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VSEPR Model and Shape of Molecular Compounds

Valence Shell Electron Pair Repulsion (VSEPR) Model provides a simple method to predict the shapes of simple covalent molecules or polyatomic ions. This model was developed earlier by Sidgwick and Powell in 1940, and it was further improved by Gillespie and Nyholm in 1957. It is primarily based upon the fact that in a polyatornic molecule the direction of bonds around the central atom depends upon the total number of electron pairs (bond pairs as well as lone pairs) in its valence shell. These electron pairs place themselves as far apart as possible in space so as to have minimum repulsive interactions between them. The minimum repulsions correspond to the state of minimum energy and maximum stability of the molecule. The main points of this model are:


  • The geometry and shape of the molecule depend upon the number of electron pairs (bonded as well as nonbonded) in the valence shell of the central atom. The geometrical arrangement of different numbers of electron pairs around the central atom A is shown in Table 9.2.


  • The electron pairs surrounding the central atom repel one another as their orbitals/electron clouds are negatively charged.


  • In order to minimise repulsion, these electron pairs tend to occupy such positions in space where the distance between them is maximum. In other words, if valence shell of central atom is taken as sphere, the electron pairs are localised on this spherical surface at maximum distance from one another.



Table 9,2. Number of Electron Pairs and their Geometrical Arrangement

• The repulsive interactions between two lone pairs (lp) are different from those between two bond pairs (bp) or those between a lone pair and a bond pair. The repulsive interactions between various electron pairs decrease in order as:

lp-lp > lp-bp > bp-bp

Nyholm and Gillespie postulated that bond pair of electrons is shared by the two atoms whereas lone pair is under the influence of only central atom. Hence, in a molecule, the electron cloud containing lone pair is more spread out and occupy more space as compared to the electron cloud

containing bond pair. This causes relatively greater repulsive interactions between the lone pairs in comparison to lone pair bond pair and bond pair-bond pair repulsive interactions.


  • When all the electron pairs around the central atom are bond pairs only and the surrounding atoms are similar, the molecule is said to possess regular geometry. Table 9.3 gives the geometry of molecules in which central atom does not have any lone pair.


  • When lone pairs are also present in addition to bond pairs around the central atom. the repulsive interactions between the electron pairs around the central atom become unequal. This causes distortion in geometrical arrangement of electron pairs. Such a molecule is said to have an irregular or distorted geometry. Table 9.4 gives the shapes of some simple molecules and ions in which central atom has one or more lone pairs along with bond pairs.


  • It may be noted that for the purpose of applying VSEPR theory, a multiple bond is treated as a single electron pair. For example in methanol (formaldehyde) molecule (HCHO), the bond pairs around carbon atom are counted as three (2 bp for two C-H bonds and 1 bp for C = O bond).  similarly two, C = O bonds in CO2 molecule are counted as two bond pairs. The shapes of HCHO and CO2 are shown below

Table 9.3. Geometry of Molecules in which Central Atom has No Lone Pair of Electrons (here central atom is A and surrounding atoms are B)


Number of

electron pairs



Arrangement of electron pairs Examples
        2 AB2 Linear   

Bef2 ,BeCl2, HgCl2, CO2
        3 AB3 Trigonal planar







       4 AB4 Tetrahedral

      5 AB5 Trigonal bipyamidal                 



In this arrangement, three of the bond pairs lie in the same plane at an angle of 120°. These are marked as e and are called equatorial pairs. The other two bond pairs lie at right angle to the plane of equatorial pairs. These are called axial pairs

and are marked as a.

  • Axial pairs in this arrangement experience relatively larger repulsive interactions. Hence axial bonds are relatively longer than equatorial bonds.
  • • Any lone pair appearing in this arrangement tend to occupy equatorial positions due to relatively lower repulsive interactions.
PF5,PCl5, SbCl5, AsF5
     6 AB6 Octahedral

SF6, TeF6




Table 9.4. Geometry/Shapes of Molecules containing Bond Pairs and Lone Pairs (here A is central atom, B is surrounding atom, B is lone pair)



Example 9.l. Describe the shapes of following molecules on the basis of VSEPR Model


BeCl2, BF3 CH4, NH3, H2O+, CO2‘ SF

Solution. The Lewis dot structures and shapes of various species are given below in tabular form:


Example 9.2 Describe the bonding in methanal (formaldehyde) in tem1s of sigma and pi bonding

solution .  The electron dot structure of methanal (C2O) is

In methanal, carbon atom assumes sp2 hybrid state (see Figs. 9.13 and 9.14).


The two sp-hybrid orbitals of carbon atom overlap with sp2 orbital of each of the two H atoms to form two C- H σ bonds. The third sp2-hybrid orbital of carbon atom overlaps axially with similar orbital of oxygen atom to form sp2-sp2 sigma bond. The unhybridized 2p orbital of carbon overlap with similar orbital of oxygen sidewise and forms pπ-pπ bond. The 1t clouds and also sigma cloud have more concentration of electron density around oxygen atom due to its greater electronegativity.