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Chapter 4: Chemical Bonding and Molecular Structure

Q-1) Explain Chemical bond formation process.

Ans.)

“Chemical bond is an attractive force that bounds the constituents of a chemical species together.”

So many theories are suggested for chemical bond formation such as valence shell electron pair repulsion theory, electronic theory, molecular orbital theory, and valence bond theory.

Formation of chemical bond is credited to the tendency of system to achieve stability. It was noticed that inertness of noble gasses is direct result of their completely filled outermost orbitals. Consequently, it was proposed that the elements having deficiency of electrons inoutermost shells are unstable. Thus, atomscombine with one another and finish their separate octets or duplets to achieve the stable configuration of the closest inert gasses. So, thiscombination may occur either by sharing of electrons. The formed chemical bond is a result of sharing of electrons amongatoms is known as a covalent bond.  Also a formed ionic bond is a result of sharing of electrons amongatoms.

 

Q-2) Give the lewis dot symbols for the elements given below:

a) Na

b) Mg

c) O

d) N

e) B

f) Br

Ans.)

a) Na

Sodium atom contains only 1 valence electron. Thus, the lewis dot symbols for Na is \(Na\cdot\)

 

b) Mg

Magnesium atom contains only 2 valence electrons. Thus, the lewis dot symbols for Mg is

1

c) O

Oxygen atom contains only 6 valence electrons. Thus, the lewis dot symbols for O is

2

d) N

Nitrogen atom contains only 5 valence electrons. Thus, the lewis dot symbols for N is

3

e) B

Boron atom contains only 3 valence electrons. Thus, the lewis dot symbols for B is

4

f) Br

Bromine atom contains only 7 valence electrons. Thus, the lewis dot symbols for Br is

5

Q-3) Give the lewis dot symbols for the elements given below:

(a) Al and Al3+

(b) H and H

(c) S and S2-

Ans.)

(a) Al and Al3+

Aluminium atom contains only 3 valence electrons. Thus, the lewis dot symbols for Al is

6

The tri-positive charge on aluminium indicates that it has donated 3 electrons.

Thus lewis dot symbol is

\([Al]^{3+}\)

 

(b) H and H

Hydrogen atom contains only 1 valence electrons. Thus, the lewis dot symbols for H is

\(H\cdot\)

The single negative charge on hydrogen indicates that it has gained 1 electron. So, one valance electron plus one gained electron.

Thus lewis dot symbol is

7

(c) S and S2-

Sulphur atom contains only 6 valence electrons. Thus, the lewis dot symbols for S is

8

The bi-negative charge on sulphur indicates that it has gained 2 electron. So, six valance electron plus two gained electron.

Thus lewis dot symbol is

9

 

Q-4) Give the lewis structure for the ions and molecules given below:

(a) SiCl4

(b) H2S

(c) HCOOH

(d) BeF2

(e) \(CO_{3}^{2-}\)

Ans.)

(a) SiCl4

10

 

(b) H2S

11

 

(c) HCOOH

12

 

(d) BeF2

13

 

(e) \(CO_{3}^{2-}\)

14

 

Q-5) Write the definition of octect rule. State the significance and limitation of octet rule.

Ans.)

Octet rule says, “atoms can combine either by transfer of valence electrons from one atom to another or by sharing their valence electrons in order to achieve the nearest inert gas configuration by having an octet in their valence shell.”

15

Octet rule explains chemical bond formation depending upon nature of element.

Limitations:

(a) Octet rule fails to predict the relative stability and shape of the molecules.

(b) It is based on inert nature of noble gases. But, some inert gases say, krypton(Kr) and xenon(Xe) form compounds like KrF2, XeF2 etc.

(c) For elements beyond 3rd period the octet rule cannot be applied. Elements present beyond 3rd period have more than 8 valence electrons surrounding central atom. E.g. SF6, PF6 etc.

16

(d) For atoms in a molecule having odd number of electrons octet rule is not applied. E.g. For No2 and NO octet rule is not applicable.

17

(e) If an compound is having less than 8 electrons surrounding the central atom than octet rule cannot be applied to that compound. E.g. BeH2, AlCl3, LiCl etc. is not obeying octet rule.

18

 

Q-6) State the factors that are favourable for ionic bond formation.

Ans.)

Formation of ionic bond takes place by transfer of 1 or more electrons from one atom to another. Thus, ionic bond formation depends on the flexibility of neutral atoms to lose or gain electrons. Formation of ionic bond also depends on the lattice energy of compound which is formed.

The factors that are favourable for ionic bond formation:

(a) High electron affinity of atoms of non- metal.

(b) High lattice energy of compound which is formed.

(c) Low ionization enthalpy of atom of metal.

 

Q-7) Comment about the shape of the molecules given below by  using VSEPR model:

(a) BCl3

(b) SiCl4

(c) AsF5

(d) PH3

(e) BeCl2

(f) H2S

Ans.)

(a) BCl3

Central atom is not having any lone pair but have 3 bond pairs. Thus, its shape is AB3. i.e..Trigonal planar.

19

 

(b) SiCl4

Central atom is not having any lone pair but have 4 bond pairs. Thus, its shape is AB4. i.e. Tetrahedral.

20

 

(c) AsF5

Central atom is not having any lone pair but have 5 bond pairs. Thus, its shape is AB5. i.e. Trigonalbipyramidal.

21

 

(d) PH3

Central atom is having 1 lone pair andis having 3 bond pairs. Thus, its shape is AB3E. i.e. Trigonalbipyramidal.

22

 

(e) BeCl2

Central atom is not having any lone pair but have 2 bond pairs. Thus, its shape is AB2. i.e. Linear shape.

23

(f) H2S

Central atom is having 1 lone pair and is having 2 bond pairs. Thus, its shape is AB2E. i.e. Bent shape.

24

 

Q-8) Explain why the bond angle in H2O is less than NH3, even though they are having distorted tetrahedral structure.

Ans.)

The geometry of H2O and NH3:

35

Central atom(N) in ammonia is having 1 lone pair and is having 3 bond pairs.

Central atom(O) in water  is having 2 lone pair and is having 2 bond pairs.

Thus, these 2 lone pairs on O- atom in water molecule repels the 2 bond pairs. And this repulsion is between lone pair and bond pair on O- atom of H2O is stronger than the repulsion is between lone pair and bond pair on N-atom of NH3.

Thus, the bond angle in H2O is less than NH3, even though they are having distorted tetrahedral structure.

 

Q-9) Express,‘bond strength in terms of bond order’.

Ans.)

The extent of bonding which occurs between two atoms while forming a molecule is represented by bond strength. As the bond strength increases the bond becomes stronger and the bond order increases.

 

 

Q-10) Define: Bond length.

Ans.)

“Bond length is defined as the equilibrium distance between the nuclei of 2 boned atoms in a molecule.”

 

Q-11) Explain important aspects of the resonance with reference to \(CO_{3}^{2-}\) ion.

Ans.)

Experimental results shows that, all the C-O bond in \(CO_{3}^{2-}\) are equivalent.

Thus, it is inefficient to represent \(CO_{3}^{2-}\) ion by single lewis structure which is having 1 double bond and 2 single bond.

Thus, the resonance structures of \(CO_{3}^{2-}\) is :

26

 

Q-12) H3PO3 can be represented by following structures.

Can we take these 2 structures as canonical form of resonance hybrid which is representing H3PO3? Explain.

Ans.)

27

In the given structures the position of atoms are changed, so we cannot take the 2 given structures as canonical form of resonance hybrid which is representing H3PO3.

 

 

Q-13) Give the resonance structures of the following ions

(a) \(NO_{3}^{-}\)

(b) NO2

(c) SO3

Ans.)

(a) \(NO_{3}^{-}\)

28

 

(b) NO2

29

(c) SO3

30

 

Q-14) Show transfer of electrons between the atoms given below to form anions and cation.

(i) O and  Ca

(ii) N and Al

(iii) S and K.

Ans.)

(i) O and  Ca

Electronic configurations of O and Ca are:

O: 2, 6

Ca: 2, 8, 8, 2

Here, it clear that Ca has 2  more electrons then nearest inert gas. i.e. Ar, whereas O needs 2 electrons to complete its octet. Thus, the transfer of electrons takes place in following way,

31

 

(ii) N and Al

Electronic configurations of N and Al are:

N: 2, 5

Al: 2, 8, 3

Here, it clear that Al has 3 more electrons then nearest inert gas. i.e. Ne, whereas N needs 3 electrons to complete its octet. Thus, the transfer of electrons takes place in following way,

32

 

(iii)S and K

Electronic configurations of S and K are:

S: 2, 8, 6

K: 2, 8, 8, 1

Here, it clear that K has 1 more electrons then nearest inert gas. i.e. Ne, whereas S needs 2 electrons to complete its octet. Thus, the transfer of electrons takes place in following way,

33

 

Q-15) H2O and CO2 are tri-atomic molecules, then also the shape of H2O is bent and that of shape of CO2 is linear. Why?

Explain with the help of dipole moment.

Ans.)

Experimental results show that the dipole moment of CO2 is 0. And it is possible only if the shape of molecule is linear as dipole moments of bond between C-O is equal and opposite so, it nullifies each other.

34

\(∴ Resultant,\; \mu = 0\)

H2O has 1.84D dipole moment. The value of dipole moments indicates that structure of water molecule is bent as dipole moments of bond between O-H is unequal.

35

 

Q-16) Give the significance of dipole moment.

Ans.)

In heteronuclear molecule, there is difference in electro-negativities of constituents of atom, due to which polarisation arisesin it. Thus, one end of acquires positive charge and other end acquires negative charge. So, molecules are said to have a dipole moment.

The product of distance between centres of negative and positive charges and magnitude of charge is known as dipole moment (\(\mu\)). Dipole moment is a vector quantity and it is represented by an arrow in such a way that head of a tail is pointing towards negative centre and tail of arrow is pointing towards positive centre.

\(\mu\) = distance between the centres ( r) * charge (Q)

SI unit of dipole moment is ‘esu’.

1 esu = \(3.335\times 10^{-30} cm\)

It is the measure of polarity of a bond.It differentiates the non-polar and polar bonds as all the non-polar molecules have dipole moment equal to 0. It also calculates the ionic character of a molecule.

36

 

Q-17) Give definition of electronegativity.State difference between electronegativity and electron affinity.

Ans.)

“Electronegativity is the ability of an atom in a chemical compound to attract a bond pair of electrons towards itself”.

Sr. No Electronegativity Electron affinity
1 Tendency to attract the shared pairs of electrons for an atom which is in chemical compound is its electronegativity. Tendency to gain electrons for an isolated gaseous atom is its electron gain enthalpy.
2 It varies according to element to which it is bounded with. It does not vary according to element to which it is bounded with.
3 It is not constant for any element. It is constant for an element.
4 It is not a measurable quantity. It is a measurable quantity.

 

 

Q-18) Explain polar covalent bond with suitable example.

Ans.)

When two unique atoms having distinctelectronegativities join to form a covalent bond, the bond pair ofelectrons is not shared equally. The nucleus of an atom having greater electronegativityattracts the bond pair. So, the electron distribution gets distorted and an electronegativity atom attracts the electron cloud.

Thus, the electronegative element gets slightly negatively charged and on the other hand, the other atom gets slightly positively charged. As a result of this, two opposite poles are developed in a molecule and this type of bond formed is termed as ‘polar covalent bond’.

E.g. HCl is having a polar covalent bond. In HCl, Cl- atom is having more electronegativity than H- atom. Thus, bond pair shifts towards Cl- atom and because of that it acquires positive charge.

37

 

Q-19) For molecules given below, arrange the bonds in increasing order of their ionic characteristic.  

K2O, N2, ClF2, LiF, SO2

Ans.)

Ionic characteristic of a molecule depends on the difference in electronegativity between constituents atoms. So, higher the difference, the ionic characteristic of a molecule will be higher.

So, the required order of ionic characteristic of the given molecules is

N2< SO2< ClF3< K2O <LiF.

 

 

Q-20) Some of the bonds are incorrectly shown in the skeletal structure of CH3COOH. Give the correct Lewis structure of CH3COOH.

Ans.)

38

Correct Lewis structure of CH3COOH is given below:

39

 

Q-21) Ch4 can also have square planar geometry other than tetrahedral geometry, and in that geometry all the 4 H- atoms are at the corner of the square and C- atom is at the centre. Give reason why Methane’s geometry is not square planer.

Ans.)

Electronic configuration of C- atom:

6C :\(1s^{2}\;2s^{2}\;2p^{2}\)

Orbital picture of C- atom in excited state is:

40

Thus, C- atom undergoes sp3 hybridization in methane molecule and forms tetrahedral structure.

41

For square planer geometry, C-atom should have dsp2 hybridization. But as C- atom is not having d- orbital so it cannot undergo dsp2 hybridization. Thus, Methane’s geometry cannot be square planer.

Also in square planar geometry the bond angle is \(90^{\circ}\) so the stability is not there because of repulsion between bond pairs. So as per VSEPR theory methane’s tetrahedral structure is perfect.

 

Q-22) Even though Be- H bonds are polar In BeH2, the dipole moment of BeH2 is 0. Why is it so?

Ans.)

Lewis structure of BeH2 is:

42

Central atom is not having any lone pair but have 2 bond pairs. Thus, its shape is AB2. i.e. Linear shape.

43

Thus, the dipole moment of Be- H bond is equal and opposite in direction so it nullifies one another. Thus, the dipole moment of BeH2 is 0.

 

Q-23) Which is having higher dipole moment out of NF3 and NH3?

Ans.)

N- atom is the central atom of NF3 and NH3.

Central atom is having 1 lone pair and is having 3 bond pairs. Thus,  for both teh shape is AB3E. i.e. Pyramidal.

As, F-atom is more electronegativity than H- atom,NF3 should have higher dipole moment than NH3. But the dipole moment of NH3 is 1.46D which is higher than dipole moment of NF3 which is 0.24D.

It gets clear from the directions of dipole moments of individual bond in NF3 and NH3.

44

As, both the N-H bond are in same direction it adds to the bond moment of the lone pair, while N-F bond are in opposite direction so they partly cuts the bond moment of lone pair.

Thus, dipole moment of NH3 is higher than that of NF3.

 

 

Q-24) Explain ‘hybridisation of atomic orbitals’. Draw sp, sp2, sp3, hybrid orbitals. 

Ans.)

“Hybridization is defined as an intermixing of a set of atomic orbitals of slightly different energies, thereby forming a new set of orbitals having equivalent energies and shapes”.

E.g. 1 s- orbital hybridises with 3 p- orbitals to form 4 sp3 hybrid orbitals.

 

(a) sp hybrid orbital

1 s- orbital hybridises with 1 p- orbitals to form 2 sp hybrid orbitals. sp hybrid orbital is having linear shape. The formation of sp orbital is:

45

 

(b) sp2 hybrid orbital

 

1 s- orbital hybridises with 2 p- orbitals to form 3 sp2 hybrid orbitals. The shape of sp2 orbital is trigonal planar.

46

 

(c) sp3 hybrid orbital

 

1 s- orbital hybridises with 3 p- orbitals to form 4 sp3 hybrid orbitals. The shape of sp3 orbital is tetrahedron.

47

 

Q-25) Whether the hybridisation of Al- atom change for the reaction given below?

AlCl3 + \(Cl^{-}\)à\(AlCl_{4}^{-}\)

Ans.)

The ground state of valence orbital of Al –atom is:

48

In excited state the orbital picture of Al- atom is:

49

Thus, Al -atom in AlCl3 undergoes sp2 hybridisation and forms trigonal planar geometry. Foe the formation of \(AlCl_{4}^{-}\) the vacant 3pz orbital will also get involved. Thus, sp2 hybridisation is converted into sp3 hybridisation and forms a tetrahedral structure.

 

 

Q-26) Determine that there is any change in hybridisation of N-Atom and B- atom because of the reaction given below:

BF3 + NH3à F3B.NH3

Ans.)

N- atom in NH3 is having sp3 hybridization. Orbital picture of N- atom is shown below:

50

B- atom in NF3 is having sp2 hybridisation . Orbital picture of B- atom is shown below:

51

On the reaction of NH3 and BF3, F3B.NH3 is obtained as product, as hybridization of B-atom is changed to sp3. Although, hybridization of N- atomsremains unchanged.

 

 

Q-27) Explain the formation of double and triple bond in C2H2 and C2H4 with the help of diagrams.

Ans.)

  • C2H2

In formation of ethyne(C2H2) molecule, C- atom is having sp hybridization with 2 2p- orbitals in unhybridized state.

1 sp orbital of each C- atom overlaps the inter-nuclear axis and forms C-C sigma bond. The 2ndsp orbital of each C- atom overlaps half-filled 1s orbital so as to form a sigma bond.

The triple bond between the 2 C- atoms has 1 sigma bond and 2 Pie bonds. This is because 2 unhybridized 2p- orbitals overlaps with the 2p- orbital of other C- atom, thus forming 2 pie bonds.

52

 

 

  • C2H4

Electronic configuration of carbon atom in excited state is given below:

6C: \(1s^{2}2s^{1}2p_{x}^{1}2p_{y}^{1}2p_{z}^{1}\)

In the formation of C2H4 (ethane) molecule 1 sp2 orbital of C- atom overlaps sp2 orbital of other C- atom. Thus, forming a C-C sigma bond.

The 2 remaining sp2 orbital of every C- atom forms sp2-s \(\sigma\) bond with 2 H- atoms. One c- atom having unhybridized orbital overlaps with the unhybridized orbital of other C- atom and forms a pie bond.

53

 

Q-28) Find out the number of \(\pi\)  and \(\sigma\) bonds in the molecules given below:

(i) C2H4

(ii) C2H2

Ans.)

Single bond is formed as the axis of bonding orbital overlaps. Thus, it forms a \(\sigma\) bond. By sidewise overlapping of orbital double and triple bonds i.e. multiple bonds are formed.\(\pi\) bond is always present in the multiple bonds. Triple bond consist of 2 \(\pi\) and 1\(\sigma\) bond.

(i) C2H2

54

Thus, there are 2 \(\pi\) bonds and 3 \(\sigma\) in C2H2.

 

(ii) C2H4

Thus, there are 1 \(\pi\) bonds and 5 \(\sigma\) in C2H4.

55

 

Q-29) Take x- axis as inter-nuclear axis and find which of the orbital given below will not form \(\sigma\) bond and why?

(i) 1s and 2px

(ii) 1s ans 1s

(iii) 1s and 2s

(iv) 2py and 2py

Ans.)

(iv) 2py and 2py

2py and 2py orbitals won’t form a \(\sigma\) as it will undergo lateral over lapping and will form a \(\pi\) bond.

 

 

 

Q-30) Give the hybrid orbitals used by the C- atoms in the molecules given below:

(i) CH3-CH=CH2

(ii) CH3-CH3

(iii) CH3-CHO

(iv) CH3COOH

(v) CH3CH2-OH

Ans.)

(i) CH3-CH=CH2

56

Here, C3 and C2 are having sp2 hybridization and C1 is having sp3 hybridization.

 

(ii) CH3-CH3

57

Here, C1 and C2 are having sp3 hybridization.

 

(iii) CH3-CHO

58

Here, C1 is having sp3 hybridization and C2 is having sp2 hybridization.

 

(iv) CH3COOH

59

Here, C1 is having sp3 hybridization and C2 is having sp2 hybridization.

 

(v)CH3CH2-OH

60

Here, C1 and C2 are having sp3 hybridization.

 

 

Q-31) Explain lone pairs and bond pairs of electrons with an example.

Ans.)

A covalent bond is formed when 2 atoms combine with each other by sharing their valence electrons.

“The shared pairs of electrons present between the bonded atoms are called bond pairs”. Each and every electron cannot participate in bonding. “The pairs of electrons which do not participate in bonding are called lone pairs”.

E.g. a) Ethane is having 7 bond pairs but zero lone pair.

61

b) Water is having 2 bond pairs and 2 lone pairs on O- atom.

62

 

Q-32) State difference between pi bond and sigma bond.

Ans.)

Sr. No. Pi bond Sigma bond
1 Pi bond is formed by lateral overlapping of orbitals. Sigma bond is formed by end to end overlapping of orbitals.
2 It is comparatively weak bond. It is comparatively strong bond.
3 There is only one overlapping orbital is p-p. The overlapping orbitals are s-s, s-p, p-p.
4 Rotation around pi- bond is restricted. Rotation is possible around sigma bond.
5 Electron cloud is not symmetrical about the line joining 2 nuclei. Electron cloud is symmetrical about the line joining 2 nuclei.
6 It is having 2 electron clouds one above the plane of atomic nuclei and one below the plane of atomic nuclei. It is having 1 electron cloud and that is symmetrical about the inter-nuclear axis.

 

 

Q-33) Using valence bond theory, explain the formation is hydrogen molecule.

Ans.)

Assuming 2 H- atoms X and Y with nuclei NX and NY and electrons eX and eY respectively.

When X and Y are far for each other then there is no interaction between them. As soon as they come closer, the attractive force and repulsive force becomes active.

The repulsive forces are:

(i) Between electrons of both the atoms i.e. eX and eY.

(ii) Between nuclei of boththe atoms i.e. NX and NY.

The attractive forces are:

(i) Between the electro and nucleus of the same atom i.e. NX – eX and NY -eY.

(ii) Between the electron of one atom and nucleus of other atomi.e. NX–eY and NY–eX.

The repulsive force push the 2 atoms apart whereas the attractive force tend to bring them together.

 

Repulsive forces:

63

Attractive forces:

64

The values of repulsive forces are less than that of attractive forces. Thus, 2 atoms approach each other. Thus, there is decrease in potential energy. At the end a stage is reached when the repulsive forces balance the attractive forces and the system achieves the minimum energy,which leads to formation of H2 molecule.

 

 

Q-34) What are the condition that is required for linear combination of an atomic orbitals to form a molecular orbitals.

Ans.)

The condition that is required for linear combination of an atomic orbitals to form a molecular orbitals are as follows:

(i) The joining of atomic orbitals must have approximately same energy. Thisimplies in a homo-nuclear molecule, the 1s-orbital of one atom can join with the 1s- orbital of another atom, but cannot join  with the 2s-orbital.

(ii) The joining atomic orbitals must have legitimate orientations to ensure  themaximum overlap.

(iii) The overlapping must be in a large extent.

 

 

Q-35) Explain why molecule Be2doesn’t exist by using molecular orbital theory.

Ans.)

Electronic configuration of Be:

1s2 2s2

Molecular orbital electronic configuration of Be2 is:

\(\sigma _{1s}^{2}\; \sigma _{1s}^{\cdot 2}\;\sigma _{2s}^{2}\;\sigma _{2s}^{\cdot2}\)

Thus, bond order of Be2: 0.5(Nb – Na).

Na: No. of electrons in the bonding orbitals

Nb: No. of electrons in the anti-bonding orbitals

Therefore, bond order of Be2 = 0.5(4 – 4) = 0

Zero value of bond order indicates that given molecule is unstable. Thus, Be2 doesn’t exist.

 

 

Q-36) Compare relative stability of the species given below and also write their magnetic properties.

(a) O2

(b) \(O_{2}^{2-}\) (peroxide)

(c) \(O_{2}^{-}\) (superoxide)

(d) \(O_{2}^{+}\)

Ans.)

O2 contain 16 electrons i.e. 8 electrons from each O- atom.

 

  • Electronic configuration of O2 is:
\([\sigma – (1s)]^{2}[\sigma ^{*}(1s)]^{2}[\sigma (2s)]^{2}[\sigma ^{*}(2s)]^{2}[\sigma (1p_{z})]^{2}[\pi (2p_{x})]^{2}[\pi (2p_{y})]^{2}[\pi^{*} (2p_{x})]^{1}[\pi^{*} (2p_{y})]^{1}\)

As 1s- orbital of each O- atom does not involve in the bonding,

No. of bonding electrons = Nb = 8

No. of anti-bonding electrons = Na = 4

Now,

Bond order = 0.5(8 – 4) = 2

 

  • Electronic configuration of \(O_{2}^{2-}\) (peroxide) is:
\(KK[\sigma(2s)]^{2}[\sigma ^{*}(2s)]^{2}[\sigma (2p_{z})]^{2}[\pi (2p_{x})]^{2}[\pi (2p_{y})]^{2}[\pi^{*} (2p_{x})]^{2}[\pi^{*} (2p_{y})]^{2}\)

No. of bonding electrons = Nb = 8

No. of anti-bonding electrons = Na = 6

Now,

Bond order = 0.5(8 – 6) = 1

 

  • Electronic configuration of \(O_{2}^{-}\) (superoxide) is:
\(KK[\sigma(2s)]^{2}[\sigma ^{*}(2s)]^{2}[\sigma (2p_{z})]^{2}[\pi (2p_{x})]^{2}[\pi (2p_{y})]^{2}[\pi^{*} (2p_{x})]^{2}[\pi^{*} (2p_{y})]^{1}\)

No. of bonding electrons = Nb = 8

No. of anti-bonding electrons = Na = 5

Now,

Bond order = 0.5(8 – 5) = 1.5

 

  • Electronic configuration of \(O_{2}^{+}\) is:
\(KK[\sigma(2s)]^{2}[\sigma ^{*}(2s)]^{2}[\sigma (2p_{z})]^{2}[\pi (2p_{x})]^{2}[\pi (2p_{y})]^{2}[\pi^{*} (2p_{x})]^{1}\)

No. of bonding electrons = Nb = 8

No. of anti-bonding electrons = Na = 3

Now,

Bond order = 0.5(8 – 3) = 2.5

As, bonddissociation energy \(\propto\) bond order

Hence, higher the bond order, higher stability will be there.

Arrangement according to decreasing order of stability is given as:

\(O_{2}^{+} > O_{2} > O_{2}^{-} > O_{2}^{2-}\)

 

 

Q-37) What is the significance of positive (+) and negative (-) sign in representing the molecular orbitals?

Ans.)

Generally molecular orbital are represented by the ‘wave function’.

Positive (+) sign in representing a molecular orbital indicates positive wave function.

Negative (-) sign in representing a molecular orbital indicates negative wave function.

 

 

Q-38) Which hybridization is PCl5 is having? Why are equatorial bonds are shorter than axial bonds?

Ans.)

The electronic configuration of outer orbital of phosphorus in excited state and in ground state is given below:

Ground State:

65

Excited State:

66

Phosphorus atom is having sp3d hybridization. These orbitals are filled due to donation of electron pairs by 5 Cl- atoms: PCl5

67

The 5 sp3d hybrid orbitals present here are directed towards 5 corners of trigonalbipyramidals. Thus, the geometry of PCl5 is given below:

68

PCl5 contains 5 P- Cl sigma bonds. Out of which 3 P-Cl bond lie in only 1 plane and they are making \(120^{\circ}\) with each other. And as these bonds lie in 1 plane they areknown equatorial bonds.

Out of 2 remaining P-Cl bonds one bond lie above the equatorial plane and one bond lie below the equatorial bond. And they are making \(90^{\circ}\) with each other. These bonds are called axial bond.

Equatorial bond pairs repel axial bond pairs to a large extent so, equatorial bonds are slightly shorter than axial bonds.

 

 

Q-39) Give definition of H- bond (Hydrogen bond). Is H- bond stronger or weaker compared to Van der Waals force?

Ans.)

“H-bond is defined as an attractive force acting between the hydrogen attached to an electronegative atom of one molecule and an electronegative atom of a different molecule”.

As there is differencein the electronegativities between the atoms, thus, the bond pairelectronegative  atom and hydrogen atom is drifted away from H- atom. Therefore, hydrogen atom gets electropositive w.r.t. the other atom and procures a positive charge.

 

\(4^{\delta } – X^{\delta }\;………..H^{\delta +} – X^{\delta -}\;………….H^{\delta +} – X^{\delta -}\)

 

The value of H- bond is minimum in gaseous state and maximum in the solid state.

Two types of hydrogen bond are there:

(a) Intramolecular hydrogen bond e.g., o- nitrophenol

(b) Inter molecular hydrogen bond e.g., HF, H2O etc.

69

H- bonds are stronger than Van der Waals forces as H- bond are regarded as extreme form of the dipole- dipole interaction.

 

 

Q-40)What do you mean by bond order? Also calculate the bond order for:

(a) O2

(b) \(O_{2}^{+}\) (superoxide)

(c) \(O_{2}^{-}\)

(d) N2

Ans.)

Bond Order: It is defined as 0.5 times the difference between the “No. of electrons present in bonding orbitals and No. of electrons present in anti- bonding orbitals” of a molecule.

Bond Order = 0.5(Nb – Na);

Na: No. of anti-bonding electrons

Nb: No. of bonding electrons

O2 contain 16 electrons i.e. 8 electrons from each O- atom.

 

 

 

 

 

  • Electronic configuration of O2 is:

 

\([\sigma – (1s)]^{2}[\sigma ^{*}(1s)]^{2}[\sigma (2s)]^{2}[\sigma ^{*}(2s)]^{2}[\sigma (1p_{z})]^{2}[\pi (2p_{x})]^{2}[\pi (2p_{y})]^{2}[\pi^{*} (2p_{x})]^{1}[\pi^{*} (2p_{y})]^{1}\)As 1s- orbital of each O- atom does not involve in the bonding,

No. of bonding electrons = Nb = 8

No. of anti-bonding electrons = Na = 4

Now,

Bond order = 0.5(8 – 4) = 2

 

  • Electronic configuration of \(O_{2}^{-}\) (superoxide) is:
\(KK[\sigma(2s)]^{2}[\sigma ^{*}(2s)]^{2}[\sigma (2p_{z})]^{2}[\pi (2p_{x})]^{2}[\pi (2p_{y})]^{2}[\pi^{*} (2p_{x})]^{2}[\pi^{*} (2p_{y})]^{1}\)

No. of bonding electrons = Nb = 8

No. of anti-bonding electrons = Na = 5

Now,

Bond order = 0.5(8 – 5) = 1.5

 

  • Electronic configuration of \(O_{2}^{+}\) is:
\(KK[\sigma(2s)]^{2}[\sigma ^{*}(2s)]^{2}[\sigma (2p_{z})]^{2}[\pi (2p_{x})]^{2}[\pi (2p_{y})]^{2}[\pi^{*} (2p_{x})]^{1}\)

No. of bonding electrons = Nb = 8

No. of anti-bonding electrons = Na = 3

Now,

Bond order = 0.5(8 – 3) = 2.5

  • Electronic configuration of N2 is:
\([\sigma(1s)]^{2}[\sigma ^{*}(1s)]^{2}[\sigma (2s)]^{2}[\sigma ^{*}(2s)]^{2}[\pi (2p_{x})]^{2}[\pi (2p_{y})]^{2}[\pi (2p_{z})]^{2}\)

No. of bonding electrons = Nb = 10

No. of anti-bonding electrons = Na = 4

Now,

Bond order = 0.5(10 – 4) = 3



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1 Comment

  1. Tonty Tonty
    May 26, 2018    

    I don’t understand why you change questions. You should review question number 34.

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