Calcutta University Chemistry General Syllabus

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Calcutta University B.Sc. Chemistry honours course syllabus
CONTENTS
Chem. Hons. Course: module names and distribution p 2
Course contents
Hons. Course: Year 1 p 3-7
Hons. Course: Year 2 p 8-11
Hons. Course: Year 3 p 12-19
Chem. Gen. Course: module names and distribution p 20
Course contents
Gen. Course: Year 1 p 21-23
Gen. Course: Year 2 p 24-26
Gen. Course: Year 3 p 27-29
Appendices
Appendix 1: Paper-wise distribution of modules (Hons. & General) p 30
Appendix 2: Instructions about Theoretical Examinations p 31
Appendix 3: General instructions about Practical Examinations p 32
Appendix 4: Specific instructions about Practical Examinations p 33-36
Appendix 5: Recommended list of books p 37-38
Chemistry Hons: Syllabus Scheme in modular form
Course names and distribution
PART – I (Year 1), total marks = 200 (Theory = 150, Practical = 50)
CHT 11a, 11b, each 25 marks, Theory
CHT 12a, 12b, each 25 marks, Theory
CHT 13a, 13b, each 25 marks, Theory
CHP 14a+14b, 50 marks, Practical
PART – II (Year 2), total marks = 200 (Theory = 150, Practical = 50)
CHT 21a, 21b, each 25 marks, Theory
CHT 22a, 22b, each 25 marks, Theory
CHT 23a, 23b, each 25 marks, Theory
CHP 24a+24b, 50 marks, Practical
PART – III (Year 3), total marks = 400 (Theory = 250, Practical = 150)
CHT 31a, 31b, 31c, 31d, each 25 marks, Theory
CHT 32a, 32b, 32c, each 25 marks, Theory
CHT 33a, 33b, 33c, each 25 marks, Theory
CHP 34a, 34b, 25 and 50 marks, Practical
CHP 35a, 35b, 25 and 50 marks, Practical
Abbreviations:
CHP: Chem Hons Practical; CHT: Chem Hons Theory
First digit refers to year, second to paper.
Each CHT Exam = 1 hr for 25 marks, 2 hr for 50 marks, etc.
Each CHP Exam = 2-3 hr for 25 marks, 4 hr for 50 marks on each day
Notes:
1. Each Theory module of 25 marks contains units I (marks = 15) and II (marks = 10).
2. Number of class hours = 25-35 for a 25-mark Theory module, 70-80 for a 25-mark
Practical module.
Chemistry Hons: Course Description
Year 1
PART – I
CHT 11a
Unit-I. Radioactivity and Atomic Structure
Nuclear stability and nuclear binding energy. Nuclear forces: meson exchange
theory. Nuclear models (elementary idea): Concept of nuclear quantum number, magic
numbers. Nuclear Reactions: Artificial radioactivity, transmutation of elements, fission,
fusion and spallation. Nuclear energy and power generation. Separation and uses of
isotopes. Radio chemical methods: principles of determination of age of rocks and
minerals, radio carbon dating, hazards of radiation and safety measures.
Bohr’s theory to hydrogen-like atoms and ions; spectrum of hydrogen atom.
Quantum numbers. Introduction to the concept of atomic orbitals; shapes, radial and
angular probability diagrams of s, p and d orbitals (qualitative idea). Many electron atoms
and ions: Pauli’s exclusion principle, Hund’s rule, exchange energy, Aufbau principle
and its limitation. Electronic energy level diagram and electronic configurations of
hydrogen-like and polyelectronic atoms and ions. Term symbols of atoms and ions for
atomic numbers < 30.
Unit-II. Chemical periodicity I
Periodic table, group trends and periodic trends in physical properties.
Classification of elements on the basis of electronic configuration. Modern IUPAC
Periodic table. General characteristic of s, p, d and f block elements. Position of hydrogen
and noble gases in the periodic table.
Effective nuclear charges, screening effects, Slater’s rules, atomic radii, ionic
radii (Pauling’s univalent), covalent radii. Ionization potential, electron affinity and
electronegativity (Pauling’s, Mulliken’s and Allred-Rochow’s scales) and factors
influencing these properties. Inert pair effect. Group trends and periodic trends in these
properties in respect of s-, p- and d-block elements.
CHT 11b
Unit-I. Chemical Bonding and structure
Ionic bonding: Size effects, radius ratio rules and their limitations. Packing of ions
in crystals, lattice energy, Born-lande equation and its applications, Born-Haber cycle and
its applications. Solvation energy, polarizing power and polarizability, ionic potential,
Fazan’s rules. Defects in solids (elemementary idea).
Covalent bonding: Lewis structures, formal charge. Valence Bond Theory,
directional character of covalent bonds, hybridizations, equivalent and non-equivalent
hybrid orbitals, Bent’s rule, VSEPR theory, shapes of molecules and ions containing lone
pairs and bond pairs (examples from main groups chemistry), Partial ionic Character of
covalent bonds, bond moment, dipole moment and electronegativity differences. Concept
of resonance, resonance energy, resonance structures
Unit-II. Acid-Base reactions
Acid-Base concept: Arrhenius concept, theory of solvent system (in H2O, NH3,
SO2 and HF), Bronsted-Lowry’s concept, relative strength of acids, Pauling rules.
Amphoterism. Lux-Flood concept, Lewis concept. Superacids, HSAB principle. Acidbase
equilibria in aqueous solution and pH. Acid-base neutralisation curves; indicator,
choice of indicators.
CHT 12a
Unit I. Acyclic stereochemistry
Representation of molecules in saw horse, Fischer, flying-wedge and Newman
formulae and their inter translations, symmetry elements, molecular chirality.
Configuration: stereogenic units i) stereocentres: systems involving 1, 2, 3
centres, stereogenicity, chirotopicity. pseudoasymmetric (D/L and R/S descriptor,
threo/erythro and syn/anti nomenclatures (for aldols) ii) stereoaxis: chiral axis in allenes
& biphenyls, R/S descriptor; cis/trans, syn/anti, E/Z descriptors (for C=C, C=N).
Optical activity of chiral compounds: specific rotation, optical purity (enantiomeric
excess), racemic compounds, racemisation (through cationic and anionic and radical
intermediates), resolution of acids, bases and alcohols via diastereomeric salt formation.
Topicity of ligands and faces (elementary idea): Pro-R, Pro-S and Re /Si descriptors.
Conformation: Conformational nomenclature, eclipsed, staggered, gauche and
anti; dihedral angle, torsion angle, energy barrier of rotation, relative stability of
conformers on the basis of steric effect, dipole-dipole interaction, H-bonding;
conformational analysis of ethane, propane, n-butane, haloethane, 1,2-haloethane, 1,2-
glycol, 1,2-halohydrin; invertomerism of trialkylamines.
Unit II. Bonding and physical properties
Valence bond theory: concept of hybridisation, resonance (including
hyperconjugation), orbital pictures of bonding (sp3, sp2, sp: C-C, C-N & C-O system).
Inductive effect, bond polarization and bond polarizability, steric effect, steric inhibition
of resonance.
MO theory: sketch and energy levels of MOs of i) acyclic p orbital system (C=C,
conjugated diene and allyl systems) ii) cyclic p orbital system (neutral system: [4], [6]
annulenes; charged system: 3,4,5-ring system); Frost diagram, Huckel’s rules for
aromaticity & antiaromaticity; homoaromaticity.
Physical properties: bond distance, bond angles, mp/bp & dipole moment in terms
of structure and bonding (covalent & non covalent). Heat of hydrogenation and heat of
combustion.
CHT 12b
Unit I. General treatment of reaction mechanism
Mechanistic classification: ionic, radical and pericyclic; heterolytic bond cleavage
and heterogenic bond formation, homolytic bond cleavage and homogenic bond
formation; representation of mechanistic steps using arrow formalism.
Reactive intermediates: carbocations (cabenium and carbonium ions), carbanions,
carbon radicals, carbenes – structure using orbital picture, electrophilic/nucleophilic
behaviour, stability, generation and fate (elementary idea)
Reaction thermodynamics: free energy and equilibrium, enthalpy and entropy
factor, intermolecular & intramolecular reactions. Application of thermodynamic
principles in tautomeric equilibria [keto-enol tautomerism, composition of the
equilibrium in different systems (simple carbonyl, 1,3 and 1,2- dicarbonyl systems,
phenols and related system), substituent and solvent effect].
Concept of acids and bases: effect of structure, substituent and solvent on acidity
and basicity.
Reaction kinetics: transition state theory, rate const and free energy of activation,
free energy profiles for one step and two step reactions, catalyzed reactions, kinetic
control and thermodynamic control of reactions, isotope effect , primary kinetic isotopic
effect (kH/kD), principle of microscopic reversibility, Hammond postulate.
Unit II. Nucleophilic substitution reactions
Sustitution at sp3 centre - Mechanism: SN1, SN2, SN2’, SNi mechanisms, effect of
solvent, substrate structure, leaving group, nucleophiles including ambident nucleophiles
(cyanide & nitrite) substitution involving NGP; relative rate & stereochemical features
[systems: alkyl halides, allyl halides, alcohols, ethers, epoxides].
Halogenation of alkanes and carbonyls.
Substitution at sp2 carbon (carbonyl system) - Mechanism: BAC2, AAC2, AAC1, AAL1 (in
connection to acid and ester). Systems: amides, anhydrides & acyl halides [formation and
hydrolysis]
CHT 13a
Unit I. Kinetic theory and the gaseous state
Concept of pressure and temperature. Nature of distribution of velocities in one,
two and three dimensions. Maxwell's distribution of speeds. Kinetic energy distribution
in one, two and three dimensions, calculations of average, root mean square and most
probable values in each case; calculation of number of molecules having energy ≥ ε,
Principle of equipartition of energy and its application to calculate the classical limit of
molar heat capacity of gases.
Collision of gas molecules; collision diameter; collision number and mean free
path; frequency of binary collisions (similar and different molecules); wall collision and
rate of effusion.
Deviation of gases from ideal behaviour; compressibility factor; Andrew's and
Amagot's plots; van der Waals equation and its characteristic features. Existence of
critical state. Critical constants in terms of van der Waals constants. Law of
corresponding state and significance of second virial coefficient. Boyle temperature.
Intermolecular forces (Debye, Keesom and London interactions; Lennard-Jones potential,
elementary idea).
Unit II. Thermodynamics – I
Importance and scope, definitions of system and surroundings; type of systems
(isolated, closed and open). Extensive and intensive properties. Steady state and
equilibrium state. Concept of thermal equilibrium and the zeroth-law of thermodynamics.
Thermodynamic coordinates, state of a system, equation of state, state functions and path
functions. Partial derivatives and cyclic rule. Concept of heat and work (IUPAC
convention). Graphical explanation of work done during expansion and compression of
an ideal gas. Reversible and irreversible processes and work done.
First law of thermodynamics, internal energy (U) as a state function. Enthalpy as
a state function. Heat changes at constant volume and constant pressure; relation between
CP and CV using ideal gas and van der Waals equations. Joule's experiment and its
consequence. Explanation of term (δU/δV)T. Isothermal and adiabatic processes.
Thermochemistry: heat changes during physicochemical processes at constant
pressure/volume. Kirchoff's relations. Bond dissociation energies. Changes of
thermodynamic properties in different chemical changes.
CHT 13b
Unit I. Thermodynamics – II
Second law of thermodynamics – need for a Second law. Concept of heat
reservoirs and heat engines. Kelvin – Planck and Clausius statements and equivalence of
the two statements with entropic formulation. Carnot cycle and refrigerator. Carnot's
theorem; thermodynamic scale of temperature.
Physical concept of entropy. Entropy as a measure of the microscopic but not
macroscopic disorder. Values of §dQ/T and Clausius inequality. Entropy change of
systems and surroundings for various processes and transformations. Entropy change
during the isothermal mixing of ideal gases. Entropy and unavailable work. Auxiliary
state functions (G and A) and their variation with T, P and V. Criteria for spontaneity and
equilibrium.
Thermodynamic relations: Maxwell's relations, thermodynamic equation of state.
Gibbs- Helmholtz equation, Joule-Thomson experiment and its consequences; inversion
temperature. Joule-Thomson coefficient for a van der Waals gas. General heat capacity
relations.
Unit II. Chemical kinetics
Introduction of reaction rate in terms of extent of reaction; rate constants, order
and molecularity of reactions. Reactions of zero order, first order, second order and
fractional order. Pseudo first order reactions (example using acid catalyzed hydrolysis of
methyl acetate). Determination of order of a reaction by half-life and differential method.
Rate-determining and steady-state approximation – explanation with suitable examples.
Opposing reactions, consecutive reactions and parallel reactions (with explanation
of kinetic and thermodynamic control of products; all steps first order).
Temperature dependence of rate constant: Arrhenius equation, energy of
activation. Homogeneous catalysis with reference to acid-base catalysis. Enzyme
catalysis: Michaelis-Menten equation, turn-over number.
CHP 14a+14b
Qualitative inorganic analysis of mixtures containing not more than 4 radicals from the
following:
Cation Radicals: Na+, K+, Ca+2, Sr+2, Ba+2, Al+3, Cr+3, Mn+2, Fe+3, Co+3, Ni+3, Cu+2, Zn+2.
Anion Radicals: F-, Cl-, Br-, BrO3
-, I-, SCN-, S2-, SO4
2-, S2O3
2-, NO3
-, NO2
-, PO4
3-, BO3
3-,
CrO4
2-/ Cr2O7
2-, Fe(CN)6
4-, Fe(CN)6
3-.
Insoluble Materials: Al2O3, Fe2O3, Cr2O3, SnO2, SrSO4, BaSO4, CaF2.
Experiment A: Preliminary Tests for acid and basic radicals in given samples.
Experiment B: Wet tests for Acid and Basic radicals in given samples.
Experiment C: Confirmatory tests.
Notes:
At least 10 unknown samples are to be analyzed by each student during the laboratory
session. Oxide, hydroxide, carbonate and bicarbonate should not be reported as radicals.
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