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Rich Harvest Public School HOLIDAY HOMEWORK (2018-19)

CHEMISTRY : XII * Note: The holiday homework should be done on separate registers/notebooks as per the given instructions. Submit the same by 6th July 2018 CH-SOLID STATE 1. Differentiate between amorphous and crystalline solids 2. Why is glass considered a super cooled liquid? 3. Define the term 'amorphous'. Give a few examples of amorphous solids. 4. Which solid is used in photovoltaic cells? 5. Solid A is a very hard electrical insulator in solid as well as in molten state and melts at

extremely high temperature. What type of solid is it? 6. Ionic solids conduct electricity in molten state but not in solid state. Explain 7. What type of solids are electrical conductors, malleable and ductile? 8. (i) What is meant by the term 'coordination number'? (ii) What is the coordination number of atoms: (a) in a cubic close-packed structure? (b) in a body-centred cubic structure? 9. Give the significance of a ‘lattice point’. 10. Name the parameters that characterise a unit cell. 11. Distinguish between (i) Hexagonal and monoclinic unit cells(ii) Face-centred and end-centred unit cells. 12. How will you distinguish between the following pairs of terms: (i) Hexagonal close-packing and cubic close-packing?(ii) Crystal lattice and unit cell? (iii) Tetrahedral void and octahedral void? 13. Calculate the efficiency of packing in case of a metal crystal for (i) simple cubic(ii) body-centred cubic(iii) face-centred cubic (with the assumptions that atoms are

touching each other). 14. What is a semiconductor? Describe the two types of semiconductors and contrast their

conduction mechanism. 15. Which of the following lattices has the highest packing efficiency (i) simple cubic (ii) body-

centred cubic and (iii) hexagonal close-packed lattice? 16. In terms of band theory, what is the difference between: (i) a conductor and an insulator(ii) a conductor and a semiconductor? 17. Explain the following terms with suitable examples: (i) Schottky defect (ii) Frenkel defect (iii) Interstitials and (iv) F-centres. 18. Explain the following with suitable examples:

QUESTION BANK FROM THE CHAPTERS GIVEN BELOW: CH-1 :SOLID STATE CH-2 :SOLUTIONS

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(i) Ferromagnetism(ii) Para magnetism (iii) Ferrimagnetism(iv) Ant ferromagnetism (v) 12-16 and 13-15 group compounds. 19. What type of defect can arise when a solid is heated? Which physical property is affected by it

and in what way? 20. Explain how vacancies are introduced in an ionic solid when a cation of higher valence is added as an impurity in it. Explain what impurity defect is. 21. Ionic solids, which have anionic vacancies due to metal excess defect, develop colour. Explain with the help of a suitable example. 22. What type of stoichiometric defect is shown by: (i) ZnS (ii) AgBr 23. A group 14 element is to be converted into n-type semiconductor by doping it with a suitable impurity. To which group should this impurity belong? 24. Why does ZnO appear yellow on heating? Or Why does ZnO exhibit increased electrical

conductivity on heating? 25. Define the terms isotropy and anisotropy. 26. KClcrystals sometimes appear violet. Explain 27. What is the nature of of crystal defect produced when NaCl is doped with MgCl2. 28. What happens when a ferromagnetic substance is subjected to high temperature? 29. What is curie temperature? 30. Why is common salt yellow sometimes instead of white? 31. Why are Frenkel defects not found in pure alkali halides? 32. Why is Frenkel defect found in AgCl? 33. If the radius of the octahedral void is r and radius of the atoms in close packing is R, derive relation between r and R. 34. Classify each of the following as being either a p-type or a n-type semiconductor:(i) Ge doped with In (ii) B doped with Si. 35. What type of substances would make better permanent magnets, ferromagnetic or ferrimagnetic. Justify your answer. 36. Explain (i) The basis of similarities and differences between metallic and ionic crystals. (ii) Ionic solids are hard and brittle. 37. Explain how much portion of an atom located at (i) corner and (ii) body centre of a cubic unit cell

is part of its neighbouring unit cell. 38. Give the significance of a ‘lattice point’. 39. What type of solids are electrical conductors, malleable and ductile? 40. Ionic solids conduct electricity in molten state but not in solid state. Explain. 41. Classify the following solids in different categories based on the nature of intermolecular forces operating in them: Potassium sulphate, tin, benzene, urea, ammonia, water, zinc sulphide, graphite, rubidium, argon, silicon carbide. 42. Classify the following as amorphous or crystalline solids: Polyurethane, naphthalene, benzoic

acid, Teflon, potassium nitrate, cellophane, polyvinylchloride, fibre glass, copper. 43.Why do solids have a definite volume? 44. Why are solids rigid? 45. Refractive index of a solid is observed to have the same value along all directions. Comment on the nature of this solid. Would it show cleavage property? 46. What makes a glass different from a solid such as quartz? Under what conditions could quartz be

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converted into glass? 47.How can you determine atomic mass of an unknown metal if you know its density and dimension of the unit cell?

CH-Solutions 1. State Henry’s Law. What are the applications constant temperature. 2. What are azeotropes? Classify them with examples. 3. What do you understand by colligative properties? 4. How is relative lowering of vapour pressure defined for a solution containing a volatile

solvent and a non volatile solute? How is this function related to mole fraction of the solute and the solvent? (or)What is relative lowering of vapour pressure? Derive an expression for it. How is it useful in determining molecular mass of solute?

5. The solution of a non volatile solute boils at a higher temperature than the pure solvent. Show this relationship on a graphic diagram. (or) Why is the boiling point of solvent elevated when a non volatile solute is added to it?

6. With the help of a neat diagram, indicate why the solution of a non volatile solute should freeze at a temperature lower than the freezing point of the pure solvent? (or) Why is the freezing point of a solvent be depressed when a non volatile solute is added to it?

7. How can you say that osmotic pressure is a colligative property? How can it be used to determine the molecular mass of a non volatile solute?

8. What is meant by abnormal molecular mass of solute? Discuss the factors which bring abnormality in the experimentally determined molecular masses of solutes using colligative properties.

9. What is Van’t Hoff factor? 10. Define the following terms. Mass percentage, Mass percentage, Mass by volume

percentage, Parts per million, Mole fraction, Molarity, Molality. 11. What are isotonic solutions? 12. Define the term solution. How many types of solutions are formed? Write briefly about

each type with an example. 13. Suppose a solid solution is formed between two substances, one whose particles are

very large and the other whose particles are very small. What kind of solid solution is this likely to be?

14. What role does the molecular interaction play in a solution of alcohol and water? 15. Why do gases always tend to be less soluble in liquids as the temperature is raised? 16. Suggest the most important type of intermolecular attractive interaction in the following

pairs. a. n-hexane and n-octane (ii) I2 and CCl4 (iii) NaClO4 and water (iv) methanol and

acetone (v) acetonitrile (CH3CN) and acetone (C3H6O). 17. Based on solute-solvent interactions, arrange the following in order of increasing

solubility in n-octane and explain. Cyclohexane, KCl, CH3OH, CH3CN. 18. Amongst the following compounds, identify which are insoluble, partially soluble and

highly soluble in water? (i) phenol (ii) toluene (iii) formic acid (iv) ethylene glycol (v) chloroform (vi) pentanol.

19. What are the advantages of using Osmotic property to determine molecular mass of solids over other colligative properties?

20. What do you mean by boiling point elevation constant (ebullioscopy constant) and freezing point depression constant (cryoscopy constant) for a solvent.

21. What is the effect of temperature on molality and molarity? NUMERICALS 1. A solution to be used in hand lotion is made by mixing 90.0g of water, 9.2g of ethyl

alcohol and and 18.4g of glycerol

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( C3H8O3). Calculate the mole fraction of glycerol present in it. (0.037) 2. The density of 10% by mass of KCl solution is 1.06g/cc. Calculate molarity of the

solution. ( 1.42M) 3. Find the molarity of the solution obtained by miximg 100mL of 0.3M H2SO4 and 200mL

of 1.5M H2SO4 (1.1M) 4. What volume of 96% of H2SO4 solution (density = 1.83g/mL) is required to make 4 litres

of 3M H2SO4? (669 mL of diluted to 4 litres). 5. Methanol and ethanol form nearly an ideal solution at 300K. A solution is made by

mixing 32g methanol and 23g ethanol. Calculate the partial pressures of the components and total vapour pressure of solution at 300K. (90 mmHg, 17mmHg, 77mm Hg)

6. Vapou pressures of pure benzene and toluene at 293 K is 75mm Hg and 22 mm Hg repecively.

7. . Vapour pressure of pure benzene and toluene at 293 K is 75 mm Hg and 22 mm Hg respectively. 23.4 g of benzene and 64.4 g of toluene are mixed. If the two form an ideal solution, calculate the mole fraction of benzene in the vapour phase assuming that the vapours are in equilibrium with the liquid mixture at this temperature. (0.59)

8. 18.2 g of urea is dissolved in 100g of water at 50ºC. The lowering of vapour pressure produced is 5mm Hg. Calculate the molecular mass of urea. The vapour pressure of water at 50ºC is 92 mm Hg. (57.05)

9. 8. The vapour pressure of dilute aqueous solution of glucose is 750 mm Hg. Calculate (i)molality (ii) mole fraction of solute. (0.74m ; 0.0132)

10. The boiling point of water becomes becomes 100.52 50ºC. If 1.5 g of a non volatile solute is dissolved in 100 mL of it. Calculate the molecular mass of the solute. (Kb = 0.6k/m ) (17.3)

11. 10 g of a non volatile solute when dissolved in 100g of benzene raises its boiling point by 1ºC. What is the molecular mass of the solute? (253)

12. The osmotic pressure of a dilute aqueous solution of a compound X containg 0.12g per litre is twice the osmotic pressure of a dilute aqueous solution of another compound Y containing 0.18g per litre. What is the ratio of the molecular mass of X to Y if both remain in molecular form in solution? (1:3)

13. 1 litter aqueous solution of sucrose weighing 1015 g has an osmotic pressure of 4.82 atm at 293 K. What is the molality of sucrose solution? (0.2112m)

14. Assuming complete ionisation, calculate the expected freezing point of solton prepared by adding Glauber’s salt Na2SO4.10H2O in 0.1Kg of water? Kf = 1.86 K Kg/mol. (271.95)

15. A solution contains 7.45 KCl per litre of the solution, It has an osmotic pressure of 4.68 atm at 300 K. Calculate the degree of dissociation of KCl in this solution. (90%)