7.1 Discuss the general characteristics of Group 15 elements with reference to their electronic configuration, oxidation state, atomic size, ionisation enthalpy and electronegativity.
7.2 Why does the reactivity of nitrogen differ from phosphorus?
7.3 Discuss the trends in chemical reactivity of group 15 elements.
7.4 Why does NH3 form hydrogen bond but PH3 does not?
7.5 How is nitrogen prepared in the laboratory? Write the chemical equations of the reactions . involved.
7.6 How is ammonia manufactured industrially?
7.7 Illustrate how copper metal can give different products on reaction with HN03.
7.8 Give the resonating structures of N02 and N2O5.
7.9 The HNH angle value is higher than HPH, H AsH and HSbH angles. Why?
7.10 Why does R3P = O exist but R3N = O does not (R = alkyl group)?
7.11 Explain why NH3 is basic while BiH3 is only feebly basic.
7.12 Nitrogen exists as diatomic molecule and phosphorus as P4. Why?
7.13 Write main differences between the properties of white phosphorus and red phosphorus.
7.14 Why does nitrogen show catenation properties less than phosphorus?
7.15 Give the disproportionation reaction of H3 P03 .
7.16 Can PCl5 act as an oxidising as well as a reducing agent Justify.
7.17 Justify the placement of O, S, Se, Te and Po in the same group’of the periodic table in terms of electronic configuration, oxidation state and hydride formation.
7.18 Why is dioxygen a gas but sulphur a solid?
7.19 Knowing the electron gain enthalpy values of O—>O– and O—>O2- as -141 and 702 kJ mol-1 respectively, how can you account for [he formation of a large number of oxides having O2- species and not O–?
7.20 Which aerosols deplete ozone?
7.21 Describe the manufacture of H2SO4 by contact process?
7.22 How is SO2 an air pollutant?
7.23 Why are halogens strong oxidising agents?
7.24 Explain why fluorine forms only one oxoacid, HOF.
7.25 Explain why inspite of nearly the same electronegativity, nitrogen forms hydrogen bonding while chlorine does not.
7.26 Write two uses of ClO2
7.27 Why are halogens coloured?
7.28 Write the reactions of F2 and Cl2 with water.
7.29 How can-you prepare Cl2 from HCl and HCl from CI2? Write reactions only.
7.30 What inspired N. Bartlett for carrying out reaction between Xe and PtF6?
7.31 What are the oxidation states of phosphorus in the following: –
7.32 Write balanced equations for the following:
7.33 How are xenon fluorides XeF2, XeF4 and XeF6 obtained?
7.34 With what neutral molecule is CIO– isoelectronic? Is that molecule a Lewis base?
7.35 How are XeO3 and XeOF4prepared?
7.36 Arrange the following in the order of property indicated for each set: –
7.37 Which one of the following does not exist ?
7.38 Give the formula and describe the structure of a noble gas species which is isostructural with: (i) ICI4– (ii) IBr2– (iii) Br03–
7.39 Why do noble gases have comparatively large atomic sizes?
7.40 List the uses of neoirand argon gases.
NCERT Solutions For Class 12 Chemistry Chapter 6 General Principles and Processes of Isolation of Elements
6.1 Copper can be extracted by hydrometallurgy but not zinc. Explain.
6.2.What is the role of depressant in froth-floatation process?
6.3 Why is the extraction of copper from pyrites more difficult than that from its oxide ore through reduction?
6.5 Out of C and CO which is a better reducing agent at 673 K?
6.6.Name the common elements present in the anode mud in electrolytic refining of copper. Why are they so present?
6.7 Write down the reactions taking place in different zones in the blast furnace during the extraction of iron.
6.8 Write chemical reactions taking place in the extraction of zinc from zinc blende.
6.9.State the role of silica in the metallurgy of copper.
6.10.What is meant by the term “chromatography”?
6.11.What criterion is followed for the selection of the stationary phase in chromatography?
6.12 Describe a method for refining nickel.
6.13 How can you separate alumina from silica in a bauxite ore associated with silica? Give equations, if any.
6.14 Giving examples, differentiate between ‘roasting’ and ‘calcination’.
6.15 I low is ‘cast iron’ different from ‘pig iron ’?
6.16 Differentiate between “minerals” and “ores’.
6.17 Why copper matte is put in silica lined converter?
6.18 What is the role of cryolite in the metallurgy of aluminium?
6.19 How is leaching carried out in case of low grade copper ores?
6.20 Why is zinc not extracted from zinc oxide through reduction using CO?
6.21 The value of ΔfG° for formation of Cr2O3 is – 540 kJ mol-1 and that of Al203 is – 827 kJ mol-1 . Is the reduction of Cr2O3 possible with Al?
6.22 Out of C and CO, which is a better reducing agent for ZnO?
6.23 The choice of a reducing agent in a particular case depends on thermodynamic factor. How far do you agree with this statement? Support your opinion with two examples.
6.24 Name the processes from which chlorine is obtained as a by-product What will happen if an aqueous solution of NaCl is subjected to electrolysis?
6.25 What is the role of graphite rod in the electrometallurgy of aluminium?
6.26 Outline the principles of refining of metals by the following methods:
6.27 Predict conditions under which Al might be expected to reduce MgO.
5.3. Give reason why a finely divided substance is more effective as an adsorbent?
5.4. What are the factors which influence the adsorption of a gas on a solid?
5.5. What is an adsorption isotherm? Describe Freundlich adsorption isotherm.
5.6. What do you understand by activation of adsorbent? How is it achieved?
5.7. What role does adsorption play in heterogeneous catalysis?
5.8. Why is adsorption always exothermic?
5.9. How are the colloidal solutions classified on the basis of physical states of the dispersed phase and dispersion medium?
5.10. Discuss the effect of pressure and temperature on the adsorption of gases on solids.
5.11. What are lyophilic and lyophobic sols? Give one example of each type. Why are hydrophobic sols easily coagulated?
5.12. What is the difference between multimolecular and macromolecular colloids? Give one example of each.
5.13. What are enzymes? Write in brief the mechanism of enzyme catalysis.
5.14. How are colloids classified on the basis of
5.15. Explain what is observed
5.16. What are emulsions? What are their different types? Give example of each type.
5.17. How do emulsifires stabilise emulsion? Name two emulsifiers.
5.18. Action of soap is due to emulsification and micelle formation. Comment
5.19. Give four examples of heterogeneous catalysis.
5.20. What do you mean by activity and selectivity of catalysts?
4.1 From the rate expression for the following reactions determine their order of reaction and the dimensions of the rate constants:
4.2 For the reaction, 2A + B ————> A2 B, the rate = k [AJ[B]2 with k = 2.0 x 10-6 mol-2 L2 s-1. Calculate the initial rate of the reaction when [A] = 0.1 mol L-1, [B] = 0.2 mol L-1. Calculate the rate of reaction after [A] is reduced to 0.06 mol L-1.
4.3 The decomposition of NH3 on platinum surface is zero order reaction. What are the rates of production of N2 and H2 if Ar=2.5 x 10-4 mol-1 Ls-1.
4.4 The decomposition of dimethyl ether leads to the formation of CH4, H2 and CO and die reaction, rate is given by Rate=k [CH3OCH3]3/2 The rate of reaction is followed by increase in pressure in a closed vessel, so the rate can also, be expressed in terms of the partial pressure of dimethyl ether, i.e., Rate= k (PCH3OCH3)3/2
4.5 Mention the factors that affect the rate of a chemical reaction.
4.6 A reaction is second order with respect to a reactant How is the rate of reaction affected if the concentration of the reactant is (i) doubled (ii) reduced to half?
4.7 What is the effect of temperature on the rate constant of a reaction? How can this effect of temperature on rate constant be represented quantitatively?
4.8 In a pseudo first order hydrolysis of ester in water, the following results were obtained:
4.9 A reaction is first order in A and second order in B.
4.10 In a reaction between A and B, the initial rate of reaction (r0 ) was measured for different initial concentrations of A and B as given below:
4.11 The following results have been obtained during the kinetic studies of the reaction.
4.12 The reaction between A and B is first order with respect to A and zero order with respect to B. Fill in the blanks in the following table:
4.13 Calculate the half-life of a first order reaction from their rate constants given below:
4.14 The half-life for radioactive decay of 14C is 5730 years. An archaeological artifact containing wood had only 80% of the 14C found in a living tree. Estimate the age of the sample.
4.15 The experimental data for decomposition of N2O5
3.1 Arrange the following metals in the order in which they displace each other from the solution of their salts: Al, Cu, Fe, Mg and Zn.
3.2 Given the standard electrode potentials, K+/K=-2. 93 V, Ag+/Ag = 0.80 V, Hg2+/Hg =0.79V, Mg2+/Mg=-2.37V, Cr3+/Cr=0.74V.
3.3 Depict the galvanic cell in which the reaction
3.4 Calculate the standard cell potentials of galvanic cell in which the following reactions take place
3.5 Write the Nernst equation and emf of the following cells at 298 K:
3.6 In the button cells widely used in watches and other devices the following reaction takes place:
3.7 Define conductivity and molar conductivity for the solution of an electrolyte. Discuss their variation with concentration.
3.8 The conductivity of 0.20 M solution of KCl at 298 K is 0.0248 S cm-1. Calculate its molar conductivity.
3.9 The resistance of a conductivity cell containing 0.001 M KCI solution at 298 K is 1500 Ω What is the cell constant if conductivity of 0.001 M KCI solution at 298 K is 0.146 x 10-3 S cm-1?
3.10 The conductivity of NaCl at 298 K has been determined at different concentrations and the results are given below:
3.11 Conductivity of 0.00241 M acetic acid is 7.896 x 10-5 S cm-1. Calculate its molar conductivity. If Λm0, for acetic acid is 390.5 S cm2 mol-1, what is its dissociation constant?
3.12 How much charge is required for the following reductions:
3.13 How much electricity in terms of Faraday is required to produce .
3.14 How much electricity is required in coulomb for the oxidation of (i) 1 mol of H2O to 02 (ii) 1 mol of FeO to Fe203
3.15 A solution of Ni(N03)2 is electrolyzed between platinum electrodes using a current of 5 amperes for 20 minutes. What mass of Ni is deposited at the cathode?
3.16 Three electrolytic cells A, B, C containing solutions of ZnS04, AgNO3 and CuS04, respectively are connected in series. A steady current of 1.5 amperes was passed through them until 45 g of silver deposited at the cathode of call B. How long did the current flow? What mass of copper and zinc were deposited?
3.17 Predict if the reaction between the following is feasible:
3.18 Predict the products of electrolysis in each of the following.
2.1 Define the terra solution. How many types of solutions are formed? Write briefly about each type with an example.
2.2 Give an example of a solid solution in which the solute is a gas.
2.3 Define the following terms:
2.4 Concentrated nitric acid used in laboratory work is 68% nitric acid by mass in aqueous solution. What should be the molarity of such a sample of the acid if the density of the solution is 1.504 g mL-1?
2.5 A solution of glucose in water is labelled as 10% w/w, what would be the molality and mole fraction of each component in the solution? If the density of solution is 1 .2 g m L-1, then what shall be the molarity of the solution?
2.6 How many mL of 0.1 M HCl are required to react completely with 1 g mixture of Na2C03 and NaHCO3 containing equimolar amounts of both?
2.7 A solution is obtained by mixing 300 g of 25% solution and 400 g of 40% solution by mass. Calculate the mass percentage of the resulting solution.
2.8 An antifreeze solution is prepared from 222.6 g of ethylene glycol, (C2 H6O2 ) and200 g of water. Calculate the molality of the solution. If the density of the solution is 1.072 g mL-1, then what shall be the molarity of the solution?
2.9 A sample of drinking water was found to be severely contaminated with chloroform (CHCl3), supposed to be a carcinogen. The level of contamination was 15 ppm (by mass).
2.10 What role does the molecular interaction play in a solution of alcohol and water?
2.11 Why do gases always tend to be less soluble in liquids as the temperature is raised?
2.12 State Henry’s law and mention some important applications.
2.13 The partial pressure of ethane over a solution containing 6.56 × 10-3 g of ethane is 1 bar. If the solution contains 5.00 × 10-2 g of ethane, then what shall be the partial pressure of the gas?
2.14 What is meant by positive and negative deviations from Raoult’s law and how is the sign of AmixH related to positive and negative deviations from Raoult’s law?
2.15 An aqueous solution of 2% non-volatile solute exerts a pressure of 1.004 bar at the normal boiling point of the solvent. What is the molar mass of the solute?
2.16 Heptane and octane form an ideal solution. At 373 K, the vapour pressures of the two liquid components are 105.2 kPa and 46.8 kPa respectively. What will be the vapour pressure of a mixture of 26.0 g of heptane and 35.0 g of octane?
2.17 The vapour pressure of water is 12.3 kPa at 300 K. Calculate vapour pressure of 1 molal solution of a non-volatile solute in it
2.18 Calculate the mass of a non-volatile solute (molar mass 40 g mol-1) which should be dissolved in 114 g octane to reduce its vapour pressure to 80%.
2.19 A solution containing 30g of non-volatile solute exactly in 90 g of water has a vapour pressure of 2.8 kPa at 298 K. Further, 18g of water is then added to the solution and the new of vapour pressure becomes 2.9 kPa at 298 K. Calculate
2.20 A 5% solution (by mass) of cane sugar in water has freezing point of 271 K. Calculate the freezing point of 5% glucose in water if freezing point of pure water is 273.15 K.
2.21 Two elements A and B form compounds having formula AB2 and AB4. When dissolved in 20g of benzene (C6H6), 1 g of AB2 lowers the freezing point by 2.3 K whereas 1.0 g of AB4 lowers it by 1.3 K. The molar depression constant for benzene is 5.1 K kg mol-1. Calculate atomic masses of A and B.
2.22 At 300 K, 36g of glucose present in a litre of its solution has an osmotic pressure of 4.08 bar. If the osmotic pressure of the solution is 1.52 bars at the same temperature, what would be its concentration?
2.23 Suggest the most important type of intermolecular attractive interaction in the following pairs:
2.24 Based on solute-solvent interactions, arrange the following in order of increasing solubility in n-octanc and explain.
2.25 Amongst the following compounds, identify which are insoluble, partially soluble and highly soluble in water?
2.26 If the density of some lake water is 1.25 g mL-1 and contains 92g of Na+ ions per kg of water, calculate the molality of Na+ ions in the lake.
2.27 If the solubility product of CuS is 6 x 10-16, calculate the maximum molarity of CuS in aqueous solution.
1.1.Why are solids rigid?
1.2.Why do solids have a definite volume?
1.3.Classify the following as amorphous or crystalline solids: Polyurethane, naphthalene, benzoic acid, Teflon, potassium nitrate, cellophane, polyvinyl chloride, fibreglass, copper
1.4.Why is glass considered a supercooled liquid?
1.5.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?
1.6.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
1.7.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?
1.8.Ionic solids conduct electricity in molten state but not in solid state. Explain
1.9.What type of solids are electrical conductors, malleable and ductile?
1.10.Give the significance of a ‘lattice point’.
1.11. Name the parameters that characterise a unit cell.
1.12. Distinguish between
1.13. Explain how much portion of an atom located at
1.14. What is the two-dimensional coordination number of a molecule in square close-packed layer?
1.15. A compound forms hexagonal close-packed. structure. What is the total number of voids in 0. 5 mol of it? How many of these are tetrahedral voids?
1.16. A compound is formed by two elements M and N. The element N forms ccp and atoms of M occupy l/3rd of tetrahedral voids. What is the formula of the compound?
1.17. Wh ich of the following lattices has the highest packing efficiency (i) simple cubic (ii) body- centred cubic and (iii) hexagonal close-packed lattice?
1.18. An element with molar mass 2.7 x 10-2 kg mol-1 forms a cubic unit cell with edge length 405 pm. If its density is 2.7 x 103 kg m-3, what is the nature of the cubic unit cell?
1.19. What type of defect can arise when a solid is heated? Which physical property is affected by it and in what way?
1.20.What type of stoichiometric defect is shown by:
1.21. Explain how vacancies are introduced in an ionic solid when a cation of higher valence is added as an impurity in it.
1.22.Ionic solids, which have anionic vacancies due to metal excess defect, develop colour. Explain with the help of a suitable example.
1.23.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?
1.24.What type of substances would make better permanent magnets, ferromagnetic or ferrimagnetic. Justify your answer.
1.1. Define the term ‘amorphous’. Give a few examples of amorphous solids.
1.2. What makes a glass different from a solid such as quartz? Under what conditions could quartz be converted into glass?
1.3 Classify each of the following solids as ionic, metallic, modular, network (covalent) or amorphous:
1.4 (i) What is meant by the term ‘coordination number’?
1.5 How can you determine the atomic mass of an unknown metal if you know its density and the dimensions of its unit cell? Explain.
1.6 ‘Stability of a crystal is reflected in the magnitude of its melting points’. Comment. Collect melting points of solid water, ethyl alcohol, diethyl ether and methane from a data book. What can you say about the intermolecular forces between these molecules?
1.7 How will you distinguish between the following pairs of terms:
1.8 How many lattice points are there is one unit cell of each of the following lattices?
1.10 Calculate the efficiency of packing in case of a metal crystal for (i) simple cubic, (ii) body centred cubic, and (iii) face centred cubic (with the assumptions that atoms are touching each other).
1.11 Silver crystallises in fcc lattice. If edge length of the cell is 4.07 x 10-8 cm and density is 10.5 g cm-3, calculate the atomic mass of silver.
1.12 A cubic solid is made up of two elements P and Q. Atoms of Q are at the corners of the cube and P at the body centre. What is the formula of the compound? What are the coordination numbers of P and Q?
1.13 Niobium crystallises in a body centred cubic structure. If density is 8.55 g cm-3, calculate atomic radius of niobium, using its atomic mass 93u.
1.14 If the radius of the octahedral void is r and radius of the atoms in close-packing is R, derive relation between rand R.
1.15 Copper crystallises into a fee lattice with edge length 3.61 x 10-8 cm. Show that the calculated density is in agreement with its measured value of 8.92 gcm-3.
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