Chapter 3: – Chemical Bonding
B.sc 1st year Book (Page 12)
Valence-shell of electron pair repulsion theory full detail or VSEPR theory
Vsepr theory was introduced by Gillespie and Nyholm in 1957. According to these scientists, electron pairs present in the valence shell of a central atom decide the actual structure of the molecule. In molecules, the central atom consists of two types of electron pairs; σ-bonding electron pairs (bp’s) and non-bonding electron pairs (or Ip’s). The central atom arranges these bp’s and Ip’s around itself in such a way that they are as far away from each other as possible so that minimum repulsion occurs between them and the molecule has maximum stability l. θ. the molecule has minimum energy. Thus, the “VSEPR model is an extremely powerful method for predicting molecular structure”. The following rules have been proposed by the above scientists to explain the molecular structure of covalent molecules.
VSEPR theory Rule 1:
If in a covalent molecule, the central atom is surrounded only by bp’s and not by Ip’s then the geometry of the molecule is regular as it is generally explained by the concept of hybridization and Sidgwick’s concept of electronic theory, i.e. it will be linear, triangular planar, tetrahedral, trigonal bipyramidal, regular octahedral and pentagonal bipyramidal for 2, 3, 4, 5, 6 and 7bp’s respectively. Example :
- BeCl2 ( b ‘s =2, lp=0, linear)
- BF3 (bp′s=3, lp=0, triangular planar)
- CH4 (bp ‘s =4, lp=0, tetrahedral)
- PCl5 (bp′s=5, lp=0 triangular bipyramidal)
If. (bp′s=6, p=0 regular octahedral) and IF7 (bp′s=7, k=0, pentagonal bipyramidal).
VSEPR theory Rule 2:
When the central atom of the covalent molecule is surrounded by both bp’s and lp’s then the ape of the molecule will be irregular. It is because repulsion between IP-p or IP-bp on the central atom effectively influences the geometry. In fact, it is found that repulsion between =−p is greater than those between the lp-bp and bp-bp. The order of repulsive energies of Ese electron pairs is:
∣p−Ip > Ip−bp > bp−bp
Reasons:
The lone pair electrons are under the influence of only one nucleus, they would be expected a localize with a greater electron density radially close to the central atom than the bonding Fectran pairs which are under the influence of two nuclei i.e. the bonding electron cloud is polarized in between the central atom and bonded atom. Since lp does not have a second nucleus is, therefore, tends to occupy a greater angular volume.
Related Topic | Chemical Bonding |
Thus, Ip is much closer to the central atom than bp and so it is believed that IPs will exert -more repulsion on any adjacent electron pair than a bp will exert on the same adjacent s strong pair. Consequently, the greater the repulsion between the electron pairs, the greater will – the contraction in the bond angle of the molecule.
Example-1 :
| Molecule : | CH4 > | NH3 > | H2O |
| No. of bp’s and Ip’s : | (bp’s = 4, Ip = 0) | (bp’s = 3, Ip = 1) | (bp’s = 2, Ip = 2) |
| Bond angle : | 109º28′ | 107.2º | 104.5º |
| Hybridization : | sp3 | sp3 | sp3 |
Example-2 :
| Molecule : | PCl5 > | SF4 > | BrF3 |
| No. of bp’s and Ip’s : | (bp’s = 5, Ip = 0) | (bp’s = 4, Ip = 1) | (bp’s = 3, Ip = 2) |
| Bond angle : | 120º, 180º | 103º, 179º | 86º, 172º |
| Hybridization : | sp3d | sp3d | sp3d |
Rule 3:
In a covalent molecule of the type AXn, the X-A-X bond angle decreases with the increase in electronegativity of bonded atom X. If ‘A’ is the central atom. For example :
| Molecule : | Pl3 > | Plbr3 > | PCl3 > | PF3 > |
| Bond angle : | 102º | 101º | 100.3º | 97.7º |
Reasons:
The decrease in the X−A−X bond angle is mainly due to the increase in the ionic character of the A−X bond. The more electronegative bonded atom X attracts the bonding electron pairs away from the central atom’ ‘ A ‘ and allows the lone pair to expand while the X−A−X angle closes. Reduced bond angles in non-metal fluorides are commonly observed. For the small atoms of nitrogen and oxygen, where VSEPR interactions seem to be especially important, the fluorides have smaller band angles than the hydrides. e.g.
(i) NF3=102.3∘, NH3=107,2∘
(ii) OF2=103.1∘, H2O=104.5∘
(ii) OF2=103.1∘, H2O=104.5∘
VSEPR theory Rule 4 :
Bond angles involving multiple bonds are generally larger than those involving only single bonds. However, the multiple bonds do not alter the geometry of the molecule.
For example :
For example :
| Molecule | : | CH3CH3 | < H2C−CH2 | ⩽ | HC≡CH |
| Bond angle | : | 109’28’ | 120′ | 180′ | |
| Shape | : | tetrahedral | triangular planar | linear | |
| Hybridization | : | sp3 | sp2 | sp |
VSEPR theory Rule 5 :
Bond angles in the hydrides of the elements belonging to the third or higher periods of a group are smaller than the bond angles in the hydrides of the elements of the second period of the same group. For example; In hydrides of the VIA group we observe that the bond angle decreases on moving from top to bottom as follows :
| Hydrides : | H2O | H2S | H2Se | H2Te |
| Bond angle : | 104.5º | 92.5º | 91.0º | 89.5º |
Reasons:
It is because :
(i) the repulsions between bp’s around the smaller atoms of an element of a group are stronger than between the bigger atoms of the same group.
(ii) since the Ip orbital on a bigger-sized atom occupies much more space than the Ip orbital on a smaller-sized atom of an element of the same group. Therefore, the lp-bp distances are smaller, and hence /p – bp repulsions are stronger in the hydrides of bigger-sized S. Se and Te atoms as compared to the hydride of oxygen atoms i.e. H2O.
Limitations :
1- This theory is unable to explain the shapes of molecules that have extensive delocalized p-electron systems, e.g. C6H6, CH2=CH−CH=CH2, B3N3H6, etc.
2- This theory does not explain the shapes of molecules in which the inert pair effect occurs. e.g. SnCl2,PbCl2,BiCl3 etc:
3- This theory is unable to explain the geometries of transition metal complexes.
4- It can not explain the shapes of molecules that are highly polar. e.g. Li2O.
2- This theory does not explain the shapes of molecules in which the inert pair effect occurs. e.g. SnCl2,PbCl2,BiCl3 etc:
3- This theory is unable to explain the geometries of transition metal complexes.
4- It can not explain the shapes of molecules that are highly polar. e.g. Li2O.