Chemistry > Chemical Bonding and Molecular Structure > 9.0 Valence Shell Electron Pair Repulsion (VSEPR) Theory
Chemical Bonding and Molecular Structure
1.0 Ionic Bond or Electrovalent Bond
2.0 Lattice Energy
3.0 Characteristics of Electrovalent Compounds
4.0 Covalent Bond (By Mutual Sharing of Electrons)
5.0 Characteristics of Covalent Compounds
6.0 Fajan’s Rule
7.0 Hydrogen Bonding
8.0 Coordinate Bond
9.0 Valence Shell Electron Pair Repulsion (VSEPR) Theory
10.0 Valence Bond Theory
11.0 Sigma and Pi Bonds ($\sigma $ and $\pi $ Bonds)
12.0 Hybridisation
12.1 Types of hybridization and spatial orientation of hybrid orbitals
12.2 Method of predicting the Hybrid state of the central atom in covalent molecules of polyatomic ions
13.0 Molecular Orbital Theory
9.3 Some examples using the VSEPR Theory
12.2 Method of predicting the Hybrid state of the central atom in covalent molecules of polyatomic ions
$1.$ Phosphorus pentachloride $PCl_5$:
Gaseous $PCl_5$ is covalent. (The electronic structure $P$ is $1s^22s^22p^63s^23p^3)$. All five outer electrons are used to form bonds to the five $Cl$ atoms.
In the $PCl_5$ molecule the valence shell of the $P$ atom contains five electron pairs. Hence, the structure is a trigonal bipyramid. There are no lone pairs, so the structure is not distorted. However, a trigonal bipyramid is not a completely regular structure, since some bond angels are $90^°$ and others $120^°$.
Symmetrical structures are usually more stable than asymmetrical ones.
Note: Thus $PCl_5$ is highly reactive, and in the solid state it splits into $[PCl_4]^+$ and $[PCl_6]^–$ ions, which have tetrahedral and octahedral structures respectively.
$2.$ Chlorine trifluoride $ClF_3$:
The chlorine atom is at the centre of the molecule and determines its shape. The electronic configuration of $Cl$ is $1s^22s^22p^63s^23p^5$.
Three electrons form bonds to $F$, and four electrons do not take part in bonding. Thus in $ClF_3$, the $Cl$ atom has five electron pairs in the outer shell, hence the structure is a trigonal bipyramid. There are three bond pairs and two lone pairs.
It was noted previously that a trigonal bipyramid is not a regular shape since the bond angles are not all the same. It therefore follows that all the corners are not equivalent. Lone pair occupy two of the corners, and $F$ atoms occupy the other three corners. Three different arrangements are theoretically possible, as shown in figure below.
The most stable structure will be the one of lowest energy, that is the one with the minimum repulsion between the five orbitals. The great repulsion occurs between two lone pairs. Lone pair bond pair repulsions are next strongest, and bond pair-bond pair repulsions are weakest. Groups at $90^°$ repel each other strongly, whilst groups $120^°$ apart repel each other much less.
Structure $I$ is the most symmetrical, but has six $90^°$ repulsions between lone pairs and atoms. Structure $II$ has one $90^°$ repulsion between two lone pairs, plus three $90^°$ repulsions between lone pairs and atoms. These factors indicate that structure $III$ is the most probable. The observed bond angles are $80^°40'$, which is close to the theoretical $90^°$. This confirms that the correct structure is $III$, and the slight distortion from $90^°$ is caused by the presence of the two lone pairs.
As a general rule, if lone pairs occur in a trigonal bipyramid they will be located in the equatorial position (round the middle) rather than the axial positions (top and bottom), since this arrangement minimizes repulsive forces.
$3.$ Sulphur hexafluoride $SF_6$:
The electronic structure of $S$ is $1s^22s^22p^63s^23p^6$. All six of the outer electrons are used to form bonds with the $F$ atoms.
Thus in $SF_6$, the $S$ has six electron pairs in the outer shell. Hence, the structure is octahedral. There are no lone pairs, so the structure is completely regular with bond angles of $90^°$.
