CHE(MYESTRY) SYLLABUS
1. Periodic Properties and variations of
Properties – Physical and Chemical.
(i) Periodic properties and their variations in
groups and periods.
Definitions and trends of the following
periodic properties in groups and periods
should be studied:
• atomic size
• metallic character
• non-metallic character
• ionisation potential
• electron affinity
• electronegativity
(ii) Periodicity on the basis of atomic number for
elements.
• The study of modern periodic table up to
period 3 (students to be exposed to the
complete modern periodic table but no
questions will be asked on elements
beyond period 3 – Argon);
• Periodicity and other related properties
to be explained on the basis of nuclear
charge and shells (not orbitals).
(Special reference to the alkali metals and halogen
groups).
2. Chemical Bonding
Electrovalent, covalent and co-ordinate
bonding, structures of various compounds,
Electron dot structure.
(a) Electrovalent bonding:
• Electron dot structure of
Electrovalent compounds NaCl,
MgCl2, CaO.
• Characteristic properties of
electrovalent compounds – state of
existence, melting and boiling
points, conductivity (heat and
electricity), dissociation in solution
and in molten state to be linked with
electrolysis.
(b) Covalent Bonding:
• Electron dot structure of covalent
molecules on the basis of duplet
and octet of electrons (example:
hydrogen, chlorine, nitrogen,
ammonia, carbon tetrachloride,
methane.
• Polar Covalent compounds –
based on difference in
electronegativity:
Examples – HCl and H2O
including structures.
• Characteristic properties of
Covalent compounds – state of
existence, melting and boiling
points, conductivity (heat and
electricity), ionisation in solution.
Comparison of Electrovalent and
Covalent compounds.
(c) Coordinate Bonding:
• Definition
• The lone pair effect of the oxygen
atom of the water molecule and the
nitrogen atom of the ammonia
molecule to explain the formation of
H3O+ and OH- ions in water and
NH4
+ ion.
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The meaning of lone pair; the
formation of hydronium ion and
ammonium ion must be explained with
help of electron dot diagrams.
3. Study of Acids, Bases and Salts
(i) Simple definitions in terms of the molecules
and their characteristic properties.
(ii) Ions present in mineral acids, alkalis and
salts and their solutions; use of litmus and pH
paper to test for acidity and alkalinity.
• Examples with equation for the
ionisation/dissociation of ions of acids,
bases and salts.
• Acids form hydronium ions (only positive
ions) which turn blue litmus red, alkalis
form hydroxyl ions (only negative ions)
with water which turns red litmus blue.
• Salts are formed by partial or complete
replacement of the hydrogen ion of an
acid by a metal. (To be explained with
suitable examples).
• Introduction to pH scale to test for
acidity, neutrality and alkalinity by using
pH paper or Universal indicator.
(iii)Definition of salt; types of salts.
Types of salts: normal salts, acid salt, basic
salt, definition and examples.
(iv) Action of dilute acids on salts.
Decomposition of hydrogen carbonates,
carbonates, sulphites and sulphides by
appropriate acids with heating if necessary.
(Relevant laboratory work must be done).
(v) Methods of preparation of Normal salts with
relevant equations. (Details of apparatus or
procedures not required).
Methods included are:
• Direct combination
• Displacement
• Precipitation (double decomposition)
• Neutralization of insoluble base
• Neutralisation of an alkali (titration)
• Action of dilute acids on carbonates
and bi-carbonates.
4. Analytical Chemistry
(i) Action of Ammonium Hydroxide and
Sodium Hydroxide on solution of salts:
colour of salt and its solution; formation
and colour of hydroxide precipitated for
solutions of salts of Ca, Fe, Cu, Zn and
Pb; special action of ammonium
hydroxide on solutions of copper salt and
sodium hydroxide on ammonium salts.
On solution of salts:
• Colour of salt and its solution.
• Action on addition of Sodium
Hydroxide to solution of Ca, Fe, Cu,
Zn, and Pb salts drop by drop in
excess. Formation and colour of
hydroxide precipitated to be
highlighted with the help of
equations.
• Action on addition of Ammonium
Hydroxide to solution of Ca, Fe, Cu,
Zn, and Pb salts drop by drop in
excess. Formation and colour of
hydroxide precipitated to be
highlighted with the help of
equations.
• Special action of Ammonium
Hydroxide on solutions of copper
salts and sodium hydroxide on
ammonium salts.
(ii) Action of alkalis (NaOH, KOH) on
certain metals, their oxides and
hydroxides.
The metals must include aluminium, zinc
and lead, their oxides and hydroxides,
which react with caustic alkalis (NaOH,
KOH), showing the amphoteric nature of
these substances.
5. Mole Concept and Stoichiometry
(i) Gay Lussac’s Law of Combining Volumes;
Avogadro’s Law.
• Idea of mole – a number just as a dozen,
a gross (Avogadro’s number).
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• Avogadro’s Law - statement and
explanation.
• Gay Lussac’s Law of Combining
Volumes. – Statement and explanation.
• Understanding molar volume- “the mass
of 22.4 litres of any gas at S.T.P. is equal
to its molar mass”. (Questions will not be
set on formal proof but may be taught for
clear understandin).
• Simple calculations based on the molar
volume and Gay Lussac’s law.
(ii) Refer to the atomicity of hydrogen, oxygen,
nitrogen and chlorine (proof not required).
The explanation can be given using
equations for the formation of HCl, NH3, and
NO.
(iii) Vapour Density and its relation to relative
molecular mass:
• Molecular mass = 2×vapour density
(formal proof not required)
• Deduction of simple (empirical) and
molecular formula from:
(a) the percentage composition of a
compound.
(b) the masses of combining elements.
(iv) Mole and its relation to mass.
• Relating mole and atomic mass;
arriving at gram atomic mass and then
gram atom; atomic mass is a number
dealing with one atom; gram atomic
mass is the mass of one mole of atoms.
• Relating mole and molecular mass
arriving at gram molecular mass and
gram molecule – molecular mass is a
number dealing with a molecule, gram
molecular mass is the mass of one mole
of molecules.
• Simple calculations based on relation
of mole to mass, volume and
Avogadro’s number.
(v) Simple calculations based on chemical
equations
Related to weight and/or volumes of both
reactants and products.
6. Electrolysis
(i) Electrolytes and non-electrolytes.
Definitions and examples.
(ii) Substances containing molecules only, ions
only, both molecules and ions.
• Substances containing molecules only
ions only, both molecules and ions.
• Examples; relating their composition
with their behaviour as strong and weak
electrolytes as well as non-electrolytes.
(iii) Definition and explanation of electrolysis,
electrolyte, electrode, anode, cathode, anion,
cation, oxidation and reduction (on the basis
of loss and gain of electrons).
(iv)An elementary study of the migration of
ions, with reference to the factors influencing
selective discharge of ions (reference should
be made to the activity series as indicating
the tendency of metals, e.g. Na, Mg, Fe, Cu,
to form ions) illustrated by the electrolysis
of:
• Molten lead bromide
• acidified water with platinum electrodes
• Aqueous copper (II) sulphate with
copper electrodes; electron transfer at the
electrodes.
The above electrolytic processes can be
studied in terms of electrolyte used,
electrodes used, ionization reaction, anode
reaction, cathode reaction, use of selective
discharge theory, wherever applicable.
(v) Applications of electrolysis:
• Electroplating with nickel and silver,
choice of electrolyte for electroplating.
• Electro refining of copper;
Reasons and conditions for electroplating;
names of the electrolytes and the electrodes
used should be given. Equations for the
reactions at the electrodes should be given
for electroplating, refining of copper.
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7. Metallurgy
(i) Occurrence of metals in nature:
• Mineral and ore - Meaning only.
• Common ores of iron, aluminium and
zinc.
(ii) Stages involved in the extraction of metals:
(a) Dressing of the ore – hydrolytic
method, magnetic separation, froth
flotation method.
(b) Conversion of concentrated ore to its
oxide- roasting and calcination
(definition, examples with equations).
(c) Reduction of metallic oxides- some can
be reduced by hydrogen, carbon and
carbon monoxide (e.g. copper oxide,
lead (II) oxide, iron (III) oxide and zinc
oxide) and some cannot (e.g. Al2O3,
MgO) - refer to activity series). Active
metals by electrolysis e.g. sodium,
potassium and calcium. (reference only).
Equations with conditions should be
given.
(d) Electro refining – reference only
(iii) Extraction of Aluminium.
(a) Chemical method for purifying bauxite
by using NaOH – Baeyer’s Process.
(b) Electrolytic extraction – Hall Heroult’s
process:
Structure of electrolytic cell - the
various components as part of the
electrolyte, electrodes and electrode
reactions.
Description of the changes occurring,
purpose of the substances used and the
main reactions with their equations.
(iv) Alloys – composition and uses
Stainlesssteel, duralumin, brass, bronze,
fuse metal / solder.
8. Study of Compounds
A. Hydrogen Chloride
Hydrogen chloride: preparation of hydrogen
chloride from sodium chloride; refer to the
density and solubility of hydrogen chloride
(fountain experiment); reaction with
ammonia; acidic properties of its solution.
• Preparation of hydrogen chloride from
sodium chloride; the laboratory method
of preparation can be learnt in terms of
reactants, product, condition, equation,
diagram or setting of the apparatus,
procedure, observation, precaution,
collection of the gas and identification.
• Simple experiment to show the density of
the gas (Hydrogen Chloride) –heavier
than air.
• Solubility of hydrogen chloride (fountain
experiment); setting of the apparatus,
procedure, observation, inference.
• Method of preparation of hydrochloric
acid by dissolving the gas in water- the
special arrangement and the mechanism
by which the back suction is avoided
should be learnt.
• Reaction with ammonia
• Acidic properties of its solution -
reaction with metals, their oxides,
hydroxides and carbonates to give their
chlorides; decomposition of carbonates,
hydrogen carbonates, sulphides,
sulphites.
• Precipitation reactions with silver
nitrate solution and lead nitrate solution.
B. Ammonia
Ammonia: its laboratory preparation
from ammonium chloride and collection;
ammonia from nitrides like Mg3N2
and AlN and ammonium salts.
Manufacture by Haber’s Process;
density and solubility of ammonia
(fountain experiment); aqueous solution
of ammonia; its reactions with hydrogen
chloride and with hot copper (II) oxide
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and chlorine; the burning of ammonia in
oxygen; uses of ammonia.
• Laboratory preparation from
ammonium chloride and collection;
(the preparation to be studied in
terms of, setting of the apparatus and
diagram, procedure, observation,
collection and identification)
• Ammonia from nitrides like Mg3N2
and AlN using warm water.
Ammonia from ammonium salts using
alkalies.
The reactions to be studied in terms of
reactants, products, conditions and
equations.
• Manufacture by Haber’s Process.
• Density and solubility of ammonia
(fountain experiment).
• The burning of ammonia in oxygen.
• The catalytic oxidation of ammonia
(with conditions and reaction)
• Its reactions with hydrogen chloride
and with hot copper (II) oxide and
chlorine (both chlorine in excess and
ammonia in excess).
All these reactions may be studied in
terms of reactants, products,
conditions, equations and
observations.
• Aqueous solution of ammonia - reaction
with sulphuric acid, nitric acid,
hydrochloric acid and solutions of
iron(III) chloride, iron(II) sulphate,
lead nitrate, zinc nitrate and copper
sulphate.
• Uses of ammonia - manufacture of
fertilizers, explosives, nitric acid,
refrigerant gas (Chlorofluro carbon –
and its suitable alternatives which are
non-ozone depleting), and cleansing
agents.
C. Nitric Acid
Nitric Acid: one laboratory method of
preparation of nitric acid from potassium
nitrate or sodium nitrate. Large scale
preparation. Nitric acid as an oxidizing
agent.
• Laboratory preparation of nitric acid
from potassium nitrate or sodium
nitrate; the laboratory method to be
studied in terms of reactants, products,
conditions, equations, setting up of
apparatus, diagram, precautions,
collection and identification.
• Manufacture of Nitric acid by
Ostwald’s process (Only equations
with conditions where applicable).
• As an oxidising agent: its reaction with
copper, carbon, sulphur.
D. Sulphuric Acid
Large scale preparation, its behaviour as an
acid when dilute, as an oxidizing agent when
concentrated - oxidation of carbon and
sulphur; as a dehydrating agent - dehydration
of sugar and copper (II) sulphate crystals; its
non-volatile nature.
• Manufacture by Contact Process
Equations with conditions where
applicable).
• Its behaviour as an acid when dilute -
reaction with metal, metal oxide, metal
hydroxide, metal carbonate, metal
bicarbonate, metal sulphite, metal
sulphide.
• Concentrated sulphuric acid as an
oxidizing agent - the oxidation of carbon
and sulphur.
• Concentrated sulphuric acid as a
dehydrating agent- (a) the dehydration of
sugar (b) Copper (II) sulphate crystals.
• Non-volatile nature of sulphuric acid -
reaction with sodium or potassium
chloride and sodium or potassium nitrate.
9. Organic Chemistry
(i) Introduction to Organic compounds.
• Unique nature of Carbon atom – tetra
valency, catenation.
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• Formation of single, double and triple
bonds, straight chain, branched chain,
cyclic compounds (only benzene).
(ii) Structure and Isomerism.
• Structure of compounds with single,
double and triple bonds.
• Structural formulae of hydrocarbons.
Structural formula must be given for:
alkanes, alkenes, alkynes up to 5 carbon
atoms.
• Isomerism – structural (chain, position)
(iii)Homologous series – characteristics with
examples.
Alkane, alkene, alkyne series and their
gradation in properties and the relationship
with the molecular mass or molecular
formula.
(iv) Simple nomenclature.
Simple nomenclature - of the hydrocarbons
with simple functional groups – (double
bond, triple bond, alcoholic, aldehydic,
carboxylic group) longest chain rule and
smallest number for functional groups rule –
trivial and IUPAC names (compounds with
only one functional group)
(v) Hydrocarbons: alkanes, alkenes, alkynes.
• Alkanes - general formula; methane
(greenhouse gas) and ethane - methods of
preparation from sodium ethanoate
(sodium acetate), sodium propanoate
(sodium propionate), from iodomethane
(methyl iodide) and bromoethane (ethyl
bromide). Complete combustion of
methane and ethane, reaction of methane
and ethane with chlorine through
substitution.
• Alkenes – (unsaturated hydrocarbons
with a double bond); ethene as an
example. Methods of preparation of
ethene by dehydro halogenation reaction
and dehydration reactions.
• Alkynes -(unsaturated hydrocarbons
with a triple bond); ethyne as an example
of alkyne; Methods of preparation from
calcium carbide and 1,2 dibromoethane
ethylene dibromide).
Only main properties, particularly addition
products with hydrogen and halogen
namely Cl2, Br2 and I2 pertaining to alkenes
and alkynes.
• Uses of methane, ethane, ethene, ethyne.
(vi) Alcohols: ethanol – preparation, properties
and uses.
• Preparation of ethanol by hydrolysis of
alkyl halide.
• Properties – Physical: Nature, Solubility,
Density, Boiling Points. Chemical:
Combustion, action with sodium, ester
formation with acetic acid, dehydration
with conc. Sulphuric acid to prepare
ethene.
• Denatured and spurious alcohol.
• Important uses of Ethanol.
(vii) Carboxylic acids (aliphatic - mono
carboxylic acid): Acetic acid – properties
and uses of acetic acid.
• Structure of acetic acid.
• Properties of Acetic Acid: Physical
properties – odour (vinegar), glacial
acetic acid (effect of sufficient cooling to
produce ice like crystals). Chemical
properties – action with litmus, alkalis and
alcohol (idea of esterification).
• Uses of acetic acid |
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