Chemistry 4XCH1 — specification explorer

1. Ionic bonding

• 5.1understand how ions are formed by electron loss or gainUnit 2
• 5.2

  know the charges of these ions:

  • metals in Groups 1, 2 and 3
  • non-metals in Groups 5, 6 and 7
  • Ag + , Cu 2+ , Fe 2+ , Fe 3+ , Pb 2+ , Zn 2+
  • hydrogen (H + ), hydroxide (OH – ), ammonium (NH 4 + ), carbonate (CO 3 2– ), nitrate (NO 3 - ), sulfate (SO 4 2– ).
  Unit 2
• 5.3write formulae for compounds formed between the ions listed aboveUnit 2
• 5.4draw dot-and-cross diagrams to show the formation of ionic compounds by electron transfer, limited to combinations of elements from Groups 1, 2, 3 and 5, 6, 7 only outer electrons need be shownUnit 2
• 5.5understand ionic bonding in terms of electrostatic attractionsUnit 2
• 5.6understand why compounds with giant ionic lattices have high melting and boiling pointsUnit 2
• 5.7know that ionic compounds do not conduct electricity when solid, but do conduct electricity when molten and in aqueous solutionUnit 2

2. Covalent bonding

• 5.8know that a covalent bond is formed between atoms by the sharing of a pair of electronsUnit 2
• 5.9understand covalent bonds in terms of electrostatic attractionsUnit 2
• 5.10

  understand how to use dot-and-cross diagrams to represent covalent bonds in:

  • diatomic molecules, including hydrogen, oxygen, nitrogen, halogens and hydrogen halides
  • inorganic molecules including water, ammonia and carbon dioxide
  • organic molecules containing up to two carbon atoms, including methane, ethane, ethene and those containing halogen atoms.
  Unit 2
• 5.11explain why substances with a simple molecular structures are gases or liquids, or solids with low melting and boiling points the term intermolecular forces of attraction can be used to represent all forces between moleculesUnit 2
• 5.12explain why the melting and boiling points of substances with simple molecular structures increase, in general, with increasing relative molecular massUnit 2
• 5.13explain why substances with giant covalent structures are solids with high melting and boiling pointsUnit 2
• 5.14explain how the structures of diamond, graphite and C 60 fullerene influence their physical properties, including electrical conductivity and hardnessUnit 2
• 5.15know that covalent compounds do not usually conduct electricityUnit 2

3. Metallic bonding

• 5.16Cknow how to represent a metallic lattice by a 2-D diagramUnit 2
• 5.17Cunderstand metallic bonding in terms of electrostatic attractionsUnit 2
• 5.18Cexplain typical physical properties of metals, including electrical conductivity and malleabilityUnit 2

4. Electrolysis

• 5.19Cunderstand why covalent compounds do not conduct electricityUnit 2
• 5.20Cunderstand why ionic compounds conduct electricity only when molten or in aqueous solutionUnit 2
• 5.21Cknow that anion and cation are terms used to refer to negative and positive ions respectivelyUnit 2
• 5.22Cdescribe experiments to investigate electrolysis, using inert electrodes, of molten compounds (including lead(II) bromide) and aqueous solutions (including sodium chloride, dilute sulfuric acid and copper(II) sulfate) and to predict the productsUnit 2
• 5.23Cwrite ionic half-equations representing the reactions at the electrodes during electrolysis and understand why these reactions are classified as oxidation or reductionUnit 2
• 5.24Cpractical: investigate the electrolysis of aqueous solutionsUnit 2

1. Group 1 (alkali metals) – lithium, sodium and potassium

• 6.1understand how the similarities in the reactions of these elements with water provide evidence for their recognition as a family of elementsUnit 2
• 6.2understand how the differences between the reactions of these elements with air and water provide evidence for the trend in reactivity in Group 1Unit 2
• 6.3use knowledge of trends in Group 1 to predict the properties of other alkali metalsUnit 2
• 6.4Cexplain the trend in reactivity in Group 1 in terms of electronic configurationsUnit 2

2. Group 7 (halogens) – chlorine, bromine and iodine

• 6.5know the colours, physical states (at room temperature) and trends in physical properties of these elementsUnit 2
• 6.6use knowledge of trends in Group 7 to predict the properties of other halogensUnit 2
• 6.7understand how displacement reactions involving halogens and halides provide evidence for the trend in reactivity in Group 7Unit 2
• 6.8Cexplain the trend in reactivity in Group 7 in terms of electronic configurationsUnit 2

3. Gases in the atmosphere

• 6.9know the approximate percentages by volume of the four most abundant gases in dry airUnit 2
• 6.10understand how to determine the percentage by volume of oxygen in air using experiments involving the reactions of metals (e.g. iron) and non-metals (e.g. phosphorus) with airUnit 2
• 6.11describe the combustion of elements in oxygen, including magnesium, hydrogen and sulfurUnit 2
• 6.12describe the formation of carbon dioxide from the thermal decomposition of metal carbonates, including copper(II) carbonateUnit 2
• 6.13know that carbon dioxide is a greenhouse gas and that increasing amounts in the atmosphere may contribute to climate changeUnit 2
• 6.14practical: determine the approximate percentage by volume of oxygen in air using a metal or a non-metalUnit 2

4. Chemical tests

• 6.15

  describe tests for these gases:

  • hydrogen
  • oxygen
  • carbon dioxide
  • ammonia
  • chlorine.
  Unit 2
• 6.16describe how to carry out a flame testUnit 2
• 6.17

  know the colours formed in flame tests for these cations:

  • Li + is red
  • Na + is yellow
  • K + is lilac
  • Ca 2+ is orange-red
  • Cu 2+ is blue-green.
  Unit 2
• 6.18

  describe tests for these cations:

  • NH 4 + using sodium hydroxide solution and identifying the gas evolved
  • Cu 2+ , Fe 2+ and Fe 3+ using sodium hydroxide solution.
  Unit 2
• 6.19

  describe tests for these anions:

  • Cl – , Br – and I – using acidified silver nitrate solution
  • SO 4 2– using acidified barium chloride solution
  • CO 3 2– using hydrochloric acid and identifying the gas evolved.
  Unit 2
• 6.20describe a test for the presence of water using anhydrous copper(II) sulfateUnit 2
• 6.21describe a physical test to show whether a sample of water is pureUnit 2

1. Rates of reaction

• 7.1describe experiments to investigate the effects of changes in surface area of a solid, concentration of a solution, temperature and the use of a catalyst on the rate of a reactionUnit 2
• 7.2describe the effects of changes in surface area of a solid, concentration of a solution, pressure of a gas, temperature and the use of a catalyst on the rate of a reactionUnit 2
• 7.3explain the effects of changes in surface area of a solid, concentration of a solution, pressure of a gas and temperature on the rate of a reaction in terms of particle collision theoryUnit 2
• 7.4know that a catalyst is a substance that increases the rate of a reaction, but is chemically unchanged at the end of the reactionUnit 2
• 7.5know that a catalyst works by providing an alternative pathway with lower activation energyUnit 2
• 7.6Cdraw and explain reaction profile diagrams showing Δ H and activation energyUnit 2
• 7.7practical: investigate the effect of changing the surface area of marble chips and of changing the concentration of hydrochloric acid on the rate of reaction between marble chips and dilute hydrochloric acidUnit 2
• 7.8practical: investigate the effect of different solids on the catalytic decomposition of hydrogen peroxide solutionUnit 2

2. Reversible reactions and equilibria

• 7.9know that some reactions are reversible and this is indicated by the symbol ⇌ in equationsUnit 2
• 7.10describe reversible reactions such as the dehydration of hydrated copper(II) sulfate and the effect of heat on ammonium chlorideUnit 2
• 7.11Cknow that a reversible reaction can reach dynamic equilibrium in a sealed containerUnit 2
• 7.12C

  know that the characteristics of a reaction at dynamic equilibrium are:

  • the forward and reverse reactions occur at the same rate
  • the concentrations of reactants and products remain constant.
  Unit 2
• 7.13Cunderstand why a catalyst does not affect the position of equilibrium in a reversible reactionUnit 2
• 7.14C

  know the effect of changing either temperature or pressure on the position of equilibrium in a reversible reaction:

  • an increase (or decrease) in temperature shifts the position of equilibrium in the direction of the endothermic (or exothermic) reaction
  • an increase (or decrease) in pressure shifts the position of equilibrium in the direction that produces fewer (or more) moles of gas References to Le Chatelier's principle are not required
  Unit 2

1. Alcohols

• 8.1Cknow that alcohols contain the functional group −OHUnit 2
• 8.2Cunderstand how to draw structural and displayed formulae for methanol, ethanol, propanol (propan-1-ol only) and butanol (butan-1-ol only), and name each compound the names propanol and butanol are acceptableUnit 2
• 8.3C

  know that ethanol can be oxidised by:

  • burning in air or oxygen (complete combustion)
  • reaction with oxygen in the air to form ethanoic acid (microbial oxidation)
  • heating with potassium dichromate(VI) in dilute sulfuric acid to form ethanoic acid
  Unit 2
• 8.4C

  know that ethanol can be manufactured by:

  • reacting ethene with steam in the presence of a phosphoric acid catalyst at a temperature of about 300 ºC and a pressure of about 60–70 atm
  • the fermentation of glucose, in the absence of air, at an optimum temperature of about 30 ºC and using the enzymes in yeast
  Unit 2
• 8.5Cunderstand the reasons for fermentation, in the absence of air, and at an optimum temperatureUnit 2

2. Carboxylic acids

• 8.6Cknow that carboxylic acids contain the functional groupUnit 2
• 8.7Cunderstand how to draw structural and displayed formulae for unbranched-chain carboxylic acids with up to four carbon atoms in the molecule, and name each compoundUnit 2
• 8.8Cdescribe the reactions of aqueous solutions of carboxylic acids with metals and metal carbonatesUnit 2
• 8.9Cknow that vinegar is an aqueous solution containing ethanoic acidUnit 2

3. Esters

• 8.10Cknow that esters contain the functional groupUnit 2
• 8.11Cknow that ethyl ethanoate is the ester produced when ethanol and ethanoic acid react in the presence of an acid catalystUnit 2
• 8.12Cunderstand how to write the structural and displayed formulae of ethyl ethanoateUnit 2
• 8.13Cunderstand how to write the structural and displayed formulae of an ester, given the name or formula of the alcohol and carboxylic acid from which it is formed and vice versaUnit 2
• 8.14Cknow that esters are volatile compounds with distinctive smells and are used as food flavourings and in perfumesUnit 2
• 8.15Cpractical: prepare a sample of an ester such as ethyl ethanoateUnit 2

4. Synthetic polymers

• 8.16know that an addition polymer is formed by joining up many small molecules called monomersUnit 2
• 8.17understand how to draw the repeat unit of an addition polymer, including poly(ethene), poly(propene), poly(chloroethene) and (poly)tetrafluoroetheneUnit 2
• 8.18understand how to deduce the structure of a monomer from the repeat unit of an addition polymer and vice versaUnit 2
• 8.19

  explain problems in the disposal of addition polymers, including:

  • their inertness and inability to biodegrade
  • the production of toxic gases when they are burned.
  Unit 2
• 8.20Cknow that condensation polymerisation, in which a dicarboxylic acid reacts with a diol, produces a polyester and waterUnit 2
• 8.21Cunderstand how to write the structural and displayed formula of a polyester, showing the repeat unit, given the formulae of the monomers from which it is formed including the reaction of ethanedioic acid and ethanediol:Unit 2
• 8.22Cknow that some polyesters, known as biopolyesters, are biodegradableUnit 2