Syllabus Of STPM Chemistry

Syllabus STPM Chemistry

SYLLABUS STPM CHEMISTRY

First Term

Second Term

Third Term

First Term

1 Atoms, Molecules and Stoichiometry

1.1 Fundamental particles of an atom

Candidates should be able to:

(a) describe the properties of protons, neutrons and electrons in terms of their relative charges and relative masses;
(b) predict the behaviour of beams of protons, neutrons and electrons in both electric and magnetic fields;
(c) describe the distribution of mass and charges within an atom;
(d) determine the number of protons, neutrons and electrons present in both neutral and charged species of a given proton number and nucleon number;
(e) describe the contribution of protons and neutrons to atomic nuclei in terms of proton number and nucleon number;
(f) distinguish isotopes based on the number of neutrons present, and state examples of both stable and unstable isotopes.

1.2 Relative atomic, isotopic, molecular and formula masses

Candidates should be able to:

(a) define the terms relative atomic mass, Ar, relative isotopic mass, relative molecular mass, Mr, and relative formula mass based on 12C;
(b) interpret mass spectra in terms of relative abundance of isotopes and molecular fragments;
(c) calculate relative atomic mass of an element from the relative abundance of its isotopes or its mass spectrum.

1.3 The mole and the Avogadro constant

Candidates should be able to:

(a) define mole in terms of the Avogadro constant;
(b) calculate the number of moles of reactants, volumes of gases, volumes of solutions and concentrations of solutions;
(c) deduce stoichiometric relationships from the calculations above.

 

2 Electronic Structures of Atoms

2.1 Electronic energy levels of atomic hydrogen

Candidates should be able to:

(a) explain the formation of the emission line spectrum of atomic hydrogen in the Lyman and Balmer series using Bohr’s Atomic Model.

2.2 Atomic orbitals: s, p and d

Candidates should be able to:

(a) deduce the number and relative energies of the s, p and d orbitals for the principal quantum numbers 1, 2 and 3, including the 4s orbitals;
(b) describe the shape of the s and p orbitals.

2.3 Electronic configuration

Candidates should be able to:

(a) predict the electronic configuration of atoms and ions given the proton number (and charge);
(b) define and apply Aufbau principle, Hund’s rule and Pauli Exclusion Principle.

2.4 Classification of elements into s, p, d and f blocks in the Periodic Table

Candidates should be able to:

(a) identify the position of the elements in the Periodic Table as
(i) block s, with valence shell configurations s1 and s2,
(ii) block p, with valence shell configurations from s2p1 to s2p6,
(iii) block d, with valence shell configurations from d1s2 to d10s2;
(b) identify the position of elements in block f of the Periodic Table.

 

3 Chemical Bonding

3.1 Ionic bonding

Candidates should be able to:

(a) describe ionic (electrovalent) bonding as exemplified by  NaCl and MgCl2. other ionic compounds: NaCl, KF, BaO, Na2O, MgBr2, MgCl2, Al2O3, BaCl2, CaO

3.2 Covalent bonding

Candidates should be able to:

(a) draw the Lewis structure of covalent molecules (octet rule as exemplified by NH3, CCl4, H2O, CO2, N2O4 and exception to the octet rule as exemplified by BF3, NO, NO2, PCl5, SF6);
(b) draw the Lewis structure of ions as exemplified by SO42, CO32, NO3 and CN;
(c) explain the concept of overlapping and hybridisation of the s and p orbitals as exemplified by BeCl2, BF3, CH4, N2, HCN, NH3 and H2O molecules;
(d) predict and explain the shapes of and bond angles in molecules and ions using the principle of valence shell electron pair repulsion, e.g. linear, trigonal planar, tetrahedral, trigonal bipyramid, octahedral, V-shaped, T-shaped, seesaw and pyramidal;
(e) explain the existence of polar and non-polar bonds (including CC1, CN, CO, CMg) resulting in polar or/ and non-polar molecules;
(f) relate bond lengths and bond strengths with respect to single, double and triple bonds;
(g) explain the inertness of nitrogen molecule in terms of its strong triple bond and non-polarity;
(h) describe typical properties associated with ionic and covalent bonding in terms of bond strength, melting point and electrical conductivity;
(i) explain the existence of covalent character in ionic compounds such as A12O3, A1I3 and LiI;
(j) explain the existence of coordinate (dative covalent) bonding as exemplified by H3O+, NH4+, A12C16 and [Fe(CN)6]3.

3.3 Metallic bonding

Candidates should be able to:

(a) explain metallic bonding in terms of electron sea model.

3.4 Intermolecular forces: van der Waals forces and hydrogen bonding

Candidates should be able to:

(a) describe hydrogen bonding and van der Waals forces (permanent, temporary and induced dipole);
(b) deduce the effect of van der Waals forces between molecules on the physical properties of substances;
(c) deduce the effect of hydrogen bonding (intermolecular and intramolecular) on the physical properties of substances.

 

4 States of Matter

4.1 Gases

Candidates should be able to:

(a) explain the pressure and behaviour of ideal gas using the kinetic theory;
(b) explain qualitatively, in terms of molecular size and intermolecular forces, the conditions necessary for a gas approaching the ideal behaviour;
(c) define Boyle’s law, Charles’ law and Avogadro’s law;
(d) apply the pV nRT equation in calculations, including the determination of the relative molecular mass, Mr;
(e) define Dalton’s law, and use it to calculate the partial pressure of a gas and its composition;
(f) explain the limitation of ideality at very high pressures and very low temperatures.

4.2 Liquids

Candidates should be able to:

(a) describe the kinetic concept of the liquid state;
(b) describe the melting of solid to liquid, vaporisation and vapour pressure using simple kinetic theory;
(c) define the boiling point and freezing point of liquids.

4.3 Solids

Candidates should be able to:

(a) describe qualitatively the lattice structure of a crystalline solid which is:
(i) ionic, as in sodium chloride,
(ii) simple molecular, as in iodine,
(iii) giant molecular, as in graphite, diamond and silicon(IV) oxide,
(iv) metallic, as in copper;

(b) describe the allotropes of carbon (graphite, diamond and fullerenes), and their uses.

4.4 Phase diagrams

Candidates should be able to:

(a) sketch the phase diagram for water and carbon dioxide, and explain the anomalous behaviour of water;
(b) explain phase diagrams as graphical plots of experimentally determined results;
(c) interpret phase diagrams as curves describing the conditions of equilibrium between phases and as regions representing single phases;
(d) predict how a phase may change with changes in temperature and pressure;
(e) discuss vaporisation, boiling, sublimation, freezing, melting, triple and critical points of H2O and CO2;
(f) explain qualitatively the effect of a non-volatile solute on the vapour pressure of a solvent, and hence, on its melting point and boiling point (colligative properties);
(g) state the uses of dry ice.

 

5 Reaction Kinetics

5.1 Rate of reaction

Candidates should be able to:

(a) define rate of reaction, rate equation, order of reaction, rate constant, half-life of a first-order reaction, rate determining step, activation energy and catalyst;
(b) explain qualitatively, in terms of collision theory, the effects of concentration and temperature on the rate of a reaction.

5.2 Rate law

Candidates should be able to:

(a) calculate the rate constant from initial rates;
(b) predict an initial rate from rate equations and experimental data;
(c) use titrimetric method to study the rate of a given reaction.

5.3 The effect of temperature on reaction kinetics

Candidates should be able to:

(a) explain the relationship between the rate constants with the activation energy and temperature using Arrhenius equation k = Ae ^[-(Ea/RT)]
(b) use the Boltzmann distribution curve to explain the distribution of molecular energy.

5.4 The role of catalysts in reactions

Candidates should be able to:

(a) explain the effect of catalysts on the rate of a reaction;
(b) explain how a reaction, in the presence of a catalyst, follows an alternative path with a lower activation energy;
(c) explain the role of atmospheric oxides of nitrogen as catalysts in the oxidation of atmospheric sulphur dioxide; (d) explain the role of vanadium (V) oxide as a catalyst in the Contact process;
(e) describe enzymes as biological catalysts.

5.5 Order of reactions and rate constants

Candidates should be able to:

(a) deduce the order of a reaction (zero-, first- and second-) and the rate constant by the initial rates method and graphical methods;
(b) verify that a suggested reaction mechanism is consistent with the observed kinetics;
(c) use the half-life (t½) of a first-order reaction in calculations.

 

6 Equilibria

6.1 Chemical equilibria

Candidates should be able to:

(a) describe a reversible reaction and dynamic equilibrium in terms of forward and backward reactions;
(b) state mass action law from stoichiometric equation;
(c) deduce expressions for equilibrium constants in terms of concentrations, Kc, and partial pressures, Kp, for homogeneous and heterogeneous systems;
(d) calculate the values of the equilibrium constants in terms of concentrations or partial pressures from given data;
(e) calculate the quantities present at equilibrium from given data;
(f) apply the concept of dynamic chemical equilibrium to explain how the concentration of stratospheric ozone is affected by the photodissociation of NO2, O2 and O3 to form reactive oxygen radicals;
(g) state the Le Chatelier‟s principle and use it to discuss the effect of catalysts, changes in concentration, pressure or temperature on a system at equilibrium in the following examples:
(i) the synthesis of hydrogen iodide,
(ii) the dissociation of dinitrogen tetroxide,
(iii) the hydrolysis of simple esters,
(iv) the Contact process,
(v) the Haber process,
(vi) the Ostwald process;

(h) explain the effect of temperature on equilibrium constant from the equation, ln K= -H/RT + C

6.2 Ionic equilibria

Candidates should be able to:

(a) use Arrhenius, BrØnsted-Lowry and Lewis theories to explain acids and bases;
(b) identify conjugate acids and bases;
(c) explain qualitatively the different properties of strong and weak electrolytes;
(d) explain and calculate the terms pH, pOH, Ka, pKa, Kb, pKb, Kw and pKw from given data;
(e) explain changes in pH during acid-base titrations;
(f) explain the choice of suitable indicators for acid-base titrations;
(g) define buffer solutions;
(h) calculate the pH of buffer solutions from given data;
(i) explain the use of buffer solutions and their importance in biological systems such as the role of H2CO3 / HCO3 in controlling pH in blood.

6.3 Solubility equilibria

Candidates should be able to:

(a) define solubility product, Ksp;
(b) calculate Ksp from given concentrations and vice versa;
(c) describe the common ion effect, including buffer solutions;
(d) predict the possibility of precipitation from solutions of known concentrations;
(e) apply the concept of solubility equilibria to describe industrial procedure for water softening.

6.4 Phase equilibria

Candidates should be able to:

(a) state and apply Raoult’s law for two miscible liquids;
(b) interpret the boiling point-composition curves for mixtures of two miscible liquids in terms of ideal‟ behaviour or positive or negative deviations from Raoult’s law;
(c) explain the principles involved in fractional distillation of ideal and non ideal liquid mixtures;
(d) explain the term azeotropic mixture;
(e) explain the limitations on the separation of two components forming an azeotropic mixture;
(f) explain qualitatively the advantages and disadvantages of fractional distillation under reduced pressure.

 

Second Term

 

7 Chemical Energetics

7.1 Enthalpy changes of reaction, dH

Candidates should be able to:

(a) explain that most chemical reactions are accompanied by enthalpy changes (exothermic or endothermic);

(b) define enthalpy change of reaction, H, and state the standard conditions;

(c) define enthalpy change of formation, combustion, hydration, solution, neutralisation, atomisation, bond energy, ionisation energy and electron affinity;

(d) calculate the heat energy change from experimental measurements using the relationship: heat change, q mcT or q = mc ;

(e) calculate enthalpy changes from experimental results.

7.2 Hess’ law

Candidates should be able to:

(a) state Hess’ law, and its use to find enthalpy changes that cannot be determined directly, e.g. an enthalpy change of formation from enthalpy changes of combustion;

(b) construct energy level diagrams relating the enthalpy to reaction path and activation energy;

(c) calculate enthalpy changes from energy cycles.

7.3 Born-Haber cycle

Candidates should be able to:

(a) define lattice energy for simple ionic crystals in terms of the change from gaseous ions to solid lattice;

(b) explain qualitatively the effects of ionic charge and ionic radius on the numerical magnitude of lattice energy values;

(c) construct Born-Haber cycle for the formation of simple ionic crystals.

7.4 The solubility of solids in liquids

Candidates should be able to:

(a) construct energy cycles for the formation of aqueous solutions of ionic compounds;

(b) explain qualitatively the influence on solubility of the relationship between enthalpy change of solution, lattice energy of solid and enthalpy change of hydration or other solvent-solute interaction.

 

8 Electrochemistry

8.1 Half-cell and redox equations

Candidates should be able to:

(a) explain the redox processes and cell diagram (cell notation) of the Daniell cell;

(b) construct redox equations.

8.2 Standard electrode potential

Candidates should be able to:

(a) describe the standard hydrogen electrode;

(b) use the standard hydrogen electrode to determine standard electrode potential (standard reduction potential), Eº;

(c) calculate the standard cell potential using the Eo values, and write the redox equations;

(d) predict the stability of aqueous ions from Eº values;

(e) predict the power of oxidising and reducing agents from Eº values;

(f) predict the feasibility of a reaction from ‘E note cell’ value and from the combination of various electrode potentials: spontaneous and non-spontaneous electrode reactions.

8.3 Non-standard cell potentials

Candidates should be able to:

(a) calculate the non-standard cell potential, Ecell, of a cell using the Nernst equation.

8.4 Fuel cells

Candidates should be able to:

(a) describe the importance of the development of more efficient batteries for electric cars in terms of smaller size, lower mass and higher voltage, as exemplified by hydrogen-oxygen fuel cell.

8.5 Electrolysis

Candidates should be able to:

(a) compare the principles of electrolytic cell to electrochemical cell;

(b) predict the products formed during electrolysis;

(c) state the Faraday’s first and second laws of electrolysis;

(d) state the relationship between the Faraday constant, the Avogadro constant and the electronic charge;

(e) calculate the quantity of electricity used, the mass of material and/ or gas volume liberated during electrolysis.

8.6 Applications of electrochemistry

Candidates should be able to:

(a) explain the principles of electrochemistry in the process and prevention of corrosion (rusting of iron);

(b) describe the extraction of aluminium by electrolysis, and state the advantages of recycling aluminium;

(c) describe the process of anodisation of aluminium to resist corrosion;

(d) describe the diaphragm cell in the manufacture of chlorine from brine;

(e) describe the treatment of industrial effluent by electrolysis to remove Ni2+, Cr3+ and Cd2+;

(f) describe the electroplating of coated plastics.

 

9 Periodic Table: Periodicity

9.1 Physical properties of elements of Period 2 and Period 3

Candidates should be able to:

(a) interpret and explain the trend and gradation of atomic radius, melting point, boiling point, enthalpy change of vaporisation and electrical conductivity in terms of structure and bonding;

(b) explain the factors influencing ionisation energies;

(c) explain the trend in ionisation energies across Period 2 and Period 3 and down a group;

(d) predict the electronic configuration and position of unknown elements in the Periodic Table from successive values of ionisation energies.

9.2 Reactions of Period 3 elements with oxygen and water

Candidates should be able to:

(a) describe the reactions of Period 3 elements with oxygen and water;

(b) interpret the ability of elements to act as oxidising and reducing agents.

9.3 Acidic and basic properties of oxides and hydrolysis of oxides

Candidates should be able to:

(a) explain the acidic and basic properties of the oxides of Period 3 elements;

(b) describe the reactions of the oxides of Period 3 elements with water;

(c) describe the classification of the oxides of Period 3 elements as basic, amphoteric or acidic based on their reactions with water, acid and alkali;

(d) describe the use of sulphur dioxide in food preservation.

 

10 Group 2

10.1 Selected Group 2 elements and their compounds

Candidates should be able to:

(a) describe the trends in physical properties of Group 2 elements: Mg, Ca, Sr, Ba;

(b) describe the reactions of Group 2 elements with oxygen and water;

(c) describe the behaviour of the oxides of Group 2 elements with water;

(d) explain qualitatively the thermal decomposition of the nitrates, carbonates and hydroxides of Group 2 elements in terms of the charge density and polarisability of large anions;

(e) explain qualitatively the variation in solubility of sulphate of Group 2 elements in terms of the relative magnitudes of the enthalpy change of hydration for the relevant ions and the corresponding lattice energy.

10.2 Anomalous behaviour of beryllium

Candidates should be able to:

(a) explain the anomalous behaviour of beryllium as exemplified by the formation of covalent compounds;

(b) describe the diagonal relationships between beryllium and aluminium;

(c) explain the similarity of aqueous beryllium salts to aqueous aluminium salts in terms of their acidic property.

10.3 Uses of Group 2 compounds

Candidates should be able to:

(a) state the uses of Group 2 compounds in agriculture, industry and medicine.

 

11 Group 14

11.1 Physical properties of Group 14 elements

Candidates should be able to:

(a) explain the trends in physical properties (melting points and electrical conductivity) of Group 14 elements: C, Si, Ge, Sn, Pb.

11.2 Tetrachlorides and oxides of Group 14 elements

Candidates should be able to:

(a) explain the bonding and molecular shapes of the tetrachlorides of group 14 elements;

(b) explain the volatility, thermal stability and hydrolysis of tetrachlorides in terms of structure and bonding;

(c) explain the bonding, acid-base nature and the thermal stability of the oxides of oxidation states +2 and +4.

11.3 Relative stability of +2 and +4 oxidation states of Group 14 elements

Candidates should be able to:

(a) explain the relative stability of +2 and +4 oxidation states of the elements in their oxides, chlorides and aqueous cations.

11.4 Silicon, silicone and silicates

Candidates should be able to:

(a) describe the structures of silicone and silicates (pyroxenes and amphiboles), sheets (mica) and framework structure (quartz) (general formulae are not required);

(b) explain the uses of silicon as a semiconductor and silicone as a fluid, elastomer and resin;

(c) describe the uses of silicates as basic materials for cement, glass, ceramics and zeolites.

11.5 Tin alloys

Candidates should be able to:

(a) describe the uses of tin in solder and pewter.

 

12 Group 17

12.1 Physical properties of selected Group 17 elements

Candidates should be able to:

(a) state that the colour intensity of Group 17 elements: Cl2, Br2, I2, increase down the group;

(b) explain how the volatility of Group 17 elements decreases down the group.

12.2 Reactions of selected Group 17 elements

Candidates should be able to:

(a) deduce and explain the relative reactivities of Group 17 elements as oxidising agents from Eº values;

(b) explain the order of reactivity of F2, Cl2, Br2, I2 with hydrogen, and compare the relative thermal stabilities of the hydrides;

(c) explain the reactions of chlorine with cold and hot aqueous sodium hydroxide.

12.3 Reactions of selected halide ions

Candidates should be able to:

(a) explain and write equations for reactions of Group 17 ions with aqueous silver ions followed by aqueous ammonia;

(b) explain and write equations for reactions of Group 17 ions with concentrated sulphuric acid.

12.4 Industrial applications of halogens and their compounds

Candidates should be able to:

(a) describe the industrial uses of the halogens and their compounds as antiseptic, bleaching agent and in black-and-white photography;

(b) explain the use of chlorine in water treatment.

 

13 Transition Elements

13.1 Physical properties of first row transition elements

Candidates should be able to:

(a) define a transition element in terms of incomplete d orbitals in at least one of its ions;

(b) describe the similarities in physical properties such as atomic radius, ionic radius and first ionisation energy;

(c) explain the variation in successive ionisation energies;

(d) contrast qualitatively the melting point, density, atomic radius, ionic radius, first ionisation energy and conductivity of the first row transition elements with those of calcium as a typical s-block element.

13.2 Chemical properties of first row transition elements

Candidates should be able to:

(a) explain variable oxidation states in terms of the energies of 3d and 4s orbitals;

(b) explain the colours of transition metal ions in terms of a partially filled 3d orbitals;

(c) state the principal oxidation numbers of these elements in their common cations, oxides and oxo ions;

(d) explain qualitatively the relative stabilities of these oxidation states;

(e) explain the uses of standard reduction potentials in predicting the relative stabilities of aqueous ions;

(f) explain the terms complex ion and ligand;

(g) explain the formation of complex ions and the colour changes by exchange of ligands. (Examples of ligands: water, ammonia, cyanide ions, thiocyanate ions, ethanedioate ions, ethylenediaminetetraethanoate, halide ions; examples of complex ions: [Fe(CN)6]4, [Fe(CN)6]3, [Fe(H2O)5(SCN)]2+);

(h) explain the use of first row transition elements in homogeneous catalysis, as exemplifed by Fe2+ or Fe3+ in the reaction between I and S2O82;

(i) explain the use of first row transition elements in heterogeneous catalysis, as exemplifed by Ni and Pt in the hydrogenation of alkenes.

13.3 Nomenclature and bonding of complexes

Candidates should be able to:

(a) name complexes using International Union of Pure and Applied Chemistry (IUPAC) nomenclature;

(b) discuss coordinate bond formation between ligands and the central metal atom/ion, and state the types of ligands, i.e. monodentate, bidentate and hexadentate.

13.4 Uses of first row transition elements and their compounds

Candidates should be able to:

(a) describe the use of chromium (in stainless steel), cobalt, manganese, titanium (in alloys) and TiO2 (in paints).

Third Term

14 Introductions to Organic Chemistry

14.1 Bonding of the carbon atoms: the shapes of ethane, ethene, ethyne and benzene molecules

Candidates should be able to:

(a) use the concept of sp3, sp2 and sp hybridisations in carbon atoms to describe the bonding and shapes of molecules as exemplified by CH4, C2H4, C2H2 and C6H6;

(b) explain the concept of delocalisation of electrons in benzene ring.

14.2 General, empirical, molecular and structural formulae of organic compounds

Candidates should be able to:

(a) state general, empirical, molecular and structural formulae of organic compounds;

(b) determine empirical and molecular formulae of organic compounds.

14.3 Functional groups: classification and nomenclature

Candidates should be able to:

(a) describe the classification of organic compounds by functional groups and the nomenclature of classes of organic compounds according to the IUPAC rules of the following classes of compounds:

(i) alkanes, alkenes, alkynes and arenes,

(ii) haloalkanes,

(iii) alcohols (including primary, secondary and tertiary) and phenols,

(iv) aldehydes and ketones,

(v) carboxylic acids and their derivatives (acyl chlorides, amides and esters),

(vi) primary amines, amino acids and protein.

14.4 Isomerism: structural and stereoisomerism

Candidates should be able to:

(a) define structural and stereoisomerism (geometrical and optical);

(b) explain the meaning of a chiral centre in optical isomerism;

(c) classify isomers as structural, cis-trans and optical isomers;

(d) identify chiral centres and/or cis-trans isomerism in a molecule of given structural formula;

(e) deduce the possible isomers for an organic compound of known molecular formula.

14.5 Free radicals, nucleophiles and electrophiles

Candidates should be able to:

(a) describe homolytic and heterolytic fissions;

(b) define the terms free radical, nucleophile and electrophile;

(c) explain that nucleophiles such as OH, NH3, H2O, Br, I and carbanion are Lewis bases;

(d) explain that electrophiles such as H+, NO2+, Br2, A1C13, ZnC12, FeBr3, BF3 and carbonium ion are Lewis acids.

14.6 Molecular structure and its effect on physical properties

Candidates should be able to:

(a) describe the relationship between the size of molecules in the homologous series and the melting and boiling points;

(b) explain the forces of attraction between molecules (van der Waals forces and hydrogen bonding).

14.7 Inductive and resonance effect

Candidates should be able to:

(a) explain inductive effect which can determine the properties and reactions of functional groups;

(b) use inductive effect to explain why functional groups such as NO2, CN, COOH, COOR, >C=O, SO3H, X (halogen), OH, OR, NH2, C6H5 are electron acceptors whereas R(alkyl) is an electron donor;

(c) explain how the concept of induction can account for the differences in acidity between CH3COOH, C1CH2COOH, C12CHCOOH and Cl3CCOOH; between C1CH2CH2CH2COOH and CH3CH2CHClCOOH;

(d) use the concept of resonance to explain the differences in acidity between CH3CH2OH and C6H5OH, as well as the differences in basicity between CH3NH2 and C6H5NH2.

 

15 Hydrocarbons

15.1 Can understand and describe the alkanes

Candidates should be able to:

(a) write the general formula for alkanes;

(b) explain the construction of the alkane series (straight and branched), and IUPAC nomenclature of alkanes for C1 to C10;

(c) describe the structural isomerism in aliphatic alkanes and cis-trans isomerism in cycloalkanes;

(d) state the physical properties of alkanes;

(e) define alkanes as saturated aliphatic hydrocarbons;

(f) name alkyl groups derived from alkanes and identify primary, secondary, tertiary and quartenary carbons;

(g) explain the inertness of alkanes towards polar reagents;

(h) describe the mechanism of free radical substitution as exemplified by the chlorination of methane (with particular reference to the initiation, propagation and termination reactions);

(i) describe the oxidation of alkane with limited and excess oxygen, and the use of alkanes as fuels;

(j) explain the use of crude oil as a source of aliphatic hydrocarbons;

(k) explain how cracking reactions can be used to obtain alkanes and alkenes of lower Mr from larger hydrocarbon molecules;

(l) discuss the role of catalytic converters in minimising air pollution by oxidising CO to CO2 and reducing NOx to N2;

(m) explain how chemical pollutants from the combustion of hydrocarbon affect air quality and rainwater as exemplified by acid rain, photochemical smog and greenhouse effect.

15.2 Can understand and describe the alkenes

Candidates should be able to:

(a) write the general formula for alkenes;

(b) name alkenes according to the IUPAC nomenclature and their common names for C1 to C5;

(c) describe structural and cis-trans isomerism in alkenes;

(d) state the physical properties of alkenes;

(e) define alkenes as unsaturated aliphatic hydrocarbons with one or more double bonds;

(f) describe the chemical reactions of alkenes as exemplified by the following reactions of ethene:

(i) addition of hydrogen, steam, hydrogen halides, halogens, bromine water and concentrated sulphuric acid,

(ii) oxidation using KMnO4, O2/Ag,

(iii) ozonolysis,

(iv) polymerisation;

(g) describe the mechanism of electrophilic addition in alkenes with reference to Markovnikov‟s rule;

(h) explain the use of bromination reaction and decolourisation of MnO4 ions as simple tests for alkenes and unsaturated compounds;

(i) explain briefly the importance of ethene as a source for the preparation of chloroethane, epoxyethane, ethane-1,2-diol and poly(ethane).

15.3 Can understand and describe the arenes

Candidates should be able to:

(a) name aromatic compounds derived from benzene according to the IUPAC nomenclature, including the use of ortho, meta and para or the numbering of substituted groups to the benzene ring;

(b) describe structural isomerism in arenes;

(c) describe the chemical reactions of arenes as exemplified by substitution reactions of haloalkanes and acyl chloride (Friedel-Crafts reaction), halogen, conc. HNO3/conc. H2SO4 and SO3 with benzene and methylbenzene (toluene);

(d) describe the mechanism of electrophilic substitution in arenes as exemplified by the nitration of benzene;

(e) explain why benzene is more stable than aliphatic alkenes towards oxidation;

(f) describe the reaction between alkylbenzene and hot acidified KMnO4;

(g) determine the products of halogenation of methylbenzene (toluene) in the presence of

(i) Lewis acid catalysts,

(ii) light;

(h) explain the inductive effect and resonance effect of substituted groups (OH, C1, CH3, NO2, COCH3, NH2) attached to the benzene ring towards further substitutions;

(i) predict the products in an electrophilic substitution reaction when the substituted group in benzene is electron accepting or electron donating;

(j) explain the uses of arenes as solvents;

(k) recognise arenes as carcinogen.

 

16 Haloalkanes

16.1 Can understand and describe the haloalkanes

Candidates should be able to:

(a) write the general formula for haloalkanes;

(b) name haloalkanes according to the IUPAC nomenclature;

(c) describe the structural and optical isomerism in haloalkanes;

(d) state the physical properties of haloalkanes;

(e) describe the substitution reactions of haloalkanes as exemplified by the following reactions of bromoethane: hydrolysis, the formation of nitriles and the formation of primary amines;

(f) describe the elimination reactions of haloalkanes;

(g) describe the mechanism of nucleophilic substitution in haloalkanes (SN1 and SN2);

(h) explain the relative reactivity of primary, secondary and tertiary haloalkanes;

(i) compare the reactivity of chlorobenzene and chloroalkanes in hydrolysis reactions;

(j) explain the use of haloalkanes in the synthesis of organomagnesium compounds (Grignard reagents), and their use in reactions with carbonyl compounds;

(k) describe the uses of fluoroalkanes and chlorofluoroalkanes as inert substances for aerosol propellants, coolants and fire-extinguishers;

(l) state the use of chloroalkanes as insecticide such as DDT;

(m) describe the effect of chlorofluoroalkanes in the depletion of the ozone layer, and explain its mechanism.

17 Hydroxy Compounds

17.1 Introduction to hydroxy compounds

Candidates should be able to:

(a) write the general formula for hydroxy compounds;

(b) name hydroxy compounds according to the IUPAC nomenclature;

(c) describe structural and optical isomerism in hydroxy compounds;

(d) state the physical properties of hydroxy compounds.

17.2 Alcohols

Candidates should be able to:

(a) classify alcohols into primary, secondary and tertiary alcohol;

(b) classify the reactions of alcohols whereby the ROH bond is broken: the formation of an alkoxide with sodium, esterification, acylation, oxidation to carbonyl compounds and carboxylic acids;

(c) classify the reactions of alcohols whereby the ROH is broken and OH is replaced by other groups: the formation of haloalkanes and the dehydration to alkenes and ethers;

(d) explain the relative reactivity of primary, secondary and tertiary alcohols as exemplified by the reaction rate of such alcohols to give haloalkanes, and the reaction products of KMnO4/K2Cr2O7 oxidation in the presence of sulphuric acid;

(e) explain the reaction of alcohol with the structure CH3CH(OH) with alkaline aqueous solution of iodine to form triiodomethane;

(f) describe the laboratory and industrial preparation of alcohol as exemplified by ethanol from the hydration of ethane;

(g) describe the synthesis of ethanol by fermentation process;

(h) state the uses of alcohols as antiseptic, solvent and fuel.

17.3 Phenols

Candidates should be able to:

(a) explain the relative acidity of water, phenol and ethanol with particular reference to the inductive and resonance effects;

(b) describe the reactions of phenol with sodium hydroxide, sodium, acyl chlorides and electrophilic substitution in the benzene ring;

(c) describe the use of bromine water and aqueous iron(III) chloride as tests for phenol;

(d) describe the cumene process in the manufacture of phenol;

(e) explain the use of phenol in the manufacture of cyclohexanol, and hence, nylon-6,6.

 

18 Carbonyl Compounds

Candidates should be able to:

(a) write the general formula for carbonyl compounds: aliphatic and aromatic aldehydes and ketones;

(b) name aliphatic and aromatic aldehydes and ketones according to the IUPAC nomenclature;

(c) describe structural and optical isomerism in carbonyl compounds;

(d) state the physical properties of aliphatic and aromatic aldehydes and ketones;

(e) write the equations for the preparation of aldehydes and ketones;

(f) explain the reduction reactions of aldehydes and ketones to primary and secondary alcohols respectively through catalytic hydrogenation reaction and with LiA1H4;

(g) explain the use of 2,4-dinitrophenylhydrazine reagent as a simple test to detect the presence of >C=O groups;

(h) explain the mechanism of the nucleophilic addition reactions of hydrogen cyanide with aldehydes and ketones;

(i) explain the oxidation of aldehydes;

(j) differentiate between aldehyde and ketone based on the results of simple tests as exemplified by Fehling‟s solution and Tollens‟ reagent;

(k) explain the reactions of carbonyl compounds with the structure CH3C=O with alkaline aqueous solution of iodine to give triiodomethane (iodoform test);

(l) explain that natural compounds such as glucose, sucrose and other carbohydrates which have the >C=O group;

(m) explain the characteristics of glucose as a reducing sugar.

 

19 Carboxylic Acids and their Derivatives

19.1 Carboxylic acid

Candidates should be able to:

(a) write the general formula for aliphatic and aromatic carboxylic acids;

(b) name carboxylic acids according to the IUPAC nomenclature and their common names for C1 to C6;

(c) describe structural and optical isomerism in carboxylic acids;

(d) state the physical properties of carboxylic acids;

(e) write the equations for the formation of carboxylic acids from alcohols, aldehydes and nitriles;

(f) describe the acidic properties of carboxylic acids as exemplified by their reactions with metals and bases to form salts;

(g) explain the substitution of the OH in carboxylic acids by the nucleophiles OR and C1 to form esters and acyl chlorides respectively;

(h) describe the reduction of carboxylic acids to primary alcohols;

(i) describe the oxidation and dehydration of methanoic and ethanedioic acids (oxalic acid);

(j) state the uses of carboxylic acids in food, perfume, health (aspirin) and polymer industries.

19.2 Acyl chlorides

Candidates should be able to:

(a) write the general formula for acyl chlorides;

(b) name acyl chlorides according to the IUPAC nomenclature;

(c) describe structural and optical isomerism in acyl chlorides;

(d) state the physical properties of acyl chlorides;

(e) explain the ease of hydrolysis of acyl chlorides compared to chloroalkanes;

(f) describe the reactions of acyl chlorides with alcohols, phenols and primary amines.

19.3 Esters

Candidates should be able to:

(a) write the general formula for esters;

(b) name esters according to the IUPAC nomenclature;

(c) describe structural and optical isomerism in esters;

(d) state the physical properties of esters;

(e) describe the preparation of esters by the reactions of acyl chlorides with alcohols and phenols;

(f) describe the acid and base hydrolysis of esters;

(g) describe the reduction of esters to primary alcohols;

(h) state the uses of esters as flavourings, preservatives and solvents.

19.4 Amides

Candidates should be able to:

(a) write the general formula for amides;

(b) name amides according to the IUPAC nomenclature;

(c) describe structural and optical isomerism in amides;

(d) state the physical properties of amides;

(e) describe the preparation of amides by the reaction of acyl chlorides with primary amines;

(f) describe the acid and base hydrolysis of amides.

 

20 Amines, Amino Acids and Proteins

20.1 Amines

Candidates should be able to:

(a) write the general formula for amines;

(b) name amines according to the IUPAC nomenclature and their common names;

(c) describe structural and optical isomerism in amines;

(d) state the physical properties of amines;

(e) classify amines into primary, secondary and tertiary amines;

(f) explain the relative basicity of ammonia, ethanamine and phenylamine (aniline) in terms of their structures;

(g) describe the preparation of ethanamine by the reduction of nitriles, and phenylamine by the reduction of nitrobenzene;

(h) explain the formation of salts when amines react with mineral acids;

(i) differentiate primary aliphatic amines from primary aryl (aromatic) amines by their respective reactions with nitric(III) acid (nitrous acid) and bromine water;

(j) explain the formation of dyes by the coupling reaction of the diazonium salt as exemplified by the reaction of benzenediazonium chloride with phenol.

20.2 Amino acids

Candidates should be able to:

(a) write the structure and general formula for -amino acids;

(b) name -amino acids according to the IUPAC nomenclature and their common names;

(c) describe structural and optical isomerism in amino acids;

(d) state the physical properties of -amino acids;

(e) describe the acid and base properties of -amino acids;

(f) describe the formation of zwitterions;

(g) explain the peptide linkage as amide linkage formed by the condensation between two or more -amino acids as exemplified by glycylalanine and alanilglycine.

20.3 Protein

Candidates should be able to:

(a) identify the peptide linkage in the primary structure of protein;

(b) describe the hydrolysis of proteins;

(c) state the biological importance of proteins.

 

21 Polymers

Candidates should be able to:

(a) state examples of natural and synthetic polymers;

(b) define monomer, polymer, repeating unit, homopolymer and copolymer;

(c) identify the monomers in a polymer;

(d) describe condensation polymerisation as exemplified by terylene and nylon-6,6;

(e) describe addition polymerisation as exemplified by poly(ethene)/ polyethylene/ polythene, poly(phenylethene)/ polystyrene and poly(chloroethene)/ polyvinylchloride;

 

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