Teoria dos orbitais moleculares
Em química, a teoria das orbitais moleculares é um método para determinar estruturas moleculares nas quais elétrons não são atribuídos a ligações químicas individuais entre átomos, ao invés disto são tratados como movimentos sob a influência do núcleo molecular.
Nesta teoria, cada molécula possui um conjunto de orbitais moleculares, nos quais se assume que a função de onda de cada orbital ψf pode ser descrita como uma soma ponderada das n orbitais atómicas χi, de acordo com a equação:
Onde cij podem ser determinados pela substituição destas equações pela equação de Schrödinger e pela aplicação do princípio variacional. Este método é conhecido como combinação linear de orbitais atômicas e é bastante utilizada pela química computacional. Uma transformação adicional unitária pode ser aplicada ao sistema para acelerar a convergência em alguns esquemas computacionais.
A teoria das orbitais moleculares foi visto como um competidor à ligação de valência na década de 1930, hoje foi percebido que os dois métodos são relacionados e que quando generalizados eles se tornam equivalentes.
História[editar | editar código-fonte]
A teoria das orbitais moleculares foi desenvolvida alguns anos após a ligação de valência ter sido estabelecida em 1927. Os primeiros passos foram dados por Friedrich Hund, Robert Mulliken, John C. Slater, e John Lennard-Jones. Ela foi originalmente chamada de teoria de Hund-Mulliken. A palavra orbital foi introduzida por Mulliken em 1932. Em 1933, a teoria das orbitais moleculares havia se tornado uma teoria válida.
== Definição == A teoria da [[orbital molecular]] utiliza uma [[Molecular orbital]] (MO) theory uses a [[Linear combination of atomic orbitals molecular orbital method|linear combination of atomic orbitals]] (LCAO) to form molecular orbitals which cover the whole molecule. These are often divided into bonding orbitals, [[antibonding|anti-bonding]] orbitals, and [[non-bonding orbital]]s. A [[molecular orbital]] is merely a Schrödinger orbital which includes several, but often only two nuclei. If this orbital is of type in which the electron(s) in the orbital have a higher probability of being ''between'' nuclei than elsewhere, the orbital will be a bonding orbital, and will tend to hold the nuclei together. If the electrons tend to be present in a molecular orbital in which they spend more time elsewhere than between the nuclei, the orbital will function as an [[antibonding orbital|anti-bonding orbital]] and will actually weaken the bond. Electrons in non-bonding orbitals tend to be in deep orbitals (nearly [[atomic orbital]]s) associated almost entirely with one nucleus or the other, and thus they spend equal time between nuclei or not. These electrons neither contribute nor detract from bond strength. Molecular orbitals are further divided according to the types of atomic orbitals combining to form a bond. These orbitals are results of electron-[[Atomic nucleus|nucleus]] interactions that are caused by the [[fundamental interaction|fundamental]] force of [[electromagnetism]]. Chemical substances will form a bond if their orbitals become lower in energy when they interact with each other. Different chemical bonds are distinguished that differ by [[electron configuration|electron cloud shape]] and by [[energy level]]s. MO theory provides a global, delocalized perspective on chemical bonding. For example, in the MO theory for [[hypervalent]] molecules it is unnecessary to invoke a major role for d-orbitals, whereas [[valence bond theory]] normally uses [[Orbital hybridisation|hybridization]] with d-orbitals to explain hypervalency. In MO theory, ''any'' electron in a molecule may be found ''anywhere'' in the molecule, since quantum conditions allow electrons to travel under the influence of an arbitrarily large number of nuclei, so long as permitted by certain quantum rules. Although in MO theory ''some'' molecular orbitals may hold electrons which are more localized between specific pairs of molecular atoms, ''other'' orbitals may hold electrons which are spread more uniformly over the molecule. Thus, overall, bonding (and electrons) are far more delocalized (spread out) in MO theory, than is implied in [[valence bond]] (VB) theory. This makes MO theory more useful for the description of extended systems. An example is that in the MO picture of [[benzene]], composed of a hexagonal ring of 6 carbon atoms. In this molecule, 24 of the 30 total valence bonding electrons are located in 12 σ (sigma) bonding orbitals which are mostly located between pairs of atoms (C-C or C-H), similar to the valence bond picture. However, in benzene the remaining 6 bonding electrons are located in 3 π (pi) molecular bonding orbitals that are delocalized around the ring. Two are in an MO which has equal contributions from all 6 atoms. The other two orbitals have vertical nodes at right angles to each other. As in the VB theory, all of these 6 delocalized pi electrons reside in a larger space which exists above and below the ring plane. All carbon-carbon bonds in benzene are chemically equivalent. In MO theory this is a direct consequence of the fact that the 3 molecular pi orbitals form a combination which evenly spreads the extra 6 electrons over 6 carbon atoms.
In molecules such as [[methane]], the 8 valence electrons are found in 4 MOs that are spread out over all 5 atoms. However, it is possible to approximate the MOs with 4 localized orbitals similar in shape to sp3 hybrid orbitals predicted by VB theory. This is often adequate for σ (sigma) bonds, but it is not possible for the π (pi) orbitals. However, the delocalized MO picture is more appropriate for ionization and spectroscopic predictions. Upon ionization of methane, a single electron is taken from the MO which surrounds the whole molecule, weakening all 4 bonds equally. VB theory would predict that one electron is removed for an sp3 orbital, resulting in the need for resonance between four valence bond structures, each of which has a one-electron bond. As in benzene, in substances such as beta carotene, chlorophyll or heme, some electrons the π (pi) orbitals are spread out in molecular orbitals over long distances in a molecule, giving rise to light absorption in lower energies (visible colors), a fact which is observed. This and other spectroscopic data for molecules are better explained in MO theory, with an emphasis on electronic states associated with multicenter orbitals, including mixing of orbitals premised on principles of orbital symmetry matching. The same MO principles also more naturally explain some electrical phenomena, such as high electrical conductivity in the planar direction of the hexagonal atomic sheets that exist in graphite. In MO theory, "resonance" (a mixing and blending of VB bond states) is a natural consequence of symmetry. For example, in graphite, as in benzene, it is not necessary to invoke the sp2 hybridization and resonance of VB theory, in order to explain electrical conduction. Instead, MO theory simply recognizes that some electrons in the graphite atomic sheets are completely delocalized over arbitrary distances, and reside in very large molecular orbitals that cover an entire graphite sheet, and some electrons are thus as free to move and conduct electricity in the sheet plane, as if they resided in a metal. <-->
- Daintith (2004). Oxford Dictionary of Chemistry Oxford University Press [S.l.] 0-19-860918-3.
- Licker, Mark (2004). McGraw-Hill Concise Encyclopedia of Chemistry McGraw-Hill [S.l.] 0-07-143953-6.
- Coulson, Charles, A. (1952). Valence (em inglês) Oxford at the Clarendon Press [S.l.]
- Robert Mulliken (1966). «Spectroscopy, Molecular Orbitals, and Chemical Bonding» (PDF) (em inglês).
- Lennard-Jones (1929). «Foundations of Molecular Orbital Theory» (em inglês).
- [http://www.ch.ic.ac.uk/vchemlib/course/mo_theory/main.html Introduction to Molecular Orbital Theory] - Imperial College London