ISC EXAM LATEST UPDATE
The Class 11 exams depend on the schools, as they decide when to conduct the exams. Usually, the Class 11 exams are held a few months before the final board exams. This gives you the opportunity to prepare for the final board exams and also mentally prepare for them. Therefore, it is better to contact your school regarding the Class 11 exams, and also take advantage of the complete information provided here for the final board exam.
The Council for the Indian School Certificate Examinations (CISCE) has officially released the date sheet for ISC (Class 12) board examinations for 2025. This year class 12th board exam will be started from 13 February 2025 and will last upto 05 April 2025, it will take around 52 days to complete the class 12th board exam. The first exam will be held on 13 February 2025 for Environmental Science subject and the last exam will be held on 05 April 2025 for Art subject. The Indian School Certificate Examination results will be declared in the month of May.
Official Syllabus
There will be two papers in the subject :
Paper I : Theory- 3 hours ... 70 marks
Paper II: Practical - 3 hours ... 15 marks
Project work ... 10 marks
Practical File … 5 Marks
Paper 1- Theory: 70 Marks
There will be no overall choice in the paper. Candidates have to answer all the questions. Internal choice will be available in two questions of 2 marks, two questions of 3 marks and all three questions of 5 marks.
| unit | Total weight |
| Some Basic Concepts of Chemistry | Physical Chemistry 32 Marks |
| structure of atom |
| Classification of Elements and Periodicity in Properties |
| Chemical Bonding and Molecular Structure |
| States of Matter: Gases and Liquids |
| Chemical thermodynamics |
| balance |
| redox reactions | Inorganic Chemistry 15 Marks |
| hydrogen |
| s-block elements |
| some p-block elements |
| Organic Chemistry: Some Basic Principles and Techniques | Organic Chemistry 23 Marks |
| hydrocarbons |
| environmental chemistry |
| the whole | 70 marks |
Paper I – Theory – 70 Marks
1. Some Basic Concepts of Chemistry
General Introduction: Importance and Scope of Chemistry.
study of matter. Understanding the laws of chemical combination. Dalton's Atomic Theory: Concept of Elements, Atoms and Molecules.
Isotopic (atomic) and molecular masses, mole concept and molar mass, percentage composition, empirical and molecular formulas. Stoichiometry and calculations based on chemical reactions.
(i) Accuracy and precision:
Quantities in Chemistry and their measurements, significant figures, SI units.
(ii) Dimensional Analysis:
Conversion of units, numerical and applications of units.
(iii) The concept of atoms with definite properties in explaining the laws of chemical combination.
Study about atoms. Atomic theory of Dalton: Main principles of the theory; its limits.
Laws of Chemical Combination:
- Law of conservation of mass.
Law of definite proportion.
Law of multiple proportions.
- Law of reciprocal proportion.
Gelussak 's law of gaseous volumes.
Statements, explanations and simple problems based on these laws.
(iv) Atomic (isotopic mass) and molecular mass.
Relative Molecular Mass and Mole:
The atomic mass unit is one of the experimentally determined units. This is equal to 1/12 of the mass of the carbon 12 isotope .
Numerical problems based on mole concept, Avogadro's number and gram molecular mass .
(v) Empirical and Molecular Formulas: Numerical based on the above.
(vi) Chemical equivalent in terms of normality, and volumetric calculation. C = 12.00 to be taken as a standard for expressing the atomic mass.
Equivalent weight expresses the combining ability of elements with standard elements like H, Cl, O, Ag, etc. variable equivalent weight. Relationship between gram equivalent weight, gram equivalent weight, gram molecular mass and valency.
Determination of equivalent weight of acids, bases, salts, oxidizing and reducing agents. (experimental details not required).
Terms used in volumetric calculations such as percentage (w/w and w/v), normality, molarity, molality, mole fraction, etc. should be discussed. Students are required to know the formulas and normality and molarity equations.
Simple calculation on above topics.
(vii) Chemical reactions – stoichiometric calculations based on mass-mass, mass-volume, volume-volume relationships and limiting reactant.
2. Structure of the atom
Discovery of fundamental particles electron, proton and neutron), atomic numbers, isotopes and isobars. Thomson's model and its limitations. Rutherford's experimental model and its limitations. The dual nature of matter and light. Bohr's atomic model and its limitations (de Broglie's equation, Heisenberg's uncertainty principle), concept of shells, subshells, orbitals. Quantum numbers, sizes of s, p and d orbitals. Laws of filling electrons in orbitals – Aufbau principle, Pauli's exclusion principle and Hund 's law of maximum multiplicity. Electronic configuration of atoms, stability of half filled and completely filled orbitals.
(i) Subatomic particles (electrons, protons and neutrons) with their charge and mass: The concept of indivisibility of the atom as proposed by Dalton does not exist. An atom consists of subatomic fundamental particles. Production and properties of cathode rays . Production and properties of anode rays. Chadwick's experiment for the discovery of neutrons and the properties of neutrons .
(ii) Rutherford's nuclear model based on scattering experiment : Rutherford's scattering experiment. Discovery of the nucleus. Rutherford's atomic model of the atom. Defects of Rutherford's model . Electromagnetic Wave Theory and its Limitations (Black Body Radiation and Photoelectric Effect)
Planck's quantum theory.
Arithmetic based on the above.
(iii) Types of spectra: Emission and absorption spectra. Band and line spectra are to be discussed.
(iv) Bohr's atomic model.
The principles of Bohr's theory - based on Planck's quantum theory.
Properties of Bohr's atomic model and interpretation of hydrogen spectra. Calculation based on Rydberg's formula.
Bohr's points on atomic radii, velocities and energies of orbitals (derivation not necessary).
Flaws in Bohr's model.
(v) The quantum mechanical model of the atom – a simple mathematical treatment. quantum number; Size, shape and orientation of s, p and d orbitals only (no derivation). Aufbau principle, Pauli's exclusion principle, Hund's law of maximum multiplicity. Electronic configuration of elements with respect to s, p, d, f subshells .
• De Broglie's equation. Number.
• Heisenberg's Uncertainty Principle. Number.
• Schrödinger wave equation - ? and |?| Physical importance of 2.
• Quantum numbers - types , size, shape and orientation of s, p and d subshells of quantum numbers. Information obtained in terms of nucleus, node, nodal planes and radial probability curve, energy of electron, number of electrons present in an orbit and distance of electron from an orbital .
• Structure theory, (n+l) rule.
• Pauli's Exclusion Principle.
• Hund's law of maximum multiplicity.
• Electronic configuration of elements and ions in terms of s, p, d, f subshells and stability of half filled and completely filled orbitals.
3. Classification of Elements and Periodicity in Properties
Importance of classification; The modern periodic law and the current form of the periodic table lead to periodic trends in the properties of elements - atomic radii, ionic radii, valency, ionization enthalpy, electron gain enthalpy, electronegativity. Nomenclature of elements with atomic number greater than 100.
(i) Modern Periodic Law
Modern periodic law (atomic number taken as the basis of classification of elements).
(ii) Long form of periodic table.
General characteristics of groups and periods. Division of the periodic table into S, P, D and F blocks. IUPAC nomenclature for elements with Z > 100.
(iii) Periodic trends in the properties of elements.
Atomic radius, ionic radius, ionization enthalpy, electron gain enthalpy, electronegativity, metallic and non-metallic properties.
• Periodic properties like valence electrons, atomic and ionic radii and their variation in groups and periods.
• Ionization enthalpy, electron gain enthalpy and electronegativity should be given an idea and their variation across groups and periods can be discussed.
• Factors (atomic number, screening effect and shielding effect, number of electrons in the outermost orbit) that
affect these periodic properties and their variation across groups and periods.
(iv) Periodic trends in chemical properties – periodicity of valence or oxidation states. Anomalous properties of second period elements.
diagonal relationship; Acidic and basic nature of oxides.
Note: The latest IUPAC recommendations for numbering groups are followed. Numbering 1-18 in place of the old numbering of I-VIII. Details are given at the end of the course.
4. Chemical Bonding and Molecular Structure
Valence electrons, Ionic bond character, Covalent bonds of ionic bonds, Covalent bonds, Bond parameters, Lewis structure, Polar character of covalent bonds, VSEPR theory, Geometry of covalent molecules, Valence bond theory, Concept of hybridization involving s, p and d orbitals Contains and sizes of some simple molecules. coordination bond. Molecular orbital theory of homonuclear diatomic molecules (qualitative idea only). resonance and hydrogenbond.
(i) Kossel-Lewis approach to chemical bonding. Octet rule, its applications in electrovalent and covalent bonds.
(ii) Electrovalent or ionic bond: Lewis structures of NaCl, Li2O, MgO, CaO, MgF2 and Na2S . Definition of ionic bond .
Necessary conditions for the formation of ionic bond such as:
Low ionization enthalpy of metals . High electron gain enthalpy of nonmetals
- High lattice energy. - electronegativity difference between the reacting atoms.
All these points should be discussed in detail.
Formation of cations and anions of elements and their position in the periodic table.
variable electrical connectivity; Due to variable electronegativity i.e. due to inert pair effect and due to unstable core, using suitable examples.
Characteristics of electrovalent bond.
(iii) Covalent bond – bond parameters, Lewis structure, polar character of covalent bond, shape.
Formation of sigma and pi bonds such as ammonia, nitrogen, ethene, ethyne and carbon dioxide.
Definition of covalent bond, conditions for the formation of covalent bond, types of covalent bond, i.e. single, double and triple bond. Sigma and Pi bonds: H 2 , O 2 , N 2 .Classification of covalent bonds on the basis of electronegativity of atoms – polar and non-polar covalent bonds, dipole moment.
Formation of CH 4 , NH 3 , H 2 O, Ethane, Ethane, Ethyne and CO 2 etc. and their electron dot structure or Lewis structure.
Characteristics of covalent compounds. Comparison between electronegativity and covalency. Due to variable covalency such as phosphorus 3 and 5 and sulfur 2, 4, 6 and chlorine 1, 3, 5 and 7. Formal charge of anion.
(iv) Deviation from the rules of the octave and the rules of Fazan. Definition of octet rule.
Failure of the octet rule, either due to incomplete octets or excess of octets with suitable examples.
Faizan's Laws: Statement, conditions for electrovalency and covalency. Polar and non-polar bonds should be correlated by Fagen's laws.
(v) Valence Shell Electron Pair Repulsion (VSEPR) theory; Hybridization and shape of molecules: Hybridization involving only s, p and d orbitals.
Concept of electron-pair repulsion and shape of molecules using suitable examples. Hybridization and molecular shape - Definition, hybridization of orbitals containing s, p and d orbitals (using suitable examples).
(vi) Molecular orbital theory: Qualitative treatment of homonuclear diatomic molecules of first two periods (hydrogen to neon), Energy level diagrams, bonding and antibonding molecular orbitals, bond order, paramagnetism of O2 molecule. Relative stabilities of O2, O2 - , O2 2 - , O2 + and N2, N2 +, N2 - , N2 2- .
(vii) Co-ordinate or dative covalent bond, e.g. formation of oxy-acids of chlorine: Co-ordinate or dative covalent bonding: definition, formation of chlorous acid, chloric acid, perchloric acid, ammonium ion, hydronium ion, nitric acid, ozone.
(viii) Resonance in simple inorganic molecules: Resonance in simple inorganic molecules like ozone, carbon dioxide, carbonate ion and nitrate ion.
(ix) Hydrogen bonding: the examples of hydrogen fluoride, water (ice), alcohol, etc. may be considered. H-bonding – definition, types, condition for hydrogen bond formation, examples of inter-molecular hydrogen bonding in detail taking hydrogen fluoride, water and ice and ethanol into account. Intramolecular hydrogen bonding.
5. States of Matter: Gases and Liquids
States of matter and their characteristic properties to establish the concept of the molecule. Boyle's law, Charles law, GayLussac's law, Avogadro's law, Avogadro’s number, ideal behaviour of gases and derivation of ideal gas equation. Kinetic Theory of gases, kinetic energy and molecular speeds (elementary idea). Deviation from ideal behaviour, van der Waal’s equation, liquefaction of gases, critical temperature. Liquid state - vapour pressure, viscosity and surface tension (qualitative idea only, no mathematical derivatives).
(i) Intermolecular interactions (van der Waals forces), types of van der Waals forces, melting and boiling points.
(ii) The Gas Laws.
Boyle’s law, Charles’ law, Absolute temperature, pressure temperature law, Avogadro’s law and Avogadro’s constant. Relationship between the mole and Avogadro’s number.
Simple numerical problems based on the above laws.
Dalton’s law, Graham’s law of diffusion. Dalton’s law of partial pressures and it’s application Graham’s Law of diffusion and its application. Numerical problems based on the above.
(iii) Ideal gas equation and application of this equation. Ideal gas equation PV = nRT; its application in calculation of relative molecular mass and in the calculation of the value of R.
(iv) Kinetic Theory of gases. Characteristics of gases, comparison between solid, liquid and gas. Properties of gases on the basis of kinetic theory of gases. Postulates of kinetic theory must be discussed to explain gas laws. Concept of average, root mean square and most probable velocities (numericals not required). Non ideal behaviour of gases i.e. deviation from ideal gas equation may be discussed at low and at high temperature and pressure. van der Waals’ equation (P + a/V2 ) (V-b) = RT for one mole of a gas. (numericals not required). The pressure correction and volume correction may be explained. significance and units of ‘a’ and ‘b’ (van der Waals’ constant). Liquefaction of gases, critical temperature.
(v) Liquid State - vapour pressure, viscosity and surface tension. Qualitative idea only, no mathematical derivations
6. Chemical Thermodynamics
(i) Introduction, concepts, types of system, surroundings, extensive, intensive properties and state functions. Types of system – ideal system, real system, isolated system, closed system, open system. Meaning of surroundings. Properties of the system: macroscopic, intensive and extensive properties. State of the system. Main processes the system undergoes: reversible, irreversible, adiabatic, isothermal, isobaric, isochoric, cyclic. Meaning of thermodynamic equilibrium. Meaning of thermodynamic process.
(ii) First Law of Thermodynamics and its significance, work, heat, internal energy, enthalpy (?U or ?E and ?H), heat capacity and specific heat. Hess's law of constant heat summation, enthalpy of bond dissociation, combustion, formation, atomisation, sublimation, phase transition, ionisation, solution and dilution. Meaning of: internal energy of the system, work done by the system, by the surroundings at constant temperature, heat absorbed by the system and by the surroundings at constant temperature. The sign convention for change in internal energy, heat given out or gained, work done by the system or by the surroundings. State function and path function - meaning with examples. Internal energy change, work done and heat absorbed in a cyclic process. Internal energy change in an isolated system and in a non-isolated system. Total internal energy change of a system and surroundings.
Mathematical statement of the first law. Significance of first law of thermodynamics. Need for enthalpy – constant pressure or open vessel processes. Enthalpy - a thermodynamic property, state function. Mathematical form of enthalpy. Heat - the energy in transit. Conditions for the transfer of heat. Limitations in conversion of heat into work. Condition at which heat transfer ceases, unit of heat. Meaning of work, capacity to do work,types of work. Mathematical form of reversible work and irreversible work. Difference between the reversible and irreversible work done – graphically.
Relationship between Cv and internal energy change. Relationship between Cp and Cv. Definitions of the following: Heat of reaction: Heat of formation – standard heat of formation, Heat of solution, Heat of dilution, Heat of neutralization, Heat of combustion. Constancy in the heat of neutralisation: Experimental verification in case of strong acids and strong bases. Reason for that observation – ionic neutralisation and the heat evolved. Definition of Calorific value of a fuel. Statement of Hess’ Law and its application. Problems based on Hess’ Law.
(iii) Second Law of Thermodynamics and its significance, spontaneity of a chemical change; Entropy, Free Energy. Inadequacy of First Law and need for Second Law; Ideas about reversible (recapitulation), spontaneous and non-spontaneous processes Meaning of entropy – derived from Second Law – statement of Second Law in terms of entropy; Physical significance of entropy; State function and not path function.
Entropy change of the universe, reversible isothermal process and irreversible process. Meaning of thermal death, Gibb’s free energy of the system and Helmholtz free energy. Relationship between Gibb’s free energy and Helmholtz’s free energy. Relationship between change in Gibb’s free energy and equilibrium constant of a chemical reaction. Defining the criteria for spontaneity of a chemical change in terms of Gibb’s free energy. Note: Numericals based on the First Law, Second Law of Thermodynamics and Hess’ Law.
(iv) Third Law of Thermodynamics – statement only. Self-explanatory.
7. Equilibrium
(i) Chemical Equilibrium. Introduction of physical and chemical equilibrium and its characteristics Dynamic nature of equilibrium, law of mass action, equilibrium constant and factors affecting equilibrium. Le Chatelier's principle and its applications. Irreversible and reversible reactions. Physical equilibrium: solid-liquid, liquidvapour, solid-vapour; Characteristics of Physical equilibrium.
Chemical equilibrium: Characteristics of chemical equilibrium; dynamic nature. Law of mass action; Equilibrium constant in terms of concentration Kc. Gaseous reactions; Equilibrium constant in terms of partial pressures Kp. Relationship between Kp and Kc (derivation required); Characteristics of equilibrium constant; Units for equilibrium constant; Simple calculations of equilibrium constant and concentration. The following examples should be considered to show maximum yield of products:
- Synthesis of ammonia by Haber’s process. - The dissociation of dinitrogen tetra oxide.
- Hydrolysis of simple esters. - The contact process for the manufacture of sulphuric acid. Le Chatelier’s Principle. Statement and explanation. Factors affecting chemical and physical equilibria should be discussed in the light of Le Chatelier’s principle. - Change of concentration. - Change of temperature. - Change of pressure. - Effect of catalyst. - Addition of inert gas.
(ii) Ionic equilibrium Introduction, electrolyte (strong and weak), non-electrolyte, ionisation, degree of ionisation of polybasic acids, acid strength, concept of pH, pH indicators, buffer solution, common ion effect (with illustrative examples). Henderson equation, hydrolysis of salts, solubility and solubility product. Ostwald’s dilution law and its derivation. Strength of acids and bases based on their dissociation constant. Problems based on the Ostwald’s dilution law.
Arrhenius, Brönsted-Lowry and Lewis concept of acids and bases, multistage ionisation of acids and bases with examples.
Ionic product of water
– definition, pH, pOH, pKw of solutions. pH indicators and their choice in titrimetry. Numericals on the above concepts. Common ion effect
– definition, examples (acetic acid and sodium acetate; ammonium hydroxide and ammonium chloride), applications in salt analysis. Salt hydrolysis – salts of strong acids and weak bases, weak acids and strong bases, weak acids and weak bases and the pH formula of the solutions of these salts in water with suitable examples. Buffer solutions: definition, examples, action; its interpretations based on Le Chatelier’s principle. Henderson equation. Solubility product: definition and application in qualitative salt analysis (Group II, III and IV cations). Numericals on pH, buffer solutions, solubility and solubility product.
8. Redox Reactions Concept of oxidation and reduction, redox reactions, oxidation number, change in oxidation number, balancing redox reactions (in terms of loss and gain of electrons). Applications of redox in various types of chemical reactions. ? Concept of oxidation and reduction in terms of oxygen, hydrogen, electrons. ? Redox reactions – examples. ? Oxidation number: rules for calculation, simple calculations of oxidation state in molecules and ions like K2Cr2O7, S2 2 O3 ? , etc. ? Oxidation and reduction in terms of change in oxidation number. ? Balancing of redox reactions in acidic and basic medium by oxidation number and ionelectron method.
9. Hydrogen
Hydrogen and its compounds: hydrides, water, heavy water, hydrogen peroxide. Position of hydrogen in periodic table, occurrence, isotopes, preparation, properties and uses of hydrogen, hydrides (ionic covalent and interstitial); hydrogen as a fuel. Physical and chemical properties of water, soft and hard water, and removal of hardness of water, heavy water. Hydrogen peroxide: Preparation from peroxide, structure, oxidising properties: reaction with KI, PbS, acidified FeSO4; reducing properties – reaction with acidified KMnO4 and chlorine. Calculation of strength of hydrogen peroxide
10. s-Block Elements (Alkali and Alkaline Earth Metals)
(i) Group 1 and 2 elements General characterises of Group 1 and 2 should include the following: Occurrence; physical state; electronic configuration; atomic and ionic radii; common oxidation state; electropositive /electronegative character; ionisation enthalpy; reducing/oxidising nature; distinctive behaviour of first member of each group (namely Lithium, Beryllium); nature of oxides, hydroxides, hydrides, carbonates, nitrates; chlorides, sulphates.
(ii) Preparation and properties of some important compounds. Sodium chloride, sodium hydroxide, Sodium carbonate, sodium bicarbonate, sodium thiosulphate; biological importance of sodium and potassium. Magnesium chloride hexahydrate, calcium oxide, calcium hydroxide, calcium carbonate, plaster of paris and cement. Industrial uses of the above, biological importance of magnesium and calcium.
Group 1:
• Sodium chloride - uses. • Sodium hydroxide - only the principle of preparation by Castner-Kellner cell. Uses • Sodium carbonate – principal and equation of Solvay’s process. Uses. • Sodium bicarbonate - preparation from sodium carbonate. Uses. • Sodium thiosulphate - preparation from sodium sulphite and its reaction with iodine, dilute acids and silver nitrate. Uses.
Group 2:
• Magnesium chloride hexahydrate - preparation from magnesium oxide. Effect of heat. Uses • Calcium oxide - preparation from limestone; reaction with water, carbon dioxide and silica. Uses. • Calcium hydroxide – preparation from calcium oxide and uses. • Calcium carbonate – preparation from calcium hydroxide and uses. • Plaster of Paris - preparation from gypsum. Uses. • Manufacture of cement. Uses.
11. Some p -Block Elements
(i) Group 13 Elements General introduction, electronic configuration, occurrence, variation of properties, oxidation states, trends in chemical reactivity, anomalous properties of first element of the group, Boron - physical and chemical properties.
(ii)Preparation and properties of some important compounds, borax, boric acid, boron hydrides, aluminium: Reactions with acids and alkalies. Lewis acid character of boron halides; amphoteric nature of aluminium, alums. Borax- reaction with water and action of heat on hydrated compound (preparation not required). Borax Bead Test. Diborane - Preparation properties, structure and uses. Boric acid – preparation and action of heat. Aluminium: Reactions with acids and alkalies. Alums – preparation and uses
(iii) Group 14 Elements General characteristics, electronic configuration, occurrence, variation of properties, oxidation states, trends in chemical reactivity, anomalous behaviour of first elements. Carbon-catenation, allotropic forms. Structure of diamond graphite and fullerene; stability of +2 oxidation state down the group in terms of inert pair effect.
(iv) Some important compounds; oxides of carbon and silicon, silicon carbide, silicon tetra chloride, silicones, silicates and zeolites. Preparation and properties of: Carbon monoxide - preparation from incomplete combustion of carbon. Hazards of CO. Reducing nature of CO. Carbon dioxide - preparation from limestone and carbon, limewater test. Uses. Silicon dioxide - structure, comparison with carbon dioxide. Uses. Silicon carbide - preparation from silica. Uses. Silicon tetra chloride - preparation from silicon and uses. Silicones - general method of preparation. Uses. Silicates – structure and uses. Zeolites – formula and use.
12. Organic Chemistry - Some Basic Principles and Techniques General introduction, classification and IUPAC nomenclature of organic compounds and isomerism. Methods of purification, qualitative and quantitative analysis. Electron displacement in a covalent bond: inductive effect, electromeric effect, resonance and hyperconjugation. Homolytic and heterolytic bond fission of a covalent bond: free radicals, carbocations, carbanions, electrophiles and nucleophiles, types of organic reactions.
(i)Introduction to organic chemistry: Vital force theory, reason for separate study of organic chemistry and its importance, characteristics of carbon atoms (tetra valency), Reasons for large number of organic compounds: catenation, isomerism and multiple bonding, etc.
(ii)Classification of organic compounds: (definition and examples): open chain, closed chain, homocyclic, hetrocyclic, aromatic, alicyclic compounds, homologous series and its characteristics, functional groups.
(iii) IUPAC rules for naming organic compounds. Aliphatic, alicyclic and aromatic compounds.
(iv) Definition and classification of isomerism:
Structural isomerism: definition, classification, examples.
Chain isomerism, Positional isomerism, Functional isomerism, Metamerism, Tautomerism - examples for each of the above.
Stereoisomerism: definition and classification, examples.
Geometrical isomerism: Definition. Conditions for compounds to exhibit geometrical isomerism; types and examples, cis and trans, syn and anti. Examples.
Optical isomerism: Definition, Nicol prism, plane polarised light. polarimeter. Method of measuring angle of rotation. Specific rotation. Conditions for optical activity. d, l form; External compensation, Internal compensation, racemic mixture & meso form. Examples – lactic acid and tartaric acid.
(v) Analysis of organic compounds:
Detection of elements (qualitative analysis) such as carbon, hydrogen, nitrogen, halogens and sulphur should be considered by using Lassaigne’s test and reactions involved in it.
(vi) Estimation of carbon, hydrogen, nitrogen, halogens, sulphur and phosphorous:
Estimation of carbon and hydrogen – Leibig’s method.
Estimation of nitrogen - Kjeldahl’s method.
Estimation of halogens sulphur and phosphorous - Carius method. Numericals included. Experimental details required.
(vii)Types of chemical reactions and their mechanisms
Substitution, addition, elimination reactions: definition and examples.
Homolytic and heterolytic fission – definition and examples. Free radicals, carbocation, carbanion (their reactivities and stabilities).
Electrophiles and nucleophiles – definition and examples (including neutral electrophiles and nucleophiles).
Inductive, electromeric, mesomeric effect and hyperconjugation – definition, examples.
(viii)Free radicals and polar mechanisms
In terms of fission of the bonds and formation of the new bonds including SN1, SN2, E1 and E2 mechanisms. Explain with relevant examples and conditions.
13. Hydrocarbons
Classification of Hydrocarbons
I. Aliphatic Hydrocarbons
(i) Alkanes - Nomenclature, isomerism, conformation (methane and ethane), physical properties, chemical properties including free radical mechanism of halogenation, combustion and pyrolysis.
Occurrence, conformation (Sawhorse and Newman projections of ethane).
General methods of preparation: from sodium salts of carboxylic acids (decarboxylation and Kolbe’s electrolytic method); from alcohols and alkyl halides (Wurtz reaction, Coreyhouse Synthesis). From aldehydes and Grignard’s Reagent.
Physical and chemical properties of alkanes.
Physical properties: state, freezing point, melting point, boiling point, density.
Chemical properties: combustibility, reaction with chlorine (free radical mechanism), reaction with oxygen in presence of catalyst (formation of alcohol, aldehyde, and carboxylic acid). Cyclisation, aromatisation, isomerisation and pyrolysis.
Uses of alkanes.
(Markownikoff's addition and peroxide effect), ozonolysis, oxidation, mechanism of electrophilic addition. General methods of preparation – dehydration of alcohols, dehydrohalogenation of alkyl halides (from vicinal dihalides), Kolbe’s electrolytic method and from alkynes. Physical Properties: State, freezing point, melting point, boiling point, dipole moment, density. Chemical properties - addition reactions (hydrogen, halogens, hydrogen halides, sulphuric acid, water). Markownikoff’s rule and antiMarkownikoff’s rule with mechanism and examples. Oxidation: complete combustion, hot and cold alkaline KMnO4 (Baeyer’s reagent), ozonolysis. Polymerisation. Saytzeff’s rule and its application. Uses of alkenes.
(ii)Alkenes - Nomenclature, structure of double bond (ethene), isomerism; methods of preparation; physical properties, chemical properties; addition of hydrogen, halogen, water, hydrogen halides
(Markownikoff's addition and peroxide effect), ozonolysis, oxidation, mechanism of electrophilic addition.
General methods of preparation – dehydration of alcohols, dehydrohalogenation of alkyl halides (from vicinal dihalides), Kolbe’s electrolytic method and from alkynes.
Physical Properties: State, freezing point, melting point, boiling point, dipole moment, density.
Chemical properties - addition reactions (hydrogen, halogens, hydrogen halides, sulphuric acid, water).
Markownikoff’s rule and antiMarkownikoff’s rule with mechanism and examples.
Oxidation: complete combustion, hot and cold alkaline KMnO4 (Baeyer’s reagent), ozonolysis.
Polymerisation.
Saytzeff’s rule and its application.
Uses of alkenes
(iii) Alkynes - Nomenclature, structure of triple bond (ethyne), methods of preparation; physical properties, chemical properties: acidic character of alkynes, addition reactions - hydrogen, halogens, hydrogen halides and water.
General methods of preparations of alkynes. Manufacture of ethyne by calcium carbide and from natural gas. Dehydrohalogenation and Kolbe’s electrolytic method.
Physical properties of alkynes: State of existence, freezing point, melting point, boiling point, density.
Chemical properties of alkynes – addition reactions (hydrogen, halogens, hydrogen halides and water), acidic nature of alkynes, formation of acetylides.
Oxidation: complete combustion, hot and cold alkaline KMnO4 (Baeyer’s reagent), ozonolysis.
Polymerisation.
Uses of alkynes.
Distinguishing test between Alkane, Alkene and Alkyne.
II. Aromatic Hydrocarbons
Introduction, IUPAC nomenclature, benzene: resonance, aromaticity, chemical properties: mechanism of electrophilic substitution. Nitration, sulphonation, halogenation, Friedel Crafts alkylation and acylation, directive influence of functional group in monosubstituted benzene. Carcinogenicity andtoxicity.
Structure: Resonance structures (Kekule’s) of benzene.
Benzene: Preparation from sodium benzoate and from phenol.
Physical properties: State of existence, freezing point, melting point, boiling point, density.
Chemical properties:
- Electrophilic substitution reactions with mechanism (halogenation, nitration, sulphonation).
- Alkylation, acetylation – Friedel Crafts reaction.
- Directive influence (o-, p-, and m-) of substituents in electrophilic and nucleophilic substitutions (with mechanism).
- Oxidation: catalytic oxidation, reaction with ozone.
- Addition reactions with hydrogen, chlorine, bromine.
- Pyrolysis (formation of bi-phenyl). Carcinogenicity and toxicity of benzene may be discussed.
14. Environmental Chemistry
Types of environmental pollution (air, water and soil pollution); various types of pollutants: smog, acid rain; effects of depletion of ozone layer, greenhouse effect and global warming. Pollution due to industrial wastes, green chemistry as an alternative tool for reducing pollution; strategies for control of environmental pollution.
Gaseous pollutants: oxides of nitrogen, carbon, sulphur, hydrocarbons; their sources, harmful effects and prevention; Greenhouse effect and global warming; acid rain;
Particulate pollutants: smoke, dust, smog, fumes, mist; their sources, harmful effects and prevention.
Water pollutants: pathogens, organic waste, chemical pollutants; their harmful effects and prevention.
Soil Pollutants: pesticides, herbicides. Green chemistry as an alternative tool for reducing pollution
PAPER II PRACTICAL WORK- 15 Marks
Candidates are required to complete the following experiments:
1. Basic laboratory techniques:
Cutting a glass tube.
Bending a glass tube.
Drawing out a glass jet.
Boring a cork.
2. Titration: acid-base titration involving molarity.
Titrations involving:
• Sodium carbonate solution/ dil H2SO4 or dil . HCl using methyl orange indicator.
• NaOH or KOH solution/ dil H2SO4 or dil. HCl using methyl orange indicator.
• Calculations involving molarity, concentration in grams L-1 / number of ions, the water of crystallization, and percentage purity.
NOTE: Calculation of molarity must be up to 4 decimal places at least, in order to avoid an error.
OBSERVATION TABLE
| S. No. | (A) | (B) | (B – A) |
| Initial burette reading (ml) | Final burette reading (ml) | Difference (ml) |
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| 2. |
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| 3. |
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• Concordant reading is to be used for titer value. Concordant reading is two consecutive values that are exactly the same. The average will not be accepted as titer value.
• The table is to be completed in ink only. A pencil is not to be used.
• Overwriting will not be accepted in the tabular column.
Observations:
• Pipette size (should be the same for all the candidates at the center):
• Titre value (concordant).
3. Qualitative analysis: identification of single salt containing one anion and one cation:
NOTE:
• For the wet test of anions, sodium carbonate extract must be used (except for carbonate).
• Chromyl chloride test not to be performed.
(Insoluble salts, such as lead sulfate, barium sulfate, calcium sulfate, and strontium sulfate should not be given).
4. Preparation of inorganic compounds.
(a) Preparation of potash alum/Mohr’s salt.
(b) Preparation of crystalline FeSO4/CuSO4.
5. Paper Chromatography.
Preparation of chromatogram, separation of pigments from extracts of leaves and flowers/ink mixtures; determination of Rf value.
PROJECT WORK AND PRACTICAL FILE - 15 Marks
Project Work – 10 Marks
The candidate is to creatively execute one project/assignment on a selected topic of Chemistry. Teachers may assign or students may choose any one project of their choice. (Refer to the suggested topics at the end of the Class XII syllabus).
Suggested Evaluation criteria for Project Work:
• Introduction/purpose
• Contents
• Analysis/ material aid (graph, data, structure, pie charts, histograms, diagrams, etc)
• Presentation
• Bibliography
Practical File – 5 Marks
Teachers are required to assess students on the basis of the Chemistry Practical file maintained by them during the academic year.
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