Valence shell Electron Pair Repulsion Theory (VSEPR)

Valence Shell Electron Pair Repulsion (Vsepr) Theory

Table of Contents

Introduction to Valence Shell Electron Pair Repulsion

Theory (Vsepr)

In the year 1957 Gillespie developed a theory to improve the Sidgwick-Powell theory to explain molecular shapes and bond angles more accurately. This theory may be summarized in the following points:

  • Electron pairs tend to minimize repulsions and these are in the order

lone pair-lone pair > lone pair-bond pair > bond pair-bond pair. (Here bond pair refers to a single bond.) Shapes of the molecules depend upon repulsions between bond pair and lone pair electrons.

  • The double bond is in need of more space as compared to the single bond. The repulsion order in relation to the bonds is as follows:

double bond-double bond > double bond-single bond > single bond-single bond. 

  • Keeping the central atom (having lone pair) same, Increase in electronegativity of the associated atom will result in decrease of the bond angle provided no other factors like size and back bonding play any role.
  • If the surrounding atoms are kept same, increase in electronegativity of the central atom (having the lone pair) results in increase of the bond angle.
  • Sometimes the lone pair may be transferred from filled shell of an atom to unfilled shell of the adjacent bonded atom. This phenomenon of transferring electron is known as ‘back bonding’.

In this theory, no distinction is made between s-and p-electrons. Only those electrons which are present in valence shell of the central atom are taken into account. Thus, the number of electron pairs around the central atom decides geometry of a molecule. For Example, if there are two electron pairs around the central atom, the only way to keep them as far apart as possible is to arrange them at an angle of 180° to each other. The molecule in such a case will adopt linear geometry. Similarly, the molecule forms trigonal planar geometry for three electron pairs around the central atom, and for four electron pairs around the central atom, the molecule adopts tetrahedral geometry.

The molecules will form trigonal bipyramidal geometry if they five electron pairs around the central atom. The molecules having six electron pairs around the central atom have octahedral geometry.
The geometries of molecules based on the number of electron pairs is given in table below

Geometries of Molecules

Fig. No. 1 Geometries of Molecules

Let us illustrate this theory by considering a few examples:

Shapes of Molecules containing Bond Pairs only

Shape of BeF2 molecule: Linear

In BeF2, the central Be-atom (Z = 4; 1s22s2) has two electrons in the valence shell. In the formation of BeF2, each of these valence electrons is shared by two fluorine atoms. As a result, the Be atom is surrounded by two bond pairs of electrons [Fig 2]. Therefore, the geometry of BeF2 molecule is linear as shown below and the bond angle is 180°.

Shape of BeF2 moleculeFig. No. 2 Shape of BeF2 molecule

Other molecules such as BeCl2, ZnCl2, HgCl2 have linear shape.

Shape of BF3 molecule: Trigonal planar

The central atom, boron (Z = 5, 1s22s22p1) has three valence electrons. At the time of formation of BF3 molecule, each electron in the valence shell of B-atom forms a bond pair with F-atom. As a result, the central boron atom is surrounded by three bond pairs and the molecule adopts trigonal planar geometry. In this geometry, all the F-B-F bond angles are of 120°. This geometry is planar because the three F-atoms and B-atom lie in the same plane.

Shape of BF3 Fig. No. 3 Shape of BF3

Molecules such as BCl2, AlCl3, etc. have same shape.

Shape of CH4 molecule: Tetrahedral

The central atom in methane that is, carbon (Z = 6, 1s2, 2s2, 2p2) has four valence electrons. All the four valence electrons are bonded to four hydrogen atoms forming four bond pairs around the central carbon atom. These four electron pairs, trying to remain as far apart as possible, adopt tetrahedral structure. In this geometry, all the H-C-H bond angles are of 109°28’ (or approximately 109.5°).

Shape of CH4Fig. No. 4 Shape of CH4

Other examples of tetrahedral molecules are SiF4, CCl4, NH4 etc.

Shape of PCl5 molecule: Trigonal bipyramidal

In PCl5, the central atom, P (Z=15; 1s2, 2s2, 2p6, 3s2, 3p3) has five valence electrons. It forms five bond pairs with five Cl-atoms to form a molecule of PCl5. Since there are five electron pairs around the central phosphorus atom and therefore, it has trigonal bipyramidal geometry. In this geometry, all the bond angles are not equal. Three electron pairs are in the same plane at an angle of 120°, while other two are perpendicular to the plane, both making an angle of 90° with the plane. Thus, in this arrangement three bond angles are of 120° each and two are of 90° each.

Shape of PCl5Fig. No. 5 Shape of PCl5

In this geometry, all five P-Cl bonds are not equal. The three bonds lying in the trigonal plane are called equatorial bonds. Of the remaining two bonds, one lies above and the other below the equatorial plane, both making an angle of 90° with the plane. These bonds are called axial bonds. It has been observed that axial bonds are slightly longer than equatorial bonds in this geometry. PF5 has same shape.

The larger bond length of axial bonds than equatorial bonds can be explained in terms of the repulsive forces between electron pairs due to different bond angles. The axial bond pair faces greater repulsion from other bonds and therefore, the axial bond is slightly longer than equatorial bond.

Axial and Equatorial Bonds in PCl5Fig. No. 6 Axial and Equatorial Bonds in PCl5

It may be noted that the structure of PCl5 molecule is unsymmetrical. As a result, it is less stable and is therefore, highly reactive.

Shape of SF6 molecule: Octahedral

In SF6, the central S-atom (Z = 16; 1s2, 2s2, 2p6, 3s2, 3p4) has six valence electrons. Each of these six valence electrons forms bond with F-atom and therefore, the molecule has octahedral geometry. In this case, all the bond angles are same and are of 90° each. TeF6 molecule has same shape.

Shape of SF6 moleculeFig. No. 7 Shape of SF6 molecule

Other examples of octahedral molecules are SeF6, TeF6 etc.

Shape of IF7 molecule: Pentagonal bipyramidal

In IF7, the central atom I (Z = 53, 1s22s22p63s23p63d104s24p64d105s25p5) has seven valence electrons. Each of these seven valence electrons forms bond with F-atom and therefore, the molecule has pentagonal bipyramidal geometry. In this case, all the bond angles are not equal. Five electron pairs are in the same plane at an angle of 72°, while other two are perpendicular to the plane both making an angle of 90° with the plane.

Thus, in this arrangement five bond angles are of 72° each and two are of 90° each.

Shape of IF7 Fig. No. 8 Shape of IF7

Shapes of Molecules containing Lone Pairs and Bond Pairs

Now, let us consider a few molecules containing bond pairs as well as lone pairs.

Molecules containing three electron pairs (AB3 or AB2L)

If the valence shell of an atom contains three electron pairs, then the molecule has trigonal planar geometry (Example: BF3). However, the geometry gets distorted if it contains a lone pair in addition to bond pair. For Example, a molecule of the type AB2L (where L represents a lone pair), has V-shaped geometry as discussed for SO2 molecule.

Shape of Sulphur dioxide (SO2) molecule

In SO2 molecule, there are three electron pairs (two bond pairs and one lone pair).

The three electron pairs should acquire a trigonal planar arrangement with bond angle 120°. Since one of the positions is occupied by a lone pair, the geometry may be described as angular or V-shaped or bent shape.

Shape of SO2 moleculeFig. No. 9 Shape of SO2 molecule

Now, lone pair-bond pair repulsion is more than bond pair-bond pair repulsion. Therefore, bonded pairs of electrons are pushed closer and the O-S-O bond angle gets reduced to 119° from the value of 120°.

Molecules Containing Four Electron Pairs (AB4, AB3L, AB2L2)

As already learnt, the molecule AB4 has tetrahedral geometry. But if lone pairs are also present in addition to bond pairs, the geometry gets distorted. This may be illustrated by taking two examples:
(a) Molecules containing 3 bp and 1 lp AB3L. e.g. NH3
(b) Molecules containing 2 bp and 2 lp AB2L2 e.g. H2O.

(a) Shape of NH3 molecule: Pyramidal
The central nitrogen atom (z = 7, 1s2, 2s2, 2p3) of NH3 consist of five valence electrons. Hydrogen atoms forms three bond pairs around nitrogen atom and there is one lone pair because of remaining two electrons. Therefore, nitrogen is surrounded by four electron pairs which adopts tetrahedral geometry. But all the four electron pairs around nitrogen are not equivalent as there are three bond pairs and one lone pair and therefore, it has distorted tetrahedral geometry. The bond angle is 107° unlike 109.5° as in tetrahedral geometry. The reason for distortion is the presence of one lone pair in addition to bond pairs. As lone pair-bond pair repulsion is more than bond pair-bond pair repulsion, the repulsion between the lone pair and bond pairs is strong and bond angle decreases to 107°. The geometry of ammonia molecule is also considered as pyramidal (Fig. 13).

Shape of NH3 moleculeFig. No. 10 Shape of NH3 molecule 

other molecules with same shape are PCl3, NF3, H3O+, etc.

(b) Shape of H2O molecule: Bent or angular
The central oxygen atom (Z = 8, 1s2, 2s2, 2p3) of water molecule has six valence electrons. At the time of formation of water molecule,

Shape of H2OFig No. 11 Shape of H2O

the central oxygen atom adopts tetrahedral geometry because of the four electron pairs around it. But all these four electron pairs around O are not the same and therefore geometry of H2O is distorted tetrahedral. The bond angle in water molecule is 104.5° rather than is not of 109.5° (Fig. 11). The distortion is result of repulsion among two lone pairs and the bond pairs. The repulsive force between lone pair-lone pair is greater than the force of repulsion among two bond pairs of electrons. Therefore, the two lone pairs of electrons move away from each other while the two O-H bonds are forced closer to each other which decreases the H-O-H angle to 104.5°. The resulting geometry is considered as bent or angular.H2S, F2O, SCl2, are some other molecules with similar shapes.

KEY NOTE
It may be noted here that the central atoms (C, N and O) in three molecules CH4, NH3 and H2O have four electron pairs around the central atom. Therefore, these molecules adopt tetrahedral geometries. But in methane, there is no lone pair, NH3 molecule has one lone pair while H2O molecule has two lone pairs in the total of four electron pairs. Because of lone pairs, NH3 and H2O molecules will have distorted geometries, while CH4 molecule will be of tetrahedron structure that is, of regular geometry. As larger lone pair-bond pair repulsion than bond pair-bond pair in NH3, the bond angle is reduced from 109.5° to 107°. The geometry of NH3 is pyramidal.

Now, in case of H2O, two lone pairs force the O-H bonds more closely than the N-H bonds in NH3. So the bond angle decreases to a larger extent that is, to 104. 5°. The geometry of water is regarded as V-shaped or angular.

Molecules Containing Five Electron Pairs (AB5, AB4L, AB3L2, AB2L3)

When the central atom is surrounded by five electron pairs, the geometry is trigonal bipyramidal. However, the geometry gets distorted if one or more bond pairs are replaced by lone pairs. This may be illustrated by the following examples:

(a) Molecules containing 4 bp and 1 lp. e.g.  SF4

(b) Molecules containing 3 bp and 2 lp. e.g., ClF3

(c) Molecules containing 2 bp and 3 lp. e.g., XeF2

AB4L (4 bp and 1 lp) Molecules

Let us take the example of SF4. In this case, Sulphur atom (Z = 16: 3s2 3p4) has six valence electrons. In the formation of SF4 four electrons form four bond pairs and leave two electrons as one lone pair. Thus, five electron pairs around Sulphur adopt trigonal bipyramidal geometry in which one position is occupied by lone pair.
Therefore, SF4 molecule can have structure or structure as shown in Figure, in which the lone pair is present on axial or equatorial positions respectively. Nyholm - Gillespie modification has helped in predicting accurately the geometry of such molecules containing lone pair of electrons.

Axial or Equatorial PositionFig. No. 12 Axial or Equatorial Position

In arrangement (a) the lone pair is in on axial position which has 3 lp-bp repulsions at 90°. In structure (b) the lone pair is in on equatorial position and there are only two lp-bp repulsions. Hence (b) will have lesser repulsions and will be stable when compared to arrangement (a). This shape is described as distorted tetrahedron or a folded square or a see-saw. The bond angles in SF4 are 89° and 117° instead of 90° and 120° respectively.

AB3L2 (3 bp 2 lp) molecules

Let us take the example of chlorine trifluoride, ClF3 molecule which is isoelectronic with SF4

Shape of ClF3Fig. No. 13 Shape of ClF3

The central chlorine atom (Z = 17: 3s2 3p5) has seven electrons in its valence shell. In the formation of ClF3, three electrons form three bond pairs and leave four electrons as two lone pairs. Thus, the five electron pairs around chlorine atom adopt trigonal bipyramidal geometry, in which two positions are occupied by lone pairs. As already discussed, the lone pair in trigonal bipyramidal geometry experiences more repulsions at axial positions, therefore, both the lone pairs are present at equatorial positions as shown in Fig. 16. The molecule is T-shaped and bond angle is 87.6° instead of 90°.

AB2L3 (2 bp and 3 lp) molecules

Let us take the example of Xenon difluoride, XeF2 molecule. Xenon atom has (Z = 54: 5s2, 5p6) eight electrons in the valence shell, in this molecule there are two bond pairs and three lone pairs. These five electron pairs forms structure of trigonal bipyramidal geometry with three positions occupied by lone pairs. The net repulsion on the bonds due to lone pairs is zero due to the presence of three lone pairs at the corners of an equilateral triangle. Thus, it has a linear geometry.

Molecules containing six Electron Pairs (AB6, AB5L, AB4L2)

When the central atom is surrounded by six electron pairs, the geometry is octahedral. However, if one or more lone pairs are present in addition to bond pair, the geometry gets distorted. This may be illustrated by the following examples:

(a) Molecules containing 5 bp and 1 lp e.g. BrF5.

(b) Molecules containing 4 bp and 2 lp e.g. XeF5.

AB5L (5 bp and 1 lp) molecules

Consider the example of Bromine Pentafluoride. The central bromine atom (Z = 35, 4s2, 4p5 has seven valence electrons. Br F5 consists of five bond pairs and one lone pair and the six electron pairs forms octahedral geometry out of which one of the positions is occupied by a lone pair. Since all the six positions in octahedral geometry are equivalent, therefore, lone pair may be placed on any position (Fig. 14).

Structure of BrF5Fig. No. 14 Structure of BrF5

The geometry of Br F5 is termed as square pyramidal. IF5 molecule has same geometry.

AB4L4 (4 bp and 2 lp) molecules

Let us take the example of XeF4. In this case, the central xenon atom has eight electrons. XeF4 has six electron pairs and octahedral geometry as
there are four bond pairs and two lone pairs. (Fig. 15), out of which two positions are occupied by lone pairs. The structure is called as Square Planar.

Shape of XeF4Fig. No. 15 Shape of XeF4

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Valence Shell Electron Pair Repulsion (Vsepr) Theory

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