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F.6 Chemistry Teaching Schedule (2000-01)
Textbook: New Way Chemistry for Hong Kong A-level (Book1, 2 and 3)
Cycle Date Syllabus Explanatory Notes References in texts Remark 1. Atom, Molecules and
Stoichiometry1.1 The atomic structure
1.2 Radioactivity
1.3 Relative isotopic, atomic andmolecular masses
1.4 The mole concept
Protons, neutrons and electrons as constituents of the atom.
The relative masses and charges of a proton, neutron and electron.
The atomic nucleus. Relative size of the atom and atomic nucleus.
Nature of , particles, and ofradiation.
Equations for nuclear reactions.
Uses of isotopes in leak detection, radiotherapy, nuclear power and as
tracers. (Underlying principles and instrumentation are not required.)
A brief account of the mass spectrometer in determining relative
isotopic, atomic and molecular masses (instrumental details and
mathematical treatment of the mass spectrometer, and the use of
fragmentation in structure determination are not required.)
The mole and the Avogadro constant.
Molar volume of gases at R.T.P. (room temperature and pressure) and
S.T.P. (standard temperature and pressure). Ideal gas equation, pV=nRT
and its application to the relative molecular mass determination.
Chapter 1
Chapter 2
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1.5 The Faraday and the mole
1.6 Empirical and molecular formulae
1.7 Chemical equations and stoichiometry
(Non-ideal behaviour of real gases and kinetic theory are not required.)
Partial pressure of gas and its relationship to mole fraction.
The Faraday as the quantity of electricity of one mole of electrons.
Relationship between the mass liberated and the quantity of electricity
passed in electrolysis.
Derivation of empirical formula using combustion data or composition
by mass. Molecular formula derived from empirical formula and relative
molecular mass.
The stoichiometric relationship between reactants and products in a
reaction.
Calculation involving
i. reacting masses
ii. volumes of gases, and
iii. concentrations and volumes of solutions
Chapter 3
2. The Electronic Structure of Atoms
and the Periodic Table
2.1 Atomic emission spectra and
electronic structure of atomsCharacteristics of the emission spectrum of atomic hydrogen.
Interpretation of the spectrum using the relationship, E=hleading to the
idea of discrete energy levels.
Convergence limits and ionization. (Calculation are not required).
An awareness of the uniqueness of atomic emission spectra.
Chapter 4
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2.2 Electronic structure, ionization
enthalpies, electron shell
2.3 Atomic orbitals
2.4 Electronic configurations of atoms
Electronic configurations in relation to
Plots of the following graphs to introduce shells and sub-shells:
i. successive ionization enthalpies for a particular element, and
ii. first ionization enthalpies against atomic numbers (up to Z=20).
(Experimental determination of ionization enthalpy is not required).
An awareness of the wave nature of electrons, and that electrons are not
localized in fixed orbits. An atomic orbital as a representation of a region
within which there is a high probability of finding an electron. The
designation of s, p and d orbitals. The number and relative energies of
the s, p and d orbitals for the principal quantum numbers 1,2 and 3, and
also of 4s and 4p orbitals. Shapes of s and p orbitals only.
(The uncertainty principle is not required.)
Building up of electronic configurations based on three principles:i. electrons enter the orbitals in order of ascending energy
(Aufbau principle),
ii. orbitals of the same energy must be occupied singly before
pairing occurs (Hunds rule), and
iii. electrons occupying the same orbital must have opposite spins
(Paulis exclusion principle).
Electron configurations of isolated atoms from H to Kr.
Chapter 5
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the Periodic Table
2.5 The Periodic Table and the atomic
properties of the elements
Electron configurations of atoms represented by
i. Notation using 1s, 2s, 2p, etc., e.g. Fe (ground state)
1s22s2sp63s23p63d64s2
ii. electrons-in-boxes diagram
The Periodic Table, showing the s-, p-, d- and f-blocks. Interpretation of
the trends of ionization enthalpies and atomic radii of the elements in the
Periodic Table.
3. Energetics
3.1 Energy changes in chemical reactions
3.2 Standard enthalpy changes
3.3 Hesss law
Conservation of energy. Endothermic and exothermic reactions and their
relationship to the breaking and forming of bonds.
Enthalpy change, H, as heat change at constant pressure.
Standard enthalpy change of:
i. neutralization,
ii. solution,
iii. formation, and
iv. combustion
Experimental determination of enthalpy changes of reactions, limited to
simple calorimetric method. (Bomb carlorimeter is not required).
Use of Hesss law to determine enthalpy changes that are not easily
obtainable by experiment.
Enthalpy level diagrams.
Chapter 6
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Calculations involving enthalpy changes of reactions. Mid-term test
4. Bonding and Structure
4.1 The nature of forces holding atoms
together
4.2 Ionic bonding
Energetics of formation of ionic
compounds
Stoichiometry of ionic compounds
Ionic crystals
Electrostatic interactions between electrons and nuclei leading to
different types of bonding.
Formation of ions the tendency for atoms of elements in Groups I, II,
VI and VII to attain electronic configurations of noble gases.
Dot and cross diagrams for simple ionic compounds.
Born-Haber cycles for the formation of ionic compounds in terms of
enthalpy changes of atomization and ionization, electronic affinities and
lattice enthalpies.
(Electronic affinity is the enthalpy change when one mole of electrons is
added to one mole of atoms or ions in the gaseous state,
e.g. O(g) + e-O-(g) H = -141 kJ mol-1
O-(g) + e- O2-(g) H= +791 kJ mol-1;
Lattice enthalpy is the enthalpy change when one mole of an ionic
compound is formed from its constituent ions in the gaseous state,
e.g. Na+(g) + Cl-(g) NaCl(s) H = -781 kJ mol-1
Consideration in terms of electronic configurations and enthalpy changes
of formation.
Extended three-dimensional structures of ionic compounds limited to
Chapter 7
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Ionic radii
4.3 Covalent bonding
Dative covalent bonding
Bond enthalpies, bond lengths and
covalent radii
sodium chloride and caesium chloride. Unit cells and coordination
numbers.
(Calculations involving ionic radii in a unit cell are not required.)
Comparison of sizes of ions with their parent atoms.
Comparison of sizes of isoelectronic particles.
Formation of covalent bonding sharing of electron pairs.
The simple idea of the overlapping of atomic orbitals.
Dot and cross diagrams for simple molecules, e.g. CH4, NH3, H2O, HF.
Octet rule and its limitation, e.g. PCl5 and BF3.
Treated as a special example of covalent bonding, illustrated by
H3N->BF3. The simple idea of the overlapping of an empty orbital with
an orbital occupied by a lone pair of electrons.
Estimation of bond enthalpies using data from energetics.
Bond enthalpies as a comparison of the strength of covalent bonds.
Relationship between covalent bond enthalpies and bond lengths as
illustrated by hydrogen halides.
Addition of covalent radii to give approximate covalent bond lengths as
illustrated by simple molecules.
Chapter 8
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The shapes of covalent molecules
and polymeric ion
Multiple bonds
Covalent crystals
4.4 Bonding intermediate between ionic
and covalent
Incomplete electron transfer in ionic
compounds
Polarity of covalent bond
The shapes of simple molecules and polyatomic ions explained in terms
of the repulsion between electron pairs(as illustrated by BF3, CH4, H2O,
PCl5, SF6, NH4+ and NH2
-). The directional nature of covalent bonds.
Bond angles.
Comparison of bond lengths and bond enthalpies leading to the idea of
multiple bonds, illustrated by ethane and ethyne. Shapes of carbon
dioxide and sulphur dioxide molecules explained in terms of repulsion
between electron pairs.
Exemplified by diamond, graphite and quartz.
Comparison of the experimental lattice enthalpies of e.g. silver halides
and zinc sulphide, with the theoretical values calculated on a completely
ionic model leading to the idea of polarization of ions. (Calculation of
the theoretical value of lattice enthalpy is not required).
Displacement of an electron cloud leading to the formation of a polar
covalent bond. Dipole moment as evidence for bond polarization in
simple molecules. (Calculation of dipole moment is not required.)
Chapter 9
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4.5 Metallic bonding
Metallic crystals
4.6 Intermoleular forces
Van der Waals forces
Unequal sharing of bonded electron pair(s) explained in terms of the
electronegativity difference between bonded atoms.
Electronegativity (Paulings scale) introduced as an arbitrary measure of
an atoms tendency in a molecule to attract electrons. (The formal
definition of electronegativity and its experimental determination are not
required.)
Metallic bonding illustrated by a model of cationic lattice and mobile
valence electrons. Simple explanation of the metallic conduction of
electricity based on the model.
Strength of metallic bond in terms of metallic radii and the number of
valence electron(s) per atom.
Cross-packed and open structures: hexagonal and cubic close-packed,
and body-centred cubic structures. Unit cells and coordination numbers.(Calculations related to atomic radii in a unit cell are not required.)
Brief discussion of the origin of van der Waals forces in terms of
permanent, instantaneous and induced dipoles. Comparison of the
covalent and van der Waals radii of non-metals to indicate the relative
strength of the covalent bonds and van der Waals forces.
Chapter 10
Chapter 11
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Molecular crystals
Hydrogen bonding
4.7 The relationship between structures
and properties of materials
Exemplified by iodine and carbon dioxide.
A study of the boiling points and enthalpy changes of vaporization of the
hydrides of Groups IV, V, VI and VII and compounds like alcohols and
carboxylic acids leading to the idea of hydrogen bonding.
Nature of hydrogen bonding.
Relative strength of van der Waals forces and hydrogen bonding.
Hydrogen bonding in ice, proteins and DNA (deoxyribonucleic acid)
Differences in physical properties (viz. melting and boiling points,
electrical conductivity, hardness and solubility) between ionic
compounds, covalent substances and metals
Chapter 12
5. Chemical Kinetics
5.1 Rates of chemical reactions
5.2 Factors influencing reaction rate
5.3 Rate equations and order of reactions
The meaning of the rate of a chemical reaction.
Following a reaction by chemical and physical methods, viz. following
the change in amount of reactant/product by titration, determining the
volume of gas formed, or colorimetric measurement of light intensity at
different times. (The theory of colorimetry is not required.)
Effects of concentration, temperature, pressure, surface area, catalyst and
light on reaction rate.
Simple rate equations determined from experimental results.
Zeroth, first and second order reactions. Rate constants. Half-life of a
Chapter 13
Term examination
Chapter 14
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5.4 The effect of temperature change on
reaction rate
5.5 The interpretation of rates of gaseous
reactions at molecular level
5.6 Energy profile
first order reaction.
Radioactive decay as a typical example of a first order reaction.
Carbon-14 dating in the estimation of the age of an archaeological
specimen.
Cabon-14 dating in the estimation of the age of an archaeological
specimen.
Calculations involving rate equations. (Deriving of integral forms of rate
equations is not required.)
Explanation of the effect of temperature change on reaction rate in terms
of activation energy.
Application of the Arrhenius equation
k = A exp(-Ea/RT)
to determine the activation energy of a reaction. (Derivation of the
Arrhenius equation is not required.)
Distribution of molecular speeds in a gas. (Zartmann experiment and
calculations involving molecular speeds are not required.)
Graphical representation of the Maxwell-Boltzmann distribution and its
variation with temperature. Simple collision theory. (Qualitative
treatment only.)
Energy profile as a representation of the changes in potential energy
Chapter 15
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5.7 Catalysts and their effect on reaction
rates
Homogeneous and heterogeneous
catalysis
Applications of catalysts
during a reaction. Simple stage and multi-stage reactions.
The rate determining step in a multi-stage reaction.
Catalysts can change the rate of a reaction by providing an alternative
pathway for the reaction.
Acid-catalysed esterification as an example of homogeneous catalysis.
Effect of manganese(IV) oxide on the decomposition of hydrogen
peroxide as an example of heterogeneous catalysis.
The use of catalysts in Contact and Haber processes, and the
hydrogenation of unsaturated oils. Catalytic converters in c ar exhaust
systems. An awareness that enzymes are example of biological catalysts.
6. Chemical Equilibria
6.1 Dynamic equilibrium
The equilibrium law
The effect of changes in
concentration, pressure and
Reversible reactions.
Dynamic nature of chemical equilibrium.
Characteristics of chemical equilibrium.
Equilibrium constants expressed in terms of concentration (Kc) and
partial pressure (Kp). Simple calculations of Kc and Kp. (The quantitative
relationship between Kc and Kp is not required.)
Le Chateliers principle. Changes in concentration and pressure result in
the adjustment of the system without changing the value of equilibrium
Chapter 16
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temperature on equilibria
6.2 Acid-base equilibria
Concepts of acid/base
Dissociation of water
pH and its measurement
Strong and weak acids/bases
constant, K; a change in temperature results in the adjustment of the
system to a new equilibrium constant.
Reaction of temperature and the value of K for exothermic and
endothermic reactions illustrated by the equation,
ln K = constant H/RT
(Derivation of the equation is not required.)
Simple calculation on equilibrium composition involving changes in
concentration/pressure.
Bronsted-Lowry theory
Ionic product of water, Kw
The use of indicators and pH meters to measure pH. (The theory and
instrumentation of pH meters are not required.)
Dissociation constants for weak acids (Ka) and weak bases (Kb). Use of
Ka and Kb (pKa and pKb) values to compare the strength of weak acids or
weak bases. Calculations involving pH, Ka and Kb.
(For dissociation involving more than one step, calculations are limited
to one of these steps only.)
Chapter 17
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Buffers
Indicators
Acid-base titrations
6.3 Redox Equilibria
Redox reactions
Electrochemical cells
Electrode potentials
Principle of buffer action. Calculations involving the composition and
pH of buffer solutions.
Simple theory of acid-base indicators and pH range of their colour
changes.
pH titration curves and the choice of indicators.
Redox reactions in terms of electron transfer. Oxidation states. Balancing
redox equations.
E.m.f. measurement of electrochemical cells of metal-metal ion systems.
E.m.f. values to compare the relative tendencies of half cells to release or
gain electrons. Other systems involving non-metal ions.
(e.g. I2(aq) , 2I-(aq) | Pt), ions in different oxidation states(e.g. Fe3+(aq),
Fe2+(aq)|Pt) and metal-metal salt (e.g. PbSO(s), [Pb(s) + S O42-(aq)]|Pt).
Cell equations. IUPAC convention in writing cell diagrams.
The standard hydrogen electrode as a reference. The convention of
standard reduction potentials is adopted. The electrochemical series
(redox potential series). Use of the standard electrode potential (E)
values to compare the strength of oxidizing/reducing agents, and to
Chapter 18
Chapter 19
Chapter 20
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Secondary cell and fuel cell
Corrosion of iron and its prevention
calculate the e.m.f. of cells.
Prediction of the feasibility of redox reactions from electrode potential
values and the limitation of this approach due to kinetic factor.
Lead-acid accumulator and the hydrogen-oxygen fuel cell: structure,
electrochemical processes and uses.
The electrochemical process involved in rusting.
Prevention of corrosion by coating and cathodic protection.
Socioeconomic implication of corrosion and prevention.
7. Phase Equilibrium
7.1 One component systems
7.2 Two component systems
Ideal systems
Non-ideal systems
The pressure temperature diagrams of water and carbon dioxide.
(Phase rule is not required.)
Studies limited to phase diagrams for mixtures of two miscible liquids:
i. Vapour pressure against mole fraction (with temperature
constant), and
ii. Boiling point against mole fraction (with pressure constant).
Rauolts law. The characteristic properties of an ideal system explained
in terms of molecular interactions.
Positive and negative deviations from Rauolts law explained in terms of
molecular interactions. Enthalpy changes on mixing as evidence for non-
Chapter 21
Chapter 22
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Fractional distillation
7.3 Partition of a solute between two
phases
ideal behaviour. Azeotropic mixtures.
Explanation of the principle of fraction distillation using the boiling
point composition curve.
Application of fraction distillation in oil refining.
Partition coefficient of a non-volatile solute distributed between two
immiscible liquids. (Calculations involving dissociation or association of
solute are not required.) Application to solvent extraction.
Paper chromatography as an application of partition. Rfvalue.
Chapter 23
Term test
12. Fundamentals of OrganicChemistry
12.1 Natural sources of organic
compounds
12.2 The unique nature of carbon
12.3 Functional groups and homologous
series
Alkanes, alkenes and aromatic hydrocarbons from crude oil and coal.
Carbohydrates, proteins and fats in living organisms.
Ability of carbon to catenate leading to the existence of a vast number of
carbon compounds.
Studies limited to the following functional groups:
C=C, CC, -X, -OH, -O-, -CHO, C=O, -COOH, -NH2, -NHR, -NR2,
-CN, -COOR, -COX, -CONH2 and (-CO)2O.
Effects of functional groups and the length of carbon chains on physical
properties of compounds in homologous series.
Chapter 24
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12.4 Structures and shapes of
hydrocarbons
Saturated hydrocarbons
Unsaturated hydrocarbons
Aromatic hydrocarbons
12.5 Systematic nomenclature
12.6 Isomerism
Structural isomerism
Geometrical isomerism
The tetrahedral arrangement of the bond electron pairs around a carbon
atom explained in terms of repulsion between electron pairs and in terms
of sp3 hybridized orbitals. (Conformation is not required.)
Formation of the C=C and CC bonds explained in terms of sp2 and sp
hybridized orbitals respectively. andbonds. Shapes associated with sp2
and sp hybridized carbon atoms.
Shape of the benzene molecule.
Delocalization of -electrons in benzene giving rise to a unique class of
compounds which are chemically different from alkenes.
Systematic nomenclature limited to compounds containing carbon chains
of not more than eight carbon atoms.
Isomers containing the same functional group and isomers containing
different functional groups.
Rigidity of C=C bond leading to cis/trans isomers. Geometrical isomers
Chapter 25
Chapter 26
Chapter 27
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Enantiomerism
12.7 Structure determination of organic
compounds
Use of infra-red (IR) spectrum in the
identification of functional groups
limited to acyclic compounds containing one C=C.
Studies limited to structures with one chiral cabon.
(Absolute configuration and resolution of racemic mixtures are not
required.)
Calculation of empirical formula from analytical data (linked with
section 1.6). Molecular formula. Structure deduced from reactions of
functional groups and physical properties.
An awareness that spectroscopic methods such as infrared spectroscopy
and nuclear magnetic resonance (NMR) can provide information about
the structure of a molecule.
IR spectrum and its use in the identification of the following groups:
C-H, O-H, N-H, C=C, CC, C=O and CN. (Instrumentation is not
required.)
Chapter 28
13. Chemistry of Organic Compounds
13.1 Alkanes
Mechanisms other than those mentioned specifically are not required.
Crude oil as a source of alkanes.
Chemical principles and economic importance of fractional distillation
(linked with 7.2) and cracking process (Industrial detail are not
required.)
Combustion of alkanes. Chlorination of alkanes as light-initiated chain
Chapter 30
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13.2 Alkenes
Addition reactions
Ozonolysis
Polymerization of alkenes
13.3 Aromatic hydrocarbons
Substitution
reactions. Mechanism of the chlorination of methane.
Reactions of alkenes with bromine (aqueous and non-aqueous),
hydrogen bromide and sulphuric(VI) acid. Mechanism of the
electrophilic addition of hydrogen bromide to alkenes. Markownikoffs
rule.
Catalytic hydrogenation and its application in the hardening of oils.
Conditions and reaction products. Use in the determination of positions
of the carbon-carbon double bonds in alkenes.
Formation of poly(ethene), poly(propene) and poly(phenylethene).
Mechanism of free radical polymerization of ethane.
Benzene and methylbenzene.
Stability of the benzene ring: comparison of the enthalpy changes of
hydrogenation and combustion for benzene and cyclohexene leading to
the concept of increased stability in a delocalized system. Resistance of
benzene to oxidation and addition reactions.
Nitration, halogenation, sulphonation and alkylation. (limited to mono-
substitution only)
Reaction with potassium manganate(VII)
Chapter 31
Chapter 32
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Oxidation of alkylbenzene
13.4 Halogeno-compounds
Nucleophilic substitution reactions
Elimination reaction
Uses of halogeno-compounds
13.5 Hydroxy compounds
Acidic properties of hydroxy
compounds
Reactions of alcohols
Primary, secondary and tertiary haloalkanes, halobenzene.
Reactions with sodium hydroxide, potassium cyanide and ammonia.
(Experimentation involving potassium cyanide should not be attempted.)
Comparison of rates of hydrolysis of haloalkanes and halobenzene.
Mechanism of SN1 and SN2 as exemplified by substitution with OH
group. (Linked with 5.6)
Reaction of haloalkanes with alcoholic sodium hydroxide to form
alkenes and alkynes.
Halogeno-compounds as solvents in dry-cleaning and as raw materials in
the manufacture of poly(chloroethene) and poly(tetrafluoroethene).
Primary, secondary and tertiary alcohols; phenol.
Comparison of the acidic properties between alcohols and phenol.
Reactions include halide formation, alkoxide formation, oxidation,
dehydration, esterification and triiodomethane formation.
Chapter 33
Chapter 34
Term examination
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Reactions of phenol
Uses of alcohols
Distinction between primary, secondary and tertiary alcohols.
Reactions with sodium and sodium hydroxide.
Esterification.
Alcohols as solvents.
Ethanol in beverages and as a motor fuel blending agent.
Ethane-1.2-diol as an anti-freeze and a raw material in the manufacture
of terylene.
Cycle Date Syllabus Explanatory notes Reference in text Suggested ExperimentsCarbonyl compounds
Nucleophilic addition reactions
Addition- elimination (condensation) reactions
Structures of aldehydes and ketones. Benzaldehyde and phenylethanone as
aromatic carbonyl compounds.
Reactions with hydrogen cyanide and sodium hydrogensulphate( IV).
(Experimentation involving hydrogen cyanide should notbe attempted.)
Mechanism of the addition of hydrogen cyanid e to carbonyl compounds.
Use of the reaction with sodium hyd rogensulphate( IV) in the purification of
carbonyl compounds.
Reactions with hydroxylamine and 2, 4- dinitrophenylhydrazine.Identification of a carbonyl compoun
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Oxidation and reduction
Uses of carbonyl compounds
Oxidation of aldehydes with acidified dichromate( VI), Tollens' reagent and
Fehling's reagent. Resistance of ketones to oxidati on.
Reduction of aldehydes and ketones with sodium tetrahydridoborate (sodium
borohydride) and lithium tetrahydridoaluminate (lithium aluminium hydride).
Formation of triiodomethane as a test for compounds containing
a CH 3 CO group or a CH 3 CH( OH) group.
Methanal in the manufacture of urea- methanal resin.
Propanone as a solvent and a raw material in the manufacture of perspex.
preparing its derivative.
Investigation of the reactions of aldeh
ketones.
13.7
Carboxylic acids and their derivatives
The formation of carboxylic acid
Reactions of carboxylic acids
Acidity of carboxylic acids
Reactions of acyl chlorides and Anhydrides
Structures of carboxylic acids, acyl chlorides, anhydrides, amides and esters.
Hydrolysis of nitriles. Oxidation of alcohols, aldehydes and alkylbenzenes.
Formation of salts, acyl chlorides, anhydrides, amides and esters. Reduction with
lithium tetrahydridoaluminate.
Comparison of the acidity of carboxylic acids with alcohols.
Influence of substituents, viz. alkyl and chloro groups, on
acidity.
Reactions with water, alcohols, ammonia and amines.
Investigation of the reactions of carbo
acids.
Preparation of an ester.
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Reactions of amides
Reactions of esters
Uses of carboxylic acids and their derivatives
Nitrogen compounds
The formation of amines Primary amines from
nitriles and amides.
Base properties of amines
Reaction of amines
.
Hydrolysis, dehydration, Hofmann degradation and reduction with lithium
tetrahydridoaluminate.
Acid and base hydrolyses. Reduction with lithium tetrahydridoaluminate.
Benzoic acid and benzoates as food preservatives. Polyamides and polyesters as
synthetic fibres e. g. nylon 6. 6 and terylene.
Uses of esters as solvents and flavourings.
Primary, secondary and tertiary aliphatic amines, phenylamine and amino acids.
Primary, secondary and tertiary aliphatic amines, and quaternary ammonium
compounds by alkylation. Phenylamine from nitrobenzene.
Salt formation. Comparison of the basic strength of ammonia, primary aliphatic
amines and phenylamine.
Reactions with ethanoyl chloride and benzoyl chloride. Reaction with nitric( III)
acid limited to primary amines only. Coupling reaction of benzenediazonium ion
with naphthalen- 2- ol. (Test to distinguish primary, secondary and tertiary
amines is notrequired.)
Analysis of commercial aspirin tablet
Investigation of the reactions of amin
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Uses of amines and their derivatives
Amino acids
14. Chemistry and Society
14.1 Chemistry and the environment
(a) Air pollution
Some air pollutants
The effects of polluted air on the envi ronment
The ozone layer and chlorofluoro- carbons
Azo- compounds as dyes in dyeing i ndustries. Amine derivatives as drugs.
Amino acids (e. g. aminoethanoic acid and 2- amino- propanoic acid) as
bifunctional compounds having both acidi c and basic characteristics. Zwitterion.
Dipeptides and polypeptides as dimers and polymers of amino acids. (Methods
of formation of polypeptides are notrequired.)
Carbon monoxide, sulphur dioxide, nitrogen oxides, hydrocarbons, ozone and
particulates. Combustion of fossil fuels as the main source of air pollutants.
The harmful effects of pollutants depend on their concentrations and the duration
of exposure to the pollutants. Parts per millio n (ppm) as one way of indicating
concentrations of pollutants. Acid rain and photochemical smog: their formation
and effects on the environment.
Sources and properties of ozone. The desirability of ozone in the stratosphere.
Chlorofluorocarbons as aerosol propellants, solvents for the cleaning of
Project work on air pollution, e.g. aci
smog or ozone depletion.
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(b) Water pollution
The causes of water pollution and its effects
on the environment
Water quality
(c) Solid waste
(d) Pollution control in Hong Kong
electronic components and metals, refrigerants, and blowing agents in foam
plastic manufacturing. Causes for the accumulation of chlorofluorocarbons in the
stratosphere. The free radical chain reactions involved with chloro-
fluorocarbons leading to the depletion of the ozone layer. Control of the ozone
depletion problem. Possible alternatives for chlorofluorocarbons.
The adverse effects on water quality due to liv estock waste, oil spillages,
residues of pesticide, detergents in sewage, and industrial effluents.
An awareness that oxygen dissolved in water is necessary for aquatic life.
Dissolved oxygen (DO) as an indicator of oxygen content in water, expressed as
percentage saturation or mg dm 3 . Biochemical oxygen demand (BOD) as an
indicator of the extent of water pollution.
Plastics, paper and metals. Disposal of solid waste b y landfilling and ncineration.
Pollution problems associated with the disposal of plastics. Development of
degradable plastics and recycling of plastics as possible solutions to pollution
problems.
Measures to improve air quality: use of unleaded petrol and installation of
catalytic converters in car exhaust systems, limitation of sulphur content in fuels,
Determination of dissolved oxygen in
samples.
Visit to
a. the Environment Resource Centr
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14.2 Chemistry and food
(a) Principal components of food Proteins
Carbohydrates
Fats and oils
desulphurization of flue gas, installation of electrostatic precipitators and
installation of low nitrogen oxide burners in power plants.
Measures to improve water quality: screening, sedimentation and digestion of
pollutants by micro- organisms in the treatment of sewage; ph ysical and
chemical methods, and incineration in the treatment of chemical waste from
industry and laboratories. (Technical details of the above treatment processes are
notrequired.)
Measures to reduce solid waste:
reuse/ recycling of paper, plastics and metals to minimize waste and save
resources.
Proteins as macromolecules made up of amino acids via peptide linkages.
Hydrolysis of proteins. Separation of amino acids by paper chromatography.
(Linked with Sections 7.3 and 13.8)
Classification into monosaccharide, disaccharide and polysaccharide. Open chain
and ring structures of glucose and fructose. Glycosidic linkage in carbohydrates.
Hydrolysis of sucrose and starch. Fehling's test to distinguish between reducing
and non- reducing sugars.
Fats and oils as esters of propane- 1, 2,3- triol and fatty acids. Hydrolysis of fats
and oils (Link with Section 13 .7). Use of iodine value tocomparethe degree of
b. the Chemical Waste Treatment C
c. a sewage treatment plant.
Separation of amino acids by paper
chromatography.
Investigation of the hydrolysis of suc
testing for reducing sugars.
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(b) Food preservation
The need to preserve food
Principles and techniques of food preservation
(c) Food additives
unsaturation. Hardening of vegetable oils. (Link with Section 13.2) Hydrolytic
and oxidative rancidity.
Prevention of food spoilage due to microbial activities and chemical changes.
Principles of food preservation:
killing of micro- organisms, inhibition of microbial growth, retardation of
chemical changes by removing moisture, altering temperature, changing pH, and
the use of osmotic process and chemical additives. Common techniques include
heat treatment, irradiation, drying, dehydration, refrigeration, canning, sugaring,
salting and chemical preservation such as meat- curing, picklin g and the use
of food additives.
Food additives to serve as preservatives (e. g. nit rates( III), nitrates( V), sulphur
dioxide, sulphates( IV), benzoic acid and b enzoates) and antioxidants (e. g.
BHA (butylated hydroxyanisole) and BHT (butylated hydroxytoluene)), to
enhance the flavour (e. g. MSG (monosodium glutamate), saccharin), texture
(e. g. emulsifying agents), appearance (e. g. colouring agent s) or nutritional
value (e. g. vitamins) of food.
Principle of BHA/ BHT as antioxidant to retard atmospheric oxidation of oils
and fats.
Investigation of the effects of air and
preservatives on apple browning.
Library search on different functions
common food additives.
Analysis of sulphur dioxide content in
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The possible menace of
food additives
8. Periodic Properties of the Elements in
the Periodic Table
8.1 Periodic variation in physical properties of
the elements H to Ar
8.2 Periodic relationship among the oxides,
chlorides and simple hydrides of the elements
Li to Cl
9. The s- Block Elements
9.1 Characteristic properties of the s- block
elements
The side effects of MSG, the toxicity of nitrates( III) and sulphur dioxide, and
the potent carcinogenic nature of nitrates( III) and saccharin.
An awareness that the use of food additives is monitored by research findings
and by legislation.
Variations in first ionization enthalpies (linked with Section 2. 2), atomic radii,
electronegativities and melting points. Interpretation of t hese variations in terms
of structure and bonding.
Bonding and stoichiometric composition of the hydrides, oxides and chlorides of
these elements, and their behaviour with water. (Hydrides of boron are not
required.)
Metallic character and low electronegativity. Formation of basic oxides and
hydroxides. Predominantly ionic bonding with fixed oxidation state in their
Debate on the use of food additives.
Investigation of the properties of the o
chlorides of the period 3 elements.
Flame tests for Li+, Na+, K+, Ca2+, Sr2
ions.
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9.2 Variation in properties of the s- block
elements and their compounds
9.3 Uses of the compounds of th e s- block
elements
10. The p- Block Elements
10.1 The halogens
Characteristic properties of the halogens
compounds. Characteristic flame colours of salts. Weak tendency to form
complexes.
Variations in atomic radii, ionization enthalpies, melting points and hydration
enthalpies. Interpretation of these variations in terms of structure and bonding.
Reactions of the elements with hydrogen, oxygen, chlorine and water.
Reactions of the oxides, hydrides and chlorides with water, acids and alkalis.
Relative thermal stability of the carbonates and hydroxides.
Relative solubility of the sulphates( VI) and hydroxides.
Sodium carbonate in the manufacture of glass.
Sodium hydrogencarbonate in baking powder. Sodium hydroxide in making
soap. Magnesium hydroxide as an antacid. Slaked lime in neutralization of acids
in industrial effluents. Strontium compounds in fireworks.
High electronegativity and electron affinity. Ionic and covalent bonding in
oxidation state 1.
Investigation of the effect of heat on c
of Group II elements.
Investigation of the solubility of sulph
and hydroxides of Group II elements
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Variation in properties of the halogens and
their compounds
Uses of halogens and halogen containing
compounds
10.2 Nitrogen and its compounds
10.3 Sulphur and its compounds
Variations in melting and boiling points, electronegativities and electron
affinities. Interpretation of these variations in terms of structure and bonding.
Relative oxidizing power of halogens: comparative study of reactions (Cl 2 ,
Br2 and I 2 ) with sodium, iron( II) ion and ph osphorus.
Disproportionation of the halogens in alkalis.
Comparative study of the reactions of halide ions with halogens, sulphuric( VI)
acid, phosphoric( V) acid and silver ions. Acidic properties of hydrogen halides
and the anomalous behaviour of hydrogen fluoride.
Fluoride in fluoridation of water. Chlorine in the manufacture of
poly( chloroethene), bleach and disinfectant. Silver bromide in photographic
films.
Unreactive nature of nitrogen. Direct combination of nitrogen and oxygen
leading to formation of nitrogen oxides. Manufacture of ammonia by Haber
process and its underlying physicochemical principles. Ammonia as a reducing
agent and a base. Catalytic oxidation of ammonia in the manufacture of
nitric( V) acid. Nitric( V) acid as an oxidizing agent, limited to the stud y of the
reactions with copper, iron( II) ion and sulphur only.
Action of heat on nitrates( V). Brown ring test for nitrate( V) ions.
Burning of sulphur. Oxidizing and reducing properties of sulphu r dioxide as
Investigation of the reactions of
a. halogens with alkalis,
b. halides ions in solution, and
c. solid halides with sulphuric(VI) a
Investigation of the action of heat on
nitrates(V). Brown ring test for nitrate
Investigation of the redox properties o
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11. The d- Block Elements
11.1 General features of the d- block elements
from Sc to Zn
11.2 Characteristic properties of the d- block
elements and their compounds:
(a) Variable oxidation states
(b) Complex formation
exemplified by the reactions with manganate( VII) ion, dichromate( VI) ion,
bromine and magnesium metal. Manufacture of sulphuric( VI) acid by contact
process and itsunderlying physicochemical principles. Sulphuric( VI) acid as an
oxidizing agent and a dehydrating agent.
Test for sulphate( VI) ions. Uses of sulphuric( VI) acid in the manufacture of
fertilizers, detergents, paints, pigments and dyestuffs. Investigation of the redox
properties of sulphur dioxide.
Electronic configurations (linked with Section 2. 4). d- Block elements as metals.
Comparison of ionization enthalpies, electronegativities, melting points,
hardness, densities and reactions with water between d- block and s- block
metals.
Interpretation of the characteristic properties, viz. variable oxidation states,
complex formation, coloured ions, and catalytic properties in terms of electronic
structures, successive ionization enthalpies, atomic and ionic radii.
Studies limited to common oxidation states o f vanadium (+ 2, +3, +4, +5) and
manganese (+ 2, +4, +7). Interconversions of oxidation states of each element.
Studies limited to complexes of Fe( II), Fe( III), Co( II) and Cu( II) with the
following ligands: H 2 O, NH 3 ,Cl and CN .
dioxide.
Test for sulphate(VI) ions using acidi
barium chloride solution.
Investigation of the redox reactions o
or manganese compounds.
Investigation of the relative stability o
copper(II) complexes.
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(c) Coloured ions
(d) Catalytic properties of transition metals
and their compounds
Nomenclature of these complexes. Displacement of ligands and relative stability
of complex ions. (Experimentation involving cyanide ions should notbe
attempted.) (Calculations involving stability constants are notrequired). Stereo-
structures of 4- and 6- coordinated complexes.(Optical isomerism of complexes
is notrequired.)
Studies limited to the hydrated io ns of Fe( II), Fe( III), Co( II) and Cu( II).
Exemplified by the use of Fe in Haber process, Fe 2+or Fe 3+in the reaction
between peroxodisulphate( VI) and iodide ions, and MnO 2 in the decomposition
of hydrogen peroxide (linked with Section5.7).
Investigation of the catalytic action o
ions on the reaction between
peroxodisulphate(VI) and iodide ions
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