We are discussing the basics of VSEPR theory related to molecules with lone pairs. In the previous section 4.20, we completed the discussion on the last Case IV. In this section, we will see the official definition and the various features of the theory
• The VSEPR theory was first put forward by Sidwick and Powell in 1940
• The theory helps us to predict the shapes of covalent molecules
• The procedure is very simple. It is based on the interactions between electron pairs in the valence shell of atoms
• Let us first explain what this 'interactions between electron pairs in the valence shell of atoms' is. It can be written in 7 steps:
1. There are so many electrons in an atom. We need not consider all of them
• We consider only those ‘electrons in the valence shells of the atoms’
• We know that, ‘electrons in the valence shell of an atom’ are called valence electrons of that atom
• We already know the method to find the 'number of valence electrons' of any atom
2. Consider an atom A in a molecule
• Since A is part of a molecule, it would be bonded with other atom/atoms
• Consider the valence electrons of A
• Those valence electrons can be classified into two categories:
(i) Bonded electrons
(ii) Non-bonded electrons
Let us see some examples:
Example 1:
• Fig.4.122(a) below, shows the Lewis dot structure of BH3
• We see 3 red dots around B
• They are the valence electrons of Boron (Recall that, we show only the valence electrons in Lewis dot structures)
• All those 3 electrons have entered into bonding
■ Electrons which have 'entered into bonding' are called bonded electrons
Example 2:
• Fig.4.122(b) shows the Lewis dot structure of NH3
• We see 5 red dots around N
• They are the valence electrons of Nitrogen
• Out of the five valence electrons, three have entered into bonding
♦ Two electrons have not entered into any bonding
■ Electrons which have 'not entered into bonding' are called non-bonded electrons
3. Now we know 'bonded electrons' and 'non-bonded electrons'. The next two items that we have to see are:
♦ Bonded electron pair
♦ Non-bonded electron pair
They can be easily explained:
(i) Bonded electron pair:
• Consider any 'bonded electron'. It will be always present at one end of a '—'
(Recall that '—' indicates a bond)
• In a '—', there will be two electrons, one at each end
♦ Both those electrons are 'bonded electrons'
• Since there are two electrons in a bond, we can call it 'a pair of electrons'
♦ To be precise: 'a pair of bonded electrons'
■ So we can write:
♦ Whenever we see a '—', we are looking at a Bonded electron pair
♦ In 'ball and stick models', we represent the '—' using sticks
✰ So a 'stick' represent a bonded electron pair
(ii) Non-bonded electron pair:
• We have seen what 'non-bonded electrons' are
♦ Such electrons are always seen in pairs
■ So we can write:
♦ Non-bonded electrons are always seen in 'groups of two'
♦ Such a group is called a: Non-bonded electron pair
✰ A non-bonded electron pair is also called a: lone pair
4. Next we want to see what is meant by 'interactions'
We can easily guess. There will be 4 types of interactions:
(i) A 'bonded electron pair' will interact with any other 'bonded electron pair' which is close by
(ii) A 'bonded electron pair' will interact with any other 'non-bonded electron pair' which is close by
(iii) A 'non-bonded electron pair' will interact with any other 'bonded electron pair' which is close by
(iv) A 'non-bonded electron pair' will interact with any other 'non-bonded electron pair' which is close by
But (iii) is same as (ii). So we can write:
■ There will be 3 types of interactions:
(i) A 'bonded electron pair' will interact with any other 'bonded electron pair' which is close by
(ii) A 'bonded electron pair' will interact with any other 'non-bonded electron pair' which is close by
(iii) A 'non-bonded electron pair' will interact with any other 'non-bonded electron pair' which is close by
5. But why do they interact?
• The answer is simple: They interact because they are all negatively charged
• We know that, like charges repel each other
(We can expect a condition of 'no interaction', only if the particles are 'charge less'. If two charged particles are brought close to each other, interaction will definitely occur)
6. So we can write:
The interactions are all repulsive
7. Note that, the interactions are between 'pairs'
We do not consider the interactions between individual electrons
• The VSEPR theory is based on these interactions
• Now we will see the VIII postulates of the theory and their explanations
I. The shape of a molecule depends upon the number of valence shell electron pairs (bonded or non-bonded) around the central atom
Explanation of this postulate can be written in 2 steps:
(i) First we have to count the number of electron pairs (in the valence shell of the central atom)
♦ We have to count all the bonded pairs
♦ We have to count all the lone pairs
♦ Then we take the sum
(ii) The shape of the molecule will depend on this sum
Application of this postulate can be written in 2 steps:
(i) We have already seen the application. Let us see two examples: (a) and (b)
(a) AB4
♦ Number of bonded pairs = 4
♦ Number of lone pairs = 0
♦ Sum = (4+0) = 4
• So there are 4 items around A. The basic shape will be tetrahedral
(b) AB3E
♦ Number of bonded pairs = 3
♦ Number of lone pairs = 1
♦ Sum = (3+1) = 3
• So there are 4 items around A. The basic shape will be tetrahedral
(ii) The 'sum' which is the 'total number of items around A' helps us to decide the basic shape
II Pairs of electrons in the valence shell repel one another since their electron clouds are negatively charged.
Explanation of this postulate:
This postulate do not need much explanation. We obviously know that, there will be repulsion between like charges. So the pairs repel one another
Application of this postulate:
This postulate has tremendous application. The various shapes attained by the various molecules is a consequence of these repulsions
III These pairs of electrons tend to occupy such positions in space that minimize repulsion and thus maximize distance between them
Explanation of this postulate can be written in 3 steps:
(i) Consider some particles which experience repulsion between each other
♦ Naturally, each particle will try to push the other particles ‘as far away as possible’
(ii) ‘As far away as possible’ can be achieved only by ‘increased distances’ between the particles
♦ When the distance increases, the particles begin to experience lesser repulsion from each other
(iii) So the particles will try to achieve ‘maximum possible distances’ from each other
Application of this postulate can be written in 2 steps:
(i) Tendency to achieve ‘maximum possible distances’ will obviously influence the final positions of the atoms
(ii) Final positions of the atoms will give the shape of the molecules
(iii) In the final positions, the atoms will be 'as far way from each other as possible'
IV The valence shell is taken as a sphere with the electron pairs localizing on the spherical surface at maximum distance from one another
Explanation of this postulate can be written in 4 steps:
(i) We have seen that, in the final positions, the terminal atoms will be 'as far way from each other as possible'
(ii) All terminal atoms will be equidistant from the central atom A
(iii) This is similar to the definition of a sphere:
♦ The sphere will have a center point
♦ All points on the surface of the sphere will be equidistant from the center
(iv) In the case of a molecule:
♦ The central atom is the center of the sphere
♦ All terminal atoms will lie on the surface of a sphere
♦ This is because, the terminal atoms are equidistant from the central atom
♦ The surface of the sphere can be considered as the valence shell of the central atom
Application of this postulate:
• The application can be explained with the help of an example. It can be written in 3 steps:
(i) Consider a molecule with the tetrahedral shape
(ii) The center of the tetrahedron will be the center of a sphere
(iv) The four terminal atoms will lie exactly on the surface of that sphere
V A multiple bond is treated as if it is a single electron pair and the two or three electron pairs of a multiple bond are treated as a single super pair
Explanation of this postulate can be written in 4 steps:
(i) The central atom is bonded to a number of terminal atoms
• The bonds may be single, double or triple
(ii) In the VSEPR theory:
♦ A double bond is considered as a single bond
♦ A triple bond is also considered as a single bond
(iii) So we can write:
• In the VSEPR theory,
♦ A '=' is considered as a '—'
♦ A '≡' is also considered as a '—'
(iv) But what about the pairs?
A '=' will contain two pairs. How can we consider them as '—' ?
A '≡' will contain three pairs. How can we consider them as '—' ?
■ The solution is that:
• The ‘two pairs’ in a '=' is considered as a single pair
♦ Not just an ‘ordinary single pair’
♦ But a ‘single super pair’
• The ‘three pairs’ in a '≡' is also considered as a single pair
♦ Not just an ‘ordinary single pair’
♦ But a ‘single super pair’
Application of this postulate:
• The application can be explained with the help of an example. It can be written in 2 steps:
(i) In CO2, there are two double bonds
(ii) But while applying the VSEPR theory, we considered them as single bonds
(See fig.4.82 of section 4.13)
VI Where two or more resonance structures can represent a molecule, the VSEPR model is applicable to any such structure
Explanation of this postulate can be written in 3 steps:
(i) We know that, some molecules can be represented by two or more resonance structures (Details here)
(ii) The VSEPR theory is applicable to any of those structures
(iii) We will get the same result because, VSEPR theory treats double and triple bonds as single bonds
Application of this postulate:
• The application can be explained with the help of an example. It can be written in 3 steps:
(i) The resonance structures of CO2 can be seen in fig.4.51 in section 4.9
(ii) In all the structures, C is the central atom and the two O are the terminal atoms
(iii) We can apply the theory to any one of those structures. The result will be a linear shape
(See fig.4.82 of section 4.13)
VII The repulsive interactions between electron pairs decrease in the order:
(lp-lp) > (lp-bp) > (bp-bp)
• That means:
♦ The force of repulsion between two lone pairs will be the greatest
♦ The force of repulsion between two bond pairs will be the least
♦ The force of repulsion between a lone pair and a bond pair will have an intermediate value
The explanation for this postulate was given by Nyholm and Gillespie in 1957. It can be written in 3 steps:
(i) The bond pairs have a definite space within the molecule
♦ This is because, a bond pair will be belonging to two atoms
♦ So that bond pair will lie on the line between the two owner atoms
(ii) But a lone pair belongs to only one atom (the central atom)
♦ So they are spread out into a greater space than bond pairs
♦ In other words, the lone pairs occupy a greater space than bond pairs
(iii) So the lone pairs are able to apply a greater ‘push’ on others
♦ That is the reason why we get: (lp-lp) > (lp-bp) > (bp-bp)
Application of this postulate:
• The application can be explained with the help of an example. It can be written in 3 steps:
(i) We used the following double headed arrows:
♦ Cyan double headed arrow to indicate lp-lp repulsion
♦ Yellow double headed arrow to indicate lp-bp repulsion
♦ Red double headed arrow to indicate bp-bp repulsion
(ii) We know that:
♦ Cyan is stronger than yellow
♦ Yellow is stronger than red
(iii) Now consider the shape of H2O molecule
(See fig.4.103 in section 4.17)
• If all the arrows in the fig.4.103(a) are of the same color, there will be a perfect symmetry
♦ The angle will not become lesser than 109.5
• That means, using the same colored arrows, we will not be able to explain the lesser angle in a water molecule
• That is., if all repulsions are considered to be of the same magnitude, we will not be able to explain the lesser angle in a water molecule
(iv) In this way, this postulate has tremendous applications in explaining the shape of various molecules
• It explains why the actual shape deviates from the expected shape
♦ For example, in the case of water:
✰ It explains why the actual angle is less than the expected value of 109.5o
VIII For predicting the shape of a molecule using VSEPR theory, it is convenient to divide the molecules into two categories:
(i) Molecules in which central atom has no lone pair
(ii) Molecules in which central atom has one or more lone pairs
Explanation of this postulate:
• Whenever we apply the VSEPR theory, we have to first look whether the central atom has lone pairs or not
• Method of application will be different for the two categories:
♦ Molecules in which central atom has no lone pair
♦ Molecules in which central atom has one or more lone pairs
Application of this postulate:
• We have already seen the application of this postulate
• We have applied the theory separately for the two categories
(i) Molecules in which central atom has no lone pair
♦ were discussed in sections 4.13, 4.14, 4.15 and 4.16
(ii) Molecules in which central atom has one or more lone pairs
♦ were discussed in sections 4.17, 4.18, 4.19 and 4.20
We have seen the VIII postulates. Now we will see some solved examples
Solved example 4.7
Discuss the shape of the following molecules using the VSEPR model:
BeCl2, BCl3, SiCl4, AsF5, H2S, PH3
Solution:
(i) The Lewis dot structure of BeCl2 is shown in fig.4.123(a) below:
• The central atom Be has no lone pairs. So it is of the type AB2
• Molecules of the type AB2 have a linear structure (See fig.4.82 in section 4.13)
• So BeCl2 is linear
(ii) The Lewis dot structure of BCl3 is shown in fig.4.123(b) above
• The central atom B has no lone pairs. So it is of the type AB3
• Molecules of the type AB3 have a trigonal planar structure (See fig.4.84 in section 4.13)
• So BCl3 is trigonal planar
(iii) The Lewis dot structure of SiCl4 is shown in fig.4.124(a) below:
• The central atom Si has no lone pairs. So it is of the type AB4
• Molecules of the type AB4 have a tetrahedral structure (See fig.4.82 in section 4.14)
• So SiCl4 is tetrahedral
(iv) The Lewis dot structure of AsF5 is shown in fig.4.124(b) above
• The central atom As has no lone pairs. So it is of the type AB5
• Molecules of the type AB5 have a trigonal bipyramidal structure
• So AsF5 is trigonal bipyramidal
(v) The Lewis dot structure of H2S is shown in fig.4.125(a) below:
• The central atom S has two lone pairs. So it is of the type AB2E2
• Molecules of the type AB2E2 have a bent shape
• So H2S is of bent shape
(vi) The Lewis dot structure of PH3 is shown in fig.4.125(b) above
• The central atom P has one lone pair. So it is of the type AB3E
• Molecules of the type AB3E have a bent shape
• So PH3 is of trigonal pyramidal shape
Solved example 4.8
Although geometries of NH3 and H2O molecules are distorted tetrahedral, bond angle in water is less than that of ammonia. Discuss
Solution:
The solution can be written in 5 steps:
1. Shapes of both H2O and NH3 are obtained from the basic shape: Tetrahedron
2. In the tetrahedron, the angle is 109.5o. But:
♦ In NH3, the angle is 107o
♦ In H2O, the angle is 104.5o
• We want to know why the angles are different from 109.5o
• We also want to know why the angle in H2O is lesser than that in NH3
3. Diagrams that we used in our discussions:
♦ We obtained the structure of H2O using the fig.4.103 in section 4.17
♦ We obtained the structure of NH3 using the fig.4.106 in section 4.18
4. In fig.106, we see that, there is one lone pair in N atom
♦ This gives rise to lp-bp repulsions
• The lp-bp repulsion (yellow) is stronger than bp-bp repulsion (red)
♦ So the red arrows will get compressed
■ Thus the basic angle of 109.5o will decrease to 107o in NH3
5. In fig.103, we see that, there are two lone pairs in H atom
• This gives rise to lp-lp repulsions (in addition to lp-bp repulsions)
• The lp-lp repulsion (cyan) is stronger than lp-bp repulsion (yellow)
• The lp-bp repulsion (yellow) is stronger than bp-bp repulsion (red)
♦ So the cyan will compress the yellows
♦ The yellows will further compress the red
✰ As a result, the red is compressed more
■ So the angle reduces to 104.5o in H2O
• The VSEPR theory was first put forward by Sidwick and Powell in 1940
• The theory helps us to predict the shapes of covalent molecules
• The procedure is very simple. It is based on the interactions between electron pairs in the valence shell of atoms
• Let us first explain what this 'interactions between electron pairs in the valence shell of atoms' is. It can be written in 7 steps:
1. There are so many electrons in an atom. We need not consider all of them
• We consider only those ‘electrons in the valence shells of the atoms’
• We know that, ‘electrons in the valence shell of an atom’ are called valence electrons of that atom
• We already know the method to find the 'number of valence electrons' of any atom
2. Consider an atom A in a molecule
• Since A is part of a molecule, it would be bonded with other atom/atoms
• Consider the valence electrons of A
• Those valence electrons can be classified into two categories:
(i) Bonded electrons
(ii) Non-bonded electrons
Let us see some examples:
Example 1:
• Fig.4.122(a) below, shows the Lewis dot structure of BH3
• We see 3 red dots around B
• They are the valence electrons of Boron (Recall that, we show only the valence electrons in Lewis dot structures)
• All those 3 electrons have entered into bonding
■ Electrons which have 'entered into bonding' are called bonded electrons
 |
| Fig.4.122 |
• Fig.4.122(b) shows the Lewis dot structure of NH3
• We see 5 red dots around N
• They are the valence electrons of Nitrogen
• Out of the five valence electrons, three have entered into bonding
♦ Two electrons have not entered into any bonding
■ Electrons which have 'not entered into bonding' are called non-bonded electrons
3. Now we know 'bonded electrons' and 'non-bonded electrons'. The next two items that we have to see are:
♦ Bonded electron pair
♦ Non-bonded electron pair
They can be easily explained:
(i) Bonded electron pair:
• Consider any 'bonded electron'. It will be always present at one end of a '—'
(Recall that '—' indicates a bond)
• In a '—', there will be two electrons, one at each end
♦ Both those electrons are 'bonded electrons'
• Since there are two electrons in a bond, we can call it 'a pair of electrons'
♦ To be precise: 'a pair of bonded electrons'
■ So we can write:
♦ Whenever we see a '—', we are looking at a Bonded electron pair
♦ In 'ball and stick models', we represent the '—' using sticks
✰ So a 'stick' represent a bonded electron pair
(ii) Non-bonded electron pair:
• We have seen what 'non-bonded electrons' are
♦ Such electrons are always seen in pairs
■ So we can write:
♦ Non-bonded electrons are always seen in 'groups of two'
♦ Such a group is called a: Non-bonded electron pair
✰ A non-bonded electron pair is also called a: lone pair
4. Next we want to see what is meant by 'interactions'
We can easily guess. There will be 4 types of interactions:
(i) A 'bonded electron pair' will interact with any other 'bonded electron pair' which is close by
(ii) A 'bonded electron pair' will interact with any other 'non-bonded electron pair' which is close by
(iii) A 'non-bonded electron pair' will interact with any other 'bonded electron pair' which is close by
(iv) A 'non-bonded electron pair' will interact with any other 'non-bonded electron pair' which is close by
But (iii) is same as (ii). So we can write:
■ There will be 3 types of interactions:
(i) A 'bonded electron pair' will interact with any other 'bonded electron pair' which is close by
(ii) A 'bonded electron pair' will interact with any other 'non-bonded electron pair' which is close by
(iii) A 'non-bonded electron pair' will interact with any other 'non-bonded electron pair' which is close by
5. But why do they interact?
• The answer is simple: They interact because they are all negatively charged
• We know that, like charges repel each other
(We can expect a condition of 'no interaction', only if the particles are 'charge less'. If two charged particles are brought close to each other, interaction will definitely occur)
6. So we can write:
The interactions are all repulsive
7. Note that, the interactions are between 'pairs'
We do not consider the interactions between individual electrons
• Now we have a basic idea about what the 'interactions between electron pairs in the valence shell of atoms' is
• Now we will see the VIII postulates of the theory and their explanations
I. The shape of a molecule depends upon the number of valence shell electron pairs (bonded or non-bonded) around the central atom
Explanation of this postulate can be written in 2 steps:
(i) First we have to count the number of electron pairs (in the valence shell of the central atom)
♦ We have to count all the bonded pairs
♦ We have to count all the lone pairs
♦ Then we take the sum
(ii) The shape of the molecule will depend on this sum
Application of this postulate can be written in 2 steps:
(i) We have already seen the application. Let us see two examples: (a) and (b)
(a) AB4
♦ Number of bonded pairs = 4
♦ Number of lone pairs = 0
♦ Sum = (4+0) = 4
• So there are 4 items around A. The basic shape will be tetrahedral
(b) AB3E
♦ Number of bonded pairs = 3
♦ Number of lone pairs = 1
♦ Sum = (3+1) = 3
• So there are 4 items around A. The basic shape will be tetrahedral
(ii) The 'sum' which is the 'total number of items around A' helps us to decide the basic shape
II Pairs of electrons in the valence shell repel one another since their electron clouds are negatively charged.
Explanation of this postulate:
This postulate do not need much explanation. We obviously know that, there will be repulsion between like charges. So the pairs repel one another
Application of this postulate:
This postulate has tremendous application. The various shapes attained by the various molecules is a consequence of these repulsions
III These pairs of electrons tend to occupy such positions in space that minimize repulsion and thus maximize distance between them
Explanation of this postulate can be written in 3 steps:
(i) Consider some particles which experience repulsion between each other
♦ Naturally, each particle will try to push the other particles ‘as far away as possible’
(ii) ‘As far away as possible’ can be achieved only by ‘increased distances’ between the particles
♦ When the distance increases, the particles begin to experience lesser repulsion from each other
(iii) So the particles will try to achieve ‘maximum possible distances’ from each other
Application of this postulate can be written in 2 steps:
(i) Tendency to achieve ‘maximum possible distances’ will obviously influence the final positions of the atoms
(ii) Final positions of the atoms will give the shape of the molecules
(iii) In the final positions, the atoms will be 'as far way from each other as possible'
IV The valence shell is taken as a sphere with the electron pairs localizing on the spherical surface at maximum distance from one another
Explanation of this postulate can be written in 4 steps:
(i) We have seen that, in the final positions, the terminal atoms will be 'as far way from each other as possible'
(ii) All terminal atoms will be equidistant from the central atom A
(iii) This is similar to the definition of a sphere:
♦ The sphere will have a center point
♦ All points on the surface of the sphere will be equidistant from the center
(iv) In the case of a molecule:
♦ The central atom is the center of the sphere
♦ All terminal atoms will lie on the surface of a sphere
♦ This is because, the terminal atoms are equidistant from the central atom
♦ The surface of the sphere can be considered as the valence shell of the central atom
Application of this postulate:
• The application can be explained with the help of an example. It can be written in 3 steps:
(i) Consider a molecule with the tetrahedral shape
(ii) The center of the tetrahedron will be the center of a sphere
(iv) The four terminal atoms will lie exactly on the surface of that sphere
V A multiple bond is treated as if it is a single electron pair and the two or three electron pairs of a multiple bond are treated as a single super pair
Explanation of this postulate can be written in 4 steps:
(i) The central atom is bonded to a number of terminal atoms
• The bonds may be single, double or triple
(ii) In the VSEPR theory:
♦ A double bond is considered as a single bond
♦ A triple bond is also considered as a single bond
(iii) So we can write:
• In the VSEPR theory,
♦ A '=' is considered as a '—'
♦ A '≡' is also considered as a '—'
(iv) But what about the pairs?
A '=' will contain two pairs. How can we consider them as '—' ?
A '≡' will contain three pairs. How can we consider them as '—' ?
■ The solution is that:
• The ‘two pairs’ in a '=' is considered as a single pair
♦ Not just an ‘ordinary single pair’
♦ But a ‘single super pair’
• The ‘three pairs’ in a '≡' is also considered as a single pair
♦ Not just an ‘ordinary single pair’
♦ But a ‘single super pair’
Application of this postulate:
• The application can be explained with the help of an example. It can be written in 2 steps:
(i) In CO2, there are two double bonds
(ii) But while applying the VSEPR theory, we considered them as single bonds
(See fig.4.82 of section 4.13)
VI Where two or more resonance structures can represent a molecule, the VSEPR model is applicable to any such structure
Explanation of this postulate can be written in 3 steps:
(i) We know that, some molecules can be represented by two or more resonance structures (Details here)
(ii) The VSEPR theory is applicable to any of those structures
(iii) We will get the same result because, VSEPR theory treats double and triple bonds as single bonds
Application of this postulate:
• The application can be explained with the help of an example. It can be written in 3 steps:
(i) The resonance structures of CO2 can be seen in fig.4.51 in section 4.9
(ii) In all the structures, C is the central atom and the two O are the terminal atoms
(iii) We can apply the theory to any one of those structures. The result will be a linear shape
(See fig.4.82 of section 4.13)
VII The repulsive interactions between electron pairs decrease in the order:
(lp-lp) > (lp-bp) > (bp-bp)
• That means:
♦ The force of repulsion between two lone pairs will be the greatest
♦ The force of repulsion between two bond pairs will be the least
♦ The force of repulsion between a lone pair and a bond pair will have an intermediate value
The explanation for this postulate was given by Nyholm and Gillespie in 1957. It can be written in 3 steps:
(i) The bond pairs have a definite space within the molecule
♦ This is because, a bond pair will be belonging to two atoms
♦ So that bond pair will lie on the line between the two owner atoms
(ii) But a lone pair belongs to only one atom (the central atom)
♦ So they are spread out into a greater space than bond pairs
♦ In other words, the lone pairs occupy a greater space than bond pairs
(iii) So the lone pairs are able to apply a greater ‘push’ on others
♦ That is the reason why we get: (lp-lp) > (lp-bp) > (bp-bp)
Application of this postulate:
• The application can be explained with the help of an example. It can be written in 3 steps:
(i) We used the following double headed arrows:
♦ Cyan double headed arrow to indicate lp-lp repulsion
♦ Yellow double headed arrow to indicate lp-bp repulsion
♦ Red double headed arrow to indicate bp-bp repulsion
(ii) We know that:
♦ Cyan is stronger than yellow
♦ Yellow is stronger than red
(iii) Now consider the shape of H2O molecule
(See fig.4.103 in section 4.17)
• If all the arrows in the fig.4.103(a) are of the same color, there will be a perfect symmetry
♦ The angle will not become lesser than 109.5
• That means, using the same colored arrows, we will not be able to explain the lesser angle in a water molecule
• That is., if all repulsions are considered to be of the same magnitude, we will not be able to explain the lesser angle in a water molecule
(iv) In this way, this postulate has tremendous applications in explaining the shape of various molecules
• It explains why the actual shape deviates from the expected shape
♦ For example, in the case of water:
✰ It explains why the actual angle is less than the expected value of 109.5o
VIII For predicting the shape of a molecule using VSEPR theory, it is convenient to divide the molecules into two categories:
(i) Molecules in which central atom has no lone pair
(ii) Molecules in which central atom has one or more lone pairs
Explanation of this postulate:
• Whenever we apply the VSEPR theory, we have to first look whether the central atom has lone pairs or not
• Method of application will be different for the two categories:
♦ Molecules in which central atom has no lone pair
♦ Molecules in which central atom has one or more lone pairs
Application of this postulate:
• We have already seen the application of this postulate
• We have applied the theory separately for the two categories
(i) Molecules in which central atom has no lone pair
♦ were discussed in sections 4.13, 4.14, 4.15 and 4.16
(ii) Molecules in which central atom has one or more lone pairs
♦ were discussed in sections 4.17, 4.18, 4.19 and 4.20
Solved example 4.7
Discuss the shape of the following molecules using the VSEPR model:
BeCl2, BCl3, SiCl4, AsF5, H2S, PH3
Solution:
(i) The Lewis dot structure of BeCl2 is shown in fig.4.123(a) below:
 |
| Fig.4.123 |
• Molecules of the type AB2 have a linear structure (See fig.4.82 in section 4.13)
• So BeCl2 is linear
(ii) The Lewis dot structure of BCl3 is shown in fig.4.123(b) above
• The central atom B has no lone pairs. So it is of the type AB3
• Molecules of the type AB3 have a trigonal planar structure (See fig.4.84 in section 4.13)
• So BCl3 is trigonal planar
(iii) The Lewis dot structure of SiCl4 is shown in fig.4.124(a) below:
 |
| Fig.4.124 |
• Molecules of the type AB4 have a tetrahedral structure (See fig.4.82 in section 4.14)
• So SiCl4 is tetrahedral
(iv) The Lewis dot structure of AsF5 is shown in fig.4.124(b) above
• The central atom As has no lone pairs. So it is of the type AB5
• Molecules of the type AB5 have a trigonal bipyramidal structure
• So AsF5 is trigonal bipyramidal
(v) The Lewis dot structure of H2S is shown in fig.4.125(a) below:
 |
| Fig.4.125 |
• Molecules of the type AB2E2 have a bent shape
• So H2S is of bent shape
(vi) The Lewis dot structure of PH3 is shown in fig.4.125(b) above
• The central atom P has one lone pair. So it is of the type AB3E
• Molecules of the type AB3E have a bent shape
• So PH3 is of trigonal pyramidal shape
Solved example 4.8
Although geometries of NH3 and H2O molecules are distorted tetrahedral, bond angle in water is less than that of ammonia. Discuss
Solution:
The solution can be written in 5 steps:
1. Shapes of both H2O and NH3 are obtained from the basic shape: Tetrahedron
2. In the tetrahedron, the angle is 109.5o. But:
♦ In NH3, the angle is 107o
♦ In H2O, the angle is 104.5o
• We want to know why the angles are different from 109.5o
• We also want to know why the angle in H2O is lesser than that in NH3
3. Diagrams that we used in our discussions:
♦ We obtained the structure of H2O using the fig.4.103 in section 4.17
♦ We obtained the structure of NH3 using the fig.4.106 in section 4.18
4. In fig.106, we see that, there is one lone pair in N atom
♦ This gives rise to lp-bp repulsions
• The lp-bp repulsion (yellow) is stronger than bp-bp repulsion (red)
♦ So the red arrows will get compressed
■ Thus the basic angle of 109.5o will decrease to 107o in NH3
5. In fig.103, we see that, there are two lone pairs in H atom
• This gives rise to lp-lp repulsions (in addition to lp-bp repulsions)
• The lp-lp repulsion (cyan) is stronger than lp-bp repulsion (yellow)
• The lp-bp repulsion (yellow) is stronger than bp-bp repulsion (red)
♦ So the cyan will compress the yellows
♦ The yellows will further compress the red
✰ As a result, the red is compressed more
■ So the angle reduces to 104.5o in H2O
• In the next section, we will see valence bond theory
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