Chem XI Short Answers Questions Chapter 1 PDF

CHAPTER 1 Exercise Short Answers Questions 1. What are ions? Under what conditions are they produced? 2. How does a mass spectrograph show the relative abundance of isotopes of an element? 3. What is the justification of two strong peaks in the mass spectrum fo...

CHAPTER 1 Exercise Short Answers Questions 1. What are ions? Under what conditions are they produced? 2. How does a mass spectrograph show the relative abundance of isotopes of an element? 3. What is the justification of two strong peaks in the mass spectrum for bromine; while for iodine only one peak at 127 amu is indicated? 4. 23 g of sodium and 238 g of uranium have equal number of atoms in them. 5. Mg atom is twice heavier than that of carbon atom. 6. 180 g of glucose and 342 g of sucrose have same number of molecules but different number of atoms present in them. 7. 4.9 g of H2SO4 when completely ionized in water, have equal number of positive and negative charges but number of positively charged ions are twice the number of negatively charged ions. 8. One mg of K2CrO4 has thrice the number of ions than the number of formula units when ionized in water. 9. 2 g of H2, 16 g of CH4 and 44 g of CO2 occupy separately the volumes of 22.414 dm3, although the sizes and masses of molecules of three gases are very different from each other. 10. Calculate mass in grams of 2.74 moles of KMnO4. 11. Calculate moles of O atoms in 9.00 g of Mg(NO3)2. 12. Calculate number of O atoms in 10.037 g of CuSO4.5H2O. 13. Calculate mass in kilograms of 2.6  1020 molecules of SO2. 14. Calculate moles of Cl atoms in 0.822 g C2H4Cl2. 15. Calculate mass in grams of 5.136 moles of Ag2CO3. 16. Calculate mass in grams of 2.78  1021 molecules of CrO2Cl2. 17. Calculate number of moles and formula units in 100 g of KClO3. 18. Calculate number of K+ ions, ClO31− ions, Cl atoms, and O atoms in 100g of KClO3. 19. How do we calculate the percentage yield of a chemical reaction? 20. What are the factors which are mostly responsible for the low yield of the products in chemical 21. reactions? or Why actual yield is always less than theoretical yield. 22. Law of conservation of mass has to be obeyed during stoichiometric calculations. 23. Many chemical reactions taking place in our surrounding involve the limiting reactants. 24. No individual neon atom in the sample of the element has a mass of 20.18 amu. 25. One mole of H2SO4 should completely react with two moles of NaOH. How does Avogadro’s number help to explain it? 26. One mole of H2O has two moles of bonds, three moles of atoms, ten moles of electrons and twenty eight moles of total fundamental particles in it. 27. N2 and CO have the same number of electrons, protons and neutrons. Previous Boards & Additional Short Answers Questions 1. Write down main postulates of Dalton's atomic theory. 2. Differentiate between atom and molecule. 3. Define Atomicity. Give two examples. 4. Define macromolecules. Give examples. 5. Differentiate between cation and anion. 6. The removal of an electron from a neutral atom is an endothermic process. Explain with reason. 7. Define relative atomic mass. Give two examples. 8. What are molecular ions? How are they generated? 9. What are isotopes? Why they have same chemical properties? 10. Why do the isotopes have same chemical but different physical properties? 11. What are monoisotopic elements? Give name and symbol of such an element. 12. What is relative abundance of Isotopes? How is it determined? 13. Define Mass Spectrometer. 14. Differentiate between Aston’s Mass Spectrograph and Dempster’s Mass Spectrometer 15. Why is a very low pressure maintained inside the mass spectrometer? 16. How the ions in the mass spectrometer are accelerated and deflected? 17. What is the function of ionization chamber in mass spectrometer? 18. What is the function of magnetic field in mass spectrometer? 19. What is the function of electrometer in separation of isotopes in mass spectrometer? 20. Write only names of any four methods employed for the separation of Isotopes. 21. What is mass spectrum? 22. What is the reason that elements have fractional atomic masses? 23. Write importance of combustion analysis. 24. What is the function of Mg(ClO4)2 and 50% KOH in combustion analysis? 25. In combustion analysis, why the %age of oxygen cannot be measured directly? 26. Write importance of combustion analysis. 27. Differentiate between empirical and molecular formula. 28. Define empirical formula. How is it related to molecular formula? 29. Molecular formula is multiple of empirical formula. Explain with an example. 30. A compound may have same empirical formula as well as the molecular formula. Justify. 31. What is Avogadro’s number? Give equation to relate Avogadro’s number and mass of an element. 32. 58.5 a.m.u. is the formula mass of NaCl but not the molecular mass. Why? 33. One mole of different gases has different masses but occupies same volume. Explain why? 34. Calculate mass of 10−3 moles of MgSO4. 35. Calculate the number of water molecules in 10 g of ice. 36. Calculate the number of molecules present in 34 g of H3PO4. 37. Calculate moles in 0.1 gm of sodium. 38. Calculate the mass in grams of 10−3 moles of water. 39. What is the number of H+ ions in 9.8 g of H3PO4? 40. Define stoichiometry and write down its basic assumptions. 41. Write down limitations of a chemical equation. 42. Define limiting reactant. How does it control the yield of product formed? 43. Define Limiting reactants. How they can be identified? 44. Concept of limiting reactant is not applicable to the reversible reactions. 45. Why in chemical reactions one or more reactants is/are deliberately used in excess? 46. In industry, costly reactant is always taken as limiting reactant. 47. 11 g of carbon is reacted with 32 g of oxygen to give CO2. Which is the limiting reactant C + O2 ⎯→ CO2. 48. Differentiate between actual yield and theoretical yield.  CHAPTER 2 Exercise Short Answers Questions 1. Is partition chromatography a modified form of solvent extraction? If yes, justify the statement. 2. Why is there a need to crystalize the crude product? 3. A solid organic compound is soluble in water as well as in chloroform. During its preparation, it remains in aqueous layer. Describe a method of separation. 4. In solvent extraction technique, why repeated extraction using small portions of solvent are more efficient than using a single extraction but larger volume of solvent? 5. Write down the main characteristics of a solvent selected for crystallization of a compound. Previous Boards & Additional Short Answers Questions 1. Define analytical chemistry. 2. How chemical characterization of a compound is done? 3. Differentiate between qualitative and quantitative analysis. 4. Mention only steps involved in completer quantitative determination. 5. Name different methods for separation and purification of a compound. 6. How we can run the process of filtration smoothly? 7. Why fluted filter paper is used for greater rate of filtration than ordinary cone filter paper? 8. Concentrated HCl and KMnO4 solutions cannot be filtered by Gooch crucible. Give reason. 9. How does Gooch Crucible increase rate of filtration? 10. What is the difference between Gooch crucible and sintered glass crucible? 11. Differentiate between residue and filtrate. 12. What is the principle of crystallization? 13. Why NaCl cannot be purified by crystallization? 14. Mention the major steps involved in crystallization. 15. How saturated solution for crystallization can be prepared? 16. Hot saturated solution is cooled at moderate rate during crystallization. Why? 17. Explain briefly two methods for drying of the crystallized substance. 18. How are crystals dried using filter paper? Give its two disadvantages. 19. How crystals are dried in vacuum desiccator? 20. Name two drying agents used in vacuum desiccator. 21. The desiccator is a safe and reliable method for drying the crystals. Explain. 22. How can you remove undesirable colours from the crystals? 23. Define sublimation with examples. 24. Naphthalene can best be purified by sublimation. Why? 25. How mixture of NH4Cl and NaCl can be separated? 26. What is solvent extraction? 27. State Distribution / Partition law. 28. What is distribution co-efficient? To which technique it is applicable? 29. What are factors that decide solvent extraction or crystallization, a technique selected for 30. purification of a compound? 31. Iodine is more soluble in water in the presence of KI. Explain. 32. Define chromatography and give formula of distribution coefficient. 33. Differentiate between stationary and mobile phase. Give examples. 34. Differentiate between adsorption and partition chromatography. 35. Define chromatogram. How it is developed? 36. What is Rf? Why it has no unit? 37. Rf value is always less than 1.0. Comment on it. 38. Define chromatography. Give its two applications. 39. Give the four uses of paper chromatography.  CHAPTER 3 Exercise Short Answers Questions 1. What are isotherms? What happens to the positions of isotherms when they are plotted at high temperature for a particular gas? 2. How will you explain that the value of the constant k in the equation PV = k depends upon; (i) the temperature of a gas (ii) the quantity of a gas 3. What is the Charles's law? Which scale of temperature is used to verify that V/T = k? 4. A sample of carbon monoxide gas occupies 150.0 mL at 25.0°C. It is then cooled at constant pressure until it occupies 100.0 mL. What is the new' temperature? 5. Do you think that the volume of any quantity of a gas becomes zero at –273.16°C. Is it not against the law of conservation of mass? How do you deduce the idea of absolute zero from this information? 6. What is Kelvin scale of temperature? Plot a graph for one mole of a real gas to prove that a gas becomes liquid, earlier than –273.16oC. 7. Throw some light on the factor 1/273 in Charles's law. 8. Can we determine the molecular mass of an unknown gas if we know the pressure, temperature and volume along with the mass of that gas? 9. How do you justify from general gas equation that increase in temperature or decrease of pressure decreases the density of the gas? 10. Why do we feel comfortable in expressing the densities of gases in the units of g dm -3 rather than g cm- 3 , a unit which is used to express the densities of liquids and solids? 11. What is Avogadro’s law of gases? 12. Do you think that 1 mole of H2 and 1 mole of NH3 at 0oC and 1 atm pressure will have Avogadro’s number of particles? 13. Justify that 1 cm3 of H2 and 1 cm3 of CH4 at STP will have same number of molecules, when one molecule of CH4 is 8 times heavier than that of hydrogen. 14. Dalton’s law of partial pressures is only obeyed by those gases which don’t have attractive forces among their molecules. Explain it. 15. Derive an equation to find out the partial pressure of a gas knowing the individual moles of component gases and the total pressure of the mixture. 16. Explain that the process of respiration obeys the Dalton’s law of partial pressures. 17. How do you differentiate between diffusion and effusion? 18. What is critical temperature of a gas? What is its importance for liquefaction of gases? 19. What is Joule-Thomson effect? 20. Gases show non-ideal behaviour at low temperature and high pressure. Explain this with the help of a graph. 21. Do you think that some of the postulates of kinetic molecular theory of gases are faulty? Point out these postulates. 22. Hydrogen and helium are ideal at room temperature, but SO2, and Cl2 are nonideal. How will you explain this? 23. What is the physical significance of van der Waals' constants, ’a’ and ’b? Give their units. 24. The plot of PV versus P is a straight line at constant temperature and with a fixed number of moles of an ideal gas. 25. Pressure of NH3 gas at given conditions (say 20 atm pressure and room temperature) is less as calculated by van der Waals equation than that calculated by general gas equation. 26. Water vapours do not behave ideally at 273K. 27. SO2 is comparatively non-ideal at 273K but behaves ideally at 327oC. Previous Boards & Additional Short Answers Questions 1. Define pressure. Give its S.I units. 2. Write down the values of atmospheric pressure in four different units. 3. Define one atmospheric pressure. Give its two units. 4. Atmospheric pressure on the top of Mount Everest is 323 mm Hg. Convert this value to Pascal and to atmospheres. 5. State Boyle’s law. Give its expression. 6. The product of pressure and volume of a gas at constant temperature and number of moles is constant, why? 7. What is the relationship between volume and temperature at constant pressure for a given mass of a gas? 8. What do you mean by Absolute zero? 9. What is absolute zero? What happens to real gases while approaching it? 10. How kinetic energy of molecules of a gas becomes zero at –273oC? 11. Give the quantitative definition of Charles’s law. 12. –273.16oC is regarded as lowest possible temperature. Justify it. 13. Justify that volume of given mass of a gas becomes theoretically zero at –273oC? 14. Convert 80oC to Fahrenheit scale. 15. How the various scales of thermometry can be interconverted? 16. Convert -40oF temperature to Kelvin temperature. 17. Convert 100oF into oC. 18. Convert 37oC into oF scale. 19. What is Fahrenheit scale? Give formula to convert temperature in oC to oF. 20. What is the difference between centigrade and Fahrenheit scale and which relationship is used for their interconversion. 21. Calculate the density of methane at STP. 22. Derive molecular mass of a gas by general gas equation. 23. Derive expression of density of gas with help of general gas equation. Prove that d = PM/RT. 24. Why regular air cannot be used in diver’s tank? / Why deep-sea divers take oxygen mixed with an inert gas like He? 25. Why pilot feel uncomfortable breathing at higher altitude? 26. State Graham’s law of diffusion along with mathematical form. 27. Lighter gases can diffuse more rapidly than heavier gases. 28. Write expression for Kinetic equation and root mean square velocity of gases. 29. Derivation of Gas laws from KMT of gases. 30. What are faulty points in kinetic molecular theory of gases? 31. Why gases do not settle down in a vessel? 32. Ammonia gas can be liquefied quite easily than hydrogen gas. Justify. 33. What is critical temperature? It depends upon what factors? 34. Why is the critical temperature of water higher than argon? 35. Define Joule Thomson effect and give its significance. 36. The value of compressibility factor for H2 and He is always positive. Justify. 37. Why gases deviate from ideal behaviour? 38. The deviation of real gases is more at high pressure and low temperature. Why? 39. High pressure and low temperature make the gases non-ideal. Explain why? 40. Why the volume correction is done by van der Waals? 41. Derive the SI units of van der Waals constant ’a’ and ’b’. 42. Which is the fourth state of matter? How it can be obtained? 43. What is plasma state? How is plasma formed at high temperature? 44. Define Plasma. Why it is neutral? 45. What are characteristics of plasma? 46. State what is natural and artificial plasma? 47. Where is plasma found? 48. What is plasma? Give its two uses. 49. What are applications of plasma? 50. Write down any two uses of plasma.  CHAPTER 4 LIQUIDS Exercise Short Answers Questions 1. In the hydrogen bonded structure of HF which is the stronger bond, the shorter covalent bond or the longer hydrogen bond between different molecules? 2. In a very cold winter the fish in garden ponds owe their lives due to hydrogen bonding. 3. Water and ethanol can mix easily and in all proportions. 4. The origin of the intermolecular forces in water. 5. Evaporation causes cooling. 6. Evaporation takes place at all temperatures. 7. Earthenware vessels keep water cool. 8. One feels sense of cooling under the fan after bath. 9. Dynamic equilibrium is established during evaporation of a liquid in a closed vessel at constant temperature. 10. Heat of sublimation of a substance is greater than that of heat of vaporization. 11. Heat of sublimation of iodine is very high. Previous Boards & Additional Short Answers Questions 1. What are intramolecular forces of attractions. Give one example. 2. What are dipole-dipole forces? Name the properties which are affected by these forces. 3. What are Debye forces? Give example. 4. What are London Dispersion forces? 5. London dispersion forces are weaker than dipole-dipole forces. Why? 6. Why did the boiling point of noble gases increase within a group? 7. Melting and boiling points of halogens increase down the group. Explain it. 8. Ethane and Hexane has B.P –88.6oC and 68.7oC respectively, comment on this drastic change. or Ethane is a gas while hexane is a liquid. Justify. 9. Melting and boiling points of alkanes increase with increase in molar masses. Why? 10. Define Polarizability. How it affects London dispersion forces? 11. Give reason for the lowest boiling point of hydride of group IV-A elements. 12. Boiling point of water is greater than boiling point of HF, although hydrogen bonding is stronger in HF in H2O. Why? 13. Water is liquid at room temperature while H2S is a gas. Comment. 14. Why HF is a weaker acid as compared to HCl, HBr and HI? 15. Lower alcohols are soluble in H2O but hydrocarbons are insoluble. Give reason. 16. Water freezes from surface to the downward direction in ponds and lakes. Explain why? 17. Ice floats on water. Justify it. / Why the density of ice is less than water? 18. Why do fish and plants in ponds survive under blanket of ice during cold winters? 19. Explain action of soaps and detergents in light of hydrogen bonding. 20. What type of intermolecular forces will dominate in the following liquids? (i) CH3COCH3 (propanone) (b) C8H18 (octane) 21. How the rate of evaporation depends on the surface area? 22. Rate of evaporation decreases with decrease in temperature. Why? 23. How evaporation is a continuous process? 24. Define vapour pressure. Name the factors which affect vapour pressure. 25. Why vapour pressure of water, ethyl alcohol and diethyl ether are different from each other? 26. The vapour pressure of solids is far less than those of liquids. Give reason. 27. Why vapour pressure of CCl4 is 87 torr while isopentane is 580 torr at 20oC. 28. Why are the vapour pressure of solids far less than those of liquids? 29. Why different liquids evaporate at different rates even at the same temperature? 30. The vapour pressure of diethyl ether is higher than that of water at same temperature. Give reason. 31. Define boiling point. How it is effected by external pressure? 32. Boiling point of water is high as compared to boiling point of ether. Why? 33. Steam causes more severe burns than does the boiling water. Give reason. 34. Why the temperature of a boiling liquid does not rise even if heat is continuously supplied to it? 35. Why boiling needs a constant supply of heat? 36. Why the boiling point of water is different at Murree hills and at Mount Everest? 37. What is vacuum distillation? Give its advantages. 38. How vacuum distillation can be used to avoid decomposition of a sensitive liquid? 39. Why heat of vapourization of water is greater than CH4? 40. Write the importance/advantages of Vacuum Distillation. 41. Why water boils at low temperature at mountains? 42. Define molar heat of vapourization. 43. Why heat of sublimation of I2 is very high than other halogens? 44. What are liquid crystals? Who discovered it? 45. Give two uses of liquid crystals. 46. How liquid crystals can act as temperature sensors? 47. How liquid crystals are used to locate veins, arteries, infections and tumors? 48. How the liquid crystals help in the detection of the blockage in veins and arteries? SOLIDS Exercise Short Answers Questions 1. How polymorphism and allotropy are related to each other? Give examples. 2. Crystals of salts fracture easily but metals are deformed under stress without fracturing. Explain the difference. 3. Sodium is softer than copper, but both are very good electrical conductors. 4. Diamond is hard and an electrical insulator. 5. Sodium chloride and caesium chloride have different structures. 6. Iodine dissolves readily in tetrachloromethane. 7. The vapour pressures of solids are far less than those of liquids. 8. Amorphous solid like glass is also called super cooled liquid. 9. Cleavage of the crystals is itself anisotropic behaviour. 10. The crystals showing isomorphism mostly have the same atomic ratios. 11. The transition temperature is shown by elements having allotropic forms and by compounds showing polymorphism. 12. One of the unit cell angles of hexagonal crystal is 120°. 13. The electrical conductivity of the metals decreases by increasing temperature. Justify. 14. In the closest packing of atoms of metals, only 74% space is occupied. 15. Ionic crystals don’t conduct electricity in the solid state. 16. Ionic crystals are highly brittle. 17. The number of positive ions surrounding the negative ion in the ionic crystal lattice depends upon the sizes of the two ions. Previous Boards & Additional Short Answers Questions 1. How can you convert crystalline solid into amorphous solid? 2. Why graphite shows electrical conductance in one plane more than other? 3. What is habit of a crystal? 4. What do you mean by symmetry? Give elements of symmetry. 5. Define isomorphism and polymorphism in the crystalline solids and give examples. 6. Define transition temperature with two examples. 7. Define cleavage plane and unit cell. 8. Explain the term unit cells dimensions. 9. Explain tetragonal system showing angles, faces and give two examples. 10. Write down names of different types of crystalline solids. 11. Give four properties of ionic solids. 12. Why the ionic crystalline solids have high melting points? 13. Solid sodium chloride does not conduct electricity, but when electric current is passed through molten sodium chloride or its aqueous solution, electrolysis takes place. Give reason. 14. Sodium chloride and Cesium fluoride have the same geometry, comment on it. 15. Define lattice energy with an example. 16. Why NaF has higher lattice energy than NaCl? 17. The lattice energy of NaCl is greater than KCl. Why? 18. Graphite has slippery touch. Give reason. 19. Why graphite is good conductor of electricity but diamond is bad conductor of electricity? 20. Write four properties of molecular solids. 21. Why molecular solids are soft and easily compressible? 22. Why does Iodine sublime. 23. NaCl crystals are harder than I2 crystals. 24. Write four properties of metallic crystals. 25. A freshly cut metal has a shiny look. Justify. 26. Why metals are good conductor of electricity? 27. Why metals are changed into wires and sheets by applying the stress?  CHAPTER 5 ATOMIC STRUCTURE Exercise Short Answers Questions 1. Why is it necessary to decrease the pressure in the discharge tube to get the cathode rays? 2. Whichever gas is used in the discharge tube, the nature of the cathode rays remains the same. Why? 3. Why e/m value of the cathode rays is just equal to that of electron? 4. How the bending of the cathode rays in the electric and magnetic fields shows that they are negatively charged? 5. Why the positive rays are also called canal rays? 6. The e/m value of positive rays for different gases are different but those for cathode rays the e/m values are the same. Justify it. 7. The e/m value for positive rays obtained from hydrogen gas is 1836 times less than that of cathode rays. Justify it. 8. Evaluate mass of electron. 9. Which postulate Bohr’s model tells us that orbits are stationary and energy is quantized? 10. How does the Bohr’s equation tell you that; (i) radius is directly proportional to the square of the number of orbits. (ii) radius is inversely proportional to the number of protons in the nucleus. 11. How do you come to know that the velocities of electrons in higher orbits are less than those in lower orbits of hydrogen atom? 12. Justify that the distance gaps between different orbits go on increasing from the lower to the higher orbits. 13. The potential energy of the bounded electron is negative. 14. Total energy of the bounded electron is also negative. 15. Energy of an electron is inversely proportional to n2, but energy of higher orbits is always greater than those of the lower orbits. 16. The energy difference between adjacent levels goes on decreasing sharply. 17. Derive the following equations for hydrogen atom, which are related to the (i) energy difference between two levels, n1 and n2. (ii) frequency of photon emitted when an electron jumps from n2 to n1. (iii) wave number of the photon when an electron jumps from n2 to n1. 18. What is spectrum. Differentiate between continuous spectrum and line spectrum. 19. Compare line emission and line absorption spectra. 20. What is the origin of line spectrum? 21. H-atom and He+ are mono-electronic systems, but size of He+ is much smaller than H, why? 22. Do you think that the size of Li+2 is even smaller than He+? Justify with calculations. 23. What are X-rays? What is their origin? How was the idea of atomic number derived from the discovery of X-rays? 24. How does the Bohr’s model justify the Moseley’s equation? 25. Compare orbit and orbital. 26. When azimuthal quantum number has a value 3, then there are seven values of magnetic quantum number. Give reasons. 27. What is (n + l) rule. Arrange the orbitals according to this rule. Do you think that this rule is applicable to degenerate orbitals? 28. Distribute electrons in orbitals of 57La, 29Cu, 79Au, 24Cr, 53I, 86Rn. Previous Boards & Additional Short Answers Questions 1. How will you prove that cathode rays travel in straight line? 2. Give an experiment in the cathode rays tube to show that cathode rays are material particles. 3. Give reason for the production of positive rays. 4. Why the properties of positive rays depend upon the nature of the gas? 5. Justify that e/m value of positive rays is maximum for hydrogen gas. 6. How was neutron discovered by James Chadwick? Prove it by a nuclear reaction. 7. What particles are formed by the decay of free neutron? 14 11 8. How 7N is converted into 5B. Give equation. 65 66 9. Write down two equations when slow moving neutrons hit the Cu metal. or How 29Cu is converted into 30Zn. Give equation. 10. How gamma rays are produced by slow neutron? 11. How neutrons are used in the treatment of cancer? 12. How did Rutherford’s model of an atom first of all proved the existence of nucleus of the atom? 13. What are defects in Rutherford’s atomic model? 14. Write any two points of Planck’s quantum theory. 15. The energy associated with violet colour is greater than red colour in visible spectrum. Why? 16. –. Prove that E = hc 17. Justify the statement that angular momentum of an electron revolving in orbit is quantized. 18. What is the origin of hydrogen spectrum on the basis of Bohr’s model? 19. What is H line in hydrogen spectrum? Which effect explain these lines? 20. What is fine structure of Hydrogen spectrum? 21. Differentiate between Stark effect and Zeeman effect. 22. Define spectrum. Give its two types. 23. Why atomic spectrum is line spectrum? 24. Write names of spectral series of hydrogen spectrum. 25. What is Lyman series? In which region it lies? 26. Calculate wave number value of Lyman series n1 = 1 and n2 = 3. 27. What is Moseley’s law? 28. What is importance of Moseley law? 29. Derive de-Broglie equation = h/mv 30. Bohr theory versus de Broglie equation. 31. How Davission and Germer proved dual nature of matter? / How the wave or dual nature of electrons was verified experimentally? 32. State Heisenberg’s uncertainty principle. 33. Bohr’s theory is in well contradiction with Heisenberg’s uncertainty principle. Justify it. 34. What is the function of Principal quantum number? 35. Define Principal Quantum Number and give its values. 36. What will be the position of electron in an atom when (n + l) value is same for two subshells? 37. State Auf-bau principle. Write electronic configuration of Sodium (11Na) following this principle. 38. State Pauli’s exclusion principle and Hund’s rule. 39. Distribute electrons in orbitals of 29Cu and 20Ca. 40. Write Electronic Configuration of Na = 11 and Cr = 24.  CHAPTER 6 CHEMICAL BONDING Exercise Short Answers Questions – + 1. The species NH2, NH3, NH4 have bond angles of 105o, 107.5o and 109.5o respectively. Justify these values by drawing their structures. 2. Sketch the hybrid orbitals of PCl3, SF6, SiCl4 and NH+4 3. PF3 is a polar molecule which has  = 1.02 D and thus P-F bond is polar. Si is in the proximity of P in periodic table. It is expected that molecule is polar but SiF4 has no dipole moment. Explain it. 4. Bond distance is a compromise distance between two atoms. 5. The distinction between a coordinate covalent bond and a covalent bond vanishes after bond formation in NH+4 , H+3 O and CH3 NH+3. 6. The bond angles of H2O and NH3 are not 109.5o like that of CH4 although oxygen and nitrogen are sp3 hybridized. 7. -bonds are more diffused than -bond. 8. The abnormality of bond length and bond strength is less in HI and prominent in HCl. 9. Sodium chloride does not conduct electricity, but when electric current is passed through molten NaCl or its aqueous solution, electrolysis takes place. 10. The melting points, boiling points, heats of vaporization and heats of sublimation of electrovalent compounds are higher as compared with those of covalent compounds. Previous Boards & Additional Short Answers Questions 1. Why the radius of an atom cannot be determined precisely? 2. Atomic Radii increase in a group and decrease in a period, explain it. 3. What is shielding effect? How does it affect ionization energy? 4. Define ionization energy. How ionization energy varies in periodic table? 5. Why the ionization energy decreases down the group although the nuclear charge increases? 6. Explain that ionization energy is the index of metallic character. 7. Why second ionization energy of an element is always greater than first ionization energy? 8. Define Electron affinity with examples. 9. Why the second electron affinity for all the elements is positive? 10. Electron affinity of fluorine is less than chlorine although electron affinity values decrease down the group. Justify it. 11. Define electronegativity. How does it vary in periodic table? 12. How does the electronegativity difference decide the nature of chemical bond? 13. Define ionic bond with example. 14. No bond in chemistry is 100% ionic in nature. Why? 15. Differentiate between polar and non-polar covalent bond. 16. Why polar bond is stronger than non-polar bond? 17. Ionic bonds are stronger than covalent bonds. Give reason. 18. Define coordinate covalent bond with an example. 19. NH3 can form coordinate covalent bond with H+. Explain. 20. Explain the formation of coordinate covalent bond between NH3 and BF3. 21. Why NH3 can form co-ordinate covalent bond with H+ but CH4 not? 22. Why the distinction between coordinate covalent bond and covalent bond vanishes after bond formation NH+4 ? 23. Write the Lewis structures for the given compounds: (a) HCN (b) CS2 24. What does VSEPR stand for? 25. Write down basic assumption of VSEPR theory. 26. Write down any two postulates of VSEPR theory. 27. Why the lone pair of electrons on an atom occupy more space?  28. Bond angle in NF3(102o) is less than in NH3(107.5o). Justify. 29. Why the bond angle in NH3 (107.5o) is less than tetrahedral angle (109.5o)? 30. Water has bent or angular structure rather than tetrahedral. Justify. 31. Why NH3 molecule and NH+4 ion have different structures? 32. Explain geometry of H2S molecule on the basis of VSEPR theory. 33. Write two points of Valence Bond Theory. 34. A sigma () bond is stronger than a pi () bond. Justify. 35. Differentiate between  and  bond. 36. How Molecular orbital theory justifies that helium atom cannot make the He2? 37. Why MOT is superior to VBT? 38. Differentiate between MOT and VBT? 39. O2 molecule is paramagnetic in nature. Explain. 40. Define bond order. Calculate the bond order of nitrogen molecule. 41. Helium shows diamagnetic behaviour, give reason. 42. On what factors do strength of bond depend? 43. How bond length is affected by change of hybridization state? 44. Why is double bond stronger than a single bond? 45. Define dipole moment. Give its various units. 46. Why the dipole moment of CO2 is zero and that of H2O is 1.85D? 47. Why dipole moment of CO2 is zero but that of SO2 is 1.60D? 48. H2O is an angular molecule whereas CO2 is linear. Why? 49. How dipole moment is helpful to determine the molecular structure? 50. The dipole-moment of CO2 is zero but that of CO is 0.12D. Give reason. 51. Why dipole moment of CH4 is zero? 52. How the percentage ionic character of a covalent bond is determined by dipole moment? OR How the %age ionic character of the polar bond can be determined? 53. Ionic compounds are mostly soluble in water but insoluble in non-polar solvents. Give reason. 54. Reactions between ionic compounds are very rapid. Give reason. 55. How the types of bonding affects solubility of compounds? 56. Ionic compound do not show the phenomena of isomerism but covalent compounds do. Why? 57. Rate of reaction of an ionic compound is faster than covalent compound. Why?  CHAPTER 7 THERMOCHEMISTRY Exercise Short Answers Questions 1. Differentiate between the following: (i) Internal energy and enthalpy (ii) Internal energy change and Enthalpy change (iii) Exothermic reaction and Endothermic reaction (iv) Spontaneous and Non-spontaneous reactions 2. Explain that burning of a candle is a spontaneous process. 3. Is it true that a non-spontaneous process never happens in the universe? Explain it. 4. Is it true that H is equal to E for reaction taking place in solution form? 5. What is the difference between heat and temperature? Write a mathematical relationship between these two parameters. 6. Define heat of neutralization. When a dilute solution of a strong acid is neutralized by a dilute solution of a strong base, the heat of neutralization is found to be nearly the same? How do you account for this? 7. State the laws of thermochemistry and show how are they based on the first law of thermodynamics? 8. What is a thermochemical equation? Give three examples. What information do they convey? 9. Why is it necessary to mention the physical states of reactants and products in a thermo chemical reaction? Apply Hess’s law to justify your answer? 10. Hess’s law helps us to calculate the heats of those reactions which cannot be normally carried out in the laboratory. Explain it. 11. Justify that heat of formation of compound is the sum of all the other enthalpies. 12. What is the meaning of term enthalpy of ionization? 13. Explain what is meant by the terms. (i) Atomization Energy (ii) Lattice Energy Previous Boards & Additional Short Answers Questions 1. Spontaneous reactions are generally exothermic in nature. Explain. 2. Spontaneous reaction always proceeds in the forward direction. Give reason. 3. Burning of natural gas is spontaneous reaction. Justify. 4. Some spontaneous reactions require energy to start. Give example. 5. What is internal energy? What is effect of increase in internal energy on the system? 6. Define first law of thermodynamics. How it is represented? 7. Prove that change in enthalpy is equal to heat of reaction. or Prove that H = qp 8. Prove that E = qv 9. Enthalpy is a state function. Justify. 10. Why is enthalpy of neutralization of strong acid and strong base is always –57.4 kJ mol−1? 11. Differentiate between enthalpy of neutralization and enthalpy of combustion. 12. How do we determine the H in the laboratory for food, fuel etc.? 13. State Hess’s Law. 14. Differentiate between atomization energy and Lattice energy. 15. Born Haber’s Cycle is another form of Hess’s Law. Justify. 16. What are thermo chemical reactions? Give their types?  CHAPTER 8 CHEMICAL EQUILIBRIUM Exercise Short Answers Questions 1. Reversible reactions attain the position of equilibrium which is dynamic in nature and not static. Explain it. or Justify that chemical equilibrium is dynamic in nature. 2. Why do the rates of forward reactions slow down when a reversible reaction approaches the equilibrium stage? 3. The rate of decrease of concentration of any of the reactants and rate of increase of concentration of any of the products may or may not be equal, for various types of reactions, before the equilibrium time. Explain it. 4. The change in volume disturbs the equilibrium position for some of the gaseous phase but not the equilibrium constant. 5. The change of temperature disturbs both equilibrium position and equilibrium constant of a reaction. 6. The solubility of glucose in water is increased by increasing temperature. 7. What is an ionic product of water? How does this value vary with the change in temperature? Is it true that its value increase 75 times when the temperature of water is increased from 0oC to 100oC? 8. What is the justification for the increase of ionic product with temperature? 9. How would you prove that at 25oC, 1dm3of water contains 10–7 moles of H3O+ and 10–7 moles of OH–. 10. Define pH and pOH. How are they related with pKw? 11. What happens to the acidic and basic properties of aqueous solutions when pH varies from zero to 14? 12. It is true that sum of pKa and pKb is always equal to 14 at all temperatures for any acid? If not, why? 13. Acetic acid dissolves in water and gives proton to water, but when dissolved in H2SO4 , it accepts protons. Discuss the role of acetic acid in both cases. 14. In the equilibrium: PCl5(g) PCl3(g) + Cl2(g) H = 90kJ mol–1 What is the effect on; (a) the position of equilibrium (b) equilibrium constant, If: (i) temperature is increased (ii) volume of the container is decreased (iii) catalyst is added (iv) chlorine is added 15. Synthesis of ammonia by Haber’s process is an exothermic reaction. N2(g) + 3H2(g) 2NH3(g) H = –92.46 kJ What should be the possible effect of change of temperature at equilibrium state? How does the change of pressure or volume shift the equilibrium position of this reaction? What is the role of catalyst in this reaction? What happens to equilibrium position of this reaction if NH3 is removed from the reaction vessel from time to time? Sulphuric acid is the king of chemicals. It is produced by the burning of SO2 to SO3 through an exothermic reversible process. 16. Write the balanced reversible reaction. 17. What is the effect of pressure change on this reaction? 18. Reaction is exothermic but still the temperature of 400 – 500oC is required to increase the yield of SO3. Give reasons. Previous Boards & Additional Short Answers Questions 1. Write the relationship of Kp and Kc. 2. How some reactions are affected by volume at equilibrium stage? 3. What will be the effect of change in pressure on NH3 synthesis? 4. Give optimum conditions for synthesis of Ammonia gas by Haber’s process. 5. What is effect of temperature on following system at equilibrium? 2SO2(g) + O2(g) 2SO3(g) H = –194 kJ/mol 6. Briefly explain effect of pressure on the equilibrium position for dissociation of PCl5. 7. What is the effect of catalyst on equilibrium? 8. How pH and pOH are related with each other? 9. What will be nature of solution when: (a) pH = 3.0 (b) pH = 8.0 10. Prove that pKa + pKb = 14 at 25oC. 11. Why water is a weak electrolyte? 12. Write the relationship of pH and pOH with pKw. 13. What is ionic product of water? Write its value at 25oC. 14. What are the factors affecting ionization of acids? 15. What is relationship of pKb with the strength of base? 16. What is ionization constant of acids? 17. What is the formula to calculate the percentage ionization of weak acid? 18. Define pKa and pKb. 19. How common ion effect helps us to purify commercial table salt? 20. Discuss the effect of common ion on solubility. 21. Discuss effect of common ion on solubility of sparingly soluble salt with one example. 22. Explain that a mixture of NH4OH and NH4Cl gives us basic buffer. 23. Define Buffer capacity. 24. What is acidic buffer? Give one example. 25. Write down Henderson’s equation for acidic buffer and basic buffer. 26. State Le Chatelier’s Principle.  CHAPTER 9 SOLUTIONS Exercise Short Answers Questions 1. What are the concentration units of solutions? Compare molar and molal solution? 2. One has one molal NaCl solution and one molal glucose solution. 3. The concentration in terms of molality is independent of temperature but molarity depends upon temperature. 4. The sum of mole fractions of all the components in a solution is equal to unity. 5. 100 g of 98 % H2SO4 has a volume 54.34 cm3 of H2SO4 (density = 1.84 g cm−3) 6. Relative lowering of vapour pressure is independent of temperature. 7. Colligative properties are obeyed when the solute is non-electrolyte and also when the solutions are dilute? 8. The total volume of the solution by mixing 100 cm3 of water with 100 cm3 of ethyl alcohol may 9. not be equal to 200 cm3. Justify it. 10. One molal solution of urea in water is dilute as compared to one molar urea solution. Although both contains same number of solute particles. 11. Non-ideal solutions do not obey the Raoult’s law: 12. The solutions showing positive and negative deviation cannot be fractionally distilled at their specific compositions. Why? 13. What is the physical significance of Kb and Kf. 14. Boiling point of solvent increases due to the presence of solutes. 15. Freezing points are depressed due to the presence of solutes. 16. Boiling point of 1 molal urea solution is 100.52oC but boiling point of 2 molal urea solution is less than 101.04oC? 17. Beckmann’s thermometer is used to note the depression of freezing point. 18. In summer antifreeze solutions protect the liquid of radiator from boiling over? 19. NaCl and KNO3 are used to lower the melting point of ice? Or Why the NaCl and KNO3 are used to form freezing mixture? Previous Boards & Additional Short Answers Questions 1. How will you prepare 0.2 M NaOH solution? 2. Define molality. Give its formula. 3. Define molarity. How is molarity related to mass of solute? 4. What is upper consulate temperature or critical solution temperature? 5. What are conjugate solutions? Give an example. 6. Write two differences between ideal and non-ideal solutions. 7. What are azeotropic mixtures? 8. Differentiate between zeotropic and azeotropic mixtures. 9. What is discontinuous solubility curve? Give one example. 10. Define colligative properties. Name some important colligative properties. 11. Why the freezing point of the solution is always less than the freezing point of pure solvent? 12. Solid ice at 0oC can be melted by applying pressure without supply of heat from outside. Why? 13. Give two applications of colligative properties. 14. Define molal boiling point constant. Give one example. 15. Make difference between Hydration and Hydrolysis. 16. Explain why CuSO4 give acidic solution when put in water? 17. Define hydrolysis. Give chemical equation for hydrolysis of ammonium chloride. 18. What is meant by hydrate? Give formulas of any two hydrates. 19. Why heat of hydration of Li+ is greater than that of Cs+? 20. Define Heat of solution. Give example. 21. Why aqueous solution of CH3COONa is basic? 22. Why a non-volatile solute in a volatile solvent lowers the vapour pressure of solution? 23. Give any two points which show the ideality of a solution. / Write any two properties of ideal solution. 24. Define hydration energy. On what factors do it depend.  CHAPTER 10 ELECTROCHEMISTRY Exercise Short Answers Questions 1. Explain the differences between Ionization and electrolysis. 2. Explain the difference between conduction through metals and molten electrolytes. 3. Spontaneity of oxidation-reduction reaction. 4. Lead accumulator, its desirable and undesirable features. 5. A porous plate or a salt bridge is not required in lead storage cell. 6. Standard oxidation potential of Zn is 0.76 volts and reduction potential −0.76 volts. 7. Na and K can displace hydrogen from acids but Pt, Pd, Cu cannot why? 8. A salt bridge maintains the electrical neutrality in the cell. 9. Lead accumulator is a chargeable battery. 10. Impure Cu can be purified by an electrolytic process. 11. SHE acts as anode when connected with Cu-electrode and as cathode when connected to Zn-electrode. Previous Boards & Additional Short Answers Questions 1. Define oxidation state with two examples. 2. Define oxidation number. Calculate oxidation number of “Mn” in KMnO4. 3. Calculate the oxidation number of sulphur in H2SO4 and H2S. 4. The oxidation state of oxygen is +2 in OF2. Justify it. 5. Differentiate between oxidation and reduction. 6. What are redox reactions? 7. Differentiate between oxidizing agent and reducing agent. 8. Write down electrode reactions occurring during electrolysis of aqueous sodium nitrate. 9. Write down electrochemical equations involved in the electrolysis of molten sodium chloride. 10. How is aluminum anodized in an electrolytic cell? 11. How is voltaic cell represented? 12. Write two functions of salt bridge. 13. Why zinc gives up (oxidizes) electrons and copper metal takes up them (reduces) in galvanic cell? 14. How can we say that a voltaic cell is a reversible cell? 15. What is Standard electrode potential? 16. Draw the diagram of SHE. 17. Give any two applications of electrochemical series. 18. Zn can displace hydrogen from dilute acid solution but copper cannot. Justify. 19. How feasibility of reaction can be predicted from electrochemical series? 20. How does electrochemical series explain the displacement of one metal by another from its solution? 21. Why alkali metals react vigorously with water while coinage metals not react? 22. What is the difference between a cell and a battery? 23. Differentiate between primary and secondary cells. 24. What are secondary cells? Write name of any two such cells. 25. What are electrode reactions of an alkaline dry cell? 26. Write chemical reactions taking place in NICAD cell. 27. Write down chemical reactions taking place in alkaline battery. 28. Write reactions taking place at anode and cathode in silver oxide battery. 29. What is fuel cell and where it is used?  CHAPTER 11 REACTION KINETICS Exercise Short Answers Questions 1. Differentiate between rate of reaction and rate constant. 2. Differentiate between homogeneous and heterogeneous catalyses. 3. Differentiate between fast step and rate determining step: 4. Differentiate between enthalpy change of reaction and energy of activation of the reaction. 5. Rate of reaction is an ever changing parameter. 6. Reaction rate decreases every moment but the rate constant “k” of a reaction is a constant quantity in given conditions. 7. 50% of hypothetical 1st order reaction completes in one hour. The remaining 50% needs more than one hour to complete. 8. The radioactive decay is always a 1st order reaction. 9. The unit of rate constant of a 2nd order reaction is dm3 mol−1 s−1 but the unit of rate of reaction is mol dm−3 s−1. 10. The sum of coefficients of a balanced equation is not necessarily important to give order of reaction. 11. The order of reaction is obtained from the rate expression but the rate expression is obtained from experiment. 12. Change of physical state of a catalyst at the end of reaction. 13. A very small amount of a catalyst may prove sufficient to carry out a reaction. 14. A finely divided catalyst may prove more effective. 15. What is the effect of catalyst on equilibrium position? 16. A catalyst is specific in action. Previous Boards & Additional Short Answers Questions 1. Define order of reaction and specific rate constant. 2. Define order of reaction. Write the names. 3. What is a pseudo first order reaction? 4. What is a zero order reaction? 5. The order of reaction may be in fraction. Explain. 6. Differentiate between molecularity and order of reaction. 7. Define half-life period of a reaction. Give one example. 8. What is rate determining step? 9. Name four physical methods for the determination of a chemical reaction. 10. Name the factors which effect the rate of reaction. 11. How does nature of reactants affect rate of reaction? Give an example. 12. What is the effect of temperature on activation energy of a reaction? 13. How surface area affects the rate of a reaction? 14. How does the increase of temperature increase the rate of the chemical reactions? 15. Why reactions having lower energies of activation have faster rates? 16. What is energy of activation? Give its significance. 17. What is a promoter or activator? Give an example. 18. What is catalysis? Give its two types. 19. Define activation energy and activated complex. 20. Write two properties of enzyme catalysis. 21. What is negative catalyst? Give one example.