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--> In Chemistry , hybridisation or '''hybridization''' (see Spelling Differences ) is the mixing of Atomic Orbital s belonging to a same Electron Shell to form new orbitals suitable for the qualitative description of atomic bonding properties. Hybridized orbitals are very useful in explaining the shape of Molecular Orbital s for Molecule s. Hybridisation is an integral part of the Valence Shell Electron-pair Repulsion (VSEPR) Theory . EXAMPLE OF METHANE The hybridisation theory was theorised by configuration is ''1s² 2s² 2px¹ 2py¹'' or perhaps more easily read: (Note: The 1s orbital is lower in energy than the 2s orbital, and the 2s orbital is lower in energy than the 2p orbitals) The Valence Bond Theory would predict, based on the existence of two half-filled p-type orbitals (the designations px py or pz are meaningless at this point, as they do not fill in any particular order), that C forms two Covalent Bond s. CH2, however, is known as a Methylene group and cannot exist outside of a molecular system. Therefore, this theory alone cannot explain the existence of CH4. Furthermore, ground state orbitals cannot be used for bonding in CH4. 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 O2 this would imply that the various bonds of CH4 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 CH4 and many other molecules a method by which as many as 12 bonds (for Transition Metals ) of equal strength (and therefore equal length) can be created is required. MORE DETAILED DISCUSSION One approach to this is the concept of hybridisation. Historically, this concept was necessary in order to explain the bonding observed in very simple chemical systems. It was later found to be more widely applicable, and today it is considered an effective heuristic for understanding Organic Chemistry . It is less applicable to other branches of chemistry for which the heavier atoms are involved. Transition Metal Chemistry is one example. Hybridisation schemes in transition metal chemistry tend to be more sophisticated because the assumptions have to be somewhat more relaxed for hybridisation to be useful. As a consequence, this results in a more complicated theory of hybridisation for transition metals - but one in fact which is not very accurate and has little predictive power. The result of this was a development of entirely new branches of Bonding Theory , and these are an active area of theoretical chemical research today. It is important to note that orbitals are a model representation of how an electron around an atom behaves. In the case of simple hybridization, this approximation is based on the atomic orbitals of hydrogen. Hybridised orbitals are assumed to be different mixtures of these atomic orbitals, superimposed on each other in various different 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. The first step in hybridisation is the excitation of one (or more) electrons. From this point of the explanation on, it can be assumed that the subject of study is Carbon in the context of methane, for simplicity. 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 amount of charge the nucleus exerts on a given electron = Charge of Core - Charge of all electrons closer to the nucleus). 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 (core orbitals are almost never involved in bonding) mixes with the three 2p orbitals to form four sp3 hybrids (read as ess-pee-three). See graphical summary below. becomes In CH4, four sp³ hybridised orbitals are overlapped by Hydrogen 's ''1s'' orbitals, yielding four Sigma (σ) Bonds . The four bonds are of the same length and strength, and there are four of them. This theory fits our requirements. translates into Other C-compounds and other molecules may be explained similarly, for example Ethene (C2H4). Carbon will never form any less than four bonds unless it is given no other choice, which seldom occurs. Therefore, ethene has a double bond between the carbons. The Lewis structure looks like this: Carbon will sp2 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. The amount of p-character is not restricted to integer values, i.e. hybridisations like sp2.5 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. MOLECULE SHAPE Hybridisation, along with VSEPR Theory , helps to explain molecule shape.
This holds if there are no lone electron pairs on the central atom. If there are, they should be counted in the Xi number, but bond angles become smaller due to increased repulsion. For example, in Water (H2O), the Oxygen atom has two bonds with H and two lone electron pairs (as can be seen with the valence bond theory as well from the electronic configuration of oxygen), which means there are four such 'elements' on O. The model molecule is, then, AX4: sp³ hybridization is utilized, and the electron arrangement of H2O is tetrahedral. The shape, however, is non-linear bent, since lone electron pairs are not visible, and also because lone pair electrons on the bonds must be taken into account. The HOH angle is about 104.5 degrees. REFERENCES {Link without Title} L. Pauling, J. Am. Chem. Soc. 53 (1931), 1367
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