Encyclopedia
In
chemistry,
hybridisation or
hybridization is the concept of mixing
atomic orbitals to form new
hybrid orbitals suitable for the qualitative description of atomic bonding properties. Hybridised orbitals are very useful in the explanation of the shape of
molecular orbitals for
molecules. It is an integral part of valence bond theory and the
valence shell electron-pair repulsion theory .
Historical development
The hybridisation theory was promoted by
chemist Linus Pauling in order to explain the structure of
molecules such as
methane . Historically, this concept was developed for such simple chemical systems but the approach was later applied more widely, and today it is considered an effective heuristic for rationalizing the structures of
organic compounds.
Hydridization theory is, however, considered less useful and less informative than
Molecular Orbital Theory. Problems with hybridization are especially notable when the
d orbitals are involved in bonding, as in
coordination chemistry and organometallic chemistry. Although hybridisation schemes in transition metal chemistry can be used, they are not accurate and have little predictive power.
It is important to note that orbitals are a model representation of the behavior of electrons within molecules. In the case of simple hybridisation, this approximation is based on the atomic orbitals of hydrogen. Hybridised orbitals are assumed to be mixtures of these atomic orbitals, superimposed on each other in various proportions. Hydrogen orbitals are used as a basis for simple schemes of hybridisation because it is one of the few examples of orbitals for which an exact analytic solution to its Schrödinger equation is known. These orbitals are then assumed to be slightly, but not significantly distorted in heavier atoms, like carbon, nitrogen, and oxygen. Under these assumptions is the theory of hybridisation most applicable. It must be noted, that one does not need hybridisation to describe molecules, but for molecules made up from
carbon,
nitrogen and
oxygen the hybridisation theory/model makes the description much easier.
The hybridisation theory finds its use mainly in organic chemistry, and mostly concerns C, N and O . Its explanation starts with the way bonding is organized in
methane.
sp3 hybrids
Hybridisation describes the bonding atoms from an atom's point of view. That is, for a tetrahedrally coordinated carbon , the carbon should have 4 orbitals with the correct symmetry to bond to the 4 hydrogen atoms. The problem with the existence of methane is now this: Carbon's ground-state configuration is
1s² 2s² 2px¹ 2py¹ or perhaps more easily read:
The valence bond theory would predict, based on the existence of two half-filled p-type orbitals , that C forms two
covalent bonds. CH
2. However, methylene is a very reactive molecule and cannot exist outside of a molecular system. Therefore, this theory alone cannot explain the existence of CH
4.
Furthermore, ground state orbitals cannot be used for bonding in CH
4. While exciting a 2s electron into a 2p orbital would theoretically allow for four bonds, according to the valence bond theory which has been proved experimentally correct for systems like O
2 this would imply that the various bonds of CH
4 would have differing energies due to differing levels of orbital overlap. Once again, this has been experimentally disproved: any hydrogen can be removed from a carbon with equal ease.
To summarise, to explain the existence of CH
4 a method by which as many as 12 bonds of equal strength can be created was required.
The first step in hybridisation is the excitation of one electrons :
The proton that forms the nucleus of a hydrogen atom attracts one of the valence electrons on carbon. This causes an excitation, moving a 2s electron into a 2p orbital. This, however, increases the influence of the carbon nucleus on the valence electrons by increasing the effective core potential .
The combination of these forces creates new mathematical functions known as hybridised orbitals. In the case of carbon attempting to bond with four hydrogens, four orbitals are required. Therefore, the 2s orbital mixes with the three 2p orbitals to form four
sp3 hybrids . See graphical summary below.
becomes
In CH
4, four sp³ hybridised orbitals are overlapped by
hydrogen's
1s orbital, yielding four
sigma bonds. The four bonds are of the same length and strength. This theory fits our requirements.
translates into
An alternative view is: View the carbon as the C
4- anion. In this case all the orbitals on the carbon are filled:
If we now recombine these orbitals with the empty s-orbitals of 4 hydrogens and allow maximum separation between the 4 hydrogens , we see that at any orientation of the p-orbitals, a single hydrogen has an overlap of 25% with the s-orbital of the C, and a total of 75% of overlap with the 3 p-orbitals .
According to the orbital hybridization theory the valence electrons in methane should be equal in energy but its photoelectron spectrum shows two bands, one at 12.7 eV and one at 23 eV . This apparent inconsistency can be explained when one considers additional orbital mixing taking place when the sp
3 orbitals mix with the 4 hydrogen orbitals.
sp2 hybrids
Other carbon based compounds and other molecules may be explained in a similar way as methane, take for example
ethene . Ethene has a double bond between the carbons. The Lewis structure looks like this:
Carbon will sp
2 hybridise, because hybrid orbitals will form only sigma bonds and one
pi bond is required for the
double bond between the carbons. The hydrogen-carbon bonds are all of equal strength and length, which agrees with experimental data.
In
sp2 hybridization the 2s orbital is mixed with only two of the three available 2p orbitals:
forming a total of 3 sp
2 orbitals with one p-orbital remaining. In ethene the two carbon atoms form a
sigma bond by overlap of two sp
2 orbitals and each carbon atoms forms two covalent bonds with hydrogen by s - sp
3 overlap all with 120° angles. the
pi-bond between the carbon atoms perpendicular to the molecular plane is formed by 2p-2p overlap.
The amount of p-character is not restricted to integer values, i.e. hybridisations like sp
2.5 are also readily described. In this case the geometries are somewhat distorted from the ideally hybridised picture. For example, as stated in Bent's rule, a bond tends to have higher p-character when directed toward a more electronegative substituent.
sp hybrids
The chemical bonding in compounds such as
alkynes with triple bonds is explained by
sp hybridization.
In this model the 2s orbital mixes with only one of the three p-orbitals resulting in two sp orbitals and two remaining unchanged p orbitals. The chemical bonding in
acetylene consists of sp - sp overlap between the two carbon atoms forming a
sigma bond and two additional
pi bonds form by p - p overlap. Each carbon also bonds to hydrogen in a sigma s - sp overlap at 180° angles.
Hybridisation and molecule shape
Using hybridisation, along with the
VSEPR theory, helps to explain molecule shape:
- AX1 : no hybridisation; trivially linear shape
- AX2 : sp hybridisation; linear or diagonal shape; bond angles are cos-1 = 180°
- AX3 : sp² hybridisation; trigonal planar shape; bond angles are cos-1 = 120°
- AX4 : sp³ hybridisation; tetrahedral shape; bond angles are cos-1 ˜ 109.5°
- AX5 : sp³d hybridisation; trigonal bipyramidal shape
- AX6 : sp³d² hybridisation; octahedral shape
This holds if there are no lone electron pairs on the central atom. If there are, they should be counted in the X
i number, but bond angles become smaller due to increased repulsion. For example, in
water , the
oxygen atom has two bonds with H and two lone electron pairs , which means there are four such 'elements' on O. The model molecule is, then, AX
4: sp³ hybridization is utilized, and the electron arrangement of H
2O is tetrahedral. This agrees with the shape, we know water has a non-linear, bent structure, with an angle of 104.5 degrees .
Hybridization theory has been superseded by MO theory
Although the language and pictures arising from Hybridization Theory, more widely known as Valence Bond Theory, remain widespread in synthetic organic chemistry, this qualitative analysis of bonding has been largely superseded by
Molecular Orbital Theory. For example, inorganic chemistry texts have all but abandoned instruction of hybridization, except as a historical footnote. One specific problem with hybridization is that is incorrectly predicts the photoelectron spectra of many molecules, including such fundamental species such as methane and water. From a pedogical perspective, hybridization approach tends to over-emphasize localization of bonding electrons and does not effectively embrace molecular symmetry as does MO Theory.
References
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