Numericals - Chemistry XI (Five Year Papers)

1. Simplify according to the rule of significant figure .

2. The atomic mass of Zn is 65.4 a.m.u. Calculate (i) the number of moles and also the number of atoms in 10.9 gm of Zn. (ii) The mass of 1.204 x 1024 atoms of Zn in gm.

3. Adipic acid is used in the manufacture of Nylon. The acid contains 49.3%C, 6.9%H and 43.6%O by mass. The molecular mass of the acid is 146 a.m.u. Find the molecular formula of the Adipic Acid.

4. Calculate the value of R (Gas constant) with the help of Gas Equation when (i) the pressure is in atmosphere and the volume in dm3 or litre. (ii) the pressure is in Nm-2 and the volume is in cubic metre.

5. 400cm3 of helium gas effuse from a porous container in 20 seconds. How long will SO2 gas take to effuse from the same container? (Atomic Weight = S = 32, He = 4).

6. A system absorbs 200J of heat from the surroundings and does 120 J of work on the surroundings by expansions. Find the internal energy change of the system.

7. 1.2 gm of acetic acid (CH3COOH) is dissolved in water to make 200cm3 of the solution. Find the concentration of the solution in Molarity.

8. The solubility of calcium oxalate (CaC2O4) is 0.0016 g/dm3 at 25°C. Find the solubility product of calcium oxalate: CaC2O4 ® Ca2+ + C2O42-

9. Calculate H+ ion concentration of a solution whose pH = 5.6.

10. The rate constant (k) for the decomposition of nitrogen dioxide 2NO2(g) ® 2NO(g) + O2(g) is 1.8 x 103- dm3mole1-sec1-. Write down the rate expression and (i) find the initial rate when the initial concentration of NO2 is 0.75 M. (ii) Find the rate constant (k) when the initial concentration of NO2 is doubled.

11. Calculate the volume of nitrogen gas produced by heating 800 gm of ammonia at 21°C and 823 torr pressure. 2NH3 ® N2 + 3H2 (Atomic Weight = N = 14, H = 1)

12. In collection of 24 x 1025 molecules of C2H5OH. What is the number of moles. ( Atomic weight = C = 12, O = 16, H = 1)

13. Simplify using exponential notation: 43100 + 3900 + 2100.

14. A given compound contains 75. 2% carbon, 10.75% hydrogen and 14.05% oxygen. Calculate the empirical formula of the compound. (Atomic weight: C = 12, O = 16, H = 1)

15. Calculate the wave number of spectral line of hydrogen gas when an electron jumps from n = 4 to n= 2. (RH = 109678 cm-1)

16. 13.2 gm of gas occupies a volume of 0.918 dm3 at 25°C and 8 atm pressure. Calculate the molecular mass of the gas.

17. Calculate the heat of formation of benzene at 25°C when the heat of formation of CO2 and water and heat of combustion of benzene are given:

(i)


6C + 3H2 ® C6H6


DHf = ?

(ii)


C + O2 ® CO2


DH = -286KJ/mole

(iii)


H2 + ½O2 ® H2O


DH = -286KJ/mole

(iv)


C6H6 + 7.5O2 ® 6CO2 + 3H2O


DH = -3267 KJ/mole

18. The rate constant for the decomposition of nitrogen dioxide is 1.8 x 10-8 dm3 mole-1s-1. What is the initial rate when the initial concentration of NO2 is 0.50M? 2NO2 ® 2NO + O2.

19. Should AgCl precipitate from a solution prepared by mixing 400cm3 of 0.1M NaCl and 600cm3 of 0.03 M of solution of AgNO3? (Ksp for AgCl = 1.6 x 10-10 mole/dm3)

20. A sample of chlorine gas at S.T.P has a volume of 800cm3 calculate The number of moles of chlorine, the mass of the sample and the number of chlorine molecules in the sample.

21. How many atoms of carbon are present in 10 gm of coke?

22. The volume of the oxygen gas, collected over water at 24°C and 762mm pressure, is 128 ml. Calculate the mass in gm of oxygen gas obtained. The pressure of water vapour at 24°C is 22 mm.

23. Calculate the radius of orbit n = 3 for a Hydrogen atom in Armstrong unit. (h = 6.625 x 10-27 erg-sec, p = 3.14, m = 9.11 x 10-28gm, e = 4.8 x 10-10 esu)

24. For the reaction H2 + I2 ® 2HI. Kc is 49. Calculate the concentration of HI at equilibrium when initially one mole of H2 is mixed with one mole of I2 in one litre flask.

25. Determine the mass of HCl required to prepare 400 ml of 0.85M HCl solution.

26. Calculate pH value of 0.004M NaOH solution.

27. Kc for the reaction is 0.0194 and the calculated ratio of the concentration of the reactants and the product is 0.0116. Predict the direction of the reaction.

28. For the decomposition of ethyl chlorocarbonate ClCOOC2H5 ® CO2 + Cl.C2H5. Find the value of rate constant when initial concentration of Ethyl Chlorocarbonate is 0.25 M and the initial rate of the reaction is 3.25 x 10-4 mole/dm3/sec.

29. 1.0 gm of a sample of an organic substance was burnt in excess of oxygen yield 3.03 gm of CO2 and 1.55 gm of H2O. If the molecular mass of the compound is 58. Find the molecular formula.

30. Calculate the volume of the oxygen at S.T.P that may be obtained by complete decomposition of 51.3 gm of KClO3 on heating in presence of MnO2 as a catalyst. 2KClO3 ® 2KCl + 3O2. (Atomic mass of K = 39, Cl = 35.5, O = 16, Mn = 55)

31. Calculate the wave number of the Line in Lyman Series when an electron jumps from orbit 3 to orbit 1.

32. Calculate the heat of formation of ethane (C2H6) at 25°C from the following data:

(i)


2C + 3H2 ® C2H6


DHf = ?

(ii)


C + O2 ® CO2


DH = -394KJ/mole

(iii)


H2 + ½O2 ® H2O


DH = -286KJ/mole

(iv)


C2H6 + ½O2 ® 2CO2 + 3H2O


DH = -1560.632KJ/mole

33. At the equilibrium a 12 litre flask contains 0.21 mole of PCl5, 0.32 mole of Cl2 at 250°C. Find the value of Kc for the reaction. PCl5 Û PCl3 + Cl2.

34. A given compound contains C = 60%, H = 13.0% and O = 27%. Calculate its Empirical Formula.

35. How many grams of chlorine are required to prepare 7.75 dm3 of chloro benzene? The equation of the reaction is C6H6 + Cl2 ® C6H5Cl + HCl. (Atomic Number of C = 12, H = 1 and Cl = 35.5)

36. A mixture of helium and hydrogen is confined in a 12 dm3 flask at 30°C. If 0.2 mole of the helium is present, find out the partial pressure of each gas whereas the pressure of the mixture of gases is 2atm.

37. Calculate the radius by hydrogen atom by applying Bohr’s Theory. (h = 6.625 x 10-27 erg-sec, p = 3.14, m = 9.11 x 10-28gm, e = 4.8 x 10-10 esu)

38. Calculate the heat of formation of C2H2 from carbon and hydrogen from the following data:

(i)


2C + H2 ® C2H2


DHf = ?

(ii)


C + O2 ® CO2


DH = -94.05Kcal/mole

(iii)


H2 + ½O2 ® H2O


DH = -68.32Kcal/mole

(iv)


C2H2 + 5/2O2 ® 2CO2 + H2O


DH = -310Kcal/mole

39. Calculate the pH of a 2.356 x 10-3m HCl solution.

40. For the reaction N2 + 3H2 Û 2NH3. The equilibrium mixture contains 0.25 M nitrogen, 0.15M hydrogen gas at 25°C. Calculate the concentration of NH3 gas when Kc = 9.6. the volume of the container is 1dm3.

41. Determine the initial rate of the following reaction at 303°C in which its rate constant is 8.5 x 10-5 litre-mol-1 sec-1. Initial concentration of the reaction is 9.8 x 10-2 mole/litre. 2NO2 ® 2NO + O2.

Extra Numericals

1. 4.6gm of ethyl alcohol and 6.0gm of acetic acid kept at constant temperature until equilibrium was established. 2 gm of acid were present unused. Calculate Kc.

2. Kc for the dissociation of HI at 350°C is 0.01. If 0.2 mole of H2, 1.3 moles of I2 and 4 moles of HI are present. Predict the direction of reaction.

3. What is the solubility of PbCrO4 at 30°C when Ksp is 1.8 x 10-14.

4. 1.06m of an organic compound on combustion gave 1.49 gm of CO2 and 0.763gm H2O. It also has 23.73% N. Find its compercial formula.

5. 500 dm3 of moist O2 gas was collected over water at 27°C and 726torr pressure. Find the mass in gm. Of dry O2 gas at S.T.P. When the vapour pressure of water 27°C is 26 torr.

6. Atomic mass of phosphorus is 31. Calculate the mass of 45 atoms in a.m.u.

7. Methane burn in steam according the following reaction: CH4 + 2O2 ® CO2 + 2H2O. If 100 gm of each CH4 and O2 is taken, then what amount of CO2 liberated?

8. An organic compound containing C = 65.45%, H = 5.45% and O = 29.09%. If molecular weight of compound is 110, calculate molecular formula.

9. What mass of CO2 is produced by the complete combustion of 100g pentane. C5H12 + 3O2 ® 2CO2 + 2H2O.

10. One atom of an unknown element is found to have a mass of 67.8 x 10-23g. What is the atomic weight of the element?

11. The heat of combustion of glucose and alcohol is given below.

(i)


C6H12O6 + 6O2 ® 6CO2 + 6H2O


DH = -673Kcal/mole

(ii)


C2H5OH+ 3O2 ® 2CO2 + 3H2O


DH = -328Kcal/mole

Find DH for the fermentation given below:



C6H12O6 ® 2C2H5OH + 3CO2



12. At certain temperature, the equilibrium mixture contain 0.4 mole of H2, 0.4mole I2 and 1 mole of HI. If addition 2 mole of H2 are added. How many moles of HI will be present when the new equilibrium established. H2 + I2 ® 2HI.

13. A solution has pH of 8.4. Find concentration of H+ and OH-.

14. 180cm3 of a known gas diffuse in 15minutes, when 120 cm3 of SO2 diffuses in 20 minutes. What is the molecular mass of the unknown gas.

Chapter 1

Introduction to Fundamental Concepts

1. Calculate the moles of the following in 500gm, NH3, HCl, Na2CO3, H2SO4, MgBr2, CaCO3, Xe and C.

2. How many moles of Na are present in 5gm of Na?

3. Calculate the number of atoms in 12 gms of Mg.

4. 2gm diamond is studded in a ring. Diamond is a pure carbon. How many atoms of carbon are present in the ring?

5. Calculate the number of molecules in 9gms of H2O.

6. How many molecules are present in 25 gms of CaCO3?

7. Calculate the weight in gram of 3.01 x 1020 molecules of glucose (C6H12O6)

8. How many atoms of hydrogen are there in 2.57 x 10-6 gram of hydrogen?

9. A sample of oxygen contains 1.87 x 1027 atoms of oxygen. What would be the weight of the oxygen?

10. Find the weight of oxygen obtained from 49gm of KClO3.

2KClO3 ® 2KCl + 3O2

11. What weight of CO2 and CaO can be obtained by heating 12.5gm of Limestone (CaCO2)?

CaCO3 ® CaO + CO2

12. Calculate the weight of sodium chloride required to produce 142 gm of chlorine.

2NaCl ® 2Na + Cl2

13. Calculate the weight of carbon, required to produce 88gm of CO2.

C + O2 ® CO2

14. The action of CO on Fe2O3 can be represented by the following equation.

Fe2O3 + 3CO ® 2Fe + 3CO2

15. What weight of NH3 will be required to produce 100 gm of NO?

4NH3 + 5O2 ® 4NO + 6H2O

16. Find out the moles of CuSO4 which are obtained from 31.75 gm of Cu.

Cu + H2SO2 ® CuSO2 + H2

17. Calculate the number of N2 and H2 molecules, which are obtained from 8.5 gm of NH3.

N2 + 3H2 ® 2NH3

18. Find out the number of Cu and H2O molecules obtained from 7.95gm of CuO.

CuO + H2 ® Cu + H2O

19. 400gm of H2 was made to combine with 14200gm of Cl2. How much HCl will be produced?

20. 1kg of Limestone was heated 500gm of CaO was obtained. How much CO2 gas produced into air.

21. Find the weight of O2 obtained from 49 gm of KClO3.

2KClO3 ® 2KCl + 3O2

22. Chlorine is produced on the large scale by the electrolysis of NaCl aqueous solution. Chlorine the weight of NaCl required to produce 142 gm of Cl2.

2NaCl + 2H2O ® Cl2 + H2 + 2NaOH

23. How many grams of O2 are required to completely burn 18.0gm of C? How many grams of CO2 will be formed?

24. Calculate the weight of NH3, required to produce 100 gms of NO.

4NH3 + 5O2 ® 4NO + 6H2O

25. Find out the moles of H2 and N2 required producing 17gm of NH3.

26. Calculate the volume of H2 at S.T.P, which is obtained by the reaction of 120 gm Mg with MgSO4.

Mg + H2SO4 ® MgSO2 + H2

27. NH3 gas can be produced from ammonium chloride (NH4Cl) as follows:

CaO + 2NH4Cl ® CaCl2 + H2O + NH3

Calculate the volume of NH3 obtained at S.T.P by the reaction of 100 gm of NH4Cl.

28. 500gm of C2H4 on combustion in air gave CO2 and H2O. Calculate the volume of O2 and CO2 at S.T.P.

29. Find out the volume of O2, CO2 and SO2 gases at S.T.P react and obtained from 2 moles of CS2.

CS2 + 3O2 ® CO2 + 2SO2

30. Calculate the volume of CO2 gas at S.T.P obtained by the combustion of 20gm of CH4.

CH4 + 2O2 ® CO2 + 2H2O

31. Calculate the volume of O2 gas at S.T.P required to burn 600dm3 of H2S, also find the volume of SO2 gas produced at S.T.P.

32. Calculate the volume of O2 gas at S.T.P required to burn 50 gm of CH4.

33. What volume of H2 at S.T.P can be produced by the reaction of 6.54gm Zn with HCl?

Zn + 2HCl ® ZnCl2 + 2H2

34. Calculate the volume of O2 and H2 gases at S.T.P obtained from 9gm of H2O.

35. 0.264gm of Mg was burnt in pure O2. How much MgO will be formed?

2Mg + O2 ® 2MgO

36. How much H2 can be generated by passing 200gm of steam over hot iron.

4H2O + 3Fe ® Fe3O4 + 4H2

37. If 112dm3 of N2 react with 336 dm3 of H2, both at S.T.P. How many grams of NH3 would be obtained?

N2 + 3H2 ® 2NH3

38. An organic compound contains 12.8%C, 2.1% and 85.1% Br. If the mass of the compound is 188, find the molecular formula.

39. An organic compound contains 66.70%C, 7.41% H and 25.90% N2. The molecular mass of the compound is 108. Find out its molecular formula.

40. A compound contains 19.8%C, 2.5%H, 66.1%O and 11.6%N. Find out empirical formula of the compound.

41. 0.2475gm of a compound, containing C, H and O gave 0.4950gm CO2 and 0.2025gm H2O. If the molecular mass of the compound is 88. Find out the molecular formula.

42. An organic compound contains 32%C, 6.67%H, 18.66%N and 42.67%O. Its molecular mass is 75. Find out the molecular formula of the compound.

43. 1.367gm of a compound containing C, H and O on heating gave 3.002gm CO2 and 1.640gm H2O. Find out its molecular formula, when the molecular mass is 120.

44. A compound was found to contain 40%C and 6.7%H. Its molecular mass was 60. Find out its molecular formula.

45. An organic compound contains 75.2%C, 10.15%H and oxygen. Its molecular mass is 115. Find its molecular formula.

46. The empirical formula of a compound is CH2O. If the molecular mass 180. Find out the molecular formula.

47. An organic compound composed of C, H and O. On combustion of 0.94gm of this compound, 1.32gm CO2 and 0.568gm H2O were obtained. Its molecular mass is 180. Find its molecular formula.

48. An organic compound composed of C, H and O. 4.2gm of the compound on heating gave 6.21gm CO2 and 2.54gmH2O. Its molecular mass is 60. Find its molecular formula.

49. An organic compound contains C,H and 6.38gm of compound on combustion gave 9.06gm CO2 and 5.58gm H2O. Its molecular mass is 62. Find out its molecular formula.

50. 1gm of a hydrocarbon on combustion gave 3.03gm of CO2 and 1.55gm of H2O. If the molecular mass is 58, find its molecular formula.

51. 1.434gm of a compound on combustion gave 4.444gm CO2 and 2.0 gm H2O. Find out its empirical formula.

52. An organic compound composed of C, H and N. 0.225gm of compound on combustion gave 0.44gm CO2 and 0.315gm H2O. If the molecular mass of a compound is 90, find out its molecular formula.

53. An organic compound contains 40.68%C, 8.47%H, 23.73%N and 27.12%O. Find its empirical formula.

54. An organic compound composed of C, H and N. 0.419 gm of compound on combustion gave 0.88gm CO2 and 0.27gm H2O. Find out its empirical formula.

55. The analysis of a compound shows, C = 24.24%, H = 4.04% and Cl = 71.71%. If the molecular mass of the compound is 49.5, find its molecular formula.

56. An organic compound of molecular mass 90 has the empirical formula CH2O. What is its molecular formula?

57. The empirical formula of an organic compound is CH3NO2. If it’s molecular mass is 61. What is its molecular formula?

58. 0.638gm of an organic compound on combustion gave 0.594gm H2O and 1.452gm CO2.The compound is composed of C, H and O atoms. If the molecular mass is 116, find out its molecular formula.

59. The molecular formula of ethyl acetate is CH3COOC2H5. What is its empirical formula.

60. Find the empirical formulae of the following compounds from their percentage composition by mass:

· N = 26.17% H = 7.48% Cl = 66.35%

· Ca = 71.43% O = 28.57%

· Ag = 63.53% N = 8.23% O = 28.24%

· Na = 32.40% H = 45.07% Cl = 22.53%

61. A certain compound on analysis yielded 2.00gm C, 0.34gm H and 2.67gm O. If the relative molecular mass of the compound is 60, calculate its molecular formula.

62. What is the empirical formula of a compound, which contains 42.5% chlorine and 57.5 oxygen. If it’s formula mass is 167. What is its molecular formula?

63. What will be the weight of 5 moles of water in grams?

64. What is the mass of each of the following:

· 1.25 mole of NaCl

· 2.42 mole of NaNO3

· 1.5 mole of HCl

· 3.0 mole of NaOH

65. A piece of Aluminium metal weighs 70.0g. How many atoms are present in the piece.

66. How many atoms of carbon are present in 20-carat Diamond? (1 carat = 0.2g)

67. How many grams of oxygen have the same number of atoms as 16gm of sulphur?

68. A sample of oxygen gas at STP has a mass of 16gm. Calculate:

· The number of moles of oxygen

· The volume of the sample

· The number o molecules in the sample

69. Calculate the volume of CH4 gas at STP having a mass 32g.

70. What mass of zinc sulphate can be obtained from the reaction of 10.0gm of Zinc with an excess of dilute H2SO4?

Zn + H2SO4 ® ZnSO4 + H2*

71. Calculate what mass of sodium hydroxide you would need to neutralize a solution containing 7.3g hydrogen chloride by the reaction:

NaOH + HCl ® NaCl + H2O

72. Calculate how much sodium nitrate you need to give 126g of nitric acid by the reaction:

NaNO3 + H2SO4 ® HNO3 + NaHSO4

73. What volume of hydrogen at STP is evolved when 0.325g of zinc reacts will dilute hydrochloric acid.

Zn + 2HCl ® ZnCl2 + H2

74. What mass of oxygen is formed by the decomposition of a solution containing 120cm3 of H2O2 at STP?

2H2O2 ® 2H2O + O2

75. What is the mass of one molecule of water in grams?

76. 100cm3 of butane are burned in an excess of oxygen. Calculate:

· The volume of oxygen used

· The mass and volume of CO2 produced (assume all gases at STP)

2C4H10 + 13O2 ® 8CO2 + 10H2O

77. A cook is making a small cake. It needs 500cm3 at STP of CO2 to make the cake rise. The cook decides to add baking powder, which contains sodium bicarbonate. This generates CO2 by thermal decomposition.

2NaHCO3 ® CO2 + Na2CO3 + H2O

What mass of baking powder must the cook add to cake mixture?

78. What volume of ammonia at STP can be obtained by heating 0.25 mole of ammonium sulphate with calcium hydroxide?

(NH4)SO4 + Ca(OH)2 ® 2NH3 + CaSO4 + 2H2O

79. How many grams of SO2 are produced when 100g of H2S is reacted with 50g of oxygen.

2H2S + 3O2 ® 2H2O + 2SO2

80. How many grams of chlorobenzene will be produced when 100gm of each reactant is reacted?

C6H6 + Cl2 ® C6H5Cl + HCl

81. A car releases about 5g of NO into the air for each mile driven. How many molecules of NO are emitted per mile?

82. Simplify according to the rule of significant figures.

· 2.60 x 3.05

· 0.009 ¸ 0.3

·

·

Chapter 2

The Three States of Matter

1. 540cm3 of N2 at 400mm pressure are compressed to 300cm3 without changing the temperature. What will be the pressure of the gas?

2. A gas occupies 6dm3 at 1atm pressure keeping the temperature constant. If the pressure reduces to 600mm, what volume does the gas occupy?

3. At a certain temperature and 800mm pressure, the volume of H2 is 700cm3. If the pressure is increased to 1000mm at the same temperature, find the new volume of the gas.

4. 150ml of a gas at 27°C is heated to 77°C at constant pressure. Find the new volume of the gas.

5. 300ml of N2 are at 50° and the pressure is kept constant. If the temperature is doubled, what will be the volume of the gas?

6. A gas measures 5dm3 at 5°C under 0.5atm pressure. Calculate its volume at 25° and 5000mm pressure.

7. 2060ml of a gas is at 7°C and 860mm pressure. Find its volume at S.T,P.

8. 350ml of H2 was collected over water at 26°C. The pressure of the gas was 900mm. What volume will dry gas have at 30°C and 750mm pressure? The vapour pressure at 26°C is 25mm.

9. The volume of oxygen collected over water at 20°C and 1200mm pressure, is 200cm3. If aqueous at 20°C is 17.4mm, what will be the volume of the gas under S.T.P.

10. A 20dm3 flask contains H2 at 22°C under pressure of 1.2 atm. How many moles of H2 are present.

11. A gaseous mixture is at the pressure of 3000mm. The mixture contains 6 moles of N2, 0.5mole of CO2 and 2.5 moles of O2. Find the partial pressure of each gas.

12. A 5dm3 vessel contains 1.2 moles of H2 and 0.8 mole of N2 at 27°C. Find the total pressure of the mixture.

13. Composition of a sample of air by volume is, N2 = 76%, O2 = 20%, H2O = 2.5%, CO2 = 1.4% and He = 0.1%. If the pressure of the air is 760 mm, Calculate the partial pressure of these gases.

14. A 10dm3 container contains a mixture of He and Ne gases at 17°C. There are two moles of He gas and 3 moles of Ne gas. What is the partial pressure of the gases?

15. 10gm of H2, 96gm of O2 and 196gm of N2 are mixed together. The partial pressure of H2 is 0.6 atm. What is the partial pressure of O2 and N2?

16. A cylinder contains 1 mole of H2, 3 mole of He and 6 moles of N2. The total pressure in the cylinder is 15 atm. Calculate the partial pressure of H2, He and N2.

Chapter 5

Energetics of Chemical Reaction

1. Calculate the heat of formation of Acetic Acid from the following data:

(i)


2C + 2H2+ O2 ® CH3COOH


DHf = ?

(ii)


C + O2 ® CO2


DH = -394KJ/mole

(iii)


H2 + ½O2 ® H2O


DH = -286 KJ/mole

(iv)


CH3COOH + 2O2 ® 2CO2 + 2H2O


DH = -870KJ/mole

2. Calculate the heat of formation of Ethane from the following data:

(i)


2C + 3H2 ® C2H6


DHf = ?

(ii)


C + O2 ® CO2


DH = -394KJ/mole

(iii)


H2 + ½O2 ® H2O


DH = -286 KJ/mole

(iv)


C2H6 + 7/2O2 ® 2CO2 + 3H2O


DH = -1560KJ/mole

(v)


C2H5OH + 3O2 ® 2CO2 + 3H2O


DH = -327 KJ/mole

3. Calculate the heat of formation of Methane from the following data:

(i)


C + 2H2 ® CH4


DHf = ?

(ii)


C + O2 ® CO2


DH = -394KJ/mole

(iii)


H2 + ½O2 ® H2O


DH = -286 KJ/mole

(iv)


CH4 + 2O2 ® CO2 + 2H2O


DH = -890.3KJ/mole

4. Calculate the heat of formation of Ethyl Alcohol from the following data:

(i)


2C + 3H2 ½ O2® C2H5OH


DHf = ?

(ii)


C + O2 ® CO2


DH = -394KJ/mole

(iii)


H2 + ½O2 ® H2O


DH = -286 KJ/mole

(iv)


C2H5OH+ 3O2 ® 2CO2 + 3H2O


DH = -1369KJ/mole

5. Calculate the heat of formation of Ethane from the following data:

(i)


C2H6 + 7/2O2 ® 2CO2 + 3H2O


DHf = ?

(ii)


C + O2 ® CO2


DH = -394KJ/mole

(iii)


H2 + ½O2 ® H2O


DH = -286 KJ/mole

(iv)


C2H6 ® 2C + 3H2


DH = -84.68KJ/mole

6. Calculate the heat of formation of Methane from the following data:

(i)


C + 2H2 ® CH4


DHf = ?

(ii)


C + O2 ® CO2


DH = -94.1cal

(iii)


H2 + ½O2 ® H2O


DH = -68.3 cal

(iv)


CH4 + 2O2 ® CO2 + 2H2O


DH = -212.8 cal

7. Calculate the heat of formation of Ethene from the following data:

(i)


2C + 2H2 ® C2H4


DHf = ?

(ii)


C + O2 ® CO2


DH = -97kcal

(iii)


H2 + ½O2 ® H2O


DH = -65 kcal

(iv)


C2H4 + 3O2® 2CO2 + 2H2O


DH = 340 kcal

8. Calculate the heat of formation from the following data:

(i)


2C + 3H2 +1/2O2 ® C2H5O


DHf = ?

(ii)


C + O2 ® CO2


DH = -94.2Kcal/mole

(iii)


H2 + ½O2 ® H2O


DH = -68.5 Kcal/mole

9. Calculate the heat of formation of from the following data:

(i)


C + 2H2 + O2® CH3OH


DHf = ?

(ii)


C + O2 ® CO2


DH = -94.2Kcal/mole

(iii)


H2 + ½O2 ® H2O


DH = -68.32 Kcal/mole

(iv)


CH3OH + O2 ® CO2 + 2H2O


DH = -347.6Kcal/mole

10. Calculate the heat of formation of from the following data:

(i)


3C + 4H2 ® C3H8


DHf = ?

(ii)


C + O2 ® CO2


DH = -94.1Kcal/mole

(iii)


H2 + ½O2 ® H2O


DH = -68.3 Kcal/mole

(iv)


C3H8 + 5O2 ® 3CO2 + 4H2O


DH = -530.7Kcal/mole

11. Calculate the heat of formation of from the following data:

(i)


H2 + O2® H2O2


DHf = ?

(ii)


H2 + ½O2 ® H2O


DH = -68.32Kcal

(iii)


H2O + ½ O2 ® H2O2


DH = -23.48Kcal

12. Given:

(i)


NH3 + HCl ® NH4Cl


DH1 = 42.100Kcal

(ii)


H2O + ½ O2 ® H2O2


DH2 = 3.900cal

Find DH for the reaction,



NH3 + HCl ® NH4Cl


DHf = ?

Chapter 6

Chemical Equilibrium

1. 1.5 moles of acetic acid and 1.5 moles of ethyl alcohol were reacted at a certain temperature. At equilibrium, 1 mole of ethyl acetate was present in 1 litre of the equilibrium mixture. Calculate the equilibrium constant Kc.

CH3COOH + C2H5OH Û CH3COOC2H5 + H2O

2. 6.0 gm of hydrogen and 1016gm of iodine were heated in a sealed tube at a temperature, at which Kc is 50. The volume of the tube is 1 dm3. Calculate the concentration of HI.

H2 + I2 Û 2HI

3. At a certain temperature, an equilibrium mixture contains 0.4 mole H2, 0.4 mole I2 and 1 mole of HI. The volume of the reacting vessel is 4 dm3. Find out the equilibrium constant kc.

H2 + I2 Û 2HI

4. 3 moles of A and 2 moles of B are mixed in a 4dm3 flask, at a certain temperature. The following reaction occurs.

3A + 2B Û 4C

At equilibrium the flask contains 1 mole of B. Find the equilibrium constant kc.

5. At a certain temperature, 0.205 mole of H2 and 0.319 mole of I2 were reacted. The equilibrium mixture contains 0.314 mole of I2. Calculate the kc.

H2 + I2 Û 2HI

6. The kc for the reaction A + B Û C + D is 1/3. How many moles of A must be mixed with 3 moles of B to yield at equilibrium, 2 moles of C and D each. The volume of the vessel is 2 litre.

7. At a certain temperature the equilibrium mixture for the reaction A + B Û 2C, contains 2 moles A, 3 moles of B and 5 moles of C. Find the Kc for the reaction.

8. For the reaction 2A Û B + C, equilibrium constant kc is 1. If we start with 6 moles of A, how many moles of B will be formed.

9. 20 moles of SO2 and 10 moles of O2 are taken in a 20 litre flask. If at equilibrium 5 moles of SO3 are formed, Calculate kc.

2SO2 + O2 Û 2SO3

10. A quantity of PCl5 was heated in a 12 dm3 vessel at 250°C.

PCl5 Û PCl3 + Cl2

11. 2 moles of HI was introduced in a vessel held at constant temperature. When equilibrium was reached, it was found that 0.1 mole of I2 have been formed. Calculate the equilibrium constant.

H2 + I2 Û 2HI

12. When 1 mole of pure C2H5OH is mixed with 1 mole of CH3COOH at room temperature, the equilibrium mixture contains 2/3 moles of ester and water each.

· What will be the kc?

· How many moles of ester are formed at equilibrium when 3 moles of C2H5OH are mixed with 1 mole of CH3COOH?

CH3COOH + C2H5OH Û CH3COOC2H5 + H2O

13. PCl5 Û PCl3 + Cl2. Calculate the number of moles of Cl2 produced at equilibrium when 1 mole of PCl5 is heated at 250°C in a vessel having capacity of 10dm3. At 250°C, Kc is 0.041.

14. When 2.94 moles of iodine and 8.1 moles of Hydrogen were mixed and heated at 444°C and at constant volume, until the equilibrium was established. 5.64 moles of HI were formed. Calculate the value of kc.

H2 + I2 Û 2HI

15. What is the solubility of lead chromate in moles/dm3 at 25°C. The solubility product is 1.8 x 10-14.

PbCrO4 Û Pb++ + CrO4--

16. The solubility of Mg(OH)2 at 25°C is 0.00764 gm/dm3. What is the solubility product of Mg(OH)2?

Mg(OH)2 Û Mg++ + 2OH-

17. Find the solubility of AgCl in gm/dm3, when the solubility product is 1.25 x 10-10.

18. Calculate the solubility product of BaSO4. The solubility of the salt is 1.0 x 10-5 moles/dm3.

19. Calculate the solubility product of BaSO4 is 9.0 x 10-3 gm/dm3. Find its solubility product.

20. Predict whether there will be any precipitate formation by mixing 30cm3 of 0.01M NaCl with 60cm3 of 0.01M AgNO3 solution. Ksp of AgCl is 1.5 x 10-10.

21. A saturated solution of calcium fluoride was found to contain 0.0168 gm/dm3 of solute at 25°C. Calculate the ksp for CaF2.

22. A saturated solution of BaF2 at 25°C is 0.006M. Calculate Ksp of the salt.

1st Year Chemistry - Fill in the Blanks

1. The property of a crystal, which is different in different directions, is called __________.

2. 0.00051 contains __________ significant figures.

3. The oxidation number of oxygen in OF2 is __________.

4. The volume of 1 gm of hydrogen gas at S.T.P is __________.

5. The oxidation number Mn in KMnO4 is __________.

6. The product of ionic concentration in a saturated solution is called __________.

7. 16 gm of oxygen at S.T.P occupies a volume of __________ dm3.

8. The shape of the orbital for which l = 0 is __________.

9. The radius of Cl-1 is __________ than the radius of Cl0.

10. Sp2 hybridization is also known as __________.

11. The value of 1 Debye is __________.

12. The reactions catalyzed by sunlight are called __________.

13. The blue colour of CuSO4 is due to the presence of __________.

14. The force of attraction between the liquid molecules and the surface of container is called __________.

15. The heat of neutralization of a strong acid and a strong base is __________.

16. C º C triple bond is __________. C = C double bond length.

17. The ions having the same electronic configuration are called iso electronic.

18. On heating, if a solid changes directly into vapours without changing into the liquid state, the phenomenon is called __________.

19. Each orbital in an atom can be completely described by __________.

20. In a molecule of alkene, __________ restricts the rotation of the group of atoms at either end of the molecule.

21. Density, refractive index and vapour pressure are __________ properties.

22. The addition of HCl to H2 solution __________ the ionization of H2S.

23. The reaction of cation or anion (or both) with water so as to change its __________ is known as Hydrolysis.

24. A reaction with higher activation energy will start at __________ temperature.

25. 6.02 x 1023 has __________ significant figures.

26. The internal resistance in the flow of liquid is called __________.

27. A catalyst increases the velocity of a reaction but decreases the __________.

Chapter 1

Introduction to Fundamental Concepts

1. 1 mole of a gas at S.T.P occupies a volume of __________.

2. A gas occupying a volume of 22.4 dm3 at S.T.P contains __________ molecules.

3. A formula, which gives the relative number of atoms in the molecule of a compound, is called __________.

4. A formula which gives the actual number of all kinds of atoms present in the molecule of compound is termed as __________.

5. The chemical formula that not only gives the actual number of atoms but also shows the arrangement of different atoms present in the molecule is called __________.

6. Atomic weight or molecular weight expressed in grams is known as __________.

7. 2 moles of H2O contain __________ grams and __________ number of molecules.

8. Any thing that occupies space and has __________ is called matter.

9. Volume of one __________ mole of a gas at S.T.P is 22.4 cubic feet.

10. A ton mole of iron is equal to __________ tons.

11. The force with which the earth attracts a body is called the __________ of the body.

12. A pure substance contains __________ kind of molecules.

13. The smallest indivisible particle of matter is called __________.

14. The atomic number is equal to the number of __________ in nucleus.

15. The atomic mass is the total number of protons and __________ in an atom of the element.

16. The average weight of atoms of an element as compared to the weight of one atom of __________ is called the atomic mass.

17. 1.0007 contains __________ significant figures.

18. The figure 24.75 will be rounded off to __________.

19. __________ means that the readings and measurements obtained in different experiments are very close to each other.

20. __________ means that the results obtained in different experiments are very close to the accepted values.

21. The degree of a measured quantity __________ with increasing number of significant figures in it.

22. The atomic mass of sodium is __________.

23. The symbolic representation of a molecule of a compound is called __________.

24. Molecular formula of CHCl3 and its Empirical formula is __________.

25. Molecular formula of benzene is C6H6 and its empirical formula is __________.

26. 58.5 is the __________ of NaCl.

27. 4.5 gms of nitrogen will have __________ molecules.

28. 28 gms of nitrogen will have __________ molecules.

29. 2 moles of SO2 is equal to __________ gms.

30. 1000 gms of H2O is equal to __________ moles.

31. The reactions, which proceed in both directions, are called __________.

32. The reactions, which proceed in forward directions only, are called __________ reactions.

33. The __________ reactions are completed after some time.

34. 0.0006 has __________ significant figures

35. 7.40 x 108 has __________ significant figures.

36. 7 x 108 has __________ significant figures.

37. Usually Molecular formula is simple multiple of the __________.

38. 0.1 mole of H2O contains __________ molecules of H2O.

39. Mass of 3.01 x 1022 molecules of CO2 is __________.

40. __________ is the branch of science which deals with the properties, composition and structure of matter.

41. None zero digits are all __________.

42. The integer part of logarithm is called __________.

43. The decimal fraction of logarithm is called __________.

44. __________ is the amount of substance, which contains as many number of particles as there are in 12 gms of Carbon.

45. 6.02 x 1023 is called the __________.

46. The accuracy of measurement depends on the number of __________.

47. __________ is the branch of chemistry that deals with quantitative relationships among the substances undergoing chemical changes.

48. The sum of atomic weights of all the elements present in molecular formula is called the __________.

49. __________ is the sum of atomic weights of the elements represented by the Empirical formula of the compound.

50. Very small and very large quantities are expressed in terms of __________.

51. In rounding off __________ figure is dropped.

52. Mole is the quantity, which has __________ particle of the substance.

53. For three significant figures, 25.55 is rounded off to __________.

54. The S.I unit of a mass is __________.

55. Mass of 6.02 x 1023 molecules of NaCl is __________ gm.

56. 1 mole of NaOH is __________ gm of NaOH.

57. Formula weight is used for __________ substances.

58. The word S.I stands for __________.

59. 4.5 gms of water will have __________ molecules.

60. 0.0087 has __________ significant figure.

Chapter 2

The Three States of Matter

1. The intermixing of gases or liquids in a container irrespective of their densities, is called __________.

2. At constant temperature, if the pressure of a given mass of a gas is decreased, its volume will __________.

3. A volume of __________ dm3 will hold 128 gms of SO2.

4. At constant temperature of a given mass of a gas, the product of its __________ and __________ is constant.

5. The rates of diffusion of gases are __________ proportional to the square root of their densities.

6. Gases deviate from ideal behaviour more markedly at high __________.

7. Liquid diffuse __________ than gases.

8. An imaginary line passing through the centre of a crystal is called __________.

9. The temperature at which more than one crystalline forms of a substance coexist in equilibrium is called __________.

10. Two or more substances crystallizing in the same form is called __________.

11. The existence of solid substances in more than one crystalline form is known as __________.

12. Rate of diffusion of gases is __________ as compared to liquids.

13. Boiling point of a liquid __________ with the pressure.

14. Mercury in a glass tube forms __________ curvature.

15. Gases can be compressed to __________ extent.

16. Viscosity of a liquid __________ with the increase of temperature.

17. Surface tension of water __________ by adding soap solution into it.

18. The internal resistance to the flow of a liquid is called __________.

19. The rise or the fall of a liquid in a capillary tube is called __________.

20. Matter exists in __________ states.

21. The freezing point of water in Fahrenheit scale is __________.

22. Boiling point of water is __________ °K.

23. SI unit for measurement of pressure is __________.

24. The value of gas law constant R = __________ dm3 atm/°K/mole.

25. The absolute Zero is equal to __________.

26. If P is plotted against 1/V at constant temperature a __________ is obtained.

27. Gases __________ in heating.

28. The pressure of air __________ at higher altitude.

29. Standard temperature means __________.

30. Standard pressure means __________.

31. Cooling is caused by __________ of gases.

32. Rate of diffusion of O2 is __________ times more than H2.

33. H2O has __________ viscosity than CH3OH.

34. Mercury does not wet the glass surface due to its higher __________.

35. Surface tension of mercury is __________ than water.

36. Viscosity can be easily measured by an instrument called __________.

37. The pressure exerted by the vapours when these vapours are in equilibrium with the liquid is called __________.

38. Vapour pressure __________ at high temperature.

39. Boyle’s Law and Charles Law can be combined into the mathematical expression __________.

40. Equal volumes of all gases at the same temperature and pressure contain __________ number of molecules.

41. The average Kinetic energy of a gas is proportional to its __________ temperature.

42. Kinetic equation may be mathematically written as __________.

43. The temperature at which two crystalline forms of a substance can coexist in equilibrium is called __________.

44. Lighter gases diffuse __________ than heavier gases.

45. Rain drops are __________ in shape.

46. Due to surface tension, the surface area of the liquid is __________.

47. Water __________ in the capillary tube.

48. Viscosity of a solution at 10°C is __________ than at 20°C.

49. Shape of NaCl crystal is __________.

50. Gases intermix to form a mixture.

51. Pressure of a dry gas is __________ than the pressure of a moist gas.

52. 22.4 dm3 of nitrogen at S.T.P will weigh equal to __________ gm.

53. 1 mole of any gas at S.T.P is equal to __________ dm3.

54. At -273°C, volume of all gases becomes __________.

55. The gases, which strictly follows the gas Laws are called __________ gas.

56. __________ is the property that determines the direction of flow of heat.

57. __________ is defined as force per unit area.

58. __________ viscosity is defined as the viscosity of a liquid as compared to the water.

Chapter 3

Structure of Atom

1. The maximum number of electrons in 2p orbital is __________.

2. 3d orbital has __________ energy than 4s orbital.

3. __________ rays are non-material in nature.

4. Charge to mass ratio of cathode rays resembles to that of __________.

5. __________ rays are most penetrating.

6. Neutrons have mass equal to that of __________.

7. Energy is __________ when an electron jumps from higher to lower orbit.

8. Second Ionization Potential has __________ value than the First Ionization Potential.

9. Electronegativity __________ from left to right in a period of Periodic Table.

10. __________ was discovered during the course of Artificial Radioactivity.

11. The velocity of alpha rays is nearly __________ of velocity of light.

12. Natural Radioactivity is confined in __________ elements.

13. The isotopes of an element differ in their __________.

14. Two electrons with the __________ spin, can never occupy the same atomic orbital.

15. ‘Al’ has electronic configuration, 1s2, 2s2, __________.

16. In a group of Periodic Table, the ionization potential __________ from top to bottom as the size of atom increases.

17. Ionization potential values __________ from left to right in a period.

18. The energy required to remove the most loosely bond electron from an atom in gaseous state is called __________.

19. The SI unit of Ionization Potential is __________.

20. An atom of sodium possesses 11 protons and __________ neutrons.

21. The particles of Cathode rays possess __________ charge.

22. The negatively charged particles found in Cathode rays are named as __________.

23. Positive rays are emitted from __________.

24. __________ rays are also known as Canal rays.

25. __________ consists of helium ions and are doubly positively charged.

26. __________ rays consists of negatively charged particles.

27. __________ rays are light waves of very short wavelength.

28. The phenomenon in which a stable element is made radioactive by artificial disintegration is called __________.

29. The electron move around the nucleus in different circular paths called __________.

30. The maximum number of electron in a shell is determined by the formula __________.

31. A particle whose mass is equal to that of electron but carries a positive charge is called __________.

32. 2p electrons are __________ in energy that 2s electrons in the same atom.

33. Number of protons of an element also indicates its __________.

34. According to __________ Principle electrons are fed in the order of increasing orbital energy.

35. According to __________ electrons are distributed among the orbitals of a sub shell to give maximum number of unpaired electron and have same spin.

36. The specific way in which the orbitals of an atom are occupied by electrons is called __________.

37. __________ rays are stream of doubly positively charged particles.

38. Electron in the outer most shell of an atom is called __________.

39. Protons are found in the __________ of an atom and bear __________ charge.

40. The atomic number of an atom is the sum of __________ inside the nucleus.

41. __________ limits the number of electron to different shell or orbits.

42. Sir William Crookes in 1878, discovered that the cathode in high vacuum tube emit radiations what he called __________.

43. X-rays were discovered in 1895 by __________.

44. The discovery of proton was done in 1886 by __________.

45. Neutrons were discovered by __________ in 1932 by the bombardment of beryllium with alpha particles.

46. Each atom has a __________, which contains all the positive charge and practically all the mass of atom.

47. Complete the reaction: 4Be9 + 2H4 ® __________ + __________.

48. __________ have higher ionization power as compared to b-rays.

49. No dark spaces between the colours are present in __________.

50. The symbol e+ represents __________.

51. p-orbitals are __________ shaped.

52. The energy released when an electron is added to an atom in the gaseous state is called __________.

53. The power of an atom to attract a shared pair of electrons towards itself is called __________.

54. Fluorine is __________ electronegative than chlorine.

55. Lyman series of spherical lines appear in the __________ portion of spectrum.

56. The electrons with __________ spin occupy the same orbital.

57. 3d orbital has __________ energy than 4s orbital.

58. Energy and frequency are __________ proportional to each other.

59. Ionic radii of cations are __________ than the atoms from which they are formed.

60. Ionic radii of anions are __________ than the atoms from which they are formed.

Chapter 4

Chemical Bonding

1. A bond formed due to transference of electron is called __________.

2. A bond formed due to sharing of electron is called __________.

3. Sigma bond is __________ than pi bond.

4. The shape of methane molecule is __________.

5. One s and 3p orbitals overlap to produce four __________ hybrid orbitals.

6. Ethene, C2H4 is an example of __________ hybridization.

7. Water molecule has __________ structure.

8. Water molecules are inter-linked with one another due to __________.

9. Polarity of the molecule is due to the difference of __________ between the two bonded atoms.

10. A chemical bond formed between to different atoms by mutual sharing of electron is termed as __________.

11. A chemical bond formed between two similar atoms by mutual sharing of electrons is known as __________.

12. The difference between the Electronegativity values of the two atoms forming covalent bond must be __________ than 1.7.

13. When two orbitals of different atoms by hybridize with each other having their axes in the same straight lines, the bond formed is termed as __________.

14. __________ bond is formed when p-orbitals of the two atoms with their axes parallel to each other overlap with each other.

15. Melting and boiling point of ionic compounds are usually __________ than that of covalent compounds.

16. Non polar compounds are usually __________ in non polar solvent.

17. The nitrogen in NH3 is __________ hybridized.

18. A hybrid orbital is called __________ orbital.

19. Since dipole moment of CS2 is zero, it is a __________ molecule.

20. A bond formed due to the electrostatic forces of attraction between the oppositely charged ions is called __________ bond.

21. The ionic bond is formed between the atoms with low ionization potential and high __________.

22. A bond formed by the sharing of an electron pair contributed by one atom only is called a __________ bond.

23. A co-ordinate covalent bond is also known as __________ bond.

24. Polar covalent bond is __________ than a non polar covalent bond.

25. H-F bond is __________ than H-Br bond.

26. The SI unit of dipole moment is __________.

27. Commonly used unit of dipole moment is __________.

28. Dipole moment of non-polar compound is __________ D.

29. The reactions of ionic compounds are usually very __________.

30. Covalent compounds are generally __________ in nature.

31. Ionic compounds are generally __________ in nature.

32. A covalent bond is represented by a __________.

33. A co-ordinate covalent bond is represented by an __________.

34. The covalent bond between H-F is called __________ covalent bond.

35. The power of an atom to attract a shared pair of electron itself is called __________ of that atom.

36. m = d x e represents __________.

37. CO2 and SO2 molecules have __________ polar bonds.

38. NH3 molecule has __________ polar bonds.

39. A double bond has __________ bond energy than a single bond.

40. An orbital which surrounds a single nucleus is called __________ orbital.

41. An orbital which surrounds two or more atomic nuclei is called __________ orbital.

42. A molecular orbital, which is of lower energy than the atomic orbitals from which it is derived, is known as __________ orbital.

43. A molecular orbital, which has higher energy than the atomic orbitals from which it is derived, is known as __________ orbital.

44. Orbitals formed after hybridization are called __________ orbitals.

45. Bond angle in Sp3 hybridization is of __________.

46. Bond angle in Sp2 hybridization is of __________.

47. Bond angle in Sp hybridization is of __________.

48. Sp3 hybridization is also known as __________.

49. Sp2 hybridization is also known as __________.

50. Sp hybridization is also known as __________.

51. A pair of electrons residing on the central atom and which is not used in bonding is called a __________.

52. The sum of total number of electron pairs (bonding and lone pairs) is called __________ number.

53. __________ bond is usually expressed by dotted line.

54. Water molecule has dipole moment because of its __________ structure.

55. CO2 is non polar because of its __________ structure.

56. Overlapping in __________ bond is perfect.

57. Overlapping in __________ bond is not perfect.

58. H-H bond is __________ than H-Cl bond.

59. __________ hybrid orbitals are not co-planar.

60. Covalent bond is Cl2 molecule is __________.

Chapter 5

Energetics of Chemical Reaction

1. The branch of Chemistry, which deals with the heat changes that take place during chemical reaction, is called __________.

2. The branch of science which deals with energy changes accompanying physical and chemical transformation is called __________.

3. The amount of heat evolved or absorbed in a chemical reaction is called __________.

4. Such reactions in which heat is evolved are called __________ reactions.

5. Such reactions in which heat is absorbed are called __________ reactions.

6. In exothermic reactions, heat evolved is given by __________ sign of DH.

7. In endothermic reactions heat absorbed is given by __________ sign of DH.

8. The total heat change in a reaction is the same whether it takes place in one or several steps.

9. The first law of thermodynamics is also known as __________.

10. The part of universe under observation is called __________.

11. The system plus its surrounding is called __________.

12. Such properties, which give description of a system at a particular moment, is called __________.

13. The term E + PV is called __________.

14. DH represents change in __________.

15. The temperature of water is raised up when sulphuric acid is added to it. This is an __________ reaction.

16. The characteristic properties of a system which is independent of amount of material concerned is called __________ properties.

17. The characteristic properties of a system which depend on amount of substance present in it is called __________ properties.

18. Density, pressure and temperature are the examples of __________ properties.

19. Mole numbers and enthalpy are the examples of __________ properties.

20. A system, which exchange both energy and matter with its surrounding, is called __________ system.

21. A system, which only exchange energy with the surrounding but not matter is, called __________ system.

22. A system which neither exchange energy nor matter with its surrounding is called __________ system.

23. A system is __________ if it contains only one phase.

24. A system is __________ if it contains more than one phase.

25. 1 kilojoule is equal to __________ joules.

26. 1 Calorie is equal to __________ joules.

27. 1 kilo calorie is equal to __________ joules.

28. The work done (w) is mathematically denoted by __________.

29. The change in enthalpy is denoted by __________.

30. __________ law is used in calculating heat of reaction.

31. __________ is defined as the change in enthalpy when one gram mole of a compound is produced from its element.

32. Heat of formation is denoted by __________.

33. When the work is done on the system by the surrounding the sign of work done (w) is __________.

34. When the work is done by the system on surrounding the sign of work done is __________.

35. First law of Thermodynamics is mathematically represented as __________.

36. Standard enthalpies are measured at __________.

37. Hess’s Law is employed to calculate __________ of a chemical reaction.

38. Heat absorbed by the system at constant volume is completely utilize to increase the __________ of the system.

39. Heat change at constant pressure from initial to final state of the system is simply equal to the __________.

40. SI unit of measurement of heat change is __________.

Chapter 6

Chemical Equilibrium

1. The reactions, which proceed in both the directions, are called __________ reactions.

2. The reactions, which proceed to one direction only, are called __________ reactions.

3. Reversible reactions are __________ completed.

4. Irreversible reactions are __________ after some time.

5. A reversible reaction is said to be in __________ when the rate of forward reaction becomes equal to the rate of backward reaction.

6. The concentrations of reactants and products are __________ at equilibrium point.

7. The value of Kc depends upon the __________ of the reactants.

8. A increase of the value of Kc tends to move the reaction to the __________ direction.

9. A decrease of the value of Kc tends to move the reaction to the __________ direction.

10. An increase in the concentration of the reactants will move the reaction to the __________ direction.

11. A decrease in the concentration of the reactants will move the reaction to the __________ direction.

12. Equilibrium constant is denoted by __________.

13. When the equilibrium constant value is very __________, we can conclude that the forward reaction is almost completed.

14. When equilibrium constant value is very __________ we can conclude that forward reaction will occur to very little extent.

15. According to __________ principle, if system in equilibrium is subjected to a stress, the equilibrium shifts in a direction to minimize or undo the effect of the stress.

16. In exothermic reaction, the __________ of temperature favour the forward rate of reaction.

17. In endothermic reactions, the __________ of temperature favour the forward rate of reaction.

18. A __________ is a substance which effects the rate of reaction but remains unaltered at the end of the reaction.

19. A catalyst increases the velocity of the reaction by decreasing the __________.

20. The suppression of degree of ionization of a sparingly soluble weak electrolyte by the addition of a strong electrolyte containing an ion in common is called __________.

21. __________ is purified in industries by Common Ion Effect.

22. A reaction moves to the left when the concentrations of the products are __________.

23. A reaction moves to the right when the concentrations of the products are __________.

24. Increase in pressure will move the reaction in the direction of __________ volume.

25. Decrease in pressure will move the reaction in the direction of __________ volume.

26. An increase of temperature favours the formation of products in case of __________ reaction.

27. A decrease of temperature fovours the formation of products in case of __________ reaction.

28. Heating moves an endothermic reaction to the __________.

29. Cooling move an exothermic reaction to the __________.

30. The product of ionic concentration in a saturated solution is called __________ constant.

31. When HCl is added to NaCl, the concentration of __________ ion is increased.

32. Chemical reaction involving the substances in more than one phases are called __________.

33. The formation of NH3 is exothermic process hence __________ temperature will favour the formation of NH3.

34. The formation of NO from N2 and O2 is endothermic process hence __________ temperature will favour the formation of NO.

35. Chemical Equilibrium is __________ equilibrium.

36. Molar concentration is also called __________.

37. The rate at which a substance takes part in a chemical reaction depends upon its __________.

38. __________ principle is applied to all reversible reaction.

39. A common ion __________ the solubility of the salt.

40. Number of moles present per dm3 of a substance is called __________.

Chapter 7

Solutions and Electrolytes

1. A mixture of two or more substances, which are homogeneously mixed, is called a __________.

2. __________ is defined as the amount of solute dissolved in a given amount of solvent.

3. A solution is composed of two components __________ and __________.

4. A solution containing one mole of solute per dm3 of solution is called one __________ solution.

5. Molarity is denoted by __________.

6. 1M solution of NaOH contains __________ gms of it dissolved per dm3 of solution.

7. A solution containing one mole of solute dissolved by per kg of solvent is called __________ solution.

8. Molality is denoted by __________.

9. 1M solution of H2SO4 contains __________ gms of it per kg of solvent.

10. The process in which ions are surrounded by water molecules is called __________.

11. The water molecules attached with the hydrated substance are called __________.

12. Hydrated copper sulphate evolves __________ water molecules on heating.

13. The interaction between salt and water to produce acids and bases is called __________.

14. The products of ionic concentration in a saturated solution at a certain temperature are called the __________.

15. Solubility product constant expressed as __________.

16. The suppression of ionization by adding a common ion is called __________.

17. The process of dissociation of an electrolyte into ions is known as __________.

18. The chemical decomposition of a compound in a solution or in fused state brought about by a flow of electric current is known as __________.

19. Electrolysis is performed in an electrolytic cell, which is known as __________.

20. The positive electrode of a voltmeter is called __________ and negative as __________.

21. A solution, which tends to resist changes in pH is called a __________ solution.

22. A mixture of acetic acid and sodium acetate acts as a __________.

23. According to Sorenson __________ is defined as negative logarithm of the hydrogen ion concentration.

24. pH is mathematically expressed as __________.

25. The pH of a neutral solution is __________.

26. __________ substances have pH values lower than 7.

27. __________ solutions have pH values more than 7.

28. Oxidation is __________ of electron.

29. Reduction is the __________ of electron.

30. Such chemical reactions in which the oxidation number of atoms or ions is changed are called __________ reactions.

31. Oxidation number of a free element is __________.

32. Oxidation number of Oxygen in a compound is __________.

33. The sum of oxidation number of any formula of a compound is __________.

34. The oxidation number of any ion is equal to the __________ on the ion.

35. __________ is the reaction in which an acid reacts with a base to form salt and water.

36. __________ are organic compounds which change colour in accordance with the pH of the medium.

37. An indicator that changes from colourless to pink in the presence of an alkaline solution is called __________.

38. An indicator that changes from red to yellow in the presence of an alkaline solution is called __________.

39. Dissociation constant is denoted by __________.

40. According to Bronsted-Lowry Concept, __________ is the donor of proton and __________ is the acceptor of proton.

41. According to Arrhenius, acid is substance that produces __________ ions when dissolved in water.

42. According to Arrhenius, base is a substance that produces __________ ions when dissolved in water.

43. When ionic product is less than ksp, the solution will __________.

44. When ionic product is greater than ksp, the solution will __________.

45. The electrode at which oxidation takes place is called __________.

46. The electrode at which reduction takes place is called __________.

47. H3O+ ion is called __________ ion.

48. The logarithm of reciprocal of hydroxyl ion (OH)- is called __________.

49. Aqueous solution of NH4Cl is __________ while that of NaHCO3 is __________.

50. The ionic product of [H+] and [OH-] of pure water is __________.

51. An increase in the oxidation number of an element or ion during a chemical change is called __________.

52. A decrease in the oxidation number of an element or ion during a chemical change is called __________.

53. The degree of dissociation __________ with the increase in temperature.

54. The degree of dissociation __________ with the dilution of electrolytic solution.

55. A __________ consists of an electrode immersed in solution of its ion.

56. The potential difference between the electrode and the solution of its salt at equilibrium position is called __________ potential.

57. If the pH of a solution is 14, the solution is __________.

58. If the pH of a solution is 4, the solution is __________.

59. The oxidation number of Mn in KMnO4 is __________.

60. The oxidation number of Fe in FeCl3 is __________.

Chapter 8

Introduction to Chemical Kinetics

1. The branch of chemistry, which deals with the study of rates and mechanisms of chemical reactions, is known as __________.

2. Such reactions, which proceeds with very high velocities and are completed very quickly are called __________ reactions.

3. Such reactions, which take place very slowly, are called __________ reactions.

4. Reactions between silver nitrate and sodium chloride to form white precipitates of silver chloride are an example of __________ reaction.

5. Reactions of Organic compounds are slow and are called __________ reactions.

6. There are some reactions, which proceed slowly with a __________ speed.

7. The rate of __________ reaction can only be determined.

8. The amount of chemical change taking place in concentration of the per unit time is called __________ of reaction.

9. Rate of reaction is expressed in __________.

10. The rate of reaction between two specific interval of time is called __________.

11. The addition energy required to bring about a chemical reaction is called __________.

12. According to __________ theory for a chemical reaction to take place, the reacting molecules must come closed together.

13. The addition of __________ helps the reaction by lowering the energy of activation.

14. The rate of reaction __________ with the increase in concentration of the reacting molecules.

15. When the concentration of both the reacting molecules is double, the probability of collisions between them will be __________ times.

16. By __________ the surface area of the reactants, the rate of reaction is increased.

17. Rate of reaction generally __________ with the rise of temperature.

18. A __________ is a substance, which either accelerates or retards the rate of reaction without taking part in the reaction.

19. In the preparation of Oxygen from Potassium Chlorate, __________ is used as catalyst.

20. In the oxidation of SO2 to SO3 by the contact process for the manufacture of H2SO4 __________ is used as catalyst.

21. An unstable intermediate compound formed during a chemical reaction is called __________.

22. When a catalyst and the reactants are in the same phases, it is known as __________ catalyst.

23. When a catalyst and the reactants are in different phases, it is called __________.

24. When a catalyst increases the rate of reaction, it is called __________ catalyst.

25. When a catalyst retards the rate of reaction, it is called __________ catalyst.

26. A negative catalyst __________ the energy of activation, hence the rate of reaction is decreased.

27. The ratio between the rate of reaction and concentration of reactants is known as __________.

28. Velocity constant is independent of concentration but depends on __________.

29. Ionic reactions are __________ than molecular reactions.

30. The value of specific rates constant for a reaction __________ with time.

31. The sum of all exponents of concentration terms in the equation is called __________.

32. The sum of moles taking part in a chemical reaction is called __________ of the reactions.1. The property of a crystal, which is different in different directions, is called __________.

2. 0.00051 contains __________ significant figures.

3. The oxidation number of oxygen in OF2 is __________.

4. The volume of 1 gm of hydrogen gas at S.T.P is __________.

5. The oxidation number Mn in KMnO4 is __________.

6. The product of ionic concentration in a saturated solution is called __________.

7. 16 gm of oxygen at S.T.P occupies a volume of __________ dm3.

8. The shape of the orbital for which l = 0 is __________.

9. The radius of Cl-1 is __________ than the radius of Cl0.

10. Sp2 hybridization is also known as __________.

11. The value of 1 Debye is __________.

12. The reactions catalyzed by sunlight are called __________.

13. The blue colour of CuSO4 is due to the presence of __________.

14. The force of attraction between the liquid molecules and the surface of container is called __________.

15. The heat of neutralization of a strong acid and a strong base is __________.

16. C º C triple bond is __________. C = C double bond length.

17. The ions having the same electronic configuration are called iso electronic.

18. On heating, if a solid changes directly into vapours without changing into the liquid state, the phenomenon is called __________.

19. Each orbital in an atom can be completely described by __________.

20. In a molecule of alkene, __________ restricts the rotation of the group of atoms at either end of the molecule.

21. Density, refractive index and vapour pressure are __________ properties.

22. The addition of HCl to H2 solution __________ the ionization of H2S.

23. The reaction of cation or anion (or both) with water so as to change its __________ is known as Hydrolysis.

24. A reaction with higher activation energy will start at __________ temperature.

25. 6.02 x 1023 has __________ significant figures.

26. The internal resistance in the flow of liquid is called __________.

27. A catalyst increases the velocity of a reaction but decreases the __________.

Chapter 1

Introduction to Fundamental Concepts

1. 1 mole of a gas at S.T.P occupies a volume of __________.

2. A gas occupying a volume of 22.4 dm3 at S.T.P contains __________ molecules.

3. A formula, which gives the relative number of atoms in the molecule of a compound, is called __________.

4. A formula which gives the actual number of all kinds of atoms present in the molecule of compound is termed as __________.

5. The chemical formula that not only gives the actual number of atoms but also shows the arrangement of different atoms present in the molecule is called __________.

6. Atomic weight or molecular weight expressed in grams is known as __________.

7. 2 moles of H2O contain __________ grams and __________ number of molecules.

8. Any thing that occupies space and has __________ is called matter.

9. Volume of one __________ mole of a gas at S.T.P is 22.4 cubic feet.

10. A ton mole of iron is equal to __________ tons.

11. The force with which the earth attracts a body is called the __________ of the body.

12. A pure substance contains __________ kind of molecules.

13. The smallest indivisible particle of matter is called __________.

14. The atomic number is equal to the number of __________ in nucleus.

15. The atomic mass is the total number of protons and __________ in an atom of the element.

16. The average weight of atoms of an element as compared to the weight of one atom of __________ is called the atomic mass.

17. 1.0007 contains __________ significant figures.

18. The figure 24.75 will be rounded off to __________.

19. __________ means that the readings and measurements obtained in different experiments are very close to each other.

20. __________ means that the results obtained in different experiments are very close to the accepted values.

21. The degree of a measured quantity __________ with increasing number of significant figures in it.

22. The atomic mass of sodium is __________.

23. The symbolic representation of a molecule of a compound is called __________.

24. Molecular formula of CHCl3 and its Empirical formula is __________.

25. Molecular formula of benzene is C6H6 and its empirical formula is __________.

26. 58.5 is the __________ of NaCl.

27. 4.5 gms of nitrogen will have __________ molecules.

28. 28 gms of nitrogen will have __________ molecules.

29. 2 moles of SO2 is equal to __________ gms.

30. 1000 gms of H2O is equal to __________ moles.

31. The reactions, which proceed in both directions, are called __________.

32. The reactions, which proceed in forward directions only, are called __________ reactions.

33. The __________ reactions are completed after some time.

34. 0.0006 has __________ significant figures

35. 7.40 x 108 has __________ significant figures.

36. 7 x 108 has __________ significant figures.

37. Usually Molecular formula is simple multiple of the __________.

38. 0.1 mole of H2O contains __________ molecules of H2O.

39. Mass of 3.01 x 1022 molecules of CO2 is __________.

40. __________ is the branch of science which deals with the properties, composition and structure of matter.

41. None zero digits are all __________.

42. The integer part of logarithm is called __________.

43. The decimal fraction of logarithm is called __________.

44. __________ is the amount of substance, which contains as many number of particles as there are in 12 gms of Carbon.

45. 6.02 x 1023 is called the __________.

46. The accuracy of measurement depends on the number of __________.

47. __________ is the branch of chemistry that deals with quantitative relationships among the substances undergoing chemical changes.

48. The sum of atomic weights of all the elements present in molecular formula is called the __________.

49. __________ is the sum of atomic weights of the elements represented by the Empirical formula of the compound.

50. Very small and very large quantities are expressed in terms of __________.

51. In rounding off __________ figure is dropped.

52. Mole is the quantity, which has __________ particle of the substance.

53. For three significant figures, 25.55 is rounded off to __________.

54. The S.I unit of a mass is __________.

55. Mass of 6.02 x 1023 molecules of NaCl is __________ gm.

56. 1 mole of NaOH is __________ gm of NaOH.

57. Formula weight is used for __________ substances.

58. The word S.I stands for __________.

59. 4.5 gms of water will have __________ molecules.

60. 0.0087 has __________ significant figure.

Chapter 2

The Three States of Matter

1. The intermixing of gases or liquids in a container irrespective of their densities, is called __________.

2. At constant temperature, if the pressure of a given mass of a gas is decreased, its volume will __________.

3. A volume of __________ dm3 will hold 128 gms of SO2.

4. At constant temperature of a given mass of a gas, the product of its __________ and __________ is constant.

5. The rates of diffusion of gases are __________ proportional to the square root of their densities.

6. Gases deviate from ideal behaviour more markedly at high __________.

7. Liquid diffuse __________ than gases.

8. An imaginary line passing through the centre of a crystal is called __________.

9. The temperature at which more than one crystalline forms of a substance coexist in equilibrium is called __________.

10. Two or more substances crystallizing in the same form is called __________.

11. The existence of solid substances in more than one crystalline form is known as __________.

12. Rate of diffusion of gases is __________ as compared to liquids.

13. Boiling point of a liquid __________ with the pressure.

14. Mercury in a glass tube forms __________ curvature.

15. Gases can be compressed to __________ extent.

16. Viscosity of a liquid __________ with the increase of temperature.

17. Surface tension of water __________ by adding soap solution into it.

18. The internal resistance to the flow of a liquid is called __________.

19. The rise or the fall of a liquid in a capillary tube is called __________.

20. Matter exists in __________ states.

21. The freezing point of water in Fahrenheit scale is __________.

22. Boiling point of water is __________ °K.

23. SI unit for measurement of pressure is __________.

24. The value of gas law constant R = __________ dm3 atm/°K/mole.

25. The absolute Zero is equal to __________.

26. If P is plotted against 1/V at constant temperature a __________ is obtained.

27. Gases __________ in heating.

28. The pressure of air __________ at higher altitude.

29. Standard temperature means __________.

30. Standard pressure means __________.

31. Cooling is caused by __________ of gases.

32. Rate of diffusion of O2 is __________ times more than H2.

33. H2O has __________ viscosity than CH3OH.

34. Mercury does not wet the glass surface due to its higher __________.

35. Surface tension of mercury is __________ than water.

36. Viscosity can be easily measured by an instrument called __________.

37. The pressure exerted by the vapours when these vapours are in equilibrium with the liquid is called __________.

38. Vapour pressure __________ at high temperature.

39. Boyle’s Law and Charles Law can be combined into the mathematical expression __________.

40. Equal volumes of all gases at the same temperature and pressure contain __________ number of molecules.

41. The average Kinetic energy of a gas is proportional to its __________ temperature.

42. Kinetic equation may be mathematically written as __________.

43. The temperature at which two crystalline forms of a substance can coexist in equilibrium is called __________.

44. Lighter gases diffuse __________ than heavier gases.

45. Rain drops are __________ in shape.

46. Due to surface tension, the surface area of the liquid is __________.

47. Water __________ in the capillary tube.

48. Viscosity of a solution at 10°C is __________ than at 20°C.

49. Shape of NaCl crystal is __________.

50. Gases intermix to form a mixture.

51. Pressure of a dry gas is __________ than the pressure of a moist gas.

52. 22.4 dm3 of nitrogen at S.T.P will weigh equal to __________ gm.

53. 1 mole of any gas at S.T.P is equal to __________ dm3.

54. At -273°C, volume of all gases becomes __________.

55. The gases, which strictly follows the gas Laws are called __________ gas.

56. __________ is the property that determines the direction of flow of heat.

57. __________ is defined as force per unit area.

58. __________ viscosity is defined as the viscosity of a liquid as compared to the water.

Chapter 3

Structure of Atom

1. The maximum number of electrons in 2p orbital is __________.

2. 3d orbital has __________ energy than 4s orbital.

3. __________ rays are non-material in nature.

4. Charge to mass ratio of cathode rays resembles to that of __________.

5. __________ rays are most penetrating.

6. Neutrons have mass equal to that of __________.

7. Energy is __________ when an electron jumps from higher to lower orbit.

8. Second Ionization Potential has __________ value than the First Ionization Potential.

9. Electronegativity __________ from left to right in a period of Periodic Table.

10. __________ was discovered during the course of Artificial Radioactivity.

11. The velocity of alpha rays is nearly __________ of velocity of light.

12. Natural Radioactivity is confined in __________ elements.

13. The isotopes of an element differ in their __________.

14. Two electrons with the __________ spin, can never occupy the same atomic orbital.

15. ‘Al’ has electronic configuration, 1s2, 2s2, __________.

16. In a group of Periodic Table, the ionization potential __________ from top to bottom as the size of atom increases.

17. Ionization potential values __________ from left to right in a period.

18. The energy required to remove the most loosely bond electron from an atom in gaseous state is called __________.

19. The SI unit of Ionization Potential is __________.

20. An atom of sodium possesses 11 protons and __________ neutrons.

21. The particles of Cathode rays possess __________ charge.

22. The negatively charged particles found in Cathode rays are named as __________.

23. Positive rays are emitted from __________.

24. __________ rays are also known as Canal rays.

25. __________ consists of helium ions and are doubly positively charged.

26. __________ rays consists of negatively charged particles.

27. __________ rays are light waves of very short wavelength.

28. The phenomenon in which a stable element is made radioactive by artificial disintegration is called __________.

29. The electron move around the nucleus in different circular paths called __________.

30. The maximum number of electron in a shell is determined by the formula __________.

31. A particle whose mass is equal to that of electron but carries a positive charge is called __________.

32. 2p electrons are __________ in energy that 2s electrons in the same atom.

33. Number of protons of an element also indicates its __________.

34. According to __________ Principle electrons are fed in the order of increasing orbital energy.

35. According to __________ electrons are distributed among the orbitals of a sub shell to give maximum number of unpaired electron and have same spin.

36. The specific way in which the orbitals of an atom are occupied by electrons is called __________.

37. __________ rays are stream of doubly positively charged particles.

38. Electron in the outer most shell of an atom is called __________.

39. Protons are found in the __________ of an atom and bear __________ charge.

40. The atomic number of an atom is the sum of __________ inside the nucleus.

41. __________ limits the number of electron to different shell or orbits.

42. Sir William Crookes in 1878, discovered that the cathode in high vacuum tube emit radiations what he called __________.

43. X-rays were discovered in 1895 by __________.

44. The discovery of proton was done in 1886 by __________.

45. Neutrons were discovered by __________ in 1932 by the bombardment of beryllium with alpha particles.

46. Each atom has a __________, which contains all the positive charge and practically all the mass of atom.

47. Complete the reaction: 4Be9 + 2H4 ® __________ + __________.

48. __________ have higher ionization power as compared to b-rays.

49. No dark spaces between the colours are present in __________.

50. The symbol e+ represents __________.

51. p-orbitals are __________ shaped.

52. The energy released when an electron is added to an atom in the gaseous state is called __________.

53. The power of an atom to attract a shared pair of electrons towards itself is called __________.

54. Fluorine is __________ electronegative than chlorine.

55. Lyman series of spherical lines appear in the __________ portion of spectrum.

56. The electrons with __________ spin occupy the same orbital.

57. 3d orbital has __________ energy than 4s orbital.

58. Energy and frequency are __________ proportional to each other.

59. Ionic radii of cations are __________ than the atoms from which they are formed.

60. Ionic radii of anions are __________ than the atoms from which they are formed.

Chapter 4

Chemical Bonding

1. A bond formed due to transference of electron is called __________.

2. A bond formed due to sharing of electron is called __________.

3. Sigma bond is __________ than pi bond.

4. The shape of methane molecule is __________.

5. One s and 3p orbitals overlap to produce four __________ hybrid orbitals.

6. Ethene, C2H4 is an example of __________ hybridization.

7. Water molecule has __________ structure.

8. Water molecules are inter-linked with one another due to __________.

9. Polarity of the molecule is due to the difference of __________ between the two bonded atoms.

10. A chemical bond formed between to different atoms by mutual sharing of electron is termed as __________.

11. A chemical bond formed between two similar atoms by mutual sharing of electrons is known as __________.

12. The difference between the Electronegativity values of the two atoms forming covalent bond must be __________ than 1.7.

13. When two orbitals of different atoms by hybridize with each other having their axes in the same straight lines, the bond formed is termed as __________.

14. __________ bond is formed when p-orbitals of the two atoms with their axes parallel to each other overlap with each other.

15. Melting and boiling point of ionic compounds are usually __________ than that of covalent compounds.

16. Non polar compounds are usually __________ in non polar solvent.

17. The nitrogen in NH3 is __________ hybridized.

18. A hybrid orbital is called __________ orbital.

19. Since dipole moment of CS2 is zero, it is a __________ molecule.

20. A bond formed due to the electrostatic forces of attraction between the oppositely charged ions is called __________ bond.

21. The ionic bond is formed between the atoms with low ionization potential and high __________.

22. A bond formed by the sharing of an electron pair contributed by one atom only is called a __________ bond.

23. A co-ordinate covalent bond is also known as __________ bond.

24. Polar covalent bond is __________ than a non polar covalent bond.

25. H-F bond is __________ than H-Br bond.

26. The SI unit of dipole moment is __________.

27. Commonly used unit of dipole moment is __________.

28. Dipole moment of non-polar compound is __________ D.

29. The reactions of ionic compounds are usually very __________.

30. Covalent compounds are generally __________ in nature.

31. Ionic compounds are generally __________ in nature.

32. A covalent bond is represented by a __________.

33. A co-ordinate covalent bond is represented by an __________.

34. The covalent bond between H-F is called __________ covalent bond.

35. The power of an atom to attract a shared pair of electron itself is called __________ of that atom.

36. m = d x e represents __________.

37. CO2 and SO2 molecules have __________ polar bonds.

38. NH3 molecule has __________ polar bonds.

39. A double bond has __________ bond energy than a single bond.

40. An orbital which surrounds a single nucleus is called __________ orbital.

41. An orbital which surrounds two or more atomic nuclei is called __________ orbital.

42. A molecular orbital, which is of lower energy than the atomic orbitals from which it is derived, is known as __________ orbital.

43. A molecular orbital, which has higher energy than the atomic orbitals from which it is derived, is known as __________ orbital.

44. Orbitals formed after hybridization are called __________ orbitals.

45. Bond angle in Sp3 hybridization is of __________.

46. Bond angle in Sp2 hybridization is of __________.

47. Bond angle in Sp hybridization is of __________.

48. Sp3 hybridization is also known as __________.

49. Sp2 hybridization is also known as __________.

50. Sp hybridization is also known as __________.

51. A pair of electrons residing on the central atom and which is not used in bonding is called a __________.

52. The sum of total number of electron pairs (bonding and lone pairs) is called __________ number.

53. __________ bond is usually expressed by dotted line.

54. Water molecule has dipole moment because of its __________ structure.

55. CO2 is non polar because of its __________ structure.

56. Overlapping in __________ bond is perfect.

57. Overlapping in __________ bond is not perfect.

58. H-H bond is __________ than H-Cl bond.

59. __________ hybrid orbitals are not co-planar.

60. Covalent bond is Cl2 molecule is __________.

Chapter 5

Energetics of Chemical Reaction

1. The branch of Chemistry, which deals with the heat changes that take place during chemical reaction, is called __________.

2. The branch of science which deals with energy changes accompanying physical and chemical transformation is called __________.

3. The amount of heat evolved or absorbed in a chemical reaction is called __________.

4. Such reactions in which heat is evolved are called __________ reactions.

5. Such reactions in which heat is absorbed are called __________ reactions.

6. In exothermic reactions, heat evolved is given by __________ sign of DH.

7. In endothermic reactions heat absorbed is given by __________ sign of DH.

8. The total heat change in a reaction is the same whether it takes place in one or several steps.

9. The first law of thermodynamics is also known as __________.

10. The part of universe under observation is called __________.

11. The system plus its surrounding is called __________.

12. Such properties, which give description of a system at a particular moment, is called __________.

13. The term E + PV is called __________.

14. DH represents change in __________.

15. The temperature of water is raised up when sulphuric acid is added to it. This is an __________ reaction.

16. The characteristic properties of a system which is independent of amount of material concerned is called __________ properties.

17. The characteristic properties of a system which depend on amount of substance present in it is called __________ properties.

18. Density, pressure and temperature are the examples of __________ properties.

19. Mole numbers and enthalpy are the examples of __________ properties.

20. A system, which exchange both energy and matter with its surrounding, is called __________ system.

21. A system, which only exchange energy with the surrounding but not matter is, called __________ system.

22. A system which neither exchange energy nor matter with its surrounding is called __________ system.

23. A system is __________ if it contains only one phase.

24. A system is __________ if it contains more than one phase.

25. 1 kilojoule is equal to __________ joules.

26. 1 Calorie is equal to __________ joules.

27. 1 kilo calorie is equal to __________ joules.

28. The work done (w) is mathematically denoted by __________.

29. The change in enthalpy is denoted by __________.

30. __________ law is used in calculating heat of reaction.

31. __________ is defined as the change in enthalpy when one gram mole of a compound is produced from its element.

32. Heat of formation is denoted by __________.

33. When the work is done on the system by the surrounding the sign of work done (w) is __________.

34. When the work is done by the system on surrounding the sign of work done is __________.

35. First law of Thermodynamics is mathematically represented as __________.

36. Standard enthalpies are measured at __________.

37. Hess’s Law is employed to calculate __________ of a chemical reaction.

38. Heat absorbed by the system at constant volume is completely utilize to increase the __________ of the system.

39. Heat change at constant pressure from initial to final state of the system is simply equal to the __________.

40. SI unit of measurement of heat change is __________.

Chapter 6

Chemical Equilibrium

1. The reactions, which proceed in both the directions, are called __________ reactions.

2. The reactions, which proceed to one direction only, are called __________ reactions.

3. Reversible reactions are __________ completed.

4. Irreversible reactions are __________ after some time.

5. A reversible reaction is said to be in __________ when the rate of forward reaction becomes equal to the rate of backward reaction.

6. The concentrations of reactants and products are __________ at equilibrium point.

7. The value of Kc depends upon the __________ of the reactants.

8. A increase of the value of Kc tends to move the reaction to the __________ direction.

9. A decrease of the value of Kc tends to move the reaction to the __________ direction.

10. An increase in the concentration of the reactants will move the reaction to the __________ direction.

11. A decrease in the concentration of the reactants will move the reaction to the __________ direction.

12. Equilibrium constant is denoted by __________.

13. When the equilibrium constant value is very __________, we can conclude that the forward reaction is almost completed.

14. When equilibrium constant value is very __________ we can conclude that forward reaction will occur to very little extent.

15. According to __________ principle, if system in equilibrium is subjected to a stress, the equilibrium shifts in a direction to minimize or undo the effect of the stress.

16. In exothermic reaction, the __________ of temperature favour the forward rate of reaction.

17. In endothermic reactions, the __________ of temperature favour the forward rate of reaction.

18. A __________ is a substance which effects the rate of reaction but remains unaltered at the end of the reaction.

19. A catalyst increases the velocity of the reaction by decreasing the __________.

20. The suppression of degree of ionization of a sparingly soluble weak electrolyte by the addition of a strong electrolyte containing an ion in common is called __________.

21. __________ is purified in industries by Common Ion Effect.

22. A reaction moves to the left when the concentrations of the products are __________.

23. A reaction moves to the right when the concentrations of the products are __________.

24. Increase in pressure will move the reaction in the direction of __________ volume.

25. Decrease in pressure will move the reaction in the direction of __________ volume.

26. An increase of temperature favours the formation of products in case of __________ reaction.

27. A decrease of temperature fovours the formation of products in case of __________ reaction.

28. Heating moves an endothermic reaction to the __________.

29. Cooling move an exothermic reaction to the __________.

30. The product of ionic concentration in a saturated solution is called __________ constant.

31. When HCl is added to NaCl, the concentration of __________ ion is increased.

32. Chemical reaction involving the substances in more than one phases are called __________.

33. The formation of NH3 is exothermic process hence __________ temperature will favour the formation of NH3.

34. The formation of NO from N2 and O2 is endothermic process hence __________ temperature will favour the formation of NO.

35. Chemical Equilibrium is __________ equilibrium.

36. Molar concentration is also called __________.

37. The rate at which a substance takes part in a chemical reaction depends upon its __________.

38. __________ principle is applied to all reversible reaction.

39. A common ion __________ the solubility of the salt.

40. Number of moles present per dm3 of a substance is called __________.

Chapter 7

Solutions and Electrolytes

1. A mixture of two or more substances, which are homogeneously mixed, is called a __________.

2. __________ is defined as the amount of solute dissolved in a given amount of solvent.

3. A solution is composed of two components __________ and __________.

4. A solution containing one mole of solute per dm3 of solution is called one __________ solution.

5. Molarity is denoted by __________.

6. 1M solution of NaOH contains __________ gms of it dissolved per dm3 of solution.

7. A solution containing one mole of solute dissolved by per kg of solvent is called __________ solution.

8. Molality is denoted by __________.

9. 1M solution of H2SO4 contains __________ gms of it per kg of solvent.

10. The process in which ions are surrounded by water molecules is called __________.

11. The water molecules attached with the hydrated substance are called __________.

12. Hydrated copper sulphate evolves __________ water molecules on heating.

13. The interaction between salt and water to produce acids and bases is called __________.

14. The products of ionic concentration in a saturated solution at a certain temperature are called the __________.

15. Solubility product constant expressed as __________.

16. The suppression of ionization by adding a common ion is called __________.

17. The process of dissociation of an electrolyte into ions is known as __________.

18. The chemical decomposition of a compound in a solution or in fused state brought about by a flow of electric current is known as __________.

19. Electrolysis is performed in an electrolytic cell, which is known as __________.

20. The positive electrode of a voltmeter is called __________ and negative as __________.

21. A solution, which tends to resist changes in pH is called a __________ solution.

22. A mixture of acetic acid and sodium acetate acts as a __________.

23. According to Sorenson __________ is defined as negative logarithm of the hydrogen ion concentration.

24. pH is mathematically expressed as __________.

25. The pH of a neutral solution is __________.

26. __________ substances have pH values lower than 7.

27. __________ solutions have pH values more than 7.

28. Oxidation is __________ of electron.

29. Reduction is the __________ of electron.

30. Such chemical reactions in which the oxidation number of atoms or ions is changed are called __________ reactions.

31. Oxidation number of a free element is __________.

32. Oxidation number of Oxygen in a compound is __________.

33. The sum of oxidation number of any formula of a compound is __________.

34. The oxidation number of any ion is equal to the __________ on the ion.

35. __________ is the reaction in which an acid reacts with a base to form salt and water.

36. __________ are organic compounds which change colour in accordance with the pH of the medium.

37. An indicator that changes from colourless to pink in the presence of an alkaline solution is called __________.

38. An indicator that changes from red to yellow in the presence of an alkaline solution is called __________.

39. Dissociation constant is denoted by __________.

40. According to Bronsted-Lowry Concept, __________ is the donor of proton and __________ is the acceptor of proton.

41. According to Arrhenius, acid is substance that produces __________ ions when dissolved in water.

42. According to Arrhenius, base is a substance that produces __________ ions when dissolved in water.

43. When ionic product is less than ksp, the solution will __________.

44. When ionic product is greater than ksp, the solution will __________.

45. The electrode at which oxidation takes place is called __________.

46. The electrode at which reduction takes place is called __________.

47. H3O+ ion is called __________ ion.

48. The logarithm of reciprocal of hydroxyl ion (OH)- is called __________.

49. Aqueous solution of NH4Cl is __________ while that of NaHCO3 is __________.

50. The ionic product of [H+] and [OH-] of pure water is __________.

51. An increase in the oxidation number of an element or ion during a chemical change is called __________.

52. A decrease in the oxidation number of an element or ion during a chemical change is called __________.

53. The degree of dissociation __________ with the increase in temperature.

54. The degree of dissociation __________ with the dilution of electrolytic solution.

55. A __________ consists of an electrode immersed in solution of its ion.

56. The potential difference between the electrode and the solution of its salt at equilibrium position is called __________ potential.

57. If the pH of a solution is 14, the solution is __________.

58. If the pH of a solution is 4, the solution is __________.

59. The oxidation number of Mn in KMnO4 is __________.

60. The oxidation number of Fe in FeCl3 is __________.

Chapter 8

Introduction to Chemical Kinetics

1. The branch of chemistry, which deals with the study of rates and mechanisms of chemical reactions, is known as __________.

2. Such reactions, which proceeds with very high velocities and are completed very quickly are called __________ reactions.

3. Such reactions, which take place very slowly, are called __________ reactions.

4. Reactions between silver nitrate and sodium chloride to form white precipitates of silver chloride are an example of __________ reaction.

5. Reactions of Organic compounds are slow and are called __________ reactions.

6. There are some reactions, which proceed slowly with a __________ speed.

7. The rate of __________ reaction can only be determined.

8. The amount of chemical change taking place in concentration of the per unit time is called __________ of reaction.

9. Rate of reaction is expressed in __________.

10. The rate of reaction between two specific interval of time is called __________.

11. The addition energy required to bring about a chemical reaction is called __________.

12. According to __________ theory for a chemical reaction to take place, the reacting molecules must come closed together.

13. The addition of __________ helps the reaction by lowering the energy of activation.

14. The rate of reaction __________ with the increase in concentration of the reacting molecules.

15. When the concentration of both the reacting molecules is double, the probability of collisions between them will be __________ times.

16. By __________ the surface area of the reactants, the rate of reaction is increased.

17. Rate of reaction generally __________ with the rise of temperature.

18. A __________ is a substance, which either accelerates or retards the rate of reaction without taking part in the reaction.

19. In the preparation of Oxygen from Potassium Chlorate, __________ is used as catalyst.

20. In the oxidation of SO2 to SO3 by the contact process for the manufacture of H2SO4 __________ is used as catalyst.

21. An unstable intermediate compound formed during a chemical reaction is called __________.

22. When a catalyst and the reactants are in the same phases, it is known as __________ catalyst.

23. When a catalyst and the reactants are in different phases, it is called __________.

24. When a catalyst increases the rate of reaction, it is called __________ catalyst.

25. When a catalyst retards the rate of reaction, it is called __________ catalyst.

26. A negative catalyst __________ the energy of activation, hence the rate of reaction is decreased.

27. The ratio between the rate of reaction and concentration of reactants is known as __________.

28. Velocity constant is independent of concentration but depends on __________.

29. Ionic reactions are __________ than molecular reactions.

30. The value of specific rates constant for a reaction __________ with time.

31. The sum of all exponents of concentration terms in the equation is called __________.

32. The sum of moles taking part in a chemical reaction is called __________ of the reactions.