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Pure Substances and Mixtures
Gas, liquid and solid interconversions.
Pure substances and mixtures.
Solvent, solute, solution. Suspension.
Separation of mixtures by filtration, evaporation to dryness and crystallization (partial evaporation up to the point of crystallization), simple distillation, paper chromatography, sublimation and fractional distillation. The use of the separating funnel.
Boiling point and melting point of substances as criteria of purity.
Decomposition of hydrates, carbonates, hydrogencarbonates, nitrates and hydroxides by heat.
Thermally stable compounds.
Reversible reactions, including the hydration/dehydration of silica gel and of hydrated copper (ll) sulfate.
Gases in air
Handling techniques for preparation and collection of gases, including the use of the gas syringe.
Oxygen
The atmosphere – composition of: to include presence of water vapour, noble gases and carbon dioxide. An experimental determination of the percentage composition by volume of nitrogen and oxygen in air.
Principle of the extraction of oxygen by fractional distillation of liquid air
Preparation of oxygen by the catalytic decomposition of hydrogen peroxide; test for oxygen.
Reaction of metals and non-metals with oxygen in air.
Types of oxide. Oxides of the common elements and their reaction with water.
Rusting of iron and its prevention.
Ozone as an unstable, naturally occurring, allotrope of oxygen.
Water and solutions
Simple treatment of soluble and insoluble substances in water. Saturated solutions and interpretation of solubility curves.
Physical properties; test for purity. Chemical tests for the presence of water.
Water of crystallisation, deliquescence and efflorescence; hygroscopic substances.
The presence of air dissolved in water. Students should be aware that air dissolved in water has a different composition from ordinary air due to the difference in solubility of nitrogen and oxygen.
Ground water and sea water as important examples of aqueous solutions of specific solutes. Methods of purifying water, including a simple qualitative treatment of the technique of reverse osmosis for purifying brackish water and seawater.
Acids, bases and salts
Aqueous solutions of acidic and alkaline substances and their action on indicators. The pH scale treated as an arbitrary scale of acidity and alkalinity. Strong and weak acids.
Bases and alkalis.
Reaction of dilute non-oxidising acids with metals, insoluble bases and alkalis, carbonates and hydrogencarbonates, sulfites.
Normal salts and acid salts; preparation of salts.
Comparison of solutions of hydrogen chloride in water and in methylbenzene.
Standard solutions, acid-alkali titrations and related calculations in terms of moles and molar concentrations.
Hydrogen
Preparation of hydrogen from action of dilute non-oxdising acids on certain metals, exemplified by dilute hydrochloric acid or dilute sulfuric acid on magnesium, zinc or iron.
Test for hydrogen.
Combustion of hydrogen – its advantages and disadvantages as a fuel.
Reducing action of hydrogen with metal oxides.
Uses of hydrogen.
Carbon
Diamond and graphite as allotropes of carbon.
Reaction of carbon with metal oxides.
Laboratory preparation of carbon dioxide; test for carbon dioxide.
Solubility in water and in alkali. Uses of carbon dioxide.
Carbonates – some general properties limited to: solubility, reaction with dilute acids, action of heat, preparing insoluble carbonates by precipitation.
Formation of carbon dioxide and carbon monoxide from processes of complete / incomplete combustion of carbon and hydrocarbons.
Properties of carbon monoxide – neutral gas, toxic nature, combustion to form carbon dioxide; reducing action. Separation of carbon monoxide from a carbon monoxide / carbon dioxide mixture by absorption in alkali.”
Nitrogen
Nitrogen as an unreactive gas.
Principle of the industrial extraction of nitrogen from liquid air.
Manufacture of ammonia by reversible direct union of its constituent elements.
Laboratory preparation of ammonia from ammonium salts; test for ammonia gas.
Properties of ammonia gas. Use of ammonia as a reducing agent for certain metallic oxides.
Preparation and properties of aqueous ammonia – e.g. alkalinity, neutralisation of acids, precipitation of insoluble metallic hydroxides.
Ammonium salts – preparation by neutralisation; reaction with alkalis; use as fertilisers.
Sublimation / thermal dissociation of ammonium chloride.
Nitric acid as a dilute acid – exemplified by its reaction with metallic oxides, e.g. magnesium oxide, and with metallic carbonates, e.g. magnesium carbonate. Nitric acid as an oxidising agent – exemplified by its reaction with copper. Uses of nitric acid.
Nitrates – general properties, e.g. solubility, action of heat, general methods of preparation.
Nitrogen monoxide – conversion to nitrogen dioxide on exposure to air.
Nitrogen dioxide – laboratory preparation by the thermal decomposition of lead (ll) nitrate; identification due to colour. Properties of nitrogen dioxide – e.g. solubility in water and acidity (linked to acid rain).”
Allotropes of sulfur – rhombic / monoclinic. Uses of the element.
Hydrogen sulfide – its formation by the action of dilute acids on metal sulfides; toxic nature of the gas.
Laboratory preparation of sulfur dioxide by the action of a dilute acid on a sulfite; test for the gas. Reactions of sulfur dioxide which exemplify its acidic nature and its reducing action, e.g. colour change with acidified potassium dichromate solution.
Sulfur trioxide – acidic nature.
Manufacture of sulfuric acid and its uses.
Properties of dilute sulfuric acid. Properties of concentrated sulfuric acid – including its action on metallic chlorides; dehydrating action, e.g. on sugar and on ethanol; oxidizing action e.g. on copper; and hygroscopic nature / use as a drying agent.”
Chlorine, bromine and iodine – similarities and trends in properties of elements in Group 7.
Displacement reactions of one halogen by another.
Laboratory preparation of chlorine by oxidation of concentrated hydrochloric acid using manganese (lV) oxide.
Test for chlorine.
Bleaching action of chlorine water. Uses of chlorine.
Properties of dilute hydrochloric acid.
Preparation of chlorides from dilute hydrochloric acid.
Electrolysis
The effect of electricity on various solids, liquids (including molten substances), and aqueous solutions. Conductors and non-conductors; electrolytes and non-electrolytes.
Reactivity Series
Detailed description of electrolytic decomposition and the factors affecting product formation at the electrodes. Factors will include – the position of the ion in the reactivity series, concentration of ion and the nature of the electrodes.
Quantitative aspects of electrolytic cells: the Faraday as a mole of electrons.
Reactivity series: to include potassium, sodium, calcium, magnesium, aluminium, zinc, iron, lead, hydrogen, copper and silver. Reactivity of these metals with air (oxygen), water and dilute acids.
Displacement reactions involving these metals and their compounds to include: a metal displacing a less reactive metal from an aqueous metallic salt; a metal reducing an oxide of a less reactive metal.
The simple cell – as a means of transforming chemical energy into electrical energy (e.g. Zn/Cu electrodes in dilute acid).
Group 1 and Group 2 metals
Alkali metals and alkaline earths as representatives of two groups or families of elements in the Periodic Table.
Sodium and potassium. Characteristics of the metals and similarities in the group: typical physical properties; chemical properties – reaction with oxygen (to form simple oxide, M2O, only), with water and with chlorine. Trend in reactivity going down the group.
Magnesium and calcium. Similarities in the group: typical physical properties; chemical properties – reaction with oxygen, water and dilute acids. Limestone – conversion of limestone (CaCO3) to quicklime (CaO) and subsequently to slaked lime [Ca(OH)2].
Hardness in water caused by dissolved calcium and magnesium salts. Hardness in ground water associated with limestone terrains. Temporary and permanent hardness of water; softening of water.
The advantages of synthetic detergents over soap when used with hard water – simple treatment only. Scale formation, stalactites and stalagmites.
Less reactive metals: iron and copper
These metals are to be presented as typical of the transition elements, illustrating the properties of variable valency, the formation of coloured compounds and acting as catalysts.
Iron
Action of steam, hydrogen chloride and chlorine on iron.
Hydroxides of iron (ll) and iron (lll) : formation by precipitation, colour.
Oxidation of iron (ll) hydroxide to iron (lll) hydroxide by exposure to air.
Copper
Simple compounds of copper: Copper (ll) oxide as a typical basic oxide and its use in preparing copper (ll) salts by reaction with dilute acids.
Reduction of copper (ll) oxide by hydrogen.
Thermal decomposition of copper (ll) carbonate and copper (ll) nitrate to give copper (ll) oxide.
Qualitative analysis
Identification of sodium and potassium ions by flame tests.
Simple test tube reactions for the identification of the following ions in solution:
Cations – ammonium, calcium, magnesium, aluminium, lead (ll), copper (ll), iron (ll) and iron (lll);
A flame test can be used to distinguish between calcium and magnesium. Amphoteric character of aluminium and lead (ll) to be exploited in testing for these ions. Potassium iodide solution can be used to distinguish between aluminium and lead (ll) ions.
Anions – carbonate, sulfite, sulfate, chloride, bromide, iodide, and nitrate. Nitrates can be tested by reduction with aluminium and alkali to give ammonia.
Tests for the following gases: oxygen, hydrogen, carbon dioxide, ammonia, chlorine, hydrogen chloride and sulfur dioxide.
Organic chemistry
Definition of an organic compound. The unique ability of carbon to catenate leading to a large number of organic compounds.
Concept of a homologous series.
Alkanes as the first example of a homologous series: used to illustrate the terms empirical formula, molecular formula, structural formula and general formula for a homologous series.
Nomenclature limited to the first five straight chain alkanes. Chain isomerism of C4H10 and C5H12 only.
Alkyl groups limited to methyl and ethyl only.
Gradation in physical properties of straight chain alkanes linked to length of hydrocarbon chain.
Complete and incomplete combustion of hydrocarbons resulting in the formation of carbon dioxide, carbon monoxide and carbon.
Uses of alkanes as fuels.
Saturated nature of the alkanes resulting in substitution reactions with halogens, limited to monosubstitution.
Alkenes and alkynes as typical unsaturated hydrocarbons. General formula; structural formulae.
Combustion of alkenes / alkynes – sootiness of flame as an indication of unsaturation (or high percentage by mass of carbon).
Addition reactions of alkenes with hydrogen and halogens; and of ethene with hydrogen halide. Hydration of ethene.
Test for unsaturation – distinction between alkanes and alkenes/alkynes using bromine water.
Principle of addition polymerisation limited to polyethene, PTFE and PVC.
Petroleum (crude oil) as a mixture of hydrocarbons including natural gas which can be separated by fractional distillation (fractionation); limited to names of fractions and the fact that boiling point range of fraction is related to carbon number. Uses of fractions.
Cracking of long chain alkanes to form one shorter chain alkane and ethene. The cracking of diesel to obtain petrol and ethene as an industrial application of this process.
OH as the functional group. General formula, structural formulae; nomenclature.
Reactions of alcohols with sodium, with phosphorus (V) chloride to liberate misty fumes as a test for the OH group and with concentrated sulfuric acid to form alkenes.
Ethanol; manufacture from glucose (fermentation) and from petroleum derived ethene (hydration). Oxidation of ethanol to ethanoic acid using acidified potassium dichromate, limited to test-tube reaction.
Functional group isomerism limited to C2H6O.
COOH as the functional group. Naming of straight chain carboxylic acids up to pentanoic acid. Weak acidic nature.
Formation of salts by the usual methods.
Reaction of an organic acid and an alcohol to produce an ester and water in a reversible process. Role of concentrated sulfuric acid in esterification. Recognition of the esters as another homologous series characterised by their fruity smell.
Atoms, molecules, Relative atomic mass, Relative molecular and formula mass.
Atomic nature of matter. Elements and compounds. Idea of size of atoms. Avogadro’s constant and moles of atoms. Relative atomic mass with respect to 12C.
Molecules and relative molecular masses, expressed in grammes, as the mass of a mole of molecules. Mole of ions. Formula masses of ionic compounds. Mole/mass interconversions.
Molar volume of gases and Avogadro’s law. Mole/mass/volume interconversions for gases. Calculations to find the volume of a gas (measured at stp) that reacts, or is produced, in chemical reactions.
Use of PV/T = constant for converting gas volumes to and from standard conditions.
Gas to gas calculations based on Gay Lussac’s law of combining volumes.
Chemical formulae and equations
Experimental determination of chemical formulae of binary compounds, from the reacting masses of the elements. (Oxides of metals offer suitable examples). Deriving the value of xH2O in a hydrated compound by heating to constant mass. Percentage composition and related calculations.
Balanced chemical equations to represent the relative number of particles involved in chemical reactions. States of substances, [using symbols (s), (l), (g) and (aq) for solid, liquid, gas and aqueous solution respectively], should be specified where appropriate.
Experimental determinations intended to establish the combining ratios of reactants and products to include gravimetry and volumetric work, in addition to measurements involving gas volumes.
All calculations involving chemical changes are to be performed in terms of moles of substance.”
Atomic structure and the Periodic Table
Nuclear model of the atom; protons, neutrons and electrons. Isotopes and relation of isotopy to relative atomic masses (limited to two isotopes).
Electrons in shells. Electronic configuration of the first eighteen elements (hydrogen through to argon). Relation of electronic configuration to electrovalency and covalency, and to the periodic property of valency / position of element in the Periodic Table.”
Structure and Bonding
Ionic bonding – formation of simple ions by loss or gain of electrons as governed by the octet rule. Physical properties of ionic compounds – high melting points, solubility in water, conductivity. Electrostatic attractions and three dimensional lattice of ions, limited to sodium chloride.
Covalent bonding – formation of simple molecules (e.g. H2, Cl2, O2, N2, HCl, H2O, NH3, CH4, CO2) by sharing of electrons as governed by the octet rule. Single, double and triple bonds.
Giant molecular structures, limited to diamond and graphite. High sublimation temperatures of these materials explained in terms of strong covalent bonds holding the lattice structure together. Electrical conductivity in graphite explained in terms of free electrons.
Distinction between physical properties of giant covalent structures and crystals composed of simple molecules. The latter to be considered as discreet molecules held together by weak intermolecular forces.
Dry ice and iodine are suitable examples.
Metallic crystals in terms of ions in a sea of mobile electrons. Thermal and electrical conductivity, and malleability, explained in terms of this bonding model.
Kinetic molecular theory and states of matter
Diffusion and Brownian motion in terms of simple kinetic theory.
Interconversions between the three states of matter.
Application of kinetic theory to explain energy requirements for changes of state; different energy requirements for simple molecular versus giant molecular lattices in relation to the nature of bonding (simple treatment only).
Qualitative treatment of the effect of change in temperature, or pressure or volume on a fixed mass of gas.”
Ionic theory; oxidation and reduction
Ionic half equations to represent synthesis reactions involving binary compounds, displacement reactions and reactions at electrodes for electrolytes given in section
Ionic equations omitting spectator ions for – neutralization, acid on a carbonate, acid on a sulfite, alkali on an ammonium salt and precipitation reactions.
Oxidation and reduction; redox reactions in terms of loss and gain of oxygen/hydrogen; in terms of loss and gain of electrons.
The concept of oxidation number limited to simple binary compounds between metals and non-metals. Its application in simple redox reactions.
Electrolysis explained in terms of electron transfer between ions and electrodes. Tendency to lose or accept electrons related to the reactivity series.”
Energetics
The concept that energy changes accompany both physical and chemical changes. Exothermic and endothermic reactions; thermochemical equations and the ΔH notation. The Joule or kiloJoule as a unit of measuring heat energy. Definitions and calculations for energy changes accompanying combustion, solution, neutralisation and precipitation reactions. Simple experimental determination of heat of solution and heat of neutralisation and estimation of the heat of combustion of a flammable liquid.
Heats of reaction as the result of energy changes when bonds are broken and formed. Principle of conservation of energy.
Simple energy level diagrams illustrating the idea of activation energy.
Rates of reactions
Concept of reaction rate as the increase of product concentration or decrease of reactant concentration with time. Dependence of rate of heterogeneous reactions on the state of subdivision of a solid explained in terms of particle collisions involved in chemical change.
A simple kinetic picture to be used to account for the effect of concentration on the rate of homogeneous reactions. Effect of temperature. Definition of a catalyst. The effect of a catalyst on reaction rate.
Effect of light on certain reactions, e.g. photoreduction of silver halides; reaction of hydrogen with chlorine.”
Reversible reactions and chemical equilibrium
The concept of reversible change as exemplified by various processes, e.g. changes of state; hydration of copper (ll) sulfate; chromate/dichromate interconversions; esterification. Kinetic picture of dynamic equilibrium and use of the appropriate symbol to denote reversible reactions.
Le Chatelier’s principle and its application to systems in dynamic equilibrium – qualitative treatment only.
Shifting the equilibrium position of a system by changing the temperature, the total pressure or the concentration of a species.
Thermal dissociation / dynamic equilibrium of dinitrogen tetraoxide / nitrogen dioxide.
The effect of a catalyst on the rate of attainment of equilibrium and not on the equilibrium position.
Raw materials and energy requirements
Reference to the raw materials available for chemical processing as related to terrestrial abundance and ease of extraction from the source. Sources of materials: air, sea water, rock salt, limestone, iron ore, sulfur, gypsum and natural gas. Awareness that petroleum is a source of many useful organic products such as plastics, textiles, pharmaceuticals, dyes, explosives, etc.
Coal and petroleum as non-renewable energy sources.
Extraction of metals
Extraction of iron from haematite and aluminium from purified bauxite. Method of extraction related to position of metal in the activity series. Electrolytic purification of copper.
Uses of these metals as related to their properties.
The heavy chemicals industry
The Haber process for ammonia and the Contact process for sulfuric acid.
Main uses of ammonia and sulfuric acid.
Commonplace products of the chemical industry
Reference should be made to products of chemical industry which are typically found in the home and other familiar environments.
The use of the following materials and associated properties should be discussed: action of soaps and detergents on hard water; domestic bleaches – their alkalinity and oxidising action exemplified by the liberation of iodine from potassium iodide solution; domestic LPG gas (mainly butane) and its flammability; the use of organic liquids as solvents, exemplified by alcohols, liquid alkanes and ethyl ethanoate. Baking soda (sodium hydrogencarbonate) and its leavening action; vinegar (ethanoic acid) and its acidity; quicklime (calcium oxide) and slaked lime (calcium hydroxide) and their alkalinity; caustic soda (NaOH); washing soda (Na2CO3.10H2O); Milk of Magnesia (magnesium hydroxide); Plaster of Paris (CaSO4.½H2O).
Polythene, polyvinylchloride (PVC) and polytetrafluoroethene (PTFE) should be mentioned as important examples of synthetic polymers with various uses.”
Pollutant gases
An appreciation of the problems posed to the natural environment by activities which involve the use of materials and their chemical interconversions.
Carbon dioxide formation as a result of combustion processes of carbon- containing fossil fuels. The greenhouse effect and its increase leading to global warming. Recognition of the fact that most activities related to energy generation ultimately result in carbon dioxide emissions from power plants.
An awareness that incomplete combustion of fossil fuels will result in the formation of toxic carbon monoxide.
Sulfur dioxide as a pollutant gas; its source and presence in acid rain; harmful effects of acid rain.
NOx gases: nitrogen oxides formed in the internal combustion engine; contribution of nitrogen oxides to acid rain and to smog formation in urban polluted air. A simple treatment of the catalytic converter.
Importance of the ozone layer in the upper atmosphere and its depletion by
CFC’s. Awareness of harmful effect of ozone near the Earth’s surface.