ICSE Class 10 Chemistry Syllabus 2023-24 PDF Download

• atomic size
• metallic character
• non-metallic character
• ionisation potential
• electron affinity
• electronegativity

• 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).

(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.

• 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.

• Definition
• The lone pair effect of the oxygen atom of the water molecule and the nitrogen atom of the ammonia molecule explains the formation of H3O + and OH- ions in water and NH4+ ion.

The meaning of lone pair; the formation of hydronium ion and ammonium ion must be explained with the help of electron dot diagrams.

• 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.

Types of salts: normal salts, acid salt, basic salt, definition and examples.

Decomposition of hydrogen carbonates, carbonates, sulphites and sulphides by appropriate acids with heating if necessary. (Relevant laboratory work must be done).

• Direct combination
• Displacement
• Precipitation (double decomposition)
• Neutralization of insoluble base
• Neutralisation of an alkali (titration)
• The action of dilute acids on carbonates and bi-carbonates.

• Colour of salt and its solution.
• Action on addition of sodium hydroxide to a 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 the addition of ammonium hydroxide to the 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.

• 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.
• 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 understanding).
• Simple calculations based on the molar volume and Gay Lussac’s law.

The explanation can be given using equations for the formation of HCl, NH3, and NO.

• 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.

• 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 the relation of the mole to mass, volume and Avogadro’s number.

Related to the weight and/or volumes of both reactants and products.

• 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.

• 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.

• 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 and refining of copper.

• Mineral and ore – Meaning only.
• Common ores of iron, aluminium and zinc.

(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.

Stainless steel, duralumin, brass, bronze, fuse metal/solder.

• 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, and sulphites.
• Precipitation reactions with silver nitrate solution and lead nitrate solution.

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 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 are 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 are 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.

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 reacts with copper, carbon, and sulphur.

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, and 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.

• Unique nature of Carbon atom – tetra valency, catenation.
• Formation of single, double and triple bonds, straight chain, branched chain, cyclic compounds (only benzene).

• Structure of compounds with single, double and triple bonds.
• Structural formulae of hydrocarbons. The structural formula must be given for alkanes, alkenes, alkynes up to 5 carbon atoms.
• Isomerism – structural (chain, position)

Alkane, alkene, and alkyne series and their gradation in properties and the relationship with the molecular mass or molecular formula.

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).

• 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.

• 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.

• 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.