A. Very short answer question:

1.

The postulates of VSEPR theory are:

a. In a covalent molecule, there is a central atom having electron pairs (may be lone or bond pair) to which other atoms are bonded.

b. The shape of molecule depends upon the number of electron pair surrounding the central atom and their repulsion.

 

2.

In NH₃ molecule, there is one lone pair of electron in nitrogen. According to VSEPR theory we have, lone – pair – lone pair repulsion > bond pair – bond pair repulsion.sp³hybridization, pyramidal geometry with bond angle 109.5° to 107.48°.

 

3.

Shape of BeF₂ by using VSEPR theory:

BeF2: In BeF2, Be is the central atom. The electronic configuration of Be is 1s², 2s². It consists two valance electrons. The two valance electrons forms two bond pairs with two fluorine atoms. Now, the central Be is surrounded by two bond pairs which is situated in linear manner making an angle of 180°. So, BeF₂ molecule has linear geometry with (FBeF = 180°) FBeF bond angle is 180°.

The lewis dot structure is:

Electron pair = 2

Bond pair = 2

Lone pair = 0.

 

BF₃: In BF3, boron B is the central atom. The electronic configuration of B is 1s²,2s²,2p¹. It consists of three electron pairs with three fluorine atoms. The central boron atom is surrounded by three bond pairs. So, three bond pairs adopt a trigonal structure which is situated in one plane. So, BF₃ has trigonal geometry with FBF bond angle equal to 120°. The lewis structure is:

 

Election pairs = 3.

Bond pairs = 3

Lone pairs = 0

CH₄, NH₃, H₂0 all involve four pairs of electrons around the central atom and should have a tetrahedral structure like CH₄with bond angle 109.5, but as the lone pair-bond pair repulsion> bond pair-bond pair repulsion, the bond angle in NH₃ and H₂O is decreased to 107^o48^’ and 104.5^o respectively

 

4.

It is due to the presence of lone pair of electrons in NF_3. As the lone pair bond pair repulsion is more than bond pair bond pair repulsion, NF₃ is compressed and has a bond angle of 103^o rather than 120^o .

 

5.

a.

In water molecule, there are two lone pairs of electrons in oxygen. In NH₃ molecule, there is one lone pair of electron in nitrogen. According to VSEPR theory we have, lone – pair – lone pair repulsion > bond pair – bond pair repulsion. Since, the no. of lone pairs are more in oxygen, the O – H bond pairs come more closer than N – H bond pairs in NH₃. Thus, bond angle in H₂O (1o4°.27’) is less than in NH₃ (107o48′).

b.

 In water molecule, there is two lone pair of electrons in oxygen. Due to small size of oxygen, the lone pair – lone pair and the bond – pair – bond – pair repulsion is large and as a result, the electrons tend to push the molecules apart. However, the sulphur atom is greater in size and the lone pair tends to be far apart, thereby considerably reducing the repulsions and also the bond angle that is the reason why H₂O bond angle is greater than H₂S

 

6.

Postulates of valence bond theory

 Overlapping of two half filled valence orbitals of two different atoms result in the formation of covalent bond. Due to overlapping, the electron density between two bonded atoms increases and this gives stability to the molecule.

In case the atomic orbitals possess more than one unpaired electron, more than one bond can be formed and electrons paired in the valence shell cannot take part in such a bond formation.

Covalent bond is directional and it is parallel to the region of overlapping atomic orbitals.

Based on pattern of overlapping, there are two types of covalent bonds: sigma bond and pi bond. The covalent bond formed by sidewise overlapping of atomic orbitals is known as pi bond whereas the bond formed by overlapping of atomic orbital along the inter nucleus axis is known as sigma bond.

 

7.

Sigma (σ) bond	Pie (π) bond
1. Sigma bond is formed by the overlapping of half field orbital of one atom with half field orbital of another atom along the intermolecular axis.	1. Pie bond is formed by the side wiseoverlapping of half field orbital of another atom perpendicular to the intermolecular axis.
2. Overlapping takes between s – s orbital, s – p orbital and p – p orbital.	2. Overlapping takes place between only p – p orbital (py – py), (p2 – p2)
3. Sigma molecular orbital is symmetrical about the intermolecular axis.	3. π molecular orbital is discontinuous because two pie bonds are formed above the intermolecular axis.

 

 

 

8.

The shape of BeF₂ is linear structure with bond angle 180° with sp hybridized, BF₃ have sp² hybridization and similarly CH₄ have sp³ hybridization.

 

9.

In methane , the four electron pair repel each other and are directed toward the four corners of regular tetrahedron with bond angle 109.5° .Since four electron pairs are equivalent and form  covalent bond with hydrogen atoms .

 

10.

Hybridization is the process of mixing of dissimilar atomic orbitals of same atoms giving rise to equal number of a new set of orbitals having same energy contents.

The process of mixing of dissimilar atomic orbitals of same atoms giving rise to equal number of a new set of orbitals having same energy is known as hybrid orbital.

 

11.

 The mode of hybridization of the central atom whose molecular geometry is trigonal pyramid is sp³d hybridization. E.g: PCl₅, PF₅, etc.

 

12.

a. C of C₂H₄ – sp² hybridization.

b. B of BF₃ – sp² hybridization.

 c. N in NH₃   - sp³ hybridization.

d. C of C₂H₂ – sp hybridization.

 

 

13.

Sp hybrid orbital

 

b. shape of sp²  orbital

 

 

Sp³  hybrid orbital

 

14.

Hybridization is the process of mixing of dissimilar atomic orbitals of same atoms giving rise to equal number of a new set of orbitals having same energy contents.

 

15.

Limitations of valence bond theory:

1. This theory doesn’t take into account the effect of lone pair electrons during overlap between two orbitals. It can’t explain formation of co – ordinate covalent bond.

2. This theory doesn’t explain magnetic properties of compounds. For eg, it can’t explain paramagnetic behavior of oxygen molecule.

3. This theory fails to explain bonding in electron deficient compounds.

 

16.

In water molecule, there are two lone pairs of electrons in oxygen. In NH₃ molecule, there is one lone pair of electron in nitrogen. According to VSEPR theory we have, lone – pair – lone pair repulsion > bond pair – bond pair repulsion. Since, the no. of lone pairs are more in oxygen, the O – H bond pairs come more closer than N – H bond pairs in NH₃. Thus, bond angle in H₂O (1o4°.27’) is less than in NH₃ (107o48′).

The sulphur atom is greater in size and the lone pair tends to be far apart, thereby considerably reducing the repulsions and also the bond angle (92.5° )

 

17.

The central atom has Sp³ mode of the hybridization of  tetrahedral  structure E.g. CH₄,CCl₄, etc.

 

18.

Orbital picture of ethane

Orbital picture of ethene

 

 

19.

i. A central atom is located at the center with four substituent’s that are located at the corners of a tetrahedron.

ii. The bond angle is 109° 28’

 

20.

 

21.

Geometry of methane CH₄

In CH₄ carbon is the neutral atom. The electronic configuration of carbon (z = 6) is 1s²,2s²,2p². It consists of four valence electrons which forms four bond pairs with four hydrogen atoms. Now, the central carbon atom is surrounded by four electron pairs, which are situated at the four corners of the regular tetrahedral. Hence, CH₄ molecule has tetrahedral geometry with HC + l bond angle is 109.5.

 

22.

In the case of ethylene the sp² hybrid orbital of one carbon atom overlaps basically with similar sp² hybrid orbital of another C – atom, forming C – C sigma bond. The remaining two sp² hybrid orbital of each carbon overlaps with its orbital of two hydrogen atoms forming two C – H sigma bonds.

 

23.

It’s characters are:

(i) It has sp³ hybridization.

(ii) Its bond angle is 120° and 90°.

 

24.

 If a central atom in a molecule has only one bond pair it has regular geometry and if the central atom has more lone pair, molecule gets distorted to same extent giving rise to irregular geometry to the molecule. Thus, the presence of lone pair of electron causes the decrease in the bond angle and hence causes the certain distortion in regular geometry of the molecule than we except.

 

 

25.

1. In molecular orbital theory, molecular orbitals are formed by linear combinations of atomic orbitals (LACO) method.

2. Atomic orbitals of the resulting molecule lose their individual identities.

3. Molecular orbitals are poly – centric

4. The number of the molecular orbitals formed is equal to the number of the atomic orbitals theta undergo the combination.

 

26.

Bonding molecular orbitals	Antibonding molecular orbitals
1. It is formed by the addition overlap of atomic orbital.	1. It is formed by the subtraction overlap of atomic orbital.
2. It has more electron density in the region between the two nuclei and this accounts for the stability of bond.	2. It has less electron density in the region between the two nuclei and this leads to unstability of bond.
3. It may or may not have a nodal plane.	3. It always has a nodel plane between the nuclei.

 

27.

Bond order of H₂   =   $\frac{1}{2}{\rm{\: }}\left[ {{\rm{\: }}{{\rm{N}}_{\rm{b}}} - {\rm{\: }}{{\rm{N}}_{\rm{a}}}} \right]$   = $\frac{1}{2}{\rm{\: }}\left[ {2 - 0} \right] = 1$

Bond order of He₂   =   $\frac{1}{2}{\rm{\: }}\left[ {{\rm{\: }}{{\rm{N}}_{\rm{b}}} - {\rm{\: }}{{\rm{N}}_{\rm{a}}}} \right]$   = $\frac{1}{2}{\rm{\: }}\left[ {2 - 2} \right] = 0$

Since the disassociation energy of the molecule is directly proportional to the bond order, H₂ has greater dissociation energy hence more stable than He₂.  Bond order is zero for the He₂   means molecule is unstable and doesn’t exists.

 

28.

 In O₂ molecule, atoms are held by a double covalent bond. It may be noticed that it contains two unpaired electron. So, O₂molecule has paramagnetic nature. But in case of N₂ molecule, atoms are held by triple bond. Hence, the nitrogen molecule is highly stable and is confirmed by its high bond dissociation energy and small bond length. Since, there are no unpaired electron in any orbital, N₂ molecule is diamagnetic.

 

29.

Bond order is defined as the half of the difference in number of electrons between the bonding molecular orbital and anti bonding molecular orbital.

The bond order of O₂ is:

BOO2=$\frac{1}{2}$(Nb−Na)

= $\frac{1}{2}$ (10 – 6)

= 4/2 = 2

Similarly the bond order of N₂ is:

BON2=$\frac{1}{2}$ (Nb−Na)

= $\frac{1}{2}$ (10−4)

= 6/2

= 3.

 

 

30.

sp³ hybridization:

In this case, one s- and three p-orbitals hybridize to form four sp³ hybrid orbitals. These four sp³-hybrid orbitals are oriented in a tetrahedral arrangement. The common example of molecule involving sp³-hybridisation is methane (CH₄). Therefore, CH₄ has tetrahedral geometry and HCH bond angle is 109.5^o.

 

31.

 

32.

a. BF₃ à sp² hybridization with geometry of triangular planes and bond angle 120°.

b. BeCl₃ à sp hybridization with geometry of linear and bond angle 180°.

c. NCl₃ à sp³ hybridization with geometry normal tetrahedral and bond angle 109°

d. CCl₄ à sp³ hybridization with geometry normal tetrahedral and bond angle 109°

 

33.

Valence bond theory	Molecules orbital theory
1. In valence bond theory, two atomic orbitals give an inter atomic orbital obtained by space filling of two unpaired electrons one being in each of the two atomic orbitals.	1. In molecular orbital theory, molecular orbitals are formed by linear combinations of atomic orbitals (LACO) method.
2. The resulting molecule, consists of atoms and retain their individual character.	2. Atomic orbitals of the resulting molecule lose their individual identities.
3. Atomic orbitals are monocentric.	3. Molecular orbitals are poly – centric.

 

 

Short answer questions

1.

Hybridization is the process of mixing of dissimilar atomic orbitals of same atoms giving rise to equal number of a new set of orbitals having same energy contents.

Features of hybridization:

(i) Hybridization is the process of mixing of energies not the electron.

(ii) The hybrid orbitals have shape and same energy.

(iii) The number of Hybrid orbital formed is equal to the number of atomic orbitals mixing together.

(iv)The hybrid orbitals are situated in fire position in the space giving a definite geometrical shape.

(v) Only the bigger lobe takes part in overlapping.

 

 

2.

sp³ hybridization:

In this case, one s- and three p-orbitals hybridize to form four sp³ hybrid orbitals. These four sp³-hybrid orbitals are oriented in a tetrahedral arrangement. The common example of molecule involving sp³-hybridisation is methane (CH₄). Therefore, CH₄ has tetrahedral geometry and HCH bond angle is 109.5^o.

 

           Fig(i) sp hybridization

           Fig(ii) sp2 hybridization

           Fig(iii) sp3 hybridization

 

Long answer questions

1.

 Difference between bond theory and molecular orbital theory:

Valence bond theory	Molecules orbital theory
1. In valence bond theory, two atomic orbitals give an inter atomic orbital obtained by space filling of two unpaired electrons one being in each of the two atomic orbitals.	1. In molecular orbital theory, molecular orbitals are formed by linear combinations of atomic orbitals (LACO) method.
2. The resulting molecule, consists of atoms and retain their individual character.	2. Atomic orbitals of the resulting molecule lose their individual identities.
3. Atomic orbitals are monocentric.	3. Molecular orbitals are poly – centric.

 

 

The molecules contain more than one electron, the exact wave function for molecules orbitals cannot be obtained. For obtaining the wave function for molecular orbital approximately applied is that the molecular orbitals are linear combination of atomic orital. This approximate method is known as linear combination of atomic orbital (LCAO) method.

The electronic configuration of oxygen is 1s²,2s²,2p⁴. Total electrons in O₈ molecule is 16. The electronic configuration of O₂ molecule, therefore is:

O₂ : (61s)²(6 * 1s)²(62s)²(6 * 2s)²(62p_x)²(π2p_y)²(π2p_z)²(π2p_y)¹(π * 2p_z)¹

Bond orders = Nb−Na2

= (10−6)2 = 42 = 2.

In O₂ molecule, atoms are held by a double covalent bond. It maybe noticed that it contains two unpaired electron in π * 2p_y and π * 2p_z molecular orbitals. So, O₂ molecule has paramagnetic nature.

 

 

 

2.

Valence bond theory of H₂ molecule:

When two hydrogen atoms are at large separation from each other, so that there is no orbital overlap, the total energy is equal to the sum of the energies to two H – atoms. When two atoms come close to each other then attractive and repulsive force began to operate. As the attractive force exceeds the repulsive force, orbitals start to overlap and potential energy of the system begins to decrease. At a certain distance between the two atoms, the attractive and repulsive interaction because each other and energy of the system attains minimum value (which is 433) or H₂ molecule. The internuclear separation at which the orbitals overlap to the proper extent and possessing minimum energy called bond energy.

Limitations of valence bond theory:

1. This theory doesn’t take into account the effect of lone pair electrons during overlap between two orbitals. It can’t explain formation of co – ordinate covalent bond.

2. This theory doesn’t explain magnetic properties of compounds. For eg, it can’t explain paramagnetic behavior of oxygen molecule.

3. This theory fails to explain bonding in electron deficient compounds