58. Covalent Bonding(4)- VESPR model(2).

In the VSEPR model, the valence shell is considered a sphere. The electron pairs are considered to be on this spherical surface, as far away as possible from one another. The maximum distance between the electron pairs ensures that there is minimum repulsion between them.

In our last post, we discussed the geometries of molecules without lone pair of electrons. However, due to the presence of non-bonding / lone pair of electrons, molecules depart from ideal electronic geometry/bond angles. Let us discuss how these non-bonding pairs affect molecular geometries.


 The effect of non-bonding/lone pair of electrons.

The VSEPR model predicts that the presence of lone pair of electrons in a molecule results in a deviation from ideal geometries/ bond angles. Such deviations can be explained by the electron domain concept. Electron domains are the regions in space occupied by an electron pair.


The reasons for deviation from ideal geometries are –

  1. Non-bonding/lone pairs occupy more space than bonding pairs –
    The bonding pair of electrons are a part of a bond between two atoms. Thus, they are diffused through the orbitals of both A and B and between the nuclei of both atoms. Naturally, they are away from the nucleus of one single atom and are equally attracted by both nuclei. This leads to a decrease in the electron domain size.
    A lone pair of electrons is on a single atom and is only attracted by a single nucleus. So having a positive nucleus only at one end keeps the electron domain “fatter” as compared to that of a bonding pair domain.
581

2. The intensity of repulsion between different electron pairs is different.

It is as follows –
Lone pair – Lone pair > Lone pair – Bonding pair > Bonding pair – Bonding pair electrons, i.e

LP – LP > LP- BP > BP – BP.


As the lone pairs occupy more space, the electron domains of two lone pairs on a single atom come close to each other.Thus, the repulsion between them is more. This is true for LP-BP interactions too.The BP electron domains are the farthest apart from each other and thus there is minimum repulsion between them.
582.jpg

3. Repulsions occurring at angles more than 90 °are not significant.

4. The bonding pairs of the electronegative substituents occupy less space than the electropositive ones.
This is because these electrons are strongly attracted by the electronegative atom. Thus, the repulsion between two BP’s becomes less and so the bond angle becomes less.

Let us understand this with an example.

Molecule

Bond angle(H-N-H)

NH3

107.2

NF3

102.3

588

In the above example, fluorine is more electronegative than hydrogen.
So,  BP of electrons between N – F has attracted more towards the F atom → repulsion between N-F bonding electron pairs decreases → the N-F bonds come closer to each other as a  result of the decrease in repulsion between them → Bond angle decreases.

Think of these electron pairs as individuals! When the hatred(repulsion) between two people decreases, they automatically come closer to each other, and thus the bond angle decreases! Atoms, human beings… it’s all the same, isn’t it?

5. Electron pairs in a filled shell repel stronger than electron pairs in the incomplete/vacant shell.

When there is an vacant shell (of similar energy) in an atom, the electrons from the filled shell can undergo diffusion to these orbitals. The electron pairs in a vacant shell repel less. Hence, the LP – BP repulsion diminishes dramatically. Thus, the bond angles around the central metal decrease.

Less repulsion between LP -BP → the bonds come closer → decrease in bond angle.

Look at the examples below-

Example 1 gives us the difference between the bond angles in ammonia(NH3) and phosphine(PH3).
The central element in ammonia is nitrogen. N (7) 1s2 2s2 2p3.
The nitrogen atom does not have any d-orbitals and so the 2p electrons have no chance to go to any other orbital.
However, the phosphorous atom has vacant 3d orbitals. P(15) – 1s2 2s2 2p6 3s2 3p3 3d0
The energies of 3p and 3d orbitals have the same energy. Thus, the 3p electrons can easily diffuse to the 3d orbitals. This decreases the repulsive forces and leads to a decrease in the bond angle in phosphine.

Example 2

The central atom is oxygen. O(8) 1s2 2s2 2p4
Oxygen does not have any vacant orbitals of the same energy.

Sulfur however has them. S (16) 1s2 2s2 2p6 3s2 3p4 3d0

6. Triple bonds repel other bonding electrons more strongly than double bonds and double bonds repel other bonding electrons more strongly than single bonds.


Based on the above facts let us now study some examples, where lone pair affects the geometry of the molecules –

Deviation from tetrahedral geometry /109.5 °bond angle – We studied in the earlier post, that four electron domains exhibit a tetrahedral geometry with a bond angle of 109.5°. So, methane has a bond angle of 109.5°. However, when lone pairs are present in a molecule, the measure of this ideal bond angle (109.5 °) changes and so does the geometry of the molecules.
e.g.- Ammonia has a pyramidal shape(bond angle = 107.3 °) with one lone pair of electron domains and water molecule has a bent shape( bond angle = 104.5 ° ) as it has  2 lone pairs of electron domains. This change in geometry is due to the presence of one and two lone pairs of electrons on ammonia and water respectively. The LP occupies the vertices of the tetrahedron.


586.jpg

As the lone pair is ‘fatter’, it occupies more space, and thus it pushes the bonding pair domains closer to one another. Thus, the bond angle and geometry change.
(Imagine three slim people comfortably sitting in a car’s rear seat. What happens when one fat person replaces a slim one?  The two thin people are pushed closer to make room for this fatter person. A similar thing happens to the electron domains too).

The following table shows the deviations from the ideal geometries –

Steric
number(SN)
Molecular geometry
No lone pairs
Molecular geometry
1 lone pair
Molecular geometry
2 lone pairs
Molecular geometry
3 lone pairs
2AX2E0-2D.png
Linear
(CO2)
   
3AX3E0-side-2D.png
Trigonal planar (BCl3)
AX2E1-2D.png
V- shape
(SO2)
  
4AX4E0-2D.png
Tetrahedral
(CH4)
AX3E1-2D.png
Triagonal Pyramidal
(NH3)
AX2E2-2D.png
Bent
(H2O)
 
5AX5E0-2D.png
Triagonal bipyramidal
(PCl5)
AX4E1-2D.png
Seesaw
(SF4)
AX3E2-2D.png
T-shape
(ClF3)
AX2E3-2D.png
Linear
(I3)
6AX6E0-2D.png
Octahedral (SF6)
AX5E1-2D.png
Square pyramidal
(BrF5)
AX4E2-2D.png
Square planar
(XeF4)
 
7AX7E0-2D.png

 

Pentagonal bipyramidal
(
IF7)

AX6E1-2D.png
Pentagonal pyramidal
(XeOF)
AX5E2-2D.png
Pentagonal planar
(XeF5 ion)
 

The above chart shows how the geometry changes with the introduction of one and two lone pairs of electrons on the central atom.

Molecule
type
ShapeOrientation in space showing LP(yellow color)Orientation in space shown without LPExamples
AX2E0LinearAX2E0-3D-balls.pngLinear-3D-balls.pngBeCl2,HgCl2,
CO2
AX2E1BentAX2E1-3D-balls.pngBent-3D-balls.pngNO2,SO2, O3,
CCl2
AX2E2BentAX2E2-3D-balls.pngBent-3D-balls.pngH2O
AX2E3LinearAX2E3-3D-balls.pngLinear-3D-balls.pngXeF2
AX3E0Trigonal planarAX3E0-3D-balls.pngTrigonal-3D-balls.pngBF3
AX3E1Trigonal pyramidalAX3E1-3D-balls.pngPyramidal-3D-balls.pngNH3, PCl3
AX3E2T-shapedAX3E2-3D-balls.pngT-shaped-3D-balls.pngClF3, BrF3
AX4E0TetrahedralAX4E0-3D-balls.pngTetrahedral-3D-balls.pngCH4
AX4E1SeesawAX4E1-3D-balls.pngSeesaw-3D-balls.pngSF4
AX4E2Square planarAX4E2-3D-balls.pngSquare-planar-3D-balls.pngXeF4
AX5E0Trigonal bipyramidalTrigonal-bipyramidal-3D-balls.pngTrigonal-bipyramidal-3D-balls.pngPCl5
AX5E1Square pyramidalAX5E1-3D-balls.pngSquare-pyramidal-3D-balls.pngClF5, BrF5,
XeOF4
AX5E2Pentagonal planarAX5E2-3D-balls.pngPentagonal-planar-3D-balls.pngXeF5
AX6E0OctahedralAX6E0-3D-balls.pngOctahedral-3D-balls.pngSF6, WCl6
AX6E1Pentagonal pyramidalAX6E1-3D-balls.pngPentagonal-pyramidal-3D-balls.pngXeOF5
AX7E0Pentagonal bipyramidalAX7E0-3D-balls.pngPentagonal-bipyramidal-3D-balls.pngIF7

Steric numbers of 7 or greater are less common.


In our next post, we continue the discussion on VSEPR and study some examples. Till then,

Be a perpetual student of life and keep learning…

Good day!

References and Further reading –

1.https://en.wikipedia.org/wiki/VSEPR_theory

2.http://nptel.ac.in/courses/104103069/14

Image source –

1.https://en.wikipedia.org/wiki/Bent_molecular_geometry

2.https://www.google.co.in/search?q=pyramidal+shape&source=lnms&tbm=isch&sa=X&ved=0ahUKEwjJ35Lnw9zaAhXIrY8KHR32AFAQ_AUICigB&biw=1200&bih=603#imgrc=vp9rJ4wir5z3BM:

3.https://en.wikipedia.org/wiki/VSEPR_theory

4.By Benjah-bmm27 – Own work, Public Domain, https://commons.wikimedia.org/w/index.php?curid=2035653

5.http://nptel.ac.in/courses/104103069/14

Leave a Reply

Fill in your details below or click an icon to log in:

WordPress.com Logo

You are commenting using your WordPress.com account. Log Out /  Change )

Twitter picture

You are commenting using your Twitter account. Log Out /  Change )

Facebook photo

You are commenting using your Facebook account. Log Out /  Change )

Connecting to %s