will displace a less reactive halogen from a solution of one of its salts107-1082.15understand these displacement reactions as redox reactions108Oxygen and Oxides2.16recall the gases present in air and their approximate percentage by volume 542.17explain 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 air612.18describe the laboratory preparation of oxygen from hydrogen peroxide, using manganese(IV) oxide as a catalyst552.19describe the reactions of magnesium, carbon and sulfur with oxygen in air, and the acid-base character of the oxides produced55-562.20describe the laboratory preparation of carbon dioxide from calcium carbonate and dilute hydrochloric acid582.21describe the formation of carbon dioxide from the thermal decomposition of metal carbonates such as copper(II) carbonate582.22describe the properties of carbon dioxide, limited to its solubility and density582.23explain the use of carbon dioxide in carbonating drinks and in fire extinguishers, in terms of its solubility and density582.24understand that carbon dioxide is a greenhouse gas and may contribute to climate change.58Hydrogen and Water2.25describe the reactions of dilute hydrochloric and dilute sulfuric acids with magnesium, aluminium, zinc and iron 712.26describe the combustion of hydrogen922.27describe the use of anhydrous copper(II) sulfate in the chemical test for water932.28describe a physical test to show whether water is pure.93Reactivity Series2.29understand 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 602.30describe 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 copper662.31deduce 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 solutions632.32understand oxidation and reduction as the addition and removal of oxygen respectively632.33understand the terms redox, oxidising agent, reducing agent632.34describe the conditions under which iron rusts652.35describe how the rusting of iron may be prevented by grease, oil, paint, plastic and galvanising1442.36understand the sacrificial protection of iron in terms of the reactivity series.145Tests for Ions and Gases2.37describe tests for the cations:i Li+, Na+, K+, Ca2+ using flame testsii NH4+, using sodium hydroxide solution and identifying the ammonia evolvediii Cu2+, Fe2+ and Fe3+, using sodium hydroxide solution93-942.38describe tests for the anions:i Cl, Br and I, using dilute nitric acid and silver nitrate solutionii SO4, using dilute hydrochloric acid and barium chloride solutioniii CO3, using dilute hydrochloric acid and identifying the carbon dioxide evolved94-962.39describe tests for the gases:i hydrogenii oxygeniii carbon dioxideiv ammoniav chlorine.92-93Ionic Compounds (I Junior work starts here)1.28describe the formation of ions by the gain or loss of electrons 18-191.29understand oxidation as the loss of electrons and reduction as the gain of electrons18-191.30recall the charges of common ions in this specification18-191.31deduce the charge of an ion from the electronic configuration of the atom from which the ion is formed18-191.32explain, 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, 718-191.33understand ionic bonding as a strong electrostatic attraction between oppositely charged ions25-271.34understand that ionic compounds have high melting and boiling points because of strong electrostatic forces between oppositely charged ions25-271.35understand the relationship between ionic charge and the melting point and boiling point of an ionic compound25-271.36describe an ionic crystal as a giant three-dimensional latticestructure held together by the attraction between oppositelycharged ions25-271.37draw a diagram to represent the positions of the ions in a crystal of sodium chloride.25Covalent Substances1.38describe the formation of a covalent bond by the sharing of a pair of electrons between two atoms 131.39understand covalent bonding as a strong attraction between the bonding pair of electrons and the nuclei of the atoms involved in the bond131.40explain, using dot and cross diagrams, the formation of covalent compounds by electron sharing for the following substances:i hydrogenii chlorineiii hydrogen chlorideiv waterv methanevi ammoniavii oxygenviii nitrogenix carbon dioxidex ethanexi ethene14-161.41understand that substances with simple molecular structures are gases or liquids, or solids with low melting points291.42explain why substances with simple molecular structures have low melting and boiling points in terms of the relatively weak forces between the molecules291.43explain the high melting and boiling points of substances with giant covalent structures in terms of the breaking of many strong covalent bonds27-291.44draw diagrams representing the positions of the atoms in diamond and graphite27-291.45explain how the uses of diamond and graphite depend on their structures, limited to graphite as a lubricant and diamond in cutting27-29Metallic Crystals1.46understand that a metal can be described as a giant structure of positive ions surrounded by a sea of delocalised electrons24-251.47explain the electrical conductivity and malleability of a metal in terms of its structure and bonding.24-25Organic Chemistry3.1explain the terms homologous series, hydrocarbon, saturated, unsaturated, general formula and isomerism149-154Alkanes3.2recall that alkanes have the general formula CnH2n+2 1563.3draw displayed formulae for alkanes with up to five carbon atoms in a molecule, and name the straight-chain isomers1503.4recall the products of the complete and incomplete combustion of alkanes157, 1653.5describe the substitution reaction of methane with bromine to form bromomethane in the presence of UV light.157Alkenes3.6recall that alkenes have the general formula CnH2n 1583.7draw 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)1523.8describe the addition reaction of alkenes with bromine, including the decolourising of bromine water as a test for alkenes.158Ethanol3.9describe the manufacture of ethanol by passing ethene and steamover a phosphoric acid catalyst at a temperature of about 300°C and a pressure of about 60–70 atm 159-1603.10describe the manufacture of ethanol by the fermentation of sugars, for example glucose, at a temperature of about 30°C159-1603.11evaluate the factors relevant to the choice of method used in the manufacture of ethanol, for example the relative availability of sugar cane and crude oil1603.12describe the dehydration of ethanol to ethene, using aluminium oxide.161Crude Oil5.6understand that crude oil is a mixture of hydrocarbons 1635.7describe and explain how the industrial process of fractional distillation separates crude oil into fractions5.8recall the names and uses of the main fractions obtained from crude oil: refinery gases, gasoline, kerosene, diesel, fuel oil and bitumen1645.9describe the trend in boiling point and viscosity of the main fractions1645.10understand 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 oxygen1655.11understand that, in car engines, the temperature reached is high enough to allow nitrogen and oxygen from air to react, forming nitrogen oxides1655.12understand that nitrogen oxides and sulfur dioxide are pollutant gases which contribute to acid rain, and describe the problems caused by acid rain1655.13understand 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 necessary1665.14describe 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.166-167Synthetic Polymers5.15understand that an addition polymer is formed by joining up many small molecules called monomers 1695.16draw the repeat unit of addition polymers, including poly(ethene), poly(propene) and poly(chloroethene)169-1715.17deduce the structure of a monomer from the repeat unit of an addition polymer169-1715.18describe some uses for polymers, including poly(ethene), poly(propene) and poly(chloroethene)1705.19explain that addition polymers are hard to dispose of as their inertness means that they do not easily biodegrade1705.20understand that some polymers, such as nylon, form by a different process called condensation polymerisation172-1735.21understand that condensation polymerisation produces a small molecule, such as water, as well as the polymer.172-173Acids, Alkalis and Salts4.1describe the use of the indicators litmus, phenolphthalein and methyl orange to distinguish between acidic and alkaline solutions 70-714.2understand how the pH scale, from 0–14, can be used to classify solutions as strongly acidic, weakly acidic, neutral, weakly alkaline or strongly alkaline704.3describe the use of universal indicator to measure the approximate pH value of a solution70-714.4define acids as sources of hydrogen ions, H+, and alkalis as sources of hydroxide ions, OH?784.5predict 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)714.6understand the general rules for predicting the solubility of salts in water:i all common sodium, potassium and ammonium salts are solubleii all nitrates are solubleiii common chlorides are soluble, except silver chlorideiv common sulfates are soluble, except those of barium and calciumv common carbonates are insoluble, except those of sodium, potassium and ammonium81-824.7describe experiments to prepare soluble salts from acids834.8describe experiments to prepare insoluble salts using precipitation reactions85-874.9describe experiments to carry out acid-alkali titrations.84-85Calculations (latest topic to June 2017)1.16calculate relative formula masses (Mr) from relative atomic masses (Ar) 176-1781.17understand the use of the term mole to represent the amount of substance179-1821.18understand the term mole as the Avogadro number of particles (atoms, molecules, formulae, ions or electrons) in a substance1821.19carry out mole calculations using relative atomic mass (Ar) and relative formula mass (Mr)179-1821.20understand 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.190-1921.21write word equations and balanced chemical equations to represent the reactions studied in this specification 33-391.22use the state symbols (s), (l), (g) and (aq) in chemical equations to represent solids, liquids, gases and aqueous solutions respectively33-391.23understand how the formulae of simple compounds can be obtained experimentally, including metal oxides, water and salts containing water of crystallisation1841.24calculate empirical and molecular formulae from experimental data182-1841.25calculate reacting masses using experimental data and chemical equations187-1891.26calculate percentage yield1931.27carry out mole calculations using volumes and molar concentrations.209-214Rates of Reaction4.17describe 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 41-504.18describe 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 reaction41-504.19understand the term activation energy and represent it on a reaction profile49-504.20explain 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 theory41-504.21explain that a catalyst speeds up a reaction by providing an alternative pathway with lower activation energy.49-50Energetics4.10understand 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 120-1234.11describe simple calorimetry experiments for reactions such as combustion, displacement, dissolving and neutralisation in which heat energy changes can be calculated from measured temperature changes202-2074.12calculate molar enthalpy change from heat energy change205-2074.13understand the use of ΔH to represent enthalpy change for exothermic and endothermic reactions205-2074.14represent exothermic and endothermic reactions on a simple energy level diagram202-2034.15understand that the breaking of bonds is endothermic and that the making of bonds is exothermic202-2034.16use average bond energies to calculate the enthalpy change during a simple chemical reaction.203Equilibria4.22understand that some reactions are reversible and are indicated by the symbol ? in equations 1254.23describe reversible reactions such as the dehydration of hydrated copper(II) sulfate and the effect of heat on ammonium chloride1254.24understand the concept of dynamic equilibrium125-1264.25predict the effects of changing the pressure and temperature on the equilibrium position in reversible reactions.127-1295.22understand that nitrogen from air, and hydrogen from natural gas or the cracking of hydrocarbons, are used in the manufacture of ammonia 133-1345.23describe the manufacture of ammonia by the Haber process, including the essential conditions:i a temperature of about 450°Cii a pressure of about 200 atmospheresiii an iron catalyst133-1345.24understand how the cooling of the reaction mixture liquefies the ammonia produced and allows the unused hydrogen and nitrogen to be recirculated133-1345.25describe the use of ammonia in the manufacture of nitric acid and fertilisers133-1345.26recall the raw materials used in the manufacture of sulfuric acid135-1365.27describe the manufacture of sulfuric acid by the contact process, including the essential conditions:i a temperature of about 450°Cii a pressure of about 2 atmospheresiii a vanadium(V) oxide catalyst135-1365.28describe the use of sulfuric acid in the manufacture of detergents, fertilisers and paints136Electrolysis1.48understand that an electric current is a flow of electrons or ions 112-1131.49understand why covalent compounds do not conduct electricity112-1131.50understand why ionic compounds conduct electricity only when molten or in solution112-1141.51describe experiments to distinguish between electrolytes and nonelectrolytes1161.52understand that electrolysis involves the formation of new substances when ionic compounds conduct electricity112-1151.53describe experiments to investigate electrolysis, using inert electrodes, of molten salts such as lead(II) bromide and predict the products113-1141.54describe experiments to investigate electrolysis, using inert electrodes, of aqueous solutions such as sodium chloride, copper(II) sulfate and dilute sulfuric acid and predict the products116-1181.55write ionic half-equations representing the reactions at the electrodes during electrolysis1171.56recall that one faraday represents one mole of electrons1961.57calculate the amounts of the products of the electrolysis of molten salts and aqueous solutions.197-2005.29describe the manufacture of sodium hydroxide and chlorine by the electrolysis of concentrated sodium chloride solution (brine) in a diaphragm cell 136-1375.30write ionic half-equations for the reactions at the electrodes in the diaphragm cell136-1375.31describe 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.1375.1Explain how the methods of extraction of the metals in this section are related to their positions in the reactivity series 1395.2describe 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 temperatureii the need to replace the positive electrodesiii the cost of the electricity as a major factor140-1415.3write ionic half-equations for the reactions at the electrodes in aluminium extraction1415.4describe and explain the main reactions involved in the extraction of iron from iron ore (haematite), using coke, limestone and air in a blast furnace1425.5explain the uses of aluminium and iron, in terms of their properties.141+143 ................
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