ISC Class 11 Syllabus 2027 Chemistry

ISC YEAR 2027

INDIAN SCHOOL CERTIFICATE
EXAMINATION

CHEMISTRY
(862)

February 2025
____________________________________________________________________________________________

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MISSION STATEMENT

The Council for the Indian School Certificate
Examinations is committed to serving the nation's
children, through high quality educational
endeavours, empowering them to contribute towards
a humane, just and pluralistic society, promoting
introspective living, by creating exciting learning
opportunities, with a commitment to excellence.

ETHOS OF CISCE

Trust and fair play.
Minimum monitoring.
Allowing schools to evolve their own niche.
Catering to the needs of the children.
Giving freedom to experiment with new ideas
and practices.
Diversity and plurality - the basic strength for
evolution of ideas.
Schools to motivate pupils towards the
cultivation of:
Excellence - The Indian and Global
experience.
Values - Spiritual and cultural - to be the bedrock
of the educational experience.
Schools to have an 'Indian Ethos', strong roots in
the national psyche and be sensitive to national
aspirations.

CHEMISTRY (862)

Aims
1. To foster acquisition of knowledge and 5. To develop an interest in activities involving
understanding of terms, concepts, facts, usage of the knowledge of Chemistry.
processes, techniques and principles relating to
6. To develop a scientific attitude through the study
the subject of Chemistry.
of Physical Sciences.
2. To develop the ability to apply the knowledge of
7. To acquaint students with the emerging frontiers
contents and principles of Chemistry in new or
and interdisciplinary aspects of the subject.
unfamiliar situations.
8. To develop skills relevant to the discipline.
3. To develop skills in proper handling of apparatus
and chemicals. 9. To apprise students with interface of Chemistry
with other disciplines of Science, such as,
4. To develop an ability to appreciate achievements
Physics, Biology, Geology, Engineering, etc.
in the field of Chemistry and its role in nature
and society.

CLASS XI
There will be two papers in the subject: Paper II: Practical - 3 hours ... 15 marks
Paper I: Theory- 3 hours ... 70 marks Project Work … 10 marks
Practical File … 5 marks

PAPER 1- THEORY: 70 Marks
S.No. UNIT TOTAL WEIGHTAGE
1.
Some Basic Concepts of Chemistry
2.
Structure of Atom Physical Chemistry
3. 41 Marks
Classification of Elements and Periodicity in Properties
4.
Chemical Bonding and Molecular Structure
5.
Chemical Thermodynamics
6.
Equilibrium
Inorganic Chemistry
7. Redox Reactions
6 Marks
8. Organic Chemistry
Organic Chemistry: Some basic Principles and Techniques
23 Marks
9.
Hydrocarbons
70 Marks
TOTAL

1

PAPER I –THEORY – 70 Marks (v) Empirical and molecular formula:
Numericals based on the above.
1. Some Basic Concepts of Chemistry
(vi)Chemical equivalents, volumetric calculations
General introduction: Importance and scope of in terms of normality. C = 12.00 should be
chemistry. taken as a standard for expressing atomic
Study of matter. Understanding laws of masses.
chemical combination. Dalton's atomic theory: Equivalent weight expressing the combining
concept of elements, atoms and molecules. capacity of the elements with the standard
elements such as H, Cl, O, Ag, etc.
Isotopic (atomic) and molecular masses, mole
concept and molar mass, percentage Variable equivalent weight. Gram equivalent
composition, empirical and molecular weights, relationship between gram
formula. Stoichiometry and calculations based equivalent weight, gram molecular mass and
on chemical reactions. valency.
(i) Precision and accuracy: Determination of equivalent weight of acids,
alkalis, salts, oxidising and reducing agents.
Quantities and their measurements in (experimental details not required).
Chemistry, significant figures, SI units.
Terms used in volumetric calculations such
(ii) Dimensional analysis: as percentage (w/w and w/v), normality,
Conversion of units, numericals and molarity, molality, mole fraction, etc. should
applications of units. be discussed. Students are required to know
the formulae and normality and molarity
(iii) The concept of atoms having fixed properties equations.
in explaining the laws of chemical
Simple calculations on the above topics.
combination.
(vii)Chemical reactions – stoichiometric
Study about atoms. Dalton’s atomic theory: calculations based on mass-mass,
Main postulates of the theory; its limitations. mass-volume, volume-volume relationships
Laws of chemical combinations: and limiting reagent.
 Law of conservation of mass. 2. Structure of Atom
 Law of definite proportions. Discovery of fundamental particles electron,
proton and neutron), atomic number, isotopes
 Law of multiple proportions. and isobars. Thomson's model and its
 Law of reciprocal proportions. limitations. Rutherford's experimental model
and its limitations. Dual nature of matter and
 Gay Lussac’s law of gaseous volumes. light. Bohr's atomic model and its limitations
Statement, explanation and simple problems (de Broglie's equation, Heisenberg’s uncertainty
based on these laws. principle), concept of shells, subshells, orbitals.
Quantum numbers, shapes of s, p and d
(iv) Atomic (isotopic masses) and molecular mass. orbitals. Rules for filling electrons in orbitals -
Relative molecular mass and mole: aufbau principle, Pauli's exclusion principle
and Hund's rule of maximum multiplicity.
The atomic mass unit is one of the
Electronic configuration of atoms, stability of
experimentally determined unit. It is equal to
half- filled and completely filled orbitals.
1/12 of the mass of the carbon 12 isotope.
(i) Subatomic particles (electrons, protons and
Numerical problems based on mole concept,
neutrons) their charges and masses: Concept
Avogadro’s number and gram molecular
of indivisibility of atom as proposed by
volume.
Dalton does not exist. The atom consists of
subatomic fundamental particles. Production

2

of cathode rays and their properties. radial probability curve, energy of
Production of anode rays and their electron, number of electrons present in
properties. an orbit and an orbital.
Chadwick’s experiment for the discovery of • aufbau principle, (n+l) rule.
neutron and properties of neutron. • Pauli’s exclusion principle.
(ii) Rutherford’s nuclear model based on the
scattering experiment: Rutherford’s • Hund’s rule of maximum multiplicity.
scattering experiment. Discovery of nucleus. • Electronic configuration of elements and
Rutherford’s nuclear model of atom. Defects ions in terms of s, p, d, f subshells and
of Rutherford’s model. Electromagnetic wave stability of half-filled and completely
theory and its limitations (Black body filled orbitals.
radiation and photoelectric effect)
3. Classification of Elements and Periodicity in
Planck’s quantum theory. Properties
Numericals based on the above.
Significance of classification; study of
(iii) Types of spectra: emission and absorption Mendeleev’s periodic law and its limitations;
spectra. Band and line spectra to be Modern Periodic Law and the present form of
discussed. periodic table leading to periodic trends in
(iv) Bohr’s atomic model. properties of elements - atomic radii, ionic radii,
Postulates of Bohr’s theory – based on valency, ionisation enthalpy, electron gain
enthalpy, electronegativity. Nomenclature of
Planck’s quantum theory.
elements with atomic number greater than 100.
Merits of Bohr’s atomic model and
explanation of hydrogen spectra. (i) Modern Periodic Law
Calculations based on Rydberg’s formula. Mendeleev’s periodic law, defects in the
Numericals on Bohr’s atomic radii, velocity Mendeleev’s periodic table. Advantages and
and energy of orbits (derivation not disadvantages. Modern periodic law (atomic
required). number taken as the basis of classification of
Defects in Bohr’s Model. the elements).
(v) Quantum mechanical model of an atom - a (ii) Long form of Periodic Table.
simple mathematical treatment. Quantum General characteristics of groups and
numbers; shape, size and orientation of s, p periods. Division of periodic table as s, p, d
and d orbitals only (no derivation). aufbau and f blocks. IUPAC nomenclature for
principle, Pauli’s exclusion principle, elements with Z> 100.
Hund’s rule of maximum multiplicity.
Electronic configuration of elements in terms (iii)Periodic trends in properties of elements.
of s, p, d, f subshells. Atomic radius, ionic radius, ionisation
• de Broglie’s equation. Numericals. enthalpy, electron gain enthalpy,
electronegativity, metallic and non-metallic
• Heisenberg’s Uncertainty Principle. characteristics.
Numericals.
• Periodic properties such as valence
• Schrodinger Wave Equation – physical
electrons, atomic and ionic radii and their
significance of Ψ and |Ψ| 2.
variation in groups and periods.
• Quantum numbers – types of quantum
numbers, shape, size and orientation of • The idea of ionisation enthalpy, electron
the s, p and d subshells. Information gain enthalpy and electronegativity must
obtained in terms of distance of electron be given and their variation in groups and
from the nucleus, node, nodal planes and periods may be discussed.

3

• The factors (atomic number, screening Variable electrovalency; reasons for variable
effect and shielding effect, the number of electrovalency i.e, due to inert electron pair
electrons in the outermost orbit) which effect and unstable core, by using suitable
affect these periodic properties and their examples.
variation in groups and periods. Calculation of lattice enthalpy (Born-Haber
(iv) Periodic trends in chemical properties – cycle).
periodicity of valence or oxidation states. Characteristics of electrovalent bond.
Anomalous properties of second period (iii) Covalent Bond – Bond parameters, Lewis
elements. structure, polar character of covalent bond,
Diagonal relationship; acidic and basic nature shapes.
of oxides. Sigma and pi bonds e.g. formation of
NOTE: Recommendations of the latest IUPAC ammonia, nitrogen, ethene, ethyne, and
for numbering of groups to be followed. carbon dioxide.
Numbering 1 – 18 replacing old notation of Definition of covalent bond, conditions for
I – VIII. Details given at the end of the formation of covalent bonds, types of covalent
syllabus. bonds, i.e single, double and triple bonds.
4. Chemical Bonding and Molecular structure Sigma and pi bonds: H 2 , O 2 , N 2 .
Classification of covalent bonds based on
Valence electrons, ionic bond character, covalent
electronegativity of atoms - polar and non-
bond of ionic bond, covalent bond, bond
polar covalent bond, dipole moment.
parameters, lewis structure, polar character of
covalent bond, VSEPR theory, geometry of Formation of CH 4 , NH 3 , H 2 O, ethane, ethene,
covalent m o l e c u l e s , valence bond theory, ethyne and CO 2 , etc. and their electron dot
concept of hybridisation involving s, p and d structure or Lewis structure.
orbitals and shapes of some simple molecules. Characteristics of covalent compounds.
Coordinate bond. Molecular orbital theory of Comparison in electrovalency and covalency.
homonuclear diatomic molecules (qualitative
Reason for variable covalency e.g.
idea only). Resonance and hydrogen bond.
Phosphorus 3 & 5 and sulphur 2, 4, 6 &
(i) Kossel-Lewis approach to chemical bonding. chlorine 1, 3, 5 and 7.
Octet rule, its application to electrovalent and Formal charge of ions.
covalent bonds.
(iv) Deviation from octet rule and Fajan’s rules.
(ii) Electrovalent or ionic bond: Lewis structures
of NaCl, Li 2 O, MgO, CaO, MgF 2, and Na 2 S. Definition of octet rule.
Definition of ionic bond. Failure of octet rule, due to either incomplete
octet or exceeding of octet with suitable
The conditions necessary for the formation of examples.
ionic bonds such as:
Fajan’s rules: statements, conditions for
- low ionisation enthalpy of metals. electrovalency and covalency. Polar and non
- high electron gain enthalpy of non- polar bonds should be correlated with
metals. Fajan’s rules.
- high lattice energy. (v) Valence Shell Electron Pair Repulsion
- electronegativity difference between the (VSEPR) Theory; Hybridisation and shapes
reacting atoms. of molecules: hybridisation involving s, p and
All these points must be discussed in detail. d orbitals only.
The formation of cations and anions of Concept of electron-pair repulsion and
elements and their positions in the periodic shapes of molecules using suitable examples.
table.

4

Hybridisation and molecular shapes – (ii) First Law of Thermodynamics and its
definition, hybridisation of orbitals involving significance, work, heat, internal energy,
s, p and d orbitals (using suitable examples). enthalpy (∆U or ∆E and ∆H), heat capacity
(vi) Molecular orbital theory: Qualitative and specific heat. Hess's law of constant
treatment of homonuclear diatomic heat summation, enthalpy of bond
molecules of first two periods (hydrogen to dissociation, combustion, formation,
neon), Energy level diagrams, bonding and atomisation, sublimation, phase transition,
antibonding molecular orbitals, bond order, ionisation, solution and dilution.
paramagnetism of O 2 molecule. Relative Meaning of: internal energy of the system,
stabilities of O 2 , O 2 -, O 2 2 - , O 2 + and N 2 , work done by the system, by the
N 2 +, N 2 -, N 2 2-. surroundings at constant temperature, heat
(vii) Co-ordinate or dative covalent bond, e.g. absorbed by the system and by the
formation of oxy-acids of chlorine: surroundings at constant temperature.
Co-ordinate or dative covalent bonding: The sign convention for change in internal
definition, formation of chlorous acid, energy, heat given out or gained, work done
chloric acid, perchloric acid, ammonium ion, by the system or by the surroundings.
hydronium ion, nitric acid, ozone.
State function and path function - meaning
(viii) Resonance in simple inorganic molecules: with examples. Internal energy change, work
Resonance in simple inorganic molecules like done and heat absorbed in a cyclic process.
ozone, carbon dioxide, carbonate ion and
Internal energy change in an isolated system
nitrate ion.
and in a non-isolated system. Total internal
(ix) Hydrogen bonding: the examples of hydrogen energy change of a system and surroundings.
fluoride, water (ice), alcohol, etc. may be Mathematical statement of the first law.
considered.
Significance of first law of thermodynamics.
H-bonding – definition, types, condition for Need for enthalpy – constant pressure or
hydrogen bond formation, examples of open vessel processes. Enthalpy - a
inter-molecular hydrogen bonding in detail thermodynamic property, state function.
taking hydrogen fluoride, water and ice and Mathematical form of enthalpy.
ethanol into account. Intramolecular
Heat - the energy in transit. Conditions for
hydrogen bonding.
the transfer of heat. Limitations in conversion
of heat into work. Condition at which heat
5. Chemical Thermodynamics
transfer ceases, unit of heat.
(i) Introduction, concepts, types of system,
Meaning of work, capacity to do work,types
surroundings, extensive, intensive
of work. Mathematical form of reversible
properties and state functions.
work and irreversible work. Difference
Types of system – ideal system, real system, between the reversible and irreversible work
isolated system, closed system, open system. done – graphically.
Meaning of surroundings.
Relationship between C v and internal energy
Properties of the system: macroscopic, change. Relationship between C p and C v .
intensive and extensive properties.
Definitions of the following:
State of the system.
Heat of reaction: Heat of formation –
Main processes the system undergoes: standard heat of formation, Heat of solution,
reversible, irreversible, adiabatic, Heat of dilution, Heat of neutralization, Heat
isothermal, isobaric, isochoric, cyclic. of combustion.
Meaning of thermodynamic equilibrium. Constancy in the heat of neutralisation:
Meaning of thermodynamic process.
Experimental verification in case of strong
acids and strong bases. Reason for that

5

observation – ionic neutralisation and the Chemical equilibrium: Characteristics of
heat evolved. chemical equilibrium; dynamic nature. Law
Definition of Calorific value of a fuel. of mass action; Equilibrium constant in
terms of concentration K c. Gaseous
Statement of Hess’ Law and its application. reactions; Equilibrium constant in terms of
Problems based on Hess’ Law. partial pressures K p . Relationship between
(iii) Second Law of Thermodynamics and its Kp and Kc (derivation required);
significance, spontaneity of a chemical Characteristics of equilibrium constant;
change; Entropy, Free Energy. Inadequacy of Units for equilibrium constant; Simple
First Law and need for Second Law; Ideas calculations of equilibrium constant and
about reversible (recapitulation), spontaneous concentration.
and non-spontaneous processes The following examples should be considered
Meaning of entropy – derived from Second to show maximum yield of products:
Law – statement of Second Law in terms of - Synthesis of ammonia by Haber’s
entropy; Physical significance of entropy; process.
State function and not path function. Entropy
- The dissociation of dinitrogen tetra
change of the universe, reversible isothermal
oxide.
process and irreversible process.
Meaning of thermal death, Gibb’s free - Hydrolysis of simple esters.
energy of the system and Helmholtz free - The contact process for the manufacture
energy. Relationship between Gibb’s free of sulphuric acid.
energy and Helmholtz’s free energy.
Le Chatelier’s Principle. Statement and
Relationship between change in Gibb’s free explanation.
energy and equilibrium constant of a
chemical reaction. Defining the criteria for Factors affecting chemical and physical
spontaneity of a chemical change in terms of equilibria should be discussed in the light of
Gibb’s free energy. Le Chatelier’s principle.
Note: Numericals based on the First Law, - Change of concentration.
Second Law of Thermodynamics and Hess’
- Change of temperature.
Law.
(iv)Third Law of Thermodynamics – statement - Change of pressure.
only. - Effect of catalyst.
Self-explanatory. - Addition of inert gas.
6. Equilibrium (ii) Ionic equilibrium
(i) Chemical Equilibrium. Introduction, electrolyte (strong and weak),
Introduction of physical and chemical non-electrolyte, ionisation, degree of
equilibrium and its characteristics ionisation of pol ybas ic aci ds , acid
strength, concept of pH, pH indicators,
Dynamic nature of equilibrium, law of buffer solution, common ion effect (with
mass action, equilibrium constant and illustrative examples). Henderson equation,
factors affecting equilibrium. Le Chatelier's hydrolysis of salts, solubility and solubility
principle and its applications. product.
Irreversible and reversible reactions. Ostwald’s dilution law and its derivation.
Physical equilibrium: solid-liquid, liquid- Strength of acids and bases based on their
vapour, solid-vapour; Characteristics of dissociation constant. Problems based on the
Physical equilibrium. Ostwald’s dilution law.

6

Arrhenius, Brönsted-Lowry and Lewis 8. Organic Chemistry - Some Basic Principles
concept of acids and bases, multistage and Techniques
ionisation of acids and bases with examples. General introduction, classification and IUPAC
Ionic product of water – definition, pH, nomenclature of organic compounds and
pOH, pK w of solutions. isomerism.
pH indicators and their choice in titrimetry. Methods of purification, qualitative and
quantitative analysis. Electron displacement in a
Numericals on the above concepts. covalent bond: inductive effect, electromeric
Common ion effect – definition, examples effect, resonance and hyperconjugation.
(acetic acid and sodium acetate; ammonium Homolytic and heterolytic b o n d fission of a
hydroxide and ammonium chloride), covalent bond: free radicals, carbocations,
applications in salt analysis. carbanions, electrophiles and nucleophiles,
Salt hydrolysis – salts of strong acids and types of organic reactions.
weak bases, weak acids and strong bases, (i) Introduction to organic chemistry:
weak acids and weak bases and the pH Vital force theory, reason for separate study
formula of the solutions of these salts in of organic chemistry and its importance,
water with suitable examples. characteristics of carbon atoms (tetra
Buffer solutions: definition, examples, valency), Reasons for large number of
action; its interpretations based on Le organic compounds: catenation, isomerism
Chatelier’s principle. Henderson equation. and multiple bonding, etc.
Solubility product: definition and application (ii) Classification of organic compounds:
in qualitative salt analysis (Group II, III and (definition and examples): open chain, closed
chain, homocyclic, hetrocyclic, aromatic,
IV cations).
alicyclic compounds, homologous series and
Numericals on pH, buffer solutions, solubility its characteristics, functional groups.
and solubility product. (iii) IUPAC rules for naming organic compounds.
7. Redox Reactions Aliphatic, alicyclic and aromatic compounds.
Concept of oxidation and reduction, redox (iv) Definition and classification of isomerism:
reactions, oxidation number, change in Structural isomerism: definition,
oxidation number, balancing redox reactions (in classification, examples.
terms of loss and gain of electrons). Chain isomerism, Positional isomerism,
Applications of redox in various types of chemical Functional isomerism, Metamerism,
reactions. Tautomerism - examples for each of the
− Concept of oxidation and reduction in terms above.
of oxygen, hydrogen, electrons. Stereoisomerism: definition and
− Redox reactions – examples. classification, examples.
Geometrical isomerism: Definition.
− Oxidation number: rules for calculation, Conditions for compounds to exhibit
simple calculations of oxidation state in geometrical isomerism; types and examples,
molecules and ions like K 2 Cr 2 O 7 , S 2 O32 − , cis and trans, syn and anti. Examples.
etc. Optical isomerism: Definition, Nicol prism,
− Oxidation and reduction in terms of change plane polarised light. polarimeter. Method of
in oxidation number. measuring angle of rotation. Specific
rotation. Conditions for optical activity. d, l
− Balancing of redox reactions in acidic and form; External compensation, Internal
basic medium by oxidation number and ion-
electron method.

7

compensation, racemic mixture & meso form. General methods of preparation: from
Examples – lactic acid and tartaric acid. sodium salts of carboxylic acids
(v) Analysis of organic compounds: (decarboxylation and Kolbe’s electrolytic
method); from alcohols and alkyl halides
Detection of elements (qualitative analysis) (Wurtz reaction, Coreyhouse Synthesis).
such as carbon, hydrogen, nitrogen, From aldehydes and Grignard’s
halogens and sulphur should be considered Reagent.
by using Lassaigne’s test and reactions
involved in it. Physical and chemical properties of
alkanes.
(vi) Estimation of carbon, hydrogen, nitrogen,
halogens, sulphur and phosphorous: Physical properties: state, freezing point,
melting point, boiling point, density.
Estimation of carbon and hydrogen –
Leibig’s method. Chemical properties: combustibility,
reaction with chlorine (free radical
Estimation of nitrogen - Kjeldahl’s method. mechanism), reaction with oxygen in
Estimation of halogens sulphur and presence of catalyst (formation of
phosphorous - Carius method. Numericals alcohol, aldehyde, and carboxylic acid).
included. Experimental details required. Cyclisation, aromatisation, isomerisation
(vii)Types of chemical reactions and their and pyrolysis.
mechanisms. Uses of alkanes.
Substitution, addition, elimination reactions: (ii) Alkenes - Nomenclature, structure of
definition and examples. double bond (ethene), isomerism;
Homolytic and heterolytic fission – definition methods of preparation; physical
and examples. Free radicals, carbocation, properties, chemical properties;
carbanion (their reactivities and stabilities). addition of hydrogen, halogen, water,
hydrogen halides (Markownikoff's
Electrophiles and nucleophiles – definition addition and peroxide effect),
and examples (including neutral ozonolysis, oxidation, mechanism of
electrophiles and nucleophiles). electrophilic addition.
Inductive, electromeric, mesomeric effect and General methods of preparation –
hyperconjugation – definition, examples. dehydration of alcohols,
(viii) Free radicals and polar mechanisms dehydrohalogenation of alkyl halides
In terms of fission of the bonds and formation (from vicinal dihalides), Kolbe’s
of the new bonds including S N 1, S N 2, E 1 and electrolytic method and from alkynes.
E 2 mechanisms. Explain with relevant Physical Properties: State, freezing
examples and conditions. point, melting point, boiling point, dipole
moment, density.
9. Hydrocarbons Chemical properties - addition reactions
Classification of Hydrocarbons (hydrogen, halogens, hydrogen halides,
I. Aliphatic Hydrocarbons sulphuric acid, water).
Markownikoff’s rule and anti-
(i) Alkanes - Nomenclature, isomerism,
Markownikoff’s rule with mechanism and
conformation (methane and ethane),
examples.
physical properties, chemical properties
including free radical mechanism of Oxidation: complete combustion, hot
halogenation, combustion and pyrolysis. and cold alkaline KMnO 4 (Baeyer’s
reagent), ozonolysis.
Occurrence, conformation (Sawhorse
and Newman projections of ethane). Polymerisation.
Saytzeff’s rule and its application.

8

Uses of alkenes. - Electrophilic substitution reactions with
(iii)Alkynes - Nomenclature, structure of mechanism (halogenation, nitration,
triple bond (ethyne), methods of sulphonation).
preparation; physical properties, - Alkylation, acetylation – Friedel Crafts
chemical properties: acidic character of reaction.
alkynes, addition reactions - hydrogen, - Directive influence (o-, p-, and m-) of
halogens, hydrogen halides and water. substituents in electrophilic and
General methods of preparations of nucleophilic substitutions (with
alkynes. Manufacture of ethyne by mechanism).
calcium carbide and from natural gas.
- Oxidation: catalytic oxidation, reaction
Dehydrohalogenation and Kolbe’s
with ozone.
electrolytic method.
- Addition reactions with hydrogen,
Physical properties of alkynes: State of
existence, freezing point, melting point, chlorine, bromine.
boiling point, density. - Pyrolysis (formation of bi-phenyl).
Chemical properties of alkynes – Carcinogenicity and toxicity of benzene may be
addition reactions (hydrogen, halogens, discussed.
hydrogen halides and water), acidic Uses.
nature of alkynes, formation of
acetylides. PAPER II
Oxidation: complete combustion, hot
and cold alkaline KMnO 4 (Baeyer’s
PRACTICAL WORK- 15 Marks
reagent), ozonolysis. Candidates are required to complete the following
experiments:
Polymerisation.
Uses of alkynes. 1. Basic laboratory techniques:
Distinguishing test between Alkane, − Cutting a glass tube.
Alkene and Alkyne.
− Bending a glass tube.
II. Aromatic Hydrocarbons
Introduction, IUPAC nomenclature, − Drawing out a glass jet.
benzene: resonance, aromaticity, chemical − Boring a cork.
properties: mechanism of electrophilic
substitution. Nitration, sulphonation, 2. Titration: acid-base titration involving molarity.
halogenation, Friedel Crafts alkylation and Titrations involving:
acylation, directive influence of functional
group in monosubstituted benzene. • Sodium carbonate solution/ dil H 2 SO 4 or dil.
Carcinogenicity and toxicity. HCl using methyl orange indicator.
Structure: Resonance structures (Kekule’s) • NaOH or KOH solution/ dil H 2 SO 4 or dil.
of benzene. HCl using methyl orange indicator.
Benzene: Preparation from sodium benzoate • Calculations involving molarity,
and from phenol. concentration in grams L-1/ number of ions,
Physical properties: State of existence, water of crystallisation and percentage
freezing point, melting point, boiling point, purity.
density. NOTE: Calculation of molarity must be upto 4
Chemical properties: decimal places at least, in order to avoid error.

9

OBSERVATION TABLE Group V: Ba2+, Sr2+, Ca2+
S. No. (A) (B) (B – A) Group VI: Mg2+
Initial Final burette Difference NOTE:
burette reading (ml) • For wet test of anions, sodium carbonate
reading extract must be used (except for carbonate).
(ml)
(ml) • Chromyl chloride test not to be performed.
1 (Insoluble salts, such as lead sulphate, barium
2 sulphate, calcium sulphate, strontium sulphate
should not be given).
3
4. Preparation of inorganic compounds.
• Concordant reading is to be used for titre
value. Concordant reading is two consecutive (a) Preparation of potash alum/Mohr’s salt.
values which are exactly the same. Average (b) Preparation of crystalline FeSO 4 /CuSO 4 .
will not be accepted as titre value.
5. Paper Chromatography.
• The table is to be completed in ink only.
Preparation of chromatogram, separation of
Pencil is not to be used.
pigments from extracts of leaves and flowers/ink
• Overwriting will not be accepted in the mixtures; determination of R f value.
tabular column.
Observations: PROJECT WORK AND PRACTICAL
• Pipette size (should be same for all the FILE - 15 Marks
candidates at the centre):
Project Work – 10 Marks
• Titre value (concordant).
The candidate is to creatively execute one
3. Qualitative analysis: identification of single salt
project/assignment on a selected topic of Chemistry.
containing one anion and one cation:
Teachers may assign or students may choose any one
Anions: CO 3 2-, NO 2 -, S2-, SO 3 2-, SO 4 2-, NO 3 -, project of their choice. (Refer to the suggested topics
CH 3 COO-, Cl-, Br-, I-, C 2 O 4 2-, PO 4 3-. at the end of Class XII syllabus).
Cations: NH 4 +, Pb2+, Cu2+, Al3+, Fe3+, Zn2+, Mn2+ Suggested Evaluation criteria for Project Work:
Ni2+, Co2+, Ba2+, Sr2+, Ca2+, Mg2+.
 Introduction / purpose
Anions: Dilute acid group – CO 3 2-, NO 2 -, S2-,
SO 3 2-  Contents
Concentrated Acid Group – NO 3 -, Cl-, Br-, I-,  Analysis/ material aid (graph, data, structure, pie
CH 3 COO-. charts, histograms, diagrams, etc)
Special Group - SO 4 2-, PO 4 3-, C 2 O 4 2-.  Presentation
+
Cations: Group Zero: NH 4
 Bibliography
Group I: Pb2+
Group II : Cu2+, Pb2+ Practical File – 5 Marks
3+ 3+
Group III: Al , Fe Teachers are required to assess students on the basis
2+ 2+
Group IV: Zn , Mn , Ni , Co 2+ 2+ of the Chemistry Practical file maintained by them
during the academic year.

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