Inorganic Chemistry/Chemical Bonding/VSEPR theory

Valence shell electron pair repulsion (VSEPR) theory is a model used in chemistry to predict the geometry of individual molecules from the number of electron pairs surrounding their central atoms. It is also named the Gillespie-Nyholm theory after its two main developers, Ronald Gillespie and Ronald Nyholm.

The premise of VSEPR is that the valence electron pairs surrounding an atom tend to repel each other and will, therefore, adopt an arrangement that minimizes this repulsion. This in turn decreases the molecule's energy and increases its stability, which determines the molecular geometry.

History

The idea of a correlation between molecular geometry and number of valence electrons (both shared and unshared) was first presented in a Bakerian Lecture in 1940 by Nevil Sidgwick and Herbert Powell at the University of Oxford. In 1957 Ronald Gillespie and Ronald Sydney Nyholm at University College London refined this concept to build a more detailed theory capable of choosing between various alternative geometries.

Overview

VSEPR theory is used to predict the arrangement of electron pairs around central atoms in molecules, especially simple and symmetric molecules. A central atom is defined in this theory as an atom which is bonded to two or more other atoms, while a terminal atom is bonded to only one other atom. For example in the molecule methyl isocyanate (H₃C-N=C=O), the two carbons and one nitrogen are central atoms, and the three hydrogens and one oxygen are terminal atoms. The geometry of the central atoms and their non-bonding electron pairs in turn determine the geometry of the larger whole molecule.

The number of electron pairs in the valence shell of a central atom is determined after drawing the Lewis structure of the molecule, and expanding it to show all bonding groups and lone pairs of electrons. In VSEPR theory, a double bond or triple bond is treated as a single bonding group. The sum of the number of atoms bonded to a central atom and the number of lone pairs formed by its nonbonding valence electrons is known as the central atom's steric number.

The electron pairs (or groups if multiple bonds are present) are assumed to lie on the surface of a sphere centered on the central atom and tend to occupy positions that minimize their mutual repulsions by maximizing the distance between them. The number of electron pairs (or groups), therefore, determines the overall geometry that they will adopt. For example, when there are two electron pairs surrounding the central atom, their mutual repulsion is minimal when they lie at opposite poles of the sphere. Therefore, the central atom is predicted to adopt a linear geometry. If there are 3 electron pairs surrounding the central atom, their repulsion is minimized by placing them at the vertices of an equilateral triangle centered on the atom. Therefore, the predicted geometry is trigonal. Likewise, for 4 electron pairs, the optimal arrangement is tetrahedral.

As a tool in predicting the geometry adopted with a given number of electron pairs, an often used physical demonstration of the principle of minimal electron pair repulsion utilizes inflated balloons. Through handling, balloons acquire a slight surface electrostatic charge that results in the adoption of roughly the same geometries when they are tied together at their stems as the corresponding number of electron pairs. For example, five balloons tied together adopt the trigonal bipyramidal geometry, just as do the five bonding pairs of a PCl₅ molecule.

Steric number

Sulfur tetrafluoride has a steric number of 5.

The steric number of a central atom in a molecule is the number of atoms bonded to that central atom, called its coordination number, plus the number of lone pairs of valence electrons on the central atom. In the molecule SF₄, for example, the central sulfur atom has four ligands; the coordination number of sulfur is four. In addition to the four ligands, sulfur also has one lone pair in this molecule. Thus, the steric number is 4 + 1 = 5.

Degree of repulsion

The overall geometry is further refined by distinguishing between bonding and nonbonding electron pairs. The bonding electron pair shared in a sigma bond with an adjacent atom lies further from the central atom than a nonbonding (lone) pair of that atom, which is held close to its positively charged nucleus. VSEPR theory therefore views repulsion by the lone pair to be greater than the repulsion by a bonding pair. As such, when a molecule has 2 interactions with different degrees of repulsion, VSEPR theory predicts the structure where lone pairs occupy positions that allow them to experience less repulsion. Lone pair–lone pair (lp–lp) repulsions are considered stronger than lone pair–bonding pair (lp–bp) repulsions, which in turn are considered stronger than bonding pair–bonding pair (bp–bp) repulsions, distinctions that then guide decisions about overall geometry when 2 or more non-equivalent positions are possible. For instance, when 5 valence electron pairs surround a central atom, they adopt a trigonal bipyramidal molecular geometry with two collinear axial positions and three equatorial positions. An electron pair in an axial position has three close equatorial neighbors only 90° away and a fourth much farther at 180°, while an equatorial electron pair has only two adjacent pairs at 90° and two at 120°. The repulsion from the close neighbors at 90° is more important, so that the axial positions experience more repulsion than the equatorial positions; hence, when there are lone pairs, they tend to occupy equatorial positions as shown in the diagrams of the next section for steric number five.

The difference between lone pairs and bonding pairs may also be used to rationalize deviations from idealized geometries. For example, the H₂O molecule has four electron pairs in its valence shell: two lone pairs and two bond pairs. The four electron pairs are spread so as to point roughly towards the apices of a tetrahedron. However, the bond angle between the two O–H bonds is only 104.5°, rather than the 109.5° of a regular tetrahedron, because the two lone pairs (whose density or probability envelopes lie closer to the oxygen nucleus) exert a greater mutual repulsion than the two bond pairs.

A bond of higher bond order also exerts greater repulsion since the pi bond electrons contribute. For example in isobutylene, (H₃C)₂C=CH₂, the H₃C−C=C angle (124°) is larger than the H₃C−C−CH₃ angle (111.5°). However, in the carbonate ion, CO2−3, all three C−O bonds are equivalent with angles of 120° due to resonance.

AXE method

The "AXE method" of electron counting is commonly used when applying the VSEPR theory. The electron pairs around a central atom are represented by a formula AX_nE_m, where A represents the central atom and always has an implied subscript one. Each X represents a ligand (an atom bonded to A). Each E represents a lone pair of electrons on the central atom. The total number of X and E is known as the steric number. For example in a molecule AX₃E₂, the atom A has a steric number of 5.

When the substituent (X) atoms are not all the same, the geometry is still approximately valid, but the bond angles may be slightly different from the ones where all the outside atoms are the same. For example, the double-bond carbons in alkenes like C₂H₄ are AX₃E₀, but the bond angles are not all exactly 120°. Likewise, SOCl₂ is AX₃E₁, but because the X substituents are not identical, the X–A–X angles are not all equal.

Based on the steric number and distribution of Xs and Es, VSEPR theory makes the predictions in the following tables.

Main-group elements

For main-group elements, there are stereochemically active lone pairs E whose number can vary between 0 to 3. Note that the geometries are named according to the atomic positions only and not the electron arrangement. For example, the description of AX₂E₁ as a bent molecule means that the three atoms AX₂ are not in one straight line, although the lone pair helps to determine the geometry.

Steric number	Molecular geometry 0 lone pairs	Molecular geometry 1 lone pair	Molecular geometry 2 lone pairs	Molecular geometry 3 lone pairs
2	Linear
3	Trigonal planar	Bent
4	Tetrahedral	Trigonal pyramidal	Bent
5	Trigonal bipyramidal	Seesaw	T-shaped	Linear
6	Octahedral	Square pyramidal	Square planar
7	Pentagonal bipyramidal	Pentagonal pyramidal	Pentagonal planar
8	Square antiprismatic

Molecule type	Shape	Electron arrangement including lone pairs, shown in pale yellow	Geometry excluding lone pairs	Examples
AX2E0	Linear			BeCl2, CO2
AX2E1	Bent			NO2-, SO2, O3, CCl2
AX2E2	Bent			H2O, OF2
AX2E3	Linear			XeF2, I3-, XeCl2
AX3E0	Trigonal planar			BF3, CO2−3, NO3-, SO3
AX3E1	Trigonal pyramidal			NH3, PCl3
AX3E2	T-shaped			ClF3, BrF3
AX4E0	Tetrahedral			CH4, PO3−4, SO2−4, ClO4-, XeO4
AX4E1	Seesaw or disphenoidal			SF4
AX4E2	Square planar			XeF4
AX5E0	Trigonal bipyramidal			PCl5
AX5E1	Square pyramidal			ClF5, BrF5, XeOF4
AX5E2	Pentagonal planar			XeF5-
AX6E0	Octahedral			SF6
AX6E1	Pentagonal pyramidal			XeOF5-, IOF2−5
AX7E0	Pentagonal bipyramidal			IF7
AX8E0	Square antiprismatic			IF8-

Transition metals (Kepert model)

The lone pairs on transition metal atoms are usually stereochemically inactive, meaning that their presence does not change the molecular geometry. For example, the hexaaquo complexes M(H₂O)₆ are all octahedral for M = V³⁺, Mn³⁺, Co³⁺, Ni²⁺ and Zn²⁺, despite the fact that the electronic configurations of the central metal ion are d², d⁴, d⁶, d⁸ and d¹⁰ respectively. The Kepert model ignores all lone pairs on transition metal atoms, so that the geometry around all such atoms corresponds to the VSEPR geometry for AX_n with 0 lone pairs E. This is often written ML_n, where M = metal and L = ligand. The Kepert model predicts the following geometries for coordination numbers of 2 through 9:

Molecule type	Shape	Geometry	Examples
ML2	Linear		HgCl2
ML3	Trigonal planar
ML4	Tetrahedral		NiCl2−4
ML5	Trigonal bipyramidal		Fe(CO)5
	Square pyramidal		MnCl52−
ML6	Octahedral		WCl6
ML7	Pentagonal bipyramidal		ZrF3−7
	Capped octahedral		MoF7-
	Capped trigonal prismatic		TaF2−7
ML8	Square antiprismatic		ReF8-
	Dodecahedral		Mo(CN)4−8
	Bicapped trigonal prismatic		ZrF4−8
ML9	Tricapped trigonal prismatic		ReH2−9
	Capped square antiprismatic

Examples

The methane molecule (CH₄) is tetrahedral because there are four pairs of electrons. The four hydrogen atoms are positioned at the vertices of a tetrahedron, and the bond angle is cos⁻¹(−¹⁄₃) ≈ 109° 28′. This is referred to as an AX₄ type of molecule. As mentioned above, A represents the central atom and X represents an outer atom.

The ammonia molecule (NH₃) has three pairs of electrons involved in bonding, but there is a lone pair of electrons on the nitrogen atom. It is not bonded with another atom; however, it influences the overall shape through repulsions. As in methane above, there are four regions of electron density. Therefore, the overall orientation of the regions of electron density is tetrahedral. On the other hand, there are only three outer atoms. This is referred to as an AX₃E type molecule because the lone pair is represented by an E. By definition, the molecular shape or geometry describes the geometric arrangement of the atomic nuclei only, which is trigonal-pyramidal for NH₃.

Steric numbers of 7 or greater are possible, but are less common. The steric number of 7 occurs in iodine heptafluoride (IF₇); the base geometry for a steric number of 7 is pentagonal bipyramidal. The most common geometry for a steric number of 8 is a square antiprismatic geometry. Examples of this include the octacyanomolybdate (Mo(CN)4−8) and octafluorozirconate (ZrF4−8) anions. The nonahydridorhenate ion (ReH2−9) in potassium nonahydridorhenate is a rare example of a compound with a steric number of 9, which has a tricapped trigonal prismatic geometry.

Possible geometries for steric numbers of 10, 11, 12, or 14 are bicapped square antiprismatic (or bicapped dodecadeltahedral), octadecahedral, icosahedral, and bicapped hexagonal antiprismatic, respectively. No compounds with steric numbers this high involving monodentate ligands exist, and those involving multidentate ligands can often be analysed more simply as complexes with lower steric numbers when some multidentate ligands are treated as a unit.

Exceptions

There are groups of compounds where VSEPR fails to predict the correct geometry.

Some AX₂E₀ molecules

The shapes of heavier Group 14 element alkyne analogues (RM≡MR, where M = Si, Ge, Sn or Pb) have been computed to be bent.

Some AX₂E₂ molecules

One example is molecular lithium oxide, Li₂O, which is linear rather than being bent, and this has been ascribed to the bonding's being in essence ionic, leading to strong repulsion between the lithium atoms.
Another example is O(SiH₃)₂ with an Si-O-Si angle of 144.1°, which compares to the angles in Cl₂O (110.9°), (CH₃)₂O (111.7°)and N(CH₃)₃ (110.9°). Gillespie's rationalisation is that the localisation of the lone pairs, and therefore their ability to repel other electron pairs, is greatest when the ligand has an electronegativity similar to, or greater than, that of the central atom. When the central atom is more electronegative, as in O(SiH₃)₂, the lone pairs are less well-localised and have a weaker repulsive effect. This fact, combined with the stronger ligand-ligand repulsion (-SiH₃ is a relatively large ligand compared to the examples above), gives the larger-than-expected Si-O-Si bond angle.

Some AX₆E₁ and AX₈E₁ molecules

Xenon hexafluoride, which has a distorted octahedral geometry

Some AX₆E₁ molecules, e.g. xenon hexafluoride (XeF₆) and the Te(IV) and Bi(III) anions, TeCl2−6, TeBr2−6, BiCl3−6, BiBr3−6 and BiI3−6, are octahedra rather than pentagonal pyramids, and the lone pair does not affect the geometry to the degree predicted by VSEPR. One rationalization is that steric crowding of the ligands allows no room for the non-bonding lone pair; another rationalization is the inert pair effect. Similarly, the octafluoroxenate ion (XeF2−8) in nitrosonium octafluoroxenate(VI) is a square antiprism and not a bicapped trigonal prism (as predicted by VSEPR theory for an AX₈E₁ molecule), despite having a lone pair.

Square planar ML₄ complexes

The Kepert model predicts that ML₄ transition metal molecules are tetrahedral in shape, and it cannot explain the formation of square planar complexes. The majority of such complexes exhibit a d⁸ configuration as for the tetrachloroplatinate (PtCl2−4) ion. The explanation of the shape of square planar complexes involves electronic effects and requires the use of crystal field theory.

Complexes with strong d-contribution

Hexamethyltungsten, a transition metal complex whose geometry is different from main-group coordination.

Some transition metal complexes with low d electron count have unusual geometries, which can be ascribed to d subshell bonding interaction. Gillespie found that this interaction produces bonding pairs that also occupy the respective antipodal points (ligand opposed) of the sphere. This phenomenon is an electronic effect resulting from the bilobed shape of the underlying sd^x hybrid orbitals. The repulsion of these bonding pairs leads to a different set of shapes.

Molecule type	Shape	Geometry	Examples
ML2	Bent		VO+2
ML3	Trigonal pyramidal		CrO3
ML4	Tetrahedral		TiCl4
ML5	Square pyramidal		Ta(CH3)5
ML6	C3v Trigonal prismatic		W(CH3)6

The gas phase structures of the triatomic halides of the heavier members of group 2, (i.e., calcium, strontium and barium halides, MX₂), are not linear as predicted but are bent, (approximate X–M–X angles: CaF₂, 145°; SrF₂, 120°; BaF₂, 108°; SrCl₂, 130°; BaCl₂, 115°; BaBr₂, 115°; BaI₂, 105°). It has been proposed by Gillespie that this is also caused by bonding interaction of the ligands with the d subshell of the metal atom, thus influencing the molecular geometry.

Odd-electron molecules

The VSEPR theory can be extended to molecules with an odd number of electrons by treating the unpaired electron as a "half electron pair". In effect, the odd electron has an influence on the geometry which is similar to a full electron pair, but less pronounced so that the geometry may be intermediate between the molecule with a full electron pair and the molecule with one less electron pair on the central atom.

For example, nitrogen dioxide (NO₂) is an AX₂E_0.5 molecule, with an unpaired electron on the central nitrogen. VSEPR predicts a geometry similar to the NO₂- ion (AX₂E₁, bent, bond angle approx. 120°) but intermediate between NO₂- and NO+2 (AX₂E₀, linear, 180°). In fact NO₂ is bent with an angle of 134° which is closer to 120° than to 180°, in qualitative agreement with the theory.

Similarly chlorine dioxide (ClO₂, AX₂E_1.5) has a geometry similar to ClO₂- but intermediate between ClO₂- and ClO+2.

Finally the methyl radical (CH₃) is predicted to be trigonal pyramidal like the methyl anion (CH₃-) but with a larger bond angle as in the trigonal planar methyl cation (CH+3). However in this case the VSEPR prediction is not quite true as CH₃ is actually planar, although its distortion to a pyramidal geometry requires very little energy.