Preview IIT JEE main complete chemistry k l kapoor MHE mc graw hill education by k l kapoor TMH

Preview IIT JEE Main Complete Chemistry K L Kapoor MHE Mc Graw Hill Education by K L Kapoor TMH Preview IIT JEE Main Complete Chemistry K L Kapoor MHE Mc Graw Hill Education by K L Kapoor TMH Preview IIT JEE Main Complete Chemistry K L Kapoor MHE Mc Graw Hill Education by K L Kapoor TMH Preview IIT JEE Main Complete Chemistry K L Kapoor MHE Mc Graw Hill Education by K L Kapoor TMH Preview IIT JEE Main Complete Chemistry K L Kapoor MHE Mc Graw Hill Education by K L Kapoor TMH Preview IIT JEE Main Complete Chemistry K L Kapoor MHE Mc Graw Hill Education by K L Kapoor TMH

C O M P L E T E

CHEMISTRY

McGraw Hill Education (India) Private Limited

CHENNAI

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K.L Kapoor

Formerly Associate Professor,

Hindu College, University of Delhi, Delhi

C O M P L E T E

CHEMISTRY

444/1, Sri Ekambara Naicker Industrial Estate, Alapakkam, Porur, Chennai-600116

Complete Chemistry—JEE Main

Copyright © 2018, McGraw Hill Education (India) Private Limited

No part of this publication may be reproduced or distributed in any form or by any means, electronic, mechanical, photocopying, recording, or otherwise or stored in a database or retrieval system without the prior written permission of the publishers The program listings (if any) may be entered, stored and executed in a computer system, but they may not be reproduced for publication

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To Our Readers

How to Crack the JEE

To help students preparing for the JEE Main, there was need for a book which included a variety of Multiple Choice

Questions (MCQs) designed on the basis of the prescribed syllabus for this examination This book is an attempt

in this direction and will help students in developing a strong foundation and enough confidence to take the JEE Main

The various topics of chemistry may be classified into three branches—Physical, Inorganic and Organic The book covers these three branches in 29 chapters Each chapter starts with the synopsis of theory followed by MCQs along with answers and hints and solutions to arrive at correct answers Wherever needed, the chapter is divided into

sections to cover the subject in easily understandable portions to help in better grasping of the subject matter Each

section/chapter ends with MCQs from the previous years' AIEEE and JEE Main This will help students in getting an

idea about the types and levels of questions asked in this competitive examination The answers and solutions to these questions are provided separately, immediately after the questions

The analyses of these papers (provided on the next page) reveals that in most of the cases, one question is asked from each chapter and the entire syllabus is covered in the examination paper This book provides extensive coverage

of the theory as well as the associated MCQs The contents of each chapter are covered in various sections At the end

of all the sections, more extensive MCQs based on the Entire Chapter along with their solutions are also included

It will be beneficial for the students to adopt the present book as the reference book along with their main text book The MCQs included in this book should be attempted along with the class-room teaching of the subject matter A regular and periodical review of the theory and MCQs from this book will help students in gaining enough confidence to appear

in the JEE Main and enable them to face the challenge of successfully clearing this examination

From the analyses of previous years’ question papers, a pattern of predominant topics emerges on which students should pay more attention while preparing for the examination These are:

Physical Chemistry—Entire portion

Inorganic Chemistry has been thoroughly revised and updated—Chemical families—perioidic properties,

structures of compounds containing Si, N, P, S, halogens and inert gases, d-block elements and coordination chemistry

Organic Chemistry—Stereoisomerism, SN1 and S22 Reactions, Reactions involving rearrangement, Chemistry of typical reactions shown by phenols, aldehydes and ketones and amines, relative acidity/basicity of phenolic, Carboxylic acids and amines, polymers, carbohydrates, stereochemistry involved in halogenation of alkenes and dehalogenation of halogenated compounds to give alkene, reactions involving Grignard reagent and diazonium salt

K.L Kapoor

12 General Principles and

(Contd )

27 Biomolecules & Biological

About JEE Main

1 Introduction and Scheme of Examination

The Joint Entrance Examination from the year 2013 for admission to the undergraduate programmes in Engineering

is being held in two parts, JEE-Main and JEE-Advanced Only the top 1,50,000 candidates (including all categories) based on performance in JEE Main will qualify to appear in the JEE Advanced examination Admissions to IITs will be based only on category-wise All India Rank (AIR) in JEE Advanced, subject to condition that such candidates are in the top 20 percentile categories Admission to NITs will be based on 40% weightage for performance in Class XII board marks (normalized) and the remainder 60% weightage would be given to performance in JEE Main and a combined All India Rank (AIR) would be decided accordingly

In case any State opts to admit students in the engineering Colleges affiliated to state Universities where States require separate merit list to be provided based on relative weightages adopted by the states, then the merit list shall be prepared with such relative weightages as may be indicated by States

2 Eligibility Criteria and List of Qualifying Examinations for JEE(Main) Exam

The minimum academic qualification for appearing in JEE(Main) is that the candidate must have passed in final examination of 10+2 (Class XII) or its equivalent referred to as the qualifying examination (see below)

However, admission criteria in the concerned institution/university will be followed as prescribed by concerned university/institution and as per the guidelines & criteria prescribed by AICTE

Qualifying Examinations

List of Qualifying Examinations

(i) The +2 level examination in the 10+2 pattern of examination of any recognized Central/State Board of Secondary Examination, such as Central Board of Secondary Education, New Delhi, and Council for Indian School Certificate Examination, New Delhi

(ii) Intermediate or two-year Pre-University Examination conducted by a recognized Board/University

(iii) Final Examination of the two-year course of the Joint Services Wing of the National Defence Academy

(iv) Any Public School/Board/University Examination in India or in foreign countries recognized by the Association

of Indian Universities as equivalent to 10+2 system

3 Pattern of Examination

Subject combination for each paper and type of questions in each paper are given below:

Paper 1 Physics, Chemistry & Mathematics Objective type questions with equal

weightage to Physics, Chemistry &

Mathematics

3 Hours

Aptitude Test – Part II & Objective type questionsDrawing Test – Part III questions to test Drawing AptitudeRequirement of papers for different courses is given in the table below:

Scoring and Negative Marking

There will be objective type questions with four options having single correct answer For each incorrect response, one fourth (1/4) of the total marks allotted to the question would be deducted No deduction from the total score will, however, be made if no response is indicated for an item in the answer sheet

SECTION—A Physical Chemistry

Unit 1 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 2 States of Matter

Classification of matter into solid, liquid and gaseous states

Gaseous State: Measurable properties of gases; Gas laws—Boyle’s law, Charles’ 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 3 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, y and y2, concept

of atomic orbitals as one electron wave functions; Variation of y and y2 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 4 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 5 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; DS of the universe and DG of the system as criteria

for spontaneity, DGo (Standard Gibbs energy change) and equilibrium constant

Unit 6 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 7 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 (Kp and Kc) and

their significance, significance of DG and DGo 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 8 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

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; Nemst equation and its applications; Relationship between cell potential and Gibbs’ energy change; Dry cell and lead accumulator; Fuel cells; Corrosion and its prevention

Unit 9 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 temperature on rate of reactions—Arrhenius theory, activation energy and its calculation, collision theory of bimolecular gaseous reactions (no derivation)

Unit 10 Surface Chemistry

Adsorption: Physisorption and chemisorption and their characteristics, factors affecting adsorption of gases on

solids—Freundlich 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, macro-molecular and associated colloids (micelles), preparation and properties of colloids—Tyndall effect, Brownian movement, electrophoresis, dialysis, coagulation and flocculation; Emulsions and their characteristics

SECTION—B Inorganic ChemistryUnit 11 Classificaton of Elements and Periodicity in Properties

Modem periodic law and present form of the periodic table, s, p, d and f block elements, periodic trends in properties

of elements atomic and ionic radii, ionization enthalpy, electron gain enthalpy, valence, oxidation states and chemical reactivity

Unit 12 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 13 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 14 s-Block Elements (Alkali and Alkaline Earth Metals)

Group - 1 and 2 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 15 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

Groupwise study of the p-block elements

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

Occurrence and uses of noble gases; Structures of fluorides and oxides of xenon

Unit 16 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 17 Co-ordination Compounds

Introduction to co-ordination compounds, Werner’s theory; ligands, co-ordination number, denticity, chelation; IUPAC nomenclature of mononuclear co-ordination compounds, isomerism; Bonding Valence bond approach and basic ideas of Crystal field theory, colour and magnetic properties; importance of co-ordination compounds (in qualitative analysis, extraction of metals and in biological systems)

Unit 18 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

Purification: Crystallization, 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 20 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 21 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.

Alkynes: 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 22 Organic Compounds Containing Halogens

General methods of preparation, properties and reactions; Nature of C-X bond; Mechanisms of substitution reactions.Uses/environmental effects of chloroform, iodoform, freons and DDT

Unit 23 Organic Compounds Containing Oxygen

General methods of preparation, properties, reactions and uses

Alcohols, Phenols and Ethers

Alcohols: Identification of primary, secondary and tertiary alcohols; mechanism of dehydration.

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 a - hydrogen, aldol

condensation, Cannizzaro reaction, Haloform reaction; Chemical tests to distinguish between aldehydes and Ketones

Carboxylic Acids: Acidic strength and factors affecting it.

Unit 24 Organic Compounds Containing Nitrogen

General methods of preparation, properties, reactions and uses

Amines: Nomenclature, classification, structure basic character and identi-fication of primary, secondary and tertiary

amines and their basic character

Diazonium Salts: Importance in synthetic organic chemistry.

Unit 25 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 26 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 a - 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 27 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 28 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

∑ Chemistry involved in the preparation of the following:

Inorganic compounds: Mohr’s salt, potash alum

Organic compounds: Acetanilide, p-nitroacetanilide, aniline yellow, iodoform

∑ Chemistry involved in the titrimetric excercises - Acids bases and the use of indicators, oxalic-acid vs KMnO4, Mohr’s salt vs KMnO4

∑ Chemical principles involved in the qualitative salt analysis:

Cations: Pb2+, Cu2+, AI3+, Fe3+, Zn2+, Ni2+, Ca2+, Ba2+, Mg2+, NH4+

Anions: CO32–, S2–, SO42–, NO2– , NO3– , CI – , Br–, I–

(Insoluble salts excluded)

∑ Chemical principles involved in the following experiments:

1 Enthalpy of solution of CuSO4

2 Enthalpy of neutralization of strong acid and strong base

3 Preparation of lyophilic and lyophobic sols

4 Kinetic study of reaction of iodide ion with hydrogen peroxide at room temperature

S ection -1 Physical Quantities 1.1

S ection -2 Significant Figures (or Digits) 1.7

S ection -3 Atomic and Molecular Masses 1.9

S ection -4 Laws of Chemical Combination 1.14

S ection -5 Composition of a Solution 1.19

U nit -1 Gaseous State 2.1

S ection -1 Ideal Gases 2.1

S ection -2 Kinetic-Molecular Theory of Gases 2.7

S ection -3 Real Gases 2.12

U nit -2 Liquid State 2.37

U nit -3 Solid State 2.43

S ection -1 Crystal Systems 2.43

S ection -2 Closest Packings of Atoms 2.54

S ection -3 Structures of Ionic Compounds 2.64

S ection -4 Imperfection in Solids 2.74

S ection -1 Development of Structure of Atom 3.1

S ection -2 Quantum-Mechanical Model of Atom 3.8

S ection -1 Bond Formation and VSEPR Theory 4.1

S ection -2 VB and MO Theories 4.13

S ection -1 Composition of a Solution 5.1

S ection -2 Liquid Solutions 5.7

S ection -3 Colligative Properties 5.16

S ection -1 Basic Definitions and First Law of Thermodynamics 6.1

S ection -2 Thermochemistry 6.9

S ection -3 Criteria of Spontaneity 6.17

U nit -1 Chemical Equilibrium 7.1

Unit -2 Ionic Equilibrium 7.27

S ection -1 Concepts of Acids and Bases 7.27

S ection -2 The pH Scale and pH of Acid and Base Solutions 7.32

S ection -3 Hydrolysis of Salts 7.39

S ection -4 Buffer Solutions 7.44

S ection -5 Solubility Product 7.48

S ection -6 Acid-Base Indicators 7.53

U nit -1 Redox Reactions and Electrolysis 8.1

S ection -1 Redox Reactions 8.1

S ection -2 Electrolytic Cell 8.8

U nit -2 Electrolytic Conduction 8.18

U nit -3 Galvanic Cells 8.30

12 General Principles and Processes of Isolation of Metals 12.1–12.8

14 S-Block Elements (Alkali and Alkaline Earth Metals) 14.1–14.25

The Group 1 Elements – Alkali Metals 14.1

The Group 2 Element – Alkaline Earth Metals 14.12

15 Study of the p-Block Elements (Groups 13, 14 and 15) 15.1–15.60

The Group 13 Elements – Boron Family 15.1

The Group 14 Elements – Carbon Family 15.16

The Group 15 Elements 15.35

16 Study of the p-Block Elements (Groups 16, 17 and 18) 16.1–16.48

The Group 16 Elements 16.1

The Group 17 Elements 16.22

The Group 18 Elements 16.41

20 Purification and Characterization of Organic Compounds 20.1–20.9

S ection -5 Sources of Hydrocarbons 22.58

23 Organic Compounds Containing Halogens (Haloalkanes and Haloarenes) 23.1–23.14

S ection -1 Haloalkanes 23.1

S ection -2 Haloarenes 23.5

24 Organic Compounds Containing Oxygen (Alcohols, Phenols, Ethers, Aldehydes,

S ection -1 Alcohols 24.1

S ection -2 Phenols 24.16

S ection -3 Ethers 24.31

S ection -4 Aldehydes and Ketones 24.34

S ection -5 Carboxylic Acids 24.55

25 Organic Compounds Containing Nitrogen (Cyanides, Isocyanides,

S ection -1 Cyanides and Isocyanides 25.1

S ection -2 Nitro Compounds 25.2

S ection -3 Amines 25.4

Physical Quantities

Section 1

In physical sciences, we commonly deal with quantities such as pressure, volume, mass, temperature, current, etc These quantities are known as physical quantities A physical quantity has two components, namely, numerical value and its unit, and is written as

Physical quantity = (Numerical value) (Unit)

The International Union of Pure and Applied Chemistry (IUPAC) has recommended the use of seven physical quantities having their own dimensions Their dimensions are completely independent of one another and it is for this reason, these are known as dimensionally independent physical quantities These physical quantities along with their recommended symbols are given in Table 1

Seven Base Physical Quantities

Physical quantity Symbol

Seven Base Physical Quantities

Physical quantity Name of SI unit Symbol for SI unit

The SI base units stated in Table 2 have been precisely

The metre is the length of path travelled by light in vacuum during a time interval of 1/299 792 458 of a second.The kilogram is the unit of mass; it is equal to the mass of the international prototype of the kilogram This prototype is a polished cylinder of platinum-iridium alloy which was chosen for its durability and resistance to corrosion The cylinder is kept at the International Bureau of Weights and Measures in a suburb of Paris, France

The second is the duration of 9 192 631 770 periods of the radiation corresponding to the transition between

of negligible cross-section, and placed 1 metre apart in vacuum, would produce between these conductors a force equal

to 2 ¥ 10–7 newton per metre of length

The kelvin, unit of thermodynamic temperature, is the fraction 1/273.16 of the thermodynamic temperature

of the triple point of water

The mole is the amount of substance of a system which contains as many elementary entities as there are atoms

Number of moles of hydrogen = 0.5 (or moles of hydrogen = 0.5 mol)Now suppose we have 1 kg of hydrogen We never speak or write it as

Number of kg of hydrogen = 1 since we are not accustomed to speak or write it like this Instead, we say

Mass of hydrogen = 1 kg Many such examples can be cited For example, we never say

Number of kg of apples = 1 Number of litre of milk = 1Number of dozen of apples = 1 Number of metre of a line = 1

A mole of a substance is a collection of 6.022 ¥ 1023 particles of that substance It is like a dozen which is a collection

of 12 articles When we do not refer to a collection of 12 articles as

Number of dozen of articles = 1

why should we refer to a collection of 6.022 ¥ 1023 articles as

Number of moles of articles =1?

So, some inconsistency has arisen in referring to these two quantities which basically belong to the same category

To avoid this, IUPAC has recommended the use of the phrase ‘amount of substance’ for a physical quantity whose unit

is ‘mol’ More precisely, the word ‘amount’ is exclusively reserved whenever the quantity is to be expressed in terms

of moles It is like using the words mass for kg, volume for m3 (or L) and length for metre Thus, the use of the phrase

‘number of moles equal to 0.5’ should be completely replaced by ‘amount of substance equal to 0.5 mol’ It may be mentioned that to write

Amount of substance = 0.5 g

is not acceptable as the word ‘amount’ is exclusively reserved for the unit ‘mol’ and not for ‘g’

entities of that substance The proportionality factor is the same for all substances and is equal to l/NA, where NA is Avogadro constant which has a value of

where N NA is Avogadro constant For example, 3.011.5 ¥ 1023

molecules of dihydrogen will contain

Amount of dihydrogen = 3.011 ¥ 1023 / (6.022 ¥ 1023 mol–1) = 0.5 mol

Physical quantities other than base physical quantities are known as derived physical quantities These may be expressed

in terms of base physical quantities by multiplication and division Table 3 records some of the derived physical quantities

SI derived units of a few physical quantities

Physical quantity SI unit Symbol for SI unit

area (length)2 square metre m2

volume (length)3 cubic metre m3

density mass/volume kilogram per cubic metre kg m–3

speed distance/time metre per second m s–1

acceleration speed/time metre per square second m s–2

(l/m)(Dq/Dt) — J kg–1 K–1

molar heat capacity (l/n)(Dq/Dt) — J K–1 mol–1

amount concentration amount of substance/volume of solution mole per cubic metre mol m–3

molality amount of substance/mass of solvent mole per kilogram mol kg–1

Some physical quantities have been assigned special names and symbols These are described in Table 4

Special names and symbols for certain SI derived units

Physical quantity Name of SI unit Symbol for SI unit

force mass ¥ accleration Newton N kg m s–2

pressure force/(length)2 Pascal Pa kg m–1 s–2 (=N m–2)

energy force ¥ length Joule J kg m2 s–2 (= N m)

power energy/time Watt W kg m2 s–3 (= J s–1)

electric charge current ¥ time Coulomb C A s

electric potential difference — Volt V kg m2 s–3 A–1 (= J A–1 s–1)electric resistance — Ohm W kg m2 s–3 A–2 (= V A–1)

electric conductance — Siemens S kg–1 m–2 s3 A2 (= A V–1 = W–1)

— Tesla T kg s–2 A–1

Similarly, the following expressions represent one and the same thing,

Note that writing an expression of the type V/cm3 = 25.0 is very convenient while writing the headings in tables and

as labels on the axes of graphs

A physical expression should also be dimensionally correct For example, the conversion expression of Celsius temperature to kelvin temperature may be written as

T = qC + 273.15This expression is numerically correct but not dimensionally as the unit of qc is °C and that of T is K One can add

or subtract two physical quantities if they have the same unit Thus, a correct conversion equation would be

T/K = qC/°C + 273.15 For example, for 25 °C, we would have

T/K= 25 °C/°C + 273.15 = 25 + 273.15 = 298.15 or T = 298.15 K

A few other examples are

log k k∞

Ê

ËÁ ˆ¯˜ = logÊËÁk A∞ˆ¯˜ -2 303. EaRT ; ln

h h∞

Note that the division by k°, h° and p° (which stand for the corresponding unit physical quantities) make the

expression within the logarithm brackets unitless

Other examples are

pH = –log {[H+]/mol dm–3}; pK°w = –log [Kw/(mol dm–3)2}; pKa° = –log {Ka/mol dm–3}The use of standard equilibrium constant K° (= K/(mol dm–3)Ân) would avoid the division by units

The conversion of one unit to another may be carried out by It involves the following steps

Arrange the identity so as to have Desired unit

Given unit = 1Multiply the given value by the above identity and simplify the expression

Convert 57.8 m into cm unit

ˆ

¯˜= 57.8 ¥ 102 cmConvert 1.5 g cm–3 in terms of kg m–3

Volume of one mole of an ideal gas at 27 °C and 1 atm pressure

-

-¥ = 2.463 ¥ 10–2 J Pa–1

J Pa–1 = (kg m2 s–2) (kg m–1 s–2)–1 = m3 and thus V = 2.463 ¥ 10–2 m3 = 2.463 ¥ 10–2 (10 dm)3 = 24.63 dm3

MULTIPLE CHOICE QUESTIONS ON SECTION 1

Identify the correct choice in the following questions

1 The symbol for SI unit kg m–1 s–2 is

2 Which one of the following is the base physical quantity?

3 Which of the following unit represent the joule unit?

4 The value of Avogardo constant is

(a) 6.022 ¥ 1022 (b) 6.022 ¥ 1023 atoms (c) 6.022 ¥ 1023 mol–1 (d) 6.022 ¥ 1022 mol–1

5 Which of the following conversion units is correct?

(a) 1 Pa = 1 kg m s–2 (b) 1 J = 1 kg m2 s–2 (c) 1 C = 1 A s–1 (d) 1 N = 1 kg m–1 s2

18 is

–15 is

8 One gigametre stands for

= = pressure The SI symbol of pressure is pascal (Pa)

2 Electric current is the base physical quantity

3 The unit Pa m3 represents joule

4 Avogadro constant is 6.022 ¥ 1023 mol–1

¥ acceleration Its unit will be

Significant Figures (or Digits)

3 Zeros at the end of a number without a decimal point are ambiguous For example, 7500 may have two or

digit and ending with the digit that has uncertain value

Number Number

0.0380 3 3.0805 50.3800 4 0.3805 4The counting of discrete variables (such as peas, pencils, erasers, and so on) leads to an exact

An algebraic operation may involve numerical quantities of

1 In an arithmetic operation involving addition and/or subtraction, the answer should include a factor of uncertainty equal to the maximum uncertainty present in the numbers being added and/or subtracted

3 In an arithmetic operation involving mixed manipulations (addition, subtraction, multiplication and division),

If the digit following the last digit to be retained is more than 5, the last digit to be retained is increased by one

If the digit following the last digit to be retained is less than 5, the last digit is left unchanged

For

digit by 1 if it is odd and leave it unchanged if it is even.

1 6 2 2 4

¨ involves maximum uncertainty

Æ rounded off to 16.2

6 7 6 9– 2 1 1 3

6 5 5 7 7

¨ involves maximum uncertainty

1308( )( )

= ÈÎÍ

digits Hence, the answer will be 0.54

MULTIPLE CHOICE QUESTIONS ON SECTION 2

Identify the correct choice in the following questions

-mm

4 The result will carry 2 digits after decimal

Atomic and Molecular Masses

Section 3

The IUPAC has recommended the following terms while dealing with atomic and molecular masses

The relative atomic mass of an element is the ratio of the average mass per

of an element to 1/12 of the mass of an atom of the nuclide carbon-12, i.e

Ar= mass of an atom

(1/12) mass of an atom of C12 (1)The relative molecular mass of a compound is the ratio of the average carbon-12, i.e

Mr = mass of a molecule(1/12) mass of an atom of C12 (2)

1 2 3

The quantities Ar and Mr are formerly known as atomic weight and molecular weight, respectively It may be noted that Ar and Mr carry no units as these are simply the ratio of two masses.

The quantity (1/12) mass of an atom 12

mu, unit; u, also commonly abbreviated as amu) Hence

1 amu = 1 u = ma C

12

(12 ) =(0.012 kg mol mol )

12

1

- 1 ¥ 23

-6 022 10) /(

Since the unit of mass (m) is kg or g, and that of amount of substance (n) is mol, it follows that the unit of molar

mass is kg mol–1 or g mol–1

If a system has N entities (atoms or molecules),

In words, the relative atomic (or molecular) mass is the numerical value of the molar mass expressed in g mol–1

We will have

Relative atomic mass of sodium = 23

Atomic mass of sodium = 23 u = 23 (1.66 ¥ 10–27 kg) = 3.82 ¥ 10–26 kg

Molar mass of sodium = 23 g mol–1

Relative molecular mass of carbon dioxide = 44

Molecular mass of carbon dioxide = 44 u = 44 (1.66 ¥ 10–27 kg) = 7.30 ¥ 10–26 kg

Molar mass of carbon dioxide = 44 g mol–1

Quite often, we do not specify the units of atomic, molecular and molar masses Hence, our statements are not exact and precise In other words, a mere replacement of the term weight by mass without carrying the associated unit does not complete our adoption of IUPAC recommendations So, whenever we write or speak of atomic, molecular and molar masses, we must state the associated units

It may be noted that the terms gram atomic weight, gram molecular weight, gram formular weight, gram atom, gram molecule, etc., are obsolete terms and should thus be abandoned

An atom or a molecule in nature exists in its natural isotopic composition The terms mentioned above for atomic or molecular masses refer to an average value of this natural isotopic composition For example, natural magnesium consists of three isotopes as mentioned in the following.

24Mg ma = 23.985 u 78.99% 25Mg ma = 24.986 u 10.00%

26Mg ma = 25.983 u 11.01%

The average atomic mass of Mg is

ma(Mg) = (0.789 9 ¥ 23.985 + 0.100 0 ¥ 24.986 + 0.110 1 ¥ 25.983) u = (18.946 + 2.498 6 + 2.861) u = 24.306 u

As mentioned earlier, 1 mol of a substance contains 6 022 ¥ 1023 constituent particles (atoms or molecules or ions) This fact is expressed by Avogadrao constant which has value of 6.022 ¥ 1023 mol–1 It is expressed by the symbol of NA

For a given amount of substance, say n, the number of constituent particles will be given as

N = n NA

For example, The number of molecules in 0.5 mol glucose is

N = (0.5 mol) (6.022 ¥ 1023 mol–1) = 3.011 ¥ 1023

The amount (n) of a substance in a given mass (m) may be calculated

from the expression

n = m M

where M is the molar mass of the substance.

The molar mass of a compound is sum of molar masses

of its constituent atoms For example, the molar mass of H2SO4 is

MH2SO4 = 2 MH + MS + 4 MO = (2 ¥ 1 + 32 + 4 ¥ 6) g mol–1 = 98 g mol–1

The per cent of H, S and O is H2SO4 may be calculated by the expression

Mass per cent of element = No of its atom × Molar mass of atom

Molar mass of the compoound ¥100For example, in H2SO4 we have

Mass per cent of H =2 1

¥

¥

-

to determine the molecular formula as indicated in the following

1 Take the mass of each element equal to its mass percentage and divide this by the corresponding molar mass (if the mass is taken in grams) or molecular mass (if the mass is taken in terms of atomic mass unit) This gives the relative amounts (if mass is taken in gram) or number of atoms (if mass is taken in atomic mass unit) of different elements present in the compound

2 Divide the amounts (or number of atoms) of elements by the smallest amount (or number of atoms) to give a simple relative ratio of atoms present in the compound.

3 If the relative ratio of atoms involves noninteger(s), then multiply all the simple ratios by a suitable number to get a whole number ratio for each element

4 An empirical formula is written by taking as many atoms as given by the whole number ratio

The molecular formula represents the actual number of atoms of each element present in a molecule of the compound The molecular formula is either the same as the empirical formula or a simple multiple of the empirical formula, i.e.,

Molecular formula = (Empirical formula)nwhere n = 1, 2, 3 The value of n is equal to the ratio of molar mass of the compound and molar empirical formula

mass, i.e

n = Molar mass

Molar empirical formula mass

–1

To determine empirical formula from the given data, we proceed as follows

Element Per cent Mass of

mol

- = 1 41. mol 1

1.41 mol = 1 ¥ 5 = 5Hence, the required empirical formula is C14H9Cl5

The above calculations in short may be carried out as follows

47.5

12 :

.: .

2 541

:: . : :

Empirical Formula molar mass = (14 ¥ 12 + 9 ¥ 1 + 5 ¥ 35.5) g mol–1 = 354.5 g mol–1

Since, the empirical formula molar mass is the same as the molar mass, the molecular formula of DDT is C14H9Cl5

A gaseous hydrocarbon gives upon combustion 3.08 g CO2 and 0.72 g H2O Determine its empirical formula

Mass of C in the mass mCO2= M

C CO CO

2 2

¥ Amount of C in mCO2 = M

m M

C CO CO C CO CO

2 2

¥

ÊËÁ

H

H O

H O H

H O

H O

2 2

2

2

¥

ÊËÁ

ˆ

¯˜ =For the given data, we have

MULTIPLE CHOICE QUESTIONS ON SECTION 3

Identify the correct choice in the following questions

(a) 1.661 ¥ 10–27 g (b) 1.661 ¥ 10–27 kg (c) 1.661 ¥ 10–25 g (d) 1.661 ¥ 10–25 kg

2 The relative atomic mass of sodium is 23 Which of the following statements is correct about sodium?(a) Atomic mass of sodium is 23 u

(b) Atomic mass of sodium is 3.82 ¥ 10–26 kg

(c) Molar mass of sodium is 23 g mol–1

(d) The number of atoms in 24 kg of sodium is 6.022 ¥ 1023

3 Chlorine exists as 35Cl (atomic mass = 34.9688 u; 75.77%) and 37Cl (atomic mass = 36.965 9 u; 24.23%) The atomic mass of natural existing chlorine is

3 ma = (0.757 7 ¥ 34.968 8 + 0.242 3 ¥ 36.965 9) u = 35.452 7 u

13 2

44

2 7 215

¥

5 Larger the amount of an element, larger the number of atoms

n(C) = m(C) / M(C) = 16 g/12 g mol–1 = 1 333 mol n(Na) = m(Na) / M(Na) = 20 g/ 23 g mol–1= 0.870 mol

n(S) = m(S) / M(S) = 45 g/ 32 g mol–1 = 1.406 mol n(N) = m(N)/ M(N) = 15 g/14 g mol–1 = 1.071 mol

2012

47 3314

26 0116

6 661: . : . : . Empirical formula CN2OH4

Molar empirical mass is 60 g mol–1 (same as the given molar mass) Hence, Molecular formula is CN2OH4

7 Mass per cent of Na Na

1 881

1 248

Molecular formula Fe2O3

9 Molar mass of CuSO4·5H2O = (63 + 32 + 4 ¥ 16 + 5 ¥ 18) g mol–1 = 249 g mol–1

Amount of CuSO4·5H2O in the given mass is n m

g mol mol Number of oxygen atoms = 9 ¥ (0.1 mol) (6.022 ¥ 1023 mol–1 ) = 5.42 ¥ 1023

11 9914

13 716

9 25

10 8

65 0619

.: . : .

:

4

Laws of Chemical Combination

Section 4

Based on the study of chemical reactions the following laws have been established

( ) This law states that the mass is conserved

For example, in the reaction 2H2(g) + O2(g) Æ 2H2O(g), we have

2 Volume of hydrogen combines with 1 volume of oxygen to give 2 volume of water

This law follows from the fact that under similar conditions of temperature and pressure, equal volumes of all gases contain equal number of particle (Avogadro’s hypothesis) The volume occupied by 1 mol gas at standard condition of temperature (0 °C) and pressure (1 atm), abbreviated as STP, is 22.4 L Thus, in the above example, 2 ¥ 22.4 L of H2

combines of 22.4 L of O2 to give 2 ¥ 22.4 L of H2O(g)

The laws of chemical combination may be explained on the basis of Dalton

1 Matter is composed of atoms which cannot be created or destroyed

2 All the atoms of one element are alike (i.e they have the same size, shape and mass) but are different from those of other element

3 In compound, atoms combine together in the ratio of small whole numbers and are held by chemical forces The smallest entity of a compound is known as a molecule

The branch of chemistry which deals with mass relationships in chemical reactions is called stoichiometry This branch

1 Conservation of mass 2 The relative masses of atoms 3 The concept of the mole

According to the law of conservation of mass, the total mass of the products formed in a chemical equation is equal to the total mass of reactants that are consumed during the progress of the equation

The law of conservation of mass implies that the number of atoms of each kind must be the same on both sides of

a chemical equation An equation satisfying this criterion is known as a balanced chemical equation

A balanced chemical equation provides quantitative information regarding the consumption of reactants and creation of products

The numbers which appear before the chemical symbols and which balance the equation (with the understanding

proportional to the number of molecules or the amounts of the constituents that change during the reaction For

a general case

where nA,nB,nC and nD

The progress of a reaction is described in terms of a physical quantity known as (symbol, x,

pronounced xi)’ It is expressed as

-DnA = - n = n = n =

A

B B C C D D

where Dn represents the change in the amount of the substance The negative and positive signs in the above expressions

are due to the fact that DnA and DnB are negative (i.e their amounts decrease with the progress of reaction) and DnC

and DnD are postive (i.e their amounts increase with the progress of reaction) The unit of extent of reaction (x) is mol.

Decrease in the amount of A is – DnA = nA mol Increase in the amount of C is DnC = nC molDecrease in the amount of B is – DnB = nB mol Increase in the amount of D is DnD = nD mol

In other words, we say that nA mol of A on reacting with nB mol of B gives nC mol of C and nD mol of D

For example, for the reaction Pb(NO3)2 + 2KI Æ PbI2 + 2KNO3, we have

1 mol of Pb(NO3)2 on reacting with 2 mol of KI gives 1 mol of PbI2 and 2 mol of KNO3

In terms of masses consumed/produced, in the reaction

Molar mass 331.2 g mol–1 166 g mol–1 461.0 g mol–1 101.1 g mol–1

we have 331.2 g of Pb(NO3)2 on reacting with 2 ¥ 166 g of KI gives 461.0 g of PbI2 and 2 ¥ 101.1 g of KNO3

In terms of molecules, we have

1 molecule of Pb(NO3)2 on reacting 2 molecules of KI gives 1 molecule of PbI2 and 2 molecules of KNO3

In terms of molecular masses, we have

331.2 u of Pb(NO3)2 on reacting with 2 ¥ 166 u of KI gives 461.0 u of PbI2 and 2 ¥ 101.1 u of KNO3

A given reaction may be initiated with any amounts of reactants, but the consumption of reactants and production of products will be governed by the relation similar to that given by Eq (2), that is, the relative amounts

appeared in the balanced chemical equation.

As the reaction proceeds, the amounts of reactants continue to decrease The reaction continues to proceed till the amount of one of the reactants is exhausted This reactant is known as

To determine the limiting reagent, we proceed as follows

Determine the maximum value of extent of reaction with the initial amount of each of the reactants

xmax = Initial amount of the reactantCorresponding stoichiomeetric coefficient

The amounts of products formed will be governed by this limited reagent For example,

Let the reaction 2NaOH + H2SO4 Æ Na2SO4 + 2H2O be started with 100 g each of NaOH and H2SO4 We will have

mol

2 = mol 1 02. mol 1 02

1 = molSince, the extent of reaction is smaller for H2SO4, this reagent will act as limiting reagent The amounts of products,

Na2SO4 and H2O, will be decided by this reagent as its amount is exhausted earlier Thus, we will have

Amount of NaOH unreacted = (n0 – nxmax) = (2.50 – 2 ¥ 1.02)mol = 0.46 molAmount of Na2SO4 = nxmax = (1) (1.02 mol) = 1.02 mol

Amount of H2O = nxmax = (2) (1.02 mol) = 2.04 mol

MULTIPLE CHOICE QUESTIONS ON SECTION 4

Identify the correct choice in the following questions

(a) Lavoisier (b) Proust (c) Dalton (b) Gay Lussac

2 For the chemical reaction nAA + nBB ÆnCC + nDD, the extent of reaction is given by the expression

3 For the reaction 2A(g) + 3B(g) Æ 4C(g) + 5D(g)

n0 1.5 mol 2.0 mol

the amounts of A, B, C and D when the reaction has proceeded to the extent = 0.15 mol, respectively, are

(a) 1.0 mol, 1.25 mol, 1.0 mol and 1.25 mol (b) 1.2 mol, 1.55 mol, 0.60 mol and 0.75 mol

(c) 0.5 mol, 0.5 mol, 2.0 mol and 2.5 mol (d) 0.9 mol, 1.1 mol, 1.2 mol and 1.5 mol

4 Consider the reaction N2(g) + 3H2(g) Æ 2NH3(g)

n0 1.5 mol 1.5 mol

With the progress of reaction it is found that there is a formation of 0.8 mol of NH3 At this stage, the extent of reaction is

(a) 0.2 mol (b) 0.3 mol (c) 0.4 mol (d) 0.5 mol

5 For the reaction N2(g) + 3H2(g) Æ 2NH3(g)

m0 1400 g 250 g

which of the following statement is correct?

(a) N2(g) acts as a limiting reagent

(b) H2(g) acts as a limiting reagent

(c) The extent of reaction when the reaction is over is 41.67 mol

(d) At the end of reaction, the mass of NH3 formed is 1416.78 g

6 In the reaction N2H4 + 3O2Æ 2NO2 + 2H2O, the mass of O2 required to combine with 745 g of N2H4 will be(a) 2120 g (b) 2235 g (c) 2436 g (d) 2510 g

7 A mixture of 50.0 g of S and 100.0 g Cl2 reacts of form S2Cl2 The mass of S2Cl2 formed will be

(a) 150.0 g (b) 105.5 g (c) 121.0 g (d) 135.1 g

8 In the reaction Fe2O3(s) + 3C(s) Æ 2 Fe(s) + 3CO(g), 453 kg of iron was obtained from 752 kg of a sample

of Fe2O3 The perentage of Fe2O3 –1)

9 In the reaction 2NH4Cl(s) + Ca(OH)2(s) Æ CaCl2(s) + 2NH3(g) + 2H2O(g), the mass of NH3 formed by heating

a mixture containing 10.0 g each of NH4Cl and Ca(OH)2 –1)

(c) Oxygen is completely used in the reaction

(d) The extent of reaction at the completion of reaction is 0.025 mol

ANSWERS

7 (b) 8 (c) 9 (c) 10 (b)

HINTS AND SOLUTIONS

1 Law of conservation of mass was established by Lavoisier

2 The extent of reaction is x

= -DnA =Dn

A

C C

n = x fi –DnB = nBx = 3 ¥ 0.15 mol = 0.45 mol nB = n0,B + DnB = 2.0 mol – 0.45 mol = 1.55 mol

nC = DnC = nCx = 4 ¥ 0.15 mol = 0.60 mol and nD = DnD = nDx = 5 ¥ 0.15 mol = 0.75 mol

.N

.H

The extent of reaction when the reaction is over will be xmax(H2), i.e 41.67 mol

DnNH3 = nNH3xmax = 2 ¥41.67 mol = 83.34 mol

1 40841

=n = = 1.4084 mol

S acts as limiting reagent Hence

Dn(S2Cl2) = n(S2Cl2) xmax(S) = (1)(0.7812 mol) = 0.7812 mol

Dm(S2Cl2) = Dn(S2Cl2) M(S2Cl2) = (0.7812 mol)(135 g mol–1) = 105.5 g

8 We have Fe2O3(s) + 3C(s) Æ 2Fe(s) + 3CO(g)

Molar mass 160 g mol–1 56 g mol–1

2 ¥ 56 g of Fe(s) will be obtained from 160 g of Fe2O3 Mass of Fe2O3 from which 453 kg of Fe is obtained will

752 100 86

%kg

Since xmax(NH4Cl) < xmax (Ca(OH)2), the limiting reagent is NH4Cl

DnNH3 = nNH3xmax(NH4Cl) = 2 ¥ 0.0935 mol = 0.1870 mol

= - = 0.025 mol Since xmax(O2) < xmax(Mg), the limiting reagent is oxygen

DnMgO = nMgOxmax(O2) = 2(0.025 mol) = 0.05 mol