VSEPR is short for valence shell electron pair repulsion. The significance of VSEPR theory is unprecedented in chemical bonding. This is because it is the first theory proposed by scientists that helped us predict the shapes and geometries of covalently bonded molecules. In this article, you will learn what VSEPR theory is, why it is important and how it is applied to determine a molecule’s shape and/or geometry.
But first, let us start from the basics and build some essential background concepts before studying the VSEPR concept.
- Valence electrons: The electrons present in the outermost shell of an atom that can participate in covalent bond formation.
- Bond pair: An electron pair shared between two bonded atoms in a covalently bonded molecule.
- Lone pair: An electron pair present on an atom in a molecule that is not involved in any chemical bond. It is also known as an unbonded electron pair.
- Electron density regions or electron domains: The total number of bond pairs and lone pairs present in a molecule.
- Steric number: The total number of electron domains around the central atom of a molecule.
- Lewis structure: A Lewis structure is a simplified representation of all the valence electrons present in a molecule. In a Lewis dot structure, the bonded electron pairs are represented by a straight line. Each straight line = 1 bond pair = 2 bonding electrons. However, the dots represent non-bonded or lone pairs of electrons. 2 dots = 1 lone pair = 2 non-bonded electrons.
Lewis representation was introduced and used for representing covalent molecules by Gilbert Lewis Newton. However, it provided no information about the shape or geometry of a covalent molecule. Thus, there was a need for a better concept that could sufficiently explain molecular shapes. This concept emerged in 1940 as Sidgwick and Powell introduced the VSEPR theory of chemical bonding.
What is VSEPR theory – Definition
According to the VSEPR theory, the electron pairs (including bond pairs and lone pairs) arrange around the central atom of a covalently bonded molecule such that the overall electron repulsive effect is minimized. This three-dimensional arrangement of electron pairs in space determines the shape and/or geometry of the molecule.
Main postulates of the VSEPR theory
- In a polyatomic molecule, the central atom is the atom bonded to two or more atoms. Usually, the least electronegative atom is chosen as the central atom of a molecule.
- The ideal electronic geometry of a molecule depends on the total number of electron density regions around the central atom (whether it’s a bond pair or a lone pair).
- The molecular geometry or shape of a molecule is strongly influenced by the difference in bond pairs and lone pairs around the central atom.
- A lone pair occupies more space than a bond pair. The electronic repulsions decrease in the order: lone pair-lone pair repulsions > lone pair-bond pair repulsions > bond pair-bond pair repulsions.
- The presence of lone pairs on the central atom changes the mutual X-A-X bond angles and distorts the shape and/or geometry of a molecule such that it occupies a different shape from the ideal electron pair geometry.
- Increase in electronegativity of atoms bonded to the central atom also decreases the X-A-X bond angle.
- Bond angles involving multiple (double or triple) bonds are larger than bond angles in molecules involving only single bonds.
- Electron pairs in a double or triple bond are considered one region of electron density.
The role of VSEPR theory in predicting the shapes of molecules
AXE is a general formula used for predicting the shape of a molecule according to the VSEPR concept.
In the AXE generic formula;
- A represents the central atom of a molecule.
- X stands for the atoms bonded to the central atom of a molecule.
- E denotes the lone pairs present on the central atom.
AXE application with examples
To one atom at the center, two atoms are covalently bonded, and there is no lone pair on the central atom.
The AXE generic formula for the carbon dioxide (CO2) molecule is AX2E0 or simply AX2. In CO2, a carbon (C) atom is present at the center. Two oxygen (O) atoms are double-covalently bonded to the central C-atom, so X=2 while there is no lone pair of electrons on the central C-atom, so E=0 for CO2.
The electronic configuration of carbon (6C12) is 1s2 2s2 2p2. This shows that shell no. 2 is its outermost shell consisting of a total of 2+2 = 4 valence electrons. One double bond = 4 electrons, and two C=O double bonds = 4+4 = 8 valence electrons.
All four valence electrons consumed in chemical bonding ensure there is no lone pair on the central carbon atom in CO2. No lone pair-lone pair or lone pair-bond pair repulsions exist in the molecule. Hence according to the VSEPR chart (given at the end of the article), its molecular geometry or shape is identical to its ideal electron geometry, i.e., linear. The O=C=O bonded atoms form a mutual bond angle of 180°.
To one atom at the center, three atoms are covalently bonded, and there is no lone pair on the central atom.
The bonded atoms get symmetrically arranged along the three vertices of an equilateral triangle, forming a mutual bond angle of 120°. Thus, AX3-type molecules adopt a trigonal planar shape.
Examples include BF3, SO3, etc.
It is a sub-type of AX3-type molecules. If one of three bonded atoms (X) gets removed and is replaced with a lone pair (E), it results in an AX2E-type molecule.
The presence of a lone pair on the central atom A leads to strong bond pair-lone pair repulsions. Thus, the X-A-X bond angle decreases to 118°, and the molecule adopts a bent, angular or V-shape.
Examples include SO2, SnCl2, etc.
To one atom at the center, four atoms are covalently bonded, and there is no lone pair on the central atom.
The bonded atoms adopt a four vertices arrangement, i.e., a symmetrical tetrahedral shape, forming a mutual bond angle of 109.5°.
Examples include CH4, CCl4, SiF4, etc.
It is a sub-type of AX4-type molecules. If one of three bonded atoms (X) gets removed and is replaced with a lone pair (E), it results in an AX3E-type molecule.
The presence of a lone pair on the central atom A leads to strong bond pair-lone pair repulsions. Thus, the X-A-X bond angle decreases to 107.5°, and the molecule adopts a trigonal pyramidal shape.
Examples include NH3, PCl3, etc.
Another sub-type of AX4-type molecules, AX2E2-type molecules, are formed when 2 X-atoms get replaced by 2 lone pairs on the central atom.
Even stronger lone pair-lone pair repulsions exist in the molecule in addition to lone pair-bond pair repulsions. Thus, the molecule adopts a bent, angular or V-shape. The bond angle decreases further to 104.5°.
Examples include H2O, H2S, N2O, Cl2O, etc.
To one atom at the center, five atoms are covalently bonded, and there is no lone pair on the central atom.
The molecule adopts a trigonal bipyramidal shape with three different bond angles, i.e., 90°, 120° and 180°.
Examples include PCl5, PF5, etc.
In AX5-type molecules, the atoms occupying two axial positions are kept fixed. In contrast, the bonded atoms occupying three equatorial positions can be replaced with lone pairs to obtain AX4E (seesaw-shape), AX3E2 (T-shape) and AX2E2-type (linear shape) molecules, respectively.
To one atom at the center, six atoms are covalently bonded, and there is no lone pair on the central atom.
The bonded atoms occupy six corners of an eight vertices arrangement to produce an octahedral shape.
Examples include SF6, PCl6, etc. The 4 X-atoms forming a square at the center of AX6-type molecules stay fixed. However, the 2 atoms lying above and below the square can be replaced with lone pairs to yield AX5E-type (square pyramidal) and AX4E2-type (square planar) molecular shapes.
Examples of these types of molecules are given in the VSEPR chart below.
Importance and limitations of VSEPR theory
Despite considering its paramount importance in predicting the shapes and geometry of molecules, the VSEPR theory fails to answer fundamental questions such as:
- Why do molecules adopt different shapes while possessing the same number of valence electrons?
- How is a chemical bond technically formed, considering the different energies at which atomic orbitals of participating atoms lie?
A reasonable explanation for these questions lies in the valence bond theory (VBT) of chemical bonding, which we have explored in our next article.
So, stay connected and happy learning!
Last but not least, test your knowledge by practicing some examples of VSEPR theory here.
1. Chaudhary, Ghulam Rasool. 2006. Inorganic chemistry
2. Madhusha. 2017. “Difference Between VSEPR and Valence Bond Theory.”
3. Oriakhi, Christopher O. 2021. “102. Chemical Bonding 2: Modern Theories of Chemical Bonding.” In Chemistry in Quantitative Language: Fundamentals of General Chemistry Calculations, edited by Christopher O. Oriakhi, 0. Oxford University Press.