All of Valence Bond Theory Explained

Veence bond theory explains how a

coalent bond is actually formed. To

understand it deeply, first we need to

learn the main postulates of veence bond

theory. The first and most important

postulate says that a coalent bond is

formed by the overlapping of atomic

orbitals of two atoms. We must

understand that orbitals refer to the

space or region around the nucleus where

the probability of finding an electron

is maximum. In simple words, we can say

that an orbital is a region where an

electron is present around the nucleus.

When orbitals overlap with each other

they actually share electrons and form a

coalent bond which we will explain in

detail. Moving toward the second

postulate only those atomic orbitals

will overlap which contain unpaired

electrons. We can understand this with

the example of hydrogen. It has only one

electron and its electronic

configuration is 1 s1. As we can see in

this case the s orbital does not have

two electrons in it. So it is called an

unpaired electron. But keep in mind only

those electrons will overlap that

contain opposite spins. Moving toward

the third postulate, it explains that

the extent of overlapping determines the

strength of the bond. This means that

the strength of overlapping is directly

proportional to the strength of the

bond. If overlapping is more, the bond

will be stronger and if the extent of

overlapping is less, then the bond will

definitely be weak. The fourth and most

important postulate says that a coalent

bond formed between two atoms will be

directional. This means that the

direction of overlapping orbitals will

determine whether a sigma or pi bond

will be formed. To make it clear, let's

understand it deeply. Actually, atomic

orbitals can overlap in two ways. One is

by parallel or head-to-head overlap, and

the second one is perpendicular or side

to side overlap. To clarify, let's take

the overlapping of pxpx orbitals and py

orbitals. As we know, p orbitals have a

dumbbell shape and one px orbital will

overlap with another px orbital in a

parallel way. So the bond formed will be

a sigma bond since the orbitals are

overlapping in a parallel manner. But if

we take the overlapping of py orbitals,

they will overlap side to side or

perpendicularly resulting in the

formation of a pi bond. One important

point to remember is that every time an

s orbital overlaps with another s

orbital, it will form a sigma bond.

Similarly, whenever a px orbital

overlaps with another px orbital, it

will also form a sigma bond. Now we will

take the example of a hydrogen molecule

to understand these postulates. As we

know hydrogen has only one electron and

its veence electron is present in the 1

s orbital and is actually an unpaired

electron. Similarly the second hydrogen

atom will have the same configuration

and its electron is also present in the

1 s orbital. So in order to form a bond

these two orbitals will overlap with

each other. As I mentioned above,

whenever an s orbital overlaps with

another s orbital, they will overlap in

a parallel or head-to-head way. So this

means that in this case, a sigma bond

will be formed between two hydrogen

atoms. One thing we need to remember is

that between any two atoms, a sigma bond

will always form first and after the

formation of a sigma bond, a pi bond

will be formed later. Also remember that

these orbitals form a bond at that

position where the forces of attraction

dominate over the forces of repulsion.

We can take one more example to

understand the formation of a sigma

bond. Let's take the example of an HF

molecule. Here we can see that one atom

is hydrogen and the other atom is

florine. Since we know the valence

electron of hydrogen is present in the

1s orbital, we now need to find out the

valence electron of florine. By looking

at the electronic configuration of

florine, we can see that its veence

electron is present in p orbitals and

the unpaired electron is present in the

2pz orbital. So according to the

postulates of valence bond theory the

unpaired electrons of both atoms should

overlap with each other. So we can say

that in this case the 1s orbital of

hydrogen will overlap with the 2pz

orbital of florine. Always remember that

whenever an s orbital overlaps with any

other orbital whether it's a p orbital

or another s orbital it will always form

a sigma bond. So we can say that s and p

orbitals will overlap in a head-to-head

or parallel way and will form a sigma

bond. And as a result of this

overlapping the HF molecule will be

formed. Now let's take one example for

the formation of a pi bond. So we can

also understand it. Let's take the

example of a nitrogen molecule. As we

know nitrogen always forms a triple bond

with another nitrogen in the case of an

N2 molecule. Here one bond is sigma and

two bonds are pi bonds. First let's have

a look at the electronic configuration

of nitrogen. Its atomic number is seven

and veence electrons are present in the

p orbital. As we can see, nitrogen has

three unpaired electrons, which again

confirms that nitrogen needs to form

three bonds. Similarly, the second

nitrogen atom will also have the same

electronic configuration in order to

form a bond. First of all, the two px

orbitals of one nitrogen atom will

overlap with the two px orbitals of the

second nitrogen atom. Since px orbitals

always overlap in a parallel way, we can

say that a sigma bond will be formed by

the overlapping of px orbitals from each

nitrogen atom. So one bond is formed out

of a total of three bonds. After that

the 2p orbital of one nitrogen atom will

overlap with the 2p orbital of the

second nitrogen in a side to side or

perpendicular way resulting in the

formation of a pi bond. Again we need to

remember that after the sigma bond is

formed each additional bond formed will

be a pi bond and orbitals will overlap

in a side to side manner. Actually up to

this point two bonds are formed between

nitrogen atoms where one is sigma and

the second bond is a pi bond. For the

third bond formation, the two pz

orbitals of each nitrogen will overlap

in a side to side or perpendicular way

resulting in the formation of the third

bond. Now we can see that all three

bonds which were required to be formed

between nitrogen atoms are now complete

with one sigma and two pi bonds present.

After understanding how valance orbitals

overlap to form a bond, we need to

understand some main differences between

sigma and pi bonds. First of all, we can

define a sigma bond in this way. The

type of bond that is formed by the

head-to-head overlap of atomic orbitals

where the probability of finding the

electrons is maximum along the line

joining the nuclei. Similarly, a pi bond

is formed by the side to side overlap of

atomic orbitals where the probability of

finding electrons is maximum above and

below the line joining the nuclei. Now

moving toward the second point, a sigma

bond is stronger due to the direct

overlap of orbitals. But a pi bond is

weaker than a sigma bond because it is

formed by side to side overlap which is

less effective than head-to-head

overlap. Moving toward the third

difference. Remember that free rotation

around a sigma bond is possible because

electron density is maximum along the

line joining the nuclei. However, free

rotation around a pi bond cannot happen

due to the above and below probability

of finding electrons. Here free rotation

will only happen when pi bond breaking

takes place. A sigma bond is present in

every bond formation. However, a pi bond

is present only in double and triple

bonds. Now we will discuss concept of

hybridization in veence bond theory.

Actually hybridization is a concept in

veence bond theory that helps explain

the shapes of molecules. According to

VBT, atomic orbitals combine to form

bonds. But sometimes the observed shapes

of molecules do not match the expected

shapes based on pure atomic orbitals.

Hybridization occurs when atomic

orbitals of similar energy mix together.

The number of hybrid orbitals formed is

always equal to the number of atomic

orbitals that mix. These hybrid orbitals

then arrange themselves in a way that

minimizes repulsion between electrons

leading to the actual shapes of

molecules. There are three common types

of hybridization such as sp3, sp2 and sp

hybridization. In case of methane carbon

is sp3 hybridized. So here 2 s orbital

and 32p orbitals will intermix. These

orbitals mix to form four sp3 hybrid

orbitals. The four sp3 orbitals arrange

themselves in a tetrahedral shape with

bond angles of 109.5°.

In this case, each hybrid orbital forms

a sigma bond with a hydrogen atom. Now

moving towards sp2 hybridization. In

this case, one s and two p orbitals

intermix to form three sp2 type of

hybridized orbitals. We can understand

it with example of ethine. In ethine,

each carbon atom forms three sigma

bonds. Two with hydrogen and one with

another carbon. One 2 s and two 2p

orbitals of carbon mix to form three sp2

hybrid orbitals. The remaining

unhybridized p orbital forms a pi bond

between the two carbon atoms. The sp2

hybrid orbitals arrange themselves in

agonal planer shape with bond angles of

120°. Now let's move towards sp

hybridization. In this type of

hybridization on s and 1 p orbital

intermix to form two sp hybridized

orbitals. We will take example of

ethine. In ethine each carbon atom forms

two sigma bonds one with hydrogen and

one with another carbon. One 2 s and one

two ep orbital of carbon mix to form two

sp hybrid orbitals. The remaining two

unhybridized p orbitals form two pi

bonds between the carbon atoms. The sp

hybrid orbitals arrange themselves in a

linear shape with bond angles of 180

degrees. Now let's discuss limitations

of valance bond theory. It is useful for

explaining how atoms form bonds. But it

has some limitations. One major problem

is that it does not explain deoized

bonding. For example, in benzene, the

six carbon atoms form a ring with

alternating single and double bonds.

However, experiments show that all

carbonarbon bonds in benzene are

actually equal in length, meaning the

electrons are not fixed between two

atoms, but are deoized over the entire

ring. VBT cannot properly describe this

deoization. Another limitation is that

VBT does not explain magnetic and

spectral properties of molecules

accurately. It assumes that electrons

are localized in specific bonds, which

does not match experimental data for

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