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VERY IMPORTANT QUESTIONS FOR CBSE 2011-PHYSICS-XII Q(1) State Gauss’s theorem in electrostatics. Using this theorem , derive an expression for the electric field intensity at a point near a thin infinite plane sheet of charge density σ Cm -2 . Q. (2) Define electric dipole moment . Derive an expression for the electric field intensity at any point along the equatorial line of an electric dipole. Q.(3) Derive an expression for the electric potential at a point along the axial line of the dipole. Mention one contrasting feature of electric potential of a dipole at a point as compared to that due to a single charge. Q.(4) Derive an expression for the potential energy of a dipole in a uniform electric field. Hence discuss the conditions of its stable and unstable equilibrium. Q.(5) Derive an expression for the capacitance of a parallel plate capacitor. Q(6) Derive an expression for the energy stored in a capacitor with air as the medium between its plates. How does the stored energy change if air is replaced by a medium of dielectric constant k ? Q.(7) Give the principle of working of a Van de Graaf Generator.(Diagram) Q.(8) Derive Ohm’s law on the basis of the theory of electron drift. Q(9) Deduce the relation connecting current density (J) and the conductivity (σ) of the conductor, when an electric field E is applied to it.

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VERY IMPORTANT QUESTIONS FOR CBSE 2011-PHYSICS-XIIQ(1) State Gauss’s theorem in electrostatics. Using this theorem , derive an expression for the electric field intensity at a point near a thin infinite plane sheet of charge density σ Cm-2. Q. (2) Define electric dipole moment . Derive an expression for the electric field intensity at any point along the equatorial line of an electric dipole. Q.(3) Derive an expression for the electric potential at a point along the axial line of the dipole. M

### Transcript of Questions for Cbse 2011

VERY IMPORTANT QUESTIONS FOR CBSE 2011-PHYSICS-XII

Q(1) State Gauss’s theorem in electrostatics. Using this theorem , derive an expression for the electric field intensity at a point near a thin infinite plane sheet of charge density σ Cm-2.

Q. (2) Define electric dipole moment . Derive an expression for the electric field intensity at any point along the equatorial line of an electric dipole.

Q.(3) Derive an expression for the electric potential at a point along the axial line of the dipole. Mention one contrasting feature of electric potential of a dipole at a point as compared to that due to a single charge.

Q.(4) Derive an expression for the potential energy of a dipole in a uniform electric field. Hence discuss the conditions of its stable and unstable equilibrium.

Q.(5) Derive an expression for the capacitance of a parallel plate capacitor.

Q(6) Derive an expression for the energy stored in a capacitor with air as the medium between its plates. How does the stored energy change if air is replaced by a medium of dielectric constant k ?

Q.(7) Give the principle of working of a Van de Graaf Generator.(Diagram)

Q.(8) Derive Ohm’s law on the basis of the theory of electron drift.

Q(9) Deduce the relation connecting current density (J) and the conductivity (σ) of the conductor, when an electric field E is applied to it.

Q(10) State Kirchhoff’s law for electrical circuits.

Q(11) Define the term potential gradient .with the help of a circuit diagram, explain how a potentiometer can be used to compare of emfs of two primary cells.

Q(12) State the principle of potentiometer. With the help of a circuit diagram, describe a method to find the internal resistance of a primary cell.

Q(13) What do you mean by the sensitivity of a potentiometer ? How can we increase the sensitivity of a potentiometer ?

Q(14) Draw a circuit diagram which can be used to determine the resistance of a given wire. Explain the principle of the experiment and give the formula used.

Q(15) State Biot-Savart law . By using it find out an expression for magnetic field at a point due to a straight wire carrying current. OR Magnetic field on the axis of a circular current loop.

Q(16) State the principle and draw a labeled diagram of moving coil galvanometer.

Q(17) With the help of a labeled diagram, explain the principle, construction, theory and working of a cyclotron . state its limitations.

Q.(18) How can a moving coil galvanometer be converted into voltmeter and ammeter ?

Q(19) A bar magnet is placed in a uniform magnetic field with its magnetic moment making an angle θ with the field. (i) Write an expression for the torque acting on the magnet and hence define its magnetic moment.

(ii) Write an expression for the potential energy of the magnet in this orientation. When is this energy minimum ?

Q(20) Derive an expression for the induced emf produced by changing the area of a rectangular coil placed perpendicular to a magnetic field.

Q(21) Derive an expression for the self-inductance of a long solenoid of N turns, having a core of relative permeability µr.

Q(22) Using phasor diagram ,derive an expression for the impedance of a series LCR-circuit .what do you mean by resonance condition of such a circuit ?

Q(23) Define power in an ac circuit. Show that the average power transferred to an a.c. circuit is, in general given by Pav Erms x Irms x cosϕ

Q(24) state the principle of Transformer.

Q.(25) Hertz’s experiment for electromagnetic waves & Maxwell’s equations.

Q(26) Write any four electromagnetic spectrum their Wavelength, frequency & two application.

Q(27) Refraction formula at a convex surface when the object lies in rarer medium and the image is real. (µ2/v -µ1/u)= (µ2 -µ1)/R

Q(29) Derive the len’s maker’s formula for a double convex lens.

Q(30) Derive the expression for the refractive index of the material of the prism in terms of the angle of the prism and angle of minimum deviation.

Q(31) Compound microscope & its magnifying power , Ray diagram (i) when image formed at infinity (ii) when image formed at least distance of distinct vision point.

Q(32) Astronomical telescope & its magnifying power , Ray diagram (i) when image formed at infinity (ii) when image formed at least distance of distinct vision point.

Q(33) state Huygens’ principle and use it to construct refracted wavefront for refraction of a plane wave front at a plane refracting surface. Hence derive Snell’s law.

Q(34) Illustrate with the help of suitable diagram, action of the following on a plane wavefront incident on (i) a prism (ii) a convex lens and (iii) a concave mirror.

Q(35) Define fringe width. Derive an expression for fringe width in Young’s double slit experiment .

Q(36) Explain diffraction at a single slit. Derive relation for the linear width of central maximum.

Q(37) Describe an experimental arrangement to study photoelectric effect & derive Einstein’s Photo electric equation.

Q(38) Davisson and Germer experiment & related graphs

Q(39) Spectral series of hydrogen atom related question

Q(40) Draw a graph showing the variation of potential energy as a function of their separation. What is the significance of negative potential energy in this graph ?

Q(41) Binding energy per nucleon Graph or Numerical Problem

Q(42) Numerical related Half –life

Q(43) Explain the formation of energy bands in solids through graph.

Q(44) P-N junction diode & its V-I characteristics curve (forward & Reverse Bias)

Q(45) Half –Wave Rectifier or Full wave Rectifier used as a Junction diode

Q(46) Common Emitter Characteristics curve (N-P-N Input & Output)

Q(47) Common Emitter Amplifier & its various Gains or Use as an Oscillator.

Q(48) One diagram based question on Gates. (digital)

Q(49) Block diagram of Communication System or Modulation Index

Q.(50) Deduce an expression for the distance upto which the TV signals can directly be received from TV tower of Height h. OR Numerical related this formula OR for the maximum line of sight (LOS).

Paper Submitted By: Name: Mr. NAVEEN KUMAR GARG Email: [email protected] No: 09412584464

Q. 1. The unit of equivalent conductance is:

Q. 2. Give the relationship between free energy change and EMF of a cell.

Q. 3. How much amount of a substance is deposited by 1 coulomb?

Q. 4. Write the relationship between molar conductivity and specific conductivity.

Q. 5. What is the reference electrode in determining the standard electrode potential?

Q. 6. Write the use of platinum foil in the hydrogen electrode.

Q. 7. How will you identify whether the given electrolyte is a strong or a weak?

Q. 8. Name the metal that can be used in cathodic protection of iron against rusting.

Q. 9. Which cells were used in the Apollo space program?

Q. 10. What are the units of molar conductivity?

Q. 11. Why it is necessary to use a salt bridge in a galvanic cell?

Q. 12. Define-(i) Faraday’s constant (ii) Electrochemical equivalent

Q. 13. What is corrosion?

Q. 14. What is Galvanisation ?

Q. 15. Define Kohlrausch’s law.

Q. 16. When 3 ampere of electricity is passed for 45 minutes 2.0 g of metal is deposited. Find equivalent weight of metal.

Q. 17. Find the value of equilibrium constant from the following data-

Q. 18. Calculate standard free energy change for the following chemical reaction –

Q. 19. Represent the following cell reactions as galvanic cell–

1.

2.

Q. 20. Define and give one example of each of the following-

1. Primary Cells 2. Secondary Cell

3. Fuel Cell

Q. 21. What is Nernst equation? Write its expression for single electrode and cell.

Q. 22. Find the emf of following cell -

Q. 23. Predict if the following reaction is feasible or not , and

Paper Submitted By: Name: Mr. NAVEEN KUMAR GARG Email: [email protected] No: 09412584464

Q. 1. Define the term bio molecules?

Q. 2. Amino acids are the building blocks of……………….

Q. 3. The molecular weight of proteins is…………………..

Q. 4. Deficiency of vitamin A results in……………………

Q. 5. Charring of sugar is due to

Q. 6. Define Monosaccharides?

Q. 7. What is the difference between Reducing and non-reducing sugars or carbohydrates?

Q. 8. Explain the term mutarotation?

Q. 10. What are the main sources of vitamins?

Q. 11. Give two methods for the preparation of glucose?

Q. 12. Define Carbohydrates? Give their basic classification depending upon their behaviour towards hydrolysis.

Q. 13. What is Milk sugar? Give its characteristics.

Q. 14. Define the term vitamins? State its importance.

Q. 15. What do you understand by denaturation of proteins?

Q. 16. Write a short note on cellulose and give its chemical structure.

Q. 17. Give a short note on Zwitter ion?

Q. 18. How are peptides formed. Show the formation of a peptide bond with the help of a diagram.

Paper Submitted By: Name: Mr. NAVEEN KUMAR GARG Email: [email protected] No: 09412584464

Q. 1. (a) What is a Difference between a double salt and a complex.

(b). Explain the following with suitable examples: Homoleptic and Heteroleptic ligands, Coordination number.

Q. 2. What is meant by unidentate, didentate and ambidentate ligands? Give examples for each.

Q. 3. (a)Specify the oxidation numbers of the metals in the following coordination entities

i. [Co(H2O)(CN)(en)2]2+ ii. [PtCl4]2–

iii. [Cr(NH 3)3Cl3]

iv. [CoBr2(en)2]+

v. K3[Fe(CN)6]

Q. 4. (a) Using IUPAC norms write the formulae for the following:

i. Tetrahydroxozincate(II) ion ii. pentaamminenitrito-N-cobalt (III)

iii. Potassiumtri(oxalato) chromate (III)

iv. Diamminedichloridoplatinum (II)

v. Hexaammine cobalt (III) sulphate

vii. Hexaammineplatinum (IV)

viii. potassiumtetracyanonickelate (II)

ix. Tetrabromido cuprate (II)

x. pentaamminenitrito-O-cobalt (III)

xi. Tetraammineaquachloridocobalt (III) chloride

xii. Potassiumtetrahydroxozincate (II)

xiii. Potassiumtrioxalatoaluminate (III)

xiv. Dichloride bis (ethane-1, 2-diamine) cobalt (III)

xv. Tetracarbonylnickel (0)

(b) Write the formulas for the following coordination compounds:

1. Tetraamminediaquacobalt(III) chloride 2. Potassium tetracyanonickelate(II)

3. Tris(ethane–1, 2–diamine) chromium(III) chloride

4. Amminebromidochloridonitrito-N-platinate(II)

5. Dichloridobis(ethane–1, 2–diamine)platinum(IV) nitrate

6. Iron(III) hexacyanoferrate(II)

Q. 5. (a) Write the IUPAC names of the following coordination compounds:

i. [Pt (NH3)4Cl(NO2)]+ ii. K3[Cr (C2O4)3]

iii. [CoCl2 (en) 2]Cl

iv. [Co (NH3)5(CO3)]Cl

v. Hg [Co (SCN)4]

vi. [Co (NH3)6)]Cl3

vii. [Co (NH3)5Cl]Cl2

viii. K3 [Fe (CN) 6]

ix. K3 [Fe (C2O4)3]

x. K2 [PdCl4]

xi. [Pt (NH3)2Cl (NH2CH3)]Cl

xii. [NiCl4]2-

xiii. [Ni (CO) 4]

xiv. [Co (en) 3]3+

xv. [Co (NH3)6] Cl3

xvi. [Co (NH3)4Cl (NO2)]Cl

xvii. [Ni (NH3)6]Cl2

xviii. [Mn (H2O)6]+

xix. [Co (en) 3]3+

xx. [Ti(H2O)6]3+

(b) Write the IUPAC names of the following coordination compounds:

i. [Co(NH3)6]Cl3 ii. [Co(NH3)5Cl]Cl2

iii. K3[Fe(CN)6]

iv. K3[Fe(C2O4)3]

v. K2[PdCl4]

vi. [Pt(NH3)2Cl(NH2CH3)]Cl

Q. 6. The oxidation number of cobalt in K[Co(CO)4] is

a. +1 b. +3

c. -1

d. -3

Q. 7. What is understood by following, explain giving examples:

i. Linkage Isomerism ii. Ionization isomerism

iii. Coordination isomerism

iv. Solvate(Hydrate)isomerism

Q. 7. What is understood by geometrical and optical isomers?

Q. 8. Draw all the structures of optical isomers of

a. [CrCl(en)2(NH3)]2+ b. [CrCl2(en)2]+

c. [Cr(C2O4)3]3-

d. [PtCl 2(en)2]+

e. [CrCl2 (en) (NH3)2]2+

f. [CoCl (en) 2(NH3)] 2+

g. K[Cr (H2O) 2(C2O4)2]

h. [Co(en)3]Cl3

Q. 9. Draw all the structures of geometrical isomers of:

a. [Pt(NH3)2Cl2] b. [ Co(NH3)4Cl2]+

c. [ CrCl(en)2(NH3)]2+

d. [ CrCl2 (en) 2] +

e. [ Pt (NH3) (H2O) Cl2]

f. [ CoCl (en) 2(NH3)] 2+

g. [ CrCl2 (en)(NH3)2]2+

Q. 10. Write all the geometrical isomers of [Pt(NH3)(Br)(Cl)(py)] and how many of these will exhibit optical isomers?

Q. 11. Why is geometrical isomerism not possible in tetrahedral complexes having two different types of unidentate ligands coordinated with the central metal ion ?

Q. 12. Draw structures of geometrical isomers of [Fe(NH3)2(CN)4]–

Q. 13. Indicate the types of isomerism exhibited by the following complexes and draw the structures for these isomers:

i. K[Cr(H2O)2(C2O4)2 ii. [Co(en)3]Cl3

iii. [Co(NH3)5(NO2)](NO3)2

iv. [Pt(NH3)(H2O)Cl2]

Q. 14. Give evidence that [Co(NH3)5Cl] SO4 and [Co(NH3)5SO4]Cl are ionization isomers.

Q. 15. Using VBT Predict the Magnetic behaviour, Hybridization, Shape of following

i. [ Fe (CN)6]4- ii. [ Fe F6]3-

iii. [ Co F6]3-

iv. [ Co(C2O4)3]3-

v. [ Ni(CN)4]2-

vi. [ NiCl4]2-

vii. [ Ni (CO)4]

viii. [ Fe (H2O) 6]3+

ix. [ Fe (CN)6]3-

x. [ Co (NH3)6] +3

xi. [ Ni (NH3)6] +2

xii. [ Ti(H2O)6]3+

Q. 16. Account for the Following:

a. [ Ni (CN)4]2- ion with square planar structure is diamagnetic and the [ NiCl4]2- ion with tetrahedral structure is paramagnetic .

b. [ NiCl4]2- is paramagnetic while [ Ni (CO)4] is diamagnetic though both are tetrahedral.

c. [ Fe (H2O) 6]3+is strongly paramagnetic while [Fe(CN)6]3- is weakly paramagnetic.

d. [ Co (NH3)6]+3 is an inner orbital complex whereas [Ni(NH3)6]+2 is an outer complex.

e. [ Cr (NH3)6] +3is paramagnetic while [Ni (CN)4]2- is diamagnetic.

f. A solution of [Ni (H2O) 6]2+ is green but a solution of [Ni (CN)4]2- is colourless.

g. [ Fe (CN)6]3- and [Fe(H2O) 6]3+ are of different colours in dilute solutions.

Q. 17. The spin only magnetic moment of [MnBr4]2– is 5.9 BM. Predict the geometry of the complex ion.

Q. 18. 4. Predict the number of unpaired electrons in the square planar [Pt(CN)4]2– ion.

Q. 19. Draw diagram to show splitting of d – orbital in octahedral crystal field. Explain the two patterns of filling d4 in octahedral crystal Field.

Q. 20. Draw diagram to show splitting of d – orbital in a tetrahedral crystal field.

Q. 21. The Hexaaquomanganese (II) ion contains five unpaired electrons, while the Hexacyano manganese (II) ion contains only one unpaired electron. Explain using CFT.

Q. 22. (a)Write down the IUPAC name for each of the following complexes and indicate the oxidation state, electronic configuration and coordination number. Also give stereochemistry and magnetic moment of the complex:

i. K[Cr(H2O)2(C2O4)2]3H2O ii. [CrCl3(py)3] (v) K4[Mn(CN)6]

iii. [Co(NH3)5Cl]Cl2

iv. Cs[FeCl4]

(b) Give the oxidation state, d orbital occupation and coordination number of the central metal ion in the following complexes:

i. K3[Co(C2O4)3] ii. (NH4)2[CoF4]

iii. cis-[Cr(en)2Cl2]Cl

iv. [Mn(H2O)6]SO4

Q. 23. Explain the bonding in coordination compounds in terms of Werner’s postulates

Q. 24. Discuss the role of coordination chemistry in

a. Biological systems b. Analytical Chemistry

c. Medicinal Chemistry

d. metallurgy / Extraction of metals

Q. 25. What is the coordination entity formed when excess of aqueous KCN is added to an aqueous solution of Copper sulphate? Why is it that no precipitates of Copper sulphide is obtained when H2S (g) is passed through this solution.

Q. 26. What is spectrochemical series? Explain the difference between a weak field ligand and a strong field ligand.

Q. 27. What is crystal field splitting energy? How does the magnitude of ∆0 decide the actual configuration of d-orbital in a coordination entity?

Q. 28. What is meant by stability of a coordination compound in solution? State the factors which govern stability of complexes.

Q. 29. What is meant the chelate effect? Give an example.

Q. 30. Discuss the nature of Bonding in metal carbonyls.

Q. 31. Multiple choice questions

(a). How many ions are produced from the complex Co(NH3)6Cl2 in solution?

i. 6 ii. 4

iii. 3

iv. 2

(b). Amongst the following ions which one has the highest magnetic moment value?

i. [Cr(H2O)6]3+ ii. [Fe(H2O)6]2+

iii. [Zn(H2O)6]2+

(c). Amongst the following, the most stable complex is

i. [Fe(H2O)6]3+ ii. [Fe(NH3)6]3+

iii. [Fe(C2O4)3]3–

iv. [FeCl6]3–

Q. 32. What will be the correct order for the wavelengths of absorption in the visible region for the following: [Ni(NO2)6]4–, [Ni(NH3)6]2+, [Ni(H2O)6]2+ ?

Paper Submitted By: Name: Mr. NAVEEN KUMAR GARG Email: [email protected] No: 09412584464

Q. 1. Among the isomeric alkanes of molecular formula C5H12, identify the one that on photochemical chlorination yields (i) A single monochloride (ii) Three isomeric monochlorides. (1 mark)

Q. 2. While separating a mixture of ortho and para nitrophenols by steam distillation, name the isomer which will be steam volatile. Give reason. (1 mark)

Q. 3. Write the IUPAC name of the following. (2 mark)

(i) (ii) (iii) (iv)

Q. 4. With the help of an example discuss the stereo chemistry involved in SN1 and SN 2 mechanism. (2 marks)

Q. 5. Explain Saytzeff rule  by taking suitable examples. (2 marks)

Q. 6. Draw the structures of all isomeric alcohols of molecular formula C5H12O and give their IUPAC names. Classify these isomers of alcohols as primary, secondary and tertiary alcohols. (2 marks)

Q. 7. Nitro group increases the reactivity of chlorobenzene when attched to o and p- position but not at meta position? Explain? Write the mechanism also. (3 marks)

Q. 8. Account for the following. (3 marks)

a. It is necessary to avoid even traces of moisture from a Grignard reagent. b. Chloroform is stored in closed dark coloured bottles completely filled.

c. The carbon– oxygen bond length in phenol is slightly less than that in methanol.

Q. 9. (a) Explain the fact that in aryl alkyl ethers (i) the alkoxy group activates the benzene ring towards electrophilic substitution and (ii) it directs the incoming substituents to ortho and para positions in benzene ring. (2 mark)

(b) Arrange the following compounds in increasing order of their acid strength: (1 mark)

Propan-1-ol, 2,4,6-trinitrophenol, 3-nitrophenol, 3,5-dinitrophenol, phenol, 4-methylphenol

Q. 10. Give equations of the following reactions: (3 mark)

a. Preparation of Sec-Butyl alcohol from ethanal b. Preparation of Phenol from Aniline.

c. Treating phenol with chloroform in presence of aqueous NaOH.

Q. 11. How will you distinguish between the following: (3 marks)

a. Propanoic acid and propanal b. Ethanal and Benzaldehyde

c. Pentan-3-one and Pentan-2-one

Q. 12. How will you bring about the following conversions in not more than two steps. (5 marks)

a. Propanone to Propene b. Benzene to m-Nitroacetophenone

c. Chlorobenzene to benzoic acid

d. Propanoic acid to propene

e. Propyne to propan-2-ol

Q. 13. Explain the following with the help of suitable examples: (5 marks)

(a) Limitation of Williamsons synthesis (b) Stephen reaction. (c) Hell-Volhard-Zelinsky reaction (d) Aldol Condensation (e) Wolff-Kishner reduction.

Q. 14. (a) An organic compound (A) (molecular formula C8H16O2) was hydrolysed with dilute sulphuric acid to give a carboxylic acid (B) and an alcohol (C). Oxidation of (C) with chromic acid produced (B). (C) on dehydration gives but-1-ene. Write equations for the reactions involved

(b) Predict the products formed when cyclohexanecarbaldehyde reacts with following reagents.

i. PhMgBr and then H3O+ ii. Tollens ’ reagent

Q. 15. Primary alkyl halide C4H9Br (a) reacted with alcoholic KOH to give compound (b). Compound (b) is reacted with HBr to give (c) which is an isomer of (a). When (a) is reacted with sodium metal it gives compound (d), C8H18 which is different from the compound formed when n-butyl bromide is reacted with sodium. Give the structural formula of (a) and write the equations for all the reactions.

(ii) Out of C6H5CH2Cl and C6H5CHClC6H5, which is more easily hydrolysed by aqueous KOH.

Q. 16. (a) Write the reactions of glucose with (i) HI (ii)HNO3 (iii)(CH3CO)2O

(b) Discuss the amphoteric nature of amino acid with suitable example.

Q. 17. (a) Discuss the following

i. Denaturation ii. invert sugar

iii. Globular proteins

(b) What is glycogen? How is it different from starch?

(c) Draw the cyclic structures of anomers of glucose.

Q. 18. (a) Compound A of molecular formula C3H7Br, yields a compound B of molecular formula C3H8O when treated with aq. NaOH. On oxidation the compound B yields a ketone C. Compound C on treating with methyl magnesium bromide and then H2O in acidic medium gives tertiary butyl alcohol. Deduce the structures of A, B and C and write the reactions involved. (3 marks)

(b) Draw the structures of major mono halo products in each of the following reactions. (2 marks)

(i) (ii)

(2 marks)

Or

i. Carboxylic acid is a stronger acid than phenol. Why? 2 ii. How will you convert ethanal into the following compounds? 2

(i) Butane-1,3-diol(ii) But-2-enal

iii. Predict the products formed when cyclohexanecarbaldehyde reacts with  Semicarbazide and weak acid       (1 marks)

Q. 19. (a) Discuss the chemistry involved in Brown ring test for nitrate ion. Give equations also. 2 marks

(b) Complete the following equations: 3

i. P4 + SOCl2 ---------------à

ii. 4 AgNO3 + 2H2O + H3PO4 -------------à

iii. NaCl + MnO2 + H2SO4 ------------------à

Or

(a) Account For the Following: 3

i. H3PO2 act as monobasic acid. ii. Interhalogen compounds are more reactive than halogens

iii. Bond dissociation energy of F2 is smaller than Cl2.

(b) Discuss the quantitative method for estimating O3 gas. (2 marks)

Q. 20. (i) What are interstitial compounds? Why are such compounds well known for transition metals? (1 marks)

(ii) How is the variability in oxidation states of transition metals different from that of the non transition metals? Illustrate with examples. (1 marks)

(iii)Describe the preparation of potassium dichromate from iron chromite ore. What is the effect of increasing pH on a solution of potassium dichromate? (2 + 1 marks)

Or

(i) Account for the following. (1 x 3 marks)

a. Sc, the first member of first transition series doesn’t exhibit variable oxidation state. Account for the following

b. Transition metals have a strong tendency to form complexes.

c. Zirconium (atomic no 40) and hafnium (atomic no 72) exhibit similar properties.

(ii) How does acidified solution of potassium dichromate react with (a) FeSO4 and (b) H2S?

Paper Submitted By: Name: Mr. NAVEEN KUMAR GARG Email: [email protected] No: 09412584464

Q. 1. There is a considerable increase in covalent radius from N to P but from As to Bi only a small change is observed.

Q. 2. Ionisation enthalpy of group 15 elements is much higher than that of group 14 elements.

Q. 3. Ionic radius of Sb and Bi are very less when compared to the ionic radius of N,P and As.

Q. 4. Metallic character of group 15 elements decreases on going down the group.

Q. 5. Tendency to show â€“ 3 oxidation states in group 15 decreases on going down the group.

Q. 6. Nitrogen canâ€™t form penta halides.

Q. 7. Nitrogen exhibits bonding while heavier members exhibit bonding.

Q. 8. N2 is a gas while P4 is a solid.

Q. 9. Catenation tendency is weaker in nitrogen.

Q. 10. N2 molecule is chemically inert while white phosphorus is more reactive.

Q. 11. In group 15, +3 oxidation state is more stable than +5 oxidation state on going down the group .

Q. 12. R3 P=O is known but R3 N=O is unknown.

Q. 13. Basicity of hydrides NH3> PH3 > AsH3 > SbH3 > BiH3

Q. 14. Stability of hydrides NH3> PH3 > AsH3 > SbH3 > BiH3

Q. 15. Reducing character of hydrides NH3<>3 <>3 <>3 3

Q. 16. The oxides in higher oxidation states of group 15 elements are more acidic than that of lower oxidation state.

Q. 17. Basicity of group 15 oxides increases on going down the group.

Q. 18. PCl5 is more covalent than PCl3.

Q. 19. PCl5 is more covalent than PF5.

Q. 20. All the five bonds in PCl5 are not equivalent.(Or) PCl5 is more reactive than PCl3.

Q. 21. Both PCl3 and PCl5 fumes in air.

Q. 22. PH3 has lower boiling point than NH3.

Q. 23. NH3 acts as a lewis base.

Q. 24. NO2 molecule dimerise to become N2O4.

Q. 25. Aluminium is rendered passive in concentrated HNO3.

Q. 26. Concentrated HNO3 becomes yellow when exposed to light.(Or) concentrated HNO3 is an oxidizing agent.

Q. 27. White phosphorus is more reactive than red phosphorus. Black phosphorus is least reactive.

Q. 28. Bond angle in PH3+ is higher than that of PH3.

Q. 29. HNH bond angle in NH3 is less than the tetra hedral bond angle of 109.50.

Q. 30. Bond angles of HPH,HAsH and HSbH are closer to 900.

Q. 31. H3PO4 is tri protic, H3PO3is diprotic while H3PO2 is mono protic.

Q. 32. H3PO2 is a good reducing agent.

Q. 33. H3PO2 is a stronger reducing agent than H3PO3.

Q. 34. NO is an odd electron molecule but does not dimerise to give N2O2.

Q. 35. Sulphur has very high boiling and melting point when compared to oxygen.

Q. 36. In group 16 tendencies to show -2 oxidation state decreases on going down the group.

Q. 37. In group 16 +4 oxidation state become more stable than +6 oxidation state on going down the group.

Q. 38. Oxygen can show a maximum covalency of 4 and it can not form hexa valent compound.

Q. 39. Acidity of group 16 hydrides H2O2 S <>2Se <>2Te.

Q. 40. Reducing character of group 16 hydrides H2O2 S <>2Se <>2Te

Q. 41. Boiling point of H2O is higher than that of H2 S.

Q. 42. Sulphur exhibit +6 oxidation state when it combines with fluorine.

Q. 43. SF6 is exceptionally stable or it can not be hydrolysed easily.

Q. 44. SF6is known while SCl6 is unknown.

Q. 45. SF6 is known while SH6 is unknown.

Q. 46. H2O is a liquid while H2S is a gas.

Q. 47. MnO is basic while Mn2O7 is acidic.

Q. 48. O3 is thermo dynamically unstable than O2 (or) . O3 in higher concentration is explosive.

Q. 49. NO gas depletes ozone layer.

Q. 50. Sulphur in vapour state is paramagnetic.

Q. 51. HCl and HNO3 are prepared by reacting NaCl and NaNO3 respectively with H2SO4 while HBr and HI canâ€™t be prepared by this method.

Q. 52. Cane sugar chars in concentrated sulphuric acid.

Q. 53. Concentrated sulphuric acid is a good oxidizing agent.

Q. 54. Two S-O bonds in SO2 are equivalent.

Q. 55. Ka2 of H2SO4 is <<>1.

Q. 56. Halogens have maximum negative electron gain enthalpy in each period.

Q. 57. Fluorine has lesser negative value of electron gain enthalpy than chlorine.

Q. 58. All halogens are colored.

Q. 59. F2 has smaller enthalpy of dissociation than Cl2.

Q. 60. Fluorine has lesser negative value of electron gain enthalpy than chlorine but fluorine is a stronger oxidizing agent than chlorine.

Q. 61. Fluorine shows only â€“ 1 oxidation state. Other halogens can exhibit positive oxidation state.

Q. 62. Halogens show positive oxidation state when they combine with oxygen and fluorine atoms.

Q. 63. Halogens are good oxidizing agent and oxidizing power (reactivity) decreases with the increase in atomic number.

Q. 64. Most of reactions of fluorine are exothermic.

Q. 65. HF is a liquid while other hydrogen halides are gases.

Q. 66. HF has highest boiling point while HCl has lowest boiling point among hydrogen halides.

Q. 67. Acidity of hydrogen halides HF <>

Q. 68. Thermal stability of hydrogen halides HF > HCl > HBr > HI

Q. 69. Thermal stability of group 16 hydrides H6O > H6S > H6Se > H6Te

Q. 70. OF2 is fluoride of oxygen and not oxide of fluorine.

Q. 71. Oxygen and chlorine has similar electro negativity. Oxygen form hydrogen bonding but not chlorine.

Q. 72. Ionic character of halides MF>MCl>MBr>MI

Q. 73. Electron gain enthalpy of O O- is â€“ 141 KJ/mole and O O --> O2- is + 702 KJ/mole. Large number of oxides having O2- is known and not O-.

Q. 74. In metal halides, halides in higher oxidation state of the metal is more covalent than the one in lower oxidation state.(PbCl4 is more covalent than PbCl2)

Q. 75. Inter halogen compound is more reactive than the halogens from which it is formed.

Q. 76. Chlorine is a powerful bleaching agent.

Q. 77. HCl reacts with Fe to give FeCl2 and not FeCl3

Q. 78. Fluorine forms only one oxo acid HOF

Q. 79. Acidity of oxo acids HClO4> HClO3> HClO2> HC1O.

Q. 80. Acidity of oxo acids HOCl>HOBr>HOI

Q. 81. Oxidizing power of HClO4> HClO3> HClO2> HClO

Q. 82. Fluorine form fluoride of oxygen while other halogens form oxides of halogen.

Q. 83. Group 18 elements are chemically unreactive.

Q. 84. Group 18 elements have very high ionization enthalpy and it decreases on going down the group.

Q. 85. Group 18 elements have positive value of electron gain enthalpy.

Q. 86. Group 18 elements have lower value of boiling and melting point and it increases on going down the group.

Q. 87. Group 18 elements have larger atomic radius.

Q. 88. Bartlett synthesized XePtF6 from his knowledge of earlier known compound O2PtF6.

Q. 89. Xenon forms noble gas compounds.

Q. 90. Xenon forms compounds only with oxygen and fluorine.

Q. 91. Helium is used in diving apparatus.

Q. 92. Oxygen has lesser negative value of electron gain enthalpy than Sulphur.

Q. 93. What type of defect reduces the density of unit cell?

Q. 94. What is the effect of temperature on osmatic pressure of the solution?

Q. 95. Determine the order of the reaction if its half life is inversely proportional to theinitial concentration of the reactant.

Q. 96. Mention two uses of formaldehyde.

Q. 97. Write the IUPAC name of the compound CH3 OCOCH2 CH2 COO CH3

Q. 98. What do you understand by radial probability function 4pr2R2? Draw theradial probability function curve for 2s and 2p orbital.

Q. 99. Comment on the validity of the following statements:

i. A reaction with DG0< 0 has always equilibrium constant greater than 1. ii. An exothermic reaction is always spontaneous.

Q. 100. The outer electronic configuration of two elements of lanthanoids are given below

i. 4f1 5d1 6s2 ii. 4f7 5d0 6s2. Determine the atomic number of these elements.

Determine the stable oxidation state of these elements.

Q. 101. Give two differences between conformational isomerism and configurational isomerism.

OR A compound having two chiral centers does not always have four stereo isomers. Justify this statement with a suitable example.

Q. 102. Mention the differences between thermo setting and thermo plastic polymers.

Q. 103. Distinguish chemically between

i. Phenol and methanol ii. Propan-1-ol andpropan-2-ol.

Q. 104. Complete the following chemical equations:

i. U+ClF3® ii. Ca3(PO4)2 +SiO2 +C®

Q. 105. Write the ground state electronic configuration of N2.

i. Determine the bond order of N2. ii. N2 and CO has same bond order, but CO is more reactive than N2. Why?

Q. 106. An element A crystallizes in FCC structure. 200g of this element has 24x1023 atoms. Density of unit cell is 7.2g/cm3. Calculate the radius of A. Avagadro number= 6x 1023 mole- 1

Q. 107. The activation energy of a first order reaction at 270C is 54 KJ/mole. Activation energy of the same reaction at the same temperature in the presence of a catalyst is 44 KJ/mole. How many times the reaction rate changes in the presence of catalyst at this temperature? (R=8.314J/K/mole)

Q. 108. 0.85% solution of NaNO3 is 90% dissociated at 300K.Determine the osmatic pressure of the solution.( R= 0.0821 L atm/K/mole. Molar mass of NaNO3=85u) OR Determine the boiling point of 1M solution of KCl. Assume that KCl is 90% dissociated. Density of KCl solution is 1.05g/cm3. Molar mass of KCl=74.5u. Kb= 0.52 K Kg mole- 1

Q. 109.

a. Predict the sign of entropy change for the reaction CaCO3(s)®CaO(s) +CO2(g)

b. You are provided with the following DrG0 valuesS2 +2 O2 ® 2SO2 DrG0 = - 544 KJ/mole2Zn +O2® 2ZnO DrG0 = - 480 KJ/mole2Zn+ S2 ® 2ZnS DrG0 = - 293 KJ/mole

Show that roasting of ZnS to ZnO is a spontaneous process.

Q. 110. Explain the terms activity and selectivity of catalyst with examples.

Q. 111

a. How is potassium dichromate prepared from chromite ore? b. Write the ionic equation of the reaction involved when KMnO4 is treated

withferrous sulphate solution in acid medium.

Q. 112. Using Valence bond theory compare the structure and magnetic behavior of

i. Ni (CO)4 ii. [Ni (CN)4]2- Atomic number of Ni=28

Q. 113.

i. A radio active element of group 18 undergoes a decay. Determine the position ofthe new nuclide.

ii. Calculate the mass of 140 La in a sample whose activity is 3.7x1010Bq. Given half life of 140 La = 40 hours.

Q. 114. Carry out the following conversions:

i. Benzoyl chloride to benzaldehyde. ii. Hex-1-ene to pentanal

iii. Hexane nitrile to 1-amino pentane.

Q. 115. Account for the following:

i. In ammonalysis of halo alkanes primary amine is the only product when NH3is taken in large excess.

ii. Tert amine has lower boiling point than primary amine of comparable molarmass,

iii. Amide formed in acylation reaction of amine, does not react further withacid halide.

Q. 116. Describe the following with suitable examples:

i. Double base propellant ii. Mordant dyes

Q. 117.

i. Explain corrosion of iron as an electro chemical process. ii. The electrolysis of a metal salt solution was carried out by passing 4

amperesfor 45 minutes. It resulted in the deposition of 2.977g of the metal. If atomicmass of the metal is 106.4g/mole, determine the charge carried by the metal ion.

OR

i. How does conductivity and molar conductivity of an electrolyte solution vary with the dilution of the solution.

ii. The E0 potentials of two reduction electrodes are Cu+/ Cu = + 0.52V and Cu2+ / Cu+ = + 0.16 V. Calculate the work obtainable from the cell.

Q. 118. (a) Account for the following:

i. PCl5 solid is ionic in nature. ii. SF6 is resistant to hydrolysis.

iii. Inter halogen compounds are more reactive than halogens from which it ismade.

(b) Draw the structures of the molecules

i. P4O10 ii. per oxo mono sulphuric acid.

OR

(a) Account for the following:

i. PCl5 fumes in air. ii. Ga is smaller in size than Al.

iii. PbO2 is a good oxidizing agent.

(b) Draw the structures of the molecules

i. SF4 ii. IF4¯

Q. 119.

a. what are lipids? Based on their chemical composition, present a classification of lipids.

b. In reference to DNA molecule what do you under stand by the terms replicationand transcription.

OR

a. Define the terms (i) codon (ii) native state of protein (iii) Denaturation of protein.

b. Write the name of the nucleoside which is present only in (i) DNA (ii) RNA.

Q. 120. Define activation energy of a reaction.

Q. 121. A cubic solid is made of two elements X and Y. Atoms Y are at the corners of the cube and X at the body centre. What is the formula of the compound?

Q. 122. State two main functions of carbohydrates in sugarcane.

Q. 123. Why is HF not stored in plain glass bottles?

Q. 124. State one use of acetonitrile.

Q. 125. What values of quantum number, m are permitted for an electron having angular quantum number, 1 = 2 ?

Q. 126. Which types of crystals exhibit piezoelectricity?

Q. 127. What is Tyndall effect?

Q. 128. How many effective sodium ions are located at the centres of faces of a unit cell in a sodium chloride crystal?

Q. 129. Name the first element of 3 d transition metal series.

Q. 130. What are inner transition metals?

Q. 131. State second law of thermodynamics.

Q. 132. Name a direct dye.

Q. 133. Sketch the zwitter ion form of amino acetic acid.

Q. 134. Mention an industrial product manufactured from methanal.

Q. 135. How does a fuel cell operate?

Q. 136. Mention two important uses of methanol.

Q. 137. What is the importance of amino acids to us?

Q. 138. Give an example of associated colloids.

Q. 139. Define order of a reaction.

Q. 140. Zn, Cd and Hg are not considered as transition metals.

Q. 141. Cu,Ag &Au are considered as transition metal though it has 3d,10 configuration.

Q. 142. Zn, Cd and Hg are volatile and Hg is a liquid metal.

Q. 143. Transition metals have high enthalpy of atomisation.

Q. 144. 4d and 5d elements have higher enthalpy of atomisation than 3d elements.

Q. 145. Density of 3d elements increases from Sc to Ni.

Q. 146. Atomic and ionic radii generally decrease along the period.

Q. 147. Zr and Hf have similar size.

Q. 148. Transition metals do not show regular variation of ionisation enthalpies.

Q. 149. 5d elements have higher ionisation enthalpy than 3d and 4d elements.

Q. 150. Generally first ionisation enthalpy increases along the period.

Q. 151. Cr &Cu has higher second ionisation enthalpy (Cr = 24,Cu = 29)

Q. 152. Mn &Zn has higher 3rd ionisation enthalpy (Mn = 25 Zn = 30)

Q. 153. Cu, Ni and Zn normally do not exhibit oxidation state higher than +2.

Q. 154. Transition metals show variable oxidation state.

Q. 155. Transition metals do not show regular variation of E0 values.

Q. 156. E0 Mn3+/Mn2+ have higher +ve value than E0 Cr3+/Cr2+.(Cr=24 Mn=25).

Q. 157. E0 Zn2+/Zn has higher â€“ value. ( Zn = 30).

Q. 158. Transition metals form complexes easily.

Q. 159. Transition metals act as catalysts.

Q. 160. Transition metal compounds are paramagnetic.

Q. 161. Sc3+,Cu+ and Zn2+ are diamagnetic .Sc=21 Cu=29 Zn=30).

Q. 162. Oxidation state of transition metals increases by one unit.

Q. 163. E0 Ni2+/Ni has higher â€“ ve value.

Q. 164. Transition metal compounds are coloured.

Q. 165. Sc3+, Cu+ and Zn2+ are colourless .[Sc = 21 Cu = 29 Zn = 30].

Q. 166. Transition metals form interstitial compounds easily.

Q. 167. Transition metals form alloys easily.

Q. 168. Higher oxides of transition metals are acidic.

Q. 169. Enthalpy of atomization of transition metals reaches a maximum in the middle of each series

Q. 170. Lanthanoid ions are coloured and paramagnetic.

Q. 171. La3+ and Lu3+ are colourless and diamagnetic. (La=57 Lu=71).

Q. 172. Ln are paramagnetic.

Q. 173. La3+ is diamagnetic.

Q. 174. Ce is stable in +4 oxidation state. E0 Ce4+/Ce3+ has higher positive value.

Q. 175. Eu2+ is a good reducing agent.(Eu=63).

Q. 176. Ce4+ is a good oxidising agent .

Q. 177. Yb2+ is stable. It acts as a reducing agent. (Yb=70).

Q. 178. In lanthanoid hydroxide M(OH)3 basic character decreases with increase in atomic number.

Q. 179. Actinoids show greater range of oxidation state than lanthanoids.

Q. 180. Actinoid contraction is greater than lanthanoid contraction.

Q. 181. Ionisation enthalpies of early actinoids are lesser than that of early lanthanoids.

Q. 182. In the d4 species Mn3+ is an oxidising agent where as Cr2+ is a reducing agent

Q. 183. Co(II) is stable in aqueous solution but get oxidised in the presence of complexing agents.

Q. 184. d1 configuration is unstable in transition metals.

Q. 185. Transition metals exhibit higher oxidation state in oxides and flourides.

Q. 186. Zn has lowest enthalpy of atomisation.

Q. 187. First ionization enthalpy of Cr is lower while Zn is higher.(Cr=24, Zn=30)

Q. 188. Zn has lower second ionization enthalpy.

Q. 189. Number of oxidation states at the extreme ends of 3d transition metals are very few.

Q. 190. Transition metals in higher oxidation states are more stabilized in oxides than in fluorides.

Q. 191. E0 Cu2+/Cu is positive.

Q. 192. E0 Sc3+/Sc2+ has lower value (Sc = 21)

Q. 193. E0 Zn3+/Zn2+ have higher positive value.

Q. 194. E0 Fe3+/Fe2+ has comparatively lower positive value than E0 Mn3+/Mn2+ (Mn=25, Fe=26)

Q. 195. E0 V3+/V2+ have negative value (V=23)

Q. 196. Cu2+ (aq) is more stable than Cu+ (aq) despite the d10 configuration of Cu+

Q. 197. Highest fluoride of Mn is MnF4 while highest oxide is Mn2O7.

Q. 198. Oxidising power of Oxides VO2+ <>2O72- <>4-

Q. 199. KMnO4 titration is not carried out using HCl as acid medium.

Q. 200. Third ionization enthalpy of La, Gd and Lu are abnormally low( La = 57, Gd = 64, Lu = 71)

Q. 201. 5f electrons can take part in bonding to a greater extent than 4f electron though both have similar angular wave function.

Q. 202. Highest oxidation state is exhibited by oxoanion in transition metals.

Q. 203. Generally negative value of E0 M2+/M decreases across the series of 3d elements.

Q. 204. It is unsatisfactory to review the chemistry of actinoids in terms of oxidation state.

Q. 1. Draw the resonating structures of

a. NO b. NO2

c. N2O

d. N2O3

e. N2O4

f. N2O5.

g. O3.

Also draw the structures of each clearly depicting the bond parameters.

Q. 2. Draw the resonance structures of SO2.

Q. 3. Draw the structures of

a. NH3 b. HNO3

c. White phosphorus

d. Red phosphorus

e. PCl3

f. PCl5

g. Phosphoric acid

h. Phosphorus acid

i. Hypo phosphorus acid

j. Pyro phosphoric acid

k. cyclic tri meta phosphoric acid

l. Poly meta phosphoric acid.

m. S8

n. S6

o. Sulphuric acid

p. sulphurus acid

q. Peroxo di sulphuric acid

r. Pyro sulphuric acid (oleum)

s. HOCl

t. HClO2

u. HClO3

v. HClO4

w. BrF3

x. IF5

y. IF7

z. IF4- aa) SF4 bb) SF6 cc) XeOF4 dd) BrO3- ee) XeF2 ff) XeF4 gg) XeF6 hh) XeOF4 ii) XeO3

Q. 4. Give the formula and structure of noble gas species which is iso structural with

a. ICl4- b. IBr2-

c. IF6-

d. BrO3-

Q. 5. Why does nitrogen shows anomalous behavior? Give examples to show the anomalous behavior of nitrogen.

Q. 6. Why does oxygen shows anomalous behavior? Give examples to show the anomalous behavior of oxygen.

Q. 7. Why does fluorine shows anomalous behavior? Give examples to show the anomalous behavior of fluorine.

Q. 8. Describe the method of preparation of

a. NH3 by Haber procssssess b. HNO3 by Ostwald process

c. H2SO4 by contact process.

Give three uses of each.

Q. 9. Explain brown ring test for nitrate with suitable equations.

Q. 10. What is disproportionation reaction? Give equation of the reactions involved in the disproportionation of

a. HNO2 b. Se2Cl2

c. H3PO2

Q. 11. How is ozone estimated quantitatively?

Q. 12. Give two uses each of

a. N2 b. PH3

c. O2

d. Ozone

e. SO2

f. Cl2

g. ClO2

h. BrO3

i. I2O5

j. Cl2 k) HCl

k. ClF3

l. He

m. Ne

n. Ar

o. Kr

p. Xe

Q. 13. How is NH3(aq) used in salt analysis to determine the presence of

a. Fe3+ b. Zn2+

c. Ag+

in salt analysis. Write the equations of the reactions involved.

Q. 14. How is the presence of SO2 detected?

Q. 15. What is aqua regia? How does it dissolve noble metals like Au and Pt? Write the equations of the reactions involved.

Q. 16. Give differences between white phosphorus and red phosphorus.

Q. 17. How is

a. N2 b. O2

c. Cl2 prepared in the laboratory? Write the equations of the reactions involved.

Q. 18. How is

a. N2 b. O2

c. Cl2 manufactured in the industry? Write the equations of the reactions involved in the manufacture of Cl2.

Q. 19. How is ammonia prepared in laboratory? Write the equation of the reaction involved.

Q. 20. How is HNO3 prepared in laboratory? Write the equation of the reaction involved.

Q. 21. How is phosphine prepared from

a. Calcium phosphide b. White phosphorus.

Write the equations of the reactions involved.

Q. 22. How is PCl5 prepared from

a. Cl2 b. SO2Cl2?

Write the equations of the reactions involved

Q. 23. How is PCl3 prepared from

1. Cl2 2. SOCl2?

Write the equations of the reactions involved

Q. 24. How is HCl prepared from NaCl? Write the equations of the reactions involved

Q. 25. Write the chemical formula of

a. Chile saltpetre b. Indian saltpetre

c. Fluorapatite

d. Gypsum salt

e. Epsom salt

f. Baryte

g. Galena

h. Zinc blende

i. Copper pyrite

j. Florospar

k. cryolite

l. Fluoroapatite

m. carnalite.

n. Tear gas

o. mustard gas

p. Phosgene

Q. 26. With what neutral molecule ClO- is iso electronic? Is that molecule a lewis base?

Q. 27. Compare the chemistry of Sulphur.What is the transition temperature of Sulphur.

Q. 28. How is phosphine purified? Write the equations of the reactions involved.

Q. 1. There is a considerable increase in covalent radius from N to P but from As to Bi only a small change is observed.

Q. 2. Ionisation enthalpy of group 15 elements is much higher than that of group 14 elements.

Q. 3. Ionic radius of Sb and Bi are very less when compared to the ionic radius of N,P and As.

Q. .4. Metallic character of group 15 elements decreases on going down the group.

Q. .5. Tendency to show – 3 oxidation states in group 15 decreases on going down the group.

Q. .6. Nitrogen can’t form penta halides.

Q. 7. Nitrogen exhibits bonding while heavier members exhibit bonding.

Q. 8. N2 is a gas while P4 is a solid.

Q. 9. Catenation tendency is weaker in nitrogen.

Q. 10. N2 molecule is chemically inert while white phosphorus is more reactive.

Q. 11. In group 15, +3 oxidation state is more stable than +5 oxidation state on going down the group .

Q. 12. R3 P=O is known but R3 N=O is unknown.

Q. 13. Basicity of hydrides NH3> PH3 > AsH3 > SbH3 > BiH3

Q. 14. Stability of hydrides NH3> PH3 > AsH3 > SbH3 > BiH3

Q. 15. Reducing character of hydrides NH3< PH3 < AsH3 < SbH3 <BiH3

Q. 16. The oxides in higher oxidation states of group 15 elements are more acidic than that of lower oxidation state.

Q. 17. Basicity of group 15 oxides increases on going down the group.

Q. 18. PCl5 is more covalent than PCl3.

Q. 19. PCl5 is more covalent than PF5.

Q. 20. All the five bonds in PCl5 are not equivalent.(Or) PCl5 is more reactive than PCl3.

Q. 21. Both PCl3 and PCl5 fumes in air.

Q. 22. PH3 has lower boiling point than NH3.

Q. 23. NH3 acts as a lewis base.

Q. 24. NO2 molecule dimerise to become N2O4.

Q. 25. Aluminium is rendered passive in concentrated HNO3.

Q. 26. Concentrated HNO3 becomes yellow when exposed to light.(Or) concentrated HNO3 is an oxidizing agent.

Q. 27. White phosphorus is more reactive than red phosphorus. Black phosphorus is least reactive.

Q. 28. Bond angle in PH3+ is higher than that of PH3.

Q. 29. HNH bond angle in NH3 is less than the tetra hedral bond angle of 109.50.

Q. 30. Bond angles of HPH,HAsH and HSbH are closer to 900.

Q. 31. H3PO4 is tri protic, H3PO3 is diprotic while H3PO2 is mono protic.

Q. 32. H3PO2 is a good reducing agent.

Q. 33. H3PO2 is a stronger reducing agent than H3PO3.

Q. 34. NO is an odd electron molecule but does not dimerise to give N2O2.

Q. 35. Sulphur has very high boiling and melting point when compared to oxygen.

Q. 36. In group 16 tendencies to show -2 oxidation state decreases on going down the group.

Q. 37. In group 16 +4 oxidation state become more stable than +6 oxidation state on going down the group.

Q. 38. Oxygen can show a maximum covalency of 4 and it can not form hexa valent compound.

Q. 39. Acidity of group 16 hydrides H2O <H2 S < H2Se < H2Te.

Q. 40. Reducing character of group 16 hydrides H2O <H2 S < H2Se < H2Te

Q. 41. Boiling point of H2O is higher than that of H2 S.

Q. 42. Sulphur exhibit +6 oxidation state when it combines with fluorine.

Q. 43. SF6 is exceptionally stable or it can not be hydrolysed easily.

Q. 44. SF6 is known while SCl6 is unknown.

\Q. 45. SF6 is known while SH6 is unknown.

Q. 46. H2O is a liquid while H2S is a gas.

Q. 47. MnO is basic while Mn2O7 is acidic.

Q. 48. O3 is thermo dynamically unstable than O2 (or) . O3 in higher concentration is explosive.

Q. 49. NO gas depletes ozone layer.

Q. 50. Sulphur in vapour state is paramagnetic.

Q. 51. HCl and HNO3 are prepared by reacting NaCl and NaNO3 respectively with H2SO4 while HBr and HI can’t be prepared by this method.

Q. 52. Cane sugar chars in concentrated sulphuric acid.

Q. 53. Concentrated sulphuric acid is a good oxidizing agent.

Q. 54. Two S-O bonds in SO2 are equivalent.

Q. 55. Ka2 of H2SO4 is << Ka1.

Q. 56. Halogens have maximum negative electron gain enthalpy in each period.

Q. 57. Fluorine has lesser negative value of electron gain enthalpy than chlorine.

Q. 58. All halogens are colored.

Q. 59. F2 has smaller enthalpy of dissociation than Cl2.

Q. 60. Fluorine has lesser negative value of electron gain enthalpy than chlorine but fluorine is a stronger oxidizing agent than chlorine.

Q. 61. Fluorine shows only – 1 oxidation state. Other halogens can exhibit positive oxidation state.

Q. 62. Halogens show positive oxidation state when they combine with oxygen and fluorine atoms.

Q. 63. Halogens are good oxidizing agent and oxidizing power (reactivity) decreases with the increase in atomic number.

Q. 64. Most of reactions of fluorine are exothermic.

Q. 65. HF is a liquid while other hydrogen halides are gases.

Q. 66. HF has highest boiling point while HCl has lowest boiling point among hydrogen halides.

Q. 67. Acidity of hydrogen halides HF < HCl < HBr < HI

Q. 68. Thermal stability of hydrogen halides HF > HCl > HBr > HI

Q. 69. Thermal stability of group 16 hydrides H6O > H6S > H6Se > H6Te

Q. 70. OF2 is fluoride of oxygen and not oxide of fluorine.

Q. 71. Oxygen and chlorine has similar electro negativity. Oxygen form hydrogen bonding but not chlorine.

Q. 72. .Ionic character of halides MF>MCl>MBr>MI

Q. 73. Electron gain enthalpy of O O- is – 141 KJ/mole and O O --> O2- is + 702 KJ/mole. Large number of oxides having O2- is known and not O-.

Q. 74. In metal halides, halides in higher oxidation state of the metal is more covalent than the one in lower oxidation state.(PbCl4 is more covalent than PbCl2)

Q. 75. Inter halogen compound is more reactive than the halogens from which it is formed.

Q. 76. Chlorine is a powerful bleaching agent.

Q. 77. HCl reacts with Fe to give FeCl2 and not FeCl3

Q. 78. Fluorine forms only one oxo acid HOF

Q. 79. Acidity of oxo acids HClO4> HClO3> HClO2> HC1O.

Q. 80. Acidity of oxo acids HOCl>HOBr>HOI

Q. 81. Oxidizing power of HClO4> HClO3> HClO2> HClO

Q. 82. Fluorine form fluoride of oxygen while other halogens form oxides of halogen.

Q. 83. Group 18 elements are chemically unreactive.

Q. 84. Group 18 elements have very high ionization enthalpy and it decreases on going down the group.

Q. 85. Group 18 elements have positive value of electron gain enthalpy.

Q. 86. Group 18 elements have lower value of boiling and melting point and it increases on going down the group.

Q. 87. Group 18 elements have larger atomic radius.

Q. 88. Bartlett synthesized XePtF6 from his knowledge of earlier known compound O2PtF6.

Q. 89. Xenon forms noble gas compounds.

Q. 90. Xenon forms compounds only with oxygen and fluorine.

Q. 91. Helium is used in diving apparatus.

Q. 92. Oxygen has lesser negative value of electron gain enthalpy than Sulphur.

Q. 1. Why do amines act as nucleophile?

Q. 2. Name the follwong according to IUPAC system:CH3COCH2COCH3.

Q. 3. What are azeotropes?

Q. 4. State condition resulting in reverse osmosis.

Q. 5. Write the IUPAC names of BHC OR DDT.

Q. 6. Convert: i) Propene to propan-2-ol (ii) Anisole to phenol

Q. 7. How would you account for following:

i. Sulpher hexafluoride is less reactive than sulphur tetrafluoride.

ii. Of the noble gases only Xe forms chemical compounds.

Q. 8. Arrange: CH3OH, (CH3)2CHOH, (CH)3-C-OH in Increasing order of

i. Acidic strength ii. Reactivity towards HCl.

Q. 9. Arrange RF, RCl, RBr, RI in increasing order of

i. BP ii. Polarity

Q. 10. Arrange the following:

i. MF, MCl, MBr, MI (Increasing order of ionic character) ii. NH3, PH3, AsH3, SbH3 (Increasing order of basic strength)

Q. 11. Write short notes on i) diazotization ii) Coupling.

Q. 12. Account for the following:

1. Methyl amine has lower BP than methanol. 2. Methyl amine is a stronger base than ammonia.

Q. 13. Write chemical equation to illustrate:

a. Rosenmund Reduction b. Cannizaro Reaction

c. Fischer Esterification.

Q. 14. Account for any two of the following:

a. Amines are basic substances while amides are neutral. b. Alkyl halides have higher BP than hydrocarbons with almost same Molecular

mass.

c. Aromatic amines are weaker bases than aliphatic amines.

Q. 15. Why aniline is more reactive than benzene? Why for halogenation of aniline, it is first treated with acetic anhydride?

Or

Arrange primary, secondary, tertiary amines in increasing order of basic strength.

Q. 16. Arrange:

a. o-methyl aniline, m-methyl aniline, p-methyl aniline, aniline and m-nitro aniline in increasing order of basic strength.

b. o-hydroxy benzoic acid, p-hydroxy benzoic acid, m-hydroxy benzoic acid and benzoic acid (increasing order of acidic strength)

Q. 17. Explain:

i. Gabriel pthalamide synthesis ii. Aldol Condensation

iii. Arrange: Acetic acid, 1-chloroacetic acid, 1,1-dichloroaceic acid, 1,1,1-trichloroacetic acid in increasing order of acidic strength.

Q. 18. What is Van’t Hoff Factor? Why do some solutes show abnormal molecular masses? What is the value of Van’t Hoff Factor for

a. Association b. Dissociation

c. Non-electrolytes.

Q. 19. Account for following:

a. Oxygen is a gas while sulphur is a solid at room temp. b. Nitrogen does not form NCl5 but P forms PCl5.

c. Halogens have the smallest size in the group.

Q. 20. Distinguish between any three:

a. Propanone and Benzaldehyde b. Benzaldehyde and Acetaldehyde

c. ethanol and phenol

d. Chlorobenzene and chloro cyclohexane

Q. 21. Convert:

a. Ethanamine to methanamine b. Methanamine to ethanamide

c. Phenol to phenolphthalein

Q. 22. Why haloalkanes undergo nucleophilic substitution rxn while haloarenes undergo electrophilic substitution rxn?

Q. 23. A 0.1539 molal aq solution of cane sugar (MM = 342gm) has a FP of 271 K while FP of pure water is 273.15K. What will be FP of a aq solution containing 5 gm of glucose (MM = 180gm) per 100 gm of solution.

Q. 24. Write short notes on

a. Carbylamine or Riemer Teimann Rxn. b. Preparation of carboxylic acid from Grignard Reagent.

Q. 25. Assign Reasons for following:

1. Acidic Strength of acids is HF<HCl<HBr<HI 2. SnCl4 is a liquid while SnCl2 is a solid.

3. SnCl4is less stable than SnCl2.

4. H3PO2 behaves as a monoprotic acid.

5. Noble gases exhibit low chemical reactivity.

Or

a. What are interhalogen compounds? What are their types? Give examples. Why are interhalogen compounds more recative than corresponding halogens? (3)

b. Arrange: i) HClO, HClO2, HClO3, HClO4 (Increasing Acidic Strength)

c. Hydrides of group 17 (Thermal stability and Reducing Character)

Q. 26. Convert:

a. Chlorobenzene to phenol b. Phenol to 2-acetoxy benzoic acid

c. Benzene to m-nitro benzene

d. Toluene to Benzyl Chloride

e. 2-propanol to 1-bromopropane.

Q. 27.

a. Define Roult’s Law for a solution containing non volatile solute.(1) b. Explain why a mixture of ethyl alcohol and cyclohexane shows +ve deviation

from Roult’s Law. (2)

c. Arrange:1m NaCl, 1m CaCl2, 1m glucose (increasing order of elevation in BP)

d. Give one use of reverse osmosis.

2 Marks questions

Q. 1. Why water from the soil rises to the top of the trees?

Q. 2. Which substance is added to water in car radiators to act as antifreeze?

Q. 3. What is the expected value of Van’t Hoff Factor for K4[Fe(CN)6]?

Q. 4. How is depression in FP related to molecular mass?

Q. 5. Give relation between normality and molarity of i) 1 M H2SO4 ii) 1M Ca(OH)2.

Q. 6. BP of a solution is always________________ than that of pure solvent.

Q. 7. 2 gm each of solute A and B (molar mass A<B) are dissolved separately in 50 gm of solvent. Which will show greater elevation in BP?

Q. 8. Why is CaCl2 used to clear snow on roads?

Q. 9. The boiling points of ethyl alcohol and methyl alcohol are 78.3ºC and 64.5ºC. Which of them has a higher vapour pressure?

Q. 10. What is an azeotrope?

2 Marks questions

Q. 11. Define Osmotic Pressure. Arrange the following in increasing order of OP:

1. 34.2 g/L of sucrose (MM = 342) 2. 60 g/L Urea (MM = 60)

3. 90 g/L glucose (MM = 180) iv) 58.5 gm/L of NaCl (MM = 58.5) (Hint g/L means WB/V)

Q. 12. Define i) Normality ii) Mole Fraction.

Q. 13. Define Roult’s Law.

Q. 14. Define Colligative Property. Show that relative lowering in vapour pressure is equal to mole fraction of solute.

Q. 15. Explain why addition of solute lowers freezing point of sol. Draw the graph also.

3 Marks questions

Q. 16. A commercially available sample of sulphuric acid contains 25% H2SO4 by weight (density = 1.10 g/ml) Calculate

i. Molarity ii. Normaility

iii. Molality.

Q. 17. A solution containing 0.5 gm of KCl dissolved in 100 gm of water freezes at -0.240C. Calculate the degree of dissociation of salt. Kf for water = 1.86.

Q. 18. Osmotic Pressure of Blood is 8.21 atm at 370C. How much glucose should be used per litre for an intravenous injection that is isotonic with blood?

5 Marks questions

Q. 19. Define Van’t Hoff Factor. Why do some solutes show abnormal molar masses? Give the value of Van’t Hoff Factor for i) Association ii) Dissociation iii) Non-Electrolytes. How do colligative properties get modified with help of Van’t Hoff Factor. Give equation only.

Q. 20. Define and give characteristics of ideal and non-ideal solutions. With the help of graph and examples explain types of non-ideal solutions.

Q. 1. H2O is a liquid while H2S is a gas.

Q. 2. HF is a weak acid than HI although F is more electronegative than I.

Q. 3. Nitrogen does not form NCl5 but P forms PCl5.

Q. 4. Give example of a compound in which OS of Chlorine is +7.

Q. 5. Draw the shape of

i. XeOF4 ii. XeO3.

Q. 6. Which compound led to the discovery of noble gas compounds?

Q. 7. Why noble gases forms compounds only with halogens?

Q. 8. Arrange: MF, MCl, MBr, MI in increasing order of ionic character.

Q. 9. Arrange: NH3, PH3, AsH3, SbH3 in increasing order of Bond strength.

Q. 10. Arrange: HClO, HClO2, HClO3, HClO4 in increasing order of acidic strength.

Q. 11. H3PO3 is diprotic.

Q. 12. OF2 should be called oxygen difluoride and not fluorine oxygen.

Q. 13. PbX2 is more stable than PbX4.

Q. 14. SnCl2 is a solid while SnCl2 is a liquid at room temp.

Q. 15. Noble gases exhibit low reactivity.

Q. 16. Arrange F2, Cl2, Br2, I2 in increasing order of Bond Energy.

Q. 17. Nitrogen exists as N2 gas while Phosphorous exists as P4.

Q. 18. Oxygen is a gas while Sulphur is a solid at room temp.

Or

Oxygen exists as O2 while S exists as S8.

Q. 19. Arrange the halides of group no.17 in increasing order of acidic strength, boiling point, thermal stability, reducing strength, bond dissociation energy.

Q. 20. X2 is a greenish yellow gas with an offensive smell used in water purifications. It partially dissolves in water to give a solution which turns blue litmus red. When X2 is passed through NaBr solution Br2 is obtained.Identify X2, name the group to which it belongs, write general electronic configuration, what are the products obtained when X2 reacts with H2O. Write eq. also.

Q. 21. Draw the structure of i) orthophosphoric acid ii) Sulphuric Acid

Q. 22. What is Caro’s acid?

Q. 23. Explain why halogens except F show oxidation state of +5, +3, +1 etc (3marks)

Q. 24. What are interhalogen compounds? What are their various types? Give examples of interhalogen compounds. Why are interhalogen compounds more reactive than corresponding halogens? (5)

Q. 25. Why halogens are coloured?

Q. 26. What led to the discovery of noble gases? (3)

Q. 27. How is SO2 an air pollutant?

Q. 28. Halogens have the smallest size in the group. Why?

Q. 29. Noble gases have higher atomic size than expected.

Q. 30. Fluorine has higher EA than chlorine> Explain

Q. 1. What is lanthonoid contraction? What is the reason for lanthonoid contraction? Mention the consequences of lanthonoid contraction.

Q. 2. What is mischmetal? Mention two uses of mischmetal.

Q. 3. How is a) KMnO4 prepared from pyrolusite ore? b) K2Cr2O7 prepared from chromite ore?

Q. 4. Write the ionic equations of the reactions involved when acidified KMnO4is treated with

a. FAS solution b. Oxalic acid solution

c. hydrogen sulphide

d. KI solution

e. Sn2+ solution

f. SO32-

g. NO2-

Q. 5. Write the ionic equation of the reaction involved when alkaline KMnO4 is reacted with

a. KI solution. b. S2O3

2-

Q. 6. Write the ionic equation of the reaction involved when acidified K2Cr2O7 reacts with

a. Sn2+solution b. SO2

c. hydrogen sulphide

d. Fe2+

Q. 7. What is the effect of PH on chromate and dichromate solutions?

Q. 8. Compare the chemistry of lanthonoids and actinoids with reference to

a. electronic configuration b. oxidation state

c. ionization enthalpy

d. chemical reactivity

e. magnetic behaviour f) atomic size

Q. 9. What is actinoid contraction? Why is it more pronounced than lanthonoid contraction?

Q. 10. Compare the general characteristic of first row transition metals with those of second and third series metals in the respective vertical columns with reference to

a. electronic configuration b. oxidation state

c. ionization enthalpies

d. atomic size.

Q. 11. Compare the stability of +2 oxidation state for the elements of the 3d series.

Q. 12. Cr2+(aq) is a better reducing agent than Fe2+ despite the half filled stability of Fe3+ Why?

Q. 13. Which transition metal exhibit only one oxidation state. Why?

Q. 14. For M2+/M and M3+/M2+ systems E0 values of some metals are given below.

Cr2+/Cr - 0.9V Cr3+/Cr2+ - 0.4V

Mn2+/Mn - 1.2V Mn3+/Mn2+ + 1.5V

Fe2+/Fe - 0.4V Fe3+/Fe2+ + 0.8V

Use the data to comment upon :

a. Stability of Fe3+ in acid solution when compared to Cr3+ and Mn3+ b. ease with which iron can be oxidized as compared to chromium and

manganese.

Q. 15. Outer electronic configuration of elements a) X = 4f1 5d1 6s2 b) Y=4f7 5d0 6s2 Determine the

1. atomic numbers of X and Y 2. Stable oxidation states of X and Y

3. Which is an oxidizing agent and which is a reducing agent? Why?

Q. 16. Outer electronic configuration of element X is 5f7 6d0 7s2 Determine the atomic number of the element. What is the stable oxidation state of this element?

Q. 17. Draw the structures of a) manganate ion b) permanganate ion c) chromate ion and d) dichromate ion.

Q. 18. What is disproportination reaction? Give the reaction involving disproportination of

a. Cu+ ion b. MnO4

2- ion.

Q. 19. Discuss the general properties of transition metals with reference to a) electronic configuration b) atomic radius b) ionization enthalpy c) oxidation state d) magnetic behavior e) colored compounds f) complex formation g) catalytic behavior h) interstitial compounds i) alloy formation j) electrode potential(M2+/M and M3+/M2+)

Q. 20. What are transition metals? Which three elements are not considered as transition metals though they are kept in d block? In which way the electronic configuration of transition element differ from non transition elements?

Q. 1. Zn, Cd and Hg are not considered as transition metals.

Q. 2. Cu,Ag &Au are considered as transition metal though it has 3d,10 configuration.

Q. 3. Zn, Cd and Hg are volatile and Hg is a liquid metal.

Q. 4. Transition metals have high enthalpy of atomisation.

Q. 5. 4d and 5d elements have higher enthalpy of atomisation than 3d elements.

Q. 6. Density of 3d elements increases from Sc to Ni.

Q. 7. Atomic and ionic radii generally decrease along the period.

Q. 8. Zr and Hf have similar size.

Q. 9. Transition metals do not show regular variation of ionisation enthalpies.

Q. 10. 5d elements have higher ionisation enthalpy than 3d and 4d elements.

Q. 11. Generally first ionisation enthalpy increases along the period.

Q. 12. Cr &Cu has higher second ionisation enthalpy (Cr = 24,Cu = 29)

Q. 13. Mn &Zn has higher 3rd ionisation enthalpy (Mn = 25 Zn = 30)

Q. 14. Cu, Ni and Zn normally do not exhibit oxidation state higher than +2.

Q. 15. Transition metals show variable oxidation state.

Q. 16. Transition metals do not show regular variation of E0 values.

Q. 17. E0 Mn3+/Mn2+ have higher +ve value than E0 Cr3+/Cr2+.(Cr=24 Mn=25).

Q. 18. E0 Zn2+/Zn has higher – value. ( Zn = 30).

Q. 19. Transition metals form complexes easily.

Q. 20. Transition metals act as catalysts.

Q. 21. Transition metal compounds are paramagnetic.

Q. 22. Sc3+,Cu+ and Zn2+ are diamagnetic .Sc=21 Cu=29 Zn=30).

Q. 23. Oxidation state of transition metals increases by one unit.

Q. 24. E0 Ni2+/Ni has higher – ve value.

Q. 25. Transition metal compounds are coloured.

Q. 26. Sc3+, Cu+ and Zn2+ are colourless .[Sc = 21 Cu = 29 Zn = 30].

Q. 27. Transition metals form interstitial compounds easily.

Q. 28. Transition metals form alloys easily.

Q. 29. Higher oxides of transition metals are acidic.

Q. 30. Enthalpy of atomization of transition metals reaches a maximum in the middle of each series

Q. 31. Lanthanoid ions are coloured and paramagnetic.

Q. 32. La3+ and Lu3+ are colourless and diamagnetic. (La=57 Lu=71).

Q. 33. Ln are paramagnetic.

Q. 34. La3+ is diamagnetic.

Q. 35. Ce is stable in +4 oxidation state. E0 Ce4+/Ce3+ has higher positive value.

Q. 36. Eu2+ is a good reducing agent.(Eu=63).

Q. 37. Ce4+ is a good oxidising agent .

Q. 38. Yb2+ is stable. It acts as a reducing agent. (Yb=70).

Q. 39. In lanthanoid hydroxide M(OH)3 basic character decreases with increase in atomic number.

Q. 40. Actinoids show greater range of oxidation state than lanthanoids.

Q. 41. Actinoid contraction is greater than lanthanoid contraction.

Q. 42. Ionisation enthalpies of early actinoids are lesser than that of early lanthanoids.

Q. 43. In the d4 species Mn3+ is an oxidising agent where as Cr2+ is a reducing agent

Q. 44. Co(II) is stable in aqueous solution but get oxidised in the presence of complexing agents.

Q. 45. d1 configuration is unstable in transition metals.

Q. 46. Transition metals exhibit higher oxidation state in oxides and flourides.

Q. 47. Zn has lowest enthalpy of atomisation.

Q. 48. First ionization enthalpy of Cr is lower while Zn is higher.(Cr=24, Zn=30)

Q. 49. Zn has lower second ionization enthalpy.

Q. 50. Number of oxidation states at the extreme ends of 3d transition metals are very few.

Q. 51. Transition metals in higher oxidation states are more stabilized in oxides than in fluorides.

Q. 52. E0 Cu2+/Cu is positive.

Q. 53. E0 Sc3+/Sc2+ has lower value (Sc = 21)

Q. 54. E0 Zn3+/Zn2+ have higher positive value.

Q. 55. E0 Fe3+/Fe2+ has comparatively lower positive value than E0 Mn3+/Mn2+ (Mn=25, Fe=26)

Q. 56. E0 V3+/V2+ have negative value (V=23)

Q. 57. Cu2+ (aq) is more stable than Cu+ (aq) despite the d10 configuration of Cu+

Q. 58. Highest fluoride of Mn is MnF4 while highest oxide is Mn2O7.

Q. 59. Oxidising power of Oxides VO2+ < Cr2O72- < MnO4-

Q. 60. KMnO4 titration is not carried out using HCl as acid medium.

Q. 61. Third ionization enthalpy of La, Gd and Lu are abnormally low( La = 57, Gd = 64, Lu = 71)

Q. 62. 5f electrons can take part in bonding to a greater extent than 4f electron though both have similar angular wave function.

Q. 63. Highest oxidation state is exhibited by oxoanion in transition metals.

Q. 64. Generally negative value of E0 M2+/M decreases across the series of 3d elements.

Q. 65. It is unsatisfactory to review the chemistry of actinoids in terms of oxidation state.

Q. 1. Give one example each of

a. Markwonikov’s addition. b. Kharasch effect.

c. Sand Meyer reaction

d. Diazotisation reaction

e. Finkelstein reaction

f. Swarts reaction.

g. Wurtz reaction

h. Wurtz Fittig reaction

i. Fittig reaction

j. Friediel craft’s acylation reaction of chloro benzene

k. Friediel craft’s alkylation reaction.of chloro benzene.

l. nitration of chloro benzene.

m. sulphonation of chloro benzene.

n. Dehydro halogenation

o. Zatsev rule.

p. chlorination of chloro benzene.

Q. 2. Write the IUPAC names of the following compounds.

a. n - butyl chloride b. iso butyl chloride

c. secondary butyl chloride

d. tertiary butyl chloride.

e. ethylidene chloride

f. ethylene di chloride

g. vinyl chloride.

h. p - Cl C6H4 CH2 CH (CH3)2

i. m - Cl C6H4 CH2C (CH3)3

j. ( CCl3)3 CCl

k. o- BrC6H4 CH( CH3) CH2 CH3

l. CH3-C (p-Cl C6H4)2/ CH(Br) CH3

Q. 3. Explain the classification of halo alkanes based on

a. number of halogen atoms. b. compounds having sp3 C-X bond

c. compounds having sp2 C-X bond

d. dihalides. Give one example each and their IUPAC names.

Q. 4. Account for the following:

1. Halo alkanes have higher boiling point than the corresponding parent alkane. 2. Boiling point of halo alkanes RI>RBr>RCl> RF

3. Boiling point of 1-Bromo butane >2-Bromo butane> 1-Bromo- 2-methyl propane> 2-Bromo- 2-methyl propane.

4. Melting point of p-Dichlo benzene is higher than its ortho and meta isomer.

5. Halo alkanes are polar in nature but sparingly soluble in water.

6. Iodo alkane can not be prepared by the reaction of alcohol with KI and sulphuric acid. Phosphoric acid is used in place of sulphuric acid.

7. Order of reactivity of alcohol with HX is tert alcohol> sec alcohol > primary alcohol..

8. Halo arenes can not be prepared by treating phenol with HX or NaX in the presence of sulphuric acid.

9. Iodination of benzene is carried out in the presence of HIO3 or HNO3.

10. Propane on chlorination gives 2-chloro propane as a major product and not 1-chloro propane.

11. Kharasch effect is possible only with HBr and not with HCl and HI.

12. Alcohol reacts with thionyl chloride to give pure halo alkane.

13. Finkelstein reaction of halo alkane is carried out in the presence of dry acetone.

14. Order of reactivity of halo alkanes as per substitution bimolecular nucleophilic is primary halide > secondary halide>tertiary halide.

15. Order of reaction as per substitution unimolecular is tertiary halide>secondary halide >primary halide.

16. Benzylic halides and allylic halides are more reactive towards nucleophile than halo alkanes.

17. Chloro ethene is less reactive towards nucleophile than chloro ethane.

18. Halo arenes are less reactive towards nucleophile than halo alkanes.

19. SN1 mechanism is ruled out in the reaction of halo arenes with nucleophile.

20. Electron with drawing groups like NO2 at ortho and para position with respect to halogen facilitates nucleophillic substitution reaction.

21. Electron with drawing groups like NO2 at meta position with respect to halogen has no effect on nucleophillic substitution reaction.

22. Halo arenes are less reactive towards electrophile than benzene.

23. Although chlorine atom has electron with drawing effect electrophillic substitution occur at ortho and para position.

24. Order of reactivity of alkyl halide RI > RBr > RCl > RF

25. Halo alkanes react with KCN to give alkyl cyanide as a major product while it gives alkyl isocyanide as a major product with AgCN.

26. Halo alkanes give nitrito alkane with KNO2 while nitro alkane with AgNO2.

Q. 5. Explain the following with suitable examples:

a. chiral and chirality b. enantiomers

c. racemic mixture

d. retention of configuration

e. inversion of configuration.

Q. 6. Mention the differences between SN1 and SN

2 mechanism of halo alkane.

Q. 7. Explain the mechanisms of the following reactions :

1.

2.

Q. 8. Carry out the following conversions :

1. Propene to a) Propan-1-ol b) Propan-2-ol 2. Ethanol to but-1-yne

3. 1-Bromo propane to 2-Bromo propane and vice versa.

4. Toluene to benzyl alcohol.

5. Benzene to a) 4-bromonitro benzene b) 3-bromonitro benzene.

6. Benzyl alcohol to 2-phenyl ethanoic acid.

7. Ethanol to a) Propane nitrle b) Ethyl isocyanide.

8. Aniline a) Chloro benzene b) Bromo benzene c) Iodo benzene.

9. 2-Chloro butane to 3,4- dimethylhexane.

10. 2-Methyl-1-propene to 2-chloro-2-methylpropane.

11. Ethyl chloride to propanoic acid.

12. But-1-ene to n-butyl iodide.

13. 2-chloropropane to propan-1-ol

14. Isopropyl alcohol to iodoform.

15. Chloro benzene to a) p-nitro phenol b) p-chloronitro benzene c) p-chloro methyl benzene. d) p- chloro acetophenone. e ) p-chloro benzene sulphonic acid f) 1,4-Dichloro benzene. g ) biphenyl.

16. Chloroethane to butane.

17. tert-butyl bromide to isobutyl bromide.

18. Aniline to phenylisocyanide.

19. Propene to a) 2,3-dimethyl butane b) n-hexane.

20. Tert-butyl bromide to 2-methyl prop-1-ene.

Q. 9. What happens when

a. n-butyl chloride is treated with alcoholic KOH. b. bromobenzene is treated with Mg in the presence of dry ether.

c. chlorobenzene is subjected to hydrolysis.

d. ethyl chloride is treated with aqueous KOH.

e. methyl bromide is treated with Na in the presence of dry ether.

f. methyl chloride is treated with a)KCN b)AgCN c)KNO2 d)AgNO2

Q. 10. Write the structure of the major organic product in each of the following reactions:

1.

2.

3.

4.

5.

6.

7.

Q. 11. Arrange the compounds of each set in order of decreasing reactivity towards a) SN2 displacement. b) SN1 displacement.

a. 2-bromo-2-methylbutane, 1-bromopentane, 2-bromo pentane b. 1-bromo-3-methylbutane, 2-bromo-2-methylbutane,3-bromo-2-methylbutane

c. 1-bromo butane, 1-bromo-2-methyl propane, 1-bromo-2-phenyl propane.

Q. 12. Primary halide A(C4H9Br) with alcoholic KOH gives a compound B. B on treatment with HBr gives C which is an isomer of A. A on treatment with Na in dry ether gives a compound D which is different from when n-butyl bromide is reacted with Na in dry ether. Give the structural formula of A. Write the equations of the reactions involved.

Q. 13. Distinguish chemically between a) CH3Cl, CH3Br, CHM3I b) Chloro benzene and chloro methane c) chloro benzene and benzyl chloride d) CHCl3 and CCl4.

Q. 1. Define activation energy of a reaction.

Q. 2. A cubic solid is made of two elements X and Y. Atoms Y are at the corners of the cube and X at the body centre. What is the formula of the compound?

Q. 3. State two main functions of carbohydrates in sugarcane.

Q. 4. Why is HF not stored in plain glass bottles?

Q. 5. State one use of acetonitrile.

Q. 6. What values of quantum number, m are permitted for an electron having angular quantum number, 1 = 2 ?

Q. 7. Which types of crystals exhibit piezoelectricity?

Q. 8. What is Tyndall effect?

Q. 9. How many effective sodium ions are located at the centres of faces of a unit cell in a sodium chloride crystal?

Q. 10. Name the first element of 3 d transition metal series.

Q. 11. What are inner transition metals?

Q. 12. State second law of thermodynamics.

Q. 13. Name a direct dye.

Q. 14. Sketch the zwitter ion form of amino acetic acid.

Q. 15. Mention an industrial product manufactured from methanal.

Q. 16. How does a fuel cell operate?

Q. 17. Mention two important uses of methanol.

Q. 18. What is the importance of amino acids to us?

Q. 19. Give an example of associated colloids.

Q. 20. Define order of a reaction.

Q. 1. What is the role of dimethyl and monomethyl hydrazine in rockets?

Q. 2. Name the type of structure possessed by a unit cell in CsCl.

Q. 3. Why is Copper Sulphate pentahydrate colored?

Q. 4. Give important use of eqanil and morphine.

Q. 5. Why do azo dyes not impart fast color to fabrics?

Q. 6. Why are carbohydrates generally optically active?

Q. 7. Define a ‘ligand’ and give one example.

Q. 8. What is meant by faraday constant?

Q. 9. What is basic building unit of silicates?

Q. 10. Mention the uses of formalin.

Q. 11. When is Vant’s hoff factor more than one?

Q. 12. State first law of thermodynamics.

Q. 13. Why do Zr and Hf exhibit similar property?

Q. 14. What is non stoichiometric defect in crystals?

Q. 15. State a use of enzyme Streptokinase in medicine.

Q. 16. Write composition of a rocket propellent.

Q. 17. What is the role of Bithional in soaps?

Q. 18. What are thermosetting polymers and thermoplastics?

Q. 19. Define Frenkel defect .

Q. 20. Why Boron forms electron deficient compounds?

Q. 1. State function of antihistamines along with example.

Q. 2. State the function of antioxidants along with example.

Q. 3. What are hybrid propellants?

Q. 4. State de broglie relationship.

Q. 5. Why InCl undergoes disproportionation reaction but TlCl does not?

Q. 6. Differentiate between hormones and vitamins.

Q. 7. Why Hydrogen fluoride is weaker acid than hydrogen chloride in water?

Q. 8. Name the chief ore of copper.

Q. 9. Why Nitrobenzene does not undergo Friedal craft’s alkylation.

Q. 10. What are disperse dyes?

Q. 11. Why chemistry of all lanthanides is quite similar?

Q. 12. Describe cannizzaro reaction.

Q. 13. Describe aldol condensation.

Q. 14. Give a chemical test to distinguish between etanal and propanal.

Q. 15. Why anhydrous aluminium chloride acts as a catalyst.

Q. 16. Why white phosphorus is more reactive than red phosphorus?

Q. 17. Why sulphur in vapour state exhibits paramagnetic behavior?

Q. 18. What are essential and non essential amino acids?

Q. 19. Define the term co enzyme.

Q. 20. What is meant by mutation in biomolecules?

Q. 21. State function of antihistamines along with example.

Q. 22. State the function of antioxidants along with example.

Q. 23. What are hybrid propellants?

Q. 1. Among the isomeric alkanes of molecular formula C5H12, identify the one that on photochemical chlorination yields (i) A single monochloride (ii) Three isomeric monochlorides. ( 1 mark)

Q. 2. While separating a mixture of ortho and para nitrophenols by steam distillation, name the isomer which will be steam volatile. Give reason. (1 mark)

Q. 3. Write the IUPAC name of the following.  (2 marks)

1.

2.

3.

4.

Q. 4. With the help of an example discuss the stereo chemistry involved in SN1 and SN 2  mechanism. (2 marks)

Q. 5. Explain Saytzeff rule  by taking suitable examples. (2 marks)

Q. 6. Draw the structures of all isomeric alcohols of molecular formula C5H12O and give their IUPAC names. Classify these isomers of alcohols as primary, secondary and tertiary alcohols. ( 2 marks)

Q. 7. Nitro group increases the reactivity of chlorobenzene when attched to o and p- position but not at meta position? Explain? Write the mechanism also. (3 marks)

Q. 8. Account for the following.   ( 3 marks)

1. It is necessary to avoid even traces of moisture from a Grignard reagent. 2. Chloroform is stored in closed dark coloured bottles completely filled.

3. The carbon– oxygen bond length in phenol is slightly less than that in methanol.

Q9.  (a)  Explain the fact that in aryl alkyl ethers (i) the alkoxy group activates the benzene ring towards electrophilic substitution and (ii) it directs the incoming substituents to ortho and para positions in benzene ring. (2 marks)

(b) Arrange the following compounds in increasing order of their acid strength: (1 mark)Propan-1-ol, 2,4,6-trinitrophenol, 3-nitrophenol, 3,5-dinitrophenol, phenol, 4-methylphenol

Q. 10. Give equations of the following reactions: (3 marks)

1. Preparation of Sec-Butyl alcohol from ethanal 2. Preparation of Phenol from Aniline.

3. Treating phenol with chloroform in presence of aqueous NaOH.

Q. 11. How will you distinguish between the following: (3 marks)

1. Propanoic acid and propanal 2. Ethanal and Benzaldehyde

3. Pentan-3-one and Pentan-2-one

Q. 12. How will you bring about the following conversions in not more than two steps. (5 marks)

1. Propanone to Propene 2. Benzene to m-Nitroacetophenone

3. Chlorobenzene to benzoic acid

4.  Propanoic acid to propene

5. Propyne to propan-2-ol

Q. 13. Explain the following with the help of suitable examples: (5 marks)(a) Limitation of Williamsons synthesis  (b) Stephen reaction.  (c) Hell-Volhard-Zelinsky reaction (d) Aldol Condensation (e) Wolff-Kishner reduction.

Q. 14.(a)  An organic compound (A) (molecular formula C8H16O2) was hydrolysed with dilute sulphuric acid to give a carboxylic acid (B) and an alcohol (C). Oxidation of (C) with chromic acid produced (B). (C) on dehydration gives but-1-ene. Write equations for the reactions involved (b) Predict the products formed when cyclohexanecarbaldehyde reacts with following reagents.

1. PhMgBr and then H3O+ (ii) Tollens’ reagent

Q. 15.  Primary alkyl halide C4H9Br (a) reacted with alcoholic KOH to give compound (b). Compound (b) is reacted with HBr to give (c) which is an isomer of (a). When (a) is reacted with sodium metal it gives compound (d), C8H18 which is different from the compound formed when n-butyl bromide is reacted with sodium. Give the structural formula of (a) and write the equations for all the reactions.(ii) Out of C6H5CH2Cl and C6H5 CHClC6H5, which is more easily hydrolysed by aqueous KOH.

1. What is an orbital in an atom?

2. What type of hybridization explains the trigonal bipyramidal shape of SF4

3. State deBroglie principle?

4. Which point defect lowers the density of ionic crystals?

5. Why is Frenkel defect found in AgCl ?

6. Which type of crystals exhibit piezo-electricity?

7. Explain the term mole fraction?

8. Why is benzene insoluble in water but soluble in toluene?

9. Calculate molarity and molality of a 13% solution of sulphuric acid? Its density is 1.020 gram / cm3 [Atomic mass H=1,

O=16, S=32a.m.u.]

10. KX crystals have FCC structure. What is the distance between K+and X_ in a KX crystals if the density of KX is 2.48

g/cm_3

11. With the help of molecular orbital diagram show that Ne2 can not exist as a stable molecule?

12. The ionization energy of sodium is 494.7 kj/mole calculate the wavelength of the electromagnetic radiation, which is just

sufficient to ionize the Na atom.

13. An element of atomic mass 98.5 gm/mole occurs in FCC structure if its unit cell edge length is 500 pm and its density is

5.22 gm/cc what will be the value of Avogadro constant.

14. Explain the difference between the osmotic pressure and vapor pressure in a solution. Calculate the osmotic pressure and

vapor pressure of 0.6% aqueous solution of non volatile, non electrolyte solute, urea at 25 C the vapor pressure of pure

water at 25 C is 24 mm of Hg take densities to be 1 gm/ml and assume ideal solutions behavior

15. State 2 ways in which internal energy of system can be changed.

16. What is effect of temperature on entropy

17. What is meant by Faraday constant.

18. Define the term molar conductivity.

19. What are photosensitizes.

20. Define half life of a reaction.

21. How is dialysis carried out. mention its one application.

22. What is an emulsion? Give one example.

23. What is meant by the term free energy. What is the necessity of introducing this state function.

24. How does the molar conductance of a strong electrolyte varies with its concentration in solution.

25. The rate constant of the first order reaction is 0.0005 min-1 Calculate its half life.

26. By giving suitable example distinguish between the terms adsorption and absorption assuming adsorption to be

spontaneous process show thermodynamically that it is always exothermic process.

27. Calculate the half life of a reaction where the specific rate constant is

[a] 200 sec -1 [b] 2min -1 [c] 5 year-1

28. What do you mean by the law of conservation of energy ? Derive the mathematical relation of heat, internal energy and

work and also define Ohms law?

29. Name the element present in boranes.

30. What is the state of hybridization in CO3-2

31. Why does V2O5 acts as catalyst?

32. Why is bivalent Ti+2 ion paramagnetic in nature?

33. Draw the structure of Ferrocene.

34. Name the catalyst use for the polymerization of alkenes.

35. What are particle accelerator.

36. Why are fusion reaction in nuclear chemistry referred to thermonuclear reactions.

37. Draw the structure of XeF2 and XeF4?

38. What are interstitial compound mention there two important properties.

39. Describe briefly the nature of bonding in metal carbonyl

40. Explain the principle of breeder reactor.

41. The half life period of Co-60 is 5.26 years Calculate the % activity after 8 years.

42. Describe the general characteristics of transition elements with special reference to the following

[a] enthalpy of atomization

[b] variable oxidation state

[c] complex formation

[d] formation of colored ions

[e] magnetic propert

43. Particles A & B are in motion. If the wavelength associated with particle A is 5 x 10-8m. Calculate the wavelength

associated with particle B if its momentum is half of A Calculate the wavelength of 1000Kg rocket moving with a velocity of

3000 km/hr. (h = 6.626 x 10-34Js)

44. Calculate the uncertainty in velocity of a wagon of mass 2000kg whose 10mposition is known to accuracy

45. What is the physical meaning of |x2|?

46. What do you understand by?

How do they vary with r for 1s, 2s and 2p atomic orbital of Hydrogen atom?

47. Explain why bond order of N2 is greater than bond order of N2+, but the bond order of O2 is than that of O2

+

48. What is the energy gap in the Band theory? Compare it’s size in conductors, semi-conductors and insulators

49. Which of the following substances exhibit hydrogen bonding?(a) CH3CH2OH(b) CH3CO-OH(c) CH3-CO-CH3

(d) CH3-CO-NH2

50. State 3rd law of thermodynamics. Explain its’ implications

51. Place the following systems in order of increasing randomness.

a) 1 mole of gas X.

b) 1 mole of solid X.

c) 1 mole of liquid X.

52. Why would you expect a decrease in entropy as a gas condenses into liquid? Compare it with entropy decrease when a

liquid sample is converted into solid.

53. Which of the following processes are accompanied by increase in entropy.

a) Dissolution of iodine in a solvent.

b) HCl is added to AgNO3 and a ppt of AgCl is obtained.

c) A partition is removed to allow two gases to mix.

54. Comment on the following statements.

a) An exothermic reaction is always thermodynamically spontaneous.

b) The entropy of a substance increases on going from the liquid to the vapour state at any temperature.

55. The enthalpy of vaporization of liquid diethyl ether [(C2H5)2O] is 26 kJ/mol at 350C. Calculate ∆S for conversion of Liquid to

vapour

56. For the water gas reaction

C(s) + H2O(g) --> CO(g) + H2(g)

The standard Gibb’s free energy of reaction (at 1000k) is -8.1 kJmol -1. Callculate the equilibrium constant.

57. The standard Gibb’s energy is for the reaction at 1773K are given below.

C + O2 --> CO2 , DG = -380KJ/mol

2C + O2 = 2CO DG = -120KJ/mol

Discuss the possibility of reducing Al2O3 and PbO with carbon at this temperature.

4Al + 3O2 --> 2Al2O3, ∆G = -22500KJ/mol

2Pb+O2 --> 2PbO , ∆G = -120KJ/mol

58. How is concept of coupling of reaction useful in explaining the occurrence of a non-spontaneous reaction?

59. Sodium carbonate can be obtained by heating sodium hydrogen carbonate

2NaHCO3 --> Na2CO3 + H2O + CO2

Calculate the temperature above which NaHCO3 decomposes to form product at 1bar

60. Form the rate expression for the following reactions determine their order of reaction and dimensions of the rate

constants.

a) 3NO(g) --> N2O(g) + NO2 (g) Rate = K [NO]2

b) H2O2 (aq) + 3 I - (aq) + 2H+ --> 2H2O (l) + I-1 Rate =- k [H2O] [I-]

c) CH3 CHO(g) --> CH4(g) + CO(g) Rate = k [CH3 CHO]3/2

d) CHCl3 (g) + Cl2 (g) --> CCl4 (g) + HCl (g): Rate = [CHCl3] [Cl2] 1/2

e) C2 H5 Cl(g) --> C2H4(g) = HCl(g) : Rate = k [C2H5Cl]

61. A reaction is second order with respect to a reactant. How is the rate of reaction affected if the concentration of reactant is

(I) doubled (ii) reduced to ½?

62. A reaction is first order in A and second order in B.

i) Write differential rate equation.

ii) How is the rate affected when concentration of B is tripled?

iii) How is the rate affected when the concentration of both A and B is doubled?

63. In a reaction between A and B, the initial rate of reaction was measured for different initial concentration of A and B as

given below.

[A]/M

[B]/M

r0 / Ms-5

0.2

0.3

5.07 x 10-5

0.2

0.1

5.07 x 10-5

0.4

0.05

7.6 x 10-5

64. Reaction between NO2 and F2 to give NO2 F takes place by the following mechanism.

NO2 (g) + F2(g) slow > NO2 F(g) + F(g)

NO2(g) + F(g) fast > NO2F(g)

Write the rate expression for the reaction.

65. The following result has been obtained during the kinetic studies of the reaction.

2A + B --> C+D

Experiment [A]/M [B]/M initial rate of formation of D/Mmin-1

1

2

3

4

0.1

0.3

0.3

0.4

0.1

0.2

0.4

0.1

6x 10-3

7.2x 10-2

2.88x 10-1

2.4x 10-2

66. Determine the rate law and rate constant for the reaction.

67. The reaction between A and B is first order with respect to A and zero order with respect to B. fill in the blank in the

following table.

Experiment [A]/M [B]/M initial rate / M min-1

1

2

3

4

0.1

-

0.4

-

0.1

0.2

0.4

0.2

2x 10-2

4x 10-2

-

2x 10-2

68. The rate constant for a first order reaction 60 s-1. How much time will it take to reduce the initial concentration of the

reactant to its 1/16th value?

69. The rate of most of the reactions double when their temperature is raised from 298k to 308k. Calculate their activation

energy

70. Give are example each of (i) a-emission (ii) b -emission (iii) k-capture. Write the equation for these nuclear changes.

71. What is Group Displacement Law? An element belonging to group 1 decays by b-emission. To which group of the periodic

table the daughter element will belongs.

72. How many a and b particles will be emitted when 232/90 Th Changes into 208/82 Pb.

73. Write the nuclear reactions for the following radioactive decay.

a) 238/92U Undergoes a - decay

b) 234/91Pa Undergoes b - decay

c) 22/11Na Undergoes b + decay

74. What is meant by nuclear binding energy calculate binding energy per nucleon of Li isotope, which has isotopic mass of

7.016 mu . The individual masses neutron and proton are 1.008665mu and 1.007277mu respectively and mass of electron

is 0.000548mu.

75. A sample of wood from an archeological source shows a 14C activity which is 60% of the activity found in fresh wood

today. Calculate the age of archeological sample

(t1/2 of 14c =5770 years)

76. Explain the principle of (a) activation analysis and (b) Breeder reactor

77. Complete the following nuclear reaction

a. 96/42Mo ( _______n) 97/43 Tc

b. _______ (a, 2n ) 211/85 At

c. 55/25Mn (n,c) ______

d. 246/96Cm + 12/6 C --> ________ + 4 1/0 n

e. 27/13Al (a, n) _______

f. 238/92U (a, b-)________

78. Complete the equations for the following nuclear processes

a. 3517Cl + 1/0n --> ________+ 4/2 He

b. 235/92 U + 1/0 n --> _______+ 137/54 Xe + 2 1/0 n

c. 27/13 Al + 4/2 He --> _______ + 1/0 n

d. _______ (n, p) 35/16 S

e. 239/94 Pu (a, b- ) _______

79. Describe the principle of an atom bomb. What is meant by critical mass? What is the critical mass of 235/92U?

80. Write the structures of monomers used for getting the following polymer.(a) PVC (b) Teflon (c) PMMA.

81. How does the presence of carbon tetra chloride influence the cause of vinylic free radical polymerization? Explain with an

example.

82. How does vulcanization change the character of natural rubber?

83. Why are the number 66 and 6 put in the name of nylon 66 and nylon 6?

84. Illustrate with equations how is nylon-6 prepare from caprolactam.

85. What is the difference between thermoplastic and thermosetting polymers?

86. What are reducing and non-reducing sugars? What is the structural features characterising reducing sugars?

87. Explain mutarotation. Give its mechanism in the case of D-glucose

88. Give reasons for the following.

(i) Amino acids have relative higher melting points as compared to corresponding halo acids.

(ii) Amino acids are amphoteric in behavior.

(iii) On electrolysis in acidic solution amino acids migrate towards cathode while in alkaline solution these migrate towards

anode.

89. What type of linkages are responsible for the formation of,

(i) Primary structure of proteins.

(ii) Cross linking of polypeptide chains.

90. What is denaturation and renaturation of proteins?

91. What are the products obtained on complete hydrolysis of DNA?Write down the structures of pyrimidine and purine bases

present in DNA.

92. Enumerate the structural differences between DNA and RNA. Write down the structure of a nucleoside, which is present

only in DNA.

93. What are complementary bases? Draw structure to show hydrogen bonding between adenine and thymine and between

guanine and cytosine.

94. What is the melting temperature (Tm) of DNA? A DNA molecule with more number of GC base pair than AT base pairs

has higher Tm than one with lesser number of GC base pairs than AT base pairs. Why?

95. How does DNA replicate? Give the mechanism of replication? How is the process responsible for preservation of

heredity?

i) Name the location where protein synthesis occurs.

ii) How do 64 codons code for only 20 amino acids?

iii) During translation which one of the two-end functional groups of the polypeptide is formed first?

97. How are lipids classified? Give an example of each class.

98. Name the deficiency diseases caused due to lack of vitamin A, C, E, B1, B12, B6 and K.

99. How are antiseptics distinguished from disinfectants Give two examples of each of the substances

100.List two major classes of antibiotics with an example of each class.

i) Narrow spectrum antibiotic

e.g. : penicillin

ii) Board spectrum antibiotic

e.g. :chloramphenicol and tetracycline

101.Describe the following with suitable examples.

(i) Tranquilizers (ii) antifertility drugs (iii) antihistamines

102.Give examples of (i) triphenyl methane dye (ii) azo dye (iii) anthraquinone dye

103.What is a mordant dye? How is it applied to the fabric?

104.What are carbon fibers? How are they designed? Write two important uses of carbon fibers.

105.Describe the following with suitable examples

(i) Preservative (ii) Artificial sweeteners (iii) Antioxidants (iv) Edible colors

106.What are detergents? Give their scheme of classification. Why are detergents preferred over soap?

107.What is a propellant? How are various rocket propellants classified?

108.Describe the following with examples

(i) Double base propellant

(ii) Biliquid propellant

(iii) Monoliquid propellant

(iv) Hybrid propellant

109.Silver crystallizes in FCC lattice. If edge length of unit cell is 4.077x 10-8 cm and density is 10.5 g/cm3 , calculate atomic

mass of silver.

110.A cubic solid is made of two elements P & Q. Atoms Q are at the corners of the cube and P are at the body center. What

is the formula of the compound? What are the Co-ordination numbers of P & Q?

111.If the radius of octahedral void is ‘r’ and the radius of atoms in close packing is R, derive a relation between r & R.

112.Formula mass of NaCl is 58.45 g/mol and density of its pure form is 2.167 g/cm3. Average distance between adjacent

sodium and chloride ions in the crystal is 2.814x10-8 cm.

M =5 8.45 g/mol

ρ = 2.167 g/cm3

113.Analysis shows that nickel oxide has formula Ni0.98 O1.00. What fractions of Nickel exist as Ni2+ and Ni3+ ion?

114.If the radius of the Bromide ion is 0.182 nm, how large a cation can fit in each of the tetrahedral hole?.

115.Classify each of the following as p-type and n-type semiconductors.

i) Ge doped with in : Ans : p-type

ii) B Doped with Si Ans : n-type

116.Aluminium crystalises in a cubic close packed structure. Its’ metallic radius is 125 pm.

(a) What is the length of the side of the unit cell.

(b) How many unit cells are there in 1cm3 of aluminium

117.If NaCl is doped with 10-3 mol% SrCl2 , what is the concentration of cation vacancies

118.Calculate the value of Avogadro constant from the following data. Density of

NaCl = 2.165 g/cm3. Distance b/w Na+ & Cl- is 281 pm

119.Concentrated nitric acid used in the laboratory work is 68% nitric acid by mass in aquous solution. What should be the

molarity of such a sample of the acid if the density of the solution is 1.504 g/mL.

120.An antifreeze solution is prepared from 222.6g of ethylene glycol. C2H4(OH)2 and 200g of water. Calculate molality of

solution. If the density of the solution is 1.072 g/ml then what shall be the molarity of the solution?

121.What role does the molecular interaction play in solution of alchohol and water?

122.State Henry’s law and mention some important applications.

123.What is meant by positive and negative deviation from Raoults’ law and how is the sign DH related to positive & negative

deviation from Raoults’ Law?

124.Heptane and octane form ideal solution. At 373K, 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 25g of heptane and 35g of octane.

125.A solution containing 30g of non-volatile solute exactly in 90g water has a vapour pressure of 2.8 kPa at 298K. Calculate

1) Molecular mass of solute Vapour pressure of water at 298K .

126.Two elements A & B form compounds having molecular formula AB2 & AB4. When dissolved in 20g of C6H6, 1g AB2 lowers

the freezing point by 2.3 & 1g AB4 lowers it by 1.3K.The molar depression constant for benzene is 5.1Kg mol-1. Calculate

atomic mass A & B.

127.At 300k, 36g of glucose present per litre in its solution has an osmotic pressure of 4.98 bar. If the osmotic pressure of

solution is 1.52 bar at the same temperature, what would be the concentration.

Q. 1. How many millimoles of N2 gas would dissolve in 1 litre of water at 293 K? N2 exerts a partial pressure of 0.987 bar. (Henry's

Law constant at 293 K for N2 is 76.4 K bar)

Q. 2. Na2CO3 and NaHCO3 mixture containing Ig is neutralised by 0.1M HCI. Find the volume of HCI equired if the mixture contains

equimolar amounts of Na2CO3 and NaHCO3.

Q. 3. Calculate the percentage composition in terms of mass of a solution obtained by mixing 300g of 25% and 400g of 40%

solution by mass.

Q. 4. A sample of drinking water contains 15 ppm of CHCI3 (by mass). Express this in percentage by mass. Determine the molality

of CHCI3 in solution.

Q. 5. The partial pressure of ethane over a saturated solution containing 6.56 x 10-2 g of ethane is 1 bar. If the solution containing

5.00 x 10-2 g, then what will be the partial pressure of the gas.

Q. 6. An aqueous solution @2% non-volatile solute exerts a pressure of 1.004 bar at the boiling point of the solvent. Find the molar

mass of the solute.

Q. 7. Heptane and octane form an ideal solution at 373K. Vapour pressure of heptane and octane are 105.2 kPa and 46.8 kPa

respectively. Find the vapour pressure in bar of a mixture of 25g heptane and 35g octane.

Q. 8. Vapour pressure of water is 12.3 kPa at 300K. Calculate the vapour pressure of 1 molal solute in it.

Q. 9. Two elements A and B form compounds AB2 and AB4. When dissolved in 20g benzene, Ig of AB2 lowers the freezing point by

2.3K, Whereas Ig AB4 lowers the freezing point by 1.3K. Kf for benzene = 5.1 K kg mole-1. Calculate the atomic mass of A and B.

Q. 10. At 300K, 36g glucose (molar mass 180) in 1 litre solution exerted an osmotic pressure of 4.98 bar. What would be the

concentration of the solution at 300K if the osmotic pressure is 1.52 bar.

Q. 11. Concentrated HN03 is 63% by mass. Density of the solution is 1.5g/cm3. Calculate the volume of the solution which contains

20g HNO3 (Molar mass HNO3 = 63).

Q. 12. Concentrated H2SO4 is 49% by mass. Density of the solution is 1.5g/cm3. Determine molarity, olality and normality of the

solution. (Molar mass of H2SO4 = 98)

Q. 13. Calculate the molality of 1M solution of NaNO3. Density is 1.25 g/cm3. (Molar mass of NaNO3 = 85)

Q. 14. Calculate the number of moles of CH3OH in 5 litres of 2 molal solution. Density is 0.981 g/cm3. Molar mass of CH3OH = 32)

Q. 15. Calculate the volume of 80% H2SO4 (D =1.8 g/cm3) required to prepare 1 litre of 20% H2SO4 (D =1.25 g/cm3)

Q. 16. 1.8 grams glucose (Molar mass = 180) is dissolved in 36g of water. Calculate the molality and the mole fraction of glucose in

the solution.

Q. 17. Sea water contains 5.8 x 10-3g dissolved O2 per kilogram. Express the concentration in ppm.

Q. 18. Calculate the resulting molarity of the solution prepared by adding 5g NaOH to 200ml M/4 NaOH solution. (D = 1.05g/cm3).

Density of the resulting solution is 1.08 g/cm3.

Q. 19. Benzene and toluene forms an ideal solution at 300K. po benzene = 160mm Hg. potoluene = 60mm Hg. Calculate the partial

pressure of benzene and toluene and the total pressure under the following conditions:-

i. Mixing equal number of moles of benzene and toluene

ii. Equal masses of benzene and toluene

iii. Equal amount of benzene and toluene

Q. 20. A solution containing 1 mole of X and 3 moles of Y gave a vapour pressure of 550mm Hg. The same solution containing 1

mole of X and 4 moles of Y gave a vapour pressure of 560mm Hg. Find the vapour pressure of pure X and pure Y.

Q. 21. Vapour pressure of a solution containing benzene and toluene is 180x + 120 at 200K where x is the mole fraction of toluene.

Find vapour pressure of pure benzene and pure toluene.

Q. 22. Vapour pressure of aqueous dilute solution of glucose is 750mm Hg at 373K. Find the molality and mole fraction of glucose

in the solution.

Q. 23. Vapour pressure of ethanol (Molecular mass = 46) and methanol (Molecular mass = 32) at 300K are 45mm Hg and 90mm

Hg respectively. Find the total pressure of the solution containing 92g ethanol and 16g methanol.

Q. 24. (a) Vapour pressure of pure benzene (Molecular mass = 78) at 300K is 640mm Hg. 2.175g of non-volatile solute in 39g

benzene gave a vapour pressure of 600mm Hg. Determine the molar mass of the solute.

(b). 5% Solution of Sucrose(Molarmass=342)is isotonic with 0.877%Solution of urea.Find the molar mass of urea

Q. 25. A very small amount of solute in 60cm3 benzene showed a vapour pressure of 98.88mm Hg. Vapour pressure of pure

benzene is 100mm Hg at this temperature. Find the molality of the solution. If ATf = 0.73K find Kfof benzene.

Q. 26. A solution of sucrose (Molecular mass = 342) is prepared by dissolving 68.4g in lOOg H20. Determine

i. Vapour pressure of solution at 298K.

ii. Osmatic pressure at 298K.

iii. Boiling point of the solution

iv. Freezing point of the solution

Given Kb = 0.52 K Kg mole-1, Kf = 1.86 K Kg mole-1 , R = 0.0821 I atm K-mole-1

Q. 27. 34.2g sucrose and 36g glucose are dissolved in 81g H20. Find the vapour pressure of the solution. Vapour pressure of H2O

= 30mm Hg. Molecular mass of sucrose = 342, Glucose = 180.

Q. 28. Calculate the boiling point and freezing point of 1M solution of KCI. D = 1.04 g/cm3. Molar mass of KCI = 74.5. Kb = 0.52 K

Kg mole-1 Kf = 1.86 K Kg mole-1. Assume KCI is 90% dissociated.

Q. 29. BaCI2 cand KCI mixed in 1:1 molal ration showed ΔTb = 2.6K. Determine the amount of each solute in 100 g of the solvent.

Kb = 0.52 K Kg mole-1. Molar mass KCI = 74.5, BaCI2 = 208.

Q. 30. Kf of benzene 4.90K kg mole-1. 3.26g Se (Atomic mass = 78.8) in 226g of benzene showed a freezing point of 0.1120C lower

than pure benzene. Find the molecular formula of Se.

Q. 31. A solution containing 1.017g Napthalene (Ci0H8) in 100g CCI4 gave ΔTb = 0.40C. 1.24g of an unknown< solute in 100g CCI4

gave ΔTb = 0.620C. Find molar mass of the unknown solute.

Q. 32. 1.8g glucose in 100ml solution is added to 34.2g sucrose in 100ml. Find the osmatic pressure of the resulting solution.

Q. 33. 2 grams benzoic acid (C6H6COOH) in 25g benzene gave ΔTf = 1.62K. Kf = 4.9K kg mole-1. Find the %association of benzoic

acid if it exist as a dimer in solution.

Q. 34. Which of the following aqeous solution will have:-

1. Lowest freezing point

2. Highest freezing point

3. 0.1M NaCI, O.1M BaCI2, 0.1M AI2(SO4)3, 0.1M Urea

Q. 1. Kinetic energy of a moving electron is 4.55×10-25J.Determine the wavelength of the electron.[mass of electron is 9.1×10-31 Kg,

h is 6.63×10-34 Jsec]

Q. 2. Kinetic energy of a particle is 6.63 × 10-25 J. Determine its frequency .

Q. 3. Calculate the wavelength of α particle moving with a velocity of 3×103 m/sec.

Q. 4. Calculate the momentum of an electron moving with a wavelength of 7.5 × 10-6 m.

Q. 5. Calculate the product of uncertainty in velocity and uncertainty in position of moving electron.

Q. 6. The uncertainty in the position of a moving bullet of mass 10g is 10-5m. Calculate uncertainty in its velocity.

Q. 7. On the basis of uncertainty principle, show that electron cannot exist with in the atomic nucleus (radius = 10-15m)

Q. 8. Two particles A and B are in motion. If the wavelength of A is 5 x 10-8m. Find wavelength of B if the Momentum of B is half the

momentum of A.

Q. 9. The Sodium flame has λ = 589nm.What is the mass equivalence of one photon of this wavelength.

Q. 10. Calculate the uncertainty in velocity of wagon of mass 2000kg whose position is known to an Accuracy of ± 10m.

Q. 11. Calculate the energy required to excite an electron in the hydrogen atom from the first energy level to fourth energy level.

Find the frequency, wave number and wave length of the light emitted when the

Electron moves back to the ground state. Ionisation energy of hydrogen atom is 1312KJ/Mole.

Q. 12. Calculate the change in position of an electron in its ground state of hydrogen atom when it is supplied with an energy of

2.04×10-18J.

1. H2+ 1 H3→ 2He4 + 0 n1 Calculate the energy released in (a)Mev/atom (b) Mev/mole (c)J/mole (d)J/atom

(e) J/g Given 1 H2=2.014 1 H3=3.016 2He4=4.003 0 n1 =1.009 (all in amu)

Q. 2. Mass defect of Cl35 is 0.320 amu. Calculate the binding energy/nucleon

Q. 3. 92U235+0n1 → 42 Mo98 + 54Xe136 + x −1eo +y0 n1 Find x and y and energy released per atom.

. 92U235=235.044 0 n1=1.009 Mo98=97.900 Xe136=135.907 (all in amu)

Q. 4. How many α and β particles are emitted in the reaction 90Th232 → 82Pb208

Q. 5. t½ of a radioactive element is 100seconds.Find λ How long will it take for 1g element to become 0.01g?

Q. 6. A radio active element decays in such a way that after 100 minutes ⅛th of the original amount remains.

Calculate t½ and λ.

Q. 7. Starting with 1g of a radio active element 0.25g is left after 5 days. Find the amount left after 1 day.

Q. 8. A piece of wood shows C14 activity which is 60% activity found today. Find the age of the sample.t½=5770years

Q. 9. t½ Hg 203=50 days. How much of 0.2g of sample left after 6 months.

Q. 10. A sample of pitch blend contains 0.05 gram Pb206 for every gram of U238. Find the age of the rock

t½ U238=4.5 x 109years.

Q. 11. A sample of pitch blend contain 5mg of U238 and 1.8mg of Pb208. Find the age of the rock. t½U238=4.5x109year

Q. 12. A piece of wood shows C14 activity 3.8 counts/minute/g of C. Determine the age of the wood if freshly cut wood shows an

activity of 15.2 counts/minute/g of C. t½ of C14=5770years.

Q. 13. A rock contain U238 and Pb208 in the ratio 3:2. Find the age of the rock. t½U238=4.5x109year.

Q. 14. Complete the equation: (a) 7N14 (α,−) 8O17 (b) 13Al27 (α,n)……… (c) 25Mn55 (n, gamma)…….

(d) 96Cm246 + 6C12→ 102No254 + ……..

Q. 15. 84Po218 under goes an alpha emission followed by two beta emission. Find the position of the new element in the periodic

table.

Q. 16. Calculate t½ for Am241. Given it emits 1.2 x 1011αparticles per gram per second.

Q. 17. The isotopic composition of Rb is Rb 85 = 72% and Rb 87 = 28%. Rb87 is weakly radio active and decay by β− emission

with decay constant 1 x 10−11/year. A mineral pollucite contain 450milligram Rb and 0.72 mg Sr87. Find the age of the sample.

Q. 18. Calculate the mass of La140 in sample whose activity is 3.7 x 1010Bq. t½ La 140=40 hrs.

Q. 19. β−activity of CO2 prepared from a wood give 25.5 counts/minute. The same mass of CO2 from an ancient source gave 20.5

counts/minute. Find the age of the wood. What would be the expected count rate for an identical mass of CO2 from a sample which

is 4000 years old.

Q. 20. During nuclear explosion, one of the product is Sr90 with t½ = 28 years. If 1 microgram of Sr90 was absorbed in the bones

of a new born baby instead of calcium. How much of it will remain after 10 years.

Q. 21. During nuclear fission of U- 235 , 17.5 MeV of energy is generated per nucleus fissoned. Find the energy released in joules

by fissioning 1gram of U-235.

( 1 Mev = 1.622x 10-13 J NA = 6.023X 1023 mol-1

Q. 22. 2NO+ O2→ 2NO2(g) How will the rate changes when the volume of the reaction vessel is reduced to one - thirds

(g) (g 1. 2NO2+F2--> 2NO2F. Write the rate of reaction in terms of (a) rate of formation of NO2F (b) rate of disappearance of NO2 (c)

rate of disappearance of F2.

Q. 2. The decomposition of NH3 follows zero order. 2 NH3→ N2+3H2 . Find the rate of production of N2 andH2.K = 2.5 x 10-4 MS-1.

Also find rate of decomposition of NH3.

Q. 3. 2A+B+C→A2B + C Rate=K(A)(B)2 K=2x10-6M−2S-1. Calculate the initial rate when (A)=0.1M  (B) = 0.2M  (C)=0.6M. Find the

rate when 0.04mole of (A) is consumed.

Q. 4. 2NO2+F2→2NO2F

Experiment ; (NO2)M (F2) M Rate(M/S)

1 0.2 0.05 0.006

2 0.4 0.05 0.012

3 0.8  0.10 0.048

Find the order with respect to NO2 and F2. Also find the overall order of the reaction. Deduce the mechanism of the reaction.

Q. 5. A first order reaction is 20%complete in 10minutes. Find the time taken for 80%completion of the reaction. Also find the half

life of the reaction.

Q. 6. Show that(a) 2t (1/2) = t (3/4) (b) Half life of a reaction is 10seconds.Find t (2/3),

Q. 7. The pressure of a gas decomposing at a metal surface of a solid catalyst are given below:

t/s 0 100 200 300

P/Pa 4.00x102 3.50x103 3.00x103 2.5x103

Determine the order of the reaction. Find the rate Constant  and the half life of the reaction.

Q. 8. Hydrolysis of methyl acetate in aqueous solution has been studied by titrating liberated acetic acid with NaOH rate

=K(CH3COOCH3) (H2O)

t/min 0 30 60   90

c/M 0.8500 0.8004 0.7538 0.7096

Show that it follows pseudo first order reaction as the concentration of water remains constant (1L of water=1000g Of

water=1000/18=55.5M) What is the value of K?

Q. 9. The rates of a reaction starting with initial concentrations 2 x 10-3 M and1 x 10-3 M are 2.4 x 10-4 M/s and 0.6 x 10-4 M/s

respectively. Find the order of the reaction and rate constant K.

Q. 10. A + 5B + 6C → 3L + 3M

Experiment (A)M  (B)M (C)M Rate M/minute

1. 0.02 0.02 0.02 0.00208

2. 0.01 0.02 0.02 0.00104

3. 0.02 0.04 0.02 0.00416

4. 0.02 0.02 0.04 0.00832

Determine the order with respect to each reactant. Find K .Calculate the initial rate when concentration of each reactant is

0.01M.Find the initial rate of change in concentrations of B and L

Q. 11. Rate of a reaction becomes 1.414 times when concentration of the reactant is doubled. Find the order of the reaction.

Q. 12. (a) show that for a first order reaction t½ is independent of the initial concentration of the reactant. (b) show that for a zero

order reaction t½ is directly proportional to initial concentration of the reactant and inversely proportional to rate constant.

Q. 13. Rate constant of a reaction is 2M-1 S-1 at 700K and 32 M-1 S-1 at 800K.Find Ea

Q. 14. Rate of a reaction becomes 4 times when temperature changes from 270C to 370C. Find Ea.

Q. 15. Rate constant of a reaction at 700K and 760K are 0.01 M-1 S-1 and 0.105 M-1 S-1 respectively. Find A and Ea .

Q. 16. Rate constant of a reaction increases by 7% and the equilibrium constant increases by 3% when the temperature changes

from 300K to 301k.Find Ea for the for ward and backward reaction.

Q. 17. Rate = K1(NH3)|1 + K2(NH3) Find the order when concentration of (a)NH3 is very low (b) when Pt surface is completely

covered NH3 2NH3 → N2+3H2 using Pt as a catalyst.

) of its original volume?