Chemistry Class XI

8/15/2019 Chemistry Class XI

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Chapter 1 : Some Basic Concepts

Very Short Answer Questions

1. The number of molecule in 16 g of methane is :(a) 30 1023. (b) 6 022 1023.

(c) 16 6022 1023/ . (d) 16 30 1023/ .

2. The maximum amount of BaSO4 precipitated on mixing BaCl2 (0.5 M) with H SO2 4 (1 M) will

correspond to :

(a) 0.05 M (b) 0.5 M (c) 1.0 M (d) 2.0 M

3. What is the number of significant figures in 1 050 104 ?

4. What is AZT ? To which use is it being put ?

5. What is the S.I. unit of molarity ?6. Briefly explain the difference between precision and accuracy.

7. Why atomic mass is an average value ? Explain with a suitable example.

Short Answer Questions

8. Nitrogen forms a number of oxides. Write their formulae and give their names. Give appropriate

calculations, explain the law that follows from it ?

9. Why do we regard the gaseous state of water as vapours while that of ammonia as gas ?

10. The percentages of all the elements present in a compound are 92. What does it indicate ?

11. What is the difference between molarity and molality ?12. What is the difference between the mass of a molecule and molecular mass ?

13. A compound made up of two elements A and B has A = 70%, B = 30%. Their relative number of

moles in the compound are 1.25 and 1.88. Calculate :

(a) Atomic masses of the elements A and B.

(b) Molecular formula of the compound, if its molecular mass is found to be 160.

14. A 25 cm3 of 0.2 M solution metal chloride (MCl )x reacted with 150 cm3 of 0.1 M AgNO3 solution

completely to form the precipitate of AgCl. What is the formula of metal chloride ?

Long Answer Questions15. Commercially available concentrated hydrochloric acid contains 38% HCl by mass.

(a) What is the molarity of the solution if its density if 1.19 g cm-3 ?

(b) What volume of concentrated HCl is needed to make 1.0 L of 0.2 M HCl solution ?

16. Calculate the volume of 1.0 M of aqueous solution of sodium hydroxide that is neutralised by

200 mL of 2.0 M aqueous hydrochloric acid and the mass of sodium chloride produced.

[Chemistry/XI] [ 144 ]


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17. A 2g mixture of Na CO2 3 and K CO2 3 was dissolved in water to form 100 cm3 of the solution. 20

cm3 of this solution required 40 cm3 of 0.1 N HCl solution for neutralisation. Calculate the

percentage composition of the mixture.

18. How many millilitres of 0.5 M H SO2 4 are needed to dissolve 0.5 g of copper carbonate ?

19. One litre of a solution of N/2 HCl was heated in beaker and it was observed that when the volumeof solution got reduced to 600 mL, 3.25 g of HCl was lost. Calculate the normality of the

resulting solution.



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Chapter 2 : Structure of Atom


1. As the nuclear charge increases from neon o calcium, the orbitals energies :(a) increase (b) increase very rapidly

(c) increase very slowly (d) energy remains constant

2. The de-Broglie wavelength of a particle with mass 1 g and velocity 100 m/s is :

(a) 663 10 33. m (b) 663 10 34. m

(c) 663 10 35. m (d) 663 10 36. m

3. Define an (i) Isotope, (ii) Isobar and (iii) Isotone.

4. The frequency of the strong yellow line in the spectrum of Na is 509 1014. /s. Calculate the

wavelength of the light in manometer.5. Which quantum number determines (i) energy of electrons, (ii) orientation of orbital ?

6. Why is 4 s orbital filled before 3d orbital ?

7. How many unpaired electrons are present in N(7) ? Name the principle which explains the

presence of these unpaired electrons.


8. Explain how can you say electrons and protons are fundamental particles of all the atoms ?

9. Describe Rutherford’s model of atom.

10. Calculate de-Broglie wavelength of an electron travelling with a speed equal to 10% of the speedof light.

11. Write short note on Planck’s quantum theory.

12. Calculate the energy per proton associated with the following radiations :

(a) Radiations of frequency 3 1015 /s (b) Radiation of wavelength = 40 nm

(h c 662 10 3 1034 8. ;Js m/s)

Long Answer Questions

13. Give main achievements of Bohr’s model of atom.

14. Explain why atoms half filled and completely filled orbitals have extra stability.

15. What is meant by dual nature of electron ? Calculate the wavelength of an electron having mass

91 10 31. kg and kinetic energy 455 10 25. kJ.

16. (a) Write the electronic configuration of Cu ion (Z=29).

(b) Compare the energies of the two radiations with wavelength 6000 Å and 4000 Å.



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(c) Which one of the following is not possible and why :

2 s, 2d , 4 f , 5 p

17. Nitrogen laser produces a radiation at a wavelength of 337.1 nm. If the number of photons

emitted is 56 1024. , calculate the power of this laser.

18. What were the main points of Electromagnetic wave theory ? What were its limitations ? Howhave these been overcome by Planck’s quantum theory ?

19. Write a short note on the following :

(a) Solar spectrum or continuous spectrum

(b) Atomic spectra or line spectra



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Chapter 3 : Classification of Elements

and Periodicity in Properties


1. What is meant by periodicity of properties ?

2. Why do elements with similar properties occur in the same group ?

3. With which quantum numbers does every period in periodic table begin ?

4. To which series do man made elements belong to ?

5. How do the basic character and solubility in water vary from Be(OH)2 to Ba(OH)2 ?

6. Out of Na and Mg, which has higher second ionisation energy ?

7. Atomic number (Z) of an element is 108. Write its electronic configuration and name the group

to which it belongs.8. Anything that influences the valence electrons will affect the chemistry of the following factors

does not affect the valence shell :

(a) Valence principal quantum number (n)

(b) Nuclear charge (Z)

(c) Nuclear mass

(d) Number of core electrons

9. Considering the elements F, Cl, O and N the correct order of their chemical reactivity in terms of

oxidizing property is :

(a) F > Cl > O > N (b) F > O > Cl > N(c) Cl > F > O > N (d) O > F > N > Cl


10. Arrange the following in increasing order :

(a) BeCO , BaCO , CaCO , MgCO3 3 3 3 (Thermal stability)

(b) BeCl , BaCl , SrCl , CaCl2 2 2 2 (Ionic character)

11. the first (IE )1 and second (IE )2 ionisation energy (Kj/mol) of a new element designated by roman

numerical are shown below :

IE1 IE2I 2372 5251

II 520 7300

III 900 1760

IV 1680 3380

12. Give four characteristics of d -block elements ?

13. Explain why ionisation enthalpies decreases down the group of the periodic table ?



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14. Among the elements B, Al, C and Si :

(a) Which has the highest first ionisation enthalpy ?

(b) Which has most negative electron gain enthalpy ?

(c) Which has the largest atomic radius ?

(d) Which has the most metallic character ?Long Answer Questions

15. (a) Explain the second ionization energy of B is significantly higher than second ionization

energy of C, even though the first ionisation energy of B is less than B.

(b) Which has higher electron affinity F or Cl ? Why ?

(c) Why noble gases having positive electron gain enthalpy ?

16. (a) The element 119 has not been discovered. What would be IUPAC name and symbol for this

element ? On the basis of periodic table, predict the electronic configuration of this

element and also the formula of its most stable chloride and oxide.

(b) Define electro negativity.

17. (a) Predict the position of the element in the periodic table satisfying the electronic

configuration ( )n d ns 1 1 2 when n = 4.

(b) Name the species which is isoelectronic with Cl.

(c) Why f block elements are placed in a separate row at the bottom of periodic table ?

(d) I.E. of nitrogen is greater than oxygen.

(e) Write general electronic configuration of inner transition elements.

18. Predict the formulas of the stable binary compounds that would be formed by the combination of

the following pairs of elements.

(a) Lithium and Oxygen (b) Magnesium and Nitrogen

(c) Aluminium and Iodine (d) Silicon and Oxygen

(e) Phosphorus and Fluorine

19. (a) Which is largest in size Cu , Cu Cu+ 2+, and why ?

(b) Which element in periodic table has highest I.E. ?

(c) Which element is more metallic Mg or Al and why ?

(d) Give advantages of long form of periodic table.



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Chapter 4 : States of Matter


1. At constant volume for a fixed number of moles of a gas, the pressure of the gas increases withrise of temperature due to :

(a) Increase in average molecular speed

(b) Increased rate of collision amongst molecules

(c) Increase in molecular attraction

(d) Decrease in mean free path.

2. Which of the following exhibits the weakest intermolecular forces ?

(a) NH3 (b) HCl (c) He (d) H O23. How is molar mass of gas related to rate of diffusion ?

4. Why is moist air lighter than dry air ?5. What is the effect of temperature on viscosity and why ?

6. Under what conditions of T and P, most of gases deviate from ideal gas behaviour ?

7. State and explain Dalton’s law of partial pressure. Prove that partial pressure of a gas is equal to

the product of its mole fraction and total pressure in gaseous ?


8. Explain the physical significance of vander Waal’s parameters.

9. The drain cleaner, Drainex contains small bits of aluminium which react with caustic soda to

produce dihydrogen. What volume of dihydrogen at 20°C and 1 bar will be released when 0.15 gof aluminium reacts ? [Atomic mass of Al is 27 g/mol]

10. Butane gas is burnt in oxygen to give CO2 and H O2 .5 L of C H4 10 is burnt in excess of O2 at 27°C

and 1 atm. Calculate the volume of CO2 formed at same temperature and pressure. Also

calculate volume of CO2 at 67°C and 2 atm pressure.

11. What will happen to volume of fixed amount of gas at a certain T and P if :

(a) T is kept constant but pressure is decreased to 1/4th of the original value ?

(b) Pressure is halved and temperature in Kelvin is doubled ?

12. (a) Why is Boyle’s law is obeyed by N , O2 2 and CO2 only at low pressure and high

temperature ?

(b) Compare the rate of diffusion of HCl and NH4 (Atomic Masses of H = 1u, Cl = 35.5 u, N =

14 u )


13. Give various postulates of kinetic theory of gases.



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14. Calculate the total pressure in a mixture of 8 g of dioxygen and and 4 g of dihydrogen confined in

a vessels of 1 dm3 at 27°C. (R = 0.083 bar dm3 /Kmol)

15. Calculate the pressure exerted by 1.00 mol of CO2 (g) at 298 K that occupies 65.4 ml using van

der Waal’s equations.

‘a’ for CO2 is 3.592 L2

bar/mol2

, ‘b’ = 0.0427 L/mol.Compare it with the pressure predicted by ideal gas equation for same conditions of T and P ?

16. Account for the following :

(a) The size of weather balloon becomes larger and larger as it ascends up into higher

altitudes.

(b) Copper is malleable and ductile while sulphur is not.

17. What will be the pressure excited by a mixture of 3.2 g of methane and 4.4 g of carbon dioxide

contained in a 9 dm3 flask at 27°C ?

18. N O2

and CO2 have the same rate of diffusion under same conditions of temperature and

pressure. Why ?

19. Discuss the nature of the gas constant ‘R’. Derive its value in terms of different units.



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Chapter 5 : Chemical Bonding and Molecular Structure


1. NH3 and BF3 form adduct readily because they form :(a) Ionic bond (b) Covalent bond

(c) Co-Ordinate bond (d) Hydrogen bond

2. The ion which is iso-electronic with CO is :

(a) CN (b) O2 (c) N 2

(d) O2

3. Name two conditions which must be satisfied for hydrogen bonding to take place in a molecule.

4. What is valence bond approach for the formation of covalent bond and a coordinate bond ?

5. Define the term bond order and find the bond order of O 2.

6. Give the structure of sulphur tetra fluoride.

7. With what neutral molecule is ClO isoelectronic ?


8. Define lattice energy. On what factors doe sit depend ? How does it help to predict the stability

of the ionic compound formed ?

9. What are sigma and pi bonds ? Explain the difference ways of their formation diagrammatically.

Which one of them is stronger and why ?

10. Explain why carbon has a valency of four and not two and why are the four C—H bonds in

methane identical ?

11. The boiling and melting points of water are abnormally higher than those of other hydrides of

group 16 of the periodic table. Give reasons.

12. Discuss the shapes of molecules orbitals formed by the combination of the following atomic

orbitals :

(a) 2p + 2px x (b) 2p + 2pz z13. Using VSEPR theory, draw the molecular structure of OSF4 and XeF4 indicating the location of

lone pair (s) of electrons and hybridisation of central atoms.

14. Which of the following species has the shortest bond length ?

NO, NO , NO , NO+ 2+


15. Explain the formation of covalent bond on the basis of (a) Lewis concept (b) Valence Bond

Theory taking at least three examples in each case.

16. What is Resonance ? Explain with a suitable example. Define resonance energy.



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17. What are bonding and anti bonding molecular orbitals ? Describe LCAO method for their

formation. What are the important characteristics in each case ?

18. (a) Briefly explain Kossel-Lewis approach of chemical bonding.

(b) NaCl is a better conductor of electricity in a molten condition than in the solid state.

Explain.(c) Which of the following hydrogen halides has the most polar molecules and why ?

HI, HBr, HCl, HF

19. (a) Define dipole moment. Draw dipole diagram of H O2 and BF3.

(b) Explain the term hybridization taking the example of methane.'



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Chapter 6 : Thermodynamics


1. Under what conditions H and U are equal ?2. The heat of combustion of H2 (g) at constant pressure and 300 K is –280 Kj mol

1. What will be

heat of combustion at constant volume and at 300 K ?

3. What is the limitation of first law of thermodynamics ?

4. What is the value of G when ice and water are in equilibrium ?

5. How does T S determine the spontaneity of process ?

6. A reaction A + B C + D + q is found to have a positive entropy change. The reaction will

be :

(a) Possible at high temperature

(b) Possible at any temperature

(c) Not possible at any temperature

(d) Possible only at low temperature

7. The enthalpies of all elements in their standard states are :

(a) Unity (b) Zero

(c) < 0 (d) Different for each element


8. Calculate the heat of combustion of glucose from the following data :

C (graphite) + O (g) CO (g); H = 395.0 K 2 2

H (g) + ½ O (g) H O (g); H = – 269.4K 2 2 2

6C (graphite) + 6H (g) + O (g) C H O (s); H = 1169.8 K 2 2 6 12 6

9. A swimmer coming out from a pool is covered with a film of water weighing about 18g. How

much heat must be supplied to evaporate this water at 298 K ? Calculate the internal energy of

vaporisation at 100°C. vap H for water at 373 = 40.66 Kjmol1.

10. The equilibrium constant for a reaction is 10. What will be the value of G ?

R = 8.314 JK mol , T = 300 K 1 1

11. For oxidation of iron :

4Fe(s) + 3O (g) 2Fe O (s)2 2 3

Entropy change is 549 JK mol 1 1 at 298 K. Inspite of n of this reaction, why is the reaction

spontaneous ? H for this reaction is = 1648 103 J/mol)

12. For the reaction at 298 K

2A + B C



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H = 400 KJ mol 1 and S = 0.2 KJ K mol1 1

At what temperature will the reaction become spontaneous considering H and S to be constant

over the temperature range ?

13. Acetic acid and hydrochloric acid react with KOH solution. The enthalpy of neutralization of

acetic acid is

558 1. kJ mol . While of hydrochloric acid is acid is

573 1. Kj mol . Why ?

14. For the equilibrium, PCl (g) = PCl (g) + Cl (g)5 3 2 at 298 K.

K = 1.8 10 7 Calculate G for the reaction (R = 8.31 JK mol1 1 )


15. (a) Define standard enthalpy of formation. Explain why the enthalpy changes for the reaction

given below are not enthalpies of formation of CaCO3 and HBr.

(i) CaO (s) + CO (g) CaCO (s)2 3 r 1H = 178.3 KJmol

(ii) H (g) + Br (g) 2HBr (g)2 2

r

1

H = 72.8 KJmol

(b) Calculate the standard enthalpy change( H )r and standard internal energy change (r U )

for the following reaction at 300K ?

OF (g) + H O (g) O (g) + 2HF (g)2 2 2

Standard enthalpy of formation ( H )f of various species are given as below :

f 2 2H KJ mol; OF (g) = 23.0, H O (g) = 241.8; HF (g) = 268.8; / HF (g) = – 268.8;

R = 8.31 J K mol 1 1)

16. (a) State first law of thermodynamics. Heat (q) and work done (w) individually are not state

functions but their sum is always a state function. Explain why ?

(b) Use the bond enthalpies listed in the table given below to determine the enthalpy of

reaction :

H C H (g) + 2 O = O (g) O = C = O (g) + 2H O H (g)

Bond enthalpy ( H ) KJ mol; / / C=O=741; C–H = 414, H—O = 464, O = O = 498.

17. (a) Derive the mathematical expression for 1st law of the thermodynamics.

(b) q and w are not state functions but their sum is state function why.

(c) Calculate the r H for the reaction :

H (g) + Br (g) 2HBr (g)2 2

Bond enthalpy are given as under :H H KJ mol; Br Br 192 KJ mol; H Br 3 436 / / 68 KJ mol /

18. (a) Define the following terms :

(i) Bond enthalpy (ii) Zeroth law of thermodynamics

(b) Estimate change in enthalpy H for the following reaction :

C H (g) + 5O (g) 3CO (g) + 4H O (g)3 8 2 2 2



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The bond enthalpies of C—C; C—H; C=O; O=I and O—H are 347, 414, 741, 498 and 464

KJ/mol respectively.

19. (a) Define enthalpy of neutralization. The enthalpy of neutralisation of strong acid and strong

base is constant why ?

(b) Calculate the enthalpy of formation of acetic acid if it’s enthalpy of combustion is –867KJ/mol.

The enthalpy of formation of CO2 (g) and H O2 (l) are – 393.5 Kj/mol and – 285.9 Kj/mol

respectively.

(c) What is the basis of Hess’s law ?



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Chapter 7 : Equilibrium


1. In which of the following solvents, AgBr has the maximum solubility ?(a) 103 M NaBr (b) 10 3

(c) Pure water (d) 10 3 M HBr

2. The strongest conjugate base is :

(a) NO3 (b) Cl (c) SO4

2 (d) CH COO3

3. Why does the boiling point decrease at high altitudes ?

4. State Le chatellier’s principle.

5. If Q < K c c in which direction the reaction will proceed ?

6.

For the following reaction what is the effect of increasing concentration of CO2 on the directionof reaction :

CaCO (s) CaO (s) + CO (s)3 2

7. State Henry’s law and give an example.


8. What volume of 0.10 M sodium formate solution should be added to 50 ml of 0.05 M formic acid

to produce a buffer solution of pH 4.0 ? pK a for formic acid is 3.80.

9. For reaction N (g) 3H (g) 2NH (g)2 2 3

K = 3.6 10 p

2 at 500K. Calculate K

c for R = 0.083 L k mol1 1.

What is the relationship between pKa and pkb ?

10. For hypothetical reaction, K is given as :

A B, K = 21B C, K =2 4

C D, K =3 3

What will be the value of K for A D ?

11. The solubility of CO2 in water decreases with the increases in the temperature. Explain.

12. The following can act as both Bronsted acid and Bronsted base. Write the formula of Bronsted

acid and Bronsted base :

(a) HCO3 (b) H PO2 4

(c) NH3


13. Which is more soluble ? Given the values of k sp of two sparingly soluble salts Ni(OH)2 and

AgCN are 20 10 15. and 6 10 17 respectively ? Explain.



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Chapter 8 : Redox Reaction


1. Calculate the oxidation num ber of P in PO43

.2. How will you identify cathode and anode in electrochemical cell ?

3. Why do we need salt bridge ?

4. Write the name of the cell in which chemical energy is converted into chemical energy.

5. What is meant by electrode potential ?

6. Oxidation number of C in CH COOH3 is :

(a) 4 (b) 3 (c) 2 (d) 1

7. Which one is correct :

(a) Oxidation is addition of Hydrogen

(b) Oxidation is gain of electron(c) Oxidation is addition of electropositive part

(d) Oxidation is addition of electronegative part Short answer type questions:

8. Which of the following is best reducing reagent and why ?

Li, Cu, Br 2, F2, H2, K

9. While sulphur dioxide and hydrogen peroxide can act as oxidising as well as reducing agents in

their reactions. While ozone and nitric acid act only as oxidants. Why ?

10. The compound AgF2 is unstable compound.However, if formed the compound acts as a very

strong oxidising agent why ?

11. Consider the following cell notation :Al (s) | Al (aq) || Ni (aq) | Ni(s)3+ 2+

What substance act as anode ? Which of them act as cathode ? Write the net ionic equation for the

cell reaction.

12. The standard reduction potentials of Zn Mg2+, 2+ and Na + are 0 76V, 2 37V and 2 71V

respectively. Which of the following is the strong oxidising agent ?

13. (a) Give one use of heavy water in nuclear reactor ?

(b) Write down balanced chemical equations of the reaction of conc. nitric acid with (i) copper

(ii) Iodine.

14. Consider the reactions :

2S O (aq) + I (s) S O + 2I (aq)2 32

2 4 62

S O (aq) + 2Br (l) + 5H O 2SO (aq) + 4Br (aq) + 10H (a2 32

2 2 42 +

q)


15. Balance the following reactions in basic medium by ion electron method and oxidation number

methods and identify the oxidising agent and the reducing agent :



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(a) P (s) + OH (aq) PH (g) + HPO (aq)4 3 2

(b) N H (l) + ClO (aq) NO (g) + Cl (g)2 4 3

(c) Cl O (g) + H O (aq) ClO (aq) + O (g) + H2 7 2 2 2 2+

16. (a) Why does the following reaction occur ?

XeO + 2F (aq) + 6H (aq) XeO (g) + F (g) + 3H O (l)64 +

3 2 2

What conclusion about the compound Na XeO4 6 ( of which XeO64 is a part ) can be drawn from

the reaction ?

(b) Balance the following equations by ion electron method :

(i) MnO (aq) + Br (aq) + H Mn (aq) + Br (aq) + H O4+ 2+

2 2

(ii) Cl + OH ClO + Cl + H O2 3 2

17. (a) Balance the following equation by oxidation number method :

Bi + NO Bi + NO3 3+ 2 (acidic medium)

(b) A cell is prepared by dipping a copper rod in 1 MCuSO4 solution and a nickel rod in 1 M

NiSO4 solution. The standard reduction potentials of copper and nickel are 0 34 V and

–0.25 V respectively.

(i) Which electrode will work as anode and which as cathode ?

(ii) What will be the cell reaction ?

(iii) How is cell represented ?

(iv) Calculate emf of the cell.

18. How do you account for the following observations?

(a) Though alkaline potassium permanganate and acidic potassium permanganate both are

used as oxidants. Yet in the manufacture of benzoic acid from toluene we use alcoholic

potassium permanganate as an oxidant why? Write a balanced redox equation for the

reaction

(b) When concentrated sulphuric acid is added to an inorganic mixture containing chloride, we

get colourless pungent smelling gas HCl, but if the mixture contains bromide then we get

red vapour of bromine. Why?

19. Predict the product of electrolysis in each of the following :

(a) An aqueous solution of AgNO3with silver electrodes(b) An aqueous solution of AgNO3 with platinum electrodes

(c) A dilute solution of H SO42 with platinum electrodes

(d) An aqueous solution of CuCl2with platinum electrodes

(e) Why, it is not possible to store copper sulphate solution in a zinc vessel ?



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Chapter 9 : Hydrogen


1. The adsorption of hydrogen by palladium is called :(a) Hydrogenation (b) Hydration

(c) Reduction (d) Occlusion

2. Proton is likely to have high :

(a) Hydration enthalpy (b) Electron affinity

(c) Atomic size (d) Atomic mass

3. What is nascent hydrogen?

4. What is hydride gap?

5. Water molecule is bent and not linear in structure. Explain

6. What is the role of liquid hydrogen in space rockets ?7. What is clathrate ?


8. When a small dry piece of sodium metal is thrown in water, it immediately catches fire. What

actually happens ?

9. Describe the industrial use of dihydrogen which depends upon its ability to unite with nitrogen.

10. Explain why water has high melting and boiling points as compared to H S2 .

11. Explain why is hydrogen peroxide stored in coloured/plastic bottles.

12. Give four points in which hydrogen and halogen resemble.13. Explain why water acts as excellent solvent for polar and ionic substances ?

14. How will you concentrate a dilute solution of H O2 2 ?

Long Answer Question

15. What happens when hydrogen peroxide is reacted with :

(a) Lead sulphide

(b) Acidified ferrous sulphate solution

(c) Alkaline potassium ferrocyanide solution.

16. (a) Discuss the bleaching action of H O2 2.(b) Write the important uses of H O2 2.

(c) How is concentration of H O2 2 expressed ? How will you convert the same into strength in

grams per litre ?

17. Discuss the preparation of hydrogen peroxide :

(a) From sodium peroxide

(b) From 2-ethylanthraquinone.



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17. Give a brief account of the different methods for the commercial preparation of H O2 2.

18. Discuss the position of hydrogen in the periodic table. Is the present position of hydrogen

satisfactory ?

19. Give a brief account of different isotopes of hydrogen. Which out of them is radioactive in

nature ?



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Chapter 10 : s-Block Elements


1. Which is a correct statement for CsBr 3 :(a) It is a covalent compound

(b) It contains Cs+ and Br ions

(c) It contains Cs+ and Br ions

(d) It contains Cs+, Br and Bra molecule

2. Least stable amongst the following is :

(a) Li+ (b) B (c) C (d) Be

3. Arrange the following alkali metal ions in decreasing order of their mobility :

Li , Na , K , Rb , Cs

+ + + + +

4. Why beryllium chloride has a 0 dipole moment although Be-Cl bonds are polar ?

5. Washing soda contains ten water molecules as water of crystallization but it is colourless. Why ?

6. Why are lithium salts commonly hydrated and those of the other alkali metal ions usually

anhydrous ?

7. Why is LiF almost insoluble in water whereas LiCl soluble not only in water but also in acetone?


8. Potassium carbonate cannot be prepared by Solvay process. Why ?

9. Why does the salts of alkaline earth metals form stronger hydrates of higher deliquescence thanthat of those formed by the alkali metal salts ?

10. Why Li does not shows photo-electric effect ?

11. Explain why alkali and alkaline earth metals cannot be obtained by chemical reduction methods.

12. Why is Li CO2 3 decomposed at a lower temperature whereas Na CO2 3 at higher temperature ?.

Long answer type questions

13. Explain :

(a) LiCl is more covalent than KC1

(b) In aqueous solution Li+ has lowest mobility.

(c) Explain the basis of the diagonal relationship between lithium and magnesium

14. Comment on each of the following observations :

(a) Lithium forms a nitride directly like magnesium. Give equation involved.

(b) BaO is soluble but BaSO4 is insoluble in water.

15 Explain :

(a) NaOH is a stronger base than LiOH. Explain



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(b) Why are alkali metals kept in paraffin or kerosene ?

(c) Why does lithium show properties uncommon to the rest of the alkali metals ?

16. Give the names and formulae of the compounds indicated by the following statements :

(a) A compound of Ca used in setting fractured bones.

(b) A compound of Mg, S, O and H used as purgative in medicines.(c) A compound of Ca and C used for the production of acetylene.

(d) A compound of Ca, C and N used as fertilizer.

(e) A compound of Ca and H which upon reacting with water gives Ha.

17. (a) The crystalline salts of alkaline earth metals contain more water of crystallization

molecules than the corresponding alkali metal salts. Why ?

(b) Why does the solubility of alkaline earth metal carbonates and sulphates in water decrease

down the group ?

18. (a) Discuss the various reactions that occur in the Solvay process.

(b) Draw the structure of beryllium chloride in vapour and solid state

19. What happens when :

(a) Sodium metal is dropped in water

(b) Sodium peroxide dissolves in water

(c) Magnesium is burnt in air

(d) Quick lime is heated with silica

(e) Chlorine reacts with slaked lime



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Chapter 11 : p-Block Elements


1. Which of the following is not hydrolysed by water and why(a) BF3 (b) BCl3 (c) BBr 3 (d) SiCl4

2. Which of the following oxides is not an acidic oxide ?

(a) CO (b) SiO2 (c) CO2 (d) Al O2 33. What is borax bead test ?

4. What type of hybridization of central atom is possible in [GeCl ]62 ?

5. Diamond is a covalent compound, yet it has high melting point. Why ?

6. What is popularly known as Bucky balls ?

7. Suggest reason as to why CO is poisonous ?


8. Like CO, its analog SiO is not stable. Why ?

9. PbO2 is a stronger oxidizing agent than SnO2. Explain.

10. Explain why BF3 exists but BH3 does not ?

11. Explain why is there a phenomenal decrease in ionization enthalpy from carbon to silicon?

12. Write the resonating structures of CO2.

13. On heating a borate with ethyl alcohol and cone. H SO2 4 , a green edged flame is noticed. How

will you account for it ?

14. What happens when (write chemical equations) :

(a) Orthoboric acid is heated gradually.

(b) B O2 3 and ferric oxide are heated in oxidising flame.

(c) A mixture of boron and sodium hydroxide is fused at high temperature.

(d) A mixture of B O2 3 and Mg powder is fused and the product is boiled with dilute.


15. Explain the following :

(a) PbCl4 is less stable than PbCl2(b) CO2 is a gas while SiO2 is a solid

(c) AlCl3 exists as a dimer but BCl3does not.

(d) [SiF ]62 is known but [SiCl ]6

2 is not.

(e) Boron forms no compounds in uni positive state but thallium in uni positive state is quite

stable. Why ?



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16. Write the balanced chemical equation for :

(a) SiO2 + HF

(b) AlCl3 + NaOH

(c) Na B O2 4 7 + H O2

(d) BF3 + LiAlH4(e) BH3 + CO

17. When metal X is treated with sodium hydroxide, a white precipitate A is obtained, which is

soluble in excess of NaOH to give soluble complex B. Compound A is soluble in dilute HC1 to

form compound C. the compound A when heated strongly gives D which is used to extract metal.

Identify X,A,B,C,D. Write suitable equations to support their identities.

18. (a) How does BF3 reacts with :

(i) Ammonia

(ii) Lithium hydride

(iii) Water (b) How do you account for :

(i) Anhydrous AlCl3 is covalent but hydrated AlCl3 is ionic

(ii) BBr 3 is stronger lewis acid than BF3

(iii) Boron does not form B3+ ions

19. A certain salt X, gives the following results :

(a) Its aqueous solution is alkaline to litmus

(b) It swells up to a glossy material Y on strong heating.

(c) When cone. H SO2 4

is added to a hot solution of X, white crystals of an acid Z separates

out.

Write equations for all the above reactions and identify X Y and Z.



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Chapter 12 : Some Basic Principle and Techniques of

Organic Chemistry


1. What is the type of hybridisation of C in :

(i) CH Cl3 (ii) (CH ) CO3 22. How will you test presence of nitrogen in organic compounds ?

3. Give IUPAC name of the following compound :

CH CH

Br

C

O

CH

CH

CH3 3 | || |

3

4.

Write the structural formula of 4-chloro-2-pentene.5. Write IUPAC name of :

(a) CH CH

CHO

CH CH

CH

COOH3 2 | |

3

(b) CH C

O

CH CH CHO3 2 2 ||

6. How will you separate mixture of naphthalene and NaCl :

(a) Filtration (b) Crystallisation

(c) Chromatography (d) Sublimation

7. Match the following :

List (A) List (B)

(a) Molecular mass of benzene (i) Volumetric method

(b) Molecular mass of benzoic acid method (ii) Victor Meyer’s

(c) Percentage of element(other than C and H) (iii) Carius method In gammaxene

(d) Percentage of element(other than C and H) (iv) Kjeldahl’s method In aniline

Short Answer Question

8. An organic compound contains 69% carbon, 4.8% hydrogen, the remainder is oxygen. Calculate

the masses of CO2 and H O2 produced when 0.20g of this substance is subjected to complete

combustion.

9. Write the structural formulae of the following :

(a) o-ethyl anisole (b) p-nitro aniline

(c) 2, 3-dibromo-1-phenyl pentane (d) 4-ethyl-1-fluoro-2-nitrobenzene



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10. (a) Write geometrical isomers of but-2-ene.

(b) Write open chain isomers of C H O3 6 and cyclic isomers.

11. Discuss the chemistry of Lassaigne’s test.

12. How will you proceed to detect the presence of halogens in organic compounds by Lassaigne’s

test? Why is it necessary to boil the Lassaigne’s extract with conc. HNO 3 ?Long Answer Questions

13. (a) What do you understand by :

(i) Homolytic fission

(ii) Heterolytic fission

(b) What are carbanions ? Give an example.

(c) Distinguish between position and functional isomerism with one example.

14. (a) 0.5 gram of an organic compound produced ammonia in Kjeldahl’s method which was

absorb in 50 cm

3

of 1 N H SO2 4 .The residual acid required 60 cm

3

of semi normal KOHsolution. What is the % of N in the compound ?

(b) 0.395 g of an organic compound by Carius method for the estimation of Sulphur gave

0.582 g of BaSO4 . Calculate the % of S in the compound.

15. (a) A liquid organic compound ‘A’ having 92.30% carbon and 7.7% hydrogen decolourises

KMnO4and on ozonolysis gives methanal and other compound ‘B’. The vapour density of

A was found to be 52. On treatment with suitable catalyst ‘A’ gave a high molecular weight

solid products ‘C’ having the same empirical formula as that of compound ‘A’. Compound

‘C’ could be used in making toys and household goods. Identify A, Band C, explain the

reactions.(b) 0.45 g of an organic compound gave on combustion 0.792 g of CO2 and 0.324 g of water

0.24 g of same substances was Kjeldahlised and the NH3 formed was absorbed in 50.0 cm3

of N/4 H SO2 4 . The excess acid required 77.0 cm3 of N/10 NaOH for complete

neutralisation. Calculate the empirical formulae of the compound.

16. (a) Write short notes on following giving one example of each :

(i) Free radical substitution reaction (ii) Addition reaction

(iii) Elimination reaction (iv) Resonance

(v) Hyper conjugation

17. (a) In a Dumas nitrogen estimation method, 0.30 g of an organic compound gave 50 cm3 of N 2

collected at 300 K and 715 mm Hg pressure. Calculate the percentage composition of

nitrogen in the compound. (Vapour pressure of water at 300 K is 15 mm Hg)

(b)) An organic compound has only one bromine atom in its molecule. 1.57 g of the compound

gave 1.88 g of AgBr in the Carius method. Find the molecular weight of the compound

[At.wt. of Ag = 108 and Br = 80]



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(c) An organic compound with the molecular formula C H O Px 15 3 has 18.6% of phosphorus.

Calculate the Value of ‘x’ [At.wt. of P = 31]

18. (a) What are nucleophilic substitution reactions ? Give one example.

(b) What are electrohiles ? Explain electrophilic substitution reaction with the help of an

example.(c) What are rearrangement reactions?? Illustrate with the help of an example giving its

mechanism.

19. Classify the following in one of the reaction type :

(a) CH CH Br + CN CH CH CN + Br 3 2 3 2

(b) CH CH OH CH = CH + H O3 2 3 2 2

(c) CH CH = CH + HBr CH CH(Br) CH3 2 3 3

(d) CH COONa + NaOH (CaO) CH + Na CO3 4 2 3 /

(e) CH CH CH CH CH CH

CH

CH3 2 2 3 3 3

|3



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Chapter 13 : Hydrocarbons


1. Butane and 2-methyl propane are which type of isomers2. Why eclipsed and staggered forms of ethane cannot be isolated at room temperature ?

3. What are conformations ?

4. Which of the following exhibits geometrical isomerism :

(i) 2-methyl propene (ii) 1-butene

(iii) 2-butene (iv) 2,3 dibromo-2-butene.

5. Alcohols gives carboxylic acids on :

(i) Reduction (ii) Oxidation

(iii) Dehydration (iv) Hydrogenation

6. Convert benzene to Acetophenone.7. How will you distinguish between but-1-yne and but-2-yne ?


8. What is meant by :

(i) Delocalisation (ii) Resonance energy

9. Explain the following with examples :

(i) Wurtz reaction (ii) Kolbe’s electrolytic method

(iii) Hydrgenation

10. What happens when ?(i) 2-Propanol is heated with alumina at 630 K

(ii) Benzene is treated with mixture of cone. H SO2 4 and cone. HNO3.

(iii) Ethene is treated with cold alkaline KMnO4 solution.

11. An alkene ‘A’ on ozonolysis gives a mixture of ethanal and pentan-3-one. Write structure and

IUPAC name of ‘A’.

12. Addition of HBr to propene yields -2-bromopropane, while in the presence of benzoyl peroxide,

the same reactions yields-1- bromopropane.Explain and give mechanism.

13. Out of benzene, m-dinitrobenzene and toluene which will undergo nitration most easily and

why?

14. What happens when :

(a) Sodium acetate is heated with soda lime.

(b) Acetylene is passed through ammonical silver nitrate solution.


15. (a) Carry out the following conversions : (i) Acetylene to propyne (ii) n-Hexane to benzene



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(b) Draw main conformations of n-butane obtained by rotation around C-2 aad C-3. Also give

the names of these conformations. Which of these conformations is most stable and which

is the least stable and why ?

(c) How will you bring about the following conversions :

(i) Ethene to bromoethane (ii) Toluene to Benzene16. (a) What type of isomerism is shown by methoxy methane and ethanol ?

(b) Complete the following reactions :

(i) CH CH CH Br 3 2 2Alc. KOH

(ii) C H + CH COCl6 6 3Anhy. AlCl3

(iii) CH CH = CH3 2(ii) H O Zn

(i) O

2

3

/

17. Complete the following reactions :

(a) Isopropyl bromide

Alc. KOH HBr/peroxideA B

(b) n-propyl alcohol

443 K

conc. H SO KMnO OH2 4 4C D

/

(c) 1,1,2,2-tetrachloroethane

Zn, alcohol

tube

RedhotE F

(d) Acetylene H , Pd BaSO

NaNHH

O

CH CH Br

2 4

2 3 2

G H/ 3

(e) Propyne Quinoline H O Zn

I J2 /

18. (a) Give a chemical test to distinguish between but-1-yne and but-2-yne.

(b) An alkene C H4 6 upon ozonolysis gives 2 moles of formaldehyde and 1-mole of glyoxal.

What is the name of alkene ?

(c) Explain why benzene not undergo addition reaction ?

19. (a) Why terminal alkynes are acidic in nature ?

(b) Complete the following :

(i) CH CH = CH + HBr 3 2Peroxide

(ii) CH CH CH

Cl

CH + KOH (alc)3 2 3 |

(c) How will you convert :

(i) Sodium acetate to ethane

(ii) 2-chloropropane to 1-bromopropane



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Chapter 14 : Environmental chemistry


1. Why is CO more toxic than CO2 ?2. What is PAN ?

3. What is meant by photochemical smog ?

4. Name two gases which are responsible for green house effect.

5. Which gas cause of Bhopal gas tragedy? Give its formula.

6. Which acid is not present in acid rain :

(a) HNO3 (b) Acetic acid

(c) Sulphuric acid (d) Carbonic acid

7. Full form of BOD :

(i) Biochemical oxygen demand (ii) Basic oxygen demand(iii) Boron oxygen demand (iv) Biological oxygen demand

8. “Write the name of gas produced in Mathura refineries which can damage the great historical

monument Tajmahal”.

9. What types of radiations are absorbed by CO2 in the atmosphere ?


10. Carbon monoxide gas is more dangerous than carbon dioxide gas. Why ?

11. What is smog ? How is classical smog different from photochemical smog ?

12. What do you mean by green chemistry? How will it help in decreasing environmental pollution ?13. A large number offish are suddenly found floating dead on a lake. There is no evidence of toxic

dumping but you find an abundance of phytoplankton. Suggest a reason for the fish kill.

14. Explain giving reasons: “The presence of carbon monoxide reduces the amount of haemoglobin

available in the blood for carrying oxygen to the body cells.”


15. (i) Define eutrophication and pneumoconiosis.

(ii) Write difference between photochemical and classical smog.

16. (i) What is the cause of Acid rain ? How is it harmful to the environment ?(ii) What do you understand by greenhouse effect ? What are the major greenhouse gases ?

17. Describe the following in brief :

(i) Ozone depletion over Antarctica (do not write reactions)

(ii) BOD and COD

18. (i) How can domestic waste be used as manure ?



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(ii) Discuss the importance of dissolved oxygen in water. What processes are generally

responsible of de oxygenation of water ?

19. For your agricultural field or garden you have developed a compost producing pit. Discuss the

process in the light of bad odour, flies and recycling of waste for a good produce.

Documents

Chemistry Class XI