BOOK I. Some Basic Concepts of Chemistry

Transcription

1 BOOK I Some Basic Concepts of Chemistry 1. What is stoichiometry? [1] 2. How much potassium chlorate should be heated to produce 2.24L of oxygen at NTP? 3. Write an expression for molarity and molality of a solution. [2] 4. Calculate the weight of lime (CaO) obtained by heating 2000kg of 95% purelime stone (CaCO3)[2] 5. The substance which gets used up in any reaction is called [1] 6. What is 1molal solution? [1] 7. 4 litres of water are added to 2L of 6 molar HCl solutions. What is the molarity of resulting solution? [2] 8. What volume of 10M HCl and 3M HCl should be mixed to obtain 1L of 6M HCl solution? 9 What is the value of one mole? [1] 10 At NTP, what will be the volume of molecules of H2? [1] 11 Calculate the number of molecules present in 0.5 moles of CO2? [1] 12 Give one example each of a molecule in which empirical formula and molecular formula are (i) same (ii) Different. [2] 13. 1L of a gas at STP weighs 1.97g. What is molecular mass? [1] 14. Write empirical formula of following [4] CO, Na2CO3, KCl, C6H12, H2O2, H3PO4, Fe2O3, N2O Calculate the number of moles in the following masses [2] (i) 7.85g of Fe (ii) 7.9mg of Ca 16. Vitamin C is essential for the prevention of scurvy. Combustion of g of vitamin C gives g of CO2 and 0.819g of H2O. What is the empirical formula of vitamin C? [3] 17. What is the difference between precision and accuracy? [1] 18. What do you understand by significant figures? [1] 19. How many significant figures are present in [3] (a) (b) (c) State law of definite proportions. [1] 21. Explain law of multiple proportions with an example. [2] 22. State Avogadro s law. [1] 23. Write Postulates of Dalton s atomic theory. [2] 24. Define one atomic mass unit (amu).[1] 25. Calculate molecular mass of [2] C2H6, C12H22O11, H2SO4, H3PO4 26. What is formula mass? [1] 27. How are physical properties different from chemical properties? [1] 28. What are the two different system of measurement? [1] 29. Write seven fundamental quantities & their units. [2] 30. What is the difference between mass & weight? How is mass measured in laboratory? 31. How is volume measured in laboratory? Convent 0.5L into ml and 30cm3 to dm3 [2] 32. What is the SI unit of density? [1] 33. Convert 350C to of & K. [2] Class XI Page 1

2 34. What are the reference points in thermometer with Celsius scale? [1] 35. What is the SI unit of volume? What is the other common unit which in not an SI unit of volume. [1] 36. What does the following prefixes stand for [2] (a) pico (b) nano (c) centi (d) deci 37. What is chemistry? [1] 38. How has chemistry contributed towards nation s development? [1] 39. How can we say that sugar is solid and water is liquid? [2] 40. Differentiate solids, liquids & gases in terms of volume & shapes. [1] 41. How is matter classified at macroscopic level? [2] 42. Classify following substances as element, compounds and mixtures water, tea, silver, steel, carbondioxide and platinum [2] 43. Name the different methods that can be used for separation of components of a mixture. 44. Classify following as pure substances and mixtures Air, glucose, gold, sodium and milk. [1] 45. What is the difference between molecules and compounds? Give examples of each. [1] 46. How can we separate the components of a compound? [1] Lesson 2 - Structure of Atom 1. Which orbital is non directional? [1] 2. What is the meaning of quantization of energy? [1] 3. Why is energy of 1s electron lower than 2s electron? [1] 4. Which quantum number determines (i) energy of electron (ii) Orientation of orbitals. 5. What is nodal surface or nodes? [1] 6. How many spherical nodal surfaces are there in 4s sub-shell? [1] 7. Arrange the electrons represented by the following sets of quantem number in decreasing order of energy. 1. n = 4, l = 0, m = 0, s = +1/2 2. n = 3, l = 1, m = 1, s = -1/2 3. n = 3, l = 2, m = 0, s = +1/2 4. n = 3, l = 0, m = 0, s = -1/2 8. What designations are given to the orbitals having (i) n = 2, l = 1 (ii) n = 2, l = 0 (iii) n = 4, l = 3 (iv) n = 4, l = 2 (v) n = 4, l = 1? 9. Write the electronic configuration of (i) Mn4+, (ii) Fe3+ (iii) Cr2+ and Zn2+ Mention the number of unpaired electrons in each case. 10. States Heisenberg s Uncertainty Principle. [1] 11. Give the mathematical expression of uncertainty principle. [2] 12. How would the velocity be effected if the position is known? [1] 13. We don not see a car moving as a wave on the road why? [1] 14. Give the de Broglie s relation. [1] 15. Why cannot the motion of an electron around the nucleus be determined accurately? [2] 16. Calculate the uncertainty in the velocity of a wagon of mass 4000kg whose position is known accurately of ±10m [1] 17. What is the physical significance of 2 ψ up? [1] 18. Define photoelectric effect. [1] 19. How does the intensity of light effect photoelectrons? Class XI Page 2

3 20. What is threshold frequency? [1] 21. Name the scientist who demonstrated photoelectric effect experiment. [1] 22. What did Einstein explain about photoelectric effect? [1] 23. What is the relation between kinetic energy and frequency of the photoelectrons? [2] 24. Calculate energy of 2mole of photons of radiation whose frequency is Hz [1] 25. What is emission and absorption spectra? [2] 26. What transition in the hydrogen spectrum would have the same wavelength as the Balmer transition, n = 4 to n = 2 of He+ spectrum? [2] 27. Spectral lines are regarded as the finger prints of the elements. Why? [2] 28. Give the range of wavelength of the visible spectrum. [1] 29. State the two developments that led to the formation of Bohr s model of atom. [1] 30. What is an electromagnetic radiation? [1] 31. Calculate the wavelength corresponding to a frequency of 98.8MHz. [2] 32. Define black body radiation. [1] 33. Define quantum. [1] 34. Give the relation of energy (E) and frequency (v) as given by Planck. [2] 35. Calculate the frequency and energy of a photon of radiation having wavelength 30000A.[2] 36. What did Planck s theory explain? [1] 37. On what frequency does the frequency from a black body depend? [1] 38. Name the scientist who first gave the atomic model. [1] 39. What is an isotope? [1] 40. What are isobars? [1] 41. What are isotones? [1] 42. What is an atomic number? [1] 43. What is a mass number? [1] 44. Find out atomic number, mass number, number of electron and neutron in an element having atomic mass 40 and atomic number Give the main features of Thomson s Model for an atom. [2] 46. Give the drawbacks of J.J. Thomson s experiment. [1] 47. What did Rutherford conclude from the observations of α ray scattering experiment? [2] 48. Why Rutherford s model could not explain the stability of an atom? [1] 49. Name the sub atomic particles of an atom. [1] 50. Name the scientist who first formulated the atomic structure. [1] 51. What is the e/m ratio of an electron? [1] 52. What is the charge (e) of an electron? [1] 53. What is the mass (m) of an electron? [2] 54. (i) What is the mass of a proton? [1] (ii) What is the charge of a proton? [1] 55. (i) What is the mass of a neutron? [1] (ii) What is the charge of a neutron? [1] 56. Which experiment led to the discovery of electrons and how? [2] 57. Give the main properties of canal ray experiment. [2] Lesson 3 - Periodic Table 1. What is the general outer electronic configuration of f block elements? [1] 2. Why do Na and K have similar properties? [1] 3. Arrange the following elements in the increasing order of metallic character : Si, Be, Mg, Na, P. 4. The atomic number of an element is 16. Determine its position in accordance to its electronic configuration. 5. Why are elements at the extreme left and extreme right the most reactive? [2] 6. Why does the ionization enthalpy gradually decreases in a group? [1] Class XI Page 3

4 7. Why does electronegativity value increases across a period and decreases down period? 8. How does electronegativity and non metallic character related to each other? 9. Define valency. [1] 10. How does valency vary in a group and period in the periodic table? [1] 11. What is the valency of noble gases? [1] 12. How do metals react in a period? [1] 13. How do metals react in a group? [1] 14. How does the reactivity of non metals changes in a period and group? [2] 15. Give the properties of the oxides in a particular period. [2] 16. What is an amphoteric oxide? [1] 17. Define a neutral oxide. [1] 18. Why does lithium form covalent bond unlike other alkali which forms ionic bond? 19. Predict the position of the element in the periodic table satisfying the electronic configuration (n-1) d1 ns2 for n=4, 20. How does atomic size change in a group? [1] 21. Why Li and Mg show resemblance in chemical behaivour? [1] 22. The atomic radius of elements decreases along the period but Neon has highest size among III period element? Why 23. Explain why cations are smaller and anions are larger in radii than their parent atom? 24. Define ionization enthalpy and electron gain enthalpy? [2] 25. How does atomic size change in a group? [2] 26. The size of an atom can be expressed by three radii. Name them. Which of these given the highest, and the lowest value of the atomic radius of an element? 27. Among the elements B, Al, C and Si (a) Which has the highest first ionization enthalpy? (b) Which has the largest atomic radius? 28. Na+ has higher value of ionization enthalpy than Ne, though both have same electronic configuration. 29. Give the general characteristics of the long form of Modern periodic table? [1] 30. In short give the features of the seven periods. [1] 31. Define electronic configuration. [1] 32. What is the electronic configuration when elements are classified group wise? 33. Give the main features of s-block elements. [2] 34. Give the main features of p-block elements. [2] 35. Give the main features of d-block elements. [2] 36. Give the main features of f-block elements. [2] 37. How many elements are known at present? [1] 38. Who was the first scientist to classify elements according to their properties? [1] 39. What is the basis of triad formation of elements? [1] 40. Stale the modern Periodic law? [1] 41. Define and state Mendeleev s periodic law. [1] 42. How did Mendeleev arrange the elements? [2] 43. Name the two elements whose existence and properties were predicted by Mendeleev though they did not exist then. 44. Describe the main features of Mendeleev s periodic table? [3] Lesson 4 Chemical Bonding 1. Define bonding molecular orbital. [1] 2. Define antibonding molecular orbital. [1] 3. Explain diagrammatically the formation of molecular orbital by LCAO. [1] 4. Which one 2-22 O and O, may exhibit paramagnetism? [1] 5. Why are bonding molecular orbitals more stable than antibonding molecular orbitals? 6. He2 does not exist. Explain in terms of LCAO. [2] 7. Define bond order. [1] 8. Define hydrogen bonding [1] Class XI Page 4

5 9. What are the types of H-bonding? Which of them is stronger? [1] 10. NH3 has higher boiling point than PH3. Give reason. [1] 11. Define hybridisation. [1] 12. Give the features of hybridisation. [2] 13. What are the important consolations for hybridisation? [2] 14. Describe the shape of sp, sp2 and sp3 hybrid orbital? [2] 15. State the hybrid orbitals associated with B in BCl3 and C in C2H4 [1] 16. What is the state of hybridization of carbon atoms in diamond and graphite? [1] 17. Ethylene is a planar molecule whereas acetylene is a linear molecule. Give reason. 18. In H2O, H2S, H2Se, H2Te, the bond angle decreases though all have the same bent shape. Why? 19. What type of hybridisation takes place in (i) p in PCL5 and (ii) S in S F6? [1] 20. Out of p-orbital and sp-hybrid orbital which has greater directional character and Why? 21. What is sigma bond? [1] 22. What is pi bond? [1] 23. Why is σ bond stronger than π - bond? [2] 24. How many σ and π - bond are there in a molecule of C2H4 (ethane)? [1] 25. How many σ - and π - bonds are there in a molecule of CH2 = CH CH = CH2?[1] 26. What type of bond exists in multiple bond (double/ triple)? [1] 27. What are the different types of σ - bond formation? [2] 28. What type of bond are formed due to orbital overlap? [1] 29. How do covalent bonds form due to orbital overlapping? [1] 30. What is zero over lap? [2] 31. Give the main features of VSEPR Theory. [2] 32. What s difference between lone pair and bonded pair of electrons? [2] 33. CO2 is linear whereas SO2 is bend shaped. Give reason. [2] 34. Why does H2O have bent structure? [2] 35. For the molecule, Why is structure (b) more stable than structure (a)? 36. How would you attribute the structure of PH3 molecule using VSEPR model? [2] 37. In SF4 molecule, the lp electrons occupies an equatorial position in the trigonal bipyramidal arrangement to an axial position. Give reason. 38. How is VBT different from Lewis concept? [2] 39. S orbital does not show any preference for direction. Why? [2] 40. Define dipole moment. [1] 41. Give the mathematical expression of dipole moment. [1] 42. Dipole moment is a scalar or a vector quantity? [2] 43. Why NH3 has high dipole moment than NF3 though both are pyramidal? [2] 44. Why is dipole moment of CO2, BF3, CCl4 is zero? [1] 45. Why is BF3 non polar? [1] 46. Write the resonating structure of O3 molecule. [1] 47. Draw the resonating structure of NO3 48. On which factor does dipole moment depend in case of polyatomic molecules. 49. Dipole moment of Be F2 is zero. Give reason. [2] 50. Define an ionic bonding. [?] 51. Define dipole moment. [1] 52. What changes are observed in atoms undergoing ionic bonding? [2] 53. Mention the factors that influence the formation of an Ionic bond. [2] 54. Which one of the following has the highest bond order? N2, N2+ or N2-55. Define bond order. [1] 56. Give reason why H2+ ions are more stable than H2- though they have the same bond order. 57. How would the bond lengths vary in the following species? C2, C2- C22-. [2] 58. What type of bond is formed when atoms have high difference of electornegativity? 59. Out of covalent and hydrogen bonds, which is stronger. [2] Class XI Page 5

6 60. Define covalent radius. [2] 61. Define a chemical bond. [1] 62. Give the main feature of Lewis approach of chemical bonding. [1] 63. Write electron dot structure (Lewis structure) of Na, Ca, B, Br, Xe, As, Ge, N3-. [1] 64. Give the main feature of Kossel s explanation of chemical bonding. [2] 65. How can you explain the formation of NaCl according to kossel concept? [2] 66. Define electrovalent bond. [1] 67. Give the octet rule in short. [1] 68. Write the significance of octet rule. [2] 69. Write the Lewis structure for CO molecule [2] 70. Give the Lewis dot structure of HNO3 [2] Lesson 5 - STATES OF MATTER Questions based on Intermolecular forces,boyle s Law, Charles s Law, Gay Lussac s Law, Avogadro s Law 1. What are Intermolecular forces? Explain its different types with suitable example 2. State Boyle s Law. Give its mathematical Expression & graphical representation. 3. A balloon is filled with hydrogen at room temperature. It will burst if pressure exceeds 0.2 bar. If at 1 bar pressure the gas occupies 2.27 L volume, up to what volume can the balloon be expanded? 4. What will be the minimum pressure required to compress 500 dm3 of air at 1 bar to 200 dm3 at 30 C? 5. A vessel of 120 ml capacity contains a certain amount of gas at 35 C and 1.2 bar pressure. The gas is transferred to another vessel of volume 180 ml at 35 C. What would be its pressure? 6. State Charles s Law. Give its mathematical Expression & graphical representation. 7. On a ship sailing in pacific ocean where temperature is 23.4 C, a balloon is filled with 2 L air. What will be the volume of the balloon when the ship reaches Indian ocean, where temperature is 26.1 C? 8. A student forgot to add the reaction mixture to the round bottomed flask at 27 C but instead he/she placed the flask on the flame. After a lapse of time, he realized his mistake, and using a pyrometer he found the temperature of the flask was 477 C. What fraction of air would have been expelled out? 9. State Gay Lussac s Law. Give its mathematical Expression & graphical representation. 10. State Avogadro s Law. Give its mathematical Expression. Questions based on Ideal gas equation., combined gas law equation,dalton s Law of Partial Pressures 11. Derive Ideal gas equation. Give combined gas law equation. 12. Using the Ideal gas equation Show that the density of the gas is proportional to the gas pressurep. 13. At 25 C and 760 mm of Hg pressure a gas occupies 600 ml volume. What will be its pressure at a height where temperature is 10 C and volume of the gas is 640ml 14. At 0 C, the density of a certain oxide of a gas at 2 bar is same as that of dinitrogen at 5 bar. What is the molecular mass of the oxide? 15. Pressure of 1 g of an ideal gas A at 27 C is found to be 2 bar. When 2 g of another ideal gas B is introduced in the same flask at same temperature the pressure becomes 3 bar. Find a relationship between their molecular masses. 16. What will be the pressure exerted by a mixture of 3.2 g of methane and 4.4 g of carbon dioxide contained in a 9 dm3 flask at 27 C? 17. What will be the pressure of the gaseous mixture when 0.5 L of H2 at 0.8 bar and 2.0 L of dioxygen at 0.7 bar are introduced in a 1L vessel at 27 C? 18. Density of a gas is found to be 5.46 g/dm3 at 27 C at 2 bar pressure. What will be its density at STP? ml of phosphorus vapour weighs g at 546 C and 0.1 bar pressure. What is the molar mass of phosphorus? 20. Calculate the temperature of 4.0 mol of a gas occupying 5 dm3 at 3.32 bar. (R = bar dm3 Class XI Page 6

7 K 1 mol 1). 21. Calculate the total pressure in a mixture of 8 g of dioxygen and 4 g of dihydrogen confined in a vessel of 1 dm3 at 27 C. R = bar dm3 K 1 mol Calculate the volume occupied by 8.8 g of CO2 at 31.1 C and 1 bar pressure. R = bar L K 1 mol g of a gas at 95 C occupied the same volume as g of dihydrogen at 17 C, at the same pressure. What is the molar mass of the gas? 24. A mixture of dihydrogen and dioxygen at one bar pressure contains 20% by weight of dihydrogen. Calculate the partial pressure of dihydrogen State Dalton s Law of Partial Pressures. Express partial pressure in terms of mole fraction. Questions based on Kinetic molecular theory of gases, Vander Waal modified the ideal gas equation, Surface Tension & Viscosity 26. Critical temperature for CO2 and CH4 are 31.0 C and 81.9 C respectively. Which of these has stronger intermolecular forces and why? Lesson 6 - Thermodynamics 1. Questions based on system, different types of system surroundings, First law of thermodynamics, internal energy 1. Define the term system and surroundings. Explain the different types of system. 2. Explain the terms: state variables, adiabatic process, work, heat & internal energy. 3. Define First law of thermodynamics. Give its mathematical expression. 4. (a) In a process, 701 J of heat is absorbed by a system and 394 J of work is done by the system. What is the change in internal energy for the process? (b) Calculate the internal energy change when the system absorbs 5 KJ of heat and 1KJ of work 5. Express the change in internal energy of a system when(i) No heat is absorbed by the system from the surroundings, but work (w) is done on the system. What type of wall does the system have?(ii) No work is done on the system, but q amount of heat is taken out from the system and given to the surroundings. 6. What type of wall does the system have?(iii) w amount of work is done by the system and q amount of heat is Supplied to the system. What type of system would it be? Questions based on system Enthalpy & its types. 6. Explain the term Enthalpy. Give its mathematical expression. 7. If water vapour is assumed to be a perfect gas, molar enthalpy change for vapourisation of 1 mol of water at 1bar and 100 C is 41kJ mol 1. Calculate the internal energy change, when (i) 1 mol of water is aporized at 1 bar pressure and 100 C. (ii) 1 mol of water is converted into ice. 8. The reaction of cyanamide, NH2CN (s), with Dioxygen was carried out in a bomb calorimeter, and U was found to be KJ/mol at 298K. Calculate Enthalpy change for the reaction at 298K NH2 CN (s)+ 3 /2 O2 (g) N2 (g) +CO 2(g) +H2O(l) 9. Enthalpies of formation of CO (g), CO2 (g), N2O (g) and N2O4 (g) are -110, -393, 81 and 9.7 KJ/mol respectively. Find the value of rh for the reaction N2O4 (g) + 3 CO (g) N2O (g) + 3CO2 (g) 10. Given: N2 (g) + 3H2 (g) 2NH3 (g) rh0 = KJ/mol. Calculate f H0NH3 (g). 11. The enthalpy of combustion of methane, graphite and dihydrogen at 298 K are, kj mol kj mol 1, and kj mol 1 respectively. Calculate the Enthalpy of formation of CH4 (g). 12. For the process to occur under adiabatic conditions, the correct condition is i) ΔT = 0 (ii) Δp = 0 (iii) q = 0 (iv) w = What is the enthalpies of all elements in their standard states. 14. Enthalpy of combustion of carbon to CO2 is kj mol 1. Calculate the heat released upon formation of 35.2 g of CO2 from carbon and dioxygen gas. Class XI Page 7

8 15. The combustion of 1 mol of benzene takes place at 298K.After combustion CO2 and H2O are formed and 3267KJ/mol of heat is liberated.calculate f H0(C6H6) Given: f H0 = -286 KJ/mol, f H0 = -393 KJ/mol 16. Calculate the standard enthalpy of formation of CH3OH (l) from the following data:ch3oh (l) + 3/2 O2 (g) CO2 (g) + 2H2O (l) rh0 = -726 KJ/mol C (g) + O2 (g) CO2 (g) C H0 = -393 KJ/mol H2 (g) + ½ O2 (g) H2O (l) f H0 = -286 KJ/mol 17. Calculate the enthalpy change for the process : CCl4 (g) C (g) + 4Cl (g) and calculate the bond enthalpy of C-Cl in CCl4 ( g) vap H0 (CCl4 ) = 30.5 KJ/mol, f H0 (CCl4 ) = KJ/mol a H0 (C ) = 715 KJ/mol a H0 (Cl2 ) = 242 KJ/mol 18. Define the Extensive, intensive properties & Heat capacity. 19 (a) Give the relationship between Cp and Cv. (b)write a note on Bomb Calorimeter. 20. Explain the following terms with suitable examples: (a) Standard enthalpy of reaction b)standard enthalpy of formation (c) enthalpy of fusion (d) enthalpy of vaporization (e) enthalpy of Sublimation (f) enthalpy of Combustion (g) )enthalpy of Hydration (h) ) enthalpy of Atomization (i)bond enthalpy Questions based on Hess s law of constant heat summation, Born Haber cycle, Entropy. Gibbs Energy 21. Explain Hess s law of constant heat summation with an example. 22. Explain Born Haber cycle & lattice enthalpy. 23. Define Entropy. Give mathematical expressions related to it. 24. Predict in which of the following, entropy increases/decreases: (i) A liquid crystallizes into a solid. (ii) Temperature of a crystalline solid is raised from 0 K to 115 K. (iii) 2NaHCO3 (s) Na2CO3( s)+co2 (g) + H2O( g) (iv) H2 (g) 2H(g) 25. Define Gibbs Energy. Give its mathematical expression. What is Gibb s energy criteria of Spontaneity. 26. For the reaction at 298K, 2A + B C, H = 400 KJ/mol and S = 0.2 KJ/mol K At what temperature will the reaction become spontaneous? 27. For the reaction, 2Cl (g) Cl2 (g), What are the signs of H and S? 28. For the reaction: 2A (g) + B (g) 2D (g), U0 = KJ and S0 = J/K Calculate G0 for the reaction, and predict whether the reaction will occur spontaneously. 29. The equilibrium constant for a reaction is 10.Caculate G0, T =300K, R = J/K mol 30. Calculate the value of G0 for the conversion of Oxygen to Ozone, 3/2 O2 (g) O3 (g) at 298 K, if Kp for this conversion is Find out the value of equilibrium constant for the following reaction at 298 K. 2NH3 (g) + CO2(g) NH2CONH2(aq ) + H2O(l ) Standard Gibbs energy change, ΔrG0 at\ the given temperature is 13.6 kj mol Calculate the entropy change in surroundings when 1.00 mol of H2O(l) is formed under standard conditions. Δf H0 = 286 kj mol What is meant by entropy? Predict the sign of entropy change ( S) in each of the following: (a) Temperature of crystalline solid is raised from 0K to 115 K (b) A liquid crystallizes into solid (c) 2 NaHCO3 (s) Na2CO3 (s) + H2O (g) + CO2 (g) (d) 2 SO2(g) + O2 (g) <==>2SO3(g) (e) H2(at 298K,1 atm) H2(at 298K,10 atm) (f) H2O(at 298K,1 atm) H2O (at 330K,1 atm) (g) 2 NH4NO3 (s) at 1 atm & 373K 2 N2(g) + 4 H2O + O2 (g) (h) When rubber band is stretched. (i) When an egg is boiled Class XI Page 8

9 (j) C(graphite) C(diamond) (k) I2(g) I2(s) (l)hg(l) Hg(g) (m) AgNO3(s) AgNO3(aq) (n) Dissolution of iodine in a solvent (o) A partition is removed to allow two gases to mix (p) HCl is added to AgNO3 solution and precipitate of AgCl is obtained (q) crystallization of copper sulphate from its saturated solution 34 Give reasons for the following: (a) A real crystal has more entropy than an ideal crystal (b) The dissolution of NH4Cl in water is endothermic still it dissolves in water. (c ) Why does a mole of water at 00C have greater entropy than a mole of ice at 00C? (d) Neither q nor w is a state function but q+w is a state function. (e) Thermodynamically an exothermic reaction is sometimes not spontaneous. (f) The entropy of steam is more than that of water at its boiling point. (g) The equilibrium constant for a reaction is greater than one if rg0 for it is less than zero. (h) Endothermic reactions are carried out at higher temperature. (i) Evaporation of water is endothermic process but spontaneous. (j) When an ideal gas expands in vacuum, there is neither absorption nor evolution of heat. (k) Why does entropy of a solid increase on fusion? (l) Why a non-spontaneous reaction becomes spontaneous when coupled with a suitable spontaneous reaction? (m) why for predicting the spontaneity of a reaction, free energy criteria is better than the entropy criteria? (n) Why internal energy is a state function but work is not? (o) Why is standard heat of formation of diamond not zero although it is an element? (p) Why is entropy of a solution higher than that of pure liquid? (q) acetic acid and hydrochloric acid react with KOH solution. The enthalpy of neutralization of acetic acid is kj per mole while that of hydrochloric acid is kj/mol. Why? 35 Justify the following statements (a) An exothermic reaction is always thermodynamically spontaneous.why? (b)the entropy of a substance increases on going from liquid to vapour state at any temperature. ( c) At low temperatures enthalpy change dominates ΔG expression and at high temperatures it is the entropy which dominate the value of ΔG. (d) Many thermodynamically feasible reactions do not occur under ordinary conditions. (e) Reactions with ΔG0 < 0 always have an equilibrium constant greater than one. 36 Define / Discuss the following terms:(give examples and chemical equation wherever necessary ) (a) Standard enthalpy of combustion (b) Lattice enthalpy (c) Enthalpy of solution (d) Standard enthalpy fusion / vapourisation / sublimation (e) Enthalpy of atomization (f) Bond enthalpy for diatomic and polyatomic molecule. (g) Calorific value (h) Enthalpy of Neutralisation (i) State variables/ state functions (j) System (k) work (l) isobaric and isochoric process. 37 Differentiate between the following (with examples ) (a) Intensive and extensive properties (b) Enthalpy of formation and Enthalpy of reaction. (c) Enthalpy and Internal energy (d) Heat capacity and specific heat capacity. (e) Reversible and Irreversible process (f) Adiabatic and Isothermal process. (g) State function and path function (h) Exothermic and Endothermic reaction. (i) Isolated, Open, Closed, Adiabatic Systems (j) Heat and Work. Class XI Page 9

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11 Unit No. 9 Hydrogen Class XI Page 11

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13 Unit No 10 S Block Elements Class XI Page 13

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16 Unit No. 11 P Blocks Elements Class XI Page 16

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18 Unit No. 12?& 13 - Organic Chemistry : Some basic Principles and Techniques & Hydrocarbons Class XI Page 18

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20 ( Q. 8, Q. 9 [ a,b,c ] ) Class XI Page 20

21 Q.11 (a,b,c & d) Class XI Page 21

22 Unit No. 14 Environmental Chemistry Class XI Page 22