Valence Shell Electron Pair Repulsion Theory

What is the VSEPR Theory?

The Valence Shell Electron Pair Repulsion Theory is often shortened as VSEPR (pronounced “vesper”). It is generally a theory to anticipate the geometry of molecules. Particularly, the VSEPR theory looks at the bonding and molecular geometry of organic molecules and polyatomic ions. It is useful for nearly all substances that have a central atom that is not a metal.

Sidgwick and Powell (1940) explained that the shapes of molecules could be analyzed in terms of electron pairs in the external orbit of the central atom. Nyholm and Gillespie established the VSEPR theory, which discusses the shapes of molecules for non- transition elements.

Basic Assumption

The valence electron pairs (lone pairs and the bond pairs) are set up around the central atom to remain at a maximum range apart to keep repulsions at a minimum.

Postulates of VSEPR Theory

• (i) Both the lone pairs along with the bond pairs participate in figuring out the geometry of the particles.
• (ii) The electron pairs are arranged around the central polyvalent atom so as to remain at a maximum range apart to avoid repulsions.
• (iii) The electron pairs of lone pairs occupy more area than the bond pairs.

A bonding electron pair is brought in by both nuclei of atoms while non-bonding by only one nucleus. Because a lone pair experiences less nuclear attraction, its electronic charge is expanded more in space than that for a bonding pair. As a result, the non-bonding electron pair exerts higher repulsive forces on bonding electron pairs and therefore tends to compress the bond pairs.

The magnitude of repulsions between the electron sets in an offered molecule decreases in the following order:

Lone pair- lone pair > lone pair-bond pair > bond pair-bond pair

These repulsions are called van der Waals repulsions.

• (iv) The two-electron pairs of a double bond and 3 electron pairs of a triple bond, contain a greater electronic charge density. For that reason, they inhabit more space than one electron pair of a single bond but act like a single electron pair in identifying the geometry of the molecule. This is because they tend to occupy the same region between the two nuclei like a single bond.

Shapes of Molecules According to VSEPR Theory

1. Molecules Containing Two-Electron Pairs (AB₂ type)

In such, molecules two electrons, pairs around the central atom are set up at a farther range apart at an angle of 180 °, in order to lessen repulsions between them. Therefore, they form a linear geometry.

Beryllium chloride is a common linear molecule, which includes two electrons pairs. MgCl₂, CaCl₂, SrCl₂, CdCL₂ and HgCl₂ are also linear molecules. The central atoms have 2 electrons in the outermost orbitals.

2. Molecules Containing Three Electron Pairs — (AB₃ type)

(a) AB₃ Type with no Lone Pairs

In such molecules, the central atom includes three bonding electron pairs, which are organized at a maximum distance apart at a shared angle of 120 °, providing a triangular planar geometry. The boron atom in BH₃ is surrounded by three charge clouds, which remain farthest apart in one plane, each pointing towards the corners of an equilateral triangle.

Therefore, BH₃, molecules have a trigonal planar geometry, with each H- B-H bond angles of 120 °. We expect similar geometries in hydrides of group III-A (AlH₃, GaH₃, InH_3, and TlH₃) and their halides (BF₃, AlCl₃, and so on).

(b) AB₃-Type with One Lone Pair and Two Bond Pairs

In SnCl₂, one of the corners of the triangle is inhabited by a lone pair, giving rise to a distorted triangular structure in vapor phase.

(c) AB₃-Type with MultipleBonds

In SO₂, one corner of the triangle is occupied by a lone pair and two corners each by S =O double bond, while in SO₃ all 3 corners, each is occupied by S = O bonds. This structure of SO₃ is completely triangular.

 

3. Molecules Containing Four Electron Pairs (AB₄– Type)

(a) AB₄ Type with no Lone Pairs

The charge clouds due to 4 electron pairs avoid their electrostatic repulsions by drifting apart, so to preserve a mutual bond angle of 109.5 °. Such geometry makes it possible for a kinda shape of the regular tetrahedron.

Examples:

Each of the four valence electrons of carbon pair up with the sole electron of hydrogen in methane.

₆C = 1s², 2s¹, 2px¹, 2py¹, 2py¹

The 4 electron pairs are directed from the center towards the corners of a regular tetrahedron, with each apex representing a hydrogen nucleus. The arrangement permits a non-planar plane of electron pairs. Each H-C-H bond is perfectly 109.5 °. On the very same premises, SiH₄, GeH₄, CCl₄ type similar geometries. This structure has 4 corners, four faces, six edges, and six bond angles.

(b) AB₄ – Type with One Lone Pair and Three Bond Pairs

In such cases, the charge cloud of lone pair electrons (non-bonding electrons) expands more than that of bonding electrons. As a result, a somewhat large lone pair charge cloud tends to compress the bond angles in the remaining of the molecules.

Ammonia, NH₃ is a common example in this case.

The non-bonding electron in 2s orbital uses up more space and exerts a strong repulsive force on the bonding electron pairs. Consequently, to avoid a larger repulsion, the bonding electron pairs move closer that minimizes the ideal bond angle from 109.50 to 107.5 °. This effect forces ammonia to assume a triangular pyramidal geometry instead of a tetrahedral, as in methane. Similar, effects are evident in the geometries of molecules like PH₃, AsH₃, SbH_3, and BiH₃.

The substitution of hydrogen with electronegative atoms like F or Cl further decreases the bond angle.

In NF3, the strong polarity of the N-F bond pulls the lone pair of N atom closer to its nucleus, which in turn puts in a stronger repulsion over bonding electrons. Hence, the angle further diminishes to 102 °. Furthermore, the bond pairs N-F bonds are nearer to F atoms than N atoms. The increased distances in these bond sets make their repulsions less operative.

(c) AB₄-Type with Two lone Pairs and Two Bond Pairs:

The presence of two lone pairs introduces three types of repulsion i.e. lone pair-lone pair, lone pair-bond pair, and bond pair-bond pair repulsion. For instance: water (H₂O), a triatomic molecule is anticipated to be an AB₂ type linear molecule like BeCl₂ and CO₂.

But speculative proofs confirm a bent or angular geometry. VSEPR theory, effectively justifies the speculative results by arguing the participation of lone pairs, in addition to bond pairs in figuring out the general geometry of water molecules.

2 of the corners of a tetrahedron are inhabited by each of the two lone pairs and staying by bond pairs. But owing to the spatial arrangement of lone pairs and their repulsive action among themselves and on bond pairs, the bond angle is further decreased to 104.5 °. H₂S, H₂Se, H₂Te have similar geometries.

Multiple-Choice Questions (MCQs) with Answers:

 

Frequently Asked Questions (FAQs) – Valence Shell Electron Pair Repulsion Theory