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Chapterwise Topicwise Solved Papers 2021-1979 IITJEE JEE Main & Advanced Chemistry Ranjeet Shahi Arihant Prakashan (Series), Meerut Arihant Prakashan (Series), Meerut All Rights Reserved © Author Administrative & Production Offices Regd Office ‘Ramchhaya’ 4577/15, Agarwal Road, Darya Ganj, New Delhi -110002 Tele: 011- 47630600, 43518550 Head Office Kalindi, TP Nagar, Meerut (UP) - 250002, Tel: 0121-7156203, 7156204 Sales & Support Offices Agra, Ahmedabad, Bengaluru, Bareilly, Chennai, Delhi, Guwahati, Hyderabad, Jaipur, Jhansi, Kolkata, Lucknow, Nagpur & Pune ISBN 978-93-25796-14-0 PO No : TXT-XX-XXXXXXX-X-XX Published by Arihant Publications (India) Ltd For further information about the books published by Arihant, log on to www.arihantbooks.com or e-mail at info@arihantbooks.com Follow us on CONTENTS 1-22 19 Extraction of Metals 282-293 Atomic Structure 23-40 20 Qualitative Analysis 294-306 Periodic Classification and Periodic Properties 21 Organic Chemistry Basics 307-331 41-47 22 Hydrocarbons 332-349 Chemical Bonding 48-66 23 Alkyl Halides 350-363 States of Matter 67-83 24 Alcohols and Ethers 364-377 25 Aldehydes and Ketones 378-396 26 Carboxylic Acids and their Derivatives 397-412 Some Basic Concepts of Chemistry Chemical and Ionic Equilibrium 84-108 Thermodynamics and Thermochemistry 109-129 Solid State 130-139 Solutions and Colligative Properties 140-155 27 Aliphatic Compounds Containing Nitrogen 413-422 10 Electrochemistry 156-177 28 Benzene and Alkyl Benzene 423-440 11 Chemical Kinetics 178-195 12 Nuclear Chemistry 196-199 29 Aromatic Compounds Containing Nitrogen 441-457 13 Surface Chemistry 200-206 30 Aryl Halides and Phenols 458-470 14 s-Block Elements 207-217 31 Aromatic Aldehydes, Ketones and Acids 471-484 15 p-Block Elements-I 218-227 16 p-Block Elements-II 228-248 32 Biomolecules and Chemistry in Everyday Life 485-501 33 Environmental Chemistry 502-504 17 Transition and Inner-Transition Elements 249-258 18 Coordination Compounds 259-281 JEE Advanced Solved Paper 2021 1-16 SYLLABUS JEE MAIN Section A : PHYSICAL CHEMISTRY UNIT I Some Basic Concepts in Chemistry Matter and its nature, Dalton's atomic theory; Concept of atom, molecule, element and compound; Physical quantities and their measurements in Chemistry, precision and accuracy, significant figures, S.I Units, dimensional analysis; Laws of chemical combination; Atomic and molecular masses, mole concept, molar mass, percentage composition, empirical and molecular formulae; Chemical equations and stoichiometry UNIT II States of Matter Classification of matter into solid, liquid and gaseous states Gaseous State Measurable properties of gases; Gas laws - Boyle's law, Charle's law, Graham's law of diffusion, Avogadro's law, Dalton's law of partial pressure; Concept of Absolute scale of temperature; Ideal gas equation, Kinetic theory of gases (only postulates); Concept of average, root mean square and most probable velocities; Real gases, deviation from Ideal behaviour, compressibility factor, van der Waals' equation, liquefaction of gases, critical constants Liquid State Properties of liquids - vapour pressure, viscosity and surface tension and effect of temperature on them (qualitative treatment only) Solid State Classification of solids: molecular, ionic, covalent and metallic solids, amorphous and crystalline solids (elementary idea); Bragg's Law and its applications, Unit cell and lattices, packing in solids (fcc, bcc and hcp lattices), voids, calculations involving unit cell parameters, imperfection in solids; electrical, magnetic and dielectric properties UNIT III Atomic Structure Discovery of sub-atomic particles (electron, proton and neutron); Thomson and Rutherford atomic models and their limitations; Nature of electromagnetic radiation, photoelectric effect; spectrum of hydrogen atom, Bohr model of hydrogen atom - its postulates, derivation of the relations for energy of the electron and radii of the different orbits, limitations of Bohr's model; dual nature of matter, de-Broglie's relationship, Heisenberg uncertainty principle Elementary ideas of quantum mechanics, quantum mechanical model of atom, its important features, ψ and ψ2, concept of atomic orbitals as one electron wave functions; Variation of ψ and ψ2 with r for 1s and 2s orbitals; various quantum numbers (principal, angular momentum and magnetic quantum numbers) and their significance; shapes of s, p and d - orbitals, electron spin and spin quantum number; rules for filling electrons in orbitals – aufbau principle, Pauli's exclusion principle and Hund's rule, electronic configuration of elements, extra stability of half-filled and completely filled orbitals UNIT IV Chemical Bonding and Molecular Structure Kossel Lewis approach to chemical bond formation, concept of ionic and covalent bonds Ionic Bonding Formation of ionic bonds, factors affecting the formation of ionic bonds; calculation of lattice enthalpy Covalent Bonding Concept of electronegativity, Fajan's rule, dipole moment; Valence Shell Electron Pair Repulsion (VSEPR) theory and shapes of simple molecules Quantum mechanical approach to covalent bonding Valence bond theory - Its important features, concept of hybridization involving s, p and d orbitals; Resonance Molecular Orbital Theory Its important features, LCAOs, types of molecular orbitals (bonding, antibonding), sigma and pi-bonds, molecular orbital electronic configurations of homonuclear diatomic molecules, concept of bond order, bond length and bond energy Elementary idea of metallic bonding Hydrogen bonding and its applications UNIT V Chemical Thermodynamics Fundamentals of thermodynamics System and surroundings, extensive and intensive properties, state functions, types of processes First law of thermodynamics Concept of work, heat internal energy and enthalpy, heat capacity, molar heat capacity, Hess's law of constant heat summation; Enthalpies of bond dissociation, combustion, formation, atomization, sublimation, phase transition, hydration, ionization and solution Second law of thermodynamics Spontaneity of processes; ΔS of the universe and ΔG of the system as criteria for spontaneity, ΔGo (Standard Gibb's energy change) and equilibrium constant UNIT VI Solutions Different methods for expressing concentration of solution - molality, molarity, mole fraction, percentage (by volume and mass both), vapour pressure of solutions and Raoult's Law - Ideal and non-ideal solutions, vapour pressure - composition plots for ideal and non-ideal solutions Colligative properties of dilute solutions - relative lowering of vapour pressure, depression of freezing point, elevation of boiling point and osmotic pressure; Determination of molecular mass using colligative properties; Abnormal value of molar mass, van't Hoff factor and its significance UNIT VII Equilibrium Meaning of equilibrium, concept of dynamic equilibrium Equilibria involving physical processes Solid -liquid, liquid - gas and solid - gas equilibria, Henry's law, general characteristics of equilibrium involving physical processes Equilibria involving chemical processes Law of chemical equilibrium, equilibrium constants (K and K) and their significance, significance of ΔG and ΔGo in chemical equilibria, factors affecting equilibrium concentration, pressure, temperature, effect of catalyst; Le -Chatelier's principle Ionic equilibrium Weak and strong electrolytes, ionization of electrolytes, various concepts of acids and bases (Arrhenius, Bronsted Lowry and Lewis) and their ionization, acid-base equilibria (including multistage ionization) and ionization constants, ionization of water, pH scale, common ion effect, hydrolysis of salts and pH of their solutions, solubility of sparingly soluble salts and solubility products, buffer solutions UNIT VIII Redox Reactions and Electrochemistry Electronic concepts of oxidation and reduction, redox reactions, oxidation number, rules for assigning oxidation number, balancing of redox reactions Eectrolytic and metallic conduction, conductance in electrolytic solutions, specific and molar conductivities and their variation with concentration: Kohlrausch's law and its applications temperature on rate of reactions - Arrhenius theory, activation energy and its calculation, collision theory of bimolecular gaseous reactions (no derivation) Electrochemical cells - Electrolytic and Galvanic cells, different types of electrodes, electrode potentials including standard electrode potential, half - cell and cell reactions, emf of a Galvanic cell and its measurement; Nernst equation and its applications; Relationship between cell potential and Gibbs' energy change; Dry cell and lead accumulator; Fuel cells; Corrosion and its prevention UNIT X Surface Chemistry Adsorption - Physisorption and chemisorption and their characteristics, factors affecting adsorption of gases on solidsFreundlich and Langmuir adsorption isotherms, adsorption from solutions Catalysis Homogeneous and heterogeneous, activity and selectivity of solid catalysts, enzyme catalysis and its mechanism Colloidal state distinction among true solutions, colloids and suspensions, classification of colloids - lyophilic, lyophobic; multi molecular, macromole-cular and associated colloids (micelles), preparation and properties of colloids Tyndall effect, Brownian movement, electrophoresis, dialysis, coagulation and flocculation; Emulsions and their characteristics UNIT IX Chemical Kinetics Rate of a chemical reaction, factors affecting the rate of reactions concentration, temperature, pressure and catalyst; elementary and complex reactions, order and molecularity of reactions, rate law, rate constant and its units, differential and integral forms of zero and first order reactions, their characteristics and half - lives, effect of Section B : INORGANIC CHEMISTRY UNIT XI Classification of Elements and Periodicity in Properties Periodic Law and Present Form of the Periodic Table, s, p, d and f Block Elements, Periodic Trends in Properties of Elementsatomic and Ionic Radii, Ionization Enthalpy, Electron Gain Enthalpy, Valence, Oxidation States and Chemical Reactivity UNIT XII General Principles and Processes of Isolation of Metals Modes of occurrence of elements in nature, minerals, ores; steps involved in the extraction of metals - concentration, reduction (chemical and electrolytic methods) and refining with special reference to the extraction of Al, Cu, Zn and Fe; Thermodynamic and electrochemical principles involved in the extraction of metals UNIT XIII Hydrogen Position of hydrogen in periodic table, isotopes, preparation, properties and uses of hydrogen; physical and chemical properties of water and heavy water; Structure, preparation, reactions and uses of hydrogen peroxide; Classification of hydrides ionic, covalent and interstitial; Hydrogen as a fuel UNIT XIV s - Block Elements (Alkali and Alkaline Earth Metals) Group and Elements General introduction, electronic configuration and general trends in physical and chemical properties of elements, anomalous properties of the first element of each group, diagonal relationships Preparation and properties of some important compounds - sodium carbonate, sodium chloride, sodium hydroxide and sodium hydrogen carbonate; Industrial uses of lime, limestone, Plaster of Paris and cement; Biological significance of Na, K, Mg and Ca UNIT XV p - Block Elements Group 13 to Group 18 Elements General Introduction Electronic configuration and general trends in physical and chemical properties of elements across the periods and down the groups; unique behaviour of the first element in each group.Group wise study of the p – block elements Group 13 Preparation, properties and uses of boron and aluminium; structure, properties and uses of borax, boric acid, diborane, boron trifluoride, aluminium chloride and alums Group 14 Tendency for catenation; Structure, properties and uses of allotropes and oxides of carbon, silicon tetrachloride, silicates, zeolites and silicones Group 15 Properties and uses of nitrogen and phosphorus; Allotrophic forms of phosphorus; Preparation, properties, structure and uses of ammonia nitric acid, phosphine and phosphorus halides,(PCl3, PCl5); Structures of oxides and oxoacids of nitrogen and phosphorus Group 16 Preparation, properties, structures and uses of dioxygen and ozone; Allotropic forms of sulphur; Preparation, properties, structures and uses of sulphur dioxide, sulphuric acid (including its industrial preparation); Structures of oxoacids of sulphur Group 17 Preparation, properties and uses of chlorine and hydrochloric acid; Trends in the acidic nature of hydrogen halides; Structures of Interhalogen compounds and oxides and oxoacids of halogens Group 18 Occurrence and uses of noble gases; Structures of fluorides and oxides of xenon UNIT XVI d–and f–Block Elements Transition Elements General introduction, electronic configuration, occurrence and characteristics, general trends in properties of the first row transition elements - physical properties, ionization enthalpy, oxidation states, atomic radii, colour, catalytic behaviour, magnetic properties, complex formation, interstitial compounds, alloy formation; Preparation, properties and uses of K2 Cr2 O7 and KMnO4 Inner Transition Elements Lanthanoids - Electronic configuration, oxidation states, chemical reactivity and lanthanoid contraction Actinoids - Electronic configuration and oxidation states UNIT XVII Coordination Compounds Introduction to coordination compounds, Werner's theory; ligands, coordination number, denticity, chelation; IUPAC nomenclature of mononuclear coordination compounds, isomerism; Bonding Valence bond approach and basic ideas of Crystal field theory, colour and magnetic properties; importance of coordination compounds (in qualitative analysis, extraction of metals and in biological systems) UNIT XVIII Environmental Chemistry Environmental pollution Atmospheric, water and soil Atmospheric pollution - Tropospheric and stratospheric Tropospheric pollutants Gaseous pollutants Oxides of carbon, nitrogen and sulphur, hydrocarbons; their sources, harmful effects and prevention; Green house effect and Global warming; Acid rain; Particulate pollutants Smoke, dust, smog, fumes, mist; their sources, harmful effects and prevention Stratospheric pollution Formation and breakdown of ozone, depletion of ozone layer - its mechanism and effects Water pollution Major pollutants such as, pathogens, organic wastes and chemical pollutants their harmful effects and prevention Soil pollution Major pollutants such as: Pesticides (insecticides, herbicides and fungicides), their harmful effects and prevention Strategies to control environmental pollution Section C : ORGANIC CHEMISTRY UNIT XIX Purification & Characterisation of Organic Compounds Purification Crystallisation, sublimation, distillation, differential extraction and chromatography principles and their applications Qualitative analysis Detection of nitrogen, sulphur, phosphorus and halogens Quantitative analysis (basic principles only) Estimation of carbon, hydrogen, nitrogen, halogens, sulphur, phosphorus Calculations of empirical formulae and molecular formulae; Numerical problems in organic quantitative analysis UNIT XX Some Basic Principles of Organic Chemistry Tetravalency of carbon; Shapes of simple molecules hybridization (s and p); Classification of organic compounds based on functional groups: —C=C—,—C=C— and those containing halogens, oxygen, nitrogen and sulphur, Homologous series; Isomerism - structural and stereoisomerism Nomenclature (Trivial and IUPAC) Covalent bond fission Homolytic and heterolytic free radicals, carbocations and carbanions; stability of carbocations and free radicals, electrophiles and nucleophiles Electronic displacement in a covalent bond Inductive effect, electromeric effect, resonance and hyperconjugation Common types of organic reactions Substitution, addition, elimination and rearrangement UNIT XXI Hydrocarbons Classification, isomerism, IUPAC nomenclature, general methods of preparation, properties and reactions Alkanes Conformations: Sawhorse and Newman projections (of ethane); Mechanism of halogenation of alkanes Alkenes Geometrical isomerism; Mechanism of electrophilic addition: addition of hydrogen, halogens, water, hydrogen halides (Markownikoff's and peroxide effect); Ozonolysis, oxidation, and polymerization Alkenes acidic character; addition of hydrogen, halogens, water and hydrogen halides; polymerization Aromatic hydrocarbons Nomenclature, benzene structure and aromaticity; Mechanism of electrophilic substitution: halogenation, nitration, Friedel – Craft's alkylation and acylation, directive influence of functional group in mono-substituted benzene UNIT XXIV Organic Compounds Containing Nitrogen General methods of preparation, properties, reactions and uses Amines Nomenclature, classification, structure basic character and identification of primary, secondary and tertiary amines and their basic character Diazonium Salts Importance in synthetic organic chemistry UNIT XXV Polymers General introduction and classification of polymers, general methods of polymerization-addition and condensation, copolymerization; Natural and synthetic rubber and vulcanization; some important polymers with emphasis on their monomers and uses - polythene, nylon, polyester and bakelite UNIT XXVI Biomolecules General introduction and importance of biomolecules Carbohydrates Classification aldoses and ketoses; monosaccharides (glucose and fructose), constituent monosaccharides of oligosacchorides (sucrose, lactose, maltose) and polysaccharides (starch, cellulose, glycogen) Proteins Elementary Idea of α-amino acids, peptide bond, polypeptides; proteins: primary, secondary, tertiary and quaternary structure (qualitative idea only), denaturation of proteins, enzymes Vitamins Classification and functions Nucleic Acids Chemical constitution of DNA and RNA Biological functions of Nucleic acids UNIT XXVII Chemistry in Everyday Life Chemicals in medicines Analgesics, tranquilizers, antiseptics, disinfectants, antimicrobials, antifertility drugs, antibiotics, antacids, antihistamins - their meaning and common examples Chemicals in food Preservatives, artificial sweetening agents - common examples Cleansing agents Soaps and detergents, cleansing action Unit XXVIII Principles Related to Practical Chemistry — Detection of extra elements (N, S, halogens) in organic compounds; Detection of the following functional groups: hydroxyl (alcoholic and phenolic), carbonyl (aldehyde and ketone), carboxyl and amino groups in organic compounds UNIT XXII Organic Compounds Containing Halogens General methods of preparation, properties and reactions; Nature of C—X bond; Mechanisms of substitution reactions — Inorganic compounds Mohr's salt, potash alum Uses/environmental effects of chloroform, iodoform, freons and DDT — UNIT XXIII Organic Compounds Containing Oxygen General methods of preparation, properties, reactions and uses Alcohols, Phenols and Ethers Organic compounds Acetanilide, p-nitroacetan ilide, aniline yellow, iodoform — Chemistry involved in the titrimetric excercises - Acids bases and the use of indicators, oxali acid vs KMnO4, Mohr's salt vs KMnO4 Alcohols Identification of primary, secondary and tertiary alcohols; mechanism of dehydration Chemistry involved in the preparation of the following — Chemical principles involved in the qualitative salt analysis — Cations — Pb2+ , Cu2+, Al3+, Fe3+, Zn2+, Ni2+, Ca2+, Ba2+ , Mg2+ NH4+ Anions – CO32-, S2-, SO42-, NO2, NO3, Cl -, Br-, I- (Insoluble salts excluded) Phenols Acidic nature, electrophilic substitution reactions: halogenation, nitration and sulphonation, Reimer - Tiemann reaction Ethers: Structure Aldehyde and Ketones Nature of carbonyl group; Nucleophilic addition to >C=O group, relative reactivities of aldehydes and ketones; Important reactions such as - Nucleophilic addition reactions (addition of HCN, NH3 and its derivatives), Grignard reagent; oxidation; reduction (Wolff Kishner and Clemmensen); acidity of α hydrogen, aldol condensation, Cannizzaro reaction, Haloform reaction; Chemical tests to distinguish between aldehydes and Ketones Carboxylic Acids Acidic strength & factors affecting it — Chemical principles involved in the following experiments Enthalpy of solution of CuSO4 Enthalpy of neutralization of strong acid and strong base Preparation of lyophilic and lyophobic sols Kinetic study of reaction of iodide ion with hydrogen peroxide at room temperature JEE ADVANCED PHYSICAL CHEMISTRY General Topics Concept of atoms and molecules, Dalton's atomic theory, Mole concept, Chemical formulae, Balanced chemical equations, Calculations (based on mole concept) involving common oxidation-reduction, neutralisation, and displacement reactions, Concentration in terms of mole fraction, molarity, molality and normality Gaseous and Liquid States Absolute scale of temperature, ideal gas equation, Deviation from ideality, van der Waals' equation, Kinetic theory of gases, average, root mean square and most probable velocities and their relation with temperature, Law of partial pressures, Vapour pressure, Diffusion of gases Atomic Structure and Chemical Bonding Bohr model, spectrum of hydrogen atom, quantum numbers, Wave-particle duality, de-Broglie hypothesis, Uncertainty principle, Qualitative quantum mechanical picture of hydrogen atom, shapes of s, p and d orbitals, Electronic configurations of elements (up to atomic number 36), Aufbau principle, Pauli's exclusion principle and Hund's rule, Orbital overlap and covalent bond; Hybridisation involving s, p and d orbitals only, Orbital energy diagrams for homonuclear diatomic species, Hydrogen bond, Polarity in molecules, dipole moment (qualitative aspects only), VSEPR model and shapes of molecules (linear, angular, triangular, square planar, pyramidal, square pyramidal, trigonal bipyramidal, tetrahedral and octahedral) Energetics First law of thermodynamics, Internal energy, work and heat, pressure-volume work, Enthalpy, Hess's law, Heat of reaction, fusion and vaporization, Second law of thermodynamics, Entropy, Free energy, Criterion of spontaneity Chemical Equilibrium Law of mass action, Equilibrium constant, Le-Chatelier's principle (effect of concentration, temperature and pressure), Significance of DG and DGo in chemical equilibrium, Solubility product, common ion effect, pH and buffer solutions, Acids and bases (Bronsted and Lewis concepts), Hydrolysis of salts Electrochemistry Electrochemical cells and cell reactions, Standard electrode potentials, Nernst equation and its relation to DG, Electrochemical series, emf of galvanic cells, Faraday's laws of electrolysis, Electrolytic conductance, specific, equivalent and molar conductivity, Kohlrausch's law, Concentration cells Chemical Kinetics Rates of chemical reactions, Order of reactions, Rate constant, First order reactions, Temperature dependence of rate constant (Arrhenius equation) Solid State Classification of solids, crystalline state, seven crystal systems (cell parameters a, b, c), close packed structure of solids (cubic), packing in fcc, bcc and hcp lattices, Nearest neighbours, ionic radii, simple ionic compounds, point defects Solutions Raoult's law, Molecular weight determination from lowering of vapour pressure, elevation of boiling point and depression of freezing point Surface Chemistry Elementary concepts of adsorption (excluding adsorption isotherms), Colloids, types, methods of preparation and general properties, Elementary ideas of emulsions, surfactants and micelles (only definitions and examples) Nuclear Chemistry Radioactivity, isotopes and isobars, Properties of rays, Kinetics of radioactive decay (decay series excluded), carbon dating, Stability of nuclei with respect to proton-neutron ratio, Brief discussion on fission and fusion reactions INORGANIC CHEMISTRY Isolation/Preparation and Properties of the following Non-metals Boron, silicon, nitrogen, phosphorus, oxygen, sulphur and halogens, Properties of allotropes of carbon (only diamond and graphite), phosphorus and sulphur Preparation and Properties of the following Compounds Oxides, peroxides, hydroxides, carbonates, bicarbonates, chlorides and sulphates of sodium, potassium, magnesium and calcium, Boron, diborane, boric acid and borax, Aluminium, alumina, aluminium chloride and alums, Carbon, oxides and oxyacid (carbonic acid), Silicon, silicones, silicates and silicon carbide, Nitrogen, oxides, oxyacids and ammonia, Phosphorus, oxides, oxyacids (phosphorus acid, phosphoric acid) and phosphine, Oxygen, ozone and hydrogen peroxide, Sulphur, hydrogen sulphide, oxides, sulphurous acid, sulphuric acid and sodium thiosulphate, Halogens, hydrohalic acids, oxides and oxyacids of chlorine, bleaching powder, Xenon fluorides Transition Elements (3d series) Definition, general characteristics, oxidation states and their stabilities, colour (excluding the details of electronic transitions) and calculation of spin-only magnetic moment; Coordination compounds: nomenclature of mononuclear coordination compounds, cis-trans and ionisation isomerisms, hybridization and geometries of mononuclear coordination compounds (linear, tetrahedral, square planar and octahedral) Preparation and Properties of the following Compounds Oxides and chlorides of tin and lead, Oxides, chlorides and sulphates of Fe2+, Cu2+ and Zn2+, Potassium permanganate, potassium dichromate, silver oxide, silver nitrate, silver thiosulphate Ores and Minerals Commonly occurring ores and minerals of iron, copper, tin, lead, magnesium, aluminium, zinc and silver Extractive Metallurgy Chemical principles and reactions only (industrial details excluded), Carbon reduction method (iron and tin), Self reduction method (copper and lead), Electrolytic reduction method (magnesium and aluminium), Cyanide process (silver and gold) Principles of Qualitative Analysis Groups I to V (only Ag+, Hg2+, Cu2+, Pb2+, Bi3+, Fe3+, Cr3+, Al3+, Ca2+, Ba2+, Zn2+, Mn2+ and Mg2+), Nitrate, halides (excluding fluoride), sulphate and sulphide ORGANIC CHEMISTRY Concepts Hybridisation of carbon, Sigma and pi-bonds, Shapes of simple organic molecules, Structural and geometrical isomerism, Optical isomerism of compounds containing up to two asymmetric centres, (R,S and E,Z nomenclature excluded), IUPAC nomenclature of simple organic compounds (only hydrocarbons, mono-functional and bi-functional compounds), Conformations of ethane and butane (Newman projections), Resonance and hyperconjugation, Keto-enol tautomerism, Determination of empirical and molecular formulae of simple compounds (only combustion method), Hydrogen bonds, definition and their effects on physical properties of alcohols and carboxylic acids, Inductive and resonance effects on acidity and basicity of organic acids and bases, Polarity and inductive effects in alkyl halides, Reactive intermediates produced during homolytic and heterolytic bond cleavage, Formation, structure and stability of carbocations, carbanions and free radicals Preparation, Properties and Reactions of Alkanes Homologous series, physical properties of alkanes (melting points, boiling points and density), Combustion and halogenation of alkanes, Preparation of alkanes by Wurtz reaction and decarboxylation reactions Preparation, Properties and Reactions of Alkenes and Alkynes Physical properties of alkenes and alkynes (boiling points, density and dipole moments), Acidity of alkynes, Acid catalysed hydration of alkenes and alkynes (excluding the stereochemistry of addition and elimination), Reactions of alkenes with KMnO4 and ozone, Reduction of alkenes and alkynes, Preparation of alkenes and alkynes by elimination reactions, Electrophilic addition reactions of alkenes with X2, HX, HOX and H2O (X=halogen), Addition reactions of alkynes, Metal acetylides Reactions of Benzene Structure and aromaticity, Electrophilic substitution reactions, halogenation, nitration, sulphonation, Friedel-Crafts alkylation and acylation Effect of o-, m- and p-directing groups in monosubstituted benzenes Phenols Acidity, electrophilic substitution reactions (halogenation, nitration and sulphonation), ReimerTiemann reaction, Kolbe reaction Characteristic Reactions of the following (including those mentioned above) Alkyl halides, rearrangement reactions of alkyl carbocation, Grignard reactions, nucleophilic substitution reactions, Alcohols, esterification, dehydration and oxidation, reaction with sodium, phosphorus halides, ZnCl2/concentrated HCl, conversion of alcohols into aldehydes and ketones, Ethers, Preparation by Williamson's Synthesis, Aldehydes and Ketones, oxidation, reduction, oxime and hydrazone formation, aldol condensation, Perkin reaction, Cannizzaro reaction, haloform reaction and nucleophilic addition reactions (Grignard addition), Carboxylic acids, formation of esters, acid chlorides and amides, ester hydrolysis Amines, basicity of substituted anilines and aliphatic amines, preparation from nitro compounds, reaction with nitrous acid, azo coupling reaction of diazonium salts of aromatic amines, Sandmeyer and related reactions of diazonium salts, carbylamine reaction, Haloarenes, nucleophilic aromatic substitution in haloarenes and substituted haloarenes (excluding Benzyne mechanism and Cine substitution) Carbohydrates Classification, mono and disaccharides (glucose and sucrose), Oxidation, reduction, glycoside formation and hydrolysis of sucrose Amino Acids and Peptides General structure (only primary structure for peptides) and physical properties Properties and Uses of Some Important Polymers Natural rubber, cellulose, nylon, teflon and PVC Practical Organic Chemistry Detection of elements (N, S, halogens), Detection and identification of the following functional groups, hydroxyl (alcoholic and phenolic), carbonyl (aldehyde and ketone), carboxyl, amino and nitro, Chemical methods of separation of mono-functional organic compounds from binary mixtures Some Basic Concepts of Chemistry The percentage composition of carbon by mole in methane is Topic Mole Concept (2019 Main, April II) Objective Questions I (Only one correct option) moles of AB2 weight 125 × 10−3 kg and 10 moles of A2 B2 weight 300 × 10−3 kg The molar mass of A ( M A ) and molar mass of B ( M B ) in kg mol −1 are (2019 Main, 12 April I) (a) M A = 10 × 10−3 and M B = × 10−3 (b) M A = 50 × 10−3 and M B = 25 × 10−3 (d) M A = × 10 and M B = 10 × 10 reactant is for the reaction (Given atomic mass : Fe = 56, O = 16, Mg = 24, P = 31, C = 12, H = 1) (2019 Main, 10 April II) C3 H8 ( g ) + 5O2 ( g ) → 3CO2 ( g ) + 4H2 O( l ) P4 ( s ) + 5O2 ( g ) → P4 O10 ( s ) 4Fe( s ) + 3O2 ( g ) → 2Fe2 O3 ( s ) 2Mg ( s ) + O2 ( g ) → 2MgO( s ) 10 mL of a hydrocarbon required 55 mL of O2 for complete combustion and 40 mL of CO2 is formed The formula of the hydrocarbon is (2019 Main, 10 April I) (c) C4H10 (d) C4H (a) C4H7Cl (b) C4H 10 mL of mM surfactant solution forms a monolayer covering 0.24 cm on a polar substrate If the polar head is approximated as a cube, what is its edge length? (2019 Main, April II) (b) 0.1 nm (c) 1.0 pm (d) 80% g of NaOH is dissolved in 18 g of H2 O Mole fraction of NaOH in solution and molality (in mol kg− ) of the solution respectively are (2019 Main, 12 Jan II) (a) 0.2, 11.11 (b) 0.167, 22.20 (c) 0.2, 22.20 (d) 0.167, 11.11 (d) 2.0 nm For a reaction, N ( g ) + 3H ( g ) → 2NH ( g ), identify dihydrogen (H ) as a limiting reagent in the following reaction mixtures (2019 Main, April I) (a) 56 g of N + 10 g of H (b) 35 g of N + g of H (c) 14 g of N + g of H (d) 28 g of N + g of H (2019 Main, 12 Jan II) (d) 5.6 The amount of sugar (C12 H22 O11 ) required to prepare L of its 0.1 M aqueous solution is (2019 Main, 10 Jan II) (a) 17.1 g (b) 68.4 g (c) 136.8 g (d) 34.2 g 10 For the following reaction, the mass of water produced from 445 g of C57 H110 O6 is : 2C57 H110 O6 ( s ) + 163O2 ( g ) → 114CO2 ( g ) + 110 H2 O ( l ) (2019 Main, Jan II) (a) 490 g At 300 K and atmospheric pressure, (a) 2.0 pm (c) 25% (Molar mass of H2 O2 = 34 g mol −1 ) (a) 16.8 (b) 22.4 (c) 11.35 −3 The minimum amount of O2 ( g ) consumed per gram of (a) (b) (c) (d) (b) 20% The volume strength of M H2 O2 is (c) M A = 25 × 10−3 and M B = 50 × 10−3 −3 (a) 75% (b) 495 g (c) 445 g (d) 890 g 11 A solution of sodium sulphate contains 92 g of Na + ions per kilogram of water The molality of Na + ions in that solution in mol kg−1 is (2019 Main, Jan I) (a) 16 (b) (c) 132 (d) 12 The most abundant elements by mass in the body of a healthy human adult are oxygen (61.4%), carbon (22.9%), hydrogen (10.0 %), and nitrogen (2.6%) The weight which a 75 kg person would gain if all Hatoms are replaced by Hatoms is (2017 JEE Main) (a) 15 kg (c) 7.5 kg (b) 37.5 kg (d) 10 kg 13 g of a carbonate (M CO3 ) on treatment with excess HCl produces 0.01186 mole of CO2 The molar mass of M CO3 in g mol −1 is (2017 JEE Main) (a) 1186 (b) 84.3 (c) 118.6 (d) 11.86 States of Matter 77 16 In the van der Waals’ equation V =1 Videal For ideal gas V = Videal 27 Compressibility factor (Z ) = n a p + (V − nb) = nRT V n2a corrects for V2 intermolecular force while b corrects for molecular volume The additional factor in pressure, i.e Q 28 Expression of rms is, urms = urms (H2 at 50 K) = urms (O2 at 800 K) ⇒ 17 Option (b) is incorrect statement because at high pressure slope of the line will change from negative to positive 18 r(He) r(CH4 ) = 16 =2 :1 = 19 Kinetic energy (E ) = kT RMS speed (u) = x 30 x Mole of hydrogen = pV pV , for positive deviation, > Q Z= nRT nRT 21 Option (b) and (d) are ruled out on the basis that at the initial point of 273 K, atm, for 1.0 mole volume must be 22.4 L, and it should increase with rise in temperature Option (a) is ruled out on the basis that initial and final points are not connected by the ideal gas equation V ∝ T , i.e V /T not have the same value at the two points In option (c), at the initial point, the volume is 22.4 L as required by ideal gas equation and (V /T ) have the same value at both initial and final points 22 Root mean square velocity (urms ) = 3RT M V < 1(given) Vid V < 22.4 L Vid (1 mol ) = 22.4 L at STP 23 Compressibility factor (Z ) = 24 Root mean square speed urms = ⇒ Mole fraction of hydrogen = 3RT M 14T (H2 ) urms (H2 ) T (H2 ) 28 ⇒ 7= = 7= × urms (N2 ) T (N2 ) T (N2 ) T (N2 ) = 2T (H2 ) i.e T (H2 ) < T (N2 ) 25 At high temperature and low pressure, the gas volume is infinitely large and both intermolecular force as well as molecular volume can be ignored Under this condition postulates of kinetic theory applies appropriately and gas approaches ideal behaviour 26 Rate of effusion ∝ pi ; pi = Partial pressure of ith component ∝ M x x x + 30 = 15 16 Partial pressure of H2 = Mole fraction of hydrogen Total pressure = 15 : 16 8RT 30 Average speed = πM ⇒ i.e at constant volume, for a fixed mass, increasing temperature increases average speeds and molecules collide more frequently to the wall of container leading to increase in gas pressure 31 The mean translational kinetic energy (∈) of an ideal gas is p ⋅ M = dRT Substituting for RT / M in urms expression gives, 3p urms = ⇒ urms ∝ d d ⇒ 50 32 × =1 800 Mole of ethane = 20 Positive deviation corresponds to Z > ⇒ Q 3R × 50 3R × 800 32 29 Let x g of each gas is mixed 3kT 2E ⇒ u= m m Also, 3RT M ∈= 32 k BT ; T = Absolute temperature, i.e ∈ ∝ T rCH rX =2= MX 16 ⇒ M X = 64 33 The ideal gas equation, pV = nRT = ⇒ w RT M w pM = RT = dRT V (d = density) pM RT i.e density will be greater at low temperature and high pressure ⇒ d= 34 The ease of liquefication of a gas depends on their intermolecular force of attraction which in turn is measured in terms of van der Waals’ constant a Hence, higher the value of a, greater the intermolecular force of attraction, easier the liquefication In the present case, NH3 has highest a, can most easily be liquefied 35 HCl will diffuse at slower rate than ammonia because rate of effusion ∝ M Therefore, ammonia will travel more distance than HCl in the same time interval and the two gas will first meet nearer to HCl end 78 States of Matter 36 In van der Waals’ equation of state 44 The two types of speeds are defined as; a p + (V − b) = RT V (For mole) Root mean square speed (urms ) = 3RT M The first factor ( p + a/V ) correct for intermolecular force while the second term (V − b) correct for molecular volume Average speed (uav ) = 8RT πM 37 Expression for average velocity is uav = 8RT πM For the same gas but at different temperature uavg (T1 ) uavg (T2 ) ⇒ = T1 300 = = T2 1200 uav (927° C) = × uav (27° C) = 0.6 ms−1 38 Rate of effusion ∝ , M x x Moles of CH4 = 16 Moles of H2 = ⇒ x Mole fraction of H2 = = x x + 16 Partial pressure of H2 = Mole fraction of H2 = Total pressure 40 According to postulates of kinetic theory, there is no intermolecular attractions or repulsions between the molecules of ideal gases 41 According to kinetic theory, average kinetic energy (E ) = k BT where, k B is Boltzmann’s constant Since, it is independent of molar mass, it will be same for He and H2 at a given temperature 42 If x g of both oxygen and methane are mixed then : x 32 x Mole of methane = 16 x 32 Mole fraction of oxygen = ⇒ = x x + 32 16 According to law of partial pressure Partial pressure of oxygen ( pO ) = Mole fraction × Total pressure pO = ⇒ p Mole of oxygen = 43 It is the Boyle temperature TB At Boyle temperature, the first virial coefficient ( B ) vanishes and real gas approaches ideal behaviour TB = = 3: = : 2.54 = 1.085 : π 45 The explanation of given statements are as follows: (a) Urms is inversely proportional to the square root of its molecular mass 39 Let x grams of each hydrogen and methane are mixed, ⇒ For the same gas, at a given temperature, M and T are same, therefore urms 3RT 8RT = : πM uav M a Rb Here, a and b are van der Waals’ constants 3RT M Hence, option (a) is correct (b) When temperature is increased four times then Urms become doubled Urms = Urms = 3R × 4T M Urms = × 3RT M Hence, option (b) is correct (c) and (d) Eav is directly proportional to temperature but does not depends on its molecular mass at a given temperature as Eav = KT If temperature raised four times than Eav becomes four time multiple Thus, option (c) is incorrect and option (d) is correct 46 Equation of state p(V − b) = RT indicates absence of intermolecular attraction or repulsion, hence interatomic potential remains constant on increasing ‘π’ in the beginning As the molecules come very close, their electronic and nuclear repulsion increases abruptly 47 (a) According to a postulate of kinetic theory of gases, collision between the molecules as well as with the wall of container is perfectly elastic in nature (b) If a gas molecule of mass m moving with speed u collide to the wall of container, the change in momentum is ∆p = – 2mu Therefore, heavier molecule will transfer more momentum to the wall as there will be greater change in momentum of the colliding gas molecule However, this is not postulated in kinetic theory (c) According to Maxwell-Boltzmann distribution of molecular speed, very few molecules have either very high or very low speeds Most of the molecules moves in a specific, intermediate speed range (d) According to kinetic theory of gases, a gas molecule moves in straight line unless it collide with another molecule or to the wall of container and change in momentum is observed only after collision States of Matter 79 48 Option (a) is correct because in the limit of large volume, both or intermolecular force and molecular volume becomes negligible in comparison to volume of gas Option (d) is wrong statement because Z can be either less or greater than unity, hence real pressure can be less or greater than ideal pressure 51 Given p1 = bar, V1 = m 3, T1 = 400 K So, n1 = 400 R Similarly, p2 = bar, V2 = m 3, T2 = 300 K, n2 = 11 20 = or 9 52 Assertion is incorrect because besides amount, pressure also depends on volume However, reason is correct because both frequency of collisions and impact are directly proportional to root mean square speed which is proportional to square root of absolute temperature 53 a is the measure of intermolecular force of attraction Greater the intermolecular force of attraction (H-bond in the present case) higher the value of a 54 X is a lighter gas than Y, hence X has greater molecular speed Due to greater molecular speed of X, it will have smaller mean free path and greater collision frequency with the incrt gas molecules As a result X will take more time to travel a given distance along a straight line Hence X and Y will meet at a distance smaller than one calculated from Graham’s law 50 Initial pressure ( p1 ) = 48 × 10−3 bar Final pressure ( p2 ) = …… × 10−6 bar Initial volume (V1 ) = π (3)3 Final volume (V2 ) = π (12)3 According to Boyle’s law p1V1 = p2V2 pV p2 = 1 V2 48 × 10−3 × π (3)3 48 × 10−3 × (3)3 p2 = = (12)3 π (12)3 48 × 10−3 × 27 = = 0.0277 × 27 × 10−3 = 750 × 10−6 bar 1728 Hence, the correct answer is 750 11 2.22 49 Pressure is inversely proportional to volume at constant temperature, hence (a) is correct Average kinetic energy of a gas is directly proportional to absolute temperature, hence (b) is correct Expansion at constant temperature cannot change the number of molecules, hence (d) is incorrect + x = 15 − 5x or x = Hence, new volume of A i.e., (1+ x ) will comes as + Option (b) is wrong statement because in the limit of large pressure Z > Option (c) is correct statement For a van der Waals’ gas, van der Waals’ constants a and b are characteristic of a gas, independent of temperature (1+ x ) = 15 − 5x or Hence, (d) is the correct choice 55 PLAN This problem can be solved by using the concept of Graham’s law of diffusion according to which rate of diffusion of non-reactive gases under similar conditions of temperature and pressure are inversely proportional to square root of their density Rate of diffusion ∝ molar weight of gas Let distance covered by X is d, then distance covered by Y is 24 – d If rX and rY are the rate of diffusion of gases X and Y, 40 rX d = = =2 10 rY 24 − d [Q Rate of diffusion ∝ distance travelled] d = 48 − 2d ⇒ 3d = 48 ⇒ d = 16 cm Hence, (c) is the correct choice 56 A At p = 200 atm, very high pressure, Z > Also, at such a high (from pV = nRT ) 300 R Let at equilibrium the new volume of A will be (1+ x ) So, the new volume of B will be (3− x ) Now, from the ideal gas equation p1V1 pV = 2 n1RT1 n2RT2 and at equilibrium (due to conduction of heat) p1 p2 = T1 T2 V1 V2 So, = or V1n2 = V2n1 n1 n2 After putting the values (3 − x ) 5 or (1+ x ) = (1+ x ) × = (3 − x ) × 300 R 400 R n2a pressure, the pressure correction factor can be ignored V in comparison to p B At p ~ 0, gas will behave like an ideal gas, pV = nRT C CO2 (p = 1atm, T = 273 K), Z < D At very large molar volume, real gas behaves like an ideal gas 57 Less; E = RT nRT At same temperature, KE (total) ∝ n 59 0.25 RT because at NTP, 5.6 L = mole 60 Inversely, time 58 : 16, KE = 61 For an ideal gas, Cp − CV = R 3 62 At 27°C, E = RT = × × 300 = 900 cal 80 States of Matter 63 An ideal gas cannot be liquefied because there exist no urms = umps 71 Given, intermolecular attraction between the molecules of ideal gas 3RT = M (X ) ⇒ 64 a is the measure of intermolecular force 65 In a close container, gas exert uniform pressure everywhere in the container 66 KE = RT where, T is absolute temperature (in Kelvin) 3R × 400 2R × 60 = 40 M (Y ) ⇒ 72 67 (DC) Diffusion coefficient ∝ λ (mean free path) ∝ U mean ⇒ Thus (DC) ∝ λ Umean T RT ⇒ λ∝ p N0 σp λ= But, rgas rO 8RT πM U mean = and 73 DC ∝ T2 T1 3/ p 4T = 1 p1 T1 92 U 238 → 82Pb 206 + 2He4 (g ) + 6− 1β n(gas)[Initial] = (air) n(gas)[Final] = (He) + 1(air) = ⇒ At constant temperature and volume; p ∝ n pf nf So, = = =9 pi ni 69 urms = 3π uav = = 1.33 = 32 Mgas (ii) Ek = a p + V = RT V pV a + =1 RT VRT a Z+ =1 ZRT RT p ⇒ ⇒ a= ∴ ⇒ pV + a = RT V pV Q =Z RT ⇒ Z+ ap =1 ZR 2T ZR 2T (1 − Z ) 0.5 (0.082 × 273)2 (1− 0.5) = p 100 a = 1.25 atm L2 mol −2 k = Boltzmann constant 75 In case of negligible molecular volume, b = and N A = Avogadro’s number −23 3 k BT = × 1.38 × 10−23 × 1000 J = 2.07 × 10−20 J 2 74 In case of negligible molecular volume, b = For mole of gas Universal gas constant, R = kN A and × 3.14 × 400 = 434 ms−1 (d) Q Z > 1, repulsive force is dominating ⇒ PLAN This problem can be solved by using the concept involved in calculation of significant figure where, 3π 18 = 50 L mol −1 0.36 pV × 50 (c) Z = = = 1.22 RT 0.082 × 500 3/ 1 = (8) = 2 68 8RT 3RT = : M πM (b) Vm = (T )3/ p p (DC)2 (x ) = p2 (DC)1 uav = urms ⇒ M (Y ) = (i) (a) Mgas = 18 g mol −1 U mean ∝ T ∴ 2RT M (Y ) 23 R = 1.380 × 10 × 6.023 × 10 J/Kmol ~ 8.312 = 8.31174 = Since, k and N A both have four significant figures, so the value of R is also rounded off upto significant figures [When number is rounded off, the number of significant figure is reduced, the last digit is increased by if following digits ≥ and is left as such if following digits is ≤ 4.] Hence, correct integer is (4) 70 Since, the external pressure is 1.0 atm, the gas pressure is also 1.0 atm as piston is movable Out of this 1.0 atm partial pressure due to unknown compound is 0.68 atm Therefore, partial pressure of He = 1.00 – 0.68 = 0.32 atm n(He)RT 0.1 × 0.082 × 273 Volume = = =7L ⇒ p(He) 0.32 ⇒ Volume of container = Volume of He van der Waals’ equation reduces to n2a p + V = nRT V RT a − (n =1 mole) V V 0.082 × 273 3.592 = 0.99 atm = − 22.4 (22.4)2 ⇒ p= 76 (i) For the same amount of gas being effused r1 t2 p1 = = r2 t1 p2 ⇒ M2 57 0.8 = ⇒ M1 38 1.6 M2 28 M = 252 g mol −1 Also, one molecule of unknown xenon-fluoride contain only one Xe atom [M (Xe) = 131], formula of the unknown gas can be considered to be XeFn ⇒ 131 + 19n = 252; n = 6.3, hence the unknown gas is XeF6 States of Matter 81 (ii) For a fixed amount and volume, p ∝ T T where, T = Kelvin temperature ⇒ = 1.1 T + 10 ⇒ ⇒ 82 Weight of butane gas in filled cylinder = 29 − 14.8 kg = 14.2 kg ⇒ During the course of use, weight of cylinder reduces to 23.2 kg ⇒ Weight of butane gas remaining now = 23.2 − 14.8 = 8.4 kg Also, during use, V (cylinder) and T remains same p1 n1 Therefore, = p2 n2 n n w 8.4 p2 = p1 = ⇒ × 2.5 Here, = 14.2 n1 n w1 T = 100 K = t + 273 t = − 173° C nRT 12 0.082 × 100 Volume = = = 0.82 L × 120 p 77 The van der Waals’ equation is = 1.48 atm Also, pressure of gas outside the cylinder is 1.0 atm ⇒ pV = nRT nRT (14.2 − 8.4 ) × 103 0.082 × 30 ⇒ V = = × L p 58 n2a p + (V − nb) = nRT V ⇒ a= (4 )2 × 0.082 × 300 V nRT − p = − 11 − ( 0.05 ) n V − nb ( ) = 6.46 atm L2 mol −2 78 Mass of liquid = 148 − 50 = 98 g 98 = 100 mL = volume of flask ⇒ Volume of liquid = 0.98 mass of gas = 50.5 − 50 = 0.50 g w Now applying ideal gas equation : pV = RT M wRT 0.5 × 0.082 × 300 ⇒ M = = = 123 g mol −1 pV × 0.1 79 False, ideal gas cannot be liquefied as there is no intermolecular attraction between the molecules of ideal gas Hence, there is no point of forming ideal solution by cooling ideal gas mixture 80 If ‘α’ is the degree of dissociation, then at equilibrium Cl r 2Cl Moles −α 2α From diffusion information r(mix) = 1.16 = r(Kr) ⇒ ⇒ 84 M (mix) pV × 40 = RT 0.082 × 400 = 1.22 81 The total moles of gaseous mixture = Let the mixture contain x mole of ethane Therefore, C2H6 + O2 → 2CO2 + 3H2O x C2H4 + 3O2 → 2CO2 + 2H2O x x + (1.22 − x ) = + 3.66 2 130 x ⇒ = + 3.66 32 ⇒ x = 0.805 mole ethane and 0.415 mole ethene 0.805 = 0.66 ⇒ Mole fraction of ethane = 1.22 Mole fraction of ethene = − 0.66 = 0.34 Total moles of O2 required = rCH4 = M CH nHe n CH4 M He = 16 =8 Initial ratio of rates of effusion gives the initial composition of mixture effusing out Therefore, n (He) : n (CH4 ) = : 84 Number of moles = × 1021 = 0.33 × 10−2 × 1023 p = 7.57 × 103 Nm −2 pV = nRT 7.57 × 103 × 10−3 pV T = = 276 K = nR 0.33 × 10−2 × 8.314 Now, ⇒ ⇒ Also, α = 0.14 1.22− x rHe 83 Total = + α M (mix) = 62.4 71 M (mix) = = 62.4 1+ α ⇒ = 2460 L = 2.46 m ⇒ urms = 3RT = M × 8.314 × 276 m s−1 = 496 ms−1 28 × 10−3 umps = 0.82 urms umps = 0.82 × urms = 0.82 × 496 ms−1 = 407 ms−1 85 First we calculate partial pressure of NO and O2 in the combined system when no reaction taken place pV = constant ⇒ p1V1 = p2V2 1.053 × 250 ⇒ p2 (NO) = = 0.752 atm 350 0.789 × 100 p2 (O2 ) = = 0.225 atm 350 Now the reaction stoichiometry can be worked out using partial pressure because in a mixture pi ∝ ni 2NO + O2 → 2NO2 → N2O4 Initial Final 0.752 atm 0.302 0.225 atm 0 0 0.225 atm Now, on cooling to 220 K, N2O4 will solidify and only unreacted NO will be remaining in the flask Q p∝T p1 T1 ∴ = p2 T2 0.302 300 = ⇒ p2 220 ⇒ p2 (NO) = 0.221 atm 82 States of Matter 86 Total moles of gas in final mixture = pV 6×3 = = 0.731 RT 0.082 × 300 Q Mole of H2 in the mixture = 0.70 ∴ Mole of unknown gas ( X ) = 0.031 Because both gases have been diffused for same time r (H2 ) 0.70 M = = ⇒ M = 1020 g mol −1 r ( X ) 0.031 87 pressure inside the bottle When the bottle is opened, there is chances of bumping of stopper To avoid bumping, bottle should be cooled that lowers the pressure inside (ii) According to Avogadro’s hypothesis, “Under identical conditions of pressure and temperature, equal volume of ideal gases contain equal number of molecules.” and nRT p For acetylene gas, g = mol 26 740 p = 740 mm = atm 760 V = ⇒ 7.6 × 10−10 = × × 6.023 × 1023 760 0.082 × 273 = 2.7 × 1010 molecules T = 50° C = 323 K 93 From the given information, it can be easily deduced that in the Substituting in ideal gas equation 0.082 × 323 V = × × 76 = 5.23 L 26 74 88 uav (average velocity ) = × 10 ms−1 = 100 ⇒ Hence, × 8.314 T1 ⇒ 3.14 × 44 × 10−3 ⇒ 8T1 × = 2T2 π 4T1 πT2 4T × 1682.5 = 2142 K T2 = = π 3.14 T1 = 1682.5 K, T2 = 2142 K 89 Volume of balloon = 4 21 πr = × 3.14 × cm 2 3 = 4847 cm ≈ 4.85 L Now, when volume of H2 (g ) in cylinder is converted into NTP volume, then p1V1 p2V2 20 × 2.82 × V2 = = ⇒ T1 T2 300 273 V2 = NTP volume ⇒ V2 = 51.324 L Also, the cylinder will not empty completely, it will hold 2.82 L of H2 (g ) when equilibrium with balloon will be established Hence, available volume of H2 (g ) for filling into balloon is 51.324 − 2.82 = 48.504 L 48.504 = 10 ⇒ Number of balloons that can be filled = 4.85 90 urms = partial pressure of A = 1.0 atm partial pressure of B = 0.5 atm pV V nA = A = RT RT pBV 0.5 V nB = = RT RT M nB wB M = = × A = × A nA M B wA M B Also 4T1 πT2 1= ⇒ final mixture, 8RT1 πM T1 = 1682.5 K Also, for the same gas 8RT1 2RT2 uav = : = umps M πM ⇒ pV RT N (Number of molecules) n= N A (Avogadro number) pV N = nNA = N RT A 92 Number of moles (n) = × 8.314 × 293 3RT = 390.2 ms−1 = M 48 × 10−3 91 (i) NH3 (l ) is highly volatile, a closed bottle of NH3 (l ) contains large number of molecules in vapour phase maintaining high ⇒ M A :M B = : p 94 Rate of effusion (r) ∝ M ⇒ r (NH3 ) = × r (HCl ) 17 ⇒ p= 36.5 ⇒ p 40 = 60 p 36.5 17 36.5 = 2.20 atm 17 3 = × 1.38 × 10−23 × 300 J = 6.21 × 10−21 J/molecule dp kp 96 Rate of effusion is expressed as − = dt M k = constant, p = instantaneous pressure dp k dt ⇒ − = p M p kt Integration of above equation gives ln = p M k 2000 47 Using first information : ln = 1500 32 95 KE = k BT : k B = Boltzmann’s constant ⇒ k= 32 4 ln 3 47 …(i) Now in mixture, initially gases are taken in equal mole ratio, hence they have same initial partial pressure of 2000 mm of Hg each After 74 : States of Matter 83 2000 74 k = ln 32 pO Also, the decomposition reaction is : MCO3 → MO + CO2 0.05 mol Substituting k from Eq (i) gives 2000 74 32 4 = × ln ln 3 p 47 32 O2 Q ∴ ⇒ ⇒ Q 2000 74 4 = ln ln 3 p 47 O2 Solving gives p (O2 ) at 74 = 1271.5 mm 2000 74 k = For unknown gas : ln 79 pg 100 The ideal gas equation : pV = nRT = Substituting k from (i) gives 2000 74 32 4 = ln ln × 3 p 47 79 g 97 First we determine empirical formula as C H Weight 10.5 Mole 10.5 = 0.875 12 Simple ratio Whole no pM = ⇒ Topic Liquid State I (CH3OH) : Surface tension decreases as concentration increases II (KCl) : Surface tension increases with concentration for ionic salt III [CH3 (CH2 )11 OSO−3 Na + ] : It is an anionic detergent There is decrease in surface tension before micelle formation, and after CMC (Critical Micelle Concentration) is attained, no change in surface tension 1/0.875 = 1.14 ⇒ Empirical formula = C7H8 w From gas equation : pV = RT M wRT 2.8 × 0.082 × 400 = = 91.84 ≈ 92 pV 1× Q Molar mass (M ) is same as empirical formula weight Molecular formula = Empirical formula = C7 H8 98 For same p and V , n ∝ T n (gas) T (H2 ) = ⇒ n (H2 ) T (gas) 0.184 n(H2 ) = = 0.092 290 n(gas) = × 0.092 = 0.0895 ⇒ 298 Q 0.0895 mole of gas weigh 3.7 g 3.7 ∴ mole of gas will weigh = 41.32 g 0.0895 99 Moles of CO2 can be calculated using ideal gas equation as : n= w RT M w RT = d RT where, ‘d’ is density V pM × 17 d= = = 3.42 g L−1 RT 0.082 × 303 ⇒ Solving gives : pg = 1500 mm ⇒ After 74 min, p (O2 ) : p (g ) = 1271.5 : 1500 Also, in a mixture, partial pressure ∝ number of moles ⇒ n (O2 ) : n (g ) = : 1.18 M = 0.05 mol 0.05 mole MCO3 = 4.21 g 4.215 1.0 mole MCO3 = = 84.3 g (molar mass) 0.05 84.3 = MW of M + 12 + 48 Molecular weight of metal = 24.3 Metal is bivalent, equivalent weight Molecular weight = = 12.15 pV 700 1336 = 0.05 = × RT 760 1000 0.082 × 300 Surface tension For O2 KCl (II) CH3OH (I) − CH3(CH2)11OSO3 Na + (III) Concentration Let us consider, 1.0 L of liquid water is converted into steam Volume of H2O (l) = 1L, mass = 1000 g 1000 Volume of 1000 g steam = ⇒ cm 0.0006 1000 cm steam = 1000 cm Q Volume of molecules in 0.0006 ∴ Volume of molecules in 1000 1000 cm steam = × 0.0006 × 1000 = 0.60 cm 1000 Critical temperature is directly proportional to intermolecular force of attraction H2O is a polar molecule, has greater intermolecular force of attraction than O2, hence higher critical temperature At liquid-vapour equilibrium at boiling point, molecules in two phase posses the same kinetic energy Chemical and Ionic Equilibrium Topic Chemical Equilibrium Objective Questions I (Only one correct option) - Consider the following reaction: N 2O4 ( g ) q 2NO2 ( g ); ∆H ° = + 58 kJ For each of the following cases (A , B ), the direction in which the equilibrium shifts is (2020 Main, Sep I) (A) temperature is decreased (B) pressure is increased by adding N at constant T (a) (A) towards product, (B) towards reactant (b) (A) towards reactant, (B) no change (d) (A) towards product, (B) no change The incorrect match in the following is (a) 10 25 (2019 Main, 12 April II) (b) ∆G ° = 0, K = (d) ∆G ° < 0, K < c c c c 2SO2 ( g ) + O2 ( g ) → 2SO3 ( g ), ∆H = − 57.2 kJ mol −1 and K c = 1.7 × 1016 Which of the following statement is incorrect? (2019 Main, 10 April II) (a) The equilibrium constant decreases as the temperature increases (b) The addition of inert gas at constant volume will not affect the equilibrium constant (c) The equilibrium will shift in forward direction as the pressure increases (d) The equilibrium constant is large suggestive of reaction going to completion and so no catalyst is required For the following reactions, equilibrium constants are given : 52 - SO (g ); K = 10 2S( s ) + 3O ( g ) - 2SO ( g ); K = 10 - 2C + D, (a) (c) (b) 16 (d) - B (g ) + C(g ); K D ( s ) - C( g ) + E ( g ); K A (s ) p1 p2 = x atm2 = y atm2 The total pressure when both the solids dissociate simultaneously is (2019 Main, 12 Jan I) For the reaction, (d) 10181 K Two solids dissociate as follows: 2CO(g ) (a) 2C(s) + O2 (g ) H2 (g ) + I2 (g ) (b) 2HI(g ) (c) NO2 (g ) + SO2 (g ) NO(g ) + SO3 (g ) (d) 2NO(g ) N2 (g ) + O2 (g ) (c) 10 the initial concentration of B was 1.5 times of the concentration of A, but the equilibrium concentrations of A and B were found to be equal The equilibrium constant (K ) for the aforesaid chemical reaction is (2019 Main, 12 April II) (2019 Main, April II) 154 In a chemical reaction, A + 2B In which one of the following equilibria, K p ≠ K c ? S( s ) + O2 ( g ) (b) 10 77 (2019 Main, 12 Jan I) (c) (A) towards reactant, (B) towards product (a) ∆G ° < 0, K > (c) ∆G ° > 0, K < The equilibrium constant for the reaction, 2SO2 ( g ) + O2 ( g ) 2SO3 ( g ) is 129 (a) x + y atm (b) x + y2 atm (c) (x + y) atm (d) 2( x + y ) atm Consider the reaction, N ( g ) + 3H ( g ) =2NH (g ) The equilibrium constant of the above reaction is K p If pure ammonia is left to dissociate, the partial pressure of ammonia at equilibrium is given by (Assume that (2019 Main, 11 Jan I) pNH < < p total at equilibrium) (a) (c) 33/ K p1/ 2P Kp (b) 1/ 2 16 P (d) 33/ K p1/ 2P 16 Kp 1/ 2 P 5.1 g NH4 SH is introduced in 3.0 L evacuated flask at 327°C 30% of the solid NH4 SH decomposed to NH3 and H2 S as gases The K p of the reaction at 327° C is Chemical and Ionic Equilibrium 85 (R = 0.082 atm mol −1 K −1 , molar mass of S = 32 g mol −1 , −1 molar mass of N = 14 g mol ) (2019 Main, 10 Jan II) (a) 242 × 10−4 atm (b) 242 atm (c) × 10−3 atm (d) × 10−4 atm 10 The values of Kp KC then the value of x is (assuming ideality) (a) 1, 24.62 dm atm mol −1 , 606.0 dm atm mol −2 (c) 24.62 dm atm mol −1, 606.0 dm6 atm −2 mol 2, 1.65 × 10−3 dm −6 atm −2 mol (d) 1, 4.1 × 10−2 dm −3atm −1 mol, 606 dm6 atm mol −2 11 Consider the following reversible chemical reactions, K1 K2 2 The relation between K and K is (a) K = K13 (c) K = K1− …(i) …(ii) (2019 Main, Jan II) (b) K 1K = (d) K 1K = 12 An aqueous solution contains 0.10 M H2S and 0.20 M HCl If the equilibrium constants for the formation of HS− from H2S is 10 × 10−7 and that of S2− from HS− ions is 12 × 10−13 then the concentration of S2− ions in aqueous solution is : (a) × 10−8 M −21 (c) × 10 M (b) × 10−20 M −19 (d) × 10 (2018 Main) M 13 The equilibrium constant at 298 K for a reaction, A + B q C + D is 100 If the initial concentrations of all the four species were M each, then equilibrium concentration of D (in mol L−1 ) will be (2016 Main) (a) 0.818 (c) 1.182 (b) 1.818 (d) 0.182 14 The standard Gibbs energy change at 300 K for the reaction, B + C is 2494 J At a given time, the composition 1 of the reaction mixture is [A]= , [ B ] = and [C ] = The 2 reaction proceeds in the (R = 8.314JK / mol, e = 2.718) (2015, Main) 2A a (a) forward direction because Q > K c (b) reverse direction because Q > K c (c) forward direction because Q < K c (d) reverse direction because Q < K c (d) CO2 , H2CO3 17 N2 + 3H2 r 2NH3 Which is correct statement if N2 is added at equilibrium condition? (2006, 3M) (b) 1, 24.62 dm atm mol −1, 1.65 × 10−3 dm −6 atm −2 mol -2 AB (g ) AB ( g ) - A ( g ) + 3B ( g ) (2006 Main) (c) HCO−3 , CO23− (2019 Main, 10 Jan I) A2 ( g ) + B2 ( g ) 16 The species present in solution when CO2 is dissolved in (b) H2CO3 , CO2− 2 (d) (a) CO2 , H2CO3 , HCO3− , CO32− =2NO(g ) =2NO (g ) N (g ) + 3H (g ) =2NH (g ) (2014 Main) (c) water are −1 respectively (At 300 K, RT = 24.62 dm atm mol ) N ( g ) + O2 ( g ) N 2O4 ( g ) (b) − (a) − for the following reactions at 300 K are, SO3 ( g ) if K p = K C ( RT )x where, the symbols have usual meaning, 15 For the reaction, SO2 ( g ) + O2 ( g ) q (a) The equilibrium will shift to forward direction because according to IInd law of thermodynamics, the entropy must increases in the direction of spontaneous reaction (b) The condition for equilibrium is G(N2) + 3G(H2) = 2G(NH3) where, G is Gibbs free energy per mole of the gaseous species measured at that partial pressure The condition of equilibrium is unaffected by the use of catalyst, which increases the rate of both the forward and backward reactions to the same extent (c) The catalyst will increase the rate of forward reaction by α and that of backward reaction by β (d) Catalyst will not alter the rate of either of the reaction 18 Ag + + NH3 s [Ag(NH3 )]+ ; K = × 10−3 [Ag (NH3 )]+ + NH3 s [Ag (NH3 )2 ]+ ; K = 1.7 ×10−3 then the formation constant of [Ag(NH3 )2 ]+ is (a) 6.08 × 10 −6 (b) 6.08 × 10 −9 (d) None of these (c) 6.08 × 10 (2006, 3M) 19 Consider the following equilibrium in a closed container N2 O4 ( g ) r 2NO2 ( g ) At a fixed temperature, the volume of the reaction container is halved For this change, which of the following statements hold true regarding the equilibrium constant ( K p ) and degree of dissociation (α ) ? (2002, 3M) (a) Neither K p nor α changes (b) Both K p and α change (c) K p changes but α does not change (d) K p does not change but α changes 20 At constant temperature, the equilibrium constant ( K p ) for the decomposition reaction, N2 O4 r 2NO2 , is expressed x2 p by K p = , where, p = pressure, x = extent of (1 − x2 ) decomposition Which one of the following statement is true? (2001, 1M) 86 Chemical and Ionic Equilibrium (a) K p increases with increase of p (b) K p increases with increase of x (c) K p increases with decrease of x (d) K p remains constant with change in p and x (c) concentration of NH3 does not change with pressure (d) concentration of hydrogen is less than that of nitrogen 28 For the reaction, H2 ( g ) + I2 ( g ) r 2HI( g ) 21 When two reactants, A and B are mixed to give products, C and D, the reaction quotient, (Q ) at the initial stages of the reaction (2000) (b) decreases with time (d) increases with time 22 For the reversible reaction, N2 ( g ) + 3H2 ( g ) r 2NH3 ( g ) at 500° C , the value of K p is 1.44 × 10–5 when partial pressure is measured in atmosphere The corresponding value of K c with concentration in mol/L is (2000, S, 1M) 1.44 × 10−5 (0.082 × 500)−2 (b) 1.44 × 10−5 (8.314 × 773)−2 (c) 1.44 × 10−5 (0.082 × 773)2 (d) 1.44 × 10–5 (0.082 × 773)−2 Objective Questions II (One or more than one correct option) 29 For a reaction, A P, the plots of [A] and [P] with time at temperatures T1 and T2 are given below 10 T2 T1 23 For the chemical reaction, X ( g ) + Y ( g ) r X 3Y ( g ) the amount of X 3Y at equilibrium is affected by (1999, 2M) 24 For the reaction , CO( g ) + H2 O( g ) r CO2 ( g ) + H2 ( g ) , at a given temperature, the equilibrium amount of CO2 ( g ) can be increased by (1998) Time If T2 > T1 , the correct statement(s) is are (Assume ∆H s and ∆S s are independent of temperature and ratio of ln K at T1 to ln K at T2 is greater than T2 / T1 Here H , S , G and K are enthalpy, entropy, Gibbs energy and equilibrium constant, respectively.) (2018 Adv.) (a) ∆H s < 0, ∆S s < (b) ∆Gs < 0, ∆H s > (c) ∆Gs < 0, ∆S s < (d) ∆Gs < 0, ∆S s > 30 The % yield of ammonia as a function of time in the reaction, N2( g )+ 3H2( g ) w NH3( g ); ∆H < % yield 25 One mole of N2 O4 ( g ) at 300 K is kept in a closed container under one atmosphere It is heated to 600 K when 20% by mass of N2 O4 ( g ) decomposes to NO2 (g) The resultant pressure is (1996, 1M) (2015 Adv.) (d) 1.0 atm If this reaction is conducted at ( p , T1 ), with T2 > T1 the % yield by of ammonia as a function of time is represented by (1985, 1M) (a) Pb(NO3 )2 (aq) + 2NaI (aq) = PbI2 (s) + 2NaNO3 (aq) (b) AgNO3 (aq) + HCl (aq) = AgCl (s) + HNO3 (aq) (c) 2Na (s) + 2H2O (l ) = 2NaOH (aq) + H2 (g ) (d) KNO3 (aq) + NaCl (aq) = KCl (aq) + NaNO3 (aq) (a) T2 T1 (b) Time T1 T2 (d) % yield % yield (c) T2 Time 27 Pure ammonia is placed in a vessel at a temperature where its dissociation constant (α ) is appreciable At equilibrium, N2 + 3H2 s 2NH3 (1984, 1M) T1 % yield 26 An example of a reversible reaction is T1 Time % yield (c) 2.0 atm at( p, T1 ) is given below (a) adding a suitable catalyst (b) adding an inert gas (c) decreasing the volume of the container (d) increasing the amount of CO(g ) (b) 2.4 atm T1 T2 Time (a) temperature and pressure (b) temperature only (c) pressure only (d) temperature, pressure and catalyst (a) 1.2 atm - 10 [A]/(mol L–1) (a) (1981, 1M) (a) total pressure (b) catalyst (c) the amount of H2 and I2 present (d) temperature [P]/(mol L–1) (a) is zero (c) is independent of time the equilibrium constant K p changes with T2 T1 (a) K p does not change significantly with pressure (b) α does not change with pressure Time Time Chemical and Ionic Equilibrium 87 31 The initial rate of hydrolysis of methyl acetate (1 M) by a weak acid (HA, 1M) is 1/100th of that of a strong acid (2013 Adv.) (HX, 1M), at 25°C The K a (HA) is (a) × 10−4 (b) × 10−5 (c) × 10−6 (d) × 10−3 I 32 The equilibrium Cu r Cu + Cu (c) SCN− (d) CN− forward reaction at constant temperature is favoured by (1991, 1M) (a) introducing an inert gas at constant volume (b) introducing chlorine gas at constant volume (c) introducing an inert gas at constant pressure (d) increasing the volume of the container (e) introducing PCl at constant volume 40 The rate of an exothermic reaction increases with increasing temperature (1993, 1M) (1987, 1M) 42 If equilibrium constant for the reaction, A2 + B2 r AB, is K, then for the backward reaction 1 AB r A2 + B2 , the equilibrium constant is 2 K (1984, 1M) 43 When a liquid and its vapour are at equilibrium and the pressure is suddenly decreased, cooling occurs 34 The equilibrium SO2 Cl ( g ) r SO2 ( g ) + Cl ( g ) is attained at 25° C in a closed container and an inert gas, helium is introduced Which of the following statements are correct? (1989, 1M) (a) Concentration of SO2 ,Cl and SO2Cl change (b) More chlorine is formed (c) Concentration of SO2 is reduced (d) None of the above 35 When NaNO3 is heated in a closed vessel, oxygen is liberated and NaNO2 is left behind At equilibrium, (1994, 1M) True/False 41 Catalyst makes a reaction more exothermic 33 For the reaction, PCl ( g ) r PCl ( g ) + Cl ( g ) the (1986, 1M) (a) addition of NaNO2 favours reverse reaction (b) addition of NaNO3 favours forward reaction (c) increasing temperature favours forward reaction (d) increasing pressure favours reverse reaction 36 For the gas phase reaction, ( ∆H = −32.7 kcal) C2 H4 + H2 r C2 H6 carried out in a vessel, the equilibrium concentration of C2 H4 can be increased by (1984, 1M) (a) increasing the temperature (b) decreasing the pressure (c) removing some H2 (1984, 1M) Subjective Questions 44 (a) In the following equilibrium N2O4 (g ) r 2NO2 (g ) when moles of each are taken, the temperature is kept at 298 K the total pressure was found to be 20 bar Given that ∆G °f (N2O4 ) = 100 kJ, ∆G °f (NO2 ) = 50 kJ (i) Find ∆G of the reaction (ii) The direction of the reaction in which the equilibrium shifts (b) A graph is plotted for a real gas which follows van der Waals’ equation with pVm taken on Y-axis and p on X-axis Find the (2004, 4M) intercept of the line where Vm is molar volume 45 When 3.06 g of solid NH4 SH is introduced into a two litre evacuated flask at 27° C, 30% of the solid decomposes into gaseous ammonia and hydrogen sulphide (i) Calculate K c and K p for the reaction at 27°C (ii) What would happen to the equilibrium when more solid NH4SH is introduced into the flask? (1999, 7M) 46 (a) The degree of dissociation is 0.4 at 400 K and 1.0 atm for the gaseous reaction PCl r PCl + Cl Assuming ideal behaviour of all the gases, calculate the density of equilibrium mixture at 400 K and 1.0 atm (relative atomic mass of P = 31.0 and Cl = 35.5) (b) Given, [Ag(NH3 )+2 ] r Ag+ + 2NH3 , K c = 6.2 × 10−8 and K sp of AgCl (d) adding some C2H6 = 1.8 × 10−10 at 298 K Fill in the Blanks 37 For a gaseous reaction 2B → A, the equilibrium constant K p is …… to/than K c (1996, 1M) equilibrium constant K p and K c are related by II (2011) (b) Cl − ten-fold increase in pressure on the reaction, N2 ( g ) + 3H2 ( g ) r 2NH3 ( g ) at equilibrium, results in in K p 39 For a given reversible reaction at a fixed temperature, in aqueous medium at 25° C shifts towards the left in the presence of (a) NO −3 38 A (1997 C, 1M) If ammonia is added to a water solution containing excess of AgCl(s) only Calculate the concentration of the complex in 1.0 M aqueous ammonia (1998, 3M+5M) 88 Chemical and Ionic Equilibrium 53 The equilibrium constant of the reaction 47 The progress of reaction, A2 ( g ) + B2 ( g ) r AB ( g ) at 100°C is 50 If a one litre flask containing one mole of A2 is connected to a two litre flask containing two moles of B2 , how many moles of AB will be formed at 373 K? (1985, 4M) (Concentration/mol L ) A r nB with time, is represented in fig use given below 0.5 54 One mole of N2 and moles of PCl are placed in a 100 L vessel heated to 227°C The equilibrium pressure is 2.05 atm Assuming ideal behaviour, calculate the degree of dissociation for PCl and K p for the reaction, 0.3 0.1 PCl ( g ) r PCl ( g ) + Cl ( g ) Time/h 55 One mole of nitrogen is mixed with three moles of hydrogen in a four litre container If 0.25 per cent of nitrogen is converted to ammonia by the following reaction Determine : (i) the value of n (ii) the equilibrium constant, K and (iii) the initial rate of conversion of A (1984, 6M) N2 ( g ) + 3H2 ( g ) r 2NH3 ( g ), then (1994, 3M) 48 0.15 mole of CO taken in a 2.5 L flask is maintained at 750 K along with a catalyst so that the following reaction can take place: CO ( g ) + 2H2 ( g ) r CH3 OH( g ) Hydrogen is introduced until the total pressure of the system is 8.5 atm at equilibrium and 0.08 mole of methanol is formed Calculate (i) K p and K c and (ii) the final pressure if the same amount of CO and H2 as before are used, but with no catalyst so that the reaction does not take place (1993, 5M) 49 For the reaction, CO( g ) + 2H2 ( g ) r CH3 OH( g ) hydrogen gas is introduced into a five litre flask at 327° C, containing 0.2 mole of CO( g ) and a catalyst, until the pressure is 4.92 atm At this point 0.1 mole of CH3 OH( g ) is formed Calculate the equilibrium constant, K p and K c (1990, 5M) 50 The equilibrium constant K p of the reaction, 2SO2 ( g ) + O2 ( g ) r 2SO3 ( g ) is 900 atm at 800 K A mixture containing SO3 and O2 having initial pressure of and atm respectively is heated at constant volume to equilibrate Calculate the partial pressure of each gas at 800 K (1989, 3M) 51 N2 O4 is 25% dissociated at 37° C and one atmosphere pressure Calculate (i) K p and (ii) the percentage dissociation at 0.1 atm and 37° C (1988, 4M) 52 At a certain temperature, equilibrium constant ( K c ) is 16 for the reaction; SO2 ( g ) + NO2 ( g ) r SO3 ( g ) + NO( g ) If we take one mole each of all the four gases in a one litre container, what would be the equilibrium concentrations of NO and NO2 ? (1987, 5M) calculate the equilibrium constant, K c in concentration units What will be the value of K c for the following equilibrium? (1981, 4M) N2 ( g ) + H2 ( g ) r NH3 ( g ) 2 Passage Based Questions Thermal decomposition of gaseous X to gaseous X at 298 K takes place according to the following equation: X (g ) s 2X (g ) The standard reaction Gibbs energy, ∆ r G°, of this reaction is positive At the start of the reaction, there is one mole of X and no X As the reaction proceeds, the number of moles of X formed is given by β Thus, βequilibrium is the number of moles of X formed at equilibrium The reaction is carried out at a constant total pressure of bar Consider the gases to behave ideally (Given, R = 0.083 L bar K −1 mol −1 ) 56 The equilibrium constant K p for this reaction at 298 K, in terms of β equilibrium is (a) (c) β 2equilibrium − β equilibrium β 2equilibrium − β equilibrium (2016 Adv.) (b) β 2equilibrium − β 2equilibrium (d) β 2equilibrium − β 2equilibrium 57 The incorrect statement among the following for this reaction, is (2016 Adv.) (a) Decrease in the total pressure will result in the formation of more moles of gaseous X (b) At the start of the reaction, dissociation of gaseous X takes place spontaneously (c) β equilibrium = 0.7 (d) KC < Chemical and Ionic Equilibrium 89 Topic Ionic Equilibrium Objective Questions I (Only one correct option) 1/ 100 mL of 0.1M HCl is taken is a beaker and to it 100 mL of Ksp (a) S = 144 Ksp (b) S = 6912 0.1 M NaOH is added in steps of mL and the pH continuously measured Which of the following graphs correctly depicts the change in pH? (2020 Main, Sep II) Ksp (c) S = 929 1/ Ksp (d) S = 216 1/ 1/ 7 If K sp of Ag CO3 is × 10− 12 , the molar solubility of (a) pH Ag CO3 in 0.1 M AgNO3 is (b) pH vol of NaOH vol of NaOH (2019 Main, 12 Jan II) (a) × 10− 12 M (b) × 10− 13 M (c) × 10− 10 M (d) × 10− 11 M 20 mL of 0.1 M H2 SO4 solution is added to 30 mL of 0.2 M NH4 OH solution The pH of the resultant mixture is [pK b of (2019 Main, Jan I) NH4 OH = 4.7] (c) pH (a) 9.3 (c) 9.0 (d) pH (b) 5.0 (d) 5.2 An aqueous solution contains an unknown concentration of vol of NaOH vol of NaOH The molar solubility of Cd (OH)2 is 184 × 10−5 m in water The expected solubility of Cd(OH)2 in a buffer solution of (2019 Main, 12 April II) pH = 12 is (a) 184 × 10 −9 M (c) 6.23 × 10−11 M 2.49 (b) × 10−9 M 184 (d) 2.49 × 10−10 M What is the molar solubility of Al(OH)3 in 0.2 M NaOH solution? Given that, solubility product of Al(OH)3 = 2.4 × 10−24 (2019 Main, 12 April II) (a) × 10−19 (c) × 10−22 (b) 12 × 10−21 (d) 12 × 10−23 The pH of a 0.02 M NH4 Cl solution will be [Given K b (NH4 OH) = 10−5 and log = 0.301] (a) 4.65 (c) 5.35 (2019 Main, 10 April II) −9 (a) × 10 −9 M(b) × 10 (c) 11 × 10−9 M M (d) 10 × 10−10 M 10 Which of the following are Lewis acids? (a) PH3 and BCl (c) PH3 and SiCl (2018 Main) (2018 Main) (b) AlCl and SiCl (d) BCl and AlCl 11 Which of the following salts is the most basic in aqueous solution? (a) Al(CN)3 (c) FeCl (2018 Main) (b) CH3COOK (d) Pb(CH3COO)2 12 pK a of a weak acid (HA) and pK b of a weak base (BOH) are (2017 Main) I The pH of a mixture containing 400 mL of 0.1 M H2SO4 and 400 mL of 0.1 M NaOH will be approximately 1.3 II Ionic product of water is temperature dependent III A monobasic acid with K a = 10−5 has a pH = The degree of dissociation of this acid is 50% IV The Le-Chatelier’s principle is not applicable to common-ion effect (a) I, II and IV (c) I and II original concentration of Ba 2+ ? 3.2 and 3.4, respectively The pH of their salt (AB) solution is (b) 2.65 (d) 4.35 Consider the following statements The correct statements are Ba2 + When 50 mL of a M solution of Na2SO4 is added, BaSO4 just begins to precipitate The final volume is 500 mL The solubility product of BaSO4 is × 10−10 What is the (2019 Main, 10 April I) (b) II and III (d) I, II and III If solubility product of Zr3 (PO4 )4 is denoted by K sp and its molar solubility is denoted by S , then which of the following relation between S and K sp is correct? (2019 Main, April I) (a) 7.2 (c) 7.0 (b) 6.9 (d) 1.0 13 How many litres of water must be added to L of an aqueous solution of HCl with a pH of to create an aqueous solution with pH of 2? (2013 Main) (a) 0.1 L (c) 2.0 L (b) 0.9 L (d) 9.0 L 14 Solubility product constant ( K sp ) of salts of types MX , MX and M X at temperature ‘T ’ are 4.0 × 10−8 , 3.2 × 10−14 and 2.7 × 10−15 , respectively Solubilities (mol dm− ) of the salts at temperature ‘T ’ are in the order (2008, 3M) (a) MX > MX > M X (b) M X > MX > MX (c) MX > M X > MX (d) MX > M X > MX 90 Chemical and Ionic Equilibrium 15 2.5 mL of 2/5 M weak monoacidic base (K b = × 10− 12 at 25°C) is titrated with 2/15 M HCl in water at 25°C The concentration of H+ at equivalence point is (K w = × 10− 14 at 25°C) (2008, 3M) (a) 3.7 × 10 (c) 3.2 × 10 − 13 −2 (d) 2.7 × 10 M M (b) 10−5 M (Ag+ ) and 10−5 M (Cl − ) M (c) 10−6 M (Ag+ ) and 10−6 M (Cl − ) HCl and the solution is diluted to one litre, resulting hydrogen ion concentration is (2005, 1M) (a) 1.6 × 10− 11 (b) × 10− 11 (c) × 10− (d) × 10− 17 HX is a weak acid ( K a = 10−5 ) It forms a salt NaX (0.1M) on reacting with caustic soda The degree of hydrolysis of NaX is (2004, 1M) (b) 0.0001% (c) 0.1% (d) 0.5% 18 A solution which is 10−3 M each in Mn 2+ , Fe2+ , Zn 2+ and Hg 2+ is treated with 10−16 M sulphide ion If K sp of MnS, FeS, ZnS and HgS are 10−15 , 10−23 , 10−20 and 10−54 respectively, which one will precipitate first? (2003, 1M) (a) FeS (b) MgS (c) HgS (d) ZnS 19 Identify the correct order of solubility of Na S, CuS and ZnS in aqueous medium (a) CuS > ZnS > Na 2S (c) Na 2S > CuS > ZnS (2002) (b) ZnS > Na 2S > CuS (d) Na 2S > ZnS > CuS 20 For a sparingly soluble salt A p Bq , the relationship of its solubility product ( Ls ) with its solubility (S) is (a) Ls = S p + q ⋅ pp ⋅ qq (b) Ls = S p + q ⋅ pq ⋅ qp (c) Ls = S (d) Ls = S pq ⋅ p ⋅q p q pq ⋅ ( p q) (2001, 1M) the order (1999, 2M) 23 Amongst the following hydroxides, the one which has the lowest value of K sp at ordinary temperature (about 25° C) is (1990, 1M) (d) Be(OH)2 24 Which of the following is the strongest acid? (d) SO2 (OH) 27 The compound that is not a Lewis acid is (a) BF3 (b) AlCl 28 The conjugate acid of (a) NH3 (c) BeCl NH–2 (1985, 1M) (d) SnCl is (1985, 1M) (c) NH+4 (b) NH2OH (d) N2H4 29 The best indicator for detection of end point in titration of a weak acid and a strong base is (1985, 1M) (a) methyl orange (3 to 4) (b) methyl red (5 to 6) (c) bromothymol blue (6 to 7.5) (d) phenolphthalein (8 to 9.6) 30 A certain weak acid has a dissociation constant of 1.0 × 10−4 The equilibrium constant for its reaction with a strong base is −4 (a) 1.0 × 10 −10 10 (c) 1.0 × 10 (b) 1.0 × 10 (d) 1.0 × 1014 31 A certain buffer solution contains equal concentration of X − and HX The K b for X − is 10−10 The pH of the buffer is (b) (c) 10 (d) 14 equal volumes of which of the following are mixed? (a) 100 mL of (M/10) HCl + 100 mL of (M/10) NaOH (b) 55 mL of (M/10) HCl + 45 mL of (M/10) NaOH (c) 10 mL of (M/10) HCl + 90 mL of (M/10) NaOH (d) 75 mL of (M/5) HCl + 25 mL of (M/5) NaOH (c) SO(OH)2 (a) unionised in the small intestine and in the stomach (b) completely ionised in the small intestine and in the stomach (c) ionised in the stomach and almost unionised in the small intestine (d) ionised in the small intestine and almost unionised in the stomach 32 The precipitate of CaF2 , ( K sp = 1.7 × 10−10 ) is obtained, when (1992, 1M) (b) ClO2 (OH) gastric juice in human stomach is about 2-3 and the pH in the small intestine is about Aspirin will be (1988, 1M) (a) 22 Which of the following solutions will have pH close to 1.0 ? (a) ClO3 (OH) 26 The pK a of acetyl salicylic acid (aspirin) is 3.5 The pH of (1984, 1M) (a) NaCl < NH4Cl < NaCN < HCl (b) HCl < NH4Cl < NaCl < NaCN (c) NaCN < NH4Cl < NaCl < HCl (d) HCl < NaCl < NaCN < NH4Cl (c) Ba(OH)2 (d) 10−10 M (Ag+ ) and 10−10 M (Cl − ) (1984, 1M) (p + q) 21 The pH of 0.1 M solution of the following salts increases in (a) Mg(OH)2 (b) Ca(OH)2 (a) 10−4 M (Ag+ ) and 10−4 M (Cl − ) −2 16 CH3 NH2 (0.1 mole, K b = × 10− ) is added to 0.08 mole of (a) 0.01% precipitation of AgCl ( K sp = 1.8 × 10−10 ) will occur only with (1988, 1M) −7 (b) 3.2 × 10 M 25 When equal volumes of the following solutions are mixed, (1989, 1M) (1982, 1M) (a) 10 −4 M Ca 2+ −4 + 10 MF − (c) 10−5 M Ca 2+ + 10−3 M F− (b) 10 −2 M Ca 2+ + 10−3 M F− (d) 10−3 M Ca 2+ + 10−5 M F− 33 An acidic buffer solution can be prepared by mixing the solution of (1981, 1M) (a) acetate and acetic acid (b) ammonium chloride and ammonium hydroxide (c) sulphuric acid and sodium sulphate (d) sodium chloride and sodium hydroxide 34 Of the given anions, the strongest base is (a) ClO− (b) ClO−2 (c) ClO−3 (1981, 1M) (d) ClO−4 35 At 90°C, pure water has [H3 O+ ] as 10−6 mol L−1 What is the value of K w at 90°C ? (a) 10−6 (b) 10−12 (1981, 1M) (c) 10−14 (d) 10−8 Chemical and Ionic Equilibrium 91 36 The pH of 10−8 M solution of HCl in water is (a) (b) −8 (c) between and (d) between and (1981, 1M) H 2S What is the minimum molar concentration (M) of H + required to prevent the precipitation of ZnS? Use K sp ( ZnS) = 25 × 10−22 Objective Questions II (One or more than one correct option) mol/L) of Ag CrO4 in a 0.1 M AgNO3 solution is (b) 1.1 × 10−10 (c) 1.1 × 10−12 (d) 1.1 × 10−9 (2013 Adv.) CH3 COONa of identical concentrations are provided The pair(s) of solutions which form a buffer upon mixing is(are) HNO3 and CH3COOH KOH and CH3COONa HNO3 and CH3COONa CH3COOH and CH3COONa (2010) (1999, 3M) (b) sodium acetate and HCl in water (c) ammonia and ammonium chloride in water (d) ammonia and sodium hydroxide in water (1998, 2M) (c) Autoprotolysis constant of water increases with temperature (d) When a solution of a weak monoprotic acid is titrated against 1 a strong base, at half-neutralisation point pH = pK a 2 Numerical Answer Type Questions 41 A solution of 0.1 M weak base (B) is titrated with 0.1 M of a strong acid (HA) The variation of pH of the solution with the volume of HA added is shown in the figure below What is the pK b of the base? The neutralisation reaction is given by 12 10 pH 0.06 M Fe2 + ( aq ) and M S2 − ( aq ) solutions are mixed, the equilibrium concentration of Fe2 + ( aq ) is found by Y × 10− 17 M The value of Y is (2019 Adv.) When equal volumes of 44 The solubility of a salt of weak acid (AB) at pH is ionisation constant of HB ( K a ) = × 10−8 Match the Column (2020 Adv.) Note Degree of dissociation (α ) of weak acid and weak base is