in-depth overview and comparison of 2006 & 2007 al chem syllabus

8/14/2019 In-Depth Overview and Comparison of 2006 & 2007 AL Chem Syllabus

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Section 5 The d-Block Elements

5.1 General Features of d-Block Elements from Sc to Zn

5.2 Characteristic Properties of d-block Elements and Their Compounds

C. Organic Chemistry

Section 1 Fundamentals of Organic Chemistry Structure & Naming

Section 2 Fundamentals of Organic Chemistry Isomerism

Section 3 Fundamentals of Organic Chemistry Organic Acids, Bases and Mechanisms

Section 4 Introduction to Practical Organic Chemistry

Section 5 Alkanes & Alkenes (Aliphatic Hydrocarbons)

Section 6 Aromatic Hydrocarbons

Section 7 Halogeno-compounds

Section 8 Hydroxy Compounds

Section 9 Carbonyl Compounds

Section 10 Carboxylic Acids & their Derivatives

Section 11 Nitrogen Compounds

D. Chemistry and the Environment (Out of Syllabus)

Air Pollution (Out of Syllabus)

Water Pollution (Out of Syllabus)

Solid Waste (Out of Syllabus)

Pollution Control in Hong Kong (Out of Syllabus)

E. Chemistry and Food (Partly Out of Syllabus)

Proteins (In Syllabus)

Carbohydrates (Out of Syllabus)

Fats and oils (In Syllabus)

Food Preservation (Out of Syllabus)


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F. Chemistry In Action (Newly Added)

Polymers

Drugs

Green Chemistry

Physical Chemistry

Section 1 Atoms, Molecules and Stoichiometry

Objectives:

1. Recognize that protons, neutrons & electrons are constituents of the atom

2. State the relative masses & charges of proton, neutron & neutron

3. Recognize that the atomic nucleus is much smaller than the atom

4. Distinguish among relative isotope mass, relative atomic mass & mass no.

5. Recognize that the relative atomic mass of an element is an average value of all the relative isotope

masses of the element

6. Calculate the relative atomic mass from the relative isotopic masses & relative abundances

7. Calculate the relative molecular mass of a compound from relative atomic masses

8. Recognize the mass spectrometer can be used to determine the mass-to-charge ratio of a particle

9. State the main functions of the various parts of the mass spectrometer: vaporization chamber,

ionization chamber, electric field, magnetic field, ion detector & the recorder

10. Recognize that a mass spectrum of an element may consist of many peaks due to the presence of

isotopes with different mass/charge ratio

11. Recognize that the number of C atoms in 12g of carbon-12 is called the Avogadros constant

12. Recognize that equal no. of moles of gases occupy equal volumes at the same temperature and

pressure (Avogadros Law).

13. Recognize that the molar volumes of a gas at R.T.P. and S.T.P. are approximately 24dm and

22.4dm respectively

14. Perform calculation involving mole, e.g. finding the

no. of moles from a given mass of substance or the volume of a gas at R.T.P. or S.T.P.

no. of particles in a given mass of substance

molarity from mass or no. of moles of substance & the volume of solution

15. Use the ideal gas equation: PV=nRT in calculation, e.g. finding density, pressure, volume, relative

molecular mass etc.

16. Recognize that the ideal gas equation is a generalization of the Boyles Law, Charles Law,

pressure law & Avogadros law

17. Define the partial pressure & mole fraction


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18. State the Daltons Law of partial pressures & its derivation from the ideal gas equation

19. Derive the equation: =PM/RT

20. Calculations of empirical and molecular formulae

21. Balancing redox equations

22. Calculations of theoretical yield and percentage yield of a reaction

23. Recognize several types of titrations e.g. acid-base titration, redox titration, back titration

24. State the procedures of volumetric analysis including the preparation of a standard solution

Out of Syllabus

1. Radioactivity

Name of , particles and radiation

Balance equations of nuclear reactions

Uses of radioactive isotopes in leak detection, radiotherapy, nuclear power and as tracers

2. 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

No. of mole of metal deposited during electrolysis = It/nF

, where I = current, t = time, n = no. of mole of electrons, F = Faradays constant (96500C)

Note: Nothing is added in this section

Section 2 The electronic structures of atoms and the periodic table

Objectives:

1. Recognize that the atomic emission spectrum of an element is unique & is an important evidence

for the electronic structure of the atom

2. Recognize that there exists discrete energy levels for the electron(s) in an atom

3. State the cause of the emission spectrum of hydrogen: an electron in a lower energy level can be

excited by the absorption of energy; when this electron falls back to the lower energy level, a

photon with specified frequency or wavelength corresponding to the energy difference is emitted

4. Recognize that the difference in energy between 2 energy levels can be found by the Planks

equation: E=hv

5. Recognize the uniqueness of atomic emission spectra which can be used to identify different

elements

6. State the experimental procedure for flame test

7. Recognize that each series of emission lines of hydrogen converges at high energy end & the


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convergence limit of the Lyman series of the hydrogen atom can be used to find the ionization

enthalpy of hydrogen

8. Recognize that electrons show particle nature and wave nature

9. Recognize that in the wave mechanics model of atom, the electrons in an atom do not localized in

fixed orbitals but move very fast in regions that extend to infinty. These regions are called orbitals,

within which the probability of finding an electron is high

10. Distinguish between shell, subshell & orbital

11. Recognize the pictorial representations of the s, p & d orbitals

12. Recognize the relative energy levels of subshells in

non hydrogen atoms: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p

13. Predict the electronic configuration of an atom or ion for elements with Z


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Energy absorbed to break the bonds of the reactants < Energy released from the formation of

bonds of the products

7. Recognize that for endothermic reactions: (H > 0)

Energy absorbed to break the bonds of the reactants > Energy released from the formation of

bonds of the products

8. State the standard conditions in thermochemistry: 1 atm, 298K & 1 moldm

9. Define the standard enthalpy changes of:

neutralization

solution

formation

combustion

10. Outline the experimental methods for the determination of H, in particular the enthalpy change of

neutralization and state the sources of errors

11. State Hesss Law the total enthalpy change of a chemical reaction is independent of the path

taken

12. Draw enthalpy cycle & enthalpy energy level diagram for givens sets of equations and perform

thermochemical calculations

Newly added objectives

1. Recognize that entropy change (S) is related to the change in the degree of disorder of a system, a

process with increase in randomness has a +ve S, and is more favorable to occur

2. Recognize that free energy change (G = H - TS) is the driving force of a reaction, a

spontaneous reaction should have a ve G

Note: Nothing is cut in this section

Section 4 Bonding and Structure

4.1 Nature of forces holding atoms together

Objectives:

1. Recognize that chemical bonds are strong electrostatic forces holding atoms or ions together,

which are formed by the rearrangement of electrons

2. Recognize that atoms tend to form chemical bonds in order to achieve a noble gas structure

3. Recognize that there are 3 types of chemical bondings: Ionic bond, covalent bond and metallic

bond

4. Sketch energy profiles of the formations of ionic bond & covalent bond

5. Recognize that covalent bonds are formed by the sharing of electrons between 2 atoms


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5. Recognize that there are 6 enthalpy terms in the Born-Haber cycle

6. Recognize that an ionic crystal consists of a large no. of regularly packed ions to form a 3-

dimensional structure

7. Describe the crystal packings & draw diagrams for the crystal structures of NaCl, CsCl

8. Recognize that the formula of an ionic compound can be deduced from the coordination numbers

of the cation & anion in the crystal

9. Define Ionic radius

10. Recognize that an anion is larger than a neutral atom and a cation is smaller than a neutral atom

11. Describe and explain the trends of ionic radii in the Periodic Table

Note: Nothing is cut or added in this section

4.4 Covalent bonding

Objectives:

1. Recognize that a covalent bond is formed by the sharing of electron pair(s)

2. Describe the formation of covalent bond using the concept of the overlapping of atomic orbitals

3. Draw dot and cross diagrams for simple molecules & ions, e.g. CH4, NH3, H2O, HF, NH4+ &

NH2-

4. Recognize that the octet rule had some limitations, e,g, in SF6, PCl5, BF3

5. Recognize that dative covalent bond is a special case of covalent bond (e.g. in H3NBF3) which

is formed by the overlapping of an empty orbital with an orbital by a lone pair of electrons

6. Recognize that bond enthalpies can be used to compare the strengths of covalent bonds

7. Define covalent radius and explain the trends of covalent radii in the Periodic Table

8. State approximate additivity rules of nond enthalpies & covalent radii. Recognize that the break

down of additivity rules indicates that a particular bond is dependent on its environment & the

inadequacy of simple bonding model (e.g.benzene)

9. State the relationship between covalent bond enthalpies and bond lengths as illustrated by

hydrogen halides

10. Recognize that covalent bonds are directional & predict the shapes of molecules & ions using the

electron pairs repulsion theory

11. Recognize that the atomic orbitals of an atom can hybridize to give hybrid orbitals, e.g. sp, sp2 &

sp3

12. Describe the formation of bond and bond by the overlapping of orbitals

13. Recognize that a double bond is made up of a bond and a bond while a triple bond is made up

of a bond and 2 bonds

14. Recognize that the delocalization of electrons exists in some molecules and ions, e.g. benzene,


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CO32- & NO3-

15. Describe the crystal packing & draw diagrams for covalent crystals, e.g. diamond, graphite and

quartz


4.5 Bonding intermediate between ionic & covalent

Objectives:

1. Recognize that pure covalent & ionic bonds are extremes of 2 bonding models

2. Recognize that there may be incomplete electron transfer in an ionic bond (polarization of ionic

bond)

3. State the conditions for the polarization of ionic bonds

4. Recognize that the difference in lattice enthalpies calculated from Born-Haber cycle & from

theoretical value is an evidence for the failure of the pure ionic model

5. Recognize that the electron density in a covalent bond may not be equally shared polar covalent

bond

6. State the evidences for the existence of polar molecules

7. Explain the cause for polar covalent bond in a molecule using the concept of electronegativity

8. Recognize that dipole moment is a vector quantity & the resultant dipole moment of a molecule is

the vector sum of all the bond dipole moments

9. Recognize that the dipole moment of a molecule is governed by

polarity of bond

shape of molecule

position of lone pair(s)

10. Recognize that a molecule with polar bonds may have zero resultant dipole moment because of the

high symmetry of the molecule, e.g. CCl4

11. Recognize that the dipole moment of a molecule can provide information about the shape of the

molecule, e.g. dipole moment of CO2 = 0 supports a linear structure, while the dipole moment of

SO2 > 0 supports a bent structure


4.6 Intermolecular forces

Objectives:

1. Recognize that intermolecular forces are weak forces holding molecules together


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2. Recognize that there are 2 types of intermolecular forces: van der Waals forces & hydrogen bonds

3. Explain the formation of van der Waals forces in terms of

permanent dipoles

instantaneous dipoles

induced dipoles

4. State the 3 types of van der Waals forces:

dipole-dipole attractions (permanent dipole-permanent dipole attractions)

dispersion forces (instaneous dipole-induced dipole attractions)

induction forces (permanent dipole-induced dipole attractions)

5. Recognize that the strength of van der Waals forces increases with the ease of distortion of

electron cloud (or polarizability) of the molecule which in turn increases with molecular size (or

number of electrons)

6. Recognize that van der Waals forces exist in ALL molecules (polar & non-polar)

7. Explain the meaning of van der Waals radius

8. Recognize that the van der Waals radius of a nonmetal is larger than the covalent radius because

the strength of covalent bond is about 100 to 200 times stronger than the van der Waals forces

9. Describe the structure of molecular crystals: iodine, carbon dioxide

10. State the meaning of triple point, critical point, vaporization curve, sublimation curve & fusion

curve in the pressure-temperature diagrams of CO2 & H2O

11. Sketch the pressure-temperature diagrams of CO2 & H2O & state the special features in these

phase diagrams negative slope of fusion curve in H2O & triple point pressure > 1 atm in CO2

12. Recognize that hydrogen bond is formed by a H situated between 2 electronegative elements: F, N

or O

13. Explain the cause & nature of hydrogen bond

14. Describe an experiment to find the approziamate strength of hydrogen bond

15. Recognize that hydrogen bond can be intermolecular or intramolecular

16. Explain the following by the existence of hydrogen bond:

abnormally high b.p. of NH3/H2O/HF over other Group 5/6/7 hydrides

high enthalpy changes of vaporization of alcohols & carboxylic acids

dimerization of carboxylic acid

open structure of ice

special properties of water: high m.p., high b.p., high heat capacity etc.

helical structure of protein & DNA base pairing

Newly added Objectives:

1. Describe the structure molecular crystal: buckminsterfullerene (C60)



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4.7 Structures and properties of substances

Objectives:

1. Recognize that the structure and bonding of a material governs its physical properties

2. Recognize that the m.p. & b.p. of a substance are related to the strengths of the

intermolecular forces (for simple molecular substances)

chemical bonds (for metals, giant ionic & giant covalent substances)

3. Recognize that the electrical conductivity of a substance is due to the presence of delocalized

electrons or mobile ions

4. Recognize that the solubility of a substance in a solvent is dependent on the polarities of the solute

and the solvent

Newly added objectives:

1. Recognize the the properties of bukminsterfullerene are related to its unique structure

2. State the impacts of modern materials, such as semiconductors, nanotubes & liquid crystals on our

daily life


Section 5 Chemical Kinetics

Objectives:

1. Explain the meaning of reaction rate & its units

2. express rate in terms of the change in concentration of reactants or products per unit time

3. Describe the various physical and chemical methods for measuring the rate:

monitor the change in amount (or concentration) of reactant or product by quenching

followed by titration

determining the volume of gas formed at different times

colorimetric measurement of light intensity at different times

4. State the various factors affecting the rate:

--concentration

--temperature

--pressure

--surface area

--temperature

--light

5. Describe experiments for studying the various factors on the reaction rate:

-- concentration: HCL/Mg


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-- temperature: HCl/Na2S2O3

-- particle size: acid/marble chips and acid/powdered CaCO3

-- catalyst: MnO2 on the decomposition of H2O2

-- light: Br2/hexane

6. Explain the meaning of rate equations for zeroth order, first order and second order reactions

7. Recognize that the unit of a rate constant is dependent on the rate equation

8. Recognize that radioactive decay is a first order process

9. State the meaning of half-life of a reaction

10. Derive the relationship for the first order reactions: k=0.6931/t where t= half-life

11. Interpret rate data of concentration ([A]) at various time (t) by plotting appropriate graphs:

-- [A] versus t for zeroth order reaction (negative slope)

-- ln[A] versus t for first order reaction (negative slope)

-- 1/[A] versus t for second order reaction (positive slope)

12. Interpret rate data of initial rates at various initial concentrations of reactant by calculation

13. Recognize that the order of a reaction reflects the molecularity of the rate determining step

14. Recognize that in order to start a chemical reaction, an amount of energy which is larger than or

equal to the activation energy must be supplied

15. Recognize that the activation energy of a reaction can be calculated using the Arrhenius equation: k

= A exp(-Ea/RT) when the rate constants at 2 different temperatures are known

16. Recognize that molecules in a gas show a wide range of speeds the Maxwell-Boltzmann

distribution of molecular speeds

17. Recognize that collision theory can be used to explain the rate of a reaction & recognize the

limitation and inadequacy

18. Explain the effect of temperature on reaction rate based on the Maxwell-Boltzmann distribution

curve

19. Recognize that the course of a reaction can be represented graphically by an energy profile a

graph of potential energy against reaction coordinate

20. Recognize the existence of transition state in an energy profile

21. Distinguish between the order of a reaction & molecularity of an elementary step

22. Recognize that some reactions involve only 1 single step (single stage reactions) while some

reactions involve more than 1 step (multi-stage reactions)

23. Recognize that the mechanism of a reaction is a step-by-step description of how the reaction

occurs

24. Recognize that the rate determining step in a multi-stage reaction is the slowest elementary step

(highest activation energy)

25. Recognize that a catalyst can change BOTH the rate AND mechanism of a reaction by providing

an alternative pathway for the reaction


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26. Explain the difference between homogeneous catalysis & heterogeneous catalysis

27. State examples of homogeneous catalysis & heterogeneous catalysis

28. State the various applications of catalysts:

V2O5 in contact process, Fe in Haber process

Ni (Pt or Pd) in the hydrogenation of unsaturated oils

Pt in catalytic converters

Enzyme


Section 6 Chemical Equilibria

6.1 Dynamic Equilibria

Objectives:

1. Recognize that a chemical equilibrium is dynamic

2. State the characteristics of a chemical equilibrium

It is dynamic

It exists only in a closed system

It can be reached from either directions, quickly or slowly

The relative amounts of the reactants & products do not change

At equilibrium: forward reaction = backward reaction rate 0

The position of equilibrium can be shifted by changes in conditions

The percentage yield for the reversible reaction is always less than 100%

3. State the factors affecting equilibrium position: concentration, pressure & temperature

4. State the Le Chateliers Principle

5. Predict qualitatively the effect of changes in pressure, concentration & temperature on the

equilibrium position using the Le Chateliers principle on the following systems:

Br2 (aq) + H2O (l) H+ (aq) + Br (aq) + HOBr (aq)

Cr2O72 (aq) + H2O (aq) 2CrO42 (aq) + 2H+ (aq)

BiCl3 (aq) + H2O (l) BiOCl (s) + 2H+ (aq) + 2Cl (aq)

N2O4 (g) 2NO2 (g)

6. State the meaning of equilibrium constants Kc & Kp

7. Perform calculation in Kc & Kp

8. Describe the experiments for the determination of Kc in

esterification by titration

Fe+ (aq) + NCS (aq) [Fe(NCS)]+ (aq) by colorimetry

9. Recognize that the equilibrium constant of a reaction is dependent only on the temperature & not


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affected by concentration or pressure

10. Recognize that the shift in equilibrium position by the change in concentration or pressure can be

explained in terms of the equilibrium law

11. Explain the effect of temperature on equilibrium using: lnK = constant (H/RT)

12. Recognize that choosing the operating conditions for the Haber Process & Contact Process can be

explained in terms of equilibrium

13. Recognize that the partition (distribution) of a non-volatile solute in 2 immiscible liquids can be

described using the partition coefficient

14. Calculations involving partition coefficient

15. State the applications of solvent extraction

16. State the principle of paper chromatography


6.2 Ionic Equilibrium

Objectives:

1. Define an acid or base using the Brnsted-Lowry theory

2. Recognize that acid-base equilibria involve the competition for protons

3. Recognize that water undergoes ionization slightly:

2H2O (l) H3O+ (aq) + OH (aq)

4. Define the ionic product of water, Kw

5. Define the pH value as log10[H3O+(aq)]

6. Recognize the dissociation constants of Ka & Kb of weak acid & weak base respectively

7. State the meanings of pKw, pKa & pKb

8. Recognize that a strong acid has a large Ka (small pKa), while a strong base has a large Kb (small

pKb)

9. Recognize that for solutions of the same concentration, a strong acid (or base) has a lower (or

higher) pH value & a higher conductivity

10. Recognize that a strong acid has a weak conjugate base, while a strong base has a weak conjugate

acid

11. Recognize that a buffer can resist the change in pH of a solution

12. Recognize that a more concentrated buffer has a higher buffering capacity (i.e. more resistant to

the change in pH)

13. State the different methods to prepare a buffer


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14. Describe experimental procedure to find Ka & kb by pH measurement

15. Explain the change in colour of acid-base indicators using acid-base equilibra

16. Describe experimental procedure & perform calculation for the titration using double indicators for

a mixture of NaHCO3 & Na2CO3 against standard HCl

17. State & explain the change in pH for acid-base titration

18. Sketch different titration curves for changes in pH & choose suitable indicator or method(s) fr an

acid-base titration

19. Perform calculation in volumetric analysis

20. Perform calculation involving

pH, Ka, Kb & buffer: pH = pKa + log[A]/[HA] or pOH = pKb + log[HB+]/[B]

hydrolysis of salt: conjugate base (Kh = Kw/Ka) & conjugate acid (Kh = Ka = Kw/Kb)

indicator: pH = pKln + log([ln]/[Hln])


1. Define the solubility product Ksp of a sparingly soluble electrolyte, e.g.

-- PbS (Ksp = [Pb2+(aq)][S2(aq)]

-- AgCl (Ksp = [Ag+(aq)][Cl(aq)]

-- Ag2CrO4 (Ksp = [Ag+(aq)][CrO4(aq)]

2. Describe experimental procedure to find Ksp

3. Recognize that for a sparingly soluble compound such as AgCl (aq), the solubility of AgCl can be

decreased by the addition of either Ag+ &/or Cl (common ion effect)


6.3 Redox Equilibria

Objectives:

1. Define oxidation, reduction, oxidant (oxidizing agent) & reductant (reducing agent) in terms of

electron transfer & the change in oxidation number (state)

2. Recognize that a redox equilibrium: oxidant + ne reductant represents the competition for e

3. Use half-equation to represent electron transfer of redox reaction

4. Identify the cathode & anode in an electrochemical cell

5. Recognize that the following are common half-cells:

metal & its ion, e.g. Cu2+ (aq) + 2e Cu (s)

non-metal & its ion, e.g. I2 (aq) + 2e 2I (aq)

ions in different oxidation numbers, e.g. Fe3+ (aq) + e Fe2+ (aq)


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metal & insoluble metal salt, e.g. PbSO4 (s) + 2e Pb (s) + SO42 (aq)

6. State the standard conditions in electrochemistry

7. Explain the meaning of standard electrode (reduction) potential and how to obtain it from

experiment using standard hydrogen electrode

8. State how the standard electrode potential can be used to compare the relative tendency for a

metal/metal ion system to gain or to lose electrons

9. E.m.f. of a cell (with a given cell diagram) E cell = Eright Eleft

10. Recognize that the electrochemical series (given in the exam) can be used to compare the strength

of ozidizing agents or reducing agents


1. Recognize that the electrode potential can be calculated from the molarities of RA & OA by the

Nerst equation:

E = E + (0.059/n)log([oxidant]/[reductant])

2. Use the standard electrode potential to predict the feasibility of redox reactions and state the

limitation of this approach due to kinetic factor

3. State the structure, electrochemical processes & uses of primary cell (e.g. Zn-C cell), secondary cell

(e.g. lead-acid accumulator) & hydrogen-oxygen fuel cell

Out of syllabus:

1. IUPAC conventions in writing cell diagrams (but you must be able to read it)

2. Corrosion of iron and its prevention

The electrochemical process involved in rusting

Prevention corrosion by coating and cathodic protection

Socioeconomic implications of corrosion and prevention

6.4 Phase Equilibrium (Out of syllabus)

1. Two component system:

-- Studies limited to phase diagrams for mixtures of two miscible liquids

(i) vapour pressure against mole fraction (with temperature constant)

(ii) boiling point against mole fraction (with pressure constant)

2. Ideal systems:

Raoults Law. The characteristic properties of an ideal system explained in terms of

molecular interactions

3. Non-ideal system:

Positive and negative deviations from raoults Law explained in terms of molecular

interactions. Enthalpy changes on mixing as evidence of non-ideal behaviour


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Azeotropic mixtures

4. Fractional distillation:

Explaination of the principle of fractional distillation using the boiling point composition

curve

Application of fractional distillation in oil refining

- End of Physical Chemistry -

Inorganic Chemistry

Section 1 Periodic Properties of Elements in the Periodic Table

(Periodic Variations in Physical Properties of the Elements from H to Ar)

Objectives:

1. Recognize that the properties of elements are periodic functions of the atomic number

2. Recognize that elements can be classified into metals (on the left of Periodic Table), Semi-metals

(middle) & non-metals (on the right)

3. Recognize that the properties of elements may be predicted using the Periodic Table

4. Define the following terms:

Ionization enthalpy

Atomic radius (metallic radius, covalent radius, van der Waals radius)

Electron affinity

Electronegativity

Polarizing power of cation

Polarizability of anion

5. Describe and explain the variation of the following physical properties of elements from Li to Ar

Structure & bonding

Melting point

Atomic radius

Ionization enthalpy

Electron affinity

Electronegativity

Polarizing power of cation & polarizability of anion

6. State and explain the diagonal relationship of the following pairs of elements: Li & Mg, Be & Al,

B & Si

7. Recognize that second period elements (Li to F) show anomalous properties from the other

members in the same group



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Section 2 Periodic Properties of Elements in the Periodic Table

(Periodic Relationships Amongst the Oxides of the Elements from Li to Cl)

Objectives:

1. Recognize that elements Li to Cl show a wide range of reactivities towards

water

oxygen

hydrogen

clorine

dilute non-oxidizing acid

2. Recognize that the oxidation states & chemical reactivities of elements are periodic functions of

the atomic numbers

3. Describe the formulae, bonding nature & structures of the oxides of Li to Cl

4. Describe the bonding in oxides & explain the acid/base behavior & hydrolytic behaviour of oxides

in terms of the structures of the oxides

5. Describe & predict the reactions between oxide and

water

dilute acids

dilute alkalis

Out of syllabus

1. Periodic Relationships Amongst the Chlorides and Hydrides of the Elements from Li to Cl

Note: Nothing is added in this section

Section 3 The s-Block elements (Group I & II Elements)

Objectives:

1. State and explain the major characteristics of s-block elements:

high metallic character & low electronegativity


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fixed oxidation state in their compounds

formation of ionic & basic oxides

show characteristic flame colour in flame test

relativity weak tendency to form complexes

2. Outline the experimental procedure for carrying out flame test & memorize the flame colour of s-

block elements, e.g. Na (yellow), K (Lilac), Ca (brick-red), Ba (apple-green)

3. Explain the similarities shown by elements within the same group and their compounds using

electronic configurations and chemical bonding

4. Explain the general trends in the variation of the following physical properties of s-block elements

in terms of bonding & structure:

m.p. & b.p. usually decrease down the group

atomic radius (& ionic radius) increases down the group

1st (or sum of 1st & 2nd) ionization enthalpy decreases down the group

standard reduction potential usually becomes more ve down the group

hydration enthalpy of cations becomes less ve down the group

5. Explain the general trends in the variation of the following chemical properties:

reducing properties

reactivities, e.g. with water & non-metals increase down the group

polarizing power of cations decreases down the group

stabilization of anion (e.g. CO3) by the cation increases down the group

tendency to form covalent bond decreases down the group

tendency to form complexes or hydrated compounds decreases down the group

basic strengths of oxides & hydroxides increase down the group

6. Write equations for the following reactions of s-Block elements

reaction with water

reaction with oxygen

reaction with acids

reaction with nitrogen

7. Describe & explain the following properties of the compounds of s-block elements

Thermal stability of carbonates & hydroxides (in terms of polarization & difference in lattice

enthalpies) for compounds with

(i) small anions e.g. HO, thermal stability decreases down the group

(ii) large anions e.g. CO3, thermal stability increases down the group

Solubility of sulphates(VI) & hydroxides (in terms of hydration enthalpy & lattice enthalpy)

for compounds with

(i) small anions e.g. HO, solubility increases down the group

(ii) large anions e.g. SO4, solubility decreases down the group

8. Recognize that Li & Be show anomalous properties and explain these properties in terms of the


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small cationic sizes and high electronegativities


Section 4 The p-Block Elements

4.1 The Halogens

Objectives:

1. State & explain the major characteristics of the halogens:

high electronegativity

formation of ionic halides & covalent halides with -1 oxidation state

formation of oxoanions & oxoacids with oxidation states from +1 to +7

2. Explain the general trends in the variation of the following physical properties of halogens:

m.p. & b.p.

atomic radii & ionic radii

first ionization enthalpies & electron affinities

standard reduction potentials

3. Explain the similarities shown by the halogens and their compounds using electronic

configurations and chemical bonding

4. Describe the laboratory preparation of the halogens, especially Cl2

5. Describe & explain the following chemical properties of F2, Cl2, Br2 & I2:

variable oxidation states (except for F)

oxidizing properties (decrease down the group)

reactions with water & alkalis (mainly disproportionation)

reactions with metals (e.g. Na + Cl2, Br2 or I2)

reactions with non-metals (e.g. H2 + Cl2, Br2 or I2)

reactions with iron(II) (Cl2 & Br2)

6. Describe the observations and write equations for the following reactions: solid ionic halide (NaX) + conc. H2SO4 formation of HX (if X = F, Cl) or X2 (if X = Br &

I)

solid ionic halide (NaX) + conc. H3PO4 formation of HX

aqueous solution of halide (X) + Ag+ (aq) formation of insoluble AgX (s) can be used

for identification

7. State the necessary conditions for disproportionation

8. Explain the variation of the acid strength of hydrogen halides (HX) and oxoacids (HXOn)

9. Recognize that HF shows anomalous behaviour extensive hydrogen bonding & weaker acidity

than other HX (explain in terms of ion-pair formation)


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10. Predict possible redox reactions involving halogens & halides

11. Recognize that F2 shows anomalous properties and explain these properties in terms of the small

atomic radius and high electornegativity


4.2 Nitrogen and its compounds (Out of syllabus)

4.3 Sulphur and its compounds (Out of syllabus)

4.4 Group IV Elements (Newly added)

Objectives:

1. Explain the general trends in the variation of the following physical properties of Group IV

elements in terms of structure and bonding:

m.p. & b.p.

enthalpies of atomization

atomic radii & ionic radii

first ionization enthalpies

2. Recognize that Group IV elements show different chemical properties down the group (while

Group I, II, VII elements show similarities in the group)

3. Explain the dissimilarities shown by elements within the same group and their compounds using

electronic configurations and chemical bonding

4. Explain the relative stabilities of the +2 and +4 oxidation no. of the Group IV elements using the

concepts of promotion energy, bond energy and inert pair effect

5. Describe the composition and structures of oxides

6. Explain the hydrolytic behaviour and relative stabilities of the chlorides

7. Describe the compositions and structures of oxides

8. Explain the relative stabilities of the +4 & +2 oxides

9. Describe & explain the variation of the acidic, amphoteric and basic properties of the oxides down

the group

10. Describe & explain the uniqueness of carbon

11. State the uses of silicon

12. Describe the bonding and structures of silicates

13. Explain the effect of structure on properties of silicates as exemplified by chain silicates, sheet

silicates and network silicates

14. Recognize the importance of silicon and its compounds such as feldspar, mica and quartz in daily

life


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Section 5 The d-Block Elements

5.1 General Features of d-Block Elements from Sc to Zn

Objectives:

1. Define transition elements & d-block elements

2. Recognize that elements in the first transition series (Sc to Zn) have electronic configurations of

[Ar]3d(1-10)4s(1-2)

3. Draw electrons-in-box diagrams for the first transition series (Sc to Zn) and their ions using

Hunds rule

4. Explain the irregularities in electronic configurations of d-block elements & their ions on terms of

the fact that the energies of 3d & 4s are very close

the extra stabilities of the half-filled d-subshell & full-filled d-subshell, e.g. Cr90), Cu(0),

Mn(II), Fe(III) etc.

5. Explain the general trends in the cariation of the following physical properties of the elements

across the period from Sc to Zn:

ionization enthalpies

electronegativities

m.p. & hardness

densities

atomic (metallic) radii & ionic radii

standard reduction potentials

6. Recognize that similarities in physical properties of transition elements are closely related to the

similarities in

electronic structure

the no. of electrons participated in the metallic bond

atomic radii

7. Explain the formation of coloured compounds & complexes using the concept of splitting of d-

orbitals & the absorption of visible light

8. Recognize that coloured compounds or ions are associated with atoms/ ions with d1 to d9

electronic configuration

9. Compare & explain the differences and similarities between Group I/II elements and transition

elements e.g

stronger metallic bond


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variable oxidation states

formation of complexes

formation of coloured compounds

catalytic behaviour


5.2 Characteristic Properties of d-block Elements and Their Compounds

Objectives:

1. State the 4 major characteristic properties of d-block elements/ transition elements & their

compounds: variable oxidation states, complex formation, coloured ions & catalytic properties

2. Recognize that transition elements can exist in variable oxidation states in various compounds

such as chlorides, oxides & complexes

3. Recognize that the variable oxidation states of transition element are caused by the low successive

ionization enthalpies of that element

4. Explain the relative stabilities of oxidation states in terms of electronic structures

5. Predict possible reactions (& write equations) using standard electrode potentials under acidic,

neutral & alkaline conditions

6. Recognize that the change in oxidation states of transition metal can lead to the change in colour

7. Describe the interconversion between oxidation states of compounds of vanadium (+2, +3, +4, +5)

& manganese (+2, +4, +7)

8. Recognize that a complex is formed from ligands and a metal ion or atom

9. Recognize that a ligand is a Lewis base which donates e pair(s) to a metla ion/atom

10. Recognize that in a complex, the metal ion or atom and the ligands are bonded by dative covalent

bonds / coordinate bonds

11. Give the IUPAC name of a complex

12. Draw the common shapes of complexes: octahedral, tetrahedral & square planar

13. Recognize that geometrical & structural isomerism may exist in complexes & draw

stereostructures of these complexes

14. Recognize that transition metal ions in aqueous solution actually exist as aquo complex whichcan undergo hydrolysis

15. Recognize that the stability of a complex can be expressed in terms of the stability constant which

is dependent on the nature of the metal, nature of the ligand and the no. of ligands attached to the

metal

16. Recognize that a stronger ligand can displace a weaker ligand from a complex due to the formation

of a more stable complex

17. Explain the ligand substitution reactions of Cu(II) complexes in terms of the relative stabilities of

the complexes

18. Explain the formation of coloured compounds & complexes using the concept of splitting of d-


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orbitals

19. State the colour of the hydrated ions of Fe(II), Fe(III), Co(II) & Cu(II)

20. Recognize that transition metals and their ions can act as catalyst

21. Explain the homogeneous catalysis & heterogeneous catalysis of transition metals and their ions

22. State the role of transition metals & their compounds as catalysts in industrial processes

Fe in Haber Process

Fe+ or Fe+ in the reaction between I2 & S2O8

MnO2 in the decomposition of H2O2


- End of Inorganic Chemistry -


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Organic Chemistry

Section 1 Fundamentals of Organic Chemistry Structure & Naming

Objectives:

1. State the natural sources of organic compounds:

Alkanes. Alkenes & aromatic hydrocarbons from crude oil & coal

Carbohydrates, proteins & fats in living organisms

2. Recognize that the uniqueness of carbon in forming a large no. of organic compounds is mainly

due to the ability of carbon to catenate

3. Recognize the 3 types of hybridization of C atom in organic compounds:

sp, e.g. in alkynes

sp. e/g/ in alkenes

sp, e.g. in alkanes

4. Recognize that most organic molecules can be represented by a carbon parent chain linked to one

or more functional group(s)

5. Recognize the function groups of the following compounds & give the systematic names of

organic compounds containing these functional groups:

alkenes

alkynes

halogenocompounds

phenols & alcohols

ethers

aldehydes

ketones

carboxylic acids

1, 2, 3 amines

nitriles

esters

acid halides/ acyl halides

1, 2, 3 amides

acid anhydrides

6. Describe the effect of functional group, relative size & branching of the carbon chain on the

physical properties of a homologous series, e.g. melting points, boiling points & solubilities in


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water

7. Recognize that the shape of an organix molecule can be explained in terms of

the hybridizations of C atoms in the molecule,

the repulsion between electron pairs

8. Describe the structures, shapes & bondings in the following hydrocarbons:

saturated hydrocarbons in alkanes, each C atom is sp hybridized

Unsaturate hydrocarbons in alkenes, each C atom involved in the C=C bond is sp

hybridized. In alkynes, each C atom involved in the CC bond is sp hybridized

Aromatic hydrocarbons In benzene, the 6 C atoms are arranged in a planar 6-membered

ring, each C atom is sp hybridized with delocalization of electrons

9. Recognize that the delocalization of electrons in the benzene ring

leads to the higher stability of benzene over other unsaturated hydrocarbons

gives rise to a unique class of hydrocarbons which are chemically different from alkenes

10. Describe simple tests to

distinguish between saturated (e.g. cyclohexane) & unsaturated compounds (e.g.

cyclohexene) by bromine, manganate(VII) & sulphuric(VI) acid

distinguish benzene from other unsaturated compounds by bromine & manganate(VII)


Section 2 Fundamentals of Organic Chemistry Isomerism

Objectives:

1. Recognize that isomerism can be classified into

structural isomerism

stereoisomerism (which can be further classified into geometrical isomerism & optical

isomerism/ enantiomerism)

2. Recognize that structural isomers are isomers different in the ways the atoms linked. They may be

isomers containing the same functional group or isomers containing different functional groups,

e.g.

branched chains & unbranched chains hydrocarbons

distributed benzenes

ethers & alcohols

ketones & aldehydes

carboxylic acids & esters


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3. Recognized that geometrical isomers in acyclic compounds are caused by the rigidity of C=C bond

(cannot freely rotate)

4. Recognized that geometrical isomerism occurs in alkenes (e.g. but-2-ene & hexan-1,3-dienes) &

its derivatives (e.g. butenedioic acid)

5. Recognize that a pair of geometrical isomers such as cis-/trrans-butenedioic acid have different

physical properties (e.g. m.p. & solubility in water) & chemical properties (e.g. acidity & ease of

dehydration on heating)

6. Recognize that enantiomerism is usually caused by the presence of chiral centre(s) in an organic

molecule. A chiral centre in organic compound is an asymmetric carbon atom linked to 4 different

atoms or groups of atoms

7. Recognize that a pair of enantiomers

are non-superimposable mirror images of each other (i.e. chiral)

have identical physical properties except different effects on plane polarized light

have identical chemical roperties towards optically inactive reagent

8. Recognize that an equimolar mixture of a pair of enantiomers does not show optical activity. This

mixture is known as a recemate or recemic mixture

9. Recognize that diastereomers (diastereoisomers) (e.g. geometrical isomers) are stereoisomers that

are not enantiomers

Newly Added Objectives:

1. Recognize that many drugs are chiral


Section 3 Fundamentals of Organic Chemistry Organic Acids, Bases and Mechanisms

Objectives:

1. Recognize that organic reactions can be rationalized & classified using the electronic theory

rearrangement of e of reactants to give products

2. Describe the 2 ways to break a covalent bond:

homolysis (symmetrical fission)

heterolysis (unsymmetrical fission)

3. Describe the 3 types of reactive species in organic chemistry & give examples:

free radical (usually with 7 outermost e)

electrophile (usually with 6 outermost e, seeks for negative centres)


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nucleophile (usually with 8 outermost e, seeks for positice centres)

4. Recognize that a negatively charged nucleophile (e.g. HO) is a stronger nucleophile than its

conjugate acid (e.g. H2O)

5. Recognize that a curly arrow is used to describe the movement of 2 es & a fishhook arrow is used

to describe the movement of 1 e

6. Recognize that the electronic effects and steric effects are major factors affecting organic reactions

7. Recognize that there are 2 types of electronic effects: inductive effects & mesomeric (resonance)

effects. Electronic effects can be electron-donating or electron-withdrawing

8. Recognize that an alkyl group is electron-donating

9. Predict & explain the relative stabilities of the carbocations

10. Recognize that the acidity of acid and the basicity of a base can be explained in terms of the

relative stabilities of the reactant & product

11. Predict & explain the following relative acidities in terms of the stability of the anion formed:

stability of anion: R-O < H-O 3) aliphatic amines > ammonia > aromatic amines (e.g. phenylamine)

13. Describe the mechanisms (using curly arrows & fishhook arrows) & give examples of the 6 major

types of organic reactions:

free radical substitution : chlorination of alkane (chain reaction mechanism)

e.g. CH4 + Cl2 CH3Cl + HCl

electrophilic addition on C=C obeys the Markownikoffs rule

e.g. RCH=CH2 + H-Br RCHBrCH3

electrophilic substitution on benzene ring : nitration of benzene (newly added)

nucleophilic substitution SN1 (2 Steps) & SN2 (1 Step)

e.g. RX + OH ROH + X (X = Halogen)

nucleophilic addition on carbonyl compounds

e.g. RCHO + HCN RCH(OH)CN

nucleophilic acyl substitution

RCOCl + H2O RCOOH + HCl


Section 4 Introduction to Practical Organic Chemistry

Objectives:


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1. Recognize that the molecular formula of a compound can be determined from analytical data (e.g.

% mass) & relative molecular mass

2. Recognize that the structure of a compound can be deduced:

from reactions of functional groups & physical properties

by spectroscopic methods such as mass spectrometry (MS), Infra-red (IR) spectroscopy &

nuclear magnetic resonance

3. Recognize that infra-red spectroscopy depends on the vibrations of atoms with respect to each

other in the molecule. The infra-red spectrum can tell us whether particular functional groups are

present in a molecule

4. Recognize that different functional groups have characteristic absorption wavenumbers

5. Use infra-red spectrum in the identification of the following groups C-H, O-H, N-H, C=C, CC,

C=O & CN in organic compounds

6. State the working principle of mass spectrometry

7. Recognize that from a mass spectrum, it is possible to obtain structural information from masses of

molecular ions & fragments

8. Recognize that nuclear magnetic resonance depends on the magnetic properties of the atomic

nuclei such as hydrogen nuclei in a molecule. The proton NMR spectrum can tell us the number of

hydrogen nuclei present in the molecule & give information avout the structural environment of

the hydrogen

9. State the uses of the following practical methods in organic chemistry

heating under reflux

purification methods such as adsorption, drying, recrystallization, solvent extraction,

distillation, fractional distillation, chromatography

10. Recognize that a pure compoundhas a sharp (or narrow range of) melting / boiling point, while an

impure compound has a wide range of melting / boiling point


Section 5 Alkanes & Alkenes (Aliphatic Hydrocarbons)

Objectives:

1. State the natural sources of hydrocarbons: crude oil, coal & natural gas

2. Recognize that crude oil is the major source of alkanes & other hydrocarbons & coal is the major

source of aromatic hydrocarbons

3. State the chemical principles & economic importance of the fraction distillation of crude oil


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4. State the main uses of alkanes: as fuels & raw material for cracking

5. State the uses of LPG, naphtha, petrol, kerosene & gas oil

6. Recognize that extensive use of fossil fuels may lead to global warming

7. Recognize that the inertness if alkanes is due to the essentially non- polar nature of the C-H bond

which does not facour heterolysis

8. State the major reactions of alkanes: combustion, free radical substitution (halogenation) &

cracking

9. Outline the free radical substitution mechanism of the chlorination/bromination of alkane

10. Explain the distribution of % yield of the monohalogenated products in terms of the stability of

free radicals

11. Describe the bonding in alkenes

12. State the major synthesis of alkenes: dehydration of ROH & dehydrohalogenation of RX at high

temperature

13. Recognize that alkenes can be obtained industrially from the cracking of alkanes followed by

fractional distillation

14. Recognize that a more substituted alkene is more stable than a less substituted on ( Sayteffs rule)

15. Recognize that the high reactivities of alkenes over alkanes are due to the presence of the

electrons rich C=C double bond which can be attacked by electrophile or radical

16. State the major reactions of alkenes:

electrophilic addition of HBr

electrophilic addition of Br2 (in aqueous & non-aqueous solvents)

electrophilic addition of conc. H2SO4 & acid-catalyzed hydration of alkenes

catalytic hydrogenation

ozonolysis

reaction with dilute neutral/alkaline KMnO4

polymerization of ethane, propene and phenylethene to give poly(ethene), poly(propene) &

poly(phenylethene) respectively

17. Outline the mechanism for the electrophilic addition of hydrogen halide on alkene

18. State the Markownikoffs rule for the electrophilc addition of hydrogen halide on alkene19. Recognize that the Markownikoffs rule can be explained in terms of the stability of the

carbocation intermediate

20. Recognize that catalytic hydrogenation is used for the hardening of oils (making margarine)

21. State the reaction conditions & products of the ozonolysis of alkenes

22. Recognize that ozonolysis can be used to determine the positions of C=C bonds in alkenes

23. Recognize that the properties of poly(alkene) are governed by the internal structure of the polymer

24. Recognize that the IR spectra of alkenes contain a peak at around 1645 cm-1 due to the stretching

of the C=C bond

25. Describe the tests for alkenes using bromine & potassium manganate (VII)


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Section 7 Halogeno-compounds

Objectives:

1. Recognize that halogenoalkanes (alkyl halide, RX) are classified into 1, 2, 3

2. State the major synthesis of halogeno-compounds (X = halogen)

Halogenoalkanes: ROH + PCl5 or SOCl2 / HCl + ZnCl2 / P + Br2 / P + I2

Halogenoalkanes: electrophilic addition of HX or X2 on C=C or CC

Halogenoarenes: arene + X2 + Fe / X2 + Al / X2 + AlX3

Halogenoalkanes or halogenoarene: R-H or Ar-R + Cl2/ Br2 + UV light

3. Explain the relative reactivities of ROH towards halogenation: 3 > 2 > 1

4. Outline the Lucas test for alcohols

5. Recognize that the presence of halogen in halogenoalkanes can increase the reactivity of

halogenoalkanes towards nucleophilic substitution & elimination

6. Recognize that the C-X bond in halogenoarene (Ar-X) possesses some double bond character due

to the overlappong of the lone pair p orbital of X with the orbitals of the benzene ring to give a

stable delocalized system

7. Describe the SN1 mechanism (for 3 RX) and SN2 mechanism (for 1, 2 RX) for the substitution

of X in alkyl halides by OH

8. Recognize & explain the relative reactivities for the hydrolysis of R-X & Ar-X:

3 R-X > 1 R-X > 2 R-X >> halogenoarenes

RCl < R-Br< R-I

9. State the major reactions of R-X:

substitution by aqueous OH to give R-OH

substitution by CN to give R-CN

substitution by NH3 (& amines) to give R-NH2 (& other amines)

elimination by alcoholic KOH/NaOH to give C=C

10. State the preparations & reactions of dihaloalkanes

11. Recognize that chloroethene undergoes addition polymerization to give poly(chloroethene) [PVC]

12. State the properties of PVC as related to the structure of the polymer

13. Describe tests to distinguish between R-X & Ar-X

14. Recognize that the IR spectra (in the range of 1400 to 4000 cm-1) of R-X (or Ar-X) are similar to

those of alkanes (or arenes)

15. State the applications of halogeno-compounds: solvents in dry-cleaning, raw materials in the


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manufacture of poly(chloroethene) & poly(tetrafluoroethene), CFC as aerosol propellant, BFC &

BTM as fire extinguishing agents

16. Recognize that halogeno-compounds may cause pollution problems


Section 8 Hydroxy Compounds

(this section is combined with halogeno-compounds, refer to section 3 & 7)

Section 9 Carbonyl Compounds

Objectives:

1. Describe structures of aldehydes & ketones & recognize that the C in C=O is sp hybridized

2. Recognize that benzaldehyde & phenylethanone are examples of aromatic carbonyl compounds

3. State the major preparations of:

aldehyde by the mild oxidation of 1 alcohol RCH2OH

ketone by the oxidation of 2 alcohol R2CHOH

aldehyde & ketone by the hydrolysis of gem-dihalides (e.g. R2Cl2)

4. Recognize that the C in C=O is a positive centre & can be attacked by nucleophiles

5. Describe the mechanism of the nucleophilic addition of HCN on carbonyl compounds

6. Explain the order of reactivities of aldehydes & ketones towards nucleophilic addition reactions in

terms of electronic & steric sffects of R & Benzene

H2C=O > RCHO > R2CHO > benzaldehyde > phenylalkanone > 1,2-diphenylmethanone

7. State the main reactions of aldehydes & ketones:

Nucleophilic addition of HCN

Nucleophilic addition of NaHSO3

Addition-elimination (condensation) with NH2OH & 2,4-dinitrophenylhydrazine

Oxidation (aldehyde can be easily oxidized to RCOOH, but ketone is resistant to oxidation

& cannot be easily oxidized)

Reduction to ROH by LiAlH4 & NaBH4 (aldehyde 1 ROH, ketone 2 ROH)

Triiodomethane reaction for CH3COR(H) (an oxidation reaction)

Oxidation by Cu(II) as in Fehlings reagent and Ag(I) as in Tollens reagent

9. Recognize that the condensation reaction is in fact an addition followed by elimination

10. Recognize that the reaction between sodium hydrogensulphate(VI) & carbonyl compounds can be

used for the purification of carbonyl compounds

11. Recognize that the formation of oximes or 2,4-dinitrophenylhydrazones of carbonyl compounds


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can be used for the identification of carbonyl compounds

12. Describe the identification of aldehydes & ketones by the

preparation of the derivative of a carbonyl compound, e.g. 2,4-dinitrophenylhydrazone

Tollens reagent (silver mirror test)

Fehlings solution

13. State the chemicals required & the intermediates for the following synthetic conversions:

RCHO RCH(OH)COOH

RCOCH3 RCH2OH

RCOCH3 RCH2COOH

14. Recognize that the IR spectra of carbonyl compounds show a strong & sharp peak at 1680 to 1750

cm-1 due to the strectching of the C=O bond

15. State the use of methanal in the manufacture of condensation polymers (phenol-methanal & urea-

methanal) & the use of propanone as a solvent and a raw material in the manufacture of the

addition of polymer Perspex


Section 10 Carboxylic Acids & their Derivatives

Objectives:

1. Describe the structures & bonding of carboxylic acids, acyl chlorides (acid chlorides), acid

anhydrides, amides & esters

2. State the 3 special features of carboxylic acids

high acidity of carboxylic acids over other organic compounds

less likely to undergo nucleophilic addition than carbonyl compounds

presence of extensive hydrogen bonding

3. State the major preparations of carboxylic acids:

hydrolysis of nitrile, amide or ester

oxidation of 1 alkanols

oxidation of aldehydes

cigourous oxidation of side chain of aromatic compounds

4. Predict & explain the acidity of the carboxylic acid in terms of

the equilibrium: RCOOH + H2O H3O+ RCOO

the stability of the carboxylate ion due to delocalization

the influence of substituents by electronic effects on the stability of the carboxylate ion


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5. State the main reactions of RCOOH

neutralization: formation of salt

formation of esters [RCOOR], acyl chlorides [RCOCl], acid anhydrides [(RCO)2O] &

amides [RCONH2]

reduction to RCH2OH by LiAlH4

6. State the procedure for the separation of carboxylic acid from other organic substances by the

extraction with Na2CO3(aq) or NaHCO3(aq)

7. Explain the differences in chemical & physical properties of cis-butenedioic acid & trans-

butenedioic acid

8. State the chemicals, intermediates & conditions for the following conversions:

RCH2Cl RCOOR

RCl RCH2Cl (1 C added)

RCOOH RCl or ROH (1 C removed)

RCH2Cl RCl (1 C removed)

ROH RCOOH (1 C added)

9. State the general mechanism of nucleophilic acyl substitution as exemplified by hydropysis of acyl

chloride

10. Explain the reactivities of carboxylic acid derivatives:

RCOCl > (RCO)2O > RCOOR > RCONH2

11. State the major preparations of acyl chlorides, acid anhydrides, amides & esters

12. State the main reactions of acyl chlorides

reaction with water hydrolysis to give RCOOH

reaction with ROH to give RCOOR

reaction with RCOO - formation of acid anhydrides [(RCO)2O]

reaction with NH3 or amines formation of amides [e.g. RCONH2]

13. State the main reactions of esters

acid hydrolysis or alkaline hydrolysis with HCl (aq) or NaOH (aq) to give RCOOH or

RCOONa+ respectively

reduction to RCH2OH ( + ROH)

14. Recognize that fats & oils are esters. On hydrolysis, fats & oils can be hydrolysed into propane-

1,2,3-triol & fatty acids

15. Recognize that oil contains a higher % of fatty acids with unsaturated hydrocarbon chains than fat

16. Define the iodine value & recognize that the degree of unsaturation of fats & oils can be compared

in terms of the iodine value

17. Recognize that the hardening of vegetable oils is the hydrogenation process of the unsaturated

hydrocarbon chains in oils

18. State the main reactions of amides

alkaline hydrolysis to give RCOONa+


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Hofmann Degradation of RCONH2 to give RNH2

Reduction to amines

19. Describe experiments on

oxidation of 1 or 2 ROH, RCHO or alkylarenes

hydrolysis of nitrile or amide (e.g. preparation of benzoic acid from benzamide)

analysis of commercial asprin tablets

the formation of acyl chlorides

esterification

acid/ alkaline hydrolysis of esters

20. Recognize that the IR spectra of carboxylic acids & their derivatives show a strong & sharp peak

at 1680 to 1750 cm-1 due to the vibration of the C=O bond. In addition, carboxylic acids show a

broad band in the region 2500 to 3300 cm-1 due to O-H stretching

21. State the uses of carboxylic acids & their derivatives

benzoic acid & benzoates as food preservatives

polyesters (e.g. terylene or Dacron) and polyamides (e.g. nylon 6.6) as synthetic fibres

esters as flavourings

esters as fat / oil

22. Recognize theat the rancidity of fats & oils is caused by the hydrolysis &/or oxidation (free

radicals reactions)

23. Describe the principle of antioxidants ( e.g. BHA (butylated hydroxyanisole) & BHT (butylated

hydroxytoluene) to prevent the autoxidation of fat/oil


Section 11 Nitrogen Compounds

(this section is combined with carboxylic acid and its derivatives, refer to section 10)

In Syllabus

1. Preparation of amines, amides

2. Reactions of amines, amides

formation of diazonium salt from R-NH2, phenylamine

coupling reaction: diazonium salt from 1 aromatic amine reacts with e.g. phenol or

naphthalein-2-ol to give azo-dyes

Hofmann degradation of amides

3. Amino acids

bifunctional compounds having both acid and basic properties (Zwitterion)


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dipeptides and polypeptides of amino acids

- End of Organic Chemistry-

D. Chemistry and the Environment (Out of Syllabus)

12.1 Air Pollution (Out of Syllabus)

12.2 Water Pollution (Out of Syllabus)

12.3 Solid Waste (Out of Syllabus)

12.4 Pollution Control in Hong Kong (Out of Syllabus)

E. Chemistry and Food (Partly Out of Syllabus)

13.1 Proteins (In Syllabus)

1. amino acids

2. Hydrolysis of proteins

13.2 Carbohydrates (Out of Syllabus)

1. Structures of glucose, fructose, sucrose

2. Hydrolysis of sucrose

3. Fehlings test to distinguish between reducing and non-reducing sugars

13.3 Fats and oils (In Syllabus)

1. Hydrolysis of fats and oils

2. Use of iodine value

3. Hydrolytic and oxidative rancidity of fats/oils

Note: refer to section 10 carboxylic acid and its derivatives

13.4 Food Preservation (Out of Syllabus)

1. Techniques of food preservation (chilling, canning , etc.)

2. Food additives (except benzoic acid, BHT, BHA)

3. Menace of food additives

-End of Chemistry and the Environment and Food-


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F. Chemistry In Action (Newly Added)

14.1 Polymers

1. Natural occurring polymers

-- proteins, polysacchrides, DNA

2. Synthetic polymers

-- Addition polymers e.g. Poly(ethene), poly(propene), polystyrene, polyvinyl chloride PVC, etc.

--Condensation polymers e.g. Nylon, urea-methanal, Dacron

3. Effect of structure on properties such as density, hardness, rigidity, elasticity and biodegradability

14.2 Drugs

1. Key stages of drug development exemplified by asprin

-- lead compound discovery

-- molecular modification

-- formulation development

-- safety tests and human trials

-- approval for marketing

2. Narcotic drugs such as morphine and heroin and their adverse effects

3. Stimulants such as ketamine and phenylethanamine and their adverse effects

14.3 Green Chemistry

1. Green chemistry practices exemplified by

-- decaffeination using superficial carbon dioxide

-- the use of H2O2 in the presence of manganese based catalyst as bleaching agent

-End of Chemistry and Action-

in-depth overview and comparison of 2006 & 2007 al chem syllabus

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