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GOVERNMENT HIGHER SECONDARY SCHOOL-KEERIPATTY
XII- CHEMISTRY – QUESTION BANK
UNIT-1
1. What are the differences between minerals and ores?
2. What are the various steps involved in extraction of pure metals from their ores?
3. What is the role of limestone in the extraction of iron from its oxide Fe2O3?
4. Which type of ores can be concentrated by froth floatation method? Give two
examples for such ores.
5. Out of coke and CO, which is better reducing agent for the reduction of ZnO?
Why?
6. Describe a method for refining nickel.
7. Explain zone refining process with an example using the Ellingham diagram
given below
8. (A) Predict the conditions under which
(i) Aluminium might be expected to reduce magnesia.
(ii) Magnesium could reduce alumina.
(B) Carbon monoxide is more effective reducing agent than carbon below
983K but, above this temperature, the reverse is true –Explain.
(c) it is possible to reduce Fe2O3 by coke at a temperature around 1200K
9. Give the uses of zinc.
10. Explain the electrometallurgy of aluminium.
11. Explain the following terms with suitable examples. (i) Gangue (ii) slag
12. Give the basic requirement for vapour phase refining.
13. Describe the role of the following in the process mentioned.
(i) Silica in the extraction of copper
(ii) Cryolite in the extraction of aluminium.
(iii) Iodine in the refining of Zirconium.
(iv) Sodium cyanide in froth floatation.
14. Explain the principle of electrolytic refining with an example.
15. The selection of reducing agent depends on the thermodynamic factor: Explain
with an example.
16. Give the limitations of Ellingham diagram.
17. Write a short note on electrochemical principles of metallurgy.
UNIT-2
1. Write a short note on anamolous properties of the first element of p-block.
2. Describe briefly allotropism in p- block elements with specific reference to carbon.
3. Boron does not react directly with hydrogen. Suggest one method to prepare
diborane from BF3.
4. Give the uses of Borax.
5. What is catenation ? describe briefly the catenation property of carbon.
6. Write a note on Fisher tropsch synthesis.
7. Give the structure of CO and CO2.
8. Give the uses of silicones.
9. AlCl3 behaves like a lewis acid. Substantiate this statement.
10. Describe the structure of diborane.
11. Write a short note on hydroboration.
12. Give one example for each of the following
(i) icosogens (ii) tetragen (iii) prictogen (iv) chalcogen
13. Write a note on metallic nature of p-block elements.
14. How will you identify borate radical?
15. Write a note on zeolites.
16. How will you convert boric acid to boron nitride?
17. A hydride of 2nd period alkali metal (A) on reaction with compound of Boron
(B) to give a reducing agent (C). identify A , B and C.
18. A double salt which contains fourth period alkali metal (A) on heating at
500K gives (B). aqueous solution of (B) gives white precipitate with BaCl2 and
gives a red colour compound with alizarin. Identify A and B.
19. CO is a reducing agent . justify with an example
UNIT-3
1.What is inert pair effect?
2. Chalcogens belongs to p-block. Give reason.
3. Explain why fluorine always exhibit an oxidation state of -1?
4. Give the oxidation state of halogen in the following.
a) OF2 b) O2F2 c) Cl2O3 d) I2O4
5. What are interhalogen compounds? Give examples.
6. Why fluorine is more reactive than other halogens?
7. Give the uses of helium.
8. What is the hybridisation of iodine in IF7? Give its structure.
9. Give the balanced equation for the reaction between chlorine with cold NaOH
and hot NaOH.
10. How will you prepare chlorine in the laboratory?
11. Give the uses of sulphuric acid.
12. Give a reason to support that sulphuric acid is a dehydrating agent.
13. Write the reason for the anamolous behaviour of Nitrogen.
14. Write the molecular formula and structural formula for the following
molecules. a) Nitric acid b) dinitrogen pentoxide c) phosphoric acid
d) phosphine
15. Give the uses of argon.
16. Write the valence shell electronic configuration of group-15 elements.
17. Give two equations to illustrate the chemical behaviour of phosphine.
18. Give a reaction between nitric acid and a basic oxide.
19. What happens when PCl5 is heated?
20. Suggest a reason why HF is a weak acid, whereas binary acids of the all
other halogens are strong acids.
21. Deduce the oxidation number of oxygen in hypofluorous acid – HOF.
22. What type of hybridisation occur in a) BrF5 b) BrF3
UNIT-4
1. What are transition metals? Give four examples.
2. Explain the oxidation states of 4d series elements.
3. What are inner transition elements?
4. Justify the position of lanthanides and actinides in the periodic table.
5. What are actinides? Give three examples.
6. Why Gd3+ is colourless?
7. Explain why compounds of Cu2+ are coloured but those of Zn2+ are colourless.
8. Describe the preparation of potassium dichromate.
9. What is lanthanide contraction and what are the effects of lanthanide
contraction?
10. What are interstitial compounds?
11. Calculate the number of unpaired electrons in Ti3+, Mn2+ and calculate the spin
only magnetic moment.
12. Write the electronic configuration of Ce4+ and Co2+.
13. Explain briefly how +2 states becomes more and more stable in the first half
of the first row transition elements with increasing atomic number.
14. Which is more stable? Fe3+ or Fe2+ explain.
15. Explain the variation in E0M
3+/M2+ 3d series.
16. Compare lanthanides and actinides.
17. Explain why Cr2+ is strongly reducing while Mn3+ is strongly oxidizing.
18. Compare the ionization enthalpies of first series of the transition elements.
19. Actinoid contraction is greater from element to element than the lanthanoid
contraction, why?
20. Out of Lu(OH)3 and La(OH)3 which is more basic and why?
21. Why europium (II) is more stable than Cerium (II)?
22. Why do zirconium and Hafnium exhibit similar properties?
23. Which is stronger reducing agent Cr2+ or Fe2+?
24. The E0M
2+/M value for copper is positive. Suggest a possible reason for this.
25. predict which of the following will be coloured in aqueous solution Ti2+ , V3+ , Sc4+, Cu+,Sc3+, Fe3+, Ni2+ and Co3+
26. Describe the variable oxidation state of 3d series elements.
27. Which metal in the 3d series exhibits +1 oxidation state most frequently and
why?
28. Why first ionization enthalpy of chromium is lower than that of zinc?
29. Transition metals show high melting points why?
UNIT-5
1. Write the IUPAC names for the following complexes.
i) Na2[Ni(EDTA)] ii) [Ag (CN)2]-
iii) [Co(en)3] (SO⁴)3
iv) [Co(ONO)(NH3)5]2+
v) [Pt(NH3)2Cl (NO²)]
2. Write the formula for the following coordination compounds.
a) potassiumhexacyanidoferrate(II) b) pentacarbonyliron(0)
c) pentaamminenitrito −ҝ −N-cobalt(III)ion d) hexaamminecobalt(III)sulphate
e) sodiumtetrafluoridodihydroxidochromate(III)
3. Arrange the following in order of increasing molar conductivity
i) Mg[Cr(NH3) Cl5] ii) [Cr(NH3)5Cl]3 [CoF6]2 iii) [Cr(NH3)3Cl3]
4. Ni2+ is identified using alcoholic solution of dimethyl glyoxime. Write the
structural formula for the rosy red precipitate of a complex formed in the
reaction.
5. [CuCl4]2-
exists while [CuI4]2- does not exist why?
8. Give an example of coordination compound used in medicine and two examples
of biologically important coordination compounds.
9. Based on VB theory explain why [Cr(NH3)6]
3+ is paramagnetic, while [Ni(CN)4]2-
is diamagnetic.
9. Draw all possible geometrical isomers of the complex [Co(en)2 Cl2]+ and identify
the optically active isomer.
10. [Ti(H2O)6]
3+ is coloured, while [Sc(H2O)6]
3+ is colourless- explain.
11. Give an example for complex of the type [Ma2b2c2] where a, b, c are
monodentate ligands and give the possible isomers.
12. Give one test to differentiate [Co(NH3)5Cl]SO4 and [Co(NH3)5SO4]Cl.
13. In an octahedral crystal field, draw the figure to show splitting of d orbitals.
14. What is linkage isomerism? Explain with an example.
15. Write briefly about the applications of coordination compounds in volumetric
analysis.
16. Classify the following ligand based on the number of donor atoms.
a) NH3 b) en c) ox2- d) triaminotriethylamine e) pyridine
17. Give the difference between double salts and coordination compounds.
18. Write the postulates of Werner’s theory.
19. [Ni(CN)4]2-
is diamagnetic, while [NiCl4]2-
is paramagnetic ,explain using
crystal field theory.
20. Why tetrahedral complexes do not exhibit geometrical isomerism.
21. Explain optical isomerism in coordination compounds with an example.
22. What are hydrate isomers? Explain with an example.
23. What is crystal field splitting energy?
24. What is crystal stabilization energy (CFSE) ?
25. A solution of field [ Ni(H2O)6]
2+ is green, whereas a solution of [Ni(CN)4]2-is colorless - Explain
26. Discuss briefly the nature of bonding in metal carbonyls.
27. What is the coordination entity formed when excess of liquid ammonia is added to an aqueous solution of copper sulphate?
28. On the basis of VB theory explain the nature of bonding in [Co(C2O4)3]3-
29. What are the limitations of VB theory?
30. Write the oxidation state, coordination number , nature of ligand, magnetic
property and electronic configuration in octahedral crystal field for the complex
K4[Mn(CN)6] .
UNIT-6
1. Define unit cell.
2. Give any three characteristics of ionic crystals.
3. Differentiate crystalline solids and amorphous solids.
4. Classify the following solids a. P4 b. Brass c. diamond d. NaCl e. Iodine
5.Explain briefly seven types of unit cell.
6. Distinguish between hexagonal close packing and cubic close packing.
7. Distinguish tetrahedral and octahedral voids.
8. What are point defects?
9. Explain Schottky defect.
10. Write short note on metal excess and metal deficiency defect with an example.
11. Calculate the number of atoms in a fcc unit cell.
12. Explain AAAA and ABABA and ABCABC type of three dimensional packing
with the help of neat diagram.
13. Why ionic crystals are hard and brittle?
14. Calculate the percentage efficiency of packing in case of body centered cubic
crystal.
15. What is the two dimensional coordination number of a molecule in square close
packed layer?
16. Experiment shows that Nickel oxide has the formula Ni0.96O1.00. What fraction
of Nickel exists as of Ni
2+ and Ni3+ ions?
17. What is meant by the term “coordination number”? What is the coordination
number of atoms in a bcc structure?
18. Write a note on Frenkel defect.
UNIT-7
1. Define average rate and instantaneous rate.
2. Define rate law and rate constant.
3. Derive integrated rate law for a zero order reaction. A→product.
4. Define half life of a reaction. Show that for a first order reaction half life is
independent of initial concentration.
5. What is an elementary reaction? Give the differences between order and
molecularity of a reaction.
6. Explain the rate determining step with an example.
7. Describe the graphical representation of first order reaction.
8. Write the rate law for the following reactions.
(a) A reaction that is 3/2 order in x and zero order in y.
(b) A reaction that is second order in NO and first order in Br2.
9. Explain the effect of catalyst on reaction rate with an example.
10. The rate law for a reaction of A, B and C has been found to be
rate =K [A]2 [B][L]
3/2 . How would the rate of reaction change when
(i) Concentration of [L] is quadrupled
(ii) Concentration of both [A] and [B] are doubled
(iii) Concentration of [A] is halved
(iv) Concentration of [A] is reduced to (1/3) and concentration of [L] is
quadrupled.
11. Explain briefly the collision theory of bimolecular reactions.
12. Write Arrhenius equation and explains the terms involved.
13. Explain pseudo first order reaction with an example.
14. How do concentrations of the reactant influence the rate of reaction?
15. How do nature of the reactant influence rate of reaction?
16. The rate constant for a first order reaction is 1.54 × 10-1 . Calculate its half life
time.
UNIT-8
1. What are lewis acids and bases? Give two examples for each.
2. Discuss the Lowery – Bronsted concept of acids and bases.
3. Indentify the conjugate acid base pair for the following reaction in aqueous
solution
4. Account for the acidic nature of HClO4 . In terms of Bronsted – Lowry theory,
identify its conjugate base.
5. When aqueous ammonia is added to CuSO4 solution, the solution turns deep
blue due to the formation of tetramminecopper (II) complex, among H2O and
NH3 Which is stronger Lewis base? [Cu(H2 O)4 ]2++4NH3 (aq) → [Cu(NH3 )4 ]2+
(aq)
6. Define solubility product
7. Define ionic product of water. Give its value at room temperature.
8. Explain common ion effect with an example
9. Derive an expression for Ostwald’s dilution law
10. Define pH
11. Derive an expression for the hydrolysis constant and degree of hydrolysis of
salt of strong acid and weak base
12. Write the expression for the solubility product of Ca3(PO4 )2 .
13. Write the expression for the solubility product of Hg2 Cl2 .
UNIT-9
1. Define anode and cathode
2. Why does conductivity of a solution decrease on dilution of the solution?
3. State Kohlrausch Law. How is it useful to determine the molar conductivity of
weak electrolyte at infinite dilution?
4. Describe the electrolysis of molten NaCl using inert electrodes.
5. State Faraday’s Laws of electrolysis.
6. Describe the construction of Daniel cell. Write the cell reaction.
7. Why is anode in galvanic cell considered to be negative and cathode positive
electrode?
8. Which of 0.1M HCl and 0.1 M KCl do you expect to have greater and why?
9. Arrange the following solutions in the decreasing order of specific conductance.
i) 0.01M KCl ii) 0.005M KCl iii) 0.1M KCl iv) 0.25 M KCl v) 0.5 M KCl
10. Why is AC current used instead of DC in measuring the electrolytic
conductance?
11. Two metals M1 and M2 have reduction potential values of -xV and +yV
respectively. Which will liberate H2 and H2SO4.
12. Derive an expression for Nernst equation
13. Write a note on sacrificial protection.
14. Explain the function of H2 – O2 fuel cell.
15.Ionic conductance at infinite dilution of Al3+ and SO42-
