emb's 2006 al chem syllabus

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CHEMISTRY

ADVANCED LEVEL

This syllabus builds on the foundation of the HKCE Chemistry syllabus, a knowledge of which is assumed. The syllabus has been designed for a two-year chemistry course at advanced level. It will adequately prepare candidates for further studies in chemistry and related disciplines. This syllabus is not a teaching syllabus and the order of the topics listed is not intended to suggest a teaching order. Explanatory notes have been included in the syllabus as appropriate to indicate the scope and depth of treatment. However, the length of the notes on any particular section should not be regarded as an indication of the time to be spend on that section. Experiments suggested in the syllabus are more than enough. Some of these experiments are alternative ones which are similar in nature. Teachers are advised to make discretionary choices as appropriate. Other activities such as visits, projects, debates, etc. are suggested to stimulate candidates’ interest in chemistry and to promote skills in reading, data collection, planning, presentation and problem solving.

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AIMS The aims of the syllabus are to enable the candidates: 1. to develop an interest in the study of chemistry; 2. to develop an understanding of the facts and patterns in the empirical world; 3. to develop an understanding of the concepts and principles of chemistry; 4. to develop experimental skills and an awareness of safety problems; 5. to develop the ability to observe, to analyze and to interpret data objectively; 6. to develop the ability to communicate using the language of chemistry; 7. to develop the ability to solve problems and to make rational decisions; 8. to develop an appreciation of chemistry and its application in daily life; and 9. to develop an awareness of the social, economic, environmental and technological implications of chemistry. ASSESSMENT OBJECTIVES The objectives of the examination are to test the following abilities: 1. to recall and understand chemical facts, patterns, principles, methods, terminology and conventions; 2. to understand the use of apparatus and materials in performing and planning experiments; 3. to handle materials, manipulative apparatus, carry out experiments safely and make accurate observations; 4. to analyze and interpret data from various sources, and draw relevant conclusions;

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5. to manipulate and translate chemical data and to perform calculations; 6. to apply chemical knowledge to explain observations and to solve problems which may involve unfamiliar situations; 7. to select and organize scientific information from appropriate sources and to communicate this information in an appropriate and

logical manner; 8. to understand the social, economic, environmental and technological implications of the applications of chemistry; and 9. to make decisions based on the examination of evidence and arguments. THE EXAMINATION The examination will consist of two written papers and one practical paper. WRITTEN EXAMINATION (80% of the Subject Mark) The duration of the two written papers is three hours each. The structure of the papers is as follows: Paper 1 Section A (Short questions) 60% Section B (Questions on practical chemistry) 20% Section C (Essay questions) 20% Paper 2 Section A (mainly based on topics 1-11 of the syllabus) 60% Section B (mainly based on topics 12-14 of the syllabus) 40%

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In Paper 1 Sections A and B, all questions are compulsory; in Section C, a choice of 1 out of 2 questions will be allowed. In Paper 2, a choice of 3 out of 4 questions is allowed in Section A, and a choice of 2 out of 3 questions is allowed in Section B. Some useful chemical information, including characteristic infra-red absorption wavenumber ranges, will be provided for both papers 1 and 2. PRACTICAL EXAMINATION (20% of the Subject Mark) Candidates should be familiar with experiments illustrating the chemical systems and chemical principles in topics 1-14 of the syllabus and the use of common laboratory apparatus. Besides, candidates should also be familiar with (a) techniques and principles involved in quantitative volumetric analysis concerning acid-base reactions, oxidation-reduction

reactions and equilibria studies in aqueous solution, (b) observations and interpretations using qualitative analysis on a test tube scale. (1) For school candidates, their practical abilities will be assessed internally by the teachers. The following areas of ability of the candidates will be assessed: Ability area A: (a) to use apparatus and to demonstrate appropriate manipulative skills in carrying out experiments; (b) to make accurate observations and measurements. Ability area B: (a) to record and to present data in an appropriate form; (b) to interpret experimental results and to draw appropriate conclusions; (c) to plan experiments.

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Ability area C : the candidate’s attitude towards practical chemistry. Scheme of Assessment: The minimum number of experiments performed in the first and second year are 18 and 10 respectively. Over the two years, not

less than 12 assessment marks are required. In each year, the number of assessment marks for each of these areas is: at least TWO assessment marks for ability area A, at least TWO assessment marks for ability area B, and ONE assessment mark for ability area C. The regulations, guidelines and methods of assessment can be found in the ‘Handbook on the Teacher Assessment Scheme for

Practical Chemistry (TAS)’ issued by the Hong Kong Examinations and Assessment Authority to participating schools. (2) For private candidates, they may opt to sit the practical examination or to use their previous TAS marks to substitute the practical

examination. The practical examination (3 hours) will include (a) an experiment on quantitative chemistry, and (b) an experiment involving an observational and deductive exercise. Candidates are assigned to take the examination in groups, and the questions for each group are not necessarily the same.

Candidates will be provided with the appropriate apparatus, and instructions will be given to candidates indicating what has to be done and which critical observation must be made.

Candidates will be expected to show their ability to work accurately within the limits of the apparatus and chemical reactions involved. Candidates will also be expected to perform chemical calculations and/or to draw conclusions from the observations.

(Note : Experiments may be set for which candidates are not expected to have had previous experience. In such cases, full instructions will be given.) No textbooks, notes etc. may be used in the examination.

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All candidates (i.e. both school and private candidates) may be required to submit their laoratory books for inspection. NOTES 1. In general, SI units will be used. 2. In naming compounds, the useful references are “Guidelines for Systematic Chemical Nomenclature (2000)” by the Hong Kong

Examinations Authority, and “Chemical Nomenclature, Symbols and Terminology for Use in School Science” (1985) by the Association for Science Education (U.K.). The overriding rule for the naming of compounds is clarity and lack of ambiguity rather than adherence to strict rules.

3. “An English-Chinese Glossary of Terms Commonly Used in the Teaching of Chemistry in Secondary Schools” (1999) prepared by the Curriculum Development Council, issued by the Education Department is a useful reference for the Chinese terms.

1. Atoms, Molecules and Stoichiometry ��Topics Explanatory Notes Suggested Experiments/Activities 1.1 The atomic structure

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.

��Topics Explanatory Notes Suggested Experiments/Activities

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1.2 Radioactivity Nature of α and β particles, and of γ radiation. Equations for nuclear reactions. Uses of isotopes in leak detection, radiotherapy, nuclear power and as tracers. (Underlying principles and instrumentation are not required.)

1.3 Relative isotopic, atomic and molecular masses

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.)


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1.4 The mole concept 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 relative molecular mass determination. (Non-ideal behaviour of real gases and kinetic theory are not required.) Partial pressure of gas and its relationship to mole fraction.

Determination of the relative molecular mass of a volatile liquid.

1.5 The Faraday and the mole

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.

A quantitative study of electrolysis.

1.6 Empirical and molecular formulae

Derivation of empirical formula using combustion data or composition by mass. Molecular formula derived from empirical formula and relative molecular mass.


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1.7 Chemical equations and stoichiometry

The stoichiometric relationship between reactants and products in a reaction. Calculations involving (i) reacting masses, (ii) volumes of gases, and (iii) concentrations and volumes of solutions.

Titrations involving (a) acid-base reactions, and (b) redox reactions.

2. The Electronic Structure of Atoms and the Periodic Table 2.1 Atomic emission spectra and

electronic structure of atoms Characteristics of the emission spectrum of atomic hydrogen. Interpretation of the spectrum using the relationship, E = hν leading to the idea of discrete energy levels.


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2.2 Electronic structure, ionization enthalpies, electron shells and subshells

Plots of the following graphs to introduce shells and subshells: (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.)


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2.3 Atomic orbitals

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.)


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2.4 Electronic configurations of atoms

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 (Hund's rule), and

(iii) electrons occupying the same orbital must have opposite spins (Pauli's exclusion principle).


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Electronic configurations in relation to the Periodic Table

Electronic configurations of isolated atoms from H to Kr. Electronic configurations of atoms represented by (i) notations using 1s, 2s,

2p, ... etc., e.g. Fe (ground state).

1s22s22p63s23p63d64s2

(ii) 'electrons-in-boxes' diagram, e.g. Fe ground state).

3d 4s [Ar] ↑↓ ↑ ↑ ↑ ↑ ↑↓ 2.5 The Periodic Table and the

atomic properties of the elements

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.

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3. Energetics ��Topics Explanatory Notes Suggested Experiments/Activities

3.1 Energy changes in chemical

reactions Conservation of energy. Endothermic and exothermic reactions and their relationship to the breaking and forming of bonds.

3.2 Standard enthalpy changes 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 calorimetry is not required).

Determination of enthalpy changes of neutralization, solution or combustion.


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3.3 Hess's law Use of Hess's law to determine enthalpy changes which are not easily obtainable by experiment. Enthalpy level diagrams. Calculations involving enthalpy changes of reactions.

Determination of the enthalpy change of formation of CaCO3, MgCO3 or MgO, or the enthalpy change of hydration of MgSO4.

4. Bonding and Structure 4.1 The nature of forces holding

atoms together

Electrostatic interactions between electrons and nuclei leading to different types of bonding.

4.2 Ionic 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.


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Energetics of formation of ionic compounds

Born-Haber cycles for the formation of ionic compounds in terms of enthalpy changes of atomization and ionization, electron affinities and lattice enthalpies. (Electron 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.)


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Stoichiometry of ionic compounds

Consideration in terms of electronic configurations and enthalpy changes of formation.

Ionic crystals Extended three-dimensional structures of ionic compounds limited to sodium chloride and caesium chloride. Unit cells and coordination numbers. (Calculations involving ionic radii in a unit cell are not required.)

Display/Build lattice models of NaCl and CsCl.

Ionic radii

Comparison of sizes of ions with their parent atoms. Comparison of sizes of isoelectronic particles.

4.3 Covalent bonding Formation of covalent bonding − sharing of electron pairs. The simple idea of the overlapping of atomic orbitals.


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Dative covalent bonding

'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.

Bond enthalpies, bond lengths and covalent radii

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.


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The shapes of covalent molecules and polyatomic ions

The shapes of simple molecules and polyatomic ions explained in terms of the repulsion between electron pairs (as illustrated by BF3, CH4, NH3, H2O, PCl5, SF6, NH4

+ and NH2

–). The directional nature of covalent bonds. Bond angles.

Display/Build models of simple molecules.

Multiple bonds Comparison of bond lengths and bond enthalpies leading to the idea of multiple bonds, illustrated by ethene and ethyne. Shapes of carbon dioxide and sulphur dioxide molecules explained in terms of repulsion between electron pairs.

Covalent crystals

Exemplified by diamond, graphite and quartz.

Display/Build models of diamond, graphite and quartz.


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4.4 Bonding intermediate between ionic and covalent

Incomplete electron transfer in ionic compounds

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.)


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Polarity of covalent bond 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.) Unequal sharing of bonded electron pair(s) explained in terms of the electronegativity difference between bonded atoms. Electronegativity (Pauling's scale) introduced as an arbitrary measure of an atom's tendency in a molecule to attract electrons. (The formal definition of electronegativity and its experimental determination are not required.)

Investigation of the effect of a non-uniform electrostatic field on a jet of liquid.


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4.5 Metallic bonding

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.

Metallic crystals Close-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.)

Display/Build models of metal crystals.


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4.6 Intermolecular forces van der Waals' forces

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 covalent bonds and van der Waals' forces.

Molecular crystals

Exemplified by iodine and carbon dioxide.


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Hydrogen bonding

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 (deoxyribo-nucleic acid).

Determination of the strength of the hydrogen bond formed between trichloromethane and ethyl ethanoate.

4.7 The relationship between structures and properties of materials

Differences in physical properties (viz. melting and boiling points, electrical conductivity, hardness and solubility) between ionic compounds, covalent substances and metals.

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5. Chemical Kinetics

Topics Explanatory Notes Suggested Experiments/Activities 5.1 Rates of chemical 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.)

5.2 Factors influencing reaction rate

Effects of concentration, temperature, pressure, surface area, catalyst and light on reaction rate.

Investigation of the following factors on the reaction rate: (a) concentration: HCl/Mg (b) temperature: HCl/Na2S2O3

(c) particle size: acid/marble chips and acid/powdered CaCO3


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5.3 Rate equations and order of

reactions

Simple rate equations determined from experimental results. Zeroth, first and second order reactions. Rate constants. Half-life of a 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. Calculations involving rate equations. (Derivation of integral forms of rate equations is not required.)

(d) catalyst: MnO2 on decomposition of H2O2

(e) light: Br2/hexane Determination of the order of a reaction.


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Explanation of the effect of

temperature change on reaction rate in terms of activation energy. Application of the Arrhenius equation to determine the activation energy of a reaction. (Derivation of the Arrhenius equation is not required.)

Determination of the activation energy of a reaction.

5.5 The interpretation of rates of gaseous reactions at molecular level

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.)

)exp(RTE

Ak a−=

5.4 The effect of temperature change on reaction rate


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5.6 Energy profile

Energy profile as a representation of the changes in potential energy during a reaction. Single stage and multi-stage reactions. The rate determining step in a multi-stage reaction.

5.7 Catalysts and their effect on reaction rates

Catalysts can change the rate of a reaction by providing an alternative pathway for the reaction.

Investigation of the effect of Mn2+ on the reaction between MnO4

– and C2O42– in

acidic medium.

Homogeneous and heterogeneous catalysis

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.


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Applications of catalysts

The use of catalysts in contact and Haber processes, and the hydrogenation of unsaturated oils. Catalytic converters in car exhaust systems. An awareness that enzymes are examples of biological catalysts.

6. Chemical Equilibria

6.1 Dynamic equilibrium

Reversible reactions. Dynamic nature of chemical equilibrium. Characteristics of chemical equilibrium.

Investigation of some reversible reactions: (a) Adding OH– and H+ alternately to

Br2(aq) (b) Adding OH– and H+ alternately to

K2Cr2O2(aq) (c) Hydrolysis of BiCl3


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The equilibrium law Equilibrium constants expressed in terms of concentrations (Kc) and partial pressures (Kp). Simple calculations of Kc and Kp. (The quantitative relationship between Kc and Kp is not required.)

Determination of Kc for a reaction: CH3CO2H(aq) + C2H5OH(aq)

� CH3CO2C2H5(�

)+ H2O(

�

) or Fe3+(aq) + NCS–(aq)

� FeNCS2+(aq)


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The effect of changes in concentration, pressure and temperature on equilibria

Le Chatelier's principle. Changes in concentration and pressure result in the adjustment of the system without changing the value of equilibrium constant, K; a change in temperature results in the adjustment of the system to a new equilibrium constant. Relation of temperature and the value of K for exothermic and endothermic reactions illustrated by the equation, (Derivation of the equation is not required.) Simple calculations on equilibrium composition involving changes in concentration/pressure.

Investigation of the effect of temperature and pressure change on the following reaction:

N2O4(g) � 2NO2(g)

RTH

K∆−= Constantln


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6.2 Acid-base equilibria Concept of acid/base

Bronsted-Lowry theory.

Dissociation of water

Ionic product of water, Kw.

pH and its measurement

The use of indicators and pH meters to measure pH. (The theory and instrumentation of pH meters are not required.)

Strong and weak acids/bases

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.)

Comparison of the strength of weak acids/weak bases by pH or by electrical conductivity measurement. Determination of Ka of a weak acid/Kb of a weak base by pH measurement.


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Buffers

Principle of buffer action. Calculations involving the composition and pH of buffer solutions.

Comparison of the effects of acid/alkali on the pH of buffered and unbuffered solutions.

Indicators

Simple theory of acid-base indicators and pH range of their colour changes.

Determination of the pH ranges of some acid-base indicators.

Acid-base titrations

pH titration curves and the choice of indicators.

Acid-base titrations using method of double indicator.

6.3 Redox Equilibria Redox reactions

Redox reactions in terms of electron transfer. Oxidation states. Balancing redox equations.


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Electrochemical cells 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. PbSO4(s), [Pb(s) + SO4

2– (aq)]

�

Pt). Cell equations. IUPAC conventions in writing cell diagrams.

Investigation of the e.m.f. of some electrochemical cells.


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Electrode potentials 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 (Eo) values to compare the strength of oxidizing/reducing agents, and to 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.

Testing predictions about the feasibility of redox reactions.

Secondary cell and fuel cell

Lead-acid accumulator and the hydrogen-oxygen fuel cell: structure, electrochemical processes and uses.


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Corrosion of iron and its prevention

The electrochemical process involved in rusting. Prevention of corrosion by coating and cathodic protection. Socioeconomic implications of corrosion and prevention.

7. Phase Equilibrium 7.1 One component systems The pressure − temperature

diagrams of water and carbon dioxide. (Phase rule is not required.)

7.2 Two component systems

Studies limited to phase diagrams for mixtures of two miscible liquids: (i) vapour pressure against mole fraction (with temperature constant), and

Investigation of the variation of the boiling point with composition for different mixtures of two miscible liquids.


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Ideal systems

(ii) boiling point against mole fraction (with pressure constant). Raoult's law. The characteristic properties of an ideal system explained in terms of molecular interactions.

Non-ideal systems

Positive and negative deviations from Raoult's law explained in terms of molecular interactions. Enthalpy changes on mixing as evidence for non-ideal behaviour. Azeotropic mixtures.

Fractional distillation

Explanation of the principle of fractional distillation using the boiling point − composition curve. Application of fractional distillation in oil refining.


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7.3 Partition of a solute between two phases

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. Rf value.

Determination of the partition coefficient of ethanoic acid between water and 2-methyl-propan-1-ol.

8. Periodic Properties of the Elements in the Periodic Table 8.1 Periodic variation in physical

properties of the elements H to Ar

Variations in first ionization enthalpies (linked with Section 2.2), atomic radii, electronegativities and melting points. Interpretation of these variations in terms of structure and bonding.


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8.2 Periodic relationship among the oxides, chlorides and simple hydrides of the elements Li to Cl

Bonding and stoichiometric composition of the hydrides, oxides and chlorides of these elements, and their behaviour with water. (Hydrides of boron are not required.)

Investigation of the properties of the oxides and chlorides of the Period 3 elements.

9. The s-Block Elements 9.1 Characteristic properties of the

s-block elements

Metallic character and low electronegativity. Formation of basic oxides and hydroxides. Predominantly ionic bonding with fixed oxidation state in their compounds. Characteristic flame colours of salts. Weak tendency to form complexes.

Flame tests for Li+, Na+, K+, Ca2+, Sr2+ and Ba2+ ions.


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9.2 Variation in properties of the s-block elements and their compounds

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.

Investigation of the effect of heat on carbonates of Group II elements. Investigation of the solubility of sulphates(VI) and hydroxides of Group I elements.


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9.3 Uses of the compounds of the s-block elements

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.

10. The p-Block Elements 10.1 The halogens

Characteristic properties of the halogens

High electronegativity and electron affinity. Ionic and covalent bonding in oxidation state –1.


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Variation in properties of the halogens and their 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 (Cl2, Br2 and I2) with sodium, iron(II) ion and phosphorus. 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.

Investigation of the reactions of (a) halogens with alkalis, (b) halide ions in solution, and (c) solid halides with sulphuric(VI) and

phosphoric(V) acids.


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Uses of halogens and halogen containing compounds

Fluoride in fluoridation of water. Chlorine in the manufacture of poly(chloroethene), bleach and disinfectant. Silver bromide in photographic films.

10.2 Nitrogen and its compounds

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 study of the reactions with copper, iron(II) ion and sulphur only.


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10.3 Sulphur and its compounds

Action of heat on nitrates(V). Brown ring test for nitrate(V) ions. Burning of sulphur. Oxidizing and reducing properties of sulphur dioxide as exemplified by the reactions with manganate(VII) ion, dichromate(VI) ion, bromine and magnesium metal. Manufacture of sulphuric(VI) acid by contact process and its underlying 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 action of heat on nitrates(V). Brown ring test for nitrate(V) ions. Investigation of the redox properties of sulphur dioxide. Test for sulphate(VI) ions using acidified barium chloride solution.

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11. The d-Block Elements ��Topics Explanatory Notes Suggested Experiments/Activities 11.1 General features of the d-block

elements from Sc to Zn 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.

11.2 Characteristic properties of the d-block elements and their

compounds:

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.


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(a) Variable oxidation states

Studies limited to common oxidation states of vanadium (+2, +3, +4, +5) and manganese (+2, +4, +7). Interconversions of oxidation states of each element.

Investigation of the redox reactions of vanadium or manganese compounds.

(b) Complex formation Studies limited to complexes of Fe(II), Fe(III), Co(II) and Cu(II) with the following ligands: H2O, NH3, Cl– and CN–. Nomenclature of these complexes. Displacement of ligands and relative stability of complex ions. (Experimentation involving cyanide ions should not be attempted.) (Calculations involving stability constants are not required). Stereo-structures of 4- and 6-coordinated complexes. (Optical isomerism of complexes is not required.)

Investigation of the relative stability of some copper(II) complexes.


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(c) Coloured ions

Studies limited to the hydrated ions of Fe(II), Fe(III), Co(II) and Cu(II).

(d) Catalytic properties of transition metals and their compounds

Exemplified by the use of Fe in Haber process, Fe2+ or Fe3+ in the reaction between peroxodisulphate(VI) and iodide ions, and MnO2 in the decomposition of hydrogen peroxide (linked with Section 5.7).

Investigation of the catalytic action of d-block ions on the reaction between peroxodisulphate(VI) and iodide ions.

12. Fundamentals of Organic Chemistry 12.1 Natural sources of organic

compounds Alkanes, alkenes and aromatic hydrocarbons from crude oil and coal. Carbohydrates, proteins and fats in living organisms.


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12.2 The unique nature of carbon Ability of carbon to catenate leading to the existence of a vast number of carbon compounds.

12.3 Functional groups and

homologous series

Studies limited to the following functional groups:

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C=C

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, −C≡C− , −X , −OH , −O− , −CHO ,

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C=O , −CO2H, −NH2 , −NHR , −NR2 , −CN , −CO2R , −COX , −CONH2 and (−CO)2O. Effects of functional groups and the length of carbon chains on physical properties of compounds in homologous series.


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12.4 Structures and shapes of hydrocarbons

Saturated hydrocarbons

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.)

Display/Build models of simple alkanes.

Unsaturated hydrocarbons

Formation of the C=C and C≡C bonds explained in terms of sp2 and sp hybridized orbitals respectively. σ and π bonds. Shapes associated with sp2 and sp hybridized carbon atoms.

Display/Build models of simple alkenes.

Aromatic hydrocarbons Shape of the benzene molecule. Delocalization of π-electrons in benzene giving rise to a unique class of compounds which are chemically different from alkenes.

Display/Build models of benzene.


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12.5 Systematic nomenclature Systematic nomenclature limited to compounds containing carbon chains of not more than eight carbon atoms.

12.6 Isomerism Structural isomerism

Isomers containing the same functional group and isomers containing different functional groups.

Geometrical isomerism Rigidity of C=C bond leading to cis/trans isomers. Geometrical isomers limited to acyclic compounds containing one C=C.

Display/Build models of but-2-enes. Investigation of some properties of cis- and trans- butenedioic acids.

Enantiomerism

Studies limited to structures with one chiral carbon. (Absolute configuration and resolution of racemic mixtures are not required.)

Illustration of optical activity using crossed polaroids or a polarimeter.


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12.7 Structure determination of organic compounds

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 infra-red spectroscopy and nuclear magnetic resonance (NMR) can provide information about the structure of a molecule.

Use of infra-red (IR) spectrum in the identification of functional groups

IR spectrum and its use in the identification of the following groups: C−H, O−H, N−H, C=C, C≡C, C=O and C≡N. (Instrumentation is not required.)

Inspection of IR spectra of organic compounds.

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13. Chemistry of the Organic Compounds NOTE: In this section, mechanisms other than those mentioned specifically are not required. ��Topics Explanatory Notes Suggested Experiments/Activities 13.1 Alkanes Crude oil as a source of

alkanes. Chemical principles and economic importance of fractional distillation (linked with Section 7.2) and cracking process. (Industrial details are not required.) Combustion of alkanes. Chlorination of alkanes as light-initiated chain reactions. Mechanism of the chlorination of methane.


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13.2 Alkenes Addition reactions

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. Markownikoff's rule. Catalytic hydrogenation and its application in the hardening of oils.

Ozonolysis Conditions and reaction products. Use in the determination of positions of the carbon-carbon double bonds in alkenes.

Polymerization of alkenes Formation of poly(ethene), poly(propene) and poly(phenylethene). Mechanism of free radical polymerization of ethene.


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13.3 Aromatic hydrocarbons 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.

Investigation of the chemical properties of cyclohexane, cyclohexene and methylbenzene.

Substitution reactions of benzene

Nitration, halogenation, sulphonation and alkylation. (Limited to mono-substitution only.)

Oxidation of alkylbenzene

Reaction with potassium manganate(VII).

13.4 Halogeno-compounds

Primary, secondary and tertiary haloalkanes; halobenzene.


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Nucleophilic substitution reactions

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.

Comparison of the rates of hydrolysis of (a) chloro-, bromo- and iodo- alkanes. (b) primary, secondary and tertiary

haloalkanes. (c) haloalkane and halobenzene.

Mechanism of SN1 and SN2 as exemplified by substitution with −OH group. (Linked with Section 5.6)

Investigation of the kinetics of the hydrolysis of 2-chloro-2-methylpropane.

Elimination reaction Reaction of haloalkanes with alcoholic sodium hydroxide to form alkenes and alkynes.

Uses of halogeno-compounds Halogeno-compounds as solvents in dry-cleaning and as raw materials in the manufacture of poly(chloroethene) and poly(tetrafluoroethene).


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13.5 Hydroxy compounds Primary, secondary and tertiary alcohols; phenol.

Acidic properties of hydroxy compounds

Comparison of the acidic properties between alcohols and phenol.

Reactions of alcohols

Reactions include halide formation, alkoxide formation, oxidation, dehydration, esterification and triiodomethane formation. Distinction between primary, secondary and tertiary alcohols.

Preparation of 1-bromobutane from butan-1-ol, or preparation of cyclohexene from cyclohexanol. Investigation of the reactions of some alcohols, and the Lucas' test for primary, secondary and tertiary alcohols.

Reactions of phenol

Reactions with sodium and sodium hydroxide. Esterification.

Investigation of the reactions of phenol.

Uses of alcohols 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.


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13.6 Carbonyl compounds Structures of aldehydes and ketones. Benzaldehyde and phenylethanone as aromatic carbonyl compounds.

Nucleophilic addition reactions Reactions with hydrogen cyanide and sodium hydrogensulphate(IV). (Experimentation involving hydrogen cyanide should not be attempted.) Mechanism of the addition of hydrogen cyanide to carbonyl compounds. Use of the reaction with sodium hydrogensulphate(IV) in the purification of carbonyl compounds.

Addition-elimination (condensation) reactions

Reactions with hydroxylamine and 2,4-dinitrophenylhydrazine.

Identification of a carbonyl compound by preparing its derivative.


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Oxidation and reduction

Oxidation of aldehydes with acidified dichromate(VI), Tollens' reagent and Fehling's reagent. Resistance of ketones to oxidation. Reduction of aldehydes and ketones with sodium tetrahydridoborate (sodium borohydride) and lithium tetrahydridoaluminate (lithium aluminium hydride).

Investigation of the reactions of aldehydes and ketones.

Formation of triiodomethane as a test for compounds containing a CH3CO− group or a CH3CH(OH)− group.

Uses of carbonyl compounds Methanal in the manufacture of urea-methanal resin. Propanone as a solvent and a raw material in the manufacture of perspex.


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13.7 Carboxylic acids and their derivatives

Structures of carboxylic acids, acyl chlorides, anhydrides, amides and esters.

The formation of carboxylic acids

Hydrolysis of nitriles. Oxidation of alcohols, aldehydes and alkylbenzenes.

Reactions of carboxylic acids Formation of salts, acyl chlorides, anhydrides, amides and esters. Reduction with lithium tetrahydridoaluminate.

Investigation of the reactions of carboxylic acids. Preparation of an ester.

Acidity of carboxylic acids Comparison of the acidity of carboxylic acids with alcohols. Influence of substituents, viz. alkyl and chloro groups, on acidity.

Reactions of acyl chlorides and anhydrides

Reactions with water, alcohols, ammonia and amines.


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Reactions of amides Hydrolysis, dehydration, Hofmann degradation and reduction with lithium tetrahydridoaluminate.

Reactions of esters Acid and base hydrolyses. Reduction with lithium tetrahydridoaluminate.

Analysis of commercial aspirin tablets.

Uses of carboxylic acids and their derivatives

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.

13.8 Nitrogen compounds Primary, secondary and tertiary aliphatic amines, phenylamine and amino acids.


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The formation of amines Primary amines from nitriles and amides. Primary, secondary and tertiary aliphatic amines, and quaternary ammonium compounds by alkylation. Phenylamine from nitrobenzene.

Base properties of amines Salt formation. Comparison of the basic strength of ammonia, primary aliphatic amines and phenylamine.

Reaction of amines 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 not required.)

Investigation of the reactions of amines.


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Uses of amines and their derivatives

Azo-compounds as dyes in dyeing industries. Amine derivatives as drugs.

Amino acids Amino acids (e.g. aminoethanoic acid and 2-amino-propanoic acid) as bifunctional compounds having both acidic and basic characteristics. Zwitterion. Dipeptides and polypeptides as dimers and polymers of amino acids. (Methods of formation of polypeptides are not required.)

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14. Chemistry and Society ��Topics Explanatory Notes Suggested Experiments/Activities 14.1 Chemistry and the environment

(a) Air pollution Some air pollutants

Carbon monoxide, sulphur dioxide, nitrogen oxides, hydrocarbons, ozone and particulates. Combustion of fossil fuels as the main source of air pollutants.

The effects of polluted air on the environment

The harmful effects of pollutants depend on their concentrations and the duration of exposure to the pollutants. Parts per million (ppm) as one way of indicating concentrations of pollutants. Acid rain and photochemical smog: their formation and effects on the environment.

Topics Explanatory Notes Suggested Experiments/Activities

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The ozone layer and chlorofluorocarbons

Sources and properties of ozone. The desirability of ozone in the stratosphere. Chlorofluorocarbons as aerosol propellants, solvents for the cleaning of 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 chlorofluorocarbons leading to the depletion of the ozone layer. Control of the ozone depletion problem. Possible alternatives for chlorofluorocarbons.

Project work on air pollution, e.g. acid rain, smog or ozone depletion.


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(b) Water pollution The causes of water

pollution and its effects on the environment

The adverse effects on water quality due to livestock waste, oil spillages, residues of pesticide, detergents in sewage, and industrial effluents.

Water quality 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.

Determination of dissolved oxygen in water samples.


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(c) Solid waste Plastics, paper and metals. Disposal of solid waste by landfilling and incineration. Pollution problems associated with the disposal of plastics. Development of degradable plastics and recycling of plastics as possible solutions to pollution problems.

(d) Pollution control in Hong Kong

Measures to improve air quality: use of unleaded petrol and installation of catalytic converters in car exhaust systems, limitation of sulphur content in fuels, desulphurization of flue gas, installation of electrostatic precipitators and installation of low nitrogen oxide burners in power plants.

Visit to (a) the Environmental Resource

Centre, (b) the Chemical Waste Treatment

Centre, or (c) a sewage treatment plant.


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Measures to improve water quality: screening, sedimentation and digestion of pollutants by micro-organisms in the treatment of sewage; physical and chemical methods, and incineration in the treatment of chemical waste from industry and laboratories. (Technical details of the above treatment processes are not required.) Measures to reduce solid waste: reuse/recycling of paper, plastics and metals to minimize waste and save resources.


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14.2 Chemistry and food (a) Principal components of

food

Proteins 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)

Separation of amino acids by paper chromatography.

Carbohydrates

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.

Investigation of the hydrolysis of sucrose and testing for reducing sugars.


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Fats and oils

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 to compare the degree of unsaturation. Hardening of vegetable oils. (Link with Section 13.2) Hydrolytic and oxidative rancidity.

(b) Food preservation

The need to preserve food

Prevention of food spoilage due to microbial activities and chemical changes.


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Principles and techniques of food preservation

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, pickling and the use of food additives.

Investigation of the effects of air and preservatives on apple browning.


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(c) Food additives Food additives to serve as preservatives (e.g. nitrates(III), nitrates(V), sulphur dioxide, sulphates(IV), benzoic acid and benzoates) 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 agents) or nutritional value (e.g. vitamins) of food. Principle of BHA/BHT as antioxidant to retard atmospheric oxidation of oils and fats.

Library search on different functions of common food additives. Analysis of sulphur dioxide content in wine.


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The possible menace of food additives

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.

Debate on the use of food additives.

emb's 2006 al chem syllabus

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