jr chemistry - chapter wise important questions - part 1

1. ATOMIC STRUCTURE

Long answer type questions

01*. What are the postulates of Bohr’s model of hydrogen atom? Discuss the importance of this model to

explain various series of line spectra in hydrogen atom.

02. Explain the salient features of quantum mechanical model of an atom.

03. What are the evidence in favour of dual behavior of electron?

04*. How are the quantum numbers n. 1 and m1 arrived at ? Explain the significance of these quantum

numbers.

05. Explain the dueal behavior of matter. Discuss its significance to microscopic par ticles like electrons.

Short answer type questions

01*. Show that the circumference of the Bohr ortbit for the hydrogen atom is an integral multiple of the de

Broglie wavelength associated with the electron revolving around the orbit.

02. What is a nodal plane? How many nodal planes are possible for 2p- and 3d- orbitals?

03*. Explain the difference between emission and absorption spectra.

04. Explain the difference between orbit and orbital.

05*. Explain photoelectric effect.

06*. Explain briefly the Plank’s quantum theory.

07. What is the wavelenth of light emitted when the electron in a hydrogen atom undergoes transition from an

energy level with n = 5 to an energy level with n = 3 ?

08. If the position of the electron is measured within an accurancy of ± 0.002 nm. Calculate the uncertainity

in the momentum of the electron.

2. CLASSIFICATION OF ELEMENTS AND PERIODICITY IN PROPERTIES


01*. Write an essay on s,p,d and f block elements.

02*. What is a periodic property? How the following properties vary in a group and in a period?

Explain a) IE , b) EN

03*. Define IE1 and IE

2. Why is IE

2 > IE

1 for a given atom? Discuss the factors that effect IE of an element.

04. How do the following properties change in group-1 and in the third period? Explain with example.

a) Atomic radius b) IE c) EA d) Nature of oxides


01. Give the outer orbit general electronic configuration of

a) Noble gases b) Representative elements

c) Transition elements d) Inner transition elements

02*. Give any four characteristic properties of transition elements.

03. What is the basic deference between the electron gain enthalpy and electropositivity?

04. What is valency of element? How does it vary with respect to hydrogen in the third period?

05*. What is diagonal relationship? Give a pair of elements having diagonal relationship.Why do they show

this relation?

SUB: CHEMISTRY

INTER FIRST YEAR CHEMISTRYChapter wise imp. Questions

Designed by SriGayatri

Jr.Inter Chemistry 2

06*. What is lanthanide contraction? What are its consequences?

07. Write a note on a) Atomic radius b) Metalic radius c) Covalent radius

3. CHEMICAL BONDING AND MOLECULAR STRUCTURE

Long answer type questions :

1. Explain the formation of Ionic Bond with a suitable example.

2. Explain the factors favourable for the formation of Ionic Compounds.

3. Give an account of VSEPR Theory, and its applications.

4*. What do you understand by Hybridisation ? Explain different types of hybridization involvings and p

orbitals.

5*. Give the Molecular Orbital Energy diagram of (a) N2 and (b) O

2. Calculate the respective bond order.

Write the magnetic nature of N2 and O

2 molecules.

Short answer type questions :

01. State Fajan’s rules and give suitable examples.

02. Write the resonance structures for NO2, and

3NO-

03*. Define Dipole moment. Write its applications.

04. Explain why BeF2, molecule has zero dipole moment although the Be-F bonds are polar.

05. Even though nitrogen in ammonia is in sp3 hybridization, the bond angle deviate from1090

06*. Explain the hybridization involved in PCl5 molecule.

07*. Explain the hybridization involved in SF6 molecule.

08*. Explain the formation of Coordinate Covalent bond with one example.

09*. What is Hydrogen bond? Explain the different types of Hydrogen bonds with ex amples.

10**. What is meant by the term Bond order? Calculate the bond orders in the following

a) 2N b) 2O c) 2O+ d)

2O-

11. How do you explain the geometry of the molecules on the basis of Valence bond Theory ?

12. How do you predict the shapes of the following molecules making use of VSEPR Theory ?

4. STATES OF MATTER : GASES AND LIQUIDS


1. Derive the van der waals equation of state. Explain the importance of vander waal’s gas equation.


1*. Derive Ideal gas equation.

2*. State and explain Graham’s law Diffusion.

3*. State and explain Dalton’s law of Partial pressures.

4*. Deduce (a) Boyle’s law and (b) Charle’s law from kinetic gas equation.

5*. Deduce (a) Graham’s law and (b) Daltons law from kinetic gas equation.

6*. Derive an expression for kinetic Energy of gas molecules.

7*. Define (a) rms (b) average and (c) most probable speeds of gas molecules. Give their interrela tionship.

8*. Write the postulates of Kinetic Molecular Theory of Gases.

9. Explain the principle underlying the liquefaction of Gases.

10. What is surface tension of liquids ? Explain the effect of temperature on the surface tension of liquids.


5. STOICHIOMETRY


1*. Chemical analysis of a carbon compounds gave the folowing percentage copmposition by weight of the

elements present, carbon = 10.06%, hydrogen = 0.84% chlorine = 89.10%,Calculate the empirical

formula of the compound.

2*. A carbon compound on analysis gave the following percentage composistion, carbon 14.5%, hydrogen

1.8%, chlorine 64.46%, oxygen 19.24%. Chlorine 64.46%, oxygen 19.24% Calculate the empirical


3*. Calculate the empirical formula of a compound having percentage composition postassium (k) = 26.57.

chromium (cr) = 35.36; oxygen (O) 38.07. (Give the atomic weights of K, cr and O as 39; 52 and 16

respectively)

4*. A carbon compound contains 12.8% carbon, 2.1% hydrogen, 85.1% bromine. The molecular weight

of the compound is 187.9. Calculate the molecular formula.

5*. Balance the following redox reactions by ion - electron method :

(a) ( )4 2( ) ( )MnO aq I aq MnO S− −+ → (in basic medium)

(b) 4 ( )aqMnO− 2

2 ( ) ( ) 4 ( )g a q a qS O M n H S O+ −+ → + (in acidic solution)

(c) 2 3

2 2 2( ) ( ) ( )H O aq Fe aq Fe aq H O+ ++ → + (in acidic solution)

(d) 2 3 2

2 7 2 4( ) ( )Cr O SO g Cr aq SO− + −+ → + (in acidic solution)

6*. Balancce the following equations in basic medium by ion-electron method and oxidation number meth-

ods and identify the oxidising agent and the redusing agent.

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

(b) 2 4( ) 3 ( ) ( ) ( )l aq g gN H ClO NO Cl− −+ → +

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

− ++ → + +

6. THERMODYNAMICS


1. Explain the spontaneity of a reaction interms of enthalpy change, entropy change and Gibb’s energy

change.


1. What are open, closed and isolated systems? Give one example for each.

2. Define the state function and state variables. Give examples

3. “Internal energy is a state function”. Expalin.

4. Derive the equation for rev

W ’ in isothermal reversible process.

5. Explain the state function ‘enthalpy,H’. What is the relationship between U∆ and H∆ ?

6. Explain extensive and intensive properties.

7. Define heat capacity. What are P

C and V

C ? Show that P V

C C R- = .

8*. State and explain the Hess’s law of constant Heat summation.

9. Explain the spontaneity of a process?


10*. What is entropy? Explain with examples.

11*. State the second law of thermodynamics and explain it.

12*. State the third law of thermodynamics. What do you understand by it ?

13. Explain spontaneity of a process in terms of Gibbs energy.

7. CHEMICAL AND IONIC EQUILIBRIUM


1*. What is Lechatlier’s principle? Discuss briefly the factor s which can influence the equilibrium.

2*. Discuss the application of Lechatelier’s principle for the industrial synthesis of Ammonia and sulphur

trioxide.

3*. Explain the concept of Bronsted acids and Bronsted bases, Illustrate the answer with suitable examples.

4*. Explain Lewis acid-base theory with suitable example. Classify the following species into Lewis acids

and Lewis bases and show these act as Lews acid / base.

a) OH- b) -F c) H + d) 3Bcl

5*. Define pH.. what is buffer solution ? Derive Henderson - Hasselbalch equation for Calculating the pH

of an acid buffer solution.

6*. Explain the term “Hydrolysis of salts” with examples.Discuss the pH of the follow ing types of salt

solutions. i) Salts of weak and strong base. ii) Salts of strong acid and weak base.


1. Write expression for the equilibrium constant, c

K , for each of the following reactions.

(i) 2( )2 2g g g

NOCI NO CI→ +←

2. How does the value of equilibrium constant predict the extent of reaction ?

3. Explian the terms :

i) extent of ionization and on what factors it depends.ii) dissociation iii) ionization

4*. Explain the Arrhenius concept of acids and bases.

5*. What is conjugate acid-base pair ? Illustrate with examples.

6. Write equation that shows 2 4H PO

- acting both as an acid and as a base.

7. Write the conjugate acid and conjugate base of each of the following :

a) OH - b) 2H O (c) 3HCO

- (d) 2 2H O

8*. Define ionic product of water. What is its value at room temperature ?

9*. What is common ion effect ? Illustrate.

10*. Define solubility product ? Write solubility product expressions for the following :

i) 2 2 7Ag Cr O ii) 3 4 4( )Zr PO

11*. Give the classification of salts. What type of salts undergo hydrolysis ?

12*. Aqueous solution of 4NH CI is acidic. Explain.

13*. Aqueous solution of 3CH COONa is basic explain.Long answer questions

8. HYDROGEN AND ITS COMPOUNDS


1*. Explain the terms hard water and soft water. Write a note on the

i) ion-exchange method and ii) Calgon method for the removal of hardness of water


2*. Write the chemical reaction to justify that hydrogen peroxide can function as on oxidizing as well as

reducing agent.

3*. Write a note on heavy water.

4*. Discuss the principle and the method of softening of hard water by synthetic, ion- exchange resins.

5. In how many ways can you express the strength of H2O

2 ? Calculate the strength of 15 volume solution

of H2O

2 in g/l. Express this strength in normality and molarity.

9. THE S-BLOCK ELEMENTS


1. Discuss the preparation of sodium carbonate.

2.* Explain the significance of sodium, postassium, magnesium and calcium in biologi cal fluids.

3*. Write a few lines about cement.


1*. What do you know about Castner-Kellner process? Write the principle involved in it.

2. Write a note on the anomalous behavior of Beryllium.

3*. What is plaster of Paris?Write a short note on it.

4. How do you obtain pure sodium chloride from a crude sample ?

5. When an alkali metal dissolves in liquid ammonia the solution can acquire different

colours. Explain the reasons for this type of colour change.

10. THE P-BLOCK ELEMENTS - GROUP 13


1*. Write any two methods of preperation of diborane. How does it react with

a) Carbon monoxide and b) Ammonia ?


1*. Explain borax bead test with a suitable example.

2*. Explain the structure of diborane.

3. Explain the reactions of aluminium with acids.

4*. Give two methods of preparation of diborane.

5*. How does diborane react with a) 2H O b) CO c) 3 3( )N CH ?

11. THE P-BLOCK ELEMENTS - GROUP 14


1*. Write note on the allotropy of carbon.


1*. Explain the difference in properties of diamond and graphite on the basis of their structure.

2*. Write a short note Zeolites.

3. Write a short note silicates.

4*. What are silicones? How are they obtained?

5. Write a short notes on fullerene.

6*. Why is diamond hard?

7. What happens when a) 2CO is passed through slaked lime b) 2CaC is heated with 2N .


VERY SHORT ANSWER TYPE QUESTIONS

1. ATOMIC STRUCTURE

1. Calculate the charge of one mole of electrons.

Ans : One electron has charge - 191.602 10−× coloumbs

one mole of electrons has charge 23 196.023 10 1.602 10−− × × ×

49.648846 10 96488.5= × = coloumbs

2. Calculate the mass of one mole of electrons.

Ans : Mass of electron = ( )31 289.1 10 kg or 9.1 10 gms.− −× ×

One moel of electrons has mass 23 316.023 10 9.1 10−× × ×

8 754.8 10 5.48 10 kg− −= × = × .

3. How many p electrons are present in sulphur atom ?

Ans : Sulphur has electronic configuration 2 2 6 2 41s 2s 2p 3s 3p

∴ Sulphur has 10 'p' electrons.

4. What are the values of principal quantum number (n) and azimuthal quantum number (l) for a 3d

electron ?

Ans : For a 3d - electron principal quantum number (n) = 3 and Azimuthal quantum number (l) = 2.

5. What is the completer symbol for the atom with the given atomic number (Z) and atomic mass

(A) ? i) Z = 4, A= 9 : ii) Z = 17, A = 35 : iii) Z = 92, A = 233.

Ans : i ) Z = 4, A = 9 Complete symbol is 9

4 Be

ii) Z = 17,A = 35 Complete symbol is 35

17 Cl

iii) Z = 92,A = 233 Complete symbol is 233

92U

6. What is the frequency of radiation of wavelength 600 nm ?

Ans : Formulae c

υλ

=

8

7

3 10

6 10−

×=

×

15110

2= ×

15 14 10.5 10 5 10 sec−= × = ×

9 7600 600 10 6 10nm m mλ − −= = × = ×

83 10 / secC m= ×

7. What is Zeeman effect ?

Ans : The splitting up of spectral lines in presence of strong exetrnal magnetic field is called as Zeeman effect .

8. What is Stark effect ?

Ans : The splitting of spectral lines in presence of strong electric field is called as Stark effect.


9. Electrons are emitted with zero velocity from a metal surface when it is exposed to radition of

wavelength 4000 0

A . What is the threshold frequency (v0) ?

Ans : Formulae

2

0

1

2hv hv mv= + ⇒ ( )

2

0

1

2hv hv m o= +

0hv hv=

0

λ = 4000 Α

3 10 74 10 10 4 10 m.− −= × × = ×

0υ = , C = 3 x 108 m/sec.

0v v⇒ =

815

7

3 10 310

4 10 4

cv

λ −

×∴ = = = ×

× 150.75 10= × 14 17.5 10 sec .−= ×

10. Which of the following orbitals are possible ? 2s,1p,3f,2p.

Ans : 2s,2p orbitals are possible among 2s,1p,3f,2p and 1p, 3f orbitals ar not possible.

11. How many electrons in an atom may have n = 4 and m = +1/2 ?

Ans : For n = l values are 0,1,2,3

0l = → s contains are electron with ms = +1/2

1l = →p contains 3 electron with ms = +1/2

2l = → d contains 5 electron with ms= +1/2

3l = → f contains 7 electron with ms= +1/2

∴ Total no.of electrons with ms= +1/2 for n = 4

1 3 5 7 16.= + + + =

2. CLASSIFICATION OF ELEMENTS

AND PERIODICITY IN PROPERTIES

1. What is the difference in the approach between the Mendeleev's periodic law and the modern

periodic law ?

Ans : → According to Mendeleev the physical and chemical properties of elements are periodic functions of

their atomic weights.

→According to modern periodic law physical and chemical properties of elements are periodic functions

of their atomic numbers.

2. In terms of period and group, where would you locate the element with Z = 114 ?

Ans : Element Z = 114 is present in 17th period and IVA group (Group - 14)

3. Write the atomic number of the element, present in the third period and seventeenth group of

the periodic table.

Ans : The Element present in 3rd period and Group - 17 (VIIA group) is chlorine (Cl) It's atomic

number is 17.

4. Which element do you think would have been named by

a) Lawrence Berkeley Laboratory


b) Seaborg's group

Ans : a) Lawrence Berkeley (Laboratory-Lanthanide)

b) Seaborg's group - Actinide (Transuranic element )

5. What are representative elements ? Give their valence shell configuration.

Ans : →Representative elements are s and p-block elements except zero group.

→ These have general electronic configuration ns1-2 np1-5

6. Give the outer shells configuration of d-block and f-block elements.

Ans : → The outer shell electronic configuration of d-block- elements is ns1-2 (n-1)d1-10

→ The outer shell electronic configuration of f-block - elements is ns1-2 (n-1)d0 (or) 1 (n-2) f1-14

7. Name the anomalous pairs of elements in the Mendaleev's periodic table.

Ans : In Mendelleev's periodic table anamalous pairs are the elements whose atomic weights

increasing order is reversed.

Eg : 1) Te I

127.6 126.92)

Co Ni

58.93 58.69

8. Among -3 -2 - + +2 +3N ,O ,F ,Na ,Mg and Al

a. What is common in them ? b. Arrange them in the increasing ionic radii.

Ans : Given ions are

-3 -2 - + +2 +3N ,O ,F ,Na ,Mg and Al .

a) The above ions have same number of electrons (All have 10 electrons ). So these are called

iso electronic species.

b) The increasing order of ionic radii among above ions is

+3 +2 + - -2 -3Al <Mg <Na <F <O <N

Reason : In case of iso electronic species as the nuclear charge increases ionic radii decreases.

9. Ionization enthalpy, (IE1) of O is less than that of N - explain.

Ans : → Oxygen has electronic configuration 2 2 41s 2s 2p

→ Nitrogen has electronic configuration 2 2 31s 2s 2p

→ Nitrogen has half filled shell and is stable so more amount of energy is required to

remove an electron , than in oxygen Hence IE, of 'O' is lessthan that of 'N'.

10. Which in each pair of elements has a more negative electron gain enthalpy ?

Ans : a. O or F b. F or Cl

a) ( )328 /KJ Mole

Fluorine−

14 2 43 has more negative electron gain enthalpy than , that of ( )141.2 /KJ Mole

Oxygen

−

14 2 43


b) ( )349 /KJ Mole

Clorine−

14 2 43 has more negative electron gain enthalpy than that of ( )328 /KJ Mole

Fluorine−

14 2 43

11. Write the increasing order of the metallic character among the elements B,Al, Mg and K.

Ans : Given elements are B,Al, Mg and K

The increasing order of metallic character is

B Al Mg K< < <

Group - ( ) ( ) ( ) ( )III III II I

12. Write the correct increasing order of non-metallic character for B,C,N F and Si.

Ans : Given elements are B,C,N,F and Si The increasing order of non-metallic character is

Si < B < C < N < F

Group - (IV) (III) (IV) (V) (VII)

13. Write the correct increasing order of chemical reactivity in terms of oxidizing property for N,

O, F and Cl

Ans : The correct increasing order of chemical reactivity interms of oxidizing property for

N,O,F and Cl is F > O > Cl > N

14. What is the valency possible to Arsenic with respect oxygen and hydrogen ?

Ans : →The valency of Arsenic with respect to hydrogen is '3'

Eg : AsH3

→ The valency of Arsenic with respect to oxygen is 'S'

Eg : As2O

5

15. Name the most electronegative element. Is it also having the highest electron gain enthalpy ?

Why or Why not ?

Ans : The most electronegative element is fluorine (F).

→It doesnot have high electron gain enthalpy.

Reasons : -

→Due to small size

→Due to high inter electronic repulsions.

* Chlorine has high electron gain enthalpy.

16. How does the nature of oxides vary in the third period ?

Ans : In 3rd period from left to right the oxide nature varies from high basic nature to high acidic nature.

→Basic nature gradually decreases and acidic nature gradually increases.

→

( ) ( ) ( ) ( ) ( ) ( ) ( )2 2 3 2 2 5 3 2 7Na O< Mgo Al O SiO PO SO Cl O

basic basic amphoteric acidic acidic acidic acidic

< < < < < <

17. What is lanthanide contraction ? Give one of its consequences.

Ans : Lanthanide contraction : Decrease in the size of the atoms or ions among the lanthanides

is known as lanthanide contraction.Consequence due to lanthanide contraction :

1) Due to lanthanide contraction, the crystall structure and other properties of the leements

become very closely similar.

2) Due to this, it becomes difficult to separate lanthanides from a mixture.


18. Why are zero group elements called noble gases or inert gases ?

Ans : →Zero group elements has general outer electronic configuration 2 6ns np (except for He)

→ These contains stable octet configuration. So these are stable and chemically inert.

Hence these are called inert gases.

→ These elements neither lose nor gain electrons. Hence these are called 'noble gases'.

19. Electron affinity of chlorine is more than that of fluorine - explain.

Ans : Chlorine (-349 KJ / Mole) has high electronegative gain enthalpy than fluorine (-328KJ/mole)

because 'F' has small size and more inter electronic repulsions.

20. Which in each has higher electron affinity ?

a. F or -Cl b. O or O− c. Na+ or F d. F or F

Ans : a) Fluorine has high electron affinity than Cl- ion because of inert gas configuration of Cl- ion.

b) Oxygen has high electron affinity than O- because O has positive of 2nd electron

affinity.

c) F has high electron affinity than Na+ because Na+ has inert gas configuration.

d) F has high electron affinity than -F because -F has inert gas configuration.

21. Arrange the following in order of increasing ionic radius :

a. 3 2, , ,Cl p s F− − − − b. 2 23, , , ,Al Mg Na O F+ + + − − c. 2, ,Na Mg K+ + +

Ans : a) The increasing order of ionic radius is 2 3F Cl S P

− −< < <

b) The increasing order of ionic radius is 3 2 2Al Mg Na F O+ + + −< < < <

c) The increasing order of ionic radius is 2Mg Na K+ + +< <

22. Mg+2 is smaller than O-2 in size, though both have same electronic configuration explain.

Ans : 2Mg + and 2O

− ions are iso electronic species. In case of iso electronic species nuclear charge

increases size of ion decreases. So 2Mg + has small size than 2O

−

23. What is diagonal relation ? Give one pair of elements, that have this relation.

Ans : The similarity in properties of 2nd period first element and 3rd period second element of

the next group is known as diagonal relation. Diagonal relation is due to same polarizing

power.

Be and Al show this relationship

3. CHEMICAL BONDING & MOLECULAR STRUCTURE

1. Write Lewis dot structures for S and S2-

Ans : → Lewis dot structure for 's' is [ ]2 2 6 2 4

16

1 2 2 2 3 3

Z

Electronic configuration s s p s p

=

−

→ Lewis dot structure for s-2 is (or) -2

2 2 2 6 2 61 2 2 2 3 3Electronic configuration of s s s p s p−


2. Predict the change, if any, in hybridization of Al atom in the following reaction

3 4AlCl Cl AlCl− −+ →

Ans : In 3AlCl Aluminium undergoes sp2 hybridisation

In 4AlCl− Aluminium undergoes sp3 hybridisation

3 4

2 3

| |

AlCl Cl AlCl

sp hybridisation sp hybridisation

− −+ →

3. Is there any change in the hybridization of Boron and Nitrogen atoms as a result of the follwing

reaction ? 3 3 3 3BF +NH F BNH→

Ans :a) Ammonia-Boron trifluoride formation ( )3 3H N BF :→

Ammonia molecule contains Nitrogen atom with a lone pair of electrons (in sp3 orbital).

BF3 has 'B' atom with an incomplete octet (with a vacant p

zorbital). Therefore, nitrogen of

ammonia donates sp3 orbital of nitrogen having a lone pair overlaps the vacant 'p' orbital

of Boron. Theequation corresponding to the reaction is written as follows :

[ ]2 3 3 3H N: :

Donor Acceptor Ammonia Boron trifluoride

BF H N BF

−

+ → →

b) Change in hybridised states N and B during [ ]3 3:H N BF→ formation

Boron in 3BF (sp2 hybridization with one vacant unhybrid 'p' orbital. This orbital also

undergoes)hybridization in presence of NH3 so that the hybridised state of 'B' changes

from sp2 to sp3 .This vacant hybrid orbital is bonded to NH3 through dative bond. During

this process there is no change in the hybridized state of Nitrogen in NH3.

4. If A and B are two different atoms when does AB molecule become Covalent ?

Ans : 1) If the difference in electronegativity values between A and B is less than 1.7, then

covalent compound formation is possible

2) If A and B are sharing one or more electron pairs mutually then AB will be a covalent

compound

5. What is meant by localized orbitals ?

Ans : The molecular orbital with bonded electron cloud localised between the two nuclei of

bonded atoms is called localized orbital. (or) The orbitals which are involved in bond

formation arelocalized orbitals.

6. Cl− has greater stability than chlorine atom. Why ?

A. It has [Ar] configuration . Hence it is stable but Cl atom has one electron short to nearest [Ar] configu-

ration. Hence it is unstable.


7. Why argon is not represented by 2Ar ?

A. Since Ar atom has only paired electrons with stable octet configuration, it can’t share its electrons with

another Ar atom and cannot form di atomic molecule.

8. What is sigma bond ? How many sigma and Pi bonds are present in Acetylene, Benzene and

Ethylene molecules ?

A. Chemical bond formed by the head - on overlapping of atomic orbitals is called sigmabond.

( )2 2C H Acetylene 3 and 2σ π

( )2 4C H Ethylene 5 and 1σ π

( )6 6C H Benzene 12 and 3σ π

9. Why the boiling point of 2H O is greater than HF ?

A. In HF molecule, hydrogen bond is present in both vapour state and liquid state where as in water only

in liquid state. Hence for HF, there is no need to break the hydrogen bonds during evaporation. In

2H O molecule, double no.of hydrogen bonds are seen than in HF,

therefore, the boiling point of H2O is greater than HF.

10. What type of hybridisation is seen in phosphorous in 5PCl and sulphur in 6SF molecules ?

A. In 5PCl , ‘P’ shows 3sp d hybridization and in 6SF molecule, ‘s’ shows 3 2sp d hybridisation.

11. What is octet rule ?

A. The demand of an atom to have 8 electrons in its valence shell to get extra stability is known as octet rule.

4. STATES OF MATTER GASES AND LIQUIDS

1. Name the different intermolecular forces experienced by the molecules of a gas.

Ans : The different inter molecular forces experienced by the molecules of a gas are ondon (or) dispersion

forces, Dipole - Dipole forces, Dipole - induced dipole forces, hydrogen bond.

2. State Boyle's law. Give its mathematical expression.

Ans : At constant temperature, the pressure of a given mass (fixed amount) of gas varies inversely with it's

volume. This is Boyle's law.

→ mathematically it can be written as

1p

vα (At constant T and no.of moles (n)

pk

v⇒ = pv k⇒ = (constant)

3. State Charle's law. Give its mathematical expression.

Ans : At constant pressure the volume of a fixed mass of a gas is directly proportional to it's

absolutetemperature. This is charle's law.

→ Mathematically it can be written as

V Tα (At constant P and no.of moles (n)

V kT⇒ = kV

T⇒ = (constant)


4. What are Isotherms ?

Ans : At constant temperature the curves which the relationship between variation of volume of

a given mass of gas and pressure are called isotherms.

5. What is Absolute Temperature ?

Ans : It is also called thermodynamic temperature (or) kelvin temperature. It is a temperature

on the absolute (or) Kelvin scale in which zero at -273.160C.

( )0T= t C+273.16 K∴

6. What are Isobars ?

Ans : The curves (or) graphs that can be drawn at constant pressure are called Isobars.

Eg : Graphs drawn between volume and temperature.

7. What is Absolute Zero ?

Ans : It is the lowest temperature theoritically possible at which volume of perfect gas is zero.

8. State Avogadro's law.

Ans : Equal volumes of all gases under the same conditions of temperature and pressure contains

equal number of molecules

( )v n mathematicallyα

v kn=

9. What are Isochores ?

Ans : At constant volume a line on a graph showing the variation of temperature of a gas with its pressure is

called Isochores.

10. What are STP Conditions ?

Ans : STP means Standard Temperature and Pressure conditions.

→ Standard temperature is 00 C = 273 K

→ Standard pressure is 1 atmosphere = 76 cm = 760 mm. of Hg.

11. What is Gram molar volume ?

Ans : The volume occupied by one gram molecular weight (or) one gram mole of an element

(or) compound in the gaseous state is called gram molar volume.

(or)

→ At STP one mole of any gas occupy 22.4 lit. of volume

This is known as gram molar volume.

12. What is an ideal gas ?

Ans : A gas which obeys gas laws i.e. Boyle's law, Charle's law and Avagadrols law exactly at all

temperatures is called an ideal gas

13. Give the values of gas constant in different units.

Ans : R = 0.0821 lit. atm. k-1 mol1 1

8.314 .J K mol− −=

( ) 1 11.987 2 .or cal K mol− −=

7 -1 -1=8.314×10 ergs.k mol

14. How are the density and molar mass of a gas related ?

Ans : Pv = n RT ⇒ Pvw

RTm

=


pw RT

v M=

Molar mass ( )MdRT w

density dp v

= ∴

P = Pressure of gas R = Universal gas constant

T = Temperature of gas in kelvins scale.

15. State Graham's law of diffusion.

Ans : The rate of diffusion of a given mass of gas at a given pressure and temperature is inversely

proportional to the square root of its density.

rate of diffusion 1

rd

α

16. Which of the gases diffuses faster among 2 2 4N ,O and CH ? Why ?

Ans : 4CH gas diffuse faster among 2 2 4N ,O and CH

Reason : 4CH gas (16) has low molecular weight than 2N (28) and 2O (32).

17. How many times methane diffuse faster than sulphurdioxide ?

Ans : According to Graham's law of diffusion.

4 2

2 4

64 42

16 1

CH SO

SO CH

r M

r M= = = =

Hence methane gas diffuses 2 times faster than SO2.

18. State Dalton's law of partial pressures.

Ans : The total pressure exerted by a mixture of chemically non-reacting gases at given

temperature and volume, is equal to the sum of partial pressures of the component gases.

19. Give the relation between the partial pressure of a gas and its mole fraction.

Ans : Partial pressure of a gas = mole fraction of the gas × Total pressue of the mixture of gases

Eg : Consider A and B in a container which are chemically non reaction.

∴ Partial pressure of ( )A A TA P X P= ×

partial pressure of ( )B B TB P X P= × , ,A BA B

A B A B

Xn n

Xn n n n

= =+ +

AX ,

BX are mole fractions

PT = Total pressure.

20. What is Boltzman's constant ? Give its value.

Ans : Boltzman's constant is the gas constant per molecule.

Boltzman's constant K=R

N

161.38 10 / ,erg k molecule−= ×


231.38 10 / ,J k molecule−= × .

21. What is RMS velocity ?

Ans : The square root of mean of the squares of the speeds of all molecules of a gas is known

as RMS velocity (URMS

)

2 2 2

1 2rms

1 2

U......

........

NU U U

n n

+ + +=

+ +

22. What is Average velocity ?

Ans : The arithematic mean of speeds of gas molecules is known as average velocity ( )avU

1 2

1 2

......

.......

n

averageU

U U U

n n

+ + +=

+ +

23. What is Most probable velocity ?

Ans : The speed velocity possessed by the maximum number of molecules of the gas is known

probable velocity ( )mpU

24. Give the ratio of RMS average and most probabale velocities of gas molecules.

Ans : 2 8 3

: : : : 1:1.128 :1.224mp av rms

RT RT RTU U U

M M M= =

Π

25. What is Compressibility factor ?

Ans : The ratio of the actual molar volume of a gas to the molar volume of a perfect gas under

the same conditions is called compressibility factor.

Compressibility factor ZPV

nRT= For a perfect Z = 1.

26. What is Boyle Temperature ?

Ans : The temperatue at which a real gas exibits ideal behaviour for a considerable range of

pressure is called Boyle's temperature.

27. What is critical temperature ? Give its value for CO2.

Ans : The temperature above which no gas can be liquified how ever high the pressure may be

applied is called critical temperature. →Critical temperature of 2CO gas is 31.980C.

28. What is critical Volume ?

Ans : The volume occupied by one mole of gas at critical temperature and critical pressure is

known as critical volume.

29. What is critical pressure ?

Ans : The pressure required to liquify a gas at critical temperature is known as critical pressure.

30. What are critical constants ?

Ans : Critical temperatue (Tc), Critical volume (V

c) and critical pressure (P

c) are called as

critical constants.

31. What is surface tension ?

Ans : The force acting at right angles to the surface of the liquid along unit length of surface is

called surface tension. → Unit : dynes/cm.


32. What is laminar flow of a liquid ?

Ans : In liquids a regular gradation of velocity for layers in passing from one layer to the next observed.

This flow of liquid is called Laminar flow.

33. What is coefficient of Viscosity ? Give its units.

Ans : The force of friction required to maintain velocity difference of 1cm. sec-1 between two

parallel layers of a liquid 1 cm. apart and each layer having an area 1cm2 is called coefficient

of viscosity.

→ It is denoted by η

→ F = ηAdu

dx

→ Units : Poise : in CGS system 1 poise = 1g. cm-1 sec-1.

34. Find the RMS speed of CO2 gas at 27°C temperature.

A. RMS velocity = 3RT

M

Given temperature T = 273 + 27 = 300 K

Molecular weight of CO2 = 44 g

Molar gas constant R = 8.314 x 107 ergs K–1 mol–1

∴ RMS velocity

73 8.314 10 300

44

× × × = 4.123 x 104 cm/s

35. Calculate the kinetic energy of 2 moles of CO2 gas at 27°C in Cals.

K.E. = 3

2nRT

Given n = 2, R = 2 cals, T = 273 + 27 = 300 K

32 2 300 1800

2∴ × × × = cals

5. STOICHIOMETRY

1. The empirical formula of a compound is 2CH O . Its molecular weight is 90. Calculate the molecular


Molecular formula = n (Emperical formula)Given

90

330

nMolecular wt

Emperical wt= = Molecular wt = 90

∴Molecular formula ( )2 3 6 33 CH O C H O= = Empericla formula = 2CH O

∴ Emperical wt = 30

2. Balance the following equation by the oxidation number method

( ) ( ) ( ) ( ) ( ) ( )3 32 3s saq aqCr Pb NO Cr NO Pb+ → + .


Ans :

So the balanced equation is 2Cr+3pb(NO3)

2 → 2Cr(NO

3)

3+ 3Pb

3. Calculate the volume of O2 at STP requied to completely burn 100 ml of acetylene.

Ans : Balanced chemical equation for combustion of acetylene is

2 2 2 2 22 5 4 2C H O CO H O+ → +

2 moles of 2 2C H require 5 moles of 2O for complete combustion at STP

2 22400 ml× of 2 2C H require ____ ?

100 5 22400 500250

2 22440 2ml

× ×= =

×

4. What volume of 2CO is obtained at STP by heating 4g of CaCO3 ?

Ans : Chemical Equation is

3 2CaCO CaO CO∆→ +

1 mole 3CaCO → 1 mole CO2 at STP

100 gms of CaCO3 → 22.4 lit of CO

2 at STP

4 gms OF CaCO3 → ?

4 22.40.896

100

×= lit

5. How many significant figures are present in the following ?

i) 0.0025 ii) 208 iii) 5005 iv) 126,000 v) 500.0 vi) 2.0034

Ans : i) 0.0025 has 2 significant figures ii) 208 has 3 significant figures

iii) 5005 has 4 significant figures iv) 126000 has 3 significant figures

v) 500.0 has 4 significant figures vi) 2.0034 has 5 significant figures

6. Round up the following upto three significant figures :

i) 34.216 ii) 10.4107 iii) 0.04597 iv) 2808

Ans : i) 34.216 becomes 34.2 ii) 10.4107 becomes 10.4

iii) 0.04597 becomes 0.046 iv) 2808 becomes 281

7. Assign oxidation number to the underlined elements in each of the following species :

a) 2 4NaH PO b) 4NaH SO c) 24 7H P O d) 2 4k MnO

e) 2CaO f) 4NaBH g) 22 7H S O h) ( )4 22

Kal SO .12H O

Ans : a ) 42NaH PO−

1 2 8 0x+ + − =


5 0x − = ,x = +5

Oxidation no.of 'p' in 2 4 5NaH PO = +

b) 4NaH OS

( ) ( ) ( )1 1 1 1 4 2 0x+ + + + + − = ⇒ 1 1 8 0x+ + − = , 6x = +

Oxidation no. of 'S' in 4 6NaHSO = +

c) 24 7H P O

( ) ( )4 1 2 7 2 0x+ + + − =

2 10 0x − = ⇒ 5x = +

Oxidation no.of 'P' in 4 2 7H P O is +5

d) 2 4K MnO

( ) ( )2 1 4 2 0x+ + + − =

2 8 0x+ − = ⇒ x = +6

Oxidation no of Mn is 2 4K MnO = + 6

e) 2CaO

+2+2x = 0, 2x = -2

x = -1

Oxidation no. of oxygen in 2CaO 1= −

f) 4NaBH

( ) ( )1 1 4 1 0x+ + + − = 1 4 0x+ − = 3x = +

Oxidation no. of 'B' in NaBH4 = +3

But 'B' most probably exhibis - 3 oxidation state.

g) 2 2 7H S O

( ) ( )2 1 1 7 2 0X+ + − =

2+2X-14=0

2 12 0x − = ⇒ x = +6

Oxidation state of 'S' in 2 2 7H S O = + 6

h) ( )4 22k Al SO 12H O :

General formula of above compound is

( )2 4 2 4 2324k SO Al SO H O (Potash alum)

Consider ( )2 4 3Al SO

Consider ( )2 4 3Al SO from the above double salt


( )2 3 2 0x + − = 2 6 0x − = 3x = +

8. What are the oxidation number of the underlined elements in each of the following and

how do you rationalise your results ?

a) 3K I b) 42 6H S O c) 3 4Fe O

Ans : 3KI

It is formed by the combining 2KI,I

∴ Oxidation no. of 'I' in KI = -1

[ 1+ x = 0 , In I2 oxidation no. of I = 0

b) 2 4 6H S O

According to 2 4 6H S O structure

( ) ( )2 1 4 6 2 0 4 10 0X X+ + − = ⇒ − =

'S' average Ox. No = 2.5 x = 2.5

c) 3 4Fe O

( )3 4 2 0x + − =

3x-8 = 0 ⇒8

3x =

In general 3 4Fe O obtained by FeO + Fe2 O

3

[ ]In FeO x=+2 2 0 2x x∴ → − = ⇒ = +

In [ ]2 3 3 2 6 0 3Fe O x x x→ = + − = ⇒ = +

c) i, 3 2CH CH -OH

2 6C H O

( ) ( )2 6 1 2 0X + + − =

2 6 2 0X + − = ⇒ x = -2

d) 3HC COOH

2 4 2C H O

( ) ( )2 4 1 2 2 0x + + + − = ⇒ 2x + 4- 4 = 0

x = 0


9. Calculate the oxidation number of sulphur, chromium and nitrogen in 2

2 5 2 7,H SO Cr O− and

3NO− .

Suggest structure of these compounds.

Ans : a) 2 5H SO Structure :

( ) ( ) ( )2 1 2 1 3 2 0x+ + − + − = (One peroxy linkage )

2 2 6 0x+ − − = ⇒ x = + 6

b) 2

2 7Cr O− Structure :

( ) ( )2 7 2 2x + − = −

2 14 2x − = − ⇒ 6x = +

c) 3NO−

( )3 2 1x + − = −

x-6=-1 ⇒ x = +5

10. How many number of moles of glucose are present in 540gm of glucose.

A: No.of moles = Given weight

Gram molecular weight

5403

180= = moles

11. Calculate the weight of 0.1 moles of sodium carbonate.

A: No.of moles Given weight

Gram molecular weight=


⇒ Weight = No.of moles x Gram mol. weight = 0.1 x 106 = 10.6gm

12. How many molecules of Glucose are present in 5.23 gm of Glucose?

A: No.of molecules = No.of moles x N.

235.236.023 10

180= × × = 1. 75 x 1022 molecules

13. Calculate the number of molecules. Present in 1.12 x 10–7 c.c of a gas of S.T.P

A: 1 mole = 6 .023 x 1023 molecules = 22,400 c.c at S.T.P

?

– 1.12 x 10–7 c.c.

723 121.12 10

6.023 10 3.01 1022400

molecules−×

× × = ×

6. THERMODYNAMICS

1. Define a system. Give an example.

Ans : System : A small part of the universe chosen for thermodynamic study is called system.

2. What is the workdone in the free expansion of an ideal gas in reversible and irreversible

processes ?

Ans : In case of free expansion of an ideal gas pressure becomes zero. (into vaccum)

∴No workdone during free expansion of an ideal gas in case to reversible as well as irreversible process.

3. What are intensive and extensive properties ?

Ans : Measurable (or) macroscopic properties such as mass, pressure, volume, temperature,

surface tension, viscosity etc, can be subdivided into two categories as below :

i) Extensive properties : The properties whose magnitude depends upon the quantity of

matter present in the system are called extensive properties. Examplels of such properties

are mass, volume, heat capacity, internal energy, entropy, heat content , gibbs free energy

etc. These properties change with quantity of matter present in the system. These properties

are additivein nature.

4. Give the equation that gives the relationship btween U and H∆ ∆ .

Ans : The equation that gives the relationship between U and H∆ ∆

U H nRT∆ = ∆ + ∆

H∆ = Change in Enthalpy

U∆ = Change in Internal energy

,P R

n n n R∆ = − = Universal gas constant

T = Temperature

5. What is the relationship between P V

C and C ?

Ans : p VC C R− = , pC = Heat capacity at constant pressure

VC = Heat capacity at constant volume

R = Universal gas constnat.

jr chemistry - chapter wise important questions - part 1

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