ICSE Class 10 Syllabus 2028 Chemistry

ICSE
INDIAN CERTIFICATE OF
SECONDARY EDUCATION
EXAMINATION
YEAR 2028

SCIENCE
CHEMISTRY
(52)

Developed by:
Research, Development and Curriculum Division (RDCD)
CISCE

January 2026
____________________________________________________________________________________________

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CLASS X
There will be one paper of two hours duration of 80 marks, and Internal Assessment of practical work carrying
20 marks.
Note: All chemical processes/reactions should be studied with reference to the reactants, products, conditions,
observations and the (balanced) equations and diagrams.

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.
• Detailed study of the modern periodic table restricted to Period 4 (up to Calcium). Elements beyond
this are to be introduced only for familiarisation of the placement of elements in the Periodic table
and are excluded from assessment.
• Periodicity and other related properties of the alkali metals, alkaline earth metals, halogen groups
and zero group elements.
(Explanation on the basis of nuclear charge and shells, not orbitals.)

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, MgCl 2 , CaO, K 2 O, Na 2 O, Na 3 N.
• 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.
• Non-polar covalent compounds (example: hydrogen, chlorine, oxygen, nitrogen, carbon
tetrachloride, methane.)
• Polar Covalent compounds – based on difference in electronegativity.
Examples – HCl, NH 3 and H 2 O 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.

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(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 H 3 O+ and OH- ions in water and NH 4 + ion.
The meaning of lone pair; the formation of hydronium ion and ammonium ion must be explained with
the help of electron dot diagrams.

3. Study of Acids, Bases and Salts
(i) Simple definitions of acids and bases and their characteristic properties.
(ii) Ions present in mineral acids, alkalis and salts and their solutions; use of common acid-base
indicator (litmus, methyl orange and phenolphthalein) 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 like formation of NaHSO 4 and Na 2 SO 4. )

• Introduction to pH scale to test for acidity, neutrality and alkalinity by using pH paper or Universal
indicator.

(iii) Definition of salt; Neutralisation and types of salts.
Types of salts: normal salts, acid salt, basic salt, complex salts (definition and examples like
Acid salt- NaHSO4, Basic Salt – Cu(OH)NO 3 , Complex salt – [Cu(NH 3 ) 4 ]SO 4 .
(iv) Classification of an acid on the basis of concentration, ionisation capacity or strength and number of
hydrogen ions.
(v) Classification of a base on the basis of solubility in water and ionisation capacity or strength and number
of hydroxyl ions.
(vi) Action of dilute acids on salts.
Decomposition of hydrogen carbonates, carbonates, hydrogen sulphites, sulphites and sulphides by
appropriate acids with heating if necessary.
(vii) Methods of preparation of Normal salts with relevant equations.
Methods included are:
• Direct combination
• Displacement
• Precipitation (double decomposition)
• Neutralisation of insoluble base
• Neutralisation of an alkali (titration)
• Action of dilute acids on carbonates and bi-carbonates

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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 Fe, Cu, Zn, Ca, Mg 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 Fe, Cu, Zn, Ca, Mg 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 Fe, Cu, Zn, Ca, Mg 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).
• Avogadro’s Law - statement and explanation.
• Gay Lussac’s Law of Combining Volumes. – Statement and explanation.
• Molar volume- “the mass of 22.4 litres of any gas at S.T.P. is equal to its molar mass.” (proof not
required).
• Simple calculations based on the molar volume and Gay Lussac’s law.
(ii) Atomicity of Hydrogen, Oxygen, Nitrogen and Chlorine with equations for the formation of HCl, NH 3
and NO (proof not required).
(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 volumes of both reactants and products.

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6. Electrolysis
(i) Electrolytes and non-electrolytes.
Definitions and examples.
(ii) Substances containing molecules only, ions only, and 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) Difference between metallic conduction (electrolysis) and electrolytic conduction.
(v) An elementary study of the migration of ions, with reference to the factors with examples 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 and platinum 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.
(vi) Applications of electrolysis.
• Electroplating with nickel and silver, choice of electrolyte for electroplating (nickel sulphate solution
and sodium argentocyanide) with equations.
• Electrorefining of copper with equations.
7. Metallurgy
(i) Occurrence of metals in nature.
• Mineral and ore – Meaning; definition of mineral, ore, gangue, matrix, Flux and slag.
• Common ores of iron, aluminium and zinc.
(ii) Stages involved in the extraction of metals.
(a) Pulverization, 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 active metal oxides can be
reduced by electrolysis only (e.g. Al 2 O 3 , MgO).
(d) Electrorefining – As a process of obtaining pure metal from an impure metal.
(iii) Extraction of Aluminium.
(a) Chemical method for purifying bauxite by using NaOH – Baeyer’s Process only.
(b) Electrolytic extraction – Hall Heroult’s process only:
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.

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(iv) Alloy and Amalgam – composition and uses.
Composition and uses of Stainless steel, magnalium, 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,
decomposition of carbonates, bicarbonates, sulphites, bisulphites, sulphides.
• Aqua regia composition and uses.
• Precipitation reactions with silver nitrate solution and lead nitrate solution.
• Tests for HCl gas, Hydrochloric acid and Chloride ions (Only AgNO 3 test for Identification of
Chloride ion.)
B. Ammonia
Ammonia: its laboratory preparation from ammonium chloride and collection; ammonia from nitrides like
Mg 3 N 2 , 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, lead (IV) oxide 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 Mg 3 N 2 , and AlN using warm water.
• Ammonia from ammonium salts using alkalies.
• Manufacture by Industrial Process- Haber’s Process (along with explanations and schematic
diagram.)
• 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, lead (IV) oxide and chlorine (both
chlorine in excess and ammonia in excess).
Conditions and observations must be emphasized.
• 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 fertilisers, explosives, nitric acid, refrigerant gas (Chlorofluro
carbon – and its suitable alternatives which are non-ozone depleting), and cleansing agents.
• Tests for ammonia.
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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 (Equations with conditions, where applicable;
explanations and schematic diagram).
• As an oxidising agent: its reaction with copper, carbon, sulphur.
• Acidic properties of its solution – reaction with magnesium and manganese, oxides, hydroxides,
carbonates, hydrogen carbonate, sulphites and hydrogen sulphites.
• Test for nitric acid and nitrate ion.
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 (with reason and
schematic diagrams only).
• Its behaviour as an acid when dilute -reaction with metal, metal oxide, hydroxide, carbonate, hydrogen
carbonate, sulphite, hydrogen sulphite and sulphide.
• Sulphuric acid as dibasic acid.
• 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.
• Tests for sulphuric acid and sulphate radical.

9. Organic Chemistry
(i) Introduction to Organic compounds.
• Unique nature of Carbon atom – tetra valency, catenation.
• Formation of single, double and triple bonds, straight chain, branched chain, cyclic compounds (only
benzene structure).
(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 – (halides, double bond, triple
bond, alcoholic, aldehydic, carboxylic ketone, ether and ester) longest chain rule and smallest number
for functional groups rule – trivial and IUPAC names (compounds with only one functional group).

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(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 and incomplete 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).
Only main properties, particularly addition products with hydrogen and halogen namely Cl 2 , Br 2 and
I 2 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, methylated alcohol or spirit 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.

INTERNAL ASSESSMENT OF PRACTICAL WORK
Candidates will be asked to observe the effect of reagents and/or of heat on substances supplied to them. The
exercises will be simple and may include the recognition and identification of certain gases and ions listed below.
The examiners will not, however, be restricted in their choice to substances containing the listed ions.
Gases: Hydrogen, Oxygen, Carbon dioxide, Chlorine, Hydrogen chloride, Sulphur dioxide, Hydrogen sulphide,
Ammonia, Water vapour, Nitrogen dioxide.
Ions: Copper, Iron, Lead, Zinc, Magnesium and Ammonium, Carbonate, Chloride, Nitrate, Calcium, Sulphide,
Sulphite and Sulphate.
Knowledge of a formal scheme of analysis is not required. Semi-micro techniques are acceptable but candidates
using such techniques may need to adapt the instructions given to suit the size of the apparatus being used.
Candidates are expected to have completed the following minimum practical work:
1. Action of heat on the following substances:
(a) Copper carbonate, zinc carbonate
(b) zinc nitrate, copper nitrate, lead nitrate
Make observations, identify the products and make deductions where possible (equations not required).

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2. Action of dilute sulphuric acid on the following substances.
(a) a metal
(b) a carbonate
(c) a sulphide
(d) a sulphite
Make observations, identify the gas evolved and make deductions.
3. Make a solution of the unknown substance: add sodium hydroxide solution or ammonium hydroxide solution,
make observations and give your deduction. Warming the mixture may be needed. Choose from substances
containing Cu2+, Fe2+, Fe3+, Pb2+, Zn2+, Mg2+, Ca2+, NH 4 +.
4. Supply a solution of a dilute acid and alkali. Determine which is acidic and which is basic, giving two tests
for each.
5. Add concentrated hydrochloric acid to each of the given substances, warm, make observations, identify any
product and make deductions: (a) copper oxide (b) manganese dioxide.
6. Revisit: Apply the flame test to identify the metal in the unknown substance.
(a) a sodium salt
(b) a potassium salt
(c) a calcium salt
(d) a copper salt

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EVALUATION
The assignments/project work are to be evaluated by the subject teacher and by an External Examiner.
(The External Examiner may be a teacher nominated by the Head of the school, who could be from the faculty,
but not teaching the subject in the section/class. For example, a teacher of Chemistry of Class VIII may be
deputed to be an External Examiner for Class X Chemistry projects.)
The Internal Examiner and the External Examiner will assess the assignments independently.

Award of Marks (20 Marks)
Subject Teacher (Internal Examiner) 10 marks
External Examiner 10 marks
The total marks obtained out of 20 are to be sent to CISCE by the Head of the school.
The Head of the school will be responsible for the online entry of marks on CISCE’s CAREERS portal by the due
date.
NOTE: According to the recommendation of International Union of Pure and Applied Chemistry (IUPAC),
the groups are numbered from 1 to 18 replacing the older notation of groups IA …. VIIA, VIII, IB …… VIIB
and 0. However, for the examination both notations will be accepted.
Old IA IIA IIIB IVB VB VIB VIIB VIII IB IIB IIIA IVA VA VIA VIIA 0
notation
New 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18
notation

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