VSEPR theory

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Valence shell electron pair repulsion (VSEPR) theory (1957) is a model in chemistry, which is used for predicting the shapes of individual molecules, based upon their extent of electron-pair electrostatic repulsion, determined using steric numbers.[1] The theory is also called the Gillespie-Nyholm theory after the two main developers, and VSEPR is sometimes pronounced as "vesper" which is easier to say.

The premise of VSEPR is that to predict the geometric shape and lone-pair behavior of a molecule one must consider : the constructed Lewis structure, expanded to show all lone pairs of electrons, alongside protruding and projecting bonds, as well as consider the total steric number of that molecule.

VSEPR theory is based on the idea that the geometry of a molecule or polyatomic ion is determined primarily by repulsion among the pairs of electrons associated with a central atom. The pairs of electrons may be bonding or nonbonding (also called lone pairs). Only valence electrons of the central atom influence the molecular shape in a meaningful way.

Contents

Basic assumptions

  1. Pairs of electrons in the valence shell of a central atom repel each other.
  2. These pairs of electrons tend to occupy positions in space that minimize repulsions and maximize the distance of separation between them.
  3. The valence shell is taken as a sphere with electron pairs localizing on the spherical surface at maximum distance from one another.
  4. A multiple bond is treated as if it is a single electron pair and the two or three electron pairs of a multiple bond are treated as a single super pair.
  5. Where two or more resonance structures can depict a molecule the VSEPR model is applicable to any such structure.

Three types of repulsion take place between the electrons of a molecule:

  • The lone pair-lone pair repulsion
  • The lone pair-bonding pair repulsion
  • The bonding pair-bonding pair repulsion.

A molecule must avoid these repulsions to remain stable. When repulsion cannot be avoided, the weaker repulsion (i.e. the one that causes the smallest deviation from the ideal shape) is preferred.

The lone pair-lone pair (lp-lp) repulsion is considered to be stronger than the lone pair-bonding pair (lp-bp) repulsion, which in turn is stronger than the bonding pair-bonding pair (bp-bp) repulsion. Hence, the weaker bp-bp repulsion is preferred over the lp-lp or lp-bp repulsion.

VSEPR theory is usually compared (but not part of) and contrasted with valence bond theory, which addresses molecular shape through orbitals that are energetically accessible for bonding. Valence bond theory concerns itself with the formation of sigma and pi bonds. Molecular orbital theory is another model for understanding how atoms and electrons are assembled into molecules and polyatomic ions.

VSEPR theory has long been criticized for not being quantitative, and therefore limited to the generation of "crude", even though structurally accurate, molecular geometries of covalent molecules. However, molecular mechanics force fields based on VSEPR have also been developed.[2]

AXE Method

The "AXE method" of electron counting is commonly used when applying the VSEPR theory. The A represents the central atom and always has an implied subscript one. The X represents how many sigma bonds are formed between the central atoms and outside atoms. Multiple covalent bonds (double, triple, etc) count as one X. The E represents the number of lone electron pairs present outside of the central atom. The sum of X and E, sometimes known as the steric number, is also associated with the total number of hybridised orbitals used by valence bond theory.

Steric
No.		 Basic Geometry
0 lone pair		 1 lone pair 2 lone pairs 3 lone pairs
2
linear
      
3
trigonal planar
bent
    
4
tetrahedral
trigonal pyramid
bent
  
5
trigonal bipyramid
seesaw
T-shaped
linear
6
octahedral
square pyramid
square planar
  
7
pentagonal bipyramid
pentagonal pyramid
    


Molecule Type Shape Electron arrangement† Geometry‡ Examples
AX1En Diatomic HF, O2
AX2E0 Linear BeCl2, HgCl2, CO2
AX2E1 Bent NO2−, SO2, O3
AX2E2 Bent H2O, OF2
AX2E3 Linear XeF2, I3−
AX3E0 Trigonal planar BF3, CO32−, NO3−, SO3
AX3E1 Trigonal pyramidal NH3, PCl3
AX3E2 T-shaped ClF3, BrF3
AX4E0 Tetrahedral CH4, PO43−, SO42−, ClO4−
AX4E1 Seesaw SF4
AX4E2 Square Planar XeF4
AX5E0 Trigonal Bipyramidal PCl5
AX5E1 Square Pyramidal ClF5, BrF5
AX6E0 Octahedral SF6
AX6E1 Pentagonal pyramidal XeF6
AX7E0 Pentagonal bipyramidal IF7
† Electron arrangement including lone pairs, shown in pale yellow
‡ Observed geometry (excluding lone pairs)

When the substituent (X) atoms are not all the same, the geometry is still approxmiately 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°. Similarly, SOCl2 is AX3E1, but because the X substituents are not identical, the XAX angles are not all equal.

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 all of the outer atoms.

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. The observed shape of the molecule is a trigonal pyramid, because the lone pair is not "visible" in experimental methods used to determine molecular geometry. The shape of a molecule is found from the relationship of the atoms even though it can be influenced by lone pairs of electrons.

A steric number of seven is possible, but it occurs in uncommon compounds such as iodine heptafluoride. The base geometry for this is pentagonal bipyramidal.

Exceptions

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

Transition metal compounds

Many transition metal compounds do not have geometries explained by VSEPR which can be ascribed to there being no lone pairs in the valence shell and the interaction of core d electrons with the ligands.[3] The structure of some of these compounds, including metal hydrides and alkyl complexes such as hexamethyltungsten, can be predicted correctly using the VALBOND theory, which is based on sd hybrid orbitals and the 3-center-4-electron bonding model.[4] [5] Crystal field theory is another theory that can often predict the geometry of coordination complexes.

Group 2 halides

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°).[6] It has been proposed by Gillespie that this is caused by interaction of the ligands with the electron core of the metal atom, polarising it so that the inner shell is not spherically symmetric, thus influencing the molecular geometry. [3][7]

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 being essentially ionic leading to strong repulsion between the lithium atoms.[8]
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°). Gillespies 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, the central atom.[3] When the central atom is more electronegative, as in O(SiH3)2, the lone pairs are less well localised, have a weaker repulsive effect and this 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.[3]

Some AX6E1molecules

Some AX6E1 molecules, e.g. the Te(IV)and Bi(III) anions, TeCl62−, TeBr62−, BiCl63−, BiBr63− and BiI63−, are regular octahedra and the lone pair does not affect the geometry.[9] One rationalisation is that steric crowding of the ligands allows no room for the non-bonding lone pair,[3], another rationalisation is the inert pair effect[10]

See also

Wikibooks

References

  1. ^ Modern Inorganic Chemistry W.L. Jolly ISBN 0-07-032760-2
  2. ^ VGS Box. Journal of Molecular Modeling, 1997, 3, 124-141.
  3. ^ a b c d e Models of molecular geometry, Gillespie R. J., Robinson E.A. Chem. Soc. Rev., 2005, 34, 396–407, doi: 10.1039/b405359c
  4. ^ Landis, C. K.; Cleveland, T.; Firman, T. K. Making sense of the shapes of simple metal hydrides. J. Am. Chem. Soc. 1995, 117, 1859-1860.
  5. ^ Landis, C. K.; Cleveland, T.; Firman, T. K. Structure of W(CH3)6. Science 1996, 272, 182-183.
  6. ^ Greenwood, Norman N.; Earnshaw, A. (1997). Chemistry of the Elements, 2nd Edition, Oxford: Butterworth-Heinemann. ISBN 0-7506-3365-4. 
  7. ^ Core Distortions and Geometries of the Difluorides and Dihydrides of Ca, Sr, and Ba Bytheway I, Gillespie R.J, Tang T.H., Bader R.F. Inorganic Chemistry, 34,9, 2407-2414, 1995 doi:10.1021/ic00113a023
  8. ^ A spectroscopic determination of the bond length of the LiOLi molecule: Strong ionic bonding, D. Bellert, W. H. Breckenridge, J. Chem. Phys. 114, 2871 (2001); doi:10.1063/1.1349424
  9. ^ Wells A.F. (1984) Structural Inorganic Chemistry 5th edition Oxford Science Publications ISBN 0-19-855370-6
  10. ^ Catherine E. Housecroft, Alan G. Sharpe (2005), Inorganic Chemistry, Pearson Education, ISBN 0130399132

External links

  • 3D Chem - Chemistry, Structures, and 3D Molecules
  • IUMSC - Indiana University Molecular Structure Center
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