std triumph chemistry - target publications › ... › std-11-chemistry-mcqs.pdf · triumph...

© Target Publications Pvt. Ltd. No part of this book may be reproduced or transmitted in any form or by any means, C.D. ROM/Audio Video Cassettes or electronic, mechanical

including photocopying; recording or by any information storage and retrieval system without permission in writing from the Publisher.

STD. XI Sci. Triumph Chemistry

Based on Maharashtra Board Syllabus

Printed at: Repro India Ltd., Mumbai

Useful for all Agricultural, Medical, Pharmacy and Engineering Entrance Examinations

held across India.

Solutions/hints to Evaluation Test available in downloadable PDF format at

www.targetpublications.org/tp10145

10145_10960_JUP

P.O. No. 30310

Salient Features

• Exhaustive subtopic wise coverage of MCQs

• Quick Review and/or Important Formulae provided for all the chapters

• Hints included for relevant questions

• Various competitive exams questions updated till the latest year

• Includes solved MCQs from JEE (Main), AIPMT / NEET P-I and P-II,

KCET 2015 and 2016

• Evaluation Test provided at the end of each chapter

Preface “Std. XI: Sci. Triumph Chemistry” is a complete and thorough guide to prepare students for

competitive level examinations. This book not only assists students with MCQs of Std. XI but also helps them to prepare for JEE, AIPMT / NEET-UG, CET and various other competitive examinations.

The content of this book is based on the Maharashtra State Board Syllabus. Quick Review which summarizes the important concepts of the entire chapter is provided for all the chapters. Formulae that form a vital part of MCQ solving are provided for relevant chapters.

MCQs in each chapter are divided into three sections: Classical Thinking: consists of straight forward questions including knowledge based questions. Critical Thinking: consists of questions that require understanding of the concept and the applications of the

same. Competitive Thinking: consists of questions from various competitive examinations like JEE, AIPMT / NEET-

UG, MH CET, KCET, CPMT, GUJ CET, AP EAMCET (Engineering, Medical), TS EAMCET (Engineering, Medical), Assam CEE, BCECE, WB JEEM, etc.

Hints (i.e., complete solutions broken down to the simplest form possible) have been provided to the relevant MCQs.

An Evaluation Test has been provided at the end of each chapter to assess the level of preparation of the student on a competitive level. In order to understand how chemistry plays an important role in our day to day life, we have made an attempt to illustrate the same in the form of images/visuals in the related chapters. The journey to create a complete book is strewn with triumphs, failures and near misses. If you think we’ve nearly missed something or want to applaud us for our triumphs, we’d love to hear from you. Please write to us at : [email protected]

Best of luck to all the aspirants!

Yours faithfully Authors

Sr. No. Topic Name Page No.1 Some Basic Concepts of Chemistry 1 2 States of Matter (Gases and Liquids) 22 3 Structure of Atom 51 4 Periodic Table 85 5 Redox Reactions 103 6 Chemical Equilibrium 119 7 Surface Chemistry 165 8 Nature of Chemical Bond 186 9 Hydrogen 220 10 s-Block Elements 240 11 p-Block Elements (Group 13 and 14) 262 12 Basic Principles and Techniques in Organic Chemistry 290 13 Alkanes 330 14 Alkenes 351 15 Alkynes 369 16 Aromatic Compounds 387 17 Environmental Chemistry 405

Index

1

Chapter 01: Some Basic Concepts of Chemistry

Subtopics

1.0 Introduction

1.1 Importance and scope of chemistry

1.2 Historical approach to particulate

nature of matter

1.3 Laws of chemical combination

1.4 Dalton’s atomic theory

1.5 Concepts of elements, atoms and

molecules

1.6 Atomic and molecular masses

1.7 Avogadro’s law, Avogadro’s number

and mole concept

1.8 Percentage composition and

empirical and molecular formula

1.9 Chemical reactions and

stoichiometry

Have you ever wondered what is the reference for 1 kilogram???? It has been defined as the mass of the International Prototype of the Kilogram. The prototype is made of platinum-iridium (Pt-Ir) cylinder that is stored in an airtight jar at International Bureau of Weights and Measures in France. Pt-Ir was chosen because its mass remains constant for an extremely long time and it is resistant to the attack of different chemicals. This reference standard is used to calibrate or standardize different measuring devices such as analytical balances.

Platinum alloy as an International Prototype of the Kilogram

Some Basic Concepts of Chemistry 01

2

Std. XI : Triumph Chemistry

Branches of chemistry: Classification of matter:

Quick Review

Chemistry Study of the composition, structure and properties of matter and the reactions by which one form of matter may be converted into another form.

Physical chemistry Deals with the structure of matter, the energy changes and the

theories, laws and principles that explain the transformation of matter from one form to another.

Inorganic chemistry Deals with chemistry of elements other than carbon and of

their compounds.

Organic chemistry Deals with the reactions of the compounds of carbon.

Analytical chemistry Deals with the separation, identification and quantitative

determination of the compositions of different substances.

Biochemistry Deals with chemistry of compounds and processes occurring in

living organisms.

Matter has mass occupies space

Pure substances Fixed composition

Elements Substances that cannot

be decomposed by a simple chemical process into two or more different substances.

eg. He (monoatomic), H2 (diatomic)

Compounds Substances of definite

compositions which can be decomposed into two or more susbtances by a simple chemical process.

eg. H2O, NaCl

Mixtures Variable composition (that can be

separated by simple physical methods)

Homogeneous mixtures Composition is

uniform. All constituents

present in one phase.

eg. air, ethyl alcohol and water

Heterogeneous mixtures Composition is not

uniform. Two or more phases

are present. eg. phenol-water

system, iron filings-sand system

3


Laws of chemical combination:

1. Atomic mass unit (1 amu) =1

12th of a 12C-atom

= 1.66 1027 kg 2. 1 Mole = 6.022 1023 particles

(atoms/molecules/ions/electrons) 3. Number of moles (n)

=Mass of thesubstance

Molar mass of thesubstance

4. Mass of an atom = 23

Atomicmass

6.022 10

5. Mass of a molecule = 23

Molecular mass

6.022 10

6. Number of molecules = n Avogadro number (NA) 7. Atomicity = number of atoms in a molecule 8. Total number of atoms in molecule = n NA Atomicity 9. Volume occupied by one mole of a gas at STP

= 22.4 L = 22.4 dm3 10. Molecular formula = r Empirical formula

11. r = Molecular mass

Empiricalmass

12. Average atomic mass

= Sumof (Isotopicmass its%abundance)

100

13. Avogadro’s law, V n (At constant T and P)

Laws of chemical combination

Law of conservation of mass First stated by Lomonosove (1765) and then by Antoine Laviosier (1783). Statement: The mass is neither created nor destroyed during chemical

combination of matter.

Law of definite composition/Law of definite proportion Stated by Joseph Proust. Statement: Any pure compound always contains the same elements in a

definite proportion by weight irrespective of its source or method ofpreparation.

Law of multiple proportion Stated by John Dalton. Statement: If two elements chemically combine with each other forming

two or more compounds with different compositions by mass then theratios of masses of two interacting elements in the two compounds aresmall whole numbers.

Gay Lussac’s law of combining volumes of gases Stated by Gay Lussac. Statement: When gases react together to produce gaseous products, the

volumes of reactants and products bear a simple whole number ratio witheach other, provided volumes are measured at same temperature andpressure.

Avogadro’s law Stated by Avogadro. Statement: Equal volumes of all gases under identical conditions of

temperature and pressure contain equal number of molecules.

Formulae

4


1.0 Introduction 1. _______ chemistry deals with the chemistry of

elements other than carbon and of their compounds.

(A) Organic (B) Physical (C) Inorganic (D) Bio 2. The branch of chemistry which deals with the

separation, identification and quantitative determination of the composition of different substances is called _______ chemistry.

(A) organic (B) inorganic (C) analytical (D) bio 1.1 Importance and scope of chemistry 3. _______ pigment acts as a photosensitizer in

plants. (A) Xanthophyll (B) Chlorophyll (C) Carotene (D) ATP 4. _______ CANNOT be carried out in a lab. (A) Photosynthesis (B) Reduction (C) Oxidation (D) Hydration 5. Solar energy can be converted into electrical

energy using _______ cell. (A) Daniel (B) lithium ion (C) photovoltaic (D) nickel cadmium 6. In computers, _______ chips are used as

microprocessors. (A) carbon (B) phosphorus (C) titanium (D) silicon 1.2 Historical approach to particulate

nature of matter 7. Which one of the following is NOT a mixture? (A) Iodized table salt (B) Gasoline (C) Liquefied Petroleum Gas (L.P.G.) (D) Distilled water 8. If two or more phases are present in a mixture

then it is called a _______ mixture. (A) heterogeneous (B) homogeneous (C) homologous (D) heterologous 9. Phenolwater system is a/an _______. (A) element (B) compound (C) homogeneous system (D) heterogeneous system

10. The phlogiston theory was suggested for _______ reaction.

(A) neutralisation (B) oxidation (C) reduction (D) combustion 11. Substances which CANNOT be decomposed

into two or more different substances by chemical process are called _______.

(A) alloys (B) molecules (C) elements (D) compounds 12. The arbitrarily decided and universally

accepted standards are called _______. (A) fundamentals (B) units (C) measures (D) symbols 13. There are _______ fundamental SI units. (A) 3 (B) 5 (C) 6 (D) 7 14. SI unit of temperature is _______. (A) K (B) C (C) F (D) D 15. SI unit of velocity is _______. (A) km s1 (B) km hr1

(C) m s2 (D) m s1

1.3 Laws of chemical combination 16. The sum of the masses of reactants and

products is equal in any physical or chemical reaction. This is in accordance with law of _______.

(A) multiple proportion (B) definite composition (C) conservation of mass (D) reciprocal proportion 17. If the law of conservation of mass was to hold

true, then 20.8 g of BaCl2, on reaction with 9.8 g of H2SO4 will produce 7.3 g of HCl and _______ of BaSO4.

(A) 11.65 g (B) 23.3 g (C) 25.5 g (D) 30.6 g 18. Pure water can be obtained from various

sources, but it always contains hydrogen and oxygen, combined in a ratio of 1:8 by weight. This is an example of _______.

(A) law of conservation of mass (B) Avogadro’s law (C) law of definite composition (D) Gay Lussac’s law

Classical Thinking

5


19. Two containers of the same size are filled separately with H2 gas and CO2 gas. Both the containers under the same T and P will contain the same _______.

(A) number of atoms (B) weight of gas (C) number of molecules (D) number of electrons 20. In SO2 and SO3, the ratio of the masses of

oxygen which combine with a fixed mass of sulphur is 2:3. This is an example of the law of _______.

(A) constant proportion (B) multiple proportion (C) reciprocal proportion (D) Gay Lussac 21. Which of the following reactions has the ratio

of volumes of reacting gases and the product as 1:2:2?

(A) 2CO(g) + O2(g) 2CO2(g) (B) O2(g) + 2H2(g) 2H2O(g) (C) H2(g) + F2(g) 2HF(g) (D) N2(g) + 3H2(g) 2NH3(g) 22. The volume of oxygen required for complete

combustion of 0.25 cm3 of CH4 at S.T.P is _______ cm3.

(A) 0.25 (B) 0.5 (C) 0.75 (D) 1 1.4 Dalton’s atomic theory 23. Greek philosopher _______ had suggested that

matter is composed of extremely small a-tomio. (A) Dalton (B) Aristotle (C) Ptolemy (D) Democritus 24. Dalton assumed that _______ are the smallest

particles of compound. (A) atoms (B) molecules (C) ions (D) elements 1.5 Concepts of elements, atoms and

molecules 25. Atoms have a mass of the order _______. (A) 1026 kg (B) 1015 kg (C) 1026 g (D) 1015 g 26. Atoms have a radius of the order _______. (A) 1026 m (B) 1015 m (C) 1015 mm (D) 1015 m

27. A/an _______ is an aggregate of two or more atoms in definite composition which are held together by chemical bonds.

(A) ion (B) molecule (C) compound (D) mixture 1.6 Atomic and molecular masses 28. Every atom of an element consists of fixed

number of _______. (A) protons (B) neutrons (C) electrons (D) all of these

29. The unit of atomic mass amu is replaced by

_______. (A) u (B) mol (C) g (D) kg 30. Mole is the SI unit of _______. (A) volume (B) pressure (C) amount of substance (D) density 31. 1 amu is equal to _______.

(A) 1

12 of C – 12 (B)

1

14 of O – 16

(C) 1 g of H2 (D) 1.66 10–23 kg 32. ________ is the sum of the atomic mass of all

the atoms as given in the molecular formula of the substance.

(A) Molecular mass (B) Atomic weight (C) Percentage weight (D) Percentage volume 1.7 Avogadro’s law, Avogadro’s number

and mole concept 33. NA = _________ atoms mol1. (A) 6.021 1021 (B) 6.024 1024

(C) 6.051 1015 (D) 6.022 1023

34. One _______ is the collection of 6.022 1023

atoms /molecules/ions. (A) kg (B) g (C) mole (D) cm 35. Avogadro’s number is ________. (A) number of atoms in one gram of element. (B) number of millilitres which one mole of

a gaseous substance occupies at N.T.P. (C) number of molecules present in one

gram molar mass of a substance. (D) number of elements in one gram of

compounds.

6


36. Which of the following law states that equal volume of all gases contain equal number of molecules?

(A) Boyle’s law (B) Charles’ law (C) Avogadro’s law (D) Gay Lussac’s law 37. According to Avogadro’s law, ________.

(A) V 1

P (B) V T

(C) V n (D) all of these 38. Volume occupied by 1 g molecular weight of

any gas is called _______. (A) gram molecular volume (B) gram atomic volume (C) gram molecular weight (D) gram atomic weight 39. Avogadro’s law distinguishes between

_______. (A) cations and anions (B) atoms and molecules (C) atoms and ions (D) molecules and ions 40. The number of atoms present in a molecule of

a substance is called its ________. (A) atomicity (B) volume (C) density (D) mass 41. How many molecules are present in one gram

of hydrogen gas? (A) 6.02 1023 (B) 3.01 1023 (C) 2.5 1023 (D) 1.5 1023

42. One mole of CO2 contains _______. (A) 6.022 1023 atoms of C (B) 6.022 1023 atoms of O (C) 18.1 1023 molecules of CO2 (D) 3 g atoms of CO2 43. One mole of H2O corresponds to _______. (A) 22.4 litres at 1 atm and 25 C (B) 6.02 1023 atoms of hydrogen and 6.02 1023 atoms of oxygen (C) 18 g of H2O (D) 1 g of H2O 44. The gram molecule of benzene is equal to

_______ g C6H6. (A) 70 (B) 72 (C) 10 (D) 78

45. 1 atom of an element weighs 1.792 10–22 g. The atomic mass of the element is _______.

(A) 1.192 (B) 17.92 (C) 64 (D) 108 46. What is the mass of 0.5 mole of ozone molecule? (A) 8 g (B) 16 g (C) 24 g (D) 48 g 47. The number of molecules in 16 g of oxygen

gas is _______. (A) 6.022 1023 (B) 3.011 1023

(C) 3.011 1022 (D) 1.5 1023 48. Which of the following weighs the least? (A) 2.0 gram mole of CO2

(B) 0.1 mole of sucrose (C12H22O11) (C) 1 gram atom of calcium (D) 1.5 mole of water 49. Which one of the following pairs of gases

contains the same number of molecules? (A) 16 g of O2 and 14 g of N2 (B) 8 g of O2 and 22 g of CO2 (C) 28 g of N2 and 22 g of CO2 (D) 32 g of O2 and 32 g of N2 50. One mole of oxygen gas weighs _______. (A) 1 g (B) 8 g (C) 32 g (D) 6.023 1023 g 51. Under similar conditions, same mass of

oxygen and nitrogen is taken. The ratio of their volumes will be _______.

(A) 7 : 8 (B) 3 : 5 (C) 6 : 5 (D) 9 : 2 1.8 Percentage composition and empirical

and molecular formula 52. Chemical formula CANNOT be determined

by using _______. (A) Raman spectroscopy (B) nuclear magnetic resonance (C) Titration (D) ultraviolet spectroscopy 53. The mass percentage of each constituent

element present in 100 g of a compound is called its _______.

(A) molecular composition (B) atomic composition (C) percentage composition (D) mass composition

7


54. _______ of a compound is the chemical formula indicating the relative number of atoms in the simplest ratio.

(A) Empirical formula (B) Molecular formula (C) Empirical mass (D) Molecular mass 55. The percentage composition of carbon in urea,

[CO(NH2)2] is _______. (A) 20% (B) 40% (C) 50% (D) 80% 56. What is the % of H2O in Fe(CNS)33H2O? (A) 19 (B) 25 (C) 30 (D) 45 57. The percentage of oxygen in NaOH is

_______. (A) 8 (B) 10 (C) 40 (D) 60 58. A compound made of two elements A and B

are found to contain 25% A (Atomic mass 12.5) and 75% B (Atomic mass 37.5). The simplest formula of the compound is _______.

(A) AB (B) AB2 (C) AB3 (D) A3B 59. _______ indicates the actual number of

constituent atoms in a molecule. (A) Empirical formula (B) Molecular formula (C) Empirical mass (D) Molecular mass 60. Which of the following has same molecular

formula and empirical formula? (A) CO2 (B) C6H12O6 (C) C2H2 (D) C2H2O4 61. Empirical formula of glucose is _______. (A) C6H12O6 (B) C6H11O6 (C) CHO (D) CH2O 1.9 Chemical reactions and stoichiometry 62. The starting material which takes part in

chemical reaction is called _______. (A) product (B) reactant (C) catalyst (D) starter 63. ________ is the quantitative relationship

between the reactants and products in a balanced chemical equation.

(A) Stoichiometry (B) Complexometry (C) Chemistry (D) Reactions

64. _______ reactant is the reactant that reacts completely but limits further progress of the reaction.

(A) Oxidizing (B) Reducing (C) Limiting (D) Excess 65. _______ reactant is the reactant which is taken

in excess than the limiting reactant. (A) Oxidizing (B) Reducing (C) Limiting (D) Excess 66. The _______ coefficients are the coefficients

of reactants and products in the balanced chemical reaction.

(A) balanced (B) chemical (C) stoichiometric (D) molar Miscellaneous 67. Which of the following relations for expressing

volume of a sample is NOT correct? (A) 1L = 103 ml (B) 1dm3 = 1L (C) 1L = 103m3 (D) 1L = 103 cm3 68. Which out of the following is NOT a

homogeneous mixture? (A) Solution of glucose in water. (B) Solution of salt in water. (C) Mixture of glucose solution and salt

solution. (D) Mixture of oil and water. 69. The molecular mass of hydrogen peroxide is

34. What is the unit of molecular mass? (A) g (B) mol (C) g mol1 (D) mol g1

Synthetic fabrics like nylon, terylene, etc., are man-made fabrics. They are very elastic and dry quickly after washing. They are mainly popular because of their crease free nature i.e., they don’t need to be ironed !!!

Synthetic fabrics

8


1.1 Importance and scope of chemistry 1. Azidothymidine drug is used for treating

_______ patients. (A) diabetes (B) AIDS (C) jaundice (D) tuberculosis 2. Which of the following indicates CORRECT

reaction for photosynthesis? (A) C + O2 + H2O Sunlight Products (B) O2 + H2O Sunlight Products (C) CO2 + H2O Sunlight Products (D) CO + H2O Sunlight Products 1.2 Historical approach to particulate

nature of matter 3. Which of the following statements is

INCORRECT? (A) Constituent substances in a mixture

retain their separate identities. (B) Composition of a mixture can be varied

to any extent. (C) Mixture of liquids are example of

homogeneous mixtures. (D) Mixtures can be separated into pure

components by simple physical methods.

4. The revised metric system in which units are

expressed is _______. (A) CGS (B) MKS (C) FPS (D) SI 5. Electrochemical equivalence has unit

_______. (A) kg m s1 (B) kg m2 s1 (C) kg C1 (D) kg m1 s2

6. Magnitude of ‘pico’ is _______. (A) 1012 (B) 1015 (C) 1012 (D) 109

1.3 Laws of chemical combination 7. Hydrogen and oxygen combine to form H2O2

and H2O containing 5.93% and 11.29% of hydrogen respectively. The data illustrates _______.

(A) law of conservation of mass (B) law of definite composition (C) law of reciprocal proportion (D) law of multiple proportion

8. Two elements, A and B, combine to form a compound in which ‘a’ g of A combines with ‘b1’ and ‘b2’ g of B respectively. According to law of multiple proportion _______.

(A) b1 = b2 (B) b1 and b2 bear a simple whole number ratio (C) a is always equal to b1 (D) no relation exists between b1 and b2 9. After a chemical reaction, the total mass of

reactants and products _______. (A) always increases (B) always decreases (C) does not change (D) always increases or decreases 10. Two samples of lead oxide were separately

reduced to metallic lead by heating in a current of hydrogen. The weight of lead from one oxide was half the weight of lead obtained from the other oxide. The data illustrates _______.

(A) law of reciprocal proportions (B) law of constant proportions (C) law of multiple proportions (D) law of equivalent proportions 11. The law of multiple proportions is illustrated

by the compounds _______. (A) carbon monoxide and carbon dioxide (B) potassium bromide and potassium chloride (C) ordinary water and heavy water (D2O) (D) calcium hydroxide and barium

hydroxide 12. How many litres of ammonia will be formed

when 2 L of N2 and 2 L of H2 are allowed to react?

(A) 0.665 (B) 1.0 (C) 1.33 (D) 4.00 1.4 Dalton’s atomic theory 13. Which of the following statements is TRUE

according to Dalton’s atomic theory? (A) A chemical reaction involves only the

separation, combination or rearrangement of integer number of atoms.

(B) Law of conservation of mass can be explained by assuming that total number of atoms in the reactants and products remain same.

(C) During chemical reactions atoms are neither created nor destroyed.

(D) All of these.

Critical Thinking

9


1.5 Concepts of elements, atoms and molecules

14. Which of the following statements is

INCORRECT? (A) Atoms may or may not have free

existence. (B) A molecule may contain atoms of same

elements or different elements. (C) A molecule can be divided into the its

constituent atoms by simple methods. (D) The properties of constituent atoms and

the compounds formed from them are completely different.

1.6 Atomic and molecular masses 15. Isotopes are the atoms of the same element

having _______. (A) different number of protons (B) different number of electrons (C) different number of neutrons (D) same number of neutrons 16. Which of the following indicates natural

abundance of neon – 20 isotope? (A) 90.92 u (B) 90.92 % (C) 90.92 gm mol–1 (D) 90.92 0.012 kg of 12C 17. In chemical scale, the relative mass of the

isotopic mixture of oxygen atoms (16O, 17O, 18O) is assumed to be equal to _______.

(A) 15.002 (B) 16.00 (C) 17.00 (D) 18.00 1.7 Avogadro’s law, Avogadro’s number

and mole concept 18. The number of moles of sodium oxide in

620 g is _______. (A) 1 mol (B) 10 moles (C) 18 moles (D) 100 moles 19. The number of water molecules in 1 litre of

water is _______. (A) 18 (B) 18 1000 (C) NA (D) 55.55 NA 20. How many atoms are contained in one mole of

sucrose (C12H22O11)? (A) 45 6.023 1023 atoms/mole (B) 5 6.623 1023 atoms/mole (C) 5 6.023 1023 atoms/mole (D) None of these

21. 1 mol of CH4 contains _______. (A) 6.02 1023 atoms of H (B) 4 g atom of H (C) 1.81 1023 molecules of CH4 (D) 3.0 g of carbon 22. The mass of 1 atom of hydrogen is _______. (A) 1 g (B) 0.5 g (C) 1.6 1024 g (D) 3.2 1024 g 23. 1 gram atom of nitrogen represents _______. (A) 6.02 1023 N2 molecules (B) 22.4 L of N2 at N.T.P. (C) 11.2 L of N2 at N.T.P. (D) 28 g of nitrogen 24. The number of molecules in 22.4 dm3 of

nitrogen gas at STP is _______. (A) 6.023 1020 (B) 6.023 1023 (C) 22.4 1020 (D) 22.4 1023 25. 27 g of Al (Atomic mass = 27) will combine

with _______ of O2 to form aluminium oxide. (A) 24 g (B) 8 g (C) 40 g (D) 10 g 26. How many moles of electrons weigh one kilogram?

(A) 6.022 1023 (B) 1

9.108 1031

(C) 6.022

9.108 1054 (D)

1

9.108 6.022 108

27. Which of the following has maximum number

of atoms? (A) 18 g of H2O (B) 16 g of O2

(C) 4.4 g of CO2 (D) 16 g of CH4 28. The number of sulphur atoms present in

0.2 moles of S8 molecules is _______. (A) 4.82 1023 (B) 9.63 1022

(C) 9.63 1023 (D) 1.20 1023

29. What will be the volume of CO2 at NTP

obtained on heating 10 grams of (90% pure) limestone?

(A) 22.4 litre (B) 2.016 litre (C) 2.24 litre (D) 20.16 litre 1.8 Percentage composition and empirical

and molecular formula 30. If two compounds have the same empirical

formula but different molecular formula, they must have _______.

(A) different percentage composition (B) different molecular weights (C) same viscosity (D) same vapour density

10


31. Which pair of species have same percentage of carbon?

(A) CH3COOH and C6H12O6 (B) CH3COOH and C2H5OH (C) HCOOCH3 and C12H22O11 (D) C6H12O6 and C12H22O11 32. Percentage of nitrogen in urea is about

_______. (A) 46 % (B) 85 % (C) 18 % (D) 28 % 33. Two elements X (Atomic mass 75) and Y

(Atomic mass 12) combine to give a compound having 75.8% X. The empirical formula of the compound is _______.

(A) XY (B) XY2

(C) X2Y2 (D) X2Y3

34. The molecular mass of an organic compound is 78. Its empirical formula is CH. The molecular formula is _______.

(A) C2H4 (B) C2H2 (C) C6H6 (D) C4H4 35. The empirical formula of a compound is

CH2O. 0.0835 moles of the compound contains 1.0 g of hydrogen. Molecular formula of the compound is ________.

(A) C6H12O6 (B) C5H10O5 (C) C4H8O4 (D) C3H6O3 1.9 Chemical reactions and stoichiometry 36. Calculate the number of moles of methane

required to produce 33 g of carbon dioxide gas on its complete combustion.

(A) 0.15 moles (B) 0.50 moles (C) 0.75 moles (D) 0.95 moles 37. The volume of ammonia obtained by the

combination of 10 mL of N2 and 30 mL H2 is _______.

(A) 20 mL (B) 40 mL (C) 30 mL (D) 10 mL 38. What mass of CaO will be obtained by heating

3 mole of CaCO3? [Atomic mass of Ca = 40] (A) 150 g (B) 168 g (C) 16.8 g (D) 15 g 39. 3 g of H2 reacts with 29 g of O2 to yield water.

Which is the limiting reactant? (A) H2 (B) O2 (C) H2O (D) none of there

Miscellaneous 40. Which of the following is a compound? (A) Diamond (B) Charcoal (C) Baking powder (D) 22 Carat Gold 41. The number of atoms in 6 amu of He is

_______. (A) 18 (B) 18 6.022 1023 (C) 54 (D) 54 6.023 1023

42. Two elements, X (Atomic mass 16) and Y

(Atomic mass 14) combine to form compounds A, B and C. The ratio of different masses of Y which combine with fixed mass of X in A, B and C is 1:3:5. If 32 parts by mass of X combine with 84 parts by mass of Y in B, then in C, 16 parts by mass of X will combine with _______.

(A) 14 parts by mass of Y (B) 42 parts by mass of Y (C) 70 parts by mass of Y (D) 82 parts by mass of Y 43. Which of the following is the value of amu?

(A) 1.57 1024 kg (B) 1.66 1024 kg

(C) 1.99 1023 kg (D) 1.66 1027 kg

The experiment for the search of the Higgs boson, known as ‘God particle’, using Large Hadron Collider (LHC) was an international effort involving thousands of people, with physicists and engineers. The LHC is the world's largest and most powerful particle collider that consists of a 27-kilometre ring of superconducting magnets with a number of accelerating structures to boost the energy of the particles along the way.

The Search for God Particle!!!

11


1.3 Laws of chemical combination 1. Which of the following is the best example of

law of conservation of mass?

[NCERT 1975]

(A) 12 g of carbon combines with 32 g of oxygen to form 44 g of CO2.

(B) When 12 g of carbon is heated in a vacuum there is no change in mass.

(C) A sample of air increases in volume when heated at constant pressure but its mass remains unaltered.

(D) The weight of a piece of platinum is the same before and after heating in air.

2. A sample of pure carbon dioxide, irrespective

of its source contains 27.27% carbon and 72.73% oxygen. The data supports _______.

[AIIMS 1992]

(A) law of definite composition

(B) law of conservation of mass

(C) law of reciprocal proportions

(D) law of multiple proportions 3. The law of multiple proportions was proposed

by _______. [IIT 1992]

(A) Lavoisier (B) Dalton

(C) Proust (D) Gay-Lussac 1.6 Atomic and molecular masses 4. Which property of an element is always a whole

number?

[CPMT 1986; MP PMT 1986]

(A) Atomic weight

(B) Atomic volume

(C) Atomic number

(D) All of these 5. The weight of a molecule of the compound

C60H122 is _______. [AIIMS 2000]

(A) 1.4 1021 g

(B) 1.09 1021 g

(C) 5.025 1023 g

(D) 16.023 1023 g

6. Boron has two stable isotopes, 10B (19%) and 11B (81%). The atomic mass that should appear for boron in the periodic table is _______.

[CBSE PMT 1990]

(A) 10.0 (B) 10.2

(C) 10.8 (D) 11.2 7. An element, X has the following isotopic

composition.

200X : 90% ; 199X : 8.0% ; 202X : 2.0%

The weighed average atomic mass of the naturally occurring element X is close to _______.

[CBSE PMT 2007]

(A) 200 amu (B) 210 amu

(C) 202 amu (D) 199 amu 1.7 Avogadro’s law, Avogadro’s number

and mole concept 8. The number of atoms in 4.25 g of NH3 is

approximately _______.

[CBSE PMT 1999; MH CET 2003]

(A) 1 1023 (B) 2 1023

(C) 4 1023 (D) 6 1023

9. The weight of 1 1022 molecules of

CuSO4.5H2O is _______. [IIT 1991]

(A) 41.59 g (B) 415.9 g

(C) 4.159 g (D) none of these 10. What amount of dioxygen (in gram) contains

1.8 1022 molecules?

[KCET 2015]

(A) 0.0960 (B) 0.960

(C) 9.60 (D) 96.0 11. The number of oxygen atoms in 4.4 g of CO2

is approximately _______.

[CBSE PMT 1990; KCET 2016]

(A) 1.2 1023 (B) 6 1022

(C) 6 1023 (D) 12 1023

12. The volume occupied by 4.4 g of CO2 at STP

is _______.

[AFMC 1997, 2004]

(A) 0.1 L (B) 0.224 L

(C) 2.24 L (D) 22.4 L

Competitive Thinking

12


13. The number of atoms in 0.1 mole of a triatomic

gas is _______. (NA = 6.02 1023 mol1)

[CBSE PMT 2010]

(A) 1.800 1022 (B) 6.026 1022

(C) 1.806 1023 (D) 3.600 1023

14. The system that contains the maximum

numbers of atoms is _______.

[WB JEEM 2014] (A) 4.25 g of NH3

(B) 8 g of O2 (C) 2 g of H2 (D) 4 g of He 15. Assuming fully decomposed, the volume of

CO2 released at STP on heating 9.85 g of BaCO3 (Atomic mass of Ba = 137) will be _______. [CBSE PMT 2000]

(A) 0.84 L (B) 2.24 L

(C) 4.06 L (D) 1.12 L 16. The numbers of moles of BaCO3 which

contain 1.5 moles of oxygen atoms is _______. [EAMCET 1991]

(A) 0.5 (B) 1

(C) 3 (D) 6.02 1023

17. The number of moles of oxygen in 1 L of air

containing 21% oxygen by volume in standard conditions is _______.

[CBSE PMT 1995] (A) 0.0093 mol (B) 0.186 mol

(C) 0.21 mol (D) 2.10 mol 18. 1 mL of water contains 20 drops. Then

number of molecules in a drop of water is _______ molecules.

[AFMC 2010]

(A) 6.023 1023 (B) 1.376 1026

(C) 1.344 1018 (D) 4.346 1020

19. What volume of hydrogen gas, at 273 K and

1 atm pressure will be consumed in obtaining 21.6 g of elemental boron (atomic mass = 10.8) from the reduction of boron trichloride by hydrogen? [AIEEE 2003]

(A) 22.4 L (B) 89.6 L

(C) 67.2 L (D) 44.8 L

20. How many moles of lead (II) chloride will be formed from a reaction between 6.5 g of PbO and 3.2 g of HCl?

[CBSE PMT 2008]

(A) 0.011 (B) 0.029

(C) 0.333 (D) 0.044 1.8 Percentage composition and empirical

and molecular formula 21. The empirical formula of an acid is CH2O2,

the probable molecular formula of acid may be _______.

[AFMC 2000]

(A) CH2O (B) CH2O2

(C) C2H4O2 (D) C3H6O4

22. A compound (80 g) on analysis gave C = 24 g, H = 4 g, O = 32 g. Its empirical formula is _______. [CPMT 1981]

(A) C2H2O2 (B) C2H2O

(C) CH2O2 (D) CH2O 23. An organic compound contains carbon hydrogen

and oxygen. Its elemental analysis gave C, 38.71% and H, 9.67 % and O, 51.62%. the empirical formula of the compound would be _______. [CBSE PMT 2008]

(A) CHO (B) CH4O

(C) CH3O (D) CH2O 1.9 Chemical reactions and stoichiometry 24. The percentage of Se in peroxidase

anhydrous enzyme is 0.5% by weight (atomic mass = 78.4). Then minimum molecular mass of peroxidase anhydrous enzyme is _______.

[CBSE PMT 2001]

(A) 1.568 104 (B) 1.568 103

(C) 15.68 (D) 3.136 104

25. During electrolysis of water the volume of O2

liberated is 2.24 dm3. The volume of hydrogen liberated, under same conditions will be _______.

[AIIMS 2008]

(A) 0.56 dm3 (B) 1.12 dm3

(C) 2.24 dm3 (D) 4.48 dm3

13


26. What volume of oxygen gas (O2) measured at

0 C and 1 atm, is needed to burn completely 1 L of propane gas (C3H8) measured under same conditions? [CBSE PMT 2008]

(A) 5 L (B) 10 L (C) 7 L (D) 6 L 27. In the reaction, 2Al(s) + 6HCl(aq) 3

(aq)2Al + (aq)6Cl + 3H2(g)

[AIEEE 2007] (A) 6 L HCl(aq) is consumed for every 3 L

H2(g) produced. (B) 33.6 L H2(g) is produced regardless of

temperature and pressure for every mole Al that reacts.

(C) 67.2 L H2(g) at STP is produced for every mole Al that reacts.

(D) 11.2 L H2(g) at STP is produced for every mole HCl(aq) consumed.

28. How many moles of magnesium phosphate,

Mg3(PO4)2 will contain 0.25 mole of oxygen atoms? [AIEEE 2006]

(A) 0.02 (B) 3.125 102

(C) 1.25 102 (D) 2.5 102

29. 10 g hydrogen is reacted with 64 g oxygen.

The amount of water formed will be (in moles) _______. [BCECE 2015]

(A) 3 (B) 4 (C) 1 (D) 2 Miscellaneous 30. The ratio of masses of oxygen and nitrogen in

a particular gaseous mixture is 1 : 4. The ratio of number of their molecule is _______.

[AFMC 2006, JEE(MAIN) 2014] (A) 1 : 4 (B) 7 : 32 (C) 1 : 8 (D) 3 : 16 31. 20.0 g of a magnesium carbonate sample

decomposes on heating to give carbon dioxide and 8.0 g magnesium oxide. What will be the percentage purity of magnesium carbonate in the sample?

[AIPMT RE-TEST 2015] (A) 60 (B) 84 (C) 75 (D) 96 (At. Wt. : Mg = 24)

32. 1.0 g of magnesium is burnt with 0.56 g O2 in a closed vessel. Which reactant is left in excess and how much?

(At wt. Mg = 24; O = 16) [AIPMT 2014]

(A) Mg, 0.16 g

(B) O2, 0.16 g

(C) Mg, 0.44 g

(D) O2, 0.28 g 33. Haemoglobin contains 0.33% of iron by

weight. The molecular weight of haemoglobin is approximately 67200. The number of iron atoms (At. wt. of Fe = 56) present in one molecule of haemoglobin is _______.

[CBSE PMT 1998]

(A) 1 (B) 2

(C) 4 (D) 6 34. The number of water molecules is maximum

in _______.

[AIPMT RE-TEST 2015]

(A) 18 gram of water

(B) 18 moles of water

(C) 18 molecules of water

(D) 1.8 gram of water 35. If Avogadro number NA, is changed from

6.022 1023 mol1 to 6.022 1020 mol1, this would change _______.

[AIPMT RE-TEST 2015]

(A) the ratio of chemical species to each other in a balanced equation

(B) the ratio of elements to each other in a compound

(C) the definition of mass in units of grams

(D) the mass of one mole of carbon 36. Suppose the elements X and Y combine to

form two compounds XY2 and X3Y2. When 0.1 mole of XY2 weighs 10 g and 0.05 mole of X3Y2 weighs 9 g, the atomic weights of X and Y are _______ .

[NEET P-II 2016]

(A) 30, 20 (B) 40, 30

(C) 60, 40 (D) 20, 30

14


Classical Thinking

1. (C) 2. (C) 3. (B) 4. (A) 5. (C) 6. (D) 7. (D) 8. (A) 9. (D) 10. (D)

11. (C) 12. (B) 13. (D) 14. (A) 15. (D) 16. (C) 17. (B) 18. (C) 19. (C) 20. (B)

21. (B) 22. (B) 23. (D) 24. (A) 25. (A) 26. (D) 27. (B) 28. (D) 29. (A) 30. (C)

31. (A) 32. (A) 33. (D) 34. (C) 35. (C) 36. (C) 37. (C) 38. (A) 39. (B) 40. (A)

41. (B) 42. (A) 43. (C) 44. (D) 45. (D) 46. (C) 47. (B) 48. (D) 49. (A) 50. (C)

51. (A) 52. (C) 53. (C) 54. (A) 55. (A) 56. (A) 57. (C) 58. (A) 59. (B) 60. (A)

61. (D) 62. (B) 63. (A) 64. (C) 65. (D) 66. (C) 67. (C) 68. (D) 69. (C)

Critical Thinking

1. (B) 2. (C) 3. (C) 4. (D) 5. (C) 6. (A) 7. (D) 8. (B) 9. (C) 10. (C)

11. (A) 12. (C) 13. (D) 14. (C) 15. (C) 16. (B) 17. (B) 18. (B) 19. (D) 20. (A)

21. (B) 22. (C) 23. (C) 24. (B) 25. (A) 26. (D) 27. (D) 28. (C) 29. (B) 30. (B)

31. (A) 32. (A) 33. (B) 34. (C) 35. (A) 36. (C) 37. (A) 38. (B) 39. (A) 40. (C)

41. (A) 42. (C) 43. (D)

Competitive Thinking 1. (A) 2. (A) 3. (B) 4. (C) 5. (A) 6. (C) 7. (A) 8. (D) 9. (C) 10. (B) 11. (A) 12. (C) 13. (C) 14. (C) 15. (D) 16. (A) 17. (A) 18. (C) 19. (C) 20. (B) 21. (B) 22. (D) 23. (C) 24. (A) 25. (D) 26. (A) 27. (D) 28. (B) 29. (B) 30. (B) 31. (B) 32. (A) 33. (C) 34. (B) 35. (D) 36. (B)

Classical Thinking

17. BaCl2 + H2SO4 HCl + BaSO4

20.8 + 9.8 = 7.3 + x

x = 23.3

22. CH4 + 2O2 CO2 + 2H2O

(1 vol.) (2 vol.) (1 vol.) (2 vol.)

1 cm3 of CH4 requires 2 cm3 of O2 for its

complete combustion

0.25 cm3 of CH4 gives 0.5 cm3 of O2. 29. u means unified mass. 37. Volume of a gas is directly proportional to its

number of moles.

41. Molecular mass of H2 = 2 g 2 g will contain 6.02 1023 molecules of H2.

1 g of H2 will contain 236.02 10

2

molecules

= 3.01 1023 molecules 44. Molecular formula of benzene is C6H6. Molecular mass = sum of atomic weight of all

the atoms Molecular mass = 12 6 + 6 1 = 72 + 6 = 78 According to Avogadro’s law, the gram

molecule of benzene is equal to 78 g of C6H6. 45. Atomic mass of the element = 1.792 1022 6.022 1023 = 108 46. 1 mole of ozone(O3) = 48 g

0.5 mole of ozone(O3) = 0.5 48

1

= 24 g.

Answer Key

Hints

15


47. Number of molecules = n 6.022 1023

Now, n = mass of oxygen

molar mass of oxygen =

16

32 = 0.5

Number of molecules = 0.5 6.022 1023 = 3.011 1023 48. 2 gram mole of CO2 88 gm of CO2 0.1 mole of sucrose (C12H22O11) 34.2 gm of

sucrose 1 gram atom of calcium 40 gm of calcium 1.5 mole of water 2.7 gm of water 1.5 mole of water weighs the least amongst

the given options.

49. 16 g O2 has number of moles 16 1

32 2

14 g N2 has number of moles 14 1

28 2

Number of moles are same, so number of molecules are same.

51. V n Number of moles (n)

= Mass of thesubstance

Molar mass of thesubstance

n = mass

atomic mass(M)

V n 1

M

Atomic Mass of O = 16 Atomic Mass of N = 14

(O) (O) (N)

(N) (N) (O)

V n M

V n M

(O)

(N)

V 14 7

V 16 8

The ratio is 7 : 8 55. Urea [H2N – CO – NH2]

60 gm of urea contains 12 g of carbon.

100 gm of urea contains 12

60 100 = 20%

56. In Fe(CNS)3.3H2O

% of H2O = 3 18

284

100 = 19%

57. 40 gm NaOH contains 16 gm of oxygen

100 gm of NaOH contains16

40100 = 40% oxygen.

58.

Elements%

Composition Atomic Mass

Moles Ratio

A 25 12.5 25

12.5= 2

2

2= 1

B 75 37.5 75

37.5= 2

2

2= 1

Hence, the simplest formula of the compound

is AB. 67. 1L = 103m3 = 103cm3 = 1dm3 = 103ml. Critical Thinking 3. Mixture of liquids may be homogeneous or

heterogeneous mixtures. 12. N2 + 3H2 2NH3 (1 vol.) (3 vol.) (2 vol.) 3 volumes of H2 gives 2 volumes of ammonia

2 L of H2 will give = 2 2

3

L of ammonia

= 1.33 L of ammonia 14. A molecule cannot be divided into its

constituent atoms by simple methods. However, under drastic conditions, it can be decomposed into constituent atoms.

15. Isotopes are the atoms of the same element

having same atomic number (i.e., containing same number of protons and electrons) but different mass number (i.e., different number of neutrons).

16. The natural abundance means percentage

occurrence. 18. Molecular weight of sodium oxide (Na2O) = 46+16 = 62 62 gm of Na2O = 1 mole 620 gm of Na2O = 10 moles.

19. d = M

V (d = density, M= mass, V = volume)

Since d = 1 So, M = V 18 gm = 18 mL 18 mL = NA molecules (NA = Avogadro's

number)

1000 mL = AN

18 1000 = 55.555 NA.

16


20. 1 mole of sucrose contains 6.023 1023 molecules

1 molecule of sucrose has 45 atoms

6.023 1023 molecule of sucrose has 45 6.023 1023 atoms/mole 22. 6.022 1023 atoms of H weighs 1 g.

Mass of 1 atom of hydrogen = 23

1

6.022 10

= 1.6 1024 g 23. 1 gram atom of nitrogen 22.4 L of N

(molar volume at N.T.P.) 11.2 L of N2 24. At S.T.P, 22.4 dm3 of any gas 6.023 1023 molecules 25. 4Al + 3O2 2Al2O3

(108 g) (96 g) (204 g)

108 g Al combines with 96 g of O2

27 g Al combines with 96 27

108

= 24 g

26. 1 mole 6.022 1023

electrons

One electron weighs 9.108 1031 kg

1 mole of electrons weighs

6.022 1023 9.108 10–31 kg

Number of moles that will weigh 1 kg

= 23 31

1

6.022 10 9.108 10 moles

1

9.108 6.022 108 moles of electrons will

weigh one kilogram.

27. 18 g of H2O 1 mole = 3 NA atoms

(Number of atoms = n NA Atomicity)

16 g of O2 1

2 mole = 2

1

2 NA atoms

4.4 g of CO2 1

10mole 3

1

10 NA atoms

16 g of CH4 1 mole = 5 NA atoms maximum number of atoms is present in 16 g

of CH4. 28. Number of atoms = n NA Atomicity Number of S atoms = 6.022 1023 0.2 8 9.63 1023

29. 3CaCO10 g

CaO + CO2

10 g of 90% pure CaCO3 9 gm of CaCO3 Number of moles of CaCO3

= 3

3

Mass of CaCO in gm

Molecular weight of CaCO in gm

Number of moles of CaCO3 = 9

100

= 0.09 mole CaCO3 CO2 = 0.09 mole At NTP Vol. CO2 = 0.09 22.4 = 2.016 L. 31. Molecular mass of CH3COOH = 60

% of C in CH3COOH = 24

60 100 = 40%

Similarly, Molecular weight of C6H12O6 = 180

% of C in C6H12O6 =72

180 100 = 40%

32. Molecular mass of Urea i.e., H2N – C – NH2 is 60. urea has 2 N atoms. mass of N = 2 14 = 28

% of N in urea = 28

60 100 = 46.6 %

33. % of X = 75.8 % of Y = 100 – 75.8 = 24.2 Element %

Composition Atomic ratio Simplest

ratio X 75.8 75.8

75= 1.011

1.011

1.011= 1

Y 24.2 24.2

12= 2.02

2.02

1.011=

1.99 2 Thus, empirical formula of the compound is

XY2. 34. Empirical formula mass = CH = 12 + 1 = 13 Molecular mass = 78

r = Molecular mass

Empiricalmass=

78

13= 6

Molecular formula = r Empirical formula

= 6 CH

= C6H6

O

17


35. Since, 0.0835 mole of compound contains 1 gm of hydrogen

1 g mole of compound contain

= 1

11.970.0835

12 g of hydrogen.

12 g of H is present in C2H12O6 36. CH4 + 2O2 CO2 + 2H2O

(12+4) (232) (12+32) (218)

16 g 62 g 44 g 36 g

16 g CH4 44 g of CO2

33g of CO2 33 16

44

= 12 g of CH4

Number of moles = 12

16

= 0.75 moles of methane

37. N2 + 3H2 2NH3

10 mL 30 mL 20 mL Thus, ammonia obtained in the reversible

reaction of NH3 is 20 mL.

38. 3CaCO2 3CaO + 3CO2

Molecular mass of CaO is = 40 + 16

= 56

3 moles of CaO is formed in the reaction.

3 56 = 168 g of CaO is formed. 39. 2H2 + O2 2H2O

Ratio of mole of reactant, H2 : O2 = 2

1 = 2

Actual amount 3 g of H2 and 29 g of O2.

3

2 = 1.5 mol of H2 and

29

32= 0.91 mol of O2.

Ratio of actual moles of H2 : O2 = 1.5

0.91 = 1.66

Theoretical ratio of moles of H2 : O2 = 2 Actual ratio of moles of H2 : O2 = 1.66 Hence, H2 is limiting reactant and O2 is excess reactant.

40. Baking powder or sodium hydrogen carbonate

(NaHCO3) is a compound. Diamond and charcoal are different physical forms of the carbon element.

41. mass of one atom =1

12 Atomic mass amu

mass of 1 He atom =1

12 4 =

1

3amu

mass of x He atoms = 6 amu x = 6 3 = 18 atoms. 42. In compound B, 32 parts of X react with 84

parts of Y. In compound B, 16 parts of X react with 42

parts of Y. In compound C, 16 parts of X react with x parts of Y. The ratio of masses of Y which combine with fixed mass of X in compounds B and C is 3:5.

B 42 3 C x 5

x = 42 5

3

= 70

Competitive Thinking 5. Molecular weight of C60H122

= 12 60 + 122 1 = 720 + 122 = 842

6 1023 molecules = 842 g

1 molecule = 23

842

6 10

= 140.33 1023

= 1.4 1021 g

6. Atomic mass =10 19 81 11

100

= 1081

100 = 10.81

7. Average atomic mass

= 200 90 199 8 202 2

100

= 199.96 200

8. Molecular mass of NH3 = 14 + (3 1) = 17

Number of moles = 4.25

17 = 0.25

Number of molecules of NH3 = 0.25 NA

= 1.506 1023 molecules One molecule of NH3 contains 4 atoms.

18


1.506 1023 molecules will contain = 1.506 1023 4 = 6.023 1023 atoms 6 1023 atoms. 9. Molecular mass of CuSO4.5H2O

= Cu + 32 + (9 O) + (10 H)

= 63.5 + 32 + (9 16) + (10 1)

= 249.5 g

6.023 1023 molecules of CuSO4.5H2O weighs 249.5 g

1 1022 molecules of CuSO4.5H2O

will weigh 22

23

249.5 1 10

6.023 10

= 4.159 g

10. 6.022 1023 dioxygen molecules are present in

1 mole i.e., 32 g of dioxygen. 1.8 1022 dioxygen molecules will be present

in 22

23

1.8 10 32

6.022 10

= 0.96 g of dioxygen.

11. Number of moles in 4.4 g of CO2

= 4.4

44 = 0.1

Number of oxygen atoms in 1 mole of CO2

= 2 NA

Number of oxygen atoms in 0.1 mole of CO2

= 0.1 2 NA

= 0.2 6.022 1023

= 1.20 1023

12. Volume occupied by 1 mole of any gas at STP

= 22.4 dm3

Volume occupied by 4.4 g of CO2 i.e., 0.1 mole of CO2 at STP = 2.24 dm3 = 2.24 L

13. Number of atoms

= n NA Atomicity

= 0.1 NA 3 ( It’s a triatomic gas)

= 6.02 1022 3

= 18.06 1022

= 1.806 1023

14. Total number of atoms in molecule

= n NA Atomicity

Atomicity’s of NH3, O2, H2 and He molecules are 4, 2, 2 and 1 respectively.

17 g of NH3 = 4 NA atoms

4.25 g of NH3 = 4.25

17

4 NA = NA atoms

32 g of O2 = 2 NA atoms

8 g of O2 = 2 8

32

NA = AN

2atoms

2 g of H2 = 2 NA atoms

= 2NA atoms 4 g of He = NA atoms Thus, the system that contains 2 g of H2 has

maximum numbers of atoms. 15. BaCO3 BaO + CO2

Molecular weight of BaCO3

= 137 + 12 + (3 16)

= 197 22.4 L of CO2 is released by 197 g of BaCO3

x L of CO2 is released by 9.85 g of BaCO3

x = 22.4 9.85

197

= 1.12 L

16. Since, 1 mole of BaCO3 contains 3 moles of

oxygen

1.5 moles of oxygen 1

3 1.5 =

1

2 = 0.5

moles of BaCO3

17. 1 L of air = 1000 mL = 1000 cc. 1000 cc of air contains 210 cc of O2 1 mole = 22.4 L = 22400 cc.

number of moles of O2 = 210

22400 = 0.0093 moles.

18. 22.4 L of water contains NA molecules 22400 mL of water contains NA molecules 1 mL of water i.e., 20 drops contains

AN

22400molecules

1 drop contains AN

22400 20=

23

3

6.022 10

448 10

= 1.344 1018

19. BCl3 + 3

2H2 B + 3HCl

3

2moles of H2 1 mole of B

19


number of moles of B = 21.6

10.8 = 2

x moles of H2 2 moles of B moles of H2 = 3

volume = 3 22.4 = 67.2 L 20. Molecular weight of PbO = 207.2 + 16 = 223.2 Molecular weight of PbCl2 = 207.2 + 2(35.5) = 278.2 223.2 g of PbO gives 278.2 g of PbCl2

6.5 g of PbO will give = 6.5 278.2

223.2

= 8.102 g of PbCl2

moles of PbCl2 obtained = 8.102

278.2 = 0.029 moles.

21. Empirical formula of an acid is CH2O2

(Empirical formula) n = Molecular formula n = whole number multiple i.e., 1,2,3,4.............. If n = 1 molecular formula CH2O2. 22. C = 24 g, H = 4 g, O = 32 g So, Molecular formula = C2H4O2 So, Empirical formula = CH2O (Simplest formula). 23. Relative number of atoms of

C = 38.71

12 = 3.22

H = 9.67

1 = 9.67

O = 51.62

16 = 3.22

Simplest ratio : C : H : O

3.22

3.22:

9.67

3.22:3.22

3.22

1 : 3 : 1

Empirical formula = CH3O 24. Since, 0.5 g Se 100 gm peroxidase anhydrous enzyme

78.4 g Se = 100 78.4

0.5

= 1.568 104

Minimum molecular mass of peroxidase anhydrous enzyme means molecule atleast contains one selenium atom.

25. H2O H2 + 1

2O2

H2 : O2 = 2 : 1

for 2.24 dm3 of O2, H2 liberated will be 4.48 dm3 26. C3H8 + 5O2 3CO2 + 4H2O

(1 vol) (5 vol)

1 L of propane gas will require 5 L of O2

27. 2Al(s) + 6HCl(aq) 3(aq)2Al + (aq)6Cl + 3H2(g)

For each mole of HCl reacted 1

2mole of H2 is

formed.

volume of H2 gas formed at STP per mole of

HCl is =1

2 22.4 = 11.2 L.

28. 1 mole of Mg3(PO4)2 gives 8 moles of O2

x moles will give 0.25 moles of O2

x = 0.25

8= 3.125 102 moles.

29. 2H2(g) + O2(g) 2H2O(g)

Ratio of moles of reactants, H2 : O2 = 2 : 1

Actual amount of reactants: 10 g H2 and 64 g O2

Actual moles of reactants: 5 mol H2 and 2 mol O2

Ratio of actual moles of reactants,

H2 : O2 = 5 : 2 = 2.5 : 1

The limiting reactant is O2.

Now, 1 mole of oxygen gives 2 moles of water. Hence, 2 moles of oxygen will give 4 moles of water.

30. Both O2 and N2 are diatomic.

ratio of their number of molecules will be equal to ratio of their number of moles.

number of moles of O2 = 2weight of O

32

number of moles of N2 = 2weight of N

28

number of moles of O2 = number of moles of N2

2weight of O

32 = 2weight of N

28

20


1

32 :

4

28

1

32 :

1

7

ratio is 7 : 32 31. MgCO3(s) MgO(s) + CO2(g)

Molar mass of MgCO3 = 84 g mol1

Number of moles of MgCO3 = 20

84

= 0.238 mol

1 mole MgCO3 gives 1 mole MgO

0.238 mole MgCO3 will give 0.238 mole MgO.

Molar mass of MgO = 40 g mol1

0.238 mole MgO = 40 0.238 = 9.52 g MgO

Theoretical yield of MgO = 9.52 g Practical yield of MgO is 8.0 g

Percentage purity = 8

9.52100

= 84% 32. 2Mg + O2 MgO

(224) (32) 48 g of Mg requires 32 g of O2

0.56 g of O2 requires =0.56 48

32

= 0.84 g of Mg

Mg left = 1 0.84 = 0.16 g 33. 100 g of haemoglobin contains 0.33 g of Fe

67200 g of haemoglobin contains

= 67200 0.33

100

= 221.76 g of Fe.

Number of atoms of Fe = 221.76

56

= 3.96 4 34. 1 mole of water = 18 g of water

= 6.022 1023 molecules of water

18 moles of water = 18 6.022 1023 molecules of water

= 1.08396 1025 molecules of water

35. When Avogadro number is 6.022 1023 mol1, the mass of 1 mol of carbon = 12 g

Mass of 1 mol of carbon when Avogadro

number is 6.022 1020 mol1

= 20

23

12 6.022 10

6.022 10

= 12 103 g

Thus, the mass of 1 mol of carbon is changed. 36. 0.1 mol of XY2 = 10 g

1 mol of XY2 = 100 g i.e, Molecular weight of XY2 = 100 0.05 mol of X3Y2 = 9 g

1 mol of X3Y2 = 180 g i.e., Molecular weight of X3Y2 = 180 Let atomic weights of X and Y be x and y

respectively.

x + 2y = 100 ….(1) 3x + 2y = 180 ….(2) Substrating (1) from (2),

2x = 180 100

x = 40 Substituting x = 40 in (1), 40 + 2y = 100

y = 30

21


1. Weight of 112 mL of oxygen at NTP on

liquefaction would be _______. (A) 0.32 g (B) 0.64 g (C) 0.16 g (D) 0.96 g 2. The largest number of molecules is present in

_______. (A) 54 g of nitrogen peroxide (B) 28 g of carbon dioxide (C) 36 g of water (D) 46 g of ethyl alcohol 3. The mass of a molecule of water is _______. (A) 3 1026 kg (B) 3 1025 kg (C) 1.5 1026 kg (D) 2.5 1026 kg 4. The modern atomic weight scale is based on

_______. (A) 12C (B) 16O (C) 1H (D) 13C 5. Which of the following gives CORRECT

order of increasing masses? (Atomic mass: N = 14, O = 16, Cu = 63). I. 1 molecule of oxygen II. 1 atom of nitrogen III. 1 1010 g molecular weight of oxygen IV. 1 1010 g atomic weight of copper (A) II < I < III < IV (B) IV < III < II < I (C) II < III < I < IV (D) III < IV < I < II 6. What is the weight of oxygen required for the

complete combustion of 2.8 kg of ethylene? (A) 2.8 kg (B) 6.4 kg (C) 9.6 kg (D) 96 kg 7. 100 mL of PH3 on decomposition produced

phosphorus and hydrogen. The change in volume is _______.

(A) 50 mL increase (B) 500 mL decrease (C) 900 mL decrease (D) nil 8. The molecular weight of a gas is 45. Its

density at STP is _______. (A) 22.4 (B) 11.2 (C) 5.7 (D) 2.0 9. A sample of phosphorus trichloride (PCl3)

contains 1.4 moles of the substance. How many atoms are there in the sample?

(A) 4 (B) 5.6 (C) 8.431 1023 (D) 3.372 1024

10. In the reaction, 4NH3(g) + 5O2(g) 4NO(g) + 6H2O(g), When

1 mole of ammonia and 1 mole of O2 are made to react to completion _______.

(A) 1.0 mole of H2O is produced (B) 1.0 mole of NO will be produced (C) all the oxygen will be consumed (D) all the ammonia will be consumed 11. What weight of SO2 can be obtained by

burning sulphur in 5.0 moles of oxygen? (A) 640 grams (B) 160 grams (C) 80 grams (D) 320 grams 12. The INCORRECT statement for 14 g of CO is

_______. (A) it occupies 2.24 L at NTP

(B) it corresponds to 1

2 mole of CO

(C) it corresponds to half mole of N2 (D) it corresponds to 3.011023 molecules of CO 13. The number of gram atom of oxygen in

6.02 1024 CO molecules is _______. (A) 1 (B) 0.5 (C) 5 (D) 10 14. A hydrocarbon contains 80% carbon. What is

the empirical formula of the compound? (A) CH2 (B) CH3 (C) CH4 (D) C2H3 15. Which of the following drug-ailment pairs is

CORRECT? (A) Tamiflu cancer (B) Cisplatin AIDS (C) L-dopa Parkinsons disease (D) Taxol diabetes

1. (C) 2. (C) 3. (A) 4. (A)

5. (A) 6. (B) 7. (A) 8. (D)

9. (D) 10. (C) 11. (D) 12. (A)

13. (D) 14. (B) 15. (C)

Evaluation Test

Answers to Evaluation Test

386


1. In which of the following, C – H bond length is minimum? (A) sp s overlapping (as in alkynes). (B) sp2 s overlapping (as in alkenes). (C) sp3 s overlapping (as in alkanes). (D) All have the same bond length. 2. A straight chain hydrocarbon has the molecular formula C8H10. The hybridisation for the carbon atoms from

one end of the chain to the other are respectively sp3, sp2, sp2, sp3, sp2, sp2, sp and sp. The structural formula of the hydrocarbon would be _______.

(A) CH3 C C CH2 CH = CH CH = CH2 (B) CH3 CH2 CH = CH CH2 C C CH = CH2 (C) CH3 CH = CH CH2 C C CH = CH2 (D) CH3 CH = CH CH2 CH = CH C CH 3. Which of the following reagents will be able to distinguish between but-1-yne and but-2-yne? (A) NaNH2 (B) HCl (C) O2 (D) Br2 4. The number of types of bonds between two carbon atoms in calcium carbide is _______ bond(s). (A) one sigma and two pi (B) one sigma and one pi (C) two sigma and one pi (D) two sigma and two pi 5. In the molecule, CH3 CH2 C CH, % p-character in C2 carbon atom is _______. (A) 50% (B) 75% (C) 25% (D) 67%

1. (A) 2. (D) 3. (A) 4. (A) 5. (A)

Answers to Evaluation Test

Evaluation Test

std triumph chemistry - target publications › ... › std-11-chemistry-mcqs.pdf · triumph...

Documents