Aims: 1. To foster acquisition of knowledge and understanding of terms, concepts, facts, processes, techniques and principles relating to the subject of Chemistry. 2. To develop the ability to apply the knowledge of contents and principles of Chemistry in new or unfamiliar situations. 3. To develop skills in proper handling of apparatus and chemicals. 4. To develop an ability to appreciate achievements in the field of Chemistry and its role in nature and society. 5. To develop an interest in activities involving usage of the knowledge of Chemistry. 6. To develop a scientific attitude through the study of Physical Sciences. 7. To acquaint students with the emerging frontiers and interdisciplinary aspects of the subject. 8. To develop skills relevant to the discipline. 9. To apprise students with interface of Chemistry with other disciplines of Science, such as, Physics, Biology, Geology, Engineering, etc.
There will be two papers in the subject. Paper I: Theory- 3 hours Paper II: Practical – 3 hours Project Work Practical File … 70 marks …20 marks … 7 marks … 3 marks Main postulates of the theory. Its limitations. Modern atomic theory. Laws of chemical combinations: Law of conservation of mass. Law of definite proportion. Law of multiple proportion. Law of reciprocal proportion. Gay-Lussac’s law of gaseous volumes. Statement, explanation and simple problems based on these laws. (ii) Atomic and isotopic masses. The atomic mass unit is one of the experimentally determined unit. It is equal to 1/12 of the mass of the carbon 12 isotope. (iii) Chemical equivalents, volumetric calculations in terms of normality. C = 12.00 should be taken as a standard for expressing atomic masses. Equivalent weight expresses the combining capacity of the elements with the standard elements such as H, Cl, O, Ag, etc. Variable equivalent weight. Gram equivalent weights, relationship between gram equivalent weight, gram molecular weight and valency. Determination of equivalent weight of acids, alkalis, salts, oxidising and reducing agents. (experimental details not required). 128
PAPER I –THEORY – 70 Marks There will be one paper of 3 hours duration divided into 2 parts. Part I (20 marks) will consist of compulsory short answer questions, testing knowledge, application and skills relating to elementary/fundamental aspects of the entire syllabus. Part II (50 marks) will be divided into 3 Sections, A, B and C. Candidates are required to answer two out of three questions from Section A (each carrying 10 marks), two out of three questions from Section B (each carrying 5 marks) and two out of three questions from Section C (each carrying 10 marks). Therefore, a total of six questions are to be answered in Part II. SECTION A 1. Atoms and Molecules (i) The concept of atoms having fixed properties in explaining the laws of chemical combination. The study about the atoms. Dalton’s atomic theory:
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 formulae. Simple calculations on the above topics. (iv) Relative molecular mass and mole. The following methods may be considered for the determination of relative molecular masses for the gases: the molar volume method; Victor Meyer’s method (experimental details not required). Numerical problems based on the above method and Victor Meyer’s method. Mole concept, Avogadro’s number and numerical problems on mole concept. Gram molecular volume. (v) Chemical Reaction calculations based mass-volume and relationships. Self explanatory. 2. Atomic Structure (i) Electrons, Protons and Neutrons as fundamental particles, their charges and masses.
Concept of indivisibility of atom as proposed by Dalton does not exist. The atom consists of subatomic fundamental particles. Production of cathode rays and their properties. Production of anode rays and their properties. Chadwick’s experiment for the discovery of neutron and properties of neutron. (ii) Rutherford’s nuclear model based on the scattering experiment. Rutherford’s nuclear model of atom. Rutherford’s scattering experiment. Discovery of nucleus. Defects of Rutherford model. (iii) Bohr’s atomic model. 1. Postulates of Bohr’s theory – based on Planck’s quantum theory. 2. Numericals on Bohr’s atomic radii, velocity and energy of orbits (derivation not required). 129 – Stoichiometric on mass-mass, volume-volume
3. Defects in the Bohr’s Model. (iv) Atomic structure: wave mechanical model- a simple mathematical treatment. Quantum numbers; shape, size and orientation of s and p orbitals only. Hund’s rule of maximum multiplicity. Pauli’s exclusion principle, Aufbau principle, electronic configuration of elements in terms of s, p, d, f subshells. • Wave mechanical model – experimental verification of wave nature of electron. • de Broglie’s equation. Numericals. • Heisenberg’s Numericals. uncertainity principle.
• Quantum numbers – types of quantum numbers, information obtained in terms of distance of electron from the nucleus, energy of electron, number of electrons present in an orbit and an orbital. • Pauli’s exclusion principle. Shape, size and orientation of the s and p subshells. • Hund’s rule of maximum multiplicity. • Aufbau principle, (n+l) rule. • Electronic configuration of elements in terms of s, p, d, f subshells. 3. Periodic Table (i) Atomic number (Proton number) as the basis for classification of the elements in the Periodic Table.
IUPAC nomenclature for elements with Z> 100. Mendeleev’s periodic law, defects in the Mendeleev’s periodic table. Advantages and disadvantages. Modern periodic law (atomic number taken as the basis of classification of the elements). Extended and long form of periodic table. General characteristics of groups and periods. Division of periodic table as s, p, d and f blocks. (ii) Extra nuclear structure as the basis of periodicity. Some idea of the following: ionisation enthalpy, electron gain enthalpy, atomic radius, atomic volume, electronegativity, etc must be given. The periodicity of electronic structure leading to
the periodicity of elements e.g the relative ease of ionisation of elements. • Periodic properties such as valence electrons, atomic volume, atomic and ionic radii and their variation in groups and periods. • The idea of ionisation enthalpy, electron gain enthalpy and electronegativity must be given and their variation in groups and periods may be discussed. • The factors (atomic number, atomic volume and shielding effect, the number of electrons in the outermost orbit) which affect these periodic properties and their variation in groups and periods. (iii) Periodicity of elements with reference to s, p, d and f block elements. Classification of elements on the basis of s, p, d, f block elements and also on the basis of their complete and incomplete electron shells.
Study of the periodicity of properties mentioned in point (ii) in terms of s, p, d, f blocks and the governing factors in terms of the block characteristics. 4. Chemical Bonding Electrovalent Bond (i) Electrovalent or ionic bond e.g formation of NaCl, Li2O, MgO, CaO, MgF2, and Na2 S. Cause of chemical combination, Octet rule, types of chemical bonds. Electrovalent formation of NaCl, Li2O, MgO, CaO, MgF2, and Na2S. Properties of ionic compounds. Electron dot structure of the following ionic compounds: NaCl, Li2O, MgO, CaO, MgF2, and Na2S must be taught in detail. (ii) Factors influencing the formation of ionic bond, e.g electron gain enthalpy, ionisation enthalpy, lattice energy and electronegativity. The conditions necessary for the formation of ionic bonds such as: low ionisation enthalpy of metals. high electron gain enthalpy of non-metals. high lattice energy.
All these points must be discussed in detail. (iii) The relation between the ionic bonding and Periodic Table. The relationship between the formation of cations and anions of the atoms and their positions in the periodic table should be discussed. Correlate the periodic property and the position of the elements in the periodic table to show the ease of formation of anions and cations and electrovalent and covalent compounds. (iv) Variable electrovalency and its causes. Variable electrovalency; reasons for variable electrovalency i.e, due to inert electron pair effect, by using suitable examples. Covalent Bond (i) Covalent bond, sigma and pi bonds e.g. formation of ammonia, nitrogen, ethene, ethyne, and carbon dioxide.
Resonance. Definition of covalent bonding, conditions for formation of covalent bonds, types of covalent bonds i.e single, double and triple bonds. Sigma and pi bonds. H2, O2, N2. Classification of covalent bonds based on electronegativity of atoms – polar and non polar covalent bond, dipole moment, formation of CH4, H2O, NH3, ethane, ethene, ethyne and CO2, etc. and their electron dot structure or Lewis structure.
Characteristics of covalent compounds. Comparison in electrovalency and covalency. Resonance in simple inorganic molecules like ozone, carbon dioxide, carbonate ion and nitrate ion. (ii) Variable valency: chlorine exhibits the valency of 1,3,5 & 7 respectively. Variable valency, cause of variable covalency e.g. chlorine exhibits the valency 1, 3, 5 and 7 respectively. Discuss in terms of atomic structure. Variable covalency of phosphorus and sulphur may be discussed. Discuss in terms of atomic structure.
(iii) Deviation from Octet rule and Fajan’s rules. Definition of Octet rule. Failure of Octet rule, due to either incomplete octet or exceeding of Octet with suitable examples. Fajan’s rules: Statements. Conditions for electrovalency and covalency must be discussed. Polar and non polar bonds should be correlated with Fajan’s rules.
(viii) Molecular orbital theory, Qualitative treatment of homonuclear diatomic molecules of first two periods. Energy level diagrams, bonding, antibonding molecular orbitals, bond order, paramagnetism of O2 molecule. Relative stabilities of O2, O2-, O2- – , O2+, O2++ Self-explanatory. 5. The Gaseous State (i) The gas laws, qualitatively. kinetic theory treated
(iv) Co-ordinate or dative covalent bond, e.g. formation of oxy-acids of chlorine. Co-ordinate or dative covalent bonding: definition, formation of hypochlorous acid, chloric acid, perchloric acid, ammonium ion, hydronium ion, nitric acid, ozone – structural formulae of the above molecules based on co-ordinate bonding. (v) Hydrogen bonding: its essential requirements, 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.
(vi) Metallic bonding, Van der Waals’ forces. Metallic bonding – Electron sea model and band model. Explanation of metallic properties in terms of metallic bonding. Van der Waals’ forces and its types. (vii)Valence Shell Electron Pair Repulsion Theory; Hybridisation and shapes of molecules: hybridisation involving s, p and d orbitals only; sigma and pi bonds. Concept of electron-pair repulsion and shapes of molecules taking methane, ammonia and water as examples. Hybridisation and molecular shapes – definition, hybridization of orbitals involving s, p and d orbitals (examples: ethane, ethene, ethyne, PCl5 and SF6).
Characteristics of gases, comparison between solid, liquid and gas. Properties of gases on the basis of kinetic theory of gases. Laws of gases – Boyle’s Law, Charles’ Law, Absolute Temperature, Pressure Temperature Law, Avogadro’s Law. Simple numerical problems based on the above laws. Postulates of Kinetic Theory must be discussed to explain gas laws. (ii) PV = nRT or PV= (w/M)RT and the application of this equation of state. Ideal gas equation PV = nRT; its application in calculation of relative molecular mass and in the calculation of the value of R.
(iii) Non ideal behaviour of gases and Van der Waals’ equation. Non ideal behaviour of gases i.e. deviation from gas laws 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. The pressure correction and volume correction may be explained. (iv) Dalton’s law, the Avogadro constant, the mole, Graham’s law of diffusion, simple numerical problems on the above. • Dalton’s Law of partial pressure. • Application of Dalton’s Law. • Numerical problems based on the above law. • Avogadro’s constant. • Relationship between the mole and Avogadro number. Graham’s Law of diffusion and its application. • Simple numerical problems on the above.
6. Colloidal Solutions Preparation and properties of colloids, both lyophilic and lyophobic colloids. Precipitation as evidence that the colloidal particles are charged. Idea of gold number is required, but application of gold number is not required. The importance of large surface area in adsorption should also be appreciated. • • • • • Thomas Graham classified the substances as crystalloid and colloid. Classification of substances on the basis of the particle size i.e. true solution, sol and suspension. Colloidal system is heterogeneous. Lyophilic and lyophobic colloids. Classification of colloidal solutions as micro, macro and associated colloids. Preparation of lyophilic colloids.
Preparation of lyophobic colloids by colloid mill, peptisation, Bredig’s arc method (procedural details not required) by oxidation, reduction, double decomposition and exchange of solvent method should be discussed. Purification of colloids (dialysis, ultra filtration, and ultracentrifugation). Properties of colloidal solutions such as Brownian movement, Tyndall effect, coagulation and protection (protective colloids), should be discussed. Gold number and Hardy Schulze rule. Application of colloids in life. Electrophoresis (movement of dispersed phase). Emulsions, surfactants, micelles (only definition and examples).
8. Chemical Energetics (i) Introduction. (a) Scope of thermodynamics- characteristics of thermodynamics. (b) Types of system – ideal system, real system, isolated system, closed system, open system. (c) Meaning of surrounding. (d) Properties of the system: macroscopic, intensive and extensive properties of the system. (e) State of the system. (f) Main processes the system undergoes: reversible, irreversible, adiabatic, isothermal, isobaric, isochoric, cyclic. (g) Meaning of thermodynamic equilibrium. (h) Meaning of thermodynamic process. (ii) First law of Thermodynamics mathematical statement. and its
(a) Idea of conservation of energy – total energy of the system and the surrounding. (b) Meaning of internal energy of the system and change in internal energy of the system. (c) Meaning of work done by the system and by the surrounding at constant temperature. (d) Meaning of heat absorbed by the system and by the surrounding at constant temperature. (e) The sign convention for change in internal energy, heat given out or gained, work done by the system or by the surrounding. (f) State function and path function- meaning with examples. (g) Internal energy change, work done and heat absorbed in a cyclic process. (h) Internal energy change in an isolated system and in non isolated system.
7. Chemical Kinetics Rate of a chemical reaction, basic idea of order and molecularity of a reaction. Rate of a chemical reaction; Relation between order and the stoichiometric coefficients in the balanced equation; Meaning of molecularity. Differences between the order and molecularity of the reaction.
Physical significance of entropy State function and not path function. Relationship between adiabatic change and entropy. Entropy change of the universe and a reversible isothermal process. Entropy change of the universe and irreversible process. Meaning of thermal death. Meaning of energy content and work content (free energy) of the system – thermodynamic quantity – state function. Types of work and meaning of the two types of work. Meaning of Helmholtz’s Free energy and Gibb’s free energy and the change in Gibb’s and Helmholtz’s free energy.
Relationship between Gibb’s free energy and Helmholtz’s free energy. Simple calculation on the change in Gibb’s free energy and Helmholtz’s free energy. Relationship between change in Gibb’s free energy and equilibrium constant of a chemical reaction. Change in Gibb’s free energy in reversible, irreversible, isobaric and isochoric processes. Based on change in Gibb’s free energy, defining the criteria for the spontaneity of a change in terms of entropy and enthalpy; defining the limits for reversible chemical reactions.
(k) Chemical change and internal energy. (l) Need for enthalpy – constant pressure or open vessel processes. (m) Enthalpy a thermodynamic property – state function. (n) Mathematical form constant pressure.
(iii) Ideas about Heat, Work and Energy. Heat – the energy in transit. Condition for the transfer of heat. Limitation in conversion of heat into work. Condition at which heat transfer ceases. Unit of heat. Meaning of energy – capacity to do work. Meaning of work – intensity factor and capacity factor. Types of work. Mathematical form of reversible work. Mathematical form of irreversible work. Difference between the reversible and irreversible work done – graphically. Adiabatic reversible expansion. Relationship between Cv and internal energy change.