Name: crystal field splitting in non-octahedral shapes
Uploaded: 2020-09-24T15:06:53.000Z
Duration: 2 min 46 s
Channel: JoVE

A teoria do campo de cristais pode ser utilizado para modelar o tetraédrico e o plano quadrado dos complexos metálicos de transição de uma forma análoga à aplicação desta teoria em complexos octaédricos. Por exemplo, para modelar o íon de tetraédrico tetracloroníquel(II)cada ligante de cloreto é substituído por uma carga de ponto negativo, resultando num campo de cristal tetraédrico. Devido à influência deste campo, os orbitais dxy, dyz e dxz são mais elevados em energia do que as orbitais dx²-y² e dz².Isto é atribuído à interação mais forte entre o campo tetraédrico de cristal e os orbitais dxy, dyz, e dxz. OS orbitais de alta energia Possuem uma simetria t₂ e são referidos como o conjunto t₂, enquanto os orbitais de baixo consumo de energia têm simetria e e compõem o conjunto e. Em comparação com a cisão dos orbitais d nos complexos octaédricos, as energias relativas dos orbitais nos complexos tetraédricos estão invertidos e a energia de divisão do campo de cristal, ou Δtet, é inferior.Em complexos planares quadrados tal como o íon de tetracitanoniquelado(II)todos os ligantes ficam no plano xy. Aqui, um campo de cristal plano quadrado é obtido através da substituição dos ligantes de cianeto com cargas de pontos negativos. Sob a influência deste campo, os orbitais d do íon metálico são divididos em quatro diferentes níveis de energia.Aqui, o orbital dx²-y² é o orbital com energia mais alta e tem lóbulos a apontar diretamente às cargas de ligante. O orbital dxy é o próximo em energia com lóbulos situados no mesmo plano como as cargas do ligante. O orbital dz² é ainda mais baixo em energia, atribuído a uma pequena sobreposição entre o orbital dz² e o campo de cristal no plano xy.O conjunto de orbitais de energia mais baixa, dxz e dyz, têm uma interação relativamente mínima com o campo de cristal. A divisão da energia campo de cristal em complexos de plano quadrado, ou Δsp, é definido como a diferença de energia entre o orbital de maior consumo de energia, dx²-y², e as orbitais de mais baixa energia, dyz e dxz. Assumindo o mesmo íon metálico e moléculas de ligante para todos os complexos, a proporção de Δtet, Δsp, e Δoct é de 0.44:1.7:1.

teoria do campo cristalino - complexos tetraédricos e quadrados planos

Crystal Field Theory - Tetrahedral and Square Planar Complexes

Complexos Tetraédricos
A teoria do campo cristalino (TCC) é aplicável a moléculas em geometrias não octaédricas. Em complexos octaédricos, os lóbulos das orbitais dx2−y2 e dz2 apontam diretamente para os ligandos. Para complexos tetraédricos, as orbitais d permanecem no lugar, mas com apenas quatro ligandos localizados entre os eixos. Nenhuma das orbitais aponta diretamente para os ligandos tetraédricos. No entanto, as orbitais dx2−y2 e dz2 (ao longo dos eixos Cartesianos) sobrepõem-se menos com os ligandos do que as orbitais dxy, dxz, e dyz. Por analogia com o caso octaédrico, o diagrama de energia para as orbitais d em um campo cristalino tetraédrico pode ser previsto como mostrado na Figura 1. Para evitar confusão, o conjunto octaédrico eg torna-se um conjunto tetraédrico e, e o conjunto octaédrico t2g torna-se um conjunto t2.
<img alt="Image1" src="/files/ftp_upload/11462/11462_Image_1.png" style="text-align: center; height: 350px;" />
Figura 1. Divisão das orbitais d do ião de metal sob campos cristalinos octaédricos e tetraédricos. Em comparação com o campo cristalino octaédrico, o padrão de divisão no campo cristalino tetraédrico é invertido. A energia da divisão de campo cristalino do complexo octaédrico, ou Δoct, é maior do que a energia de divisão de campo cristalino do complexo tetraédrico, Δtet .
Como a TCC é baseada na repulsão eletrostática, as orbitais mais próximas dos ligandos serão desestabilizadas e elevadas em energia em relação ao outro conjunto de orbitais. A divisão é menor que para complexos octaédricos porque a sobreposição é menor, então a energia de divisão de campo cristalino, ou Δtet é geralmente pequena.
Complexos Quadrados Planos
A outra geometria comum é a quadrada plana. É possível considerar uma geometria quadrada plana como uma estrutura octaédrica com um par de ligandos trans removidos. Presume-se que os ligandos removidos estejam no eixo z. Isto altera a distribuição das orbitais d, à medida que as orbitais no eixo z ou perto dele se tornam mais estáveis, e as nos eixos x ou y ou perto deles se tornam menos estáveis. Isto resulta na divisão dos conjuntos octaédricos t2g e eg e dá um padrão de divisão mais complicado (Figura 2).
<img alt="Image2" src="/files/ftp_upload/11462/11462_Image_2.png" style="text-align: center; height: 350px;" />
Figura 2. Divisão do conjunto t2g e do conjunto eg de orbitais em um campo cristalino quadrado plano. A energia de divisão de campo cristalino de complexos quadrados planos, ou Δsp, é maior que Δoct.
Este texto é adaptado de <a href="https://openstax.org/books/chemistry-2e/pages/19-3-spectroscopic-and-magnetic-properties-of-coordination-compounds">Openstax, Chemistry 2e, Section 19.3: </a><a href="https://openstax.org/books/chemistry-2e/pages/19-3-spectroscopic-and-magnetic-properties-of-coordination-compounds">S</a><a href="https://openstax.org/books/chemistry-2e/pages/19-3-spectroscopic-and-magnetic-properties-of-coordination-compounds">pectroscopic and Magnetic Properties of Coordination Compounds</a>.

Teoria do Campo Cristalino - Complexos Tetraédricos e Quadrados Planos

Complexos tetraédricos 
A teoria do campo cristalino (CFT) é aplicável a moléculas em geometrias diferentes do octaédrico. Nos complexos octaédricos, os lóbulos de ...

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crystal field splitting in non-octahedral shapes

Crystal field theory also explains d orbital splitting in tetrahedral and square planar complexes. These shapes do not follow the same pattern as an octahedral complex, so the d orbitals respond to ligand placement in different ways.

In tetrahedral complexes, the four ligands sit between the axes. None of the d orbitals points directly at the ligands, but the d x2−y2 and d z2 orbitals lie along the Cartesian axes and overlap less with the ligands than the d xy, d xz, and d yz orbitals. The d orbital energy diagram is therefore split in the opposite way from the octahedral case. The octahedral e g set becomes a tetrahedral e set, and the octahedral t 2 g set becomes a t 2 set. The splitting energy, called Δ tet, is usually small because the overlap is less than in octahedral complexes.

Square planar complexes can be viewed as octahedral structures with a pair of trans ligands removed. The removed ligands are assumed to be on the z-axis. This changes d orbital stability. Orbitals on or near the z-axis become more stable, while orbitals on or near the x- or y-axes become less stable. The octahedral t 2 g and e g sets split into a more complex pattern, and the crystal field splitting energy, Δ sp, is larger than Δ oct.

Crystal Field Splitting in Non-Octahedral Shapes

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properties of transition metals

Transition Metals in the Periodic Table

Transition metals are the d-block elements and the inner transition metals in the f-block. They are defined by partially filled d orbitals. In the periodic table, the transition elements are in groups 3 through 11, and the inner transition metals appear in the two rows below the main body of the table. The d-block is often grouped into transition series. The first transition series runs from Sc through Cu. The second series runs from Y through Ag. The third series includes La and Hf through...

coordination compounds and nomenclature

How Transition Metal Complexes Are Named

Transition metal complexes are named by following rules for ligands, metal ions, and oxidation state. These compounds form when a central metal ion or atom, often a transition metal or inner transition metal, accepts electron pairs from ligands. A ligand is a Lewis base that can be an atom, molecule, or ion with one or more lone pairs to donate. In these complexes, the bond is a coordinate covalent bond. That means both electrons in the bond come from the donor, such as ammonia, and are...

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Coordination Complexes and Chelation

Coordination complexes form when a metal ion bonds to ligands. These compounds appear in familiar examples such as hemoglobin in blood, chlorophyll in green plants, vitamin B-12, and the catalyst used to make polyethylene. Transition metal ions are especially likely to form these complexes. In a coordination complex, the metal acts as a Lewis acid. The ligand acts as a Lewis base and donates both electrons in the coordinate covalent bond. A ligand can be an atom, molecule, or ion, as long as...

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Transition Metal Shapes from Coordination Number

Transition metal complexes are often identified by their coordination number and the shape around the metal ion. The coordination number is the number of donor atoms bonded to the central metal. In many complexes, the most common shapes are octahedral, tetrahedral, and square planar. Coordination number and molecular geometry are linked in a clear pattern. A coordination number of 2 gives a linear shape, as in [Ag(NH3)2]+. A coordination number of 3 gives a trigonal planar shape, as in...

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Linkage and Ionization Isomers

Isomers are different chemical species that share the same chemical formula. In coordination compounds, structural isomerism includes linkage isomers and coordination-sphere isomers. These two forms show how the same atoms can be arranged in different ways around a metal center. Linkage isomers form when a ligand can attach to a transition metal through two different atoms. The cyanide ligand, CN−, can bind through carbon or nitrogen. The thiocyanate ligand, SCN−, can bind through sulfur or...

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Isomer Shapes in Coordination Complexes

Isomer shapes in coordination complexes can change a compound’s properties even when the chemical formula stays the same. Isomers are different chemical species with the same formula. In transition metal complexes, the atoms stay connected through the same types of bonds, but their arrangement in space can differ. One common form is geometric isomerism. In these complexes, ligands (the ions or molecules attached to the metal) can sit in cis or trans positions. The octahedral [Co(NH3)4Cl2]+ ion...

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Metal Ligand Bonds in Coordination Complexes

Coordination complexes form metal-ligand bonds through valence bond theory. In this model, Linus Pauling explained bonding with hybridization and overlap of atomic orbitals. The central metal atom or ion acts as a Lewis acid. It uses empty orbitals of suitable energy to accept electron pairs from ligands, which act as Lewis bases. These coordinate covalent bonds help explain the color, geometry, and magnetic behavior seen in coordination compounds. The type of hybridization helps determine the...

crystal field theory - octahedral complexes

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Octahedral d-Orbital Splitting in Crystal Fields

Crystal field theory explains how ligands affect the d orbitals in transition metal complexes. This electrostatic model helps predict the colors, magnetic behavior, and some structures of coordination compounds. It focuses on the nonbonding electrons on the central metal ion, not on the metal-ligand bonds. In crystal field theory, the metal ion and ligands are treated as very small point charges. Because electrons are negative, electrons donated by the ligands repel the electrons in the...

crystal field theory - tetrahedral and square planar complexes

Crystal field theory also explains d orbital splitting in tetrahedral and square planar complexes. These shapes do not follow the same pattern as an octahedral complex, so the d orbitals respond to ligand placement in different ways. In tetrahedral complexes, the four ligands sit between the axes. None of the d orbitals points directly at the ligands, but the d x2−y2 and d z2 orbitals lie along the Cartesian axes and overlap less with the ligands than the d xy, d xz, and d yz orbitals. The d...

colors and magnetism

Color and Spin in Coordination Complexes

Color in coordination complexes comes from how metal electrons absorb light. When atoms or molecules absorb light at the right frequency, electrons move to higher-energy orbitals. For many main-group atoms and molecules, the absorbed light is ultraviolet and cannot be seen. In coordination compounds, the energy gap between d orbitals often matches visible light, so the compound can absorb and emit colors that the human eye detects. Small changes in orbital energy can cause large shifts in the...

key terms and concepts

Explore how crystal field theory explains d orbital splitting in tetrahedral and square planar complexes, including inverted splitting and larger square planar splitting.

Crystal Field Splitting in Non-Octahedral Shapes | JoVE K12

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Crystal Field Splitting in Non-Octahedral Shapes

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