BASIC CONCEPTS

Objectives

MULTIPLE CHOICE QUESTIONS

1. The number of moles of CO₂ which contain 8g of oxygen: (Board 2014) 11201066

(a) 0.25 (b) 0.5

2. 27g of Al will react completely with how much mass of O₂ to produce Al₂O₃?

(a) 8g (b) 16g 11201067

3. One mole of SO₂ contains: 11201068

(a) 6.02 ´ 10²³ atoms of oxygen

(b) 18.1 ´ 10²³ molecules of SO₂

4. The largest number of molecules are present in: (Board 2013, 2014) 11201069

(a) 3.6 g of H₂O (b) 4.6 g of C₂H₅OH

5. Silver has isotopes: 11201070

(a) 9 (b) 10

6. Isotopes differ in: 11201071

(a) Properties which depend upon mass

(b) Arrangement of electrons in orbitals

(d) The extent to which they may be affected in electromagnetic field

7. The volume occupied by 1.4g of N₂ at STP is: (Board 2013, 2014,2015) 11201072

(a) 22.4 dm³ (b) 2.24 dm³

8. One mole of carbon –12 has mass:11201073

(a) 0.012 kg (b) 1 kg

9. Which of the following compounds does not have empirical formula CH₂O?

(a) HCHO (b) C₆H₁₂O₆11201074

10. Which of the following compounds contains the highest percentage by mass of nitrogen? 11201075

(a) NH₃ (b) N₂H₄

11. Which of the following sets contain only compounds? 11201076

(a) Air, water and sodium

(b) Hydrogen, oxygen and Ammonia

(d) Milk, air and sulphur

12. The mass of one mole of chlorine gas is:

(a) 35.5 g (b) 71 g 11201077

13. A compound contains 87.5% Si and 12.5% H. What is empirical formula of compound? (Relative atomic masses H=1 and Si = 28) 11201078

(a) SiH₂ (b) SiH₃

14. The ratio of number of molecules in 71g of gaseous chlorine to the number of molecules in 2g of gaseous hydrogen is:

(a) 71:1 (b) 71:2 11201079

15. The equation of burning of hydrogen in oxygen is: 11201080

2H_2(g) + O_2(g) ¾¾® 2H₂O_(g)

Which one is a correct description of this reaction?

(a) 2 atoms of hydrogen combine with 2 atoms of oxygen

(b) 2 moles of steam can be produced from 0.5 mole of oxygen

(d) 2g of hydrogen combine with 1 g of oxygen

16. A hydrocarbon contains C = 80% and H=20%. What is its empirical formula?

(a) CH₂ (b) CH₃11201081

17. What are the number of moles of hydrogen atom in 3.2g of CH₄? (Relative atomic mass of C = 12) 11201082

(a) 0.2 (b) 0.4

18. What are the number of moles of oxygen in 11g of CO₂? 11201083

(a) 0.25 (b) 0.50

19. A ring has 6g of diamond in it. What is the number of atoms in it? 11201084

(a) 6.02 ´ 10²³ (b) 12.04 ´ 10²³

20. The number of H₂O molecules in 9 grams of ice is: 11201085

(a) 3.01 ´ 10²³ (b) 6.02 ´ 10²³

21. 32g of oxygen gas contains: 11201086

(a) 6.02 ´ 10²³molecules

(b) 6.02 ´ 10²³ atoms

(d) 24.04 ´ 10²³ atoms

22. How many atoms are present in one mole of water? 11201087

(a) 3 (b) 54

23. One mole of water contains: 11201088

(a) 81g water

(b) 6.02 ´ 10²³ atoms

(d) 6.02 ´ 10²³ molecules

24. The number of molecules in 8g of oxygen gas is: 11201089

(a) 6.02 ´ 10²³(b) 3.01 ´ 10²³

25. 6g of hydrogen gas is: 11201090

(a) 1 mole (b) 2 moles

26. In SI units, the prefix “Nano” means:

(a) 10^-⁶ (b) 10^-⁹11201091

27. Which element cannot be analysed directly by combustion analysis? 11201092

(a) Hydrogen (b) Oxygen

28. A compound has empirical formula C₃H₃O and molecular mass is 110. The actual molecular formula of the compound is: 11201093

(a) C₃H₃O (b) C₆H₆O₂

29. The reactant which consumes earlier and gives least quantity of product is called: 11201094

(a) Reactant

(b) Limiting reactant

(d) Stoichiometric amount

30. The number of atoms present in 6g of Mg metal is: 11201095

(a) 6.02 ´ 10²³(b) 3.01 ´ 10²³

31. A substance that can exist independently is called: 11201096

(a) Atom (b) Molecule

32. The actual number of atoms present in a compound is called: 11201097

(a) Structural formula

(b) Empirical formula

(d) Formula mass

33. The sum of atomic masses of all the elements present in a molecule is called:

(a) Atomic mass 11201098

(b) Ionic mass

(d) Empirical formula mass

34. Which of the following substances is used as water absorber in combustion analysis? (Board 2014) 11201099

(a) Mg (CIO₄)₂

(b) 5% KOH

(d) Dilute solution of NaOH

35. Which of the following properties is always in whole number ? 11201100

(a) Atomic mass

(b) Atomic radius

(d) Atomic number

36. The total number of oxygen atoms in 22g of CO₂ gas is: (Board 2015) 11201101

(a) 6.023 ´ 10²³(b) 3.01 ´ 10²³

37. The mass of 0.5 mole of Al is: 11201102

(a) 12g (b) 13.5g

38. The mass of 1.505 ´ 10²³ atoms of S:

(a) 0.5g (b) 0.6g 11201103

39. How many moles of water are produced by burning 4 moles of H₂ with excess of oxygen? 11201104

(a) 1 mole (b) 2 moles

40. The number of moles of CO₂ in 11g of gas is: 11201105

(a) 0.2 mole (b) 0.25 mole

41. How many moles of CO are present is 12.4 ´ 10²³ molecules of CO? 11201106

(a) 0.5 mole (b) 1.0 mole

42. How many atoms are present in half mole of oxygen gas? 11201107

(a) 3.01 ´ 10²³ (b) 6.023 ´ 10²³

43. The number of Al³⁺ ions in AlCl₃ is 2.007´10²³. The number of Cl^- ions are: 11201108

(a) 6.02 ´ 10²³

(b) 3.01 ´ 10²³

(d) 1.5 ´ 10²³

44. The volume occupied by 32g of O₂ at STP is: 11201109

(a) 44.2dm³ (b) 22.4dm³

45. The ratio of number of molecules of 2g H₂ gas to number of molecules of 64g gaseous oxygen is: 11201110

(a) 1 : 1 (b) 1 : 2

46. What is the ratio of volumes of 2g of H₂ to the volume of 16g CH₄ both volumes are at STP? 11201111

(a) 1 : 8 (b) 1 : 2

47. A mixture of 8g of H₂ with 8g of O₂ is ignited 2H₂ + O₂ ® 2H₂O

What is the mass of water formed?

11201112

(a) 9 g (b) 16 g

48. What is the maximum mass of chromium, that can be extracted from 76g of Cr₂O₃? 11201113

( relative atomic mass of Cr=52;O=16 )

(a) 48g (b) 52g

49. What is the mass of oxygen obtained from 72g of pure water? 11201114

(a) 16g (b) 32g

50. In mass spectrometer, the ions are accelerated by applying potential difference of: 11201115

(a) 2000 V (b) 500 V

51. What is the number of protons in molecule of SO₃? 11201116

(a) 24 (b) 32

52. During combustion analysis of organic compounds CO₂ is absorbed in: 11201117

(a) Mg(ClO₄)₂ (b) H₂SO₄

53. Which of the following has minimum mass? 11201118

(a) 1 mole of S

(b) 79 grams of Ag

(d) 3 ´ 10²³ atoms of C

54. Which of the following has maximum mass? 11201119

(a) 2 moles of P

(b) 5 moles of H₂O

(d) 1 mole of glucose

55. Molar volume of a gas at STP is equal to: 11201120

(a) 1 gram of a gas

(b) 6.02 ´ 10²³g of a gas

(d) 1 gram molecule

56. 4g of CH₄ at STP has molecules. 11201121

(a) 6.02 x 10²³ (b) 3.01 x 10²³

57. 1 Mole of OH^-1 ion is equal to: 11201122

(a) 18g (b) 16g

58. One cm³ of H₂ at STP contains: 11201123

(a) 6.02 ´ 10²³ atoms

(b) 1 ´ 10²⁰ atoms

(d) 1.687 ´ 10²⁴ atoms

59. How many carbon atoms are present in 90g glucose? 11201124

(a) 6.02 ´ 10²³ (b) 1.8 ´ 10²³

60. The mass of one molecule of O₂ is: 11201125

(a) ´ 10²³ (b) ´ 10²³

61. Mass of 2 moles of CO₂ is: 11201126

(a) 44g (b) 88g

62. Number of molecules in 1 dm³ of steam at STP is: 11201127

(a) 55.5 ´ 6.02 ´ 10²³

(b) 1000 ´ 6.02 ´ 10²³

(d) 0.268 x 10²³

63. %age of Nitrogen in NH₃ is:11201128

(a) 14/34 ´ 100 (b) 14/17 ´ 100

64. The number of isotopes of Pd is: 11201129

(a) 4 (b) 5

65. Number of moles in 100g of KClO₃: 11201130

(a) 0.76 (b) 0.56

66. How many atoms of Oxygen are present in 90 grams of Glucose? 11201131

(a) 1.8 ´ 10²⁴ (b) 6.02 ´ 10²³

67. The number of natural isotopes is:

(a) 280 (b) 150 11201132

68. The mass of one mole of electron is:

(Board 2014) 11201133

(a) 1.088 mg (b) 0.55 mg

69. Mass spectrometer is used to determine:

(a) Mass of electrons 11201134

(b) Mass of protons

(d) Mass of isotopes

70. The number of Isotopes of Gold (Au) is:

(a) 1 (b) 3 11201135

71. The atomic mass of Fluorine is: 11201136

(a) 8 (b) 18

72. Height of peak in mass spectrum shows:

(a) Number of isotopes 11201137

(b) Mass number

(d) Number of protons

73. Molecular mass of CaCO₃ is: 11201138

(a) 100 (b) 90

74. Select the most suitable answer from the given ones: (Board 2013) 11201139

(a) Isotopes with even atomic masses are comparatively abundant.

(b) Isotopes with odd atomic masses are comparatively abundant.

(d) Isotopes with even atomic masses and odd atomic numbers are comparatively abundant.

75. Many elements have fractional atomic masses. This is because: 11201140

(a) The mass of the atom is itself fractional

(b) Atomic masses are average masses of isobars

(d) atomic masses are average masses of isotopes proportional to their relative abundance

76. A limiting reactant is the one which 11201141

(a) is taken in lesser quantity in grams as compared to other reactants

(b) is taken in lesser quantity in volume as compared to the other reactants

(d) gives the minimum amount of the product under consideration

77. A compound (60g) on analysis gave 24g C, 4g H and 32g O, its expirical formula will be: 11201142

(a) CH₂O (b) C₂H₂O

78. A mass spectrograph is a plot between: 11201143

(a) Relative abundance of isotopes and strength of electric field.

(b) Relative abundance of isotopes and strength of magnetic field.

(d) Relative abundance of isotopes and mass no.

79. In a spectrometer, which type of particles strike the detector: 11201144

(a) Monoatomic (b) Gaseous atoms

80. If 16 grams of O₂ react with excess C₂H₆, how many grams of CO₂ will be formed? 11201145

(a) 22g (b) 13g

81. How many unstable radioactive isotopes have been produced through artificial disintegration method? 11201146

(a) 280 (b) 300

SHORT QUESTIONS

Q1. Give justification of the following statements. 11201147

(i) Law of conservation of mass has to be obeyed during stoichiometric calculations. (ii) Many chemical reactions taking place in our surrounding involves the limiting reactants. 11201148 (iii) No individual Ne atom in the sample of the element has a mass of 20.18 amu.

(Board2014) 11201149

Q2. One mole of H₂SO₄ should completely react with two moles of NaOH. How does Avogadro’s number help to explain it?

11201150

Q3. One mole of H₂O has two moles of bonds, three moles of atoms, ten moles of electrons and twenty-eight moles of the total fundamental particles present in it. 11201151

Q4. N₂ and CO have the same number of electrons, protons and neutrons. Explain it.

(Board 2014) 11201152

Q5. Explain that 22.414 dm³ of each gas have different mass but the same number of molecules. 11201153

Q6. Calculate the mass of 10^-³ moles of MgSO₄. 11201154

Q7. Give reason that formation of positive ion is an endothermic process. 11201155

Q8. Prove that one atom of Mg is twice as heavy as an atom of carbon. (Board 2014) 11201156

Q.9 Calculate the gram atom in 4.0g of K. 11201157

Q10. What is Avogadro’s number? Give equation to relate the Avogadro’s number and mass of an element. 11201158

Q11. Give reasons that one mole of different compounds have different masses but have same number of molecules. 11201159

Q12. How many molecules of water are there in 12g of ice? 11201160

Q13. Calculate the number of positive and negative ions dispersed when 2.35 ´ 10^-²³ molecules of H₂SO₄ are dissolved in solution. 11201161

Q14. Define Stoichiometry. (Board 2013) 11201162

Q15. Differentiate between limiting and non-limiting reactants. 11201163

Q16. If 4 moles of hydrogen reacts with 2 moles of oxygen, then how many moles of water are produced? 11201164

Q17. Distinguish between actual yield and theoretical yield. 11201165

Q.18 Actual yield is usually less than theoretical yield. Give reasons.

(Board 2015) 11201166

Q19. Define Isotopes. 11201167

Q20. Why are relative atomic masses expressed in fractional quantities? 11201168

Q21. NaCl has 58.5 amu as formula mass and not molecular mass. Discuss 11201169

Q22. Concept of limiting reactant is not applicable to the reversible reactions. Explain it. 11201170

Q.23 What is molar volume? 11201171

Q24. 23 g of Na and 238 g of U have equal number of atoms each. How? 11201172

Q25. What is the basic principle of mass spectrometry? 11201173

Q26. Isotopes have same chemical properties but different physical properties. How? 11201174

Q27. What is difference between molecular mass and Formula mass? 11201175

Q28. What is the approximate distance between the molecules in gaseous state?

11201176

Q29. One mole of CO₂ and NO₂ has same number of molecules. Explain it. 11201177

Q30. Give assumptions of Stoichiometry.

(Board 2013, 2014,2015) 11201178

Q31. One mg of K₂CrO₄ has thrice number of formula units when ionized in water. Explain. 11201179

Q32. Why percentage of oxygen cannot be determined directly in combustion analysis? 11201180

Q33. Define molecular ion. Write its uses.

(Board 2013, 2014) 11201181

Q34. Write function of Mg(ClO₄)₂ and KOH in combustion analysis. (Board 2014) 11201182

Q35. Why do we calculate %age yield?

(Board 2014) 11201183

Q36. What is the function of Electrometer in a mass spectrometer? 11201184

Q37. A compound may have same empirical as well as molecular formula. Justify.

(Board 2015) 11201185

Q38. Define molecular formula of a compound. How is it related with its empirical formula? (Board 2015) 11201186

EXPERIMENTAL TECHNIQUES IN CHEMISTRY

Objectives

MULTIPLE CHOICE QUESTIONS

1. Solvent extraction is an equilibrium process and it is controlled by:

(Board 2014, 15) 11202018

(a) Law of mass action

(b) The amount of solvent used

(d) The amount of solute

2. The comparative rates at which the solutes move in paper chromatography, depends on. (Board 2014) 11202019

(a) The size of the paper

(b) Their R_f value

(d) The polarity of solvent used

3. A filtration process could be very time consuming if it were not aided by a gentle suction, which is developed:

11202020

(a) If the paper covers the funnel upto its circumference

(b) If the paper has got small sized pores in it

(d) If the paper fits tightly

4. Solvent extraction method is particularly useful technique for separation when the product to be separated is:

(Board 2013)11202021

(a) Non-volatile or thermally unstable

(b) Volatile or thermally stable

(d) Volatile or thermally unstable

5. During the process of crystallization, the hot saturated solution: (Board 2013) 11202022

(a) is cooled very slowly to get large sized crystals

(b) is cooled at a moderate rate to get medium-sized crystals

(d) is mixed with an Immiscible liquid to get the pure crystals of the product

6. Which of the following precautions is necessary for smooth filtration? 11202023

(a) The filter paper should be of large size

(b) The tip of funnel should not touch the side of the beaker

(d) The stem of the funnel should remain continuously full of liquid

7. Distribution law is employed in which one of the following? 11202024

(a) Paper chromatography

(b) Sublimation

(d) Solvent extraction

8. Equilibrium is established during the process of solvent extraction and the phenomenon obeys: 11202025

(a) Distribution law

(b) Le-chatelier’s principle

(d) Law of chemical equilibrium

9. When hot saturated solution is cooled very rapidly, we get: 11202026

(a) Medium-sized crystals

(b) Large-sized crystals

(d) No crystallization

10. When I₂ present in aqueous layer in the form of I goes to CCl₄ layer then the change in colour is from: 11202027

(a) Purple to brown (b) Purple to green

11. The crystallization of a solid substance is done from a hot saturated solution. The solution is: 11202028

(a) Evaporated rapidly

(b) Cooled very slowly to get good crystals

(d) Mixed with another miscible solution

12. A sintered glass crucible can be used to:

(a) Filter HCl and KMnO₄ 11202029

(b) Crystallize the solid substance

(d) Avoid premature crystallization of the solute

13. When a solute distributes between a stationary phase and a mobile phase, then the process is called: 11202030

(a) Sublimation

(b) Crystallization

(d) Chromatography

14. The Iodine present in water can be separated by which one of the following techniques? 11202031

(a) Sublimation

(b) Chromatography

(d) Solvent extraction

15. For smooth and fast filtration the filter paper should be so large so that it is full of precipitates at the end of filtration up to: 11202032

(a) to (b)

16. I₂ present in water can be extracted by using which of the following solvents?

11202033

(a) Ethanol (b) Acetic acid

17. NaCl and sand can be separated by which one of the following techniques?

11202034

(a) Formation of solution and filtration

(b) Formation of solution and evaporation without filtration

(d) Chromatography

18. Which of the following pairs can be separated by sublimation? 11202035

(a) Sand and NaCl

(b) Sand and broken pieces of glass

(d) NaCl and KCl

19. The rate of filtration in fluted filter paper is greater as compared to cone-shaped filter paper due to the reason that: 11202036

(a) It has greater surface area

(b) It has greater size of the pores

(d) The process is simple

20. Which statement about Gooch Crucible is incorrect? 11202037

(a) It helps for the quick filtration by using suction

(b) The chemicals which react with paper can be filtered

(d) It is made up of porcelain

21. Which one of the following substances is used as a drying agent in desiccators?

11202038

(a) Diethyl ether (b) Bleaching powder

22. Which one of the following substances is used as decolourizing agent ? 11202039

(a) Animal charcoal (Board2014)

(b) Concentrated H₂SO₄

23. Which one of the following properties cannot be associated with a good solvent?

11202040

(a) There should not be chemical reaction between solute and solvent

(b) The solvent should not be able to dissolve the impurities easily

(d) It should have a very low boiling point

24. Chemical characterization of the substance is studied in: 11202041

(a) Physical chemistry

(b) Applied chemistry

(d) Biochemistry

25. In order to have good crystals of a substance the temperature of the system at the time of preparation of solution should be: 11202042

(a) Around 0^oC

(b) Around room temperature

(d) Just above the room temperature

26. Which one of the following substances is not used as a drying agent in a desiccator?

(Board 2014) 11202043

(a) CaCl₂ (b) P₂O₅

27. In order to dry the crystals safely, we should: 11202044

(a) Place them in an oven

(b) Evaporate the solvent at room temperature

(d) Press the precipitates between folds

of filter paper

28. When repeated extractions are performed by using small quantities of the solvent, then this process is thought to be more:

(a) Efficient (b) Rapid 11202045

29. Which of the following substances is a sublime material? 11202046

(a) Potash alum (b) NaCl

30. Which one of the following substances does not undergo sublimation? 11202047

(a) KMnO₄ (b) Naphthalene

31. In paper chromatography the point at which the solvent rises the maximum extent is called: 11202048

(a) Eluent

(b) Chromatogram

(d) Base line

32. Chromatography is the process which involves the distribution of a solute between: 11202049

(a) Two mobile phases

(b) A stationary phase and mobile phase

(d) Two stationary phases

33. Type of chromatography, in which stationary phase is a solid substance,` is called: 11202050

(a) Thin layer chromatography

(b) Partition chromatography

(d) Paper chromatography

34. The rate at which the solute moves in the paper chromatography depends upon:

(a) Distribution co-efficient 11202051

(b) Distribution law

(d) Law of partial pressure

35. In paper chromatography mobile phase is: 11202052

(a) Gas (b) Liquid

36. Which of the following techniques is useful in organic synthesis for separation, purification and identification of products? 11202053

(a) Sublimation

(b) Filtration

(d) Solvent extraction

37. The technique which is used to separate the insoluble particles from liquid is:

11202054

(a) Crystallization (b) Sublimation

38. Size of filter paper is selected according to: 11202055

(a) Nature of solvent

(b) Quantity of solvent

(d) Size of particles

39. The solid particles left over the filter paper are called: 11202056

(a) Gel (b) Residue

40. The most common Solvent used for solvent extraction is: 11202057

(a) Water (b) Ethanol

41. Iodine is soluble in: 11202058

(a) Water

(b) CCl₄

(d) None of the above

42. Which of the following can react with filter paper? 11202059

(a) Conc. HCl

(b) Ether

(d) None of the above

43. Quantitative determination involves:

11202060

(a) 2 steps (b) 3 steps

44. Which one is not a drying agent? 11202061

(a) CaCl₂ (b) Silica gel

45. The type of analysis in which the relative amounts of the elements are determined is called: 11202062

(a) Qualitative

(b) Quantitative

(d) Volumetric

46. Which of the following techniques can separate organic compound from aqueous solution? 11202063

(a) Distillation

(b) Chromatography

(d) Solvent extraction

47. A fluted filter paper can : 11202064

(a) Increase the rate of filtration

(b) Decrease the rate of filtration

(d) Hinder the filtration process

48. The pattern of inks formed on paper in chromatography is called: 11202065

(a) Chromatophore

(b) Chromatograph

(d) Chromatograph and Chromatogram

49. Which one is not sublimable in laboratory? 11202066

(a) NH₄Cl (b) Benzoic Acid

50. The locating agent used to identify colorless components in chromatography is: 11202067

(a) H₂S (b) Rubeanic Acid

51. Which of the following is purified by sublimation? (Board 2009) 11202068

(a) Naphthalene

(b) Benzoic Acid

(d) All of these

52. Solvent extraction is a process: (Board 2014) 11202069

(a) Exothermic (b) Endothermic

53. Gooch crucible is made of: 11202070

(a) Clay (b) Asbestos (Board 2014)

54. Pure water could be obtained from sea water by: 11202071

(a) Chromatography (b) Filtration

55. When a substance having high vapour pressure is heated at a temperature below its melting point, it undergoes: 11202072

(a) Melting (b) Sublimation

56. What colour, tri-iodide ions give to the water, when they are dissolved in it? 11202073

(a) Brown (b) Blue-brown

57. If Solvent front is 10cm and distance travelled by solute is 1.2 cm, what is its R_f value? 11202074

(a) 0.83 (b) 0.12

58. Solvent used for crystallization should dissolve a large amount of substance at its:

(a) boiling point 11202075

(b) freezing point

(d) transition temperature

59. The components of a mixture which can be separated by filtration are: 11202076

(a) NaCl and CaCl₂

(b) CaCO₃ and NaCl

(d) Blue and green ink

SHORT QUESTIONS

Q.1 How do you justify that qualitative and quantitative analysis are discussed in analytical chemistry? 11202077

Q.2 Media which are used for the filtration should be selected on the basis of precipitates. Why? 11202078

Q.3 What are the various experimental techniques which are used for the purification of substances? 11202079

Q.4 Why the purification of substances by the process of crystallization requires solvents of different nature? 11202080

Q.5 How the crystals of a solid substance can be obtained from the mother liquor?

11202081

Q.6 Why is the desiccator a safe and reliable method for drying the crystals? OR

How are the crystals dried in a desiccator?

(Board 2013) 11202082

Q.7 What is Gooch Crucible? 11202083

Q.8 Concentrated HCl and KMnO₄ solutions cannot be filtered by Gooch Crucible. Give reason. 11202084

Q.9 Give the main characteristics of the solvent used for crystallization. 11202085

Q.10 Mention the major steps involved in the crystallization process. 11202086

Q.11 What is solvent extraction? 11202087

Q.12 What is ether extraction? 11202088

Q.13 Write advantages of sintered glass crucible. 11202089

Q.14 Give main uses of paper chromatography. 11202090

Q.15 How the decolourization of undesirable colours is carried out for freshly prepared crystalline substances?

11202091

Q.16 How the rate of filtration can be increased? 11202092

Q.17 What is fluted filter paper? 11202093

Q.18 How premature crystallization can be avoided? 11202094

Q.19 Which is the best method for drying the crystals? 11202095

Q.20 What is partition or distribution Law? (Board 2014, 2015) 11202096

Q.21 What is adsorption Chromato-graphy? (Board 2015) 11202097

Q.22 What is partition Chromatography? (Board 2015) 11202098

Q.23 What is principle of solvent extraction? 11202099

Q.24 Write the names of the substances, which can sublime? 11202100

Q.25 Write names of different types of paper Chromatography. 11202101

Q.26 Define sublimation with an example. (Board 2014, 2015) 11202102

Q.27 What is R_f Value? Why it has no units? 11202103

Q.28 Define sublimand and sublimate.

(Board 2014) 11202104

Q.29 Why is there a need to crystallize the crude product? (Board 2014) 11202105

Q.30 Differentiate b/w stationary and mobile phase used in chromatography.

GASES

Q.30 What is the different between qualitative and quantitative analysis?

Objectives

MULTIPLE CHOICE QUESTIONS

1. The order of the rate of diffusion of gases (NH₃, SO₂, Cl₂, and CO₂) is:11203055

(a) NH₃ > SO₂ > Cl₂ > CO₂

(b) NH₃ > CO₂ > SO₂ > Cl₂

(d) NH₃ > CO₂ > Cl₂ > SO₂

2. Pressure remaining constant, at which temperature the volume of a gas will become twice of what it is at 0^oC? 11203056

(Board 2014, 15)

(a) 546^oC (b) 200^oC

3. Equal masses of methane and oxygen are mixed in an empty container at 25^oC. The fraction of total pressure exerted by oxygen is: 11203057

(a) (b)

4. Which of the following will have the same number of molecules at STP? 11203058

(a) 280cm³ of CO₂ and 280cm³ of N₂O

(b) 1 dm³ of O₂ and 32g of O₂

(d) 28g of N₂ and 5dm³ of O₂

5. The relation used to calculate root mean square velocity of gases is: 11203059

(a) C_r.m.s =

(b) C_r.m.s =

(d) C_r.m.s =

6. If absolute temperature of a gas is doubled and the pressure is reduced to one-half, the volume of the gas will:

(a) Remain unchanged (Board 2013)11203060

(b) increase four times

(d) be doubled

7. How should the conditions be changed to prevent the volume of a given gas from expanding when its mass is increased? 11203061

(a) Temperature is lowered and pressure is increased

(b) Temperature is increased and pressure is lowered

(d) Temperature and pressure both are increased

8. The molar volume of CO₂ is maximum at: (Board 2014, 15) 11203062

(a) STP (0^oC and 1 atm)

(b) 127^oC and 1 atm

(d) 273^oC and 2 atm

9. Gases deviate from ideal behaviour at high pressure. Which of the following is correct for non-ideality? 11203063

(a) At high pressure the gas molecules move in one direction

(b) At high pressure the collisions between the gas molecules are increased manifold

(d) At high pressure, the intermolecular attractions become significant

10. The deviation of a gas from ideal behaviour is maximum at: 11203064

(a) –10^oC and 5.0 atm.

(b) –10^oC and 2.0 atm.

(d) 0^oC and 2 atm.

11. A real gas obeying Van der Waal’s equation will resemble ideal gas if:

(a) both ‘a’ and ‘b’ are large 11203065

(b) both ‘a’ and ‘b’ are small

(d) ‘a’ is large and ‘b’ is small

12. Which one of the following units is SI unit of pressure? 11203066

(a) mm of Hg

(b) Torr

(d) Pounds per square inch

13. Which of the following is false in case of gases? 11203067

(a) They diffuse easily

(b) They have mass

(d) They do not mix well

14. Normal temperature and pressure (STP) of gas refers to: 11203068

(a) 273 K and 76mm of Hg

(b) 273^oC and 760mm of Hg

(d) 273^oC and 76mm of Hg

15. 1 dm³ of O₂ at STP contains : 11203069

(a) 6.02 ´ 10²³ molecules

(b) molecules

(d) 0.602 ´ 10²² molecules

16. A real gas is one which: 11203070

(a) shows deviation from gas laws

(b) has significant volume of the molecules

(d) All of the above

17. A graph is plotted between two variables that is pressure and volume at constant temperature and fixed number of moles of the gas, the graph is called:

(a) Isotherm (b) Isobar 11203071

18. A graph obtained from Boyle’s law by plotting P versus V at constant T is :

(a) A straight line 11203072

(b) A curve with maximum

(d) A parabolic curve called isotherm

19. The isotherm of CO₂ change the position in the graph by changing the temperature in such a way that: 11203073

(a) The curve come closer to the axis by increasing pressure

(b) The curve go away from the axis by increasing the temperature

(d) The curve becomes discontinuous at 31.5^oC

20. A graph between pressure and inverse of volume for a given mass of a gas at constant temperature is: 11203074

(a) Straight line parallel to x-axis

(b) Straight line parallel to y-axis

(d) The curve showing the maximum

21. The natural plasma is extremely hot and it has minimum temperature: 11203075

(a) 2000 ⁰C (b) 20,000 ⁰C

22. If a graph is plotted between temperature on x-axis and volume of the one mole of the gas on y-axis at constant pressure then a straight line is obtained which cuts the temperature axis at: 11203076

(a) 0^oC (b) -273.16^oC

23. Which of the following pairs of gases contain same number of molecules? 11203077

(a) 11 grams of CO₂ and 14 grams of N₂

(b) 11 grams of CO₂ and 7 grams of N₂

(d) 44 grams of CO₂ and 44 grams of N₂

24. If the temperature and pressure of 2dm³ of CO₂ are doubled, then volume of CO₂ would become: 11203078

(a) 8 dm³ (b) 4 dm³

25. The gas law giving relationship between volume and temperature of gas is:11203079

(a) Dalton’s law (b) Charle’s law

26. If both temperature and volume of a gas are doubled, the pressure: 11203080

(a) Cannot be predicted

(b) Is reduced to 1/2

(d) Is doubled

27. The value of the general gas constant R in SI units is: 11203081

(a) 8.3143 kJK^-¹ mole^-¹

(b) 8.3143 JK^-¹ mole^-¹

(d) 62.4 dm³ torr K^-¹ mole^-¹

28. The product PV of a gas is a unit of:

(a) Force (b) Entropy 11203082

29. CH₄ gas is maintained at 0^oC and 1 atm pressure. Its density is 0.714 g/dm³. What is its density at 0.5 atm and 0^oC?

(a) 0.714 g dm^-³ (b) 1.428 g dm^-³11203083

30. In a closed vessel of 1000 cm³, H₂ gas is heated from 27^oC to 127^oC. Which statement is not correct? 11203084

(a) The rate of collision increases

(b) Pressure of gas increases

(d) The number of moles of a gas increases

31. If 10g of a gas at one atmospheric pressure is cooled from 273^oC to 0^oC at constant volume, its pressure would become: 11203085

(a) 1 atm (b) atm

32. One dm³ of H₂ and one dm³ of O₂ have same number of molecules at STP, their respective masses are: 11203086

(a) 0.1g and 4.3g (b) 2 g and 32g

33. Which pair of gases do not obey Dalton’s law of partial pressure? 11203087

(a) H₂ and O₂ (b) N₂ and O₂

34. Gas equation is derived by combining:

(a) Avogadro’s and Charles’s law11203088

(b) Boyle’s law and Charles’s law

(d) Avogadro’s, Boyle’s and Charles’s law

35. At 100^oC a gas has 1 atm pressure and 10 dm³ volume. Its volume at STP would be: 11203089

(a) 10 dm³

(b) Less than 10 dm³

(d) Cannot be predicted

36. A gas was compressed to half of its volume at 303K. To what temperature it should be heated so that its volume becomes double? 11203090

(a) 303 K (b) 330 K

37. A gas occupies a volume of 2 dm³ at 27^oC and 1 atm pressure. The expression for its volume at STP is: 11203091

(a) ´ 300 (b) ´ 300

38. A certain mass of a gas occupies a volume of 2 dm³ at STP. Keeping the pressure constant, at what temperature would the gas occupy a volume of 4 dm³? 11203093

(a) 300^oC (b) 546 K

39. The formula for density of a gas at a given temperature and pressure is:

(a) d = (b) d = 11203094

40. The density of CH₄ at 2 atm pressure at 27^oC is: 11203095

(a) 26 g dm³ (b) 0.26 g dm^-³

41. The volume of NH₃ obtained by the combination of 10 cm³ of N₂ and 30 cm³ of H₂ is: 11203096

(a) 20 cm³ (b) 30 cm³

42. Hydrogen gas is prepared in the laboratory and is collected over water. The pressure of the wet gas is 745 torr. The aqueous tension is 24 torr. The pressure of dry hydrogen is: 11203097

(a) 766 torr (b) 745 torr

43. Two gases H₂ and O₂ are enclosed in a porous vessel. Which statement is correct about comparative effusion rate? 11203098

(a) O₂ effuses four times faster than H₂

(b) H₂ effuses four times faster than O₂

(d) They have equal rate of effusion

44. Due to high temperature the gas can be ionized and free electrons are generated, giving us mixture of gas molecules, ions and free electrons. The collection is called: 11203099

(a) Mixture of gas

(b) Gaseous phase substance

(d) Plasma

45. The deviation of gas from ideal behaviour is due to which one of the following reasons? 11203100

(a) No forces of attraction among molecules of gases

(b) Negligible volume of gas molecules

(d) Low pressure and high temperature

46. At which of the following temperature, SO₂ gas behaves comparatively non – ideal? 11203101

(a) 327K (b) 400K

47. The free expansion of the gas from high pressure towards the low pressure causes:

(a) Increase of temperature 11203102

(b) Decrease of temperature

(d) Decrease of velocities of gas molecules

48. The molecules of a gas show more deviation from ideal behaviour at low temperature because: 11203103

(a) Attractive forces dominate at low temperature

(b) Kinetic energies are increased

(d) Densities of the gases increase

49. The liquefaction of a real gas can be carried out if: 11203104

(a) The temperature is more than critical temperature and pressure is 1000 atm.

(b) The temperature is below the critical temperature and pressure is very high.

(d) Without caring for the value of critical volume at critical stage.

50. Nitrogen and ethene have same rates of diffusion through a porous container at a fixed temperature. This is due to the reason that: 11203105

(a) Both gases have multiple bonds in them

(b) Both gases are non-polar

(d) Their molar masses are same

51. We have two gases with same molar masses at fixed temperature. Which statement is true about these two gases?

(a) Their atomicities are same 11203106

(b) Their Kinetic energies are different

(d) They have the same rates of diffusion

52. Rate of diffusion of CO and N₂ are same at room temperature due to the reason that: 11203107

(a) Both are diatomic molecules

(b) Both have same multiple bonds in them

(d) Both have same molar masses

53. The rate of diffusion of hydrogen is three times than that of an unknown gas at same temperature and pressure, then the molar mass of unknown gas is: 11203108

(a) 18 (b) 16

54. Partial pressure of oxygen in the lungs is: 11203109

(a) 760 torr (b) 116 torr

55. The gases exert the pressure on the walls of the container. This is due to: 11203110

(a) Collisions of molecules among themselves

(b) Change of direction of molecules

(d) Collisions on the walls of the vessel

56. Behaviour of gas molecules can be understood from the kinetic equation of gases. This equation was given by clausius and equation is: 11203111

(a) d = (b) PV = nRT

57. Air is a mixture of gases. The molecules of the air are not settled down due to:

(a) Different molar masses 11203112

(b) Non-polar nature of gases

(d) Collision of gas molecules

58. The highest temperature above which a gas cannot be liquefied, no matter how much the pressure is applied is known as: 11203113

(a) Boiling temperature

(b) Consulate temperature

(d) Critical temperature

59. Neon has low critical temperature and pressure as compared to other gases. The most probable reason is that: 11203114

(a) Its octet is complete

(b) It is a mono atomic gas

(d) It has least forces of attraction

60. The critical temperature of Ar gas is low as compared to NH₃ and SO₂ due to the reason that: 11203115

(a) Ar is mono atomic gas

(b) It has a small size

(d) It has four lone pairs in it

61. Volume occupied by one mole of a gas at STP is called: 11203116

(a) Normal volume

(b) Molar volume

(d) None of the above

62. All the gases are liquefied at: 11203117

(a) 373 K

(b) Zero K

(d) 546 K

63. Which of the following gases in solid state is called dry ice? 11203118

(a) H₂ (b) CO₂

64. Elastic collisions involve : 11203119

(a) Loss of energy

(b) Gain of energy

(d) No relation with energy

65. Which of the following relationships is incorrect? 11203120

(a) K = C^o + 273 (b) C^o= [F^o-32]

66. 4 gm of H₂ gas at STP occupies volume of: 11203121

(a) 60lit (b) 44.8lit

67. 760 torr pressure is equal to: 11203122

(a) 10.1325

(b) 1.01325

(d) 110.325 kilo pascal

68. Pascal in terms of SI units is equal to:

(a) 1Nm^-2 (b) 2Nm^-211203123

69. The respiration process in the animals depends on the difference of: 11203124

(a) Osmotic pressure

(b) Vapour pressure

(d) Atmospheric pressure

70. According to Kinetic Molecular Theory, Kinetic energy of molecules increases when they are: 11203125

(a) Mixed with other molecules at low temperature

(b) Frozen to solid

(d) Melted from solid to liquid

71. Body temperature of a normal person is: 11203126

(a) 97.6°F (b) 96.8°F

72. Mathematical expression of compressibility factor is: 11203127

(a) (b)

73. Inert gases are mono atomic: 11203128

(a) Molecules (b) Atoms

74. The distribution of energies among the molecules of gases was studied by:11203129

(a) Boltzmann (b) Newton

75. Kinetic theory was put forward by:

(a) Bernoulli (b) Coulomb 11203130

76. Plasma consists of mixture of neutral particles, positive ions and: 11203131

(a) Protons (b) Electrons

77. Critical temperature of NH₃ is greater than CO₂ due to its: 11203132

(a) Lesser polarity (b) Greater polarity

78. More ideal gas at room temperature is:

(a) CO₂ (b) NH₃11203133

79. If 2 moles of an ideal gas at 46K occupy a volume of 4.48 dm³, the pressure must be 11203134

(a) 1 atm (b) 2 atm

80. Which gas will diffuse more rapidly? (Board 2009, 2014) 11203135

(a) CO₂ (b) NH₃

81. Escape out of gas molecules one by one through tiny hole is: (Board 2014) 11203136

(a) Diffusion (b) Effusion

82. If “a” and “b” are zero for certain gas, then gas is: (Board 2014) 11203137

(a) Ideal

(b) Real

(d) May be any diatomic gas

83. Number of molecules in one dm³ of water is close to: (Board 2013) 11203138

(a) (b)

84. A container contains three gases; a,b,c in the molar ratio 2:3:5 respectively. If the total pressure is 500 torr then the partial pressure of the component ‘a’ 11203139

(a) 100 torr (b) 200 torr

85. Which of the following gases has lowest critical temperature (Tc)? 11203140

(a) CO₂ (b) O₂

86. Critical temperature of a gas depends on which of the following factors: 11203141

(a) size of molecules

(b) shape of molecule

(d) all of these

87. The most probable velocity is equal to: 11203142

(a) (b)

88. For an ideal gas, the compressibility factor is equal to: 11203143

(a) 0.5 (b) 1

SHORT QUESTIONS

Q1. Explain the following facts: 11203144

(a) The plot of PV versus P is a straight line at constant temperature and with a fixed number of moles of an ideal gas.

(b) The straight line in (a) is parallel to x-axis and goes away from pressure axis at higher pressure.

(d) Pressure of NH₃ gas at given conditions (say 20 atm pressure and room temperature) is less than when calculated by Van der Waal’s equation than that calculated by general gas equation.

(e) Water vapours do not behave ideally at 273K. (Board 2014)

(f) SO₂ is comparatively non-ideal at 273K but behaves ideally at 327K.

Q2. How does the position of isotherm change, when the temperature of the given gas is changed? 11203145

Q3. Why is the Boyle’s law applicable only to the ideal gases? 11203146

Q4. How is the product of pressure and volume at constant temperature and number of moles, a constant quantity? 11203147

Q5. When a gas obeys the Boyle’s law, the isotherm for the gas can be plotted. Justify it.

(Board 2014) 11203148

Q6. Charle’s law is not obeyed when the temperature is measured on celsius scale. Justify it. 11203149

Q.7 What is absolute zero? 11203150

Q.8 The volume of any gas does not become zero at –273.16^oC. Give reason.

11203151

Q.9 What factors affect the critical temperature of gases? 11203152

Q.10 How do you explain that –273^oC is a theoretical temperature and is not attainable? 11203153

Q.11 What is effect of temperature and pressure on the density of the gas? 11203154

Q.12 How the density of an ideal gas doubles by doubling the pressure or decreasing the temperature on Kelvin scale by 1/2? 11203155

Q.13 Why 99% of the matter in the universe is in the plasma state? 11203156

Q.14 Prove that the partial pressure of any gas is directly proportional to the mole fraction of that gas. OR 11203157

Prove that P_i = P_t X_i. (Board 2014)

Q.15 How do you say that the pressure of the dry gas is equal to difference of total pressure and aqueous tension of H₂O? 11203158

Q.16 The rate of diffusion of O₂ is 4 times less than that of H₂ at given temperature and pressure. Justify it. 11203159

Q.17 Why is the volume correction done by Van der Waal’s equation? 11203160

V_free = V_vessel – b

Q.18 In Joule-Thomson effect, sudden expansion of the gas molecules need energy. Why? 11203161

Q.19 Why is the excluded volume less than molar volume of the gas? 11203162

Q.20 The Van der Waal’s constant ‘a’ and ‘b’ are quantitative measurements for the non-ideality of the gas. Justify. 11203163

Q.21 The amount of pressure, which is decreased due to the forces of attraction, is given by where ‘a’ is the Van-der Waal’s constant and V is the volume of the vessel. Explain. 11203164

Q.22 How is Dalton’s law useful in determining pressure of a gas, collected over water? 11203165

Q.23 Differentiate between Diffusion and Effusion of gases. 11203166

Q.24 Why real gases deviate more at high pressure and low temperature? 11203167

Q.25 Which is the fourth state of matter? How it can be obtained? 11203168

Q.26 Lighter gas can diffuse more rapidly than heavier gas. Why? 11203169

Q.27 Rate of diffusion of NH₃ is greater than HCl, why? 11203170

Q.28 Why gases do not settle at room temperature? 11203171

Q.29 High pressure and low temperature makes the gases non-ideal. Why? 11203172

Q.30 Why at higher altitudes, one feels uncomfortable breathing or pilot cabins are pressurized? 11203173

Q.31 Deep sea Divers do not use normal air in breathing, why? (Board 2010) 11203174

Q.32 Differentiate between ideal gas and non-ideal gas. 11203175

Q.33 Value of excluded volume ‘b’ is greater for SO₂ and less for H₂, why? 11203176

Q.34 Why critical temperature of H₂ is low while that of H₂O is high? 11203177

Q.35 What is critical temperature and critical pressure? 11203178

Q.36 Isotherm is parabolic, when pressure is applied over a given mass of a gas, why?

11203179

Q.37 How the critical temperature is essential criteria to be considered for the liquefaction of gases? 11203180

Q.38 How Lind’s method is application of Joule Thomson effect? 11203181

Q.39 Describe the reasons for deviation of gases from ideality. 11203182

Q.40 Where is plasma found? 11203183

Q.41 Define Pressure. Give its S.I units.

11203184

Q.42 What are the faulty points in the Kinetic molecular theory of gases?

(Board 2009, 2015) 11203185

Q.43 State Joule Thomson effect. Write its applications. 11203186

Q.44 Define Avogadro’s law.

(Board 2014) 11203187

Q.45 Explain that the process of respiration obey’s Dalton’s law of partial pressure. (Board 2013) 11203188

Q.46 Calculate the fraction of total pressure exerted by oxygen when equal masses of CH₄ and O₂ are mixed in an empty container at 25^oC.(Board 2014) 11203189

Q.47 Calculate the value of general gas constant ‘R’ in S.I units. (Board 2014) 11203190

Q.48 Write down the values of atmospheric pressure in four different units. (Board 2015) 11203191

Q.49 Write down any two applications of plasma. (Board 2015) 11203192

Q.50 Explain Boyle’s law with the help of KMT. (Board 2013) 11203193

Q.51 What is mean square velocity and root mean square velocity? 11203194

LIQUIDS (Section-I)

Objectives

MULTIPLE CHOICE QUESTIONS

1. Gases can be converted into liquids by:

(a) Lowering the temperature only11204039

(b) Increasing the pressure only

(d) Increasing the pressure and lowering the temperature below critical point

2. London dispersion forces are the only forces present among the: (Board 2014) 11204040

(a) Molecules of water in liquid state

(b) Atoms of helium in gaseous state at high temperature

(c)Molecules of solid iodine

(d) Molecules of hydrogen chloride gas

3. Acetone and chloroform are soluble in each other due to: 11204041

(a) Intermolecular hydrogen bonding

(b) Dipole-dipole interaction

(d) All of the above

4. NH₃ shows a maximum boiling point among the hydrides of VA group elements due to: 11204042

(a) Very small size of nitrogen

(b) Lone pair of electrons present in nitrogen

(d) Pyramidal structure of NH₃

5. When water freezes at 0^oC, its density decreases due to: 11204043

(a) Cubic structure of ice

(b) Empty spaces present in the structure of ice

(d) Change of bond angles

6. The repulsions of electronic clouds of molecules are responsible for the attractive forces among the molecules. These forces are: 11204044

(a) Dipole-induced dipole forces

(b) Ion-dipole forces

(d) Dipole-dipole forces

7. The forces which are present between the ions and the polar molecules of the solvent are: 11204045

(a) Dipole-induced dipole forces

(b) Dipole-dipole forces

(d) London dispersion forces

8. The boiling point of higher alkanes are greater than those of lower alkanes due to reason that: 11204046

(a) Higher alkanes have greater number of atoms

(b) The polarizabilities of higher alkanes are greater

(d) Higher alkanes have zig-zag structure

9. The vapour pressure of water at a particular temperature is less than an alkane at same temperature. This is due to the reason that: 11204047

(a) Alkanes have no hydrogen bonding

(b) Water is a non-polar molecule

(d) London forces are present in water

10. Hydrogen bonding is extensively present in proteins between: 11204048

(a) Nitrogen and hydrogen atom

(b) Oxygen and hydrogen atom

(d) All of the above

11. Ice floats on water because: 11204049

(a) The hydrogen bonding in ice is stronger than that in water

(b) Empty spaces are left in ice

(d) The bond length of the oxygen and hydrogen bond is different in water and ice

12. The boiling point of water is greater than that of HF. This is due to: 11204050

(a) Water is more polar than HF

(b) HF has a zig-zag structure after making hydrogen bonding

(d) Water has angular structure

13. Among the hydrides of VA group NH₃ gas, PH₃ AsH₃ etc. NH₃ shows the maximum boiling point among the hydrides of the group VA due to reason that: 11204051

(a) The lone pair of nitrogen has different character than the other lone pairs on hydrides

(b) Nitrogen is a very small sized atom

(d) The electronegativity of nitrogen is maximum

14. The polarizability of elements mostly increases down the group due to the reason that: 11204052

(a) The atomic number increases

(b) Number of protons increases

(d) The behaviour of the elements remains the same

15. The long chains of amino acids are coiled about one another into a spiral by:11204053

(a) Ionic bond

(b) Van der Waal’s forces

(d) Overlapping of orbitals

16. The hydride of oxygen (H₂O) is liquid at room temperature but the hydride of sulphur (H₂S) is a gas. This is due to:

(a) Greater bond angle of water than H₂S

(b) Greater bond lengths in H₂S than H₂O

(d) Acidic character of H₂S 11204054

17. The weakest intermolecular forces present in a liquid may be: 11204055

(a) Dipole-induced dipole forces

(b) Dipole-dipole forces

(d) Electrostatic forces between ions in ionic solid

18. The nature of the attractive forces in acetone and chloroform are: 11204056

(a) Dipole-induced dipole forces

(b) Dipole-dipole forces

(d) Hydrogen bonding

19. Strong dipole-dipole forces among the liquid molecules are responsible for:

(a) Very high heat of vaporization11204057

(b) Very low heat of vaporization

(d) Negligible forces

20. Dipole-dipole interactions are present between the: 11204058

(a) Atoms of the He gas

(b) Molecules of CH₄

(d) Molecules of NH₃

21. Polarizability is responsible for the intermolecular forces and it: 11204059

(a) Increases down the group

(b) Decreases down the group

(d) Increases along a period

22. The boiling points of the halogens:

(a) Increases down the group 11204060

(b) Decreases down the group

(d) Cannot be predicted

23. The boiling point of H₂O is 100^oC while that of C₂H₅OH is 78.5^oC. The reason is that: 11204061

(a) H₂O molecules are small-sized

(b) The bond angles at oxygen atom are different

(d) The number of H-bonds are greater in H₂O than C₂H₅OH

24. Saturated hydrocarbons having carbon atoms more than 20 in a molecule are solid due to: 11204062

(a) Higher densities

(b) Higher molar masses

(d) All of the above

25. Halogens form halogen acids. HF is the weakest among all of them. This is due to the reason that: 11204063

(a) Fluorine is a very small sized atom

(b) Fluorine is highly electronegative atom

(d) The polarity of HF bond is less

26. When two ice cubes are pressed over each other they unite to form one cube. This is due to: 11204064

(a) Dipole-dipole attraction

(b) Covalent attraction

(d) Hydrogen bond formation

27. H-bonding is maximum in: 11204065

(a) Ethanol (b) Benzene

28. Which of the following can form H-Bonds? 11204066

(a) NH₃ (b) C₂H₆

29. CO₂ and SO₂ are both triatomic molecules, but heat of vaporization of SO₂ is greater than that of CO₂. This is due to: 11204067

(a) Greater electronegative character of sulphur

(b)Greater size of SO₂ molecule

(d)SO₂ is more acidic in nature than CO₂

30. To cook the food at high mountain is difficult as compared to at sea level. The reason is that: 11204068

(a) The temperature at the top of the mountain is low

(b) The density of water decreases at the mountain

(d) The hydrogen bonding of water changes with the change in height

31. Glycerin is a polar compound. It boils at 290^oC under one atmospheric pressure. It should be distilled under reduced pressure, due to the reason that:11204069

(a) There are strong intermolecular forces between molecules of glycerin

(b) It decomposes at 290^oC

(d) The reduced pressure decreases the boiling point of liquid

32. Liquid evaporates at every temperature. When the temperature becomes constant for a liquid then: 11204070

(a) Rate of evaporation is greater than rate of condensation

(b) The rate of condensation is greater than the rate of evaporation

(d) It depends upon the nature of the liquid

33. The distillation of a solution under reduced pressure is called: 11204071

(a) Fractional distillation

(b) Destructive distillation

(d)Vacuum distillation

34. Water may boil at 120^oC when external pressure is: (Board 2014) 11204072

(a) 369 torr

(b) 700 torr

(d) 1489 torr

35. The evaporation of a liquid causes:

(a) Increase in the Kinetic energy of the molecule 11204073

(b) Maintenance of a constant temperature

(d) Heating

36. The cooking time in the pressure cooker is reduced because: 11204074

(a) The vapor pressure of water decreases

(b) The external pressure on the surface of water decreases

(d) Heat is uniformly distributed inside the pressure cooker

37. Which one of the following liquids has the lowest vapour pressure at 32^oC?

(a) Ether (b) Chloroform11204075

38. Which of the following noble gases have high polarizability? 11204076

(a) Ne (b) Kr

39. The boiling point of any liquid remains constant and heat is to be supplied continuously to keep it boiling. The reason is that: 11204077

(a) Extra heat is sprout in the air

(b) High energy molecules continuously leave the liquid

(d) The vapour pressure of the boiling liquid becomes greater than the external pressure.

40. At which place the vapour pressure of H₂O at its boiling point is less than 760 torr? 11204078

(a) At sea level (b)Lower than sea level

41. Vapour pressure of liquid is measured when liquid and the vapours are in equilibrium. It means that: 11204079

(a) Liquid and vapours have same value of Kinetic energies

(b) Liquid and vapours have the same heat contents

(d) Rate of evaporation and condensation are different

42. The value of the vapour pressure of water at its boiling point at Karachi and Murree is: 11204080

(a) Same

(b) Different

(d) Depends upon the environmental conditions in both cities

43. When the distillation is performed under the reduced pressure, then process is

(a) Slow 11204081

(b) Rapid

(d) Highly dangerous

44. In order to maintain the boiling point of water at 110^oC,the external pressure should be: (Board 2015) 11204082

(a) Any value of pressure

(b) 765torr

(d) Between 760 torr& 1200 torr

45. The Hydrides of group IV have least boiling points as compared to those of group V, VI and VII A hydrides due to the reason that: 11204083

(a) They form four covalent bonds with hydrogen

(b) The elements are least electronegative

(d) The sizes of these atoms are small

46. In cold countries the water of lakes and rivers freeze earlier than seawater because. 11204084

(a) They contain more dissolved salts and Impurities

(b) They contain less dissolved salts and Impurities

(d) They contain less dissolved oxygen and Impurities

47. Which of the following does not control the rate of evaporation? 11204085

(a) Surface area

(b) Temperature

(d) Amount of liquid

48. The distillation of liquids is performed under reduced pressure due to the reason that 11204086

(a) It is a slow process

(b) The evaporation rate becomes slow

(d) Evaporation rate can be controlled

49. The cooking of meat at Murree is difficult than at sea level because11204087

(a) Density of H₂O is high at mountain

(b) Mass of meat increases at mountain

(d) Hydrogen bonding in water becomes stronger

50. The Relative Molecular Mass of Carbon tetrachloride is 154 while that of ethanol is 46. But the boiling point of Carbon tetrachloride is 76.5°C and B.P of ethanol is 78.26°C. The reason of high boiling point of ethanol is that: 11204088

(a) Its molecular mass is less than Carbon tetra Chloride

(b) Its density is higher than Carbon tetra Chloride

(d) It has hydrogen bonding

51. Which of the following has strongest hydrogen bonding? 11204089

(a) CH₄ (b) NH₃

52. Both Cl₂& I₂ belong to VII group, but I₂ is solid & Cl₂ is gas because 11204090

(a) Polarizability of I₂ is less than Cl₂

(b) Polarizability of I₂ is more than Cl₂

(d) The covalent radius of I₂ is less than Cl₂

53. Which of the following has lowest boiling point? 11204091

(a) HF (b) HBr

54. The strongest intermolecular force is:

(a) Ion – Dipole forces 11204092

(b) Electrostatic forces between ions

(d) Dipole – Induced dipole forces

55. Van-der Waal’s forces are weak intermolecular forces, they include:

(a) Dipole – Dipole forces only 11204093

(b) Ion – Dipole forces only

(d) All of the above

56. The London forces become stronger if:

(a) Size of atom is smaller 11204094

(b) Density of molecules is large

(d) Moleculesarehomo-atomic

57. The magnitude of Vapour pressure:

(a) Depends upon the amount of liquid

(b) Does not depend upon the amount of liquid 11204095

(d) Depends upon the shape of container

58. The liquid pair, which is immiscible is: (Board 2009) 11204096

(a) Acetone & H₂O (b)Phenol & H₂O

59. Which of the following is not a type of liquid crystal? 11204097

(a)Enteric (b) Cholesteric

(c)Smectic (d) Nematic

60. The lowest vapour pressure is exerted by: 11204098

(a) Ethanol (b) Water

61. The boiling point of water at the top of mount Everest is: (Board 2015) 11204099

(a) 59^oC (b) 69^oC

62. Molar heat of vapourization of water is: 11204100

(a) 40.7 kcal mol^–1 (b) 40.7kJ mol^–1

63. An organic compound cholesteryl benzoate becomes a clear liquid at: 11204101

(a) 145^oC (b) 323^oC

64. Diameter of double helical structure of DNA is: 11204102

(a) (b)

65. Which of the following forces exist in noble gases? 11204103

(a) Dipole-dipole forces

(b) Dipole-induced Dipole forces

(d) Hydrogen bonding

66. What is the boiling point of water at 23.7 torr? 11204104

(a) 25^oC (b) 69^oC

SHORT QUESTIONS

Q1. Explain the following with reasons: 11204105

In the hydrogen bonding structure of HF, indicate the stronger bond, the shorter covalent bond or the longer hydrogen bond between different molecules.

Q2. In a very cold winter, fish in garden ponds owe their lives to hydrogen bonding.

11204106

Q3.Water and ethanol can mix easily and in all proportions. 11204107

Q4. The origin of intermolecular forces in water. 11204108

Q5.Briefly consider some of the effects on our lives, if water has only very weak hydrogen bonding present among its molecules. 11204109

Q6.All gases have a characteristic critical temperature. Above the critical temperature it is impossible to liquify gas. The critical temperature of CO₂ and CH₄is31.14^oC and –81.9^oC respectively. Which gas has the stronger intermolecular forces?11204110

Q. 7.Explain the following with reasons:

(Board 2010,14)

Evaporation causes cooling.11204111

Q8. Evaporation takes place at all temperatures. (Board2014)11204112

Q9. Boiling needs a constant heat supply. (Board 2010) 11204113

Q10. Earthen-ware vessels keep water cool.

11204114

Q11.One feels sense of cooling under fan after bath. 11204115

Q12.Dynamic equilibrium is established during evaporation of a liquid in a close vessel at constant temperature. 11204116

Q13.The boiling point of water is different at Murree Hills and at Mount Everest.

11204117

Q14.Vacuum distillation can be used to avoid decomposition of sensitive liquids.

11204118

Q15.Heat of sublimation of a substance is greater than that of heat of vaporization.

11204119

Q16.Heat of sublimation of iodine is very high.(Board 2014)11204120

Q17.How dipole – dipole forces effect the thermodynamic properties? 11204121

Q18.What kind of attractive forces are created between CO₂ molecules when it changes to dry ice? 11204122

Q19.Why the overall boiling points of hydrides of group IV to VII increases down the group?

11204123

Q20.Why dirt do not resettle on cloth when it is washed with soap?11204124

Q21.In Manometric method, why liquid in the flask is repeatedly boiled & frozen?

11204125

Q22.At 0°C diethyl ether show more vapour pressure than ethyl alcohol. Why?

11204126

Q23.What is meant by energetics of phase changes? 11204127

Q24.How liquid crystals display numbers on digital watches and calculator screens?

11204128

Q25.Why DNA form double Helix structure? 11204129

Q26.What are the commercial applications of liquid crystals? 11204130

Q27.Water is liquid at room temperature but H₂S is a gas. Give reason. 11204131

Q28.How the rate of evaporation depends on the surface area? 11204132

Q29.Justify that HF is a weak acid than HCl. 11204133

Q30.Why ice floats on the surface of liquid H₂O? Explain. 11204134

Q31. Why the boiling points of noble gases increase down the group? (Board 2014)11204135

Q32. Why is boiling point of H₂O greater than HF?(Board 2014) 11204136

Q33. Gasoline evaporates much faster than water. Give reason. (Board 2014)11204137

Q34. Why the boiling point of higher alkanes is greater than that of lower alkanes? 11204138

Q35.What are intermolecular forces? Give their types. 11204139

Q36. What is the basis of cleansing action of soaps and detergents? 11204140

Q37. Define vapour pressure. On what factors does it depend? 11204141

Q38. Name the factors affecting London dispersion forces. 11204142

Q39. Why different liquids evaporate at different rates even at the same temperature? (Board 2015)11204143

Q40. How the liquid crystals help in the detection of the blockage in veins and arteries? (Board 2015) 11204119

Q41.What are dipole-dipole forces of attraction? Explain with an example.

(Board 2015) 11204119

SOLIDS (Section-II)

Objective

MULTIPLE CHOICE QUESTIONS

1. Ionic solids are characterized by:11204120

(a) Low melting points

(b) Good conductivity in solid state

(d) Solubility in polar solvents

2. Amorphous solids: 11204121

(a) Have sharp melting points

(b) Undergo clean cleavage when cut with knife

(d) Can possess small regions of orderly arrangement of atoms

3. The molecules of CO₂ in dry ice form the: (Board 2014, 2015) 11204122

(a) Ionic crystals

(b) Covalent crystals

(d) Metallic crystals

4. Which of the following is a pseudo solid? (Board 2014, 2015) 11204123

(a) CaF₂ (b) Glass

5. Diamond is a bad conductor of electricity because: 11204124

(a) It has tight structure

(b) It has high density

(d) None of the above

6. Which of the following substances is amorphous in nature? 11204125

(a) Sugar (b) Graphite

7. Plastics are amorphous solids and:

(a) Have sharp melting point 11204126

(b) Undergo clean cleavage when cut with knife

(d) Possess orderly arrangement over long distances

8. Which of the following substances is not amorphous? 11204127

(a) Polymer (b) Rubber

9. Crystals can be classified into: 11204128

(a) 7 crystal systems(b)4 crystal systems

10. Which among the following will show anisotropy? 11204129

(a) Wood (b) Paper

11. In a diamond crystal, the hybridization of carbon is: 11204130

(a) sp (b) sp²

(d) depends upon the purpose for which diamond is being used

12. Hardness of diamond is attributed to the: 11204131

(a) Strength of the ionic bond in the structure

(b) Three-dimensional network of covalent bonds

(d) Tetrahedral geometry of each carbon

13. In diamond, the carbon atoms are arranged in: 11204132

(a) Tetrahedral manner

(b) Hexagonal manner

(d) Octahedral manner

14. Which one of the following substances shows anisotropic behaviour in electrical conductivity? 11204133

(a) Diamond (b) Ice

15. Which crystal system is found in AgNO₃? 11204134

(a) Orthorhombic and rhombohedral

(b) Cubic and orthorhombic

(d) Monoclinic and hexagonal

16. Isomorphism is present in K₂SO₄ and K₂CrO₄. These two compounds:11204135

(a) Show same physical and chemical properties

(b) 100% equal ionic character

(d) The shapes of both SO and CrOion are tetrahedral and Cations are common

17. The existence of an element in more than one form is called: 11204136

(a) Isomorphism (b) Allotropy

18. Variation of a physical property in a crystal in different directions is called:

(a)Absence of symmetry 11204137

(b)Isomorphism

(d) Polymorphism

19. The crystals of Na₂SO₄ and Na₂SeO₄ should be: 11204138

(a) Isomorphs of each other

(b) Polymorphs of each other

(d) Isomorphs and Allotropes of each other

20. Most crystals show good cleavage because their atoms, ions and molecules are: 11204139

(a) Arranged in planes

(b) Weakly bonded together

(d) Strongly bonded together

21. The brittleness of the ionic compound is due to the reason that: 11204140

(a) Ions are present in the crystal

(b) The sizes of the ions are unequal

(d) The negatively charged and positively charged ions are arranged in alternate positions in layers and these positions are disturbed by stress

22. Ionic solids don’t conduct the electrical current because: 11204141

(a) Ions do not have translatory motion

(b) Free electrons are not present

(d) Strong covalent bonds are present in their structure

23. The number of Cl^- ions per unit cell of NaCl are: 11204142

(a) 6 (b) 4

24. The Cl^-ions present at the corner of the unit cell in NaCl crystal, contributes:

(a) th (b) th 11204143

25. NaCl has face centered cubic structure. The Na⁺ ion at the faces of the unit cell is shared by: 11204144

(a) 2 – unit cells (b) 4 – unit cells

26. Which one of the following statements is not true about the metallic solids?11204145

(a) Metals are good conductor of heat and electricity

(b) Metals are not malleable and ductile

(d) Metals have free electrons

27. In most of the cases the molecular crystals are: 11204146

(a) Very soft

(b)Moderately soft

(d)Sufficiently hard

28. Which one of the following is an ionic solid? 11204147

(a) Fe (b) Diamond

29. LiF is a crystalline substance and has:

(a) Ionic crystals 11204148

(b)Metallic crystals

(d) Molecular crystals

30. Some of the crystals are good conductors of heat and electricity they may be:11204149

(a) Ionic in nature

(b) Covalent in nature

(d) Of molecular nature

31. Ionic solids are characterized by which of the following properties? 11204150

(a) Moderately low pressure

(b) High pressure

(d) Solubility in polar solvents

32. SiO₂ is an example of: 11204151

(a) Metallic crystals

(b) Covalent crystals

(d) A crystal whose crystal structure depends upon the temperature

33. The forces of attraction among the H₂ molecules in solid hydrogen are:11204152

(a) Hydrogen bonds

(b) Covalent bonds

(d) Van-der Waal’s forces

34. The number of Na⁺ ions which surround each Cl^- ion in the NaCl crystal lattice is

(a) 8 (b) 12 11204153

35. The structure of NaCl crystal is:11204154

(a) Body centered cubic lattice

(b) Face centered cubic lattice

(d) Octahedral

36. The coordination number of a body-centered atom in a crystal is: 11204155

(a) 4 (b) 6

37. L. Pauling has proposed a theory about metallic bonds which is called: 11204156

(a) Molecular orbital theory

(b) Electron gas theory

(d) Valence bond theory

38. The electrical conductivity of the metals decreases with increasing temperature. This is because: 11204157

(a) The number of free electrons decreases

(b) The bonds of the metal atoms become weak

(d) The number of positive spheres increases

39. All the metals shine when they are freshly cut. The reason is that 11204158

(a) The conductivity of the metal is increased

(b) The process of cutting gives energy to the metal atoms

(d) The electrons are excited at higher energy levels and emit the photons when they fall back

40. The arrangement ABC, ABC------ is referred as: 11204159

(a) Cubic close packing

(b) Octahedral close packing

(d) Tetrahedral close packing

41. Which of the following isacovalent solid? 11204160

(a) Sugar (b) Diamond

42. Which one of the following is amorphous in nature? 11204161

(a) Glass (b) Rubber

43. The pure Crystalline substance on heating become turbid liquid. On further heating turbidity disappears. The substance is:

(a) Allotropic Crystal 11204162

(b) Liquid Crystal

(d)Isomorphic Crystal

44. Which of the following describes the hexagonal close packed arrangement of spheres? 11204163

(a) ABC ABC (b) ABC ABA

45. A solid ‘X’ melts slightly above 273K and is a poor conductor of heat and electricity. To which of the following categories does it belong? 11204164

(a) Ionic solid (b) Covalent solid

46. Which of the following solids conduct current in solid state? 11204165

(a) Iodine (b) Diamond

47. In Ionic Crystals the cleavage occurs because ions are: 11204166

(a) Weakly bonded

(b)Strongly bonded

(d) Separated by large distances

48. In how many groups, Bravislattices are arranged? 11204167

(a) 7 (b) 14

49. The crystals which show different physical properties from different directions is called 11204168

(a) Symmetry (b) Polymorphism

50. The delocalized molecular orbitals which extend over the entire crystal lattice in metallic solid is explained by: 11204169

(a) Electron gas theory

(b) Valence bond theory

(d) Metallic theory

51. The Avogadro’s number can be determined by the study of Crystalline solid if we are provided with: 11204170

(a) Density of one gram mole of crystals

(b) Volume of one gram mole of crystals

(d) The distance between particles & volume of one gram mole of crystals

52. Silicon Carbide is very hard that’s why it is also used as abrassive, its hardness is due to: 11204171

(a) Strong Ionic bond

(b) Strong Intermolecular forces

(d) Metallic bond

53. In diamond each carbon atom is sp³ hybridized and these carbon atoms arrange themselves in: 11204172

(a) One dimension (b)Two dimensions

54. The solid crystals in which neutral atoms of same or different elements are arranged, are called: 11204173

(a) Ionic solids (b) Metallic solids

55. The solids are classified on the basis of bonding into: 11204174

(a) Two types (b) Four types

56. The allotropes are those solids, which have: 11204175

(a) Same physical properties

(b) Same chemical properties

(d) Same physical & chemical properties

57. Coordination number of Na⁺ ion in NaCl is: 11204176

(a) one (b) two

58. Transition temperature of tin is:11204177

(a) 95.5^oC (b) 13.2^oC

59. The crystal of diamond is: 11204178

(a) Ionic (b) Covalent

60. The repetition of faces, angles and edges when a crystal is rotated by 360^o along its axis is called: 11204179

(a) Habit of crystal (b) Symmetry

61. Which of the following metals shows hexagonal geometry? 11204179

(a) Cu (b) Ag

62. Cadmium iodide is an example of a: 11204179

(a) covalentsolid (b) molecular solid

63. Graphite belongs to the crystal systems? 11204179

(a) hexagonal (b) monoclinic

64. Transition temperature of potassium nitrate is: 11204179

(a) 13.2^oC (b) 95.5^oC

65. Which of the following crystal systems represent the structure of sugar: 11204179

(a) triclinic (b) monoclinic

SHORT QUESTIONS

Q1.Sodium is softer than copper, but both are very good electrical conductors. Explain. 11204179

Q2.Diamond is hard and electrical insulator. Explain.11204180

Q3. Why NaCl and CsCl have different structures? 1204181

Q4.Iodine dissolves readily in tetra-chloromethane. Explain. (Board 2014) 11204182

Q5.The vapour pressures of solids are far less than those of liquids. Why? 11204183

Q6.Amorphous solid like glass is also called super cooled liquid. Why?11204184

Q7.Cleavage of crystals is anisotropic behaviour. Explain. (Board 2014)11204185

Q8.The crystals showing isomorphism mostly have the same atomic ratios. Why?

11204186

Q9.The transition temperature is given by elements having allotropic forms and by compounds showing polymorphism. Give reason. 11204187

Q10.One of the unit cell angles of hexagonal crystal is 120^oC. Explain. 11204188

Q11.The electrical conductivity of the metals decrease by increase in temperature. Why? (Board 2014)11204189

Q12.In the correct packing of atoms of metals, only 74% space is occupied. Give reason.11204190

Q13.Ionic crystals don’t conduct electricity in the solid state. Why?(Board 2014)11204191

Q14.The number of positive ions surrounding the negative ion in the ionic crystal lattice depend upon the sizes of the two ions. Give reason. 11204192

Q15.Why amorphous solids do not have sharp melting point? 11204193

Q16.What are crystallites? 11204194

Q17.Why is electrical conductivity of graphite larger from one side than other?

11204195

Q18.Why Polymorphic compounds have different physical properties?11204196

Q19.How molecular orbital theory explains the formation of metallic bond?11204197

Q20.How many Na⁺ ions and Cl^-¹ions are present in one cube of NaCl?11204198

Q21.Why freshly cut metals possess metallic luster? 11204199

Q22.What are crevices or viodes?11204200

Q23.Why ionic crystals are brittle?

(Board 2014, 2015)11204201

Q24.How habit of Crystal can be changed?

11204202

Q25.What is cleavage plane? 11204203

Q26. What are crystallographic elements?

(Board 2014)11204204

Q27.What is the difference between hexagonal close packing and cubic close packing? 11204205

Q28. Transition temperature is the term used for elements as well as compounds. Explain. (Board 2015) 11204206

Q29. What is the relationship between polymorphism and allotropy?

(Board 2015) 11204207

Q30. Define Symmetry and Anisotropy. 11204208

Q31. Define Monoclinic System. Draw its shape. 11204209

Q32. What is Isomorphism? Give an example. (Board 2015) 11204210

ATOMIC STRUCTURE

Objectives

MULTIPLE CHOICE QUESTIONS

1. Azimuthal quantum number gives us the information about: 11205022

(a) Size of orbital

(b) Shape of orbital

(d) All of the above

2. Energy associated with the electron revolving in the third orbit of H atom is: (a) 82.08 kJ mole^-1 11205023

(b) 145.92 kJ mole^-1

(d) 245.92 kJ mole^-1

3. Cathode rays are material particles having definite: 11205024

(a) Time period & amplitude

(b) Frequency & amplitude

(d) Mass

4. Charge of the canal rays is: 11205025

(a) 1.6022 ´ 10^-19 coulomb

(b) 1.6022 ´ 10^-17 coulomb

(d) 1.6022 ´ 10^-18 coulomb

5. Lines of P fund series are produced when electron jumps down to: 11205026

(a) Fifth orbit (b) Third orbit

6. When l = 2, the magnetic quantum number can have : 11205027

(a) Six values (b) Five values

7. Neutron was discovered by: 11205028

(a) Rutherford (b) Bohr

8. Alpha rays consist of : 11205029

(a) Electrons (b) Protons

9. The value of radius of first Bohr’s orbit of Hydrogen atom is: 11205030

(a) 0.229 ´ 10^-10m (b) 0.32 ´ 10^-10 m

10. Charge to mass ratio of electron was discovered by : 11205031

(a) Millikan (b) Rutherford

11. Charge on an electron is: 11205032

(a) 1.602 ´ 10^-¹⁹C (b) 9.1 ´ 10^-³⁴C

12. Mass of an electron is: 11205033

(a) 9.1 ´ 10³¹kg

(b) 9.1 ´ 10^-³⁰kg

(d) 9.1 ´ 10^-³¹kg

13. Proton was discovered by: 11205034

(a) Chadwick (b) J.J. Thomson

14. The wave number of light emitted by a certain source is 2 ´ 10⁶ m^-1, the wave length of light is: (Board 2015) 11205035

(a) 200 nm (b) 500 nm

15. The number of f-orbitals associated with n = 5: 11205036

(a) 5 (b) 6

16. In the nuclear reaction + x. The ‘x’ is: (Board 2014) 11205037

(a) electron (b) proton

17. Concept of Elliptical orbits in an atom was proposed by: 11205038

(a) Bohr (b) Heisenberg

18. The energy of an electron at infinite orbital is: 11205039

(a) Positive (b) Zero

19. The value of radius of first Bohr’s orbit is: 11205040

(a) 0.229A° (b) 0.329A°

20. Product of uncertainties of momentum & position is: 11205041

(a) 2ph (b)

21. Velocity of photon is: 11205042

(a) Dependent on its wavelength

(b) Independent of its wavelength

(d) Equal to square of amplitude

22. Orbitals with same energy are called:

(a) Hybrid orbitals (Board 2015) 11205043

(b) Molecular orbitals

(d) Degenerate orbitals

23. The mass of neutron is greater than electron by: 11205044

(a) 2000 times (b) 300 times

24. Formula for calculating the number of electrons in a sub-shell is: (Board 2014)

11205045

(a) 2n² (b) 2(2l+1)

25. The range of visible region of spectrum is from: 11205046

(a) 300 – 750 nm (b) 400 – 750 nm

26. Neutrons are present in all the atoms except: 11205047

(a) N (b) C

27. Rutherford’s experiment of scattering a particles showed for 1st time that an atom has: 11205048

(a) Electron (b) Protons

28. The number of spectral lines emitted when electron jumps from any higher orbit to fourth orbit are: 11205049

(a) 3 (b) 4

29. The spectral line obtained when an e^- jumps from n=6 to n=3 belongs to which series? 11205050

(a) Lyman series (b) Paschen series

30. The correct expression derived for the energy of an electron in the nth energy level is: 11205051

(a) E_n = (b) E_n =

31. Bohr’s model can explain: 11205052

(a) Spectrum of “H” atom only

(b) Spectrum of atom or ion containing one e^- only

(d) Spectrum of multi-electronic atoms.

32. Two electrons present in one orbital are distinguished by: 11205053

(a) Principal Quantum No.

(b) Azimuthal Quantum No.

(d) Spin Quantum No.

33. Rutherford’s alpha particle scattering experiment eventually led to the conclusion that: 11205054

(a) Mass and Energy are related

(b) Electrons occupy space around the nucleus

(d) The electrons are present in the nucleus

34. The average distance of an electron from the nucleus of an atom is: 11205055

(a) 10^-²⁰ cm (b) 10^-¹²cm

35. The mass of neutron is: 11205056

(a) 9.1095 x 10^-31kg (b) 2.67 ´ 10^-²⁷ kg

36. Which has largest frequency? 11205057

(a) X-rays (b) Microwaves

37. A quantum of light energy is called as:

(a) Proton (b) Electron 11205058

38. Which is correct relation? 11205059

(a) E = h (b) =

39. The photons of which of the following colors will be more energetic? 11205060

(a) Red (b) Violet

40. The number of space orientations in f-orbital are: 11205061

(a) 2 (b) 1

41. Which orbital lies upon the axis? 11205062

(a) (b) d_xy

42. Which pair of orbital’s can be introduced as degenerate pair? 11205063

(a) sp³, sp³ (b) sp³, sp²

43. Quantum mechanical model of atom explains that: 11205064

(a) Electron has waves according to de-broglie’s concept

(b) Heisenberg’s uncertainty principle is true

(d) All of the above

44. When a-particles is bombarded over beryllium ________ is produced: 11205065

(a) Positron (b) Proton

45. The spectrum of radiation from which particular radiation has been absorbed after passing through absorbing substance is called: 11205066

(a) Continuous spectrum

(b) Line emission spectrum

(d) Band spectrum

46. If value of Azimuthal quantum number is 3, the value of “m” will be: 11205067

(a) +3, +2, +1, 0, -1, -2, -3

(b) +2, +1, 0, -1,-2

(d) -1, 0, +1

47. The e/m ratio of cathode rays is _______ than e/m ratio of positive rays: 11205068

(a) Smaller (b) Greater

48. Which of the following is the most important factor that determines the chemical behaviour of atom? 11205069

(a) Solubility

(b) Atomic mass

(d) Electronic configuration

49. The elements that show abnormal electronic configuration are 11205070

(a) Fe and Mn (b) Cu and Cr

50. Hund’s rule states that when electrons enter to the same energy sub-levels they are: 11205071

(a) Singly occupied with same spin

(b) Doubly occupied with same spin

(d) Doubly occupied with different spin

51. Pauli exclusion principal states that no two electrons in a given orbital have:

(a) Same n, l, m quantum numbers

(b) Same four quantum numbers 11205072

(d) Same Azimuthal quantum number

52. The radiation with wavelength lesser than violet light is called : 11205073

(a) X-rays (b) Ultraviolet

53. Which of the following is the correct order of frequency of rays? 11205074

(a) UV > cosmic rays > X-rays > g-rays

(b) Cosmic rays > g-rays > X-rays > Ultraviolet

(d) g-rays > X-rays > UV > cosmic rays

54. Heisenberg’s uncertainty principle is not applicable to: 11205075

(a) Electron (b) Proton

55. An atomic orbital has l = 1, m = +1, 0, -1, n=3 then which one of the following atomic orbitals has such values? 11205076

(a) 2s (b) 2p

56. The shape of atomic orbital is:

(a) Spherical 11205077

(b) Dumb-bell

(d) Sausage

57. Which atomic orbital has lowest energy? 11205078

(a) 4f (b) 5d

58. Wave number has the unit: 11205079

(a) m (b) cm

59. Spectrum of sodium contains two lines in the region 11205080

(a) Violet region (b) Red region

60. The nature of positive rays depends upon: (Board 2014) 11205081

(a) Nature of electrode

(b) Nature of discharge tube

(d) None of the above

61. e/m value for positive rays is maximum for: 11205082

(a) hydrogen (b) helium

62. According to Bohr’s atomic model radius of second orbit of hydrogen atom is: 11205083

(a) 0.529A⁰ (b) 2.116 A⁰

63. An orbital, which is spherical and symmetrical is: 11205084

(a) s-orbital (b) p-orbital

64. Rutherford’s model of atom failed because: 11205085

(a) The atom did not have a nucleus and electron

(b) It did not account for the attraction between proton and neutrons

(d) There is actually no space between the nucleus and the electrons

65. Bohr’s model of atom is contradicted by: 11205086

(a) Planck quantum theory

(b) Pauli’s exclusion principle

(d) All of the above

66. Splitting of spectral lines when atoms are subjected to strong electric field is called: 11205087

(a) Zeeman effect

(b) Stark effect

(d) Compton effect

67. In the ground state of an atom, the electron is present: 11205088

(a) in the nucleus

(b) in the second shell

(d) farthest from the nucleus

68. Quantum number values for 2p orbitals are : 11205089

(a) n = 2, = 1

(b) n = 1, = 2

(d) n = 2, = 0

69.Orbitals having same energy are called: 11205090

(a) hybrid orbitals

(b) valence orbitals

(d) d-orbitals

70. When 6d orbital is complete, the

entering electron goes into: 11205091

(a) 7f (b) 7s

71. Lyman series is obtained when electron in an atom jumps from higher energy level to: 11205092

(a) ground level (b) 2^nd level

72. Which of the following gives correct relation: 11205093

(a) E=mc² (b)

73. Particle having the longest wavelength even if they have same speed: 11205094

(a) electron (b) proton

74. Mass of a neutron in amu is equal to: 11205095

(a) 1.0073 (b) 1.0087

75. Which of the following is correct for energy when we go from lower to higher orbits in case of Bohr’s atomic model: 11205096

(a) E₂ – E₁ < E₃ – E₂ > E₃ – E₄ > …….

(b) E₂ – E₁ > E₃ – E₂ > E₄ – E₃ > …….

(d) E₃ – E₄ > E₃ – E₂ > E₄ – E₃ > …….

74. The probability of finding the electron outside the nucleus is at a distance of: 11205097

(a) 0.83 m (b) 0.053 m

SHORT QUESTIONS

Q1. Why is it necessary to decrease the pressure in the discharge tube to produce cathode rays? 11205098

Q2. Whichever the gas is used in discharge tube the nature of cathode rays remains same why? (Board 2014) 11205099

Q3. Why e/m value of the cathode rays is just equal to that of electrons? 11205100

Q4. How the bending of cathode rays in the electric and magnetic field shows that they are negatively charged? 11205101

Q5. Why are the positive rays also called canal rays? 11205102

Q6. The e/m values of positive rays of different gases are different but of cathode rays is same. Why? 11205103

Q7. The e/m ratio for positive rays is 1836 times less than that of cathode rays. Why?

11205104

Q8. The potential energy of the bounded electron is negative. How? 11205105

Q9. Total energy of bounded electron is negative. Explain. 11205106

Q10. Energy of an electron is inversely proportional to n² but energy of higher orbits is always greater than those of lower orbits. How? (Board 2010) 11205107

Q11. The energy difference between adjacent levels goes on decreasing sharply. How?

(Board 2015) 11205108

Q12. H and He⁺ both have only one electron in their outer shells, but energy for both electrons is different. Why? 11205109

Q13. Do you think that groups of spectral lines of He⁺ and of H° are at different places?

11205110

Q14. How is Emission Spectrum obtained?

11205111

Q15. What is absorption spectrum? 11205112

Q16. What is continuous spectrum? 11205113

Q17. What is discharge tube? 11205114

Q18. What is De-Broglie’s Equation? 11205115

Q19. What is Stark effect?

(Board 2015) 11205116

Q20. What is Zeeman effect?

(Board 2015) 11205117

Q21. Write any four properties of neutrons. 11205118

Q22. What do you mean that energy of e^- is quantized? 11205119

Q23. What is the role of x-rays in the Millikans oil drop experiment for determination of charge of e^-? 11205120

Q24. How it can be proved that cathode rays have momentum? 11205121

Q25. What products are obtained by the neutron decay? OR (Board 2014) 11205122

Write a nuclear reaction for the decay of free neutron.

Q26. What are slow neutrons? Why they are more effective than the fast neutrons?

(Board 2014) 11205123

Q27. What is the limiting line of Balmer series and in which region does it fall?

11205124

Q28. What are H_a, H_b, H_g and H_s lines in the H-spectrum? 11205125

Q29. What information is obtained from principal quantum number (n)? 11205126

Q30. What is the concept of de-Broglie about duality of matter? 11205127

Q31. What information is obtained from magnetic quantum number? 11205128

Q32. Why use of X-rays increases the uncertainty in the momentum of electron?

11205129

Q33. What is Moseley’s Law and its significance? (Board 2014) 11205130

Q34. How was the wave nature of electron verified experimentally? (Board 2014) 11205131

Q35. What is the difference between continuous and line spectrum?(Board 2014)

11205132

Q36. Define Pauli’s exclusion principle.

(Board 2010, 14, 15) 11205133

Q37. Write the electronic configuration of the elements. Cu = 29 K = 19. (Board 2008)

11205134

Q38. Why is values of cathode rays same for all gases? (Board 2010) 11205135

Q39. State Heisenberg’s uncertainty principle and represent its formula. 11205136

Q40. Why is atomic spectrum, a line spectrum? 11205137

Q41. What is Aufbau Principle? 11205138

Q42. Why an electron moves faster in an orbit of smaller radius? (Board 2013) 11205139

Q43. Calculate the mass of electron when its e/m value is 1.7588 ´ 10¹¹ Ckg^-1.

(Board 2014) 11205140

Q44. Write down the electronic configuration of Fe(26) and Br(35).

(Board 2014) 11205141

Q45. Define spectrum. Give its two types. (Board 2009) 11205142

Q46. How we come to know that cathode rays are material particles with negative charge? (Board 2007) 11205143

Q47. What are defects in Rutherford’s atomic model? (Board 2015) 11205144

Q48. How will you relate energy of emitted light with its frequency and wavelength? 11205145

Q49. State Hund’s rule and give one example. (Board 2007, 2015) 11205146

2p_x

2p_y

2p_z

CHEMICAL BONDING

Objectives

MULTIPLE CHOICE QUESTIONS

1. In which of the following molecules, ionic bond is found?

(a) NH₃ (b) H – F 11206035

2. The most electronegative atom is: 11206036

(a) N (b) Cl

3. Tendency of an atom to attract the shared pair of electrons is called:

(Board 2013) 11206037

(a) Ionization Energy

(b) Electron Affinity

(d) Electropositivity

4. Electrons are filled in various orbitals according to: 11206038

(a) Auf bau principle (b) Hunds rule

5. An ionic compound will dissolve in water only if: 11206039

(a) Hydration Energy < Lattice energy

(b) Bond energy is low

(d) Bond energy is high

6. Which of the following has smallest size? 11206040

(a) Be (b) B

7. Which specie has largest size? 11206041

(a) Fe (b) Fe⁺

8. The strongest bond is: 11206042

(a) C – C (b) C = C

9. The weakest bond is: 11206043

(a) ionic (b) polar covalent

10. sp³ Hybridization is associated with structure: 11206044

(a) linear (b) trigonal

11. Which of the following has all the three characters i.e. ionic, covalent and coordinate covalent? 11206045

(a) H₂O (b) KBr

12. Highest bond order is found in: 11206046

(a) O₂ (b)

13. The molecule with linear structure is:

(a) H₂O (b) H₂S 11206047

14. Paramagnetic specie is: 11206048

(a) (b)

15. The number of lone pair of electrons in ammonium ion is: 11206049

(a) one (b) three

16. Which of the following molecules show sp hybridization of central atom? 11206050

(a) BeCl₂

(b) NH₃

(d) C₂H₄

17. Which one is a polar molecule? 11206051

(a) CO₂ (b) HCl

18. Which one is a non polar molecule?

(a) CaCl₂ (b) CaC₂ 11206052

19. Covalent bonds are: 11206053

(a) Rigid and directional

(b) Non rigid and directional

(d) Non rigid and non directional

20. The molecule with greatest dipole moment is: 11206054

(a) H₂O (b) H₂S

21. Which will have zero dipole moment?

(a) H₂O (b) H₂S (Board 2014)

22. The SI Unit for dipole moment is:

(a) Debye (b) Nm 11206056

23. One Debye is equal to: 11206057

(a) 3.336 ´ 10^-30 mC (b) 9.1 ´ 10^-31 mC

24. Dipole moment is the product of: 11206058

(a) Charge ´ displacement

(b) Charge ´ distance

(d) Charge ´ mass

25. Dipole moment of CO₂ is: 11206059

(a) 1.8 D (b) 1.94 D

26. Octet rule is not obeyed by which of the following compounds during its formation? 11206060

(a) H₂O (b) NaCl

27. The bond distance between H – H is:

(a) 436.45 pm (b) 74.5 pm 11206061

28. The bond formation energy of a compound is: 11206062

(a) less than bond dissociation energy

(b) greater than bond dissociation energy

(d) inversely proportional to bond dissociation energy

29. One of the best techniques used to measure atomic radii is: 11206063

(a) Spectrometry (b) Chromatography

30. One pico metre (pm) is equal to: 11206064

(a) 10^-2 m (b) 10^-9 m

31. The atomic radii decrease along the period due to.: 11206065

(a) Increase in nuclear charge

(b) Increase in atomic mass

(d) None of the above

32. The increase in atomic radii in a group is due to: 11206066

(a) Increase in atomic mass

(b) Increase in number of shells

(d) Increase in polarizability

33. Metals usually have: 11206067

(a) High volatility

(b) Low value of 1^st ionization energy

(d) Low thermal conductivity

34. What will be the valency of element, when there is sufficient difference in 1st and 2nd Ionization energy value?

(a) 1 (b) 2 11206068

35. Which one is phosphonium Ion? 11206069

(a) (b)

36. In Hydronium Ion each bond is: 11206070

(a) 100% covalent

(b) 50% covalent and 50% coordinate covalent

(d) 33% covalent and 66% coordinate covalent

37. Which theory explains the paramagnetic behaviour exhibited by O₂ molecule?

(a) Band theory (b) VBT 11206071

38. What is molecular geometry of Ion? 11206072

(a) Tetrahedral

(b) Trigonal planar

(d) Linear

39. Which one is not AB₄ type molecule according to VSEPR theory? 11206073

(a) SO₂ (b)

40. What is bond angle in NF₃? 11206074

(a) 109.5° (b) 107.5°

41. Which of the following hybridizations is found in ethyne? 11206075

(a) sp² (b) sp

42. Which one of the following will show paramagnetic property? 11206076

(a) O (b) O₂

43. The expected bond energy of HCl is lesser than the actual, this is because:

(a) Size of hydrogen is very small 11206077

(b) HCl is non – polar compound

(d) There exists hydrogen bonding

44. The actual bond length in a polar covalent compound is: 11206078

(a) lesser than expected

(b) greater than expected

(d) exactly half of expected

45. The bond angle in ethene molecule is:

(a) 180^o (b) 109.5^o 11206079

46. The phenomenon of Isomerism occurs when a bond in a compound is: 11206080

(a) Non rigid & non directional

(b) Non rigid but directional

(d) Rigid but non directional

47. What happens when two atoms repel each other? 11206081

(a) No change in energy occurs

(b) They lose energy

48. If the electronegativity difference between bonded atoms is more than 0.5 but less than 1.7, the bond will be: 11206082

(a) Ionic (b) Polar covalent

(d) Coordinate covalent

49. Which of the following elements has incomplete octet after making the bond?

(a) C in CH₄ (b) O in H₂O 11206083

50. When a non polar covalent compound is put in water, it will be: 11206084

(a) Slightly soluble

(b) Soluble at room temperature

51. According to VSEPR theory which of the following electron pairs show maximum repulsion? 11206085

(a) Bond pair – bond pair

(b) Bond pair – lone pair

(d) None of the above

52. The nature of hybridization in oxygen atom in H₂O is: 11206086

(a) sp³ (b) sp²

53. The bond lengths of multiple bonded atoms are: 11206087

(a) equal to single bonded atoms

(b) larger than the single bonded atoms

(d) exactly double than the single bonded atoms

54. Pi – bond is present in: 11206088

(a) H₂ (b) H₂O

55. Which one of the following molecules have angle of 120°? 11206089

(a) BeCl₂ (b) BF₃

56. Bond order of He₂is: 11206090

(a) Zero (b) 1

57. Ionic radius of Cl^- Ion is: 11206091

(a) 151 pm (b) 161 pm

58. The central atom of which of the following compounds is able to form coordinate covalent bond? 11206092

(a) BeCl₂(b) CH₄

59. The Ionic character in CsF is:

(Board 2014) 11206093

(a) 92% (b) 100%

60. The Ionic character in NaCl is: 11206094

(a) 56% (b) 66%

61. BH₃, BF₃, AlCl₃ all have: 11206095

(a) Bent structure

(b) Trigonal planar structure

(d) Different structures

62. Shielding effect of the inner electrons is responsible for: 11206096

(a) Increasing electron affinity

(b) Decreasing Ionization energy

(d) None of the above

63. When two p–orbitals overlap perpendicular to the line joining the two nuclei, the bond formed is: 11206097

(a) Sigma bond (b) p - bond

64. In SO₂, the number of p bonds are:

11206098

(a) Zero (b) One

65. Which one of the following is a vector quantity? 11206099

(a) Ionization energy (b) Electron affinity

66. The structure of Ammonia is: 11206100

(a) Linear (b) Tetragonal

(d) Triangular Pyramidal

67. Which of the following species has unpaired electrons in antibonding molecular orbitals? 11206101

(a) H₂ (b) He₂

68. A molecular orbital can contain maximum electrons equal to. OR

Maximum no. of electrons in an orbital is: (Board 2014) 11206102

(a) 6 (b) 10

69. An Angstrom is the unit of: 11206103

(a) Time (b) Length

70. The most stable elements are: 11206104

(a) Halogens (b) Lithium family

71. The hybridization of Carbon in C₂H₄ is:

(Board 2014) (B - 2009) 11206105

(a) sp (b) sp²

72. The bond order of N₂ molecule is:

(Board 2013) 11206106

(a) 1 (b) 2

73. An ionic compound A⁺B^- is most likely to be formed when:(Board 2014) 11206107

(a) The ionization energy of A is high and electron affinity of B is low

(b) The ionization energy of A is low and electron affinity of B is high

(d) Both the ionization energy of A and electron affinity of B are low

74. The number of bonds in nitrogen molecules is: 11206108

(a) one s and one p

(b) One s and two p

(d) two s and two p

75. Which of the following statements is not correct regarding bonding molecular orbitals? 11206109

(a) Bonding molecular orbitals possess less energy than atomic orbitals from which they are formed

(b) Bonding molecular orbitals have low electron density between the two nuclei

(d) Bonding molecular orbitals are formed when the electron waves undergo constructive interference

76. Which of the following molecules has zero dipole moment? 11206110

(a) NH₃ (b) CHCl₃(c) H₂O (d) BF₃

77. Which of the following hydrogen halides has the highest percentage of ionic character? 11206111

(a) HF (b) HBr

78. Which of the following molecules has unpaired electrons in antibonding molecular orbitals? 11206112

(a) (b)

79.Which of the following has highest bond energy? (Board 2015) 11206113

(a) HI (b) HBr

80.The number of bonds in oxygen molecule is: (Board 2015) 11206114

(a) one s and one p (b) one s and two p

81.Formation of a chemical bond takes place when: (Board 2013) 11206115

(a) energy is absorbed

(b) forces of repulsion overcome forces of attraction

(d) forces of attraction overcome forces of repulsion

82.The geometry of ethane is: (Board 2013)

11206116

(a) Trigonal Planar (b) Linear

83.Bond formed by mutual sharing of electrons is called: (Board 2012) 11206117

(a) Ionic bond

(b) Covalent bond

SHORT QUESTIONS

Q1. Bond distance is the compromise distance between two atoms. How? 11206118

Q2. The distinction between a coordinate covalent bond and a covalent bond vanishes after bond formation in . Why?

11206119

Q3. The bond angles of H₂O and NH₃ are not 109.5° like that of CH₄ although O, N and C are all sp³ hybridized. Why? 11206120

Q4. Pi-bonds are more diffused than sigma bonds. Why? (Board 2014) 11206121

Q5. The abnormality of bond length and bond strength in HI is less prominent than that of HCl. Why? 11206122

Q6. Solid sodium chloride does not conduct electricity but when electric current is passed through molten sodium chloride or its aqueous solution, electrolysis takes place. Why? 11206123

Q7. The melting point, boiling point, heat of vapourization and heats of sublimation of electrovalent compounds are higher as compared to those of covalent compounds. Why? 11206124

Q8. What is an octet rule? 11206125

Q9. Some atoms in certain compounds like PF₅, SF₆ and BCl₃ do not obey octet rule. Why? 11206126

Q10. Atomic radii increase down a group while decrease along a period. Why? 11206127

Q11. What is shielding effect? 11206128

Q12. A bond cannot be 100% ionic. Why?

11206129

Q13. Sigma bond is stronger than Pi bond. Why? 11206130

Q14. What is atomic radius? 11206131

Q15. Why atomic radius cannot be determined precisely? (Board 2014) 11206132

Q16. What is a chemical bond? 11206133

Q17. What is bond distance? 11206134

Q18. What is Ionic radius? 11206135

Q19. What is a covalent radius? 11206136

Q20. What is Ionization energy? 11206137

Q21. What factors effect I.E? (Board 2010)

11206138

Q22. Why I.E increases from left to right in a period? 11206139

Q23. Why second Ionization energy value of an atom is higher than the first I.E value?

11206140

Q24. What is electron affinity? 11206141

Q25. What factors effect electron affinity?

11206142

Q26. Why electron affinity of Fluorine is lesser than chlorine? 11206143

Q27. What is electronegativity? 11206144

Q28. How is sigma bond formed? 11206145

Q29. How is p - bond formed? 11206146

Q30. What is paramagnetism? 11206147

Q31. What is bond enthalpy? 11206148

Q32. How % age Ionic character can be calculated? 11206149

Q.33 The dipole moments of CO₂ & CS₂ are Zero, but that of SO₂ is 1.61 D. Why? 11206150

Q34. Why is atomic radius greater than cationic radius? OR (Board 2010)

Why is the radius of a cation smaller than its parent atom? (Board 2014) 11206151

Q35. Why is ionic radius of an ion greater than its parent atom? 11206152

Q.36 Why ionic compounds do not show phenomenon of isomerism? (Board 2010) 11206153

Q37. Write down two postulates of VSEPR theory. 11206154

Q38. Why is electron affinity of uninegative ion positive? 11206155

Q39. Why is Molecular orbital theory superior to valence bond theory? 11206156

Q40. Why the bond energies of the multiple bonds are greater than those of single bonds? 11206157

Q41. What is bond order? Give an example.

(Board 2014) 11206158

Q42. and have different bond angles. Justify. 11206159

Q43. How the electronegativity difference of two bonded atoms can be used to predict the ionic / covalent nature of a bond?

(Board 2014) 11206160

Q44. Why the ionic radius of Cl^–1 ion increases from 99 pm to 181 pm? 11206161

(Board 2013)

Q45. How a coordinate covalent bond differs from covalent bond? 11206162

Q46. Differentiate between atomic orbital and molecular orbital. 11206163

THERMOCHEMISTRY

Objectives

MULTIPLE CHOICE QUESTIONS

1. If an endothermic reaction is allowed to take place very rapidly in the air, the temperature of the surrounding air:

(a) remains constant 11207035

(b) increases

2. In an endothermic reaction, the heat content of the: 11207036

(a) products is more than that of reactants

(b) reactants is more than that of products

(d) both ‘a’ and ‘b’

3. Calorie is equivalent to:

(Board 2013, 14, 15) 11207037

(a) 0.4184 J (b) 41.84 J

4. The change in heat energy of a chemical reaction at constant temperature and pressure is called: 11207038

(a) Enthalpy change (b) Bond energy

(d) Internal energy change

5. Which of the following statements is contrary to the first law of thermodynamics? 11207039

(a) Energy can neither be created nor be destroyed

(b) One form of energy can be transferred into an equivalent amount of other kind of energy

(d) Continuous production of mechanical work without supplying an equivalent amount of heat is possible

6. For a given process, the heat changes at constant pressure q_p and at constant volume (q_v) are related to each other as:

(a) q_p = q_v (Board 2015) 11207040

(b) q_p < q_v

(d) q_p = q_v/2

7. For the reaction NaOH + HCl ® NaCl + H₂O, the change in enthalpy is called:

(a) Heat of reaction 11207041

(b) Heat of formation

(d) Heat of combustion

8. The net heat change in a chemical reaction is same whether it is brought about in two or more different ways in one or several steps. It is known as:

(a) Henry’s Law 11207042

(b) Hess’s Law

(d) Law of conservation of energy

9. Enthalpy of neutralization of all the strong acids and strong bases has the same value because: 11207043

(a) Neutralization leads to the formation of salt and water

(b) Strong acids and bases are ionic substances

(d) The net chemical change involves the combination of H⁺ and OH^- ions to form water

10. Heat of combustion can be determined by: 11207044

(a) Glass calorimeter

(b) Bomb calorimeter

(d) Glass and bomb calorimeter

11. In bomb calorimeter the reactions are carried out at: 11207045

(a) Constant pressure

(b) Constant volume

(d) All of the above

12. When coefficients of chemical equation are doubled DH: 11207046

(a) Halves (b) Doubles

13. The subject matter of thermo-chemistry is based upon: 11207047

(a) Hess’s Law

(b) Born Haber’s cycle

(d) Law of mass action

14. Which of the following is the mathematical expression of First Law of Thermodynamics? 11207048

(a) q = m.S.DT (b) DH = DE + PDV

15. Which of the following is not a state function? (Board 2014) 11207049

(a) Heat (b) Work

16. When work is done on a system then its sign is: 11207050

(a) +ve (b) -ve

17. SDH_cycleis another definition of: 11207051

(a) State function

(b) Internal Energy Change

(d) Entropy change

18. First Law of Thermodynamics is also called: 11207052

(a) Hess’s Law (b) Born Haber cycle

(d) None of the above

19. The condition of standard state of a substance can be represented as: 11207053

(a) 700 mm of Hg and 0°C

(b) 1 atm and 0°C

(d) 2 atm and 0°C

20. Which of the following formulas is used to calculate the amount of heat change in a chemical reaction? 11207054

(a) DH = DE + PDV (b) DE = DH + W

21. For a given reaction which description is correct? 11207055

CH₄+ 2O₂ ¾® CO₂ + 2H₂O

DH= -890.6kJ/mole

(a) Heat of formation of CH₄

(b) Endothermic reaction

(d) Heat of Atomization

22. Which of the following apparatus is used for the accurate determination of enthalpy of neutralization? 11207056

(a) Copper calorimeter

(b) Glass calorimeter

(d) Glass and copper calorimeter

23. The heat of neutralization for a strong acid and a strong base is 57. kJ/mole. The heat of neutralization for the reaction of NaOH and CH₃COOH is: 11207057

(a) 57.kJ/mole

(b) more than 57 kJ/mole

(d) Cannot be predicted

24. The internal energy of a system is equal to: 11207058

(a) P.E. of the particles

(b) Kinetic energy of the particles

(d) Sum of K.E and P.E.

25. Heat absorbed by the system at constant volume is equal to: 11207059

(a) Enthalpy change of system

(b) Internal energy change of a system

(d) Kinetic energy change of a system

26. Which of the following is the application of Law of conservation of energy? 11207060

(a) 1st Law of Thermodynamics

(b) Hess’s Law

27. Born Haber cycle is used to calculate the: 11207061

(a) Lattice energy (b) Enthalpy change

(d) Heat of Atomization

28. Solids which have more than one crystalline forms possess : 11207062

(a) Zero DH values (b) Same DH_f values

(d) None of the above

29. The branch of chemistry which deals with the energy changes during a chemical reaction is called: 11207063

(a) Chemical Equilibrium

(b) Stoichiometry

30. For H⁺¹₊ OH^-1 ¾® H₂O. The change in enthalpy for the reaction is called:

11207064

(a) Heat of formation of water

(b) Heat of reaction

(d) Heat of combustion

31. Standard enthalpies are measured at:

(Board 2009) 11207065

(a) 273^oC (b) 298K

32. In an endothermic reaction, ∆H is taken as: 11207066

(a) positive (b) negative

33. The enthalpy of combustion is: 11207067

(a) positive

(b) negative

(d) none of these

34. Spontaneous processes are mostly: 11207068

(a) reversible (b) irreversible

35. Which one of the following enthalpies is always an exothermic process? 11207069

(a) DH_c (b) DH_s

SHORT QUESTIONS

Q1. What are the two fundamental ways of transferring energy to or from a system?

11207070

Q2. Why is it necessary to mention the physical states of the reactants and products in thermo-chemistry? 11207071

Q3. What is the significance of q_p? 11207072

Q4. What is significance of q_v? 11207073

Q5. What is standard state of a substance?

11207074

Q6. What is the difference between Heat and Temperature? 11207075

Q7. Whether all the natural changes are reversible or irreversible? 11207076

Q8. What is the relation between DH and DE when volume of a system is decreased?

11207077

Q9. What is the relation between heat of neutralization of a weak acid and strong acid?

11207078

Q10. Whether enthalpy (H) is an extensive property or a state function? 11207079

Q11. Why are some reactions exothermic and others are endothermic? 11207080

Q12. Define thermo-chemistry and on what factors it is based upon? 11207081

Q13. What is Spontaneous and a Non-spontaneous process? (Board 2014) 11207082

Q14. Burning of coal is a spontaneous or non-spontaneous process. Explain. 11207083

Q15. All exothermic processes are spontan-eous but all spontaneous processes are not exothermic. Why? 11207084

Q16. Define: (Board 2014) 11207085

(i) System (ii) Boundary

(iii) Surrounding (iv) State function

Q17. Write a short note on Internal Energy. / What do you understand by internal energy of a system? (Board 2014) 11207086

Q18.What is First Law of Thermo-dynamics? 11207087

Q19. What is difference between DE and DH? 11207088

Q20. Prove that DH = DE 11207089

Q21. State first law of thermo chemistry.

11207090

Q22. Heat of neutralization of a strong acid with a strong base has always the same value. Explain. (Board 2014) 11207091

Q23. How would you explain that work is the product of pressure and volume? 11207092

Q24. What is Born-Haber Cycle? 11207093

Q25. Define enthalpy of solution (DH^oSol.).

11207094

Q26. Heat is evolved when stronger bonds are being formed and absorbed when weaker bonds are formed. Explain. 11207095

Q27. Hess’s Law is a special case of first law of thermodynamics. Justify it. 11207096

Q28. Hess’s Law is employed to calculate DH value of those chemical reactions which cannot be normally carried in a laboratory. Explain it. (Board 2009) 11207097

Q29. The Enthalpy of combustion of Graphite at 25°C is –393.51 kJ/mole and that of Diamond is –395.4 kJ/mole. What is Enthalpy change of the process C_graphite ® C_diamondat the same temperature? 11207098

Q30. Explain that burning of a candle is spontaneous process. (Board 2008) 11207099

Q31.What are exothermic and endothermic reactions? Give one example of each. 11207100

Q32.Define state function and give example.

11207101

e.g. T, P, V, H, E are all state functions.

Q33. Justify that heat of formation of compound is sum of all the other enthalpies.

11207102

Q34.State whyH is approximately equal toE in case of liquids and solids. 11207103

Q.35 What is thermochemical equation? Give two examples. (Board 2014) 11207104

Q.36 Define Lattice energy. Give its SI units. 11207105

CHEMICAL EQUILIBRIUM

Objectives

MULTIPLE CHOICE QUESTION

1. For which system does the equilibrium constant, K_c has unit of (concentration)^-1?

(a) N₂ + 3H₂ 2NH₃ 11208039

(b) H₂ + I₂ 2HI

(d) 2HF H₂ + F₂

2. Which statement for the following equilibrium is correct? (Board 2013)

11208040

2SO₂ + O₂ 2SO₃ DH = -188 kJ/mole.

(a) The value of K_p falls with rise in temperature

(b) The value of K_p falls with rise in pressure

(d) The value of K_p is equal to K_c

Solution:

K_p = K_c(RT)^Dⁿ Dn = n_p – n_R

Dn = 2 – 3 = - 1

K_p = K_c(RT)^-1

K_p = K_c 1/RT

3. The pH of 10^-3 mole dm^-3 of an aqueous solution of H₂SO₄ is: (Board 2014, 2015)

(a) 3.0 (b) 2.7 11208041

4. The solubility product of AgCl is 2.0´10^-10 mole² dm^-6, the maximum concentration of Ag⁺ ions in solution is:

(a) 2.0 ´ 10^-¹⁰ mole dm^-3 (Board 2013)

(b) 1.41 ´ 10^-5 mole dm^-3 11208042

(d) 4.0 ´ 10^-20 mole dm^-3

5. An excess of aqueous solution of silver nitrate is added to aqueous barium chloride and precipitate is removed by filtration. What are the main ions in the filtrate? (Board 2015) 11208043

(a) Ag⁺ and NO₃^-1 only

(b) Ag⁺, Ba⁺² and NO₃^-1

(d) Ba⁺² , NO₃^-1 and

6. The pH of 10^-4 HCl is: 11208044

(a) 2 (b) 4

7. A large value of K_c means that at equilibrium: 11208045

(a) less reactant and more products are present

(b) more reactants and less products are present

(d) none of the above

8. In the reaction H₂ +I₂ 2HI the equilibrium is disturbed by: 11208046

(a) increasing pressure

(b) decreasing pressure

(d) no affect of pressure

9. Strength of an acid can be determined by: 11208047

(a) pK_a (b) pK_b

10. In an exothermic reversible reaction, increase in temperature shifts the equilibrium to: 11208048

(a) reactant side

(b) product side

(d) none of the above

11. The units of K_w are: 11208049

(a) mole dm^-3 (b) mole² dm^-3

12. If K_c of reaction is very large, it indicates that equilibrium occurs:

11208050

(a) at a low product concentration

(b) at a high product concentration

(d) with no forward reaction

13. For what value of K_c almost forward reaction is complete? 11208051

(a) K_c = 10^-35 (b) K_c = 10³⁰

14. Law of mass action was given by:

(a) Ramsay and Reyleigh 11208052

(b) Berkeley and Hartley

(d) Goldberg and Waage

15. The unit of K_c for the reaction: 11208053

N₂ + 3H₂ 2 NH₃ will be

(a) mole dm^-3 (b) mole^-1 dm⁺³

(a) mole² dm^-6 (d) mole^-2 dm⁺⁶

16. pH of human blood is: 11208054

(a) 7.0 (b) 7.35

17. pH of pure water is: 11208055

(a) 6.2 (b) 7

18. An acidic buffer solution can be prepared by mixing: 11208056

(a) A weak Acid and weak Base

(b) A strong Acid and weak Base

(d) A weak Base and its Salt with strong acid

19. The sum of pH and pOH is: 11208057

(a) 0 (b) 7

20. In a saturated solution of AgCl, the molar concentration of Ag⁺ and Cl^-1 is 1.0 ´ 10^-5 M each. What is the value of K_sp? 11208058

(a) 1.0 ´ 10^-5 (b) 1.0 ´ 10^-10

21. The dissociation constant for water at 25°C is: 11208059

(a) 1.0 ´ 10^-7 (b) 1.0 ´ 10^-14

22. H₂SO₄ is dibasic acid. The pH of 0.005M H₂SO₄ will be: 11208060

(a) 2 (b) 3

23. To prepare a buffer with pH close to 9.0, you could use mixture of: 11208061

(a) NH₄OH and NH₄Cl

(b) CH₃COOH and CH₃COONa

(d) H₂CO₃ and NaHCO₃

24. The pH of soft drink is: 11208062

(a) > 7 (b) < 7

25. If K_sp value is greater, the salt in water is: 11208063

(a) less soluble

(b) more soluble

(d) insoluble

26. What is true about chemical equilibrium?

11208064

(a) it is established in closed container

(b) it is established in reversible system

(d) all of the above

27. A small value of K_c means that at equilibrium: 11208065

(a) Less reactant and more products are present

(b) More reactant and less products are present

(d) None of the above

28. For the reaction 2Cl ¾® Cl₂the
K_c = 1.0 ´ 10⁺³⁰ this shows that : 11208066

(a) reaction will complete more in forward direction

(b) complete more in reverse direction

(d) none of the above

29. Strength of a base can be determined by:

(a) pK_a (b) pK_b 11208067

30. In an endothermic reversible reaction, increase in temperature shifts the equilibrium to: 11208068

(a) Reactant side

(b) Product side

(d) None of the above

31. If a solution has zero pH. The hydrogen ion concentration will be: 11208069

(a) 10^-4 (b) 10^-3

32. A basic buffer solution can be prepared by mixing: 11208070

(a) Weak acid and its salt with strong base

(b) Weak base and its salt with strong acid

(d) Strong base and its salt with weak acid

33. A buffer can be explained by: 11208071

(a) Common Ion effect

(b) Law of mass action

(d) All of the above

34. Ionization of weak acid is expressed in terms of which of the following constants? 11208072

(a) K_w (b) K_n

35. Solubility of Ca(OH)₂ in water is exothermic. Its solubility will increase:

(a) At high temperature 11208073

(b) At low temperature

(d) Moderate temperature

36. Which of the following is sparingly soluble ionic solid in water? 11208074

(a) NaCl (b) AgCl

37. The rate of forward step in a reversible reaction: (Board 2015) 11208075

(a) Increases during the reaction

(b) Decreases as the reaction proceeds

(d) None of the above

38. The units of equilibrium constant K_c for the reaction H₂ + I₂ 2HI are:

11208076

(a) (b) moles^-2 dm³

39. The catalyst used in the synthesis of NH₃ by Haber’s Process is: 11208077

(a) Asbestos (b) Al₂O₃ + SiO₂

40. Which of the following solutions will have highest pH? 11208078

(a) 0.01 M H₂SO₄ (b) 0.01 M NaOH

(d) 0.01 M NaHCO₃

41. Buffers are used in: 11208079

(a) Clinical Analysis

(b) Cell Biology

(d) All of the above

42. For a reversible reaction, if the concentration of reactant is doubled, then value of equilibrium constant K_cis: 11208080

(a) Halved

(b) Doubled

(d) Not changed

43. In a gaseous reversible reaction in a closed container which is exothermic in nature, an increase in temperature changes: 11208081

(a) Pressure of gases

(b) Conc. of both reactant and product

(d) All of the above

44. The pH of rain water is: 11208082

(a) 7 (b) 7.3

45. Which solution in H₂O will have pH less than 7? 11208083

(a) NaCl (b) CuSO₄

46. What will be the pH of buffer solution containing (CH₃COOH) = 1.0M and [CH₃COONa] = 0.1M: if K_a for acid is 1.85 ´ 10^-5? 11208084

(a) 4.74 (b) 5.74

47. The sparingly soluble salt in water is:

(a) CH₃COONa (b) BiCl₃11208085

48. Which one of the following examples, indicate an irreversible reaction?

(a) N₂ + 3H₂ ® 2NH₃11208086

(b) N₂ + O₂ ® 2NO

(d) CH₃ COOH + C₂H₅ OH ®

CH₃COOC₂H₅ + H₂O

49. What will be the pH of the buffer having [CH₃COOH] = 0.09 M and [CH₃ COONa] = 0.11 M if Ka for CH₃COOH is 1.85´10^-5?

(a) 4.83 (b) 4.92 11208087

50. Which physical state of Al (27gm) will have maximum active mass? 11208088

(a) Solid state (b) Al powder

51. According to common Ion effect _____ soluble salt is precipitated out first:

(a) More (b) Less 11208089

52. For the decomposition of N₂O₄ into NO₂, if we increase the pressure, it will favour the reaction in: 11208090

(a) Forward direction

(b) Backward direction

(d) First in forward and then in backward direction

53. For the reaction PCl₅ PCl₃ + Cl₂the increase in the pressure at equilibrium will shift the equilibrium to: 11208091

(a) forward direction

(b) reverse direction

(d) equal change on both sides

54. Optimum conditions for better yield of NH₃ are: 11208092

(a) 250 atm, 800°C

(b) 100 atm, 400°C

(d) 200-300 atm, 400°C

55. In the reaction 11208093

2SO₂ + O₂ ® 2SO₃ yield of product is maximum if:

(a) Temperature is increased and pressure is kept constant

(b) Temperature is decreased and pressure is increased

(d) Both temperature and pressure are decreased

56. A chemical equilibrium in a reaction is established when: 11208094

(a) Conc. of reactants and products are equal

(b) Opposing reaction is happening

(d) Velocities of opposing reactions is equal

57. In the formation of ammonia, if pressure is decreased, then the reaction should shift to:

(a) Forward direction 11208095

(b) Reverse direction

(d) None of the above

58. A pH 7 signifies: 11208096

(a) Rain water (b) Neutral solution

59. The pH of tomatoes is: 11208097

(a) 1.2 (b) 4.2

60. The concentrations of reactants and products at equilibrium are: 11208098

(Board 2009)

(a) Equal (b) Maximum

61. When 50% reactants in a reversible reaction are converted into a product, the value of equilibrium constant K_c is: (Board 2015) 11208099

(a) 2 (b) 1

62. Ionic product of water (K_w) increases when temperature increases from 0^oC to 100^oC: 11208100

(a) 25 times (b) 75 times

63. The term pH was introduced by:

(Board 2013) 11208101

(a) Henderson (b) Sorenson

SHORT QUESTIONS

Q1. What is meant by chemical equilibrium state? 11208102

Q2. Concentration of reactants and products remain constant at equilibrium position, although reaction continues to take place. Explain. 11208103

Q3. What do you know about homogeneous and heterogeneous equilibrium? 11208104

Q4. Rate of chemical reaction does not remain constant till equilibrium is attained. Why? 11208105

Q5. Exothermic reactions are favoured in forward direction by cooling and Endothermic reactions are disfavoured by cooling. Explain. 11208106

Q6. Can the direction of a reaction be predicted during the course of reaction before equilibrium, from a knowledge of [Product/Reactant]. Why? 11208107

Q7.A catalyst does not change the position of equilibrium but this equilibrium position reaches earlier. How? 11208108

Q8. How does a catalyst increases the rate of a chemical reaction? 11208109

Q9. Le-Chatelier’s principle says that solid ice at 0°C can be melted by applying pressure without supply of heat from out-side. Justify. 11208110

Q10. The reaction will move in forward direction if products are continuously removed from the reaction mixture. Explain. 11208111

Q11. How will you differentiate between conjugate acid and conjugate base? 11208112

Q12. What are Buffers? 11208113

Q13. How can solubility of Ca(OH)₂ be calculated from the value of solubility product constant? 11208114

Q14. What are units of K_c for synthesis of ammonia? 11208115

Q15. What is Buffer Capacity? (Board 2010)

11208116

Q16. Define Common ion Effect. Give an example. (Board 2014) 11208117

Q17. How can sodium chloride be purified by common ion effect? 11208118

Q18. What are two factors on which pH of Buffer is based? 11208119

Q19. Why do we need Buffer solutions?

11208120

Q20. What is the effect of common Ion on solubility of a sparingly soluble salt? 11208121

Q21. What is Henderson’s equation and for what purpose is it used?(Board 2014) 11208122

Q22. What is Lowry Bronsted acid-base concept? 11208123

Q23. What is meant by conjugate Acid base? Relate K_a and K_b of a conjugate Acid base pair. 11208124

Q24. How the K_sp value of sparingly soluble salt can be used to calculate solubility of the salt? 11208125

Q25. What is mechanism of buffer action?

11208126

Q26. The values of few acids are given below: 11208127

Phenol = 1.3 ´ 10^-10, CH₃COOH = 1.85 ´ 10^-5, H ¾ ¾ OH = 1.66 ´ 10^-4,

H₂CO₃ = 4.4 ´ 10^-7, arrange the following in ascending order of acid strength.

Q27. A solution has pH = -1, what is the conc. of H⁺ in it, what is conc. of OH^- in it, what is its pOH at Temp = 25°C? 11208128

Q28. Define solubility product. Write a general expression for it.(Board 2014) 11208129

Q29. The reaction of active metals like Na with water is irreversible, whereas the reaction of N₂ with H₂ to form NH₃ is reversible. Why? 11208130

Q30. K_c of H₂O (dissociation of H₂O) = 1.8´10^-16 mole/dm³ at 25°C but K_w = 1.01´10^-14. Why? 11208131

Q31. Give definition of only K_p, K_c, K_x, K_n. 08(121)

11208132

Q32. How are four equilibrium constant are related with one other? When all four will have same value? 11208133

Q33.Mention one reaction where K_c has no units and one reaction where K_c has some units (indicate the units also). 11208134

Q34. Write down K_c and K_p expression for the following reactions: 11208135

PCl_5(g) PCl_3(g) + Cl₂

CaCO_3(g) CaO_(s) + CO_2(g)

Q35. HCl acts as weak acid in ethanoic acid as compared to, when dissolved in H₂O. Why? 11208136

Q36. How is the value of K_c helpful for the estimation of extent of reaction? 11208137

Q37. Explain with suitable example, how is the value of K_c helpful in detecting direction of reaction. (Board 2014) 11208138

Q38. What are basic buffer solutions?

(Board 2008) 11208139

Q39. Give statement of Le-Chatelier’s principle. (Board 2009) 11208140

Q40. Define pH of a solution. Give its mathematical equation. (Board 2014) 11208141

Q41. What is pK_a and pK_b? (Board 2014)

11208142

Q42. Give the two applications of the solubility product. (Board 2013) 11208143

SOLUTIONS

Objectives

MULTIPLE CHOICE QUESTION

1. Which one of the following is independent of temperature? 11209037

(a) molality (b) molarity

2. The unit of mole fraction is: 11209038

(a) moles/dm³ (b) moles/kg

3. Those solutions which show positive or negative deviations from Raoult’s Law are called: 11209039

(a) ideal solutions

(b) non ideal solutions

(d) saturated solutions

4. Azeotropic mixture cannot be separated into pure components by: 11209040

(a) distillation

(b) fractional distillation

(d) none of the above

5. The normal colligative properties are the properties of: 11209041

(a) dilute solutions which behave ideally

(b) concentrated solutions which behave

nonideally

(d) solutions which deviate from Raoult’s Law

6. Colligative properties are useful in calculating the: 11209042

(a) Solubility of the substance

(b) No. of water molecules of crystallization

(d) Valence of ions

7. Which one of the following gives basic solution in water? 11209043

(a) NaCl (b) Na₂CO₃

8. Depression of freezing point of equimolal aqueous solutions will be maximum for:

(a) Sucrose (b) Glucose 11209044

9. The Depression in freezing point can be measured by: 11209045

(a) Landsberger’s method

(b) Beckmann’s apparatus

(d) None ofthe above

10. The solubility of Ce₂(SO₄)₃: 11209046

(a) is independent of temperature

(b) increases with the increase in temperature

(d) decreases with the decrease in temperature

11. Na/Hg is a solution of the type: 11209047

(a) liquid in liquid (b)liquid in solid

12. The amount of NaOH required toprepare 250cm³ of 0.1M. solution is: 11209048

(a) 1g (b) 10g

13. The pressure at which water boils at 101.5^oC is : 11209049

(a) Slightly more thanoneatm

(b) 760mm Hg

(d) 2.5 atm

14. Both Ebullioscopic and cryoscopic constants depend upon: 11209050

(a) nature of solvent

(b) nature of solute

(d) none of the above

15. In discontinuous solubility curves, the sudden change in direction is due to:

(a) appearance of new phase 11209051

(b) change in no. of molecules of water of crystallization

(d) appearance of new phase and solubility

of substance

16. The molarity of solution containing 1.5g urea in 100cm³ of the solution is:

(a) 1 molar (b) 0.1 molar 11209052

17. The molality of toluene (C₇ H₈) solution in benzene is 0.22.What is the mass of toluene present in 500g of C₆H₆?11209053

(a) 267 (b) 260

18. Which of the following gives acidic solution in water? 11209054

(a) NH₄Cl (b) Na₂SO₄

19. The molality of 2% NaOH solution is approximately: 11209055

(a) 0.5 (b) 0.3

20. A solution consisting of 92 g alcohol (C₂H₅OH), 96 g of methyl alcohol (CH₃OH) and 90 g of water has the mole fraction and mole % of CH₃OH as:

(a) 0.3, 30% (b) 0.2, 30% 11209056

21. Dust particle in air is a solution of type:

(a) Liquid solute and solid solvent11209057

(b) Solid solute and liquid solvent

(d) Gas solute and solid solvent

22. Which of the following is solution in which solvent and solute both are solids? 11209058

(a) Butter (b) Mercury in silver

23. Which one of the following is partially soluble in water? 11209059

(a) NaCl (b) Urea

24. The no. of molecules of water of crystallization in borax (Na₂B₄O₇) are:

(a) 7 (b) 10 11209060

25. Mixture of water and alcohol can be separated by: 11209061

(a) Solvent extraction

(b) Crystallization

(d) Fractional distillation

26. The critical solution temperature of phenol in water is 65.9⁰C. At this temperature the phenol–water percentage is: 11209062

(a) 50% phenol + 5% H₂O

(b) 66% phenol + 34% H₂O

(d) 34% phenol + 66% H₂O

27. Which of the following is non – ideal solution? 11209063

(a) benzene – toluene

(b) ethanol – water

(d) None of the above

28. Which of the following is not a conjugate solution? 11209064

(a) Ether + Water

(b) Phenol + Water

(d) Ethanol + Water

29. Which of the following mixturesexhibits-ve deviation from Raoult’s law and azeotropic mixture with maximum boiling point? 11209065

(a) Acetone + CS₂

(b)Methanol + Benzene

(d)Water + HCl

30. 100cm³ of saturated solutionis evaporated in china dish. The mass of residue is called: 11209066

(a) Azeotropic Mixture

(b)Solubility

(d)Equilibrium constant

31. If more solvent is added to solution, the value of heat of reaction: 11209067

(a) increases

(b) decreases

(d) is affected only when the solution is infinitely diluted

32. When a crystal of solute is added into a supersaturated solution, then: 11209068

(a) the solute dissolves completely

(b) the excess solute crystallizes out

(d) the solution becomes unsaturated

33. A very dilute solution of glucose in water has: 11209069

(a) Free ions in the solution

(b) Free atoms in the solution

(d) Free atoms and molecules

34. When equal volumes of 0.2M AgNO₃ and 0.2M NaCl are mixed together, the conc. of NO ions is: 11209070

(a) 0.2M (b) 0.25M

35. The liquid pair which is not completely miscible is: 11209071

(a) CH₃OH and Water

(b) Alcohol and Water

(d) Benzene and Toluene

36. The solubility of sugar in water is due to: 11209072

(a) High dielectric constant of water

(b) High solvation energy

(d) High dipole moment of water

37. The amount of HCl required to prepare 250cm³ of 0.1M solution is: 11209073

(a) 0.91g (b) 10g

38. Molarity of solution is expressed in:

(a) Moles/kg (b) g dm^-3 11209074

39. The molality of 2% w/v NaOH solution is: 11209075

(a) 2 (b) 0.25

40. The relative lowering of vapour pressure is equal to the mole fraction of the solute. The law is known as: 11209076

(a)Ostwald’s dilution law

(b)Raoult’s law

(d) Henry’s law

41. Vapour pressure of a solution when non-volatile solute is added to a solvent is always: 11209077

(a) abovethe vapour pressure of the pure solvent

(b) equal to the vapour pressure of the solvent

(d) equal to atmospheric pressure

42. The lowest vapour pressure is exerted by: 11209078

(a) Ethanol (b) Methanol

43. The pressure under which liquid and vapour can co-exist at equilibrium is called the: 11209079

(a) Normal vapour pressure

(b) Real vapour pressure

(d) Vapour pressure at boiling point

44. The vapour pressure of a liquid in a closed container depends upon: 11209080

(a) Surface area of the container

(b) Temperature

(d) Nature of non-volatile and non- electrolyte solute

45. Those solutions, which show positive or negative deviations from Raoult’sLaw are called: 11209081

(a) True solutions (b) Non-ideal solutions

46. The mass of glucose required to prepare 1dm³ of 20% glucosesolution is:11209082

(a) 4g (b) 200g (c) 50g (d) 100g

47. Glucose is not soluble in C₆H₆ because:

(a) Glucose is non-polar compound

(b) C₆H₆ is a non polar solvent 11209083

(d) C₆H₆ can make the hydrogen bonding

48. Which one of the following solutions will have higher vapour pressure than that of H₂O? 11209084

(a) H₂O + H₂SO₄ (b) H₂O + Sucrose

49. Which of the following solutions will have the highest boiling point? 11209085

(a) 0.1M NaCl (b) 0.1M CaCl₂

50. The sum of the mole fractions of the components of a solution is equal to:

(a) Zero (b) One 11209086

51. Raoult’s Law is represented by: (Board 2009) 11209087

(a) p = p°X₁ (b) ∆p = p°X₂

52. Molarity of pure water is: 11209088

(a)1 (b) 18 (Board 2013,14,15)

53. 18 g of glucose is dissolved in 90g of water. The relative lowering of vapour pressure is equal to: 11209089

(a) 1/5 (b) 5.1

54. A solution of glucose is 10% w/v. The volume in which 1g mole of it is dissolved will be: (Board 2014)11209090

(a)1dm³ (b) 1.8dm³

55. An aqueous solution of ethanol in water has vapour pressure: 11209091

(a)equal to that of water

(b) equal to that of methanol

(d) less than that of water

56. An azeotropic mixture of two liquids boils at a lower temperature than either of them when: 11209092

(a) It is saturated

(b) It shows positive deviation from Raoult’s law

(d) It is metastable

57. In azeotropic mixture showing positive deviation from Raoult’s law, the volume of the mixture is: 11209093

(a) Slightly more than the total volume of the components

(b) Slightly less than the total volume of the components

(d) None of the above

58. Which of the following solutions has the highest boiling point? 11209094

(a) 5.85% solution of NaCl

(b) 18% solution of glucose

(d) all have the same boiling point

59. Two solutions of NaCl and KCl are prepared separately by dissolving same amount of the solute in water. Which of the following statements is true for these solutions? 11209095

(a)KCl solution will have higher boiling point than NaCl solution

(b) Both the solutions have different boiling points

(d) KCl solution possesses lower freezing point than NaCl solution

60. The molal boiling point constant is the ratio of the elevation in boiling point to:

11209096

(a)molarity

(b) molality

(d) mole fraction of solute

61. Colligative properties are the properties of: 11209097

(a) Dilute solutions which behave as nearly ideal solutions

(b) Concentrated solutions which behave as nearly non-ideal solutions

(d) Neither (a) nor (b)

62. Glycerine decomposes at its: (Board 2014)

11209098

(a) Melting point (b) Boiling point

63.Mist is an example of solution of:11209099

(a) liquid in liquid (b) gas in liquid

64.Solubility of which substance decreases by increasing temperature: 11209100

(a) NaNO₃ (b) KNO₃

65.The term “Cryoscopy" is related to the:

11209101

(a) elevation in boiling point

(b) depression in freezing point

(d) depression in boiling point

66. Benzene –ether can form: 11209102

(a) ideal solution (b) non-ideal solution

67. Butter is a solution of: 11209103

(a) liquid in liquid (b) solid in liquid

68. Which salt when dissolved in water forms a solution with a pH greater than 7: 11209104

(a) NaCl (b) CuSO₄

SHORT QUESTIONS

Q1. Justify that “like dissolves like”.11209105

Q2. Why is glucose not soluble in CCl₄ but dissolves in water? 11209106

Q3. CaCl₂.6H₂O shows discontinuous solubility curve, when plotted against temperature. Why? 11209107

Q4. What is conjugate solution?11209108

Q5. What is the effect of temperature on the conjugate solution of water and phenol?

11209109

Q6. What is consulate temperature or critical solution temperature?11209110

Q7. How does an increase in temperature may increase or decrease the solubility of a substance? 11209111

Q8. What is Raoult’s law? 11209112

Q9. What do you mean by minimum boiling point mixture? 11209113

Q10. What is ebullioscopic constant?11209114

Q11. The lowering of vapour pressure, elevation of boiling point and the depression of freezing points are called colligative properties. Comment upon it.

(Board 2014) 11209115

Q12. Why a non-volatile solute in a volatile solvent lowers the vapour pressure of solution? OR

Why is the vapour pressure of a solution less than pure solvent? (Board 2014) 11209116

Q13. Why is the freezing point of the solution always less than the freezing point of the pure solvent? 11209117

Q14. Colligative properties are obeyed when solutions are dilute. Why? 11209118

Q15. When the heat of solution is negative, then increase in temperature decreases the solubility and vice versa. Why? 11209119

Q16. How are the ions stabilized when a strong electrolyte like NaCl is dissolved in H₂O? 11209120

Q17. Why a salt produced from a strongacid and a weak base gives acidic aqueous solution? 11209121

Q18. How do you justify the given statements? 11209122

(i)Boiling points of solvents increase due to the presence of solutes.(Board 2014)

Q19. Why is sugar not soluble in benzene or petrol etc, but soluble readily in water?

(Board 2014)11209125

Q20. What is the effect of temperature increase on the two layers of phenol and water when these are mixed in equal volumes? 11209126

Q21. How can we say that a solution of two volatile liquids is an ideal solution?11209127

Q22.What is effect of rise in temperature on the solubility of NaCl, KCl and Ce₂(SO₄)₃? 11209128

Q23. What is fractional crystallization? How does it help in removing impurities from a solute? 11209129

Q24.What is advantage of adding ethylene glycol as an antifreeze to the radiator of an automobile? 11209130

Q25. How can we prepare a freezing mixture? 11209131

Q26. Why in CuSO₄.5H₂O, 4H₂O molecules are attached with Cu⁺² cation and one H₂O with ion? 11209132

Q27. Why aqueous Solution of salts NH₄Cl, AlCl₃ and CuSO₄is acidic? 11209133

Q28. Why areaqueous Solutions of Na₂CO₃and CH₃COONa basic? 11209134

Q29. Increasing the temperature increases solubility of glucose in water. Why?11209135

Q30.What will be effect on the position of equilibrium on the following system if,

i.Temperature is increased 11209136

ii.Chlorine is added

PCl_{5 (g)} PCl_{3 (g)}+ Cl_{2 (g)}H = 90kJ/mol

Q31.The sum of mole fraction of all the components is always equal to unity for any solution.Give reason. 11209137

Q32. What are azeotropic mixtures?11209138

Q33. What is water of crystallization? Give examples. (Board 2014) 11209139

Q34. Define parts per million. Give its mathematical expression. 11209140

Q35.What are colligative properties?11209141

Q36. Freezing Point of Solvents is depressed due to presence of solutes. Justify. 11209142

Q37. One molal solution of urea in water is dilute as compared to one molar solution of urea but the number of particles of the solute is same. Justify it. 11209143

Q38. The concentration in terms of molality is independent of temperature but molarity depends upon temperature. Why?11209144

Q39. What are hydrates? How are they formed? Give some examples. 11209145

Q40. Many solutions do not behave ideally. Give reason. (Board 2014) 11209146

Q41. What is meant by molality? Give its formula. (Board 2013) 11209147

ELECTROCHEMISTRY

Objectives

MULTIPLE CHOICE QUESTIONS

1. The passage of electrical current through the metal is due to the reason that 11210030

(a) metal is oxidized

(b) metal is reduced

(d) process of electrolysis takes place

2. A redox reaction is 11210031

(a) Ion combination reaction

(b) Electron transfer reaction

(d) None of the above

3. When one metal is deposited on the surface of the other metal by electric current. Then it is called 11210032

(a) Electrolytic refining

(b) Electrolytic purification

(d) Electroplating

4. Cu metal can be purified in electrolytic cell by making the impure Cu as. 11210033

(a) anode

(b) cathode

(d) SHE

5. The electrode reaction of a voltaic cell can be reversed when 11210034

(a) concentration of solutions is changed

(b) temperature is increased

(d) external circuit is employed to supply the source of electricity

6. When Pb accumulator is recharged, then the density of H₂SO₄ becomes

(a) 2.15 g cm^-3 (b) 1.81 g cm^-311210035

7. The reduction potential of Al is – 1.66 V and that of Sn is ^_0.14V. When these two electrodes are connected through a salt bridge, then which of these electrode act as a cathode? 11210036

(a) Al electrode (b) Sn electrode

8. Which battery is most likely to be used in calculators and digital watches? 11210037

(a) Alkaline battery

(b) Silver oxide battery

(d) Pb – storage battery

9. The oxidation number of Hydrogen in NaH is 11210038

(a) +1 (b) +2

10. The oxidation number of oxygen in H₂O₂ is 11210039

(a) +1 (b) -1

11. The oxidation number of oxygen in F₂O is: (Board 2013, 14) 11210040

(a) +2 (b) -2

12. The metals like Cu, Ag, Au and Pt do not liberate H₂ gas when treated with acid this is because 11210041

(a) metals have low I.P value

(b) their reduction potentials are negative

(d) high +ve value of reduction potential

13. A salt bridge contains 11210042

(a) gelatin + HCl

(b) gelatin + NaOH

(d) gelatin + KCl

14. Same element in different compounds may have 11210043

(a) same oxidation number

(b) Zero oxidation number

(d) cannot be calculated

15. The overall positive values for cell potential predicts that the process is energetically

(a) feasible (b) not feasible 11210044

16. Which one of the following is not an example of voltaic cell? 11210045

(a) Ni – Cd cell (b) Fuel cell

17. When Non-spontaneous redox reaction is carried out by using the electrical current, then the process is called 11210046

(a) Decomposition of the substances

(b) Hydrolysis (c) Electrolysis

(d) Exothermic process

18. The cathode used in alkaline battery is

(a) Cd (b) Zn 11210047

19. In electro-chemical series, the electrodes are compared with SHE and they are arranged in the decreasing order of

(a) Cell voltage 11210048

(b) Ionization potential

(d) Oxidation potential

20. The standard reduction potential of Ag and Zn are +0.80 and –0.76V respectively. Which of the following conclusions can be drawn from the data? 11210049

(a) Ag is a poor Oxidizing agent

(b) Zn has greater tendency than Ag to form positively charged ion

(d) Ag displaces Zn from a solution containing Zn ion.

21. The cathodic reaction in the electrolysis of dilute H₂SO₄ with Pt electrodes is

(a) Reduction 11210050

(b) Oxidation

(d) Neither oxidation nor reduction

22. Which of the following statements is correct about Galvanic cell? 11210051

(a) Anode is negatively charged

(b) Reduction occurs at anode

(d) Oxidation occurs at cathode

23. Stronger the oxidizing agent, greater is the: (Board 2014) 11210052

(a) Oxidation potential

(b) Reduction potential

24. If a salt bridge is not used between two half-cells, then the voltage: 11210053

(a) Decreases rapidly (b) Decreases slowly

25. If a strip of Cu metal is placed in a solution of FeSO₄ 11210054

(a) Cu will be precipitated out

(b) Fe is precipitated out

(d) No reaction takes place

26. Which of the following is not an advantage of Hydrogen-oxygen fuel cell? 11210055

(a) It produces water

(b) It is light weight and potable

(d) It is used in heavy duty automobiles

27. The cell in which electrical energy is converted into chemical energy is called:

11210056

(a) Galvanic cell (b) Electrolytic cell

28. The strong reducing agent is: 11210057

(a) Cl₂ (b) F₂

29. The least value of reduction potential is for: 11210058

(a) Li⁺¹ (b) F₂

30. Electrolysis of mixture of Na₃ AlF₆ and Al₂O₃.H₂O in the fused state using carbon as cathode, the product obtained at cathode is 11210059

(a) Sodium metal

(b) Aluminium metal

(d) The mixture of sodium and aluminium metal

31. In the rusting of iron 4Fe+3O₂®2Fe₂O₃ iron is 11210060

(a) Oxidized (b) Reduced

32. The electrode potential of standard hydrogen electrode is arbitrarily taken as 11210061

(a) Positive (b) Zero

33. Oxidizing power of an element depends upon its 11210062

(a) Oxidation potential

(b) Ionization energy

(d) Electrode potential

34. During the electrolysis of molten NaCl, the ion which is reduced is 11210063

(a) Cl^-1 (b) Na⁺

35. Positive ions are called 11210064

(a) Cations (b) Anions

36. The electrolyte used in fuel cell is (Board 2009) 11210065

(a) Aqueous NaCl (b) Molten NaCl

37. Cathode in NICAD cell is:

(Board 2009) 11210066

(a) Ag₂O (b) NiO₂

38. Voltage of NICAD cell is: 11210067

(a) 1.5 V (b) 1.4 V

39. Voltage of which of the following is about 1.5 V: 11210068

(a) Alkaline battery

(b) Silver oxide battery

(d) Both a and b

40. Which of the following metals is extracted by Hall-Beroult process? 11210069

(a) Al (b) Mg

41. Which of the following processes always involve the decrease in oxidation number?

11210070

(a) Hydrolysis (b) Decomposition

42. Which of the following redox reactions is feasible? 11210071

(a)

(b)

(c)

(d)

43. The electrochemical cell stops working after sometime because: 11210072

(a) The reaction reverse its direction

(b) One of electrode completely vanishes

electrodes equalize

(d) Electrode potentials of both the

electrodes becomes zero

44. In lead storage battery, the anode reaction is: 11210073

(a)

(b)

(c)

(d) None of above

SHORT QUESTIONS

Q1. What is difference between metallic conduction and electrolytic conduction?

(Board 2014) 11210074

Q2. What is the purpose of salt bridge?

(Board 2014) 11210075

Q3. Give the construction of SHE. 11210076

Q4. What are the advantages of fuel cell?

11210077

Q5. Define oxidation number and electrochemistry. 11210078

Q6. What are those elements, which act as oxidizing agents on the basis of electrochemical series? 11210079

Q7. Why metals like Au, Pt, Ag and Cu do not liberate H₂ gas from acids? 11210080

Q8. What are Primary Cells? 11210081

Q9. Give chemical reactions taking place at anode and cathode in a fuel cell.

(Board 2014) 11210082

Q10. Why a salt bridge or porous plate is not required in lead storage battery? 11210083

(Board 2010)

Q11.Why the standard oxidation and reduction potential of Zn is same but with opposite sign? 11210084

Q12.Why Na and K can displace hydrogen from acids but Pt, Pd and Cu cannot?11210085

Q13.Why is the equilibrium set up between metal atoms of electrodes and ions of metals in a cell? 11210086

Q14. What is the reason that a salt bridge maintains the electrical neutrality in the cell? 11210087

Q15. Why lead accumulator is a chargeable battery? 11210088

Q16. How impure copper can be purified by electrolytic process? (Board 2014) 11210089

Q17. Why SHE acts as anode when connected with copper electrode but as cathode with Zn electrode? 11210090

Q18. What is the oxidation state of Mn in KMnO₄? (Board 2014) 11210091

Q19. What are the elements which act as oxidizing agents on the basis of electrochemical series? (L.B. 2004) 11210092

Q20. What is meant by the term electrochemical series? What is the mode of electrode potential? (L.B. 2005) 11210093

Q21. What is that cell in which several kilo watts of power can be generated? What is its other use in addition to generation of power? (Board 2005) 11210094

Q22. What is meant by anodized aluminum? 11210095

Q23. What are secondary cells? 11210096

Q24.Why are the following reaction not spontaneous? Zn + MgSO₄ ® ZnSO₄ + Mg.

11210097

Q25.What is meant by EMF (Electromotive force) of cell? 11210098

Q26. What are the electrolytic products of aqueous solution of NaNO₃? 11210099

Q27. What are the electrode reactions of alkaline battery? 11210100

Q28. How electrochemical series helps to predict the feasibility of a chemical reaction?

11210102

Q29.What is the difference between an electrolytic and voltaic cell or Galvanic cell? (Board 2014) 11210103

Q30.Why in Nelson’s cell Na⁺ ions are not reduced at cathode? 11210104

Q31. What is meant by ionization? 11210105

Q32.What are the oxidation states of oxygen in different compounds? 11210106

Q33.What types of oxidation states are shown by halogens (group VIIA elements)?

11210107

Q34. How oxygen and hydrogen atoms are balanced in ion-electron method? 11210108

Q35.What is meant by electrolytic conduction? 11210109

Q36.What is meant by standard electrode potential? (Board 2014) 11210110

Q37. What are the electrode reactions of lead accumulator during discharging & recharging?

11210111

Q38. How silver oxide battery is prepared?

11210112

Q39. What are the applications of Nickel-cadmium cell? 11210113

Q40. How will you compare the reactivity of metals on the basis of their positions in electrochemical series? 11210114

Q41. What is meant by cathode and anode?

11210115

Q42. How a Galvanic cell reaction is represented? 11210116

Q43. Define Electrolytic conduction and electrolytic cell. (Board 2009) 11210117

Q44. Calculate the oxidation number of chromium in CrCl₃ and . 11210118

Q45. What is standard hydrogen electrode?

11210119

Q46. Differentiate between oxidation and reduction. (Board 2008) 11210120

Q47. How caustic soda can be prepared on industrial scale? 11210121

CHEMICAL KINETICS OR

REACTION KINETICS

Objectives

MULTIPLE CHOICE QUESTIONS

1. In zero order reaction, the rate is independent of: (Board 2013) 11211031

(a) temperature of reaction

(b) concentration of reactants

(d) None of the above

2. If the rate equation for 2A + B ® products is, Rate=k[A]²[B] and A is present in large excess, then order of reaction is: (Board 2014) 11211032

(a) 1 (b) 2

3. The rate of reaction: 11211033

(a) increases as the reaction proceeds

(b) decreases as the reaction proceeds

(d) may decrease or increase as the reaction proceeds

4. With increase of 10°C temperature, the rate of reaction doubles. This increase in rate of reaction is due to: 11211034

(a) decrease in activation energy of reaction

(b) decrease in the number of collisions between reactant molecules

(d) increase in number of effective collisions

5. The unit of the rate constant is same as that of the rate of reaction in: 11211035

(Board 2014)

(a) first order reaction

(b) second order reaction

(d) third order reaction

6. The photochemical reactions: 11211036

(a) are initiated by visible light only

(b) take place at high temperature

(d) require photons to interact with chemical species

7. Velocity constant is the rate of reaction when the concentrations of reactants are: 11211037

(a) zero (b) unity

8. Which of the following reactions is usually slow? 11211038

(a) Neutralization of acids by bases

(b) Organic substitution reactions

(d) Photochemical reactions of CH₄ and Cl₂

9. For a hypothetical reaction A + 2B ® products, the rate law is rate = k[A][B]. the order of reaction is: 11211039

(a) 1 (b) 2

10. The unit of rate constant depends upon:

(a) number of reactants 11211040

(b) concentration terms

(d) molecularity of reaction

11. When a reaction proceeds in more than one steps the overall rate is determined by: 11211041

(a) fastest step

(b) slowest step

(d) any step can be used

12. The hydrolysis of ethyl acetate (CH₃COOC₂H₅) with H₂O in the presence of acid as catalyst is a: 11211042

(a) first order reaction

(b) second order reaction

(d) fractional order reaction

13. The half life period for the decomposition of N₂O₅ is: 11211043

(a) 48 minutes

(b) 24 minutes

(d) 50 minutes

14. According to collision theory of reaction rate, the Arrhenius factor ‘A’ in Arrhenius equation: 11211044

(a) depends upon temperature

(b) depends upon order of reaction

(d) cannot be calculated experimentally

15. The minimum energy more than the average energy required for the molecules to undergo reaction is: 11211045

(a) Internal energy (b) Free energy

16. The effect of temperature on reaction rate is predicted by: 11211046

(a) change in free energy

(b) Arrhenius equation

(d) rate equation

17. A catalyst increases the rate of reaction by: 11211047

(a) reacting with reactants

(b) reacting with products

(d) increasing the activation energy

18. If the energy of the activated complex lies close to energy of reactants, it means that reaction is: 11211048

(a) slow (b) fast

19. A substance which slows down the rate of a reaction is called: 11211049

(a) inhibitor (b) activator

20. One of the best example of auto catalytic reaction is: 11211050

(a) hydrolysis of ethyl acetate

(b) hydrogenation of vegetable oil using Ni-catalyst

(d) reaction between H₂ and I₂ to form HI

21. The enzymes are basically: 11211051

(a) lipids (b) carbohydrates

22. Which of the following factors effect the enzyme activity? 11211052

(a) pH (b) temperature

23. The substance which is formed in a chemical reaction and acts as a catalyst is called: 11211053

(a) retarder (b) auto-catalyst

24. The rate of chemical reaction depends upon the: 11211054

(a) Temperature (b) Catalyst

25. Deactivation of a catalyst by small amount of impurity is called: 11211055

(a) Retardation (b) Poisoning of Catalyst

26. Which statement about k = Ae^-^Ea/RT is incorrect? 11211056

(a) It is Arrhenius equation

(b) It is equation of straight line

(d) Reaction rate increases by increasing temperature

27. Rate of a chemical reaction depends upon: 11211057

(a) The number of fruitful collisions per second

(b) The number of fruitless collisions per second

(d) The number of molecules taking part in a chemical reaction

28. A catalyst is best defined as a substance which increases the speed of chemical reaction and which: 11211058

(a) is usually a non-metal or one of its compounds

(b) is widely used in industry

(d) may be recovered unchanged chemically at the end of reaction

29. Which of the following statements about catalysed and uncatalysed path ways of a reaction is correct? 11211059

(a) DH in catalysed reaction is greater than uncatalysed reaction

(b) DH in catalysed reaction is less than uncatalysed reaction

(d) Activation energy of catalysed reaction is greater than uncatalysed reaction

30. If concentration of reactants is unity or one molar, the rate constant is called:

(a) Equilibrium constant 11211060

(b) Specific rate constant

(d) Average rate

31. With decrease in temperature the rate of reaction decreases. This is due to:

(a) Decrease in activation energy of reactants 11211061

(b) Increase in the number of collisions between reactant molecules

(d) Decrease in the number of effective collisions

32. The rate of reaction between H₂ and Cl₂ is affected by: 11211062

(a) Surface area (b) Light

33. In optical rotation method, the angle through which plane polarized light is rotated by reaction mixture is measured by: 11211063

(a) Spectrometer (b) Refractrometer

34. Rusting of iron, the chemical weathering of stone work of buildings by acidic gases in the atmosphere are the examples of: 11211064

(a) Slow reactions

(b) Very fast reactions

(d) Moderately fast reactions

35. The equation log k = + log A is called: 11211065

(a) Straight line equation

(b) Differential equation

(d) None of the above

36. A substance, which makes the catalyst more effective, is called (Board 2009)

(a) Inhibitor (b) Retarder 11211066

37. Specific rate constant is equal to rate of reaction when concentration of reactants is: 11211067

(a) zero (b) four

38. A second order rate constant can have the units: 11211068

(a) dm^–6mol²s^–1 (b) dm³mol s^–1

39. Under a given set of experimental conditions, with increase of concentration of the reactants, the rate of a chemical reaction: 11211069

(a) Always decreases

(b) Always increases

(d) None of these

40. Which of the following statements is not correct? 11211070

(a) Molecularity of a reaction cannot be fractional

(b) Molecularity of a reaction cannot be more than three

obtained from balanced chemical

equation

(d) Molecularity of a reaction may or may not be equal to the order of the reaction

41. Which of the following is a zero order reaction: 11211071

(a)

(b)

(c)

(d)

42. The Chemical reactions in which reactants require high amount of activation energy are generally: 11211072

(a) Slow (b) Fast

43. The value of activation energy is primarily determined by: 11211073

(a) Temperature

(b) Collision frequency

(d) Chemical nature of reactants and

products

44. Which of the following statements is false? 11211074

(a) Fast reactions have low activation energy

(b) Activation energy of a reaction depends on the chemical nature of reactants and products

(d) A catalyst increases the rate of reaction by decreasing the activation energy of the reaction.

45. Graph of lnk vs has slope equal to: 11211075

(a) (b)

46. of first order reactions is given by would be equal to: 11211076

(a) (b)

47. A catalyst accelerates the rate of a reaction by: 11211077

(a) Destabilising the reactants

(b) Stabilising the product

(d) Lowering the energy of activation for

the reverse reaction

SHORT QUESTIONS

Q1. The rate of a chemical reaction is an ever changing parameter under the given conditions. Comment upon the statement.

11211078

Q2. The reaction rate decreases every moment but the rate constant “k” of the reaction is a constant quantity, under the given conditions. Justify it. 11211079

Q3. 50% of a hypothetical first order reaction complete in one hour. The remaining 50% needs more than one hour to convert itself into products. Explain. 11211080

Q4. The radioactive decay is always a first order reaction. Why? (Board 2014) 11211081

Q5. The units of rate constant of second order reaction is dm³ mol^-¹ sec^-¹ but the unit of rate of reaction is mole dm^-3 sec^-1. Explain. 11211082

Q6. The sum of coefficients of a balanced chemical equation is not necessarily important to give the order of reaction. Why? 11211083

Q7. The order of reaction is obtained from the rate expression of a reaction and the rate expression is obtained from the experiments. Explain. 11211084

Q8. How is the rate of reaction determined for a reaction involving ions? 11211085

Q9. What kind of graph is obtained when it is plotted between = on x – axis and log k on y-axis? 11211086

Q10. What is homogeneous catalysis? 11211087

Q11.What happens to the rate of a chemical reaction with the passage of time? 11211088

Q12. How is the rate of acid hydrolysis of ethyl acetate ester determined chemically?

11211089

Q13. Under what conditions, zero order reaction does take place? 11211090

Q14. What is poisoning of catalyst? 11211091

Q15. How does the higher temperature increase the rate of a chemical reaction? 11211092

Q16. What is the difference between instantaneous and average rates of chemical reaction? 11211093

Q17. How will you compare instantaneous and average rate of reactions? 11211094

Q18. What is meant by specific rate constant or velocity constant? (Board 2005) 11211095

Q19. What chemical method is used to determine the rate of reaction of hydrolysis of ethyl acetate (ester)? 11211096

Q20. What are fractional order reactions? 11211097

Q21. What is meant by pseudo first order reactions? 11211098

Q22. What is meant by half life period? 11211099

Q23. What is rate determining step or rate limiting step of a chemical reaction and reaction intermediate? 11211100

Q24. What is molecularity of a chemical reaction? 11211101

Q25.How spectrometry and refractrometric methods are used to determine the rate of chemical reactions? 11211102

Q26.Define activation energy and activated complex. (Board 2014) 11211103

Q27. What is the difference between effective and ineffective collisions? 11211104

Q28. What are the various methods for finding the order of reactions? (Board 2010) 11211105

Q29. How the nature of reactants affect the rates of reactions? 11211106

Q30. How concentration of reactants affect the rate of chemical reaction? 11211107

Q31. What is the effect of surface area on reaction rate? 11211108

Q32. What is heterogeneous catalysis?

(Board 2014) 11211109

Q33. What is the effect of catalyst on equilibrium constant (K_c of a reaction)?

11211110

Q34. What is meant by autocatalyst? 11211111

Q35. How is the catalytic activity of enzyme enhanced? What are the controlling factors of the activity of enzyme? 11211112

Q36. What is catalysis? Give an example.

(Board-2008) 11211113

Q37. A finely divided catalyst may prove more affective. Give reason.(Board 2009)11211114

Q38. What are enzymes? Give two examples in which enzymes act as catalyst. 11211115

Q39. Define order of reaction with the help of an example.

Causative

11th Chemistry Full Book Objectives + SHORT QUESTIONS

MULTIPLE CHOICE QUESTIONS

SHORT QUESTIONS

MULTIPLE CHOICE QUESTIONS

SHORT QUESTIONS

MULTIPLE CHOICE QUESTIONS

MULTIPLE CHOICE QUESTIONS

MULTIPLE CHOICE QUESTIONS

SHORT QUESTIONS

Q40. Why is atomic spectrum, a line spectrum? 11205137

Q41. What is Aufbau Principle? 11205138

Q42. Why an electron moves faster in an orbit of smaller radius? (Board 2013) 11205139

Q43. Calculate the mass of electron when its e/m value is 1.7588 ´ 10¹¹ Ckg^-1.

(Board 2014) 11205140

Q44. Write down the electronic configuration of Fe(26) and Br(35).

(Board 2014) 11205141

Q45. Define spectrum. Give its two types. (Board 2009) 11205142

Q46. How we come to know that cathode rays are material particles with negative charge? (Board 2007) 11205143

Q47. What are defects in Rutherford’s atomic model? (Board 2015) 11205144

Q48. How will you relate energy of emitted light with its frequency and wavelength? 11205145

Q49. State Hund’s rule and give one example. (Board 2007, 2015) 11205146

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11th Chemistry Full Book Objectives + SHORT QUESTIONS

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Q40. Why is atomic spectrum, a line spectrum? 11205137

Q41. What is Aufbau Principle? 11205138

Q42. Why an electron moves faster in an orbit of smaller radius? (Board 2013) 11205139

Q43. Calculate the mass of electron when its e/m value is 1.7588 ´ 1011 Ckg-1.

(Board 2014) 11205140

Q44. Write down the electronic configuration of Fe(26) and Br(35).

(Board 2014) 11205141

Q45. Define spectrum. Give its two types. (Board 2009) 11205142

Q46. How we come to know that cathode rays are material particles with negative charge? (Board 2007) 11205143

Q47. What are defects in Rutherford’s atomic model? (Board 2015) 11205144

Q48. How will you relate energy of emitted light with its frequency and wavelength? 11205145

Q49. State Hund’s rule and give one example. (Board 2007, 2015) 11205146

SHORT QUESTIONS

MULTIPLE CHOICE QUESTION

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MULTIPLE CHOICE QUESTIONS

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Q43. Calculate the mass of electron when its e/m value is 1.7588 ´ 10¹¹ Ckg^-1.