| Square planar molecular geometry | |
|---|---|
 | |
| Examples | XeF4, PtCl2− 4 |
| Point group | D4h |
| Coordination number | 4 |
| Bond angle(s) | 90° |
| μ (Polarity) | 0 |
The square planar molecular geometry in chemistry describes the stereochemistry (spatial arrangement of atoms) that is adopted by certain chemical compounds. As the name suggests, molecules of this geometry have their atoms positioned at the corners.
Numerous compounds adopt this geometry, examples being especially numerous for transition metal complexes. The noble gas compound XeF4 adopts this structure as predicted by VSEPR theory. The geometry is prevalent for transition metal complexes with d8 configuration, which includes Rh(I), Ir(I), Pd(II), Pt(II), and Au(III). Notable examples include the anticancer drugs cisplatin [PtCl2(NH3)2] and carboplatin. Many homogeneous catalysts are square planar in their resting state, such as Wilkinson's catalyst and Crabtree's catalyst. Other examples include Vaska's complex and Zeise's salt. Certain ligands (such as porphyrins) stabilize this geometry.
A general d-orbital splitting diagram for square planar (D4h) transition metal complexes can be derived from the general octahedral (Oh) splitting diagram, in which the dz2 and the dx2−y2 orbitals are degenerate and higher in energy than the degenerate set of dxy, dxz and dyz orbitals. When the two axial ligands are removed to generate a square planar geometry, the dz2 orbital is driven lower in energy as electron-electron repulsion with ligands on the z-axis is no longer present. However, for purely σ-donating ligands the dz2 orbital is still higher in energy than the dxy, dxz and dyz orbitals because of the torus shaped lobe of the dz2 orbital. It bears electron density on the x- and y-axes and therefore interacts with the filled ligand orbitals. The dxy, dxz and dyz orbitals are generally presented as degenerate but they have to split into two different energy levels with respect to the irreducible representations of the point group D4h. Their relative ordering depends on the nature of the particular complex. Furthermore, the splitting of d-orbitals is perturbed by π-donating ligands in contrast to octahedral complexes. In the square planar case strongly π-donating ligands can cause the dxz and dyz orbitals to be higher in energy than the dz2 orbital, whereas in the octahedral case π-donating ligands only affect the magnitude of the d-orbital splitting and the relative ordering of the orbitals is conserved. [1]
A coordination complex consists of a central atom or ion, which is usually metallic and is called the coordination centre, and a surrounding array of bound molecules or ions, that are in turn known as ligands or complexing agents. Many metal-containing compounds, especially those of transition metals, are coordination complexes.
In coordination chemistry, a ligand is an ion or molecule that binds to a central metal atom to form a coordination complex. The bonding with the metal generally involves formal donation of one or more of the ligand's electron pairs often through Lewis Bases. The nature of metal–ligand bonding can range from covalent to ionic. Furthermore, the metal–ligand bond order can range from one to three. Ligands are viewed as Lewis bases, although rare cases are known to involve Lewis acidic "ligands".
In chemistry, the term transition metal has three possible definitions:
Crystal field theory (CFT) describes the breaking of degeneracies of electron orbital states, usually d or f orbitals, due to a static electric field produced by a surrounding charge distribution. This theory has been used to describe various spectroscopies of transition metal coordination complexes, in particular optical spectra (colors). CFT successfully accounts for some magnetic properties, colors, hydration enthalpies, and spinel structures of transition metal complexes, but it does not attempt to describe bonding. CFT was developed by physicists Hans Bethe and John Hasbrouck van Vleck in the 1930s. CFT was subsequently combined with molecular orbital theory to form the more realistic and complex ligand field theory (LFT), which delivers insight into the process of chemical bonding in transition metal complexes.
Ligand field theory (LFT) describes the bonding, orbital arrangement, and other characteristics of coordination complexes. It represents an application of molecular orbital theory to transition metal complexes. A transition metal ion has nine valence atomic orbitals - consisting of five nd, one (n+1)s, and three (n+1)p orbitals. These orbitals are of appropriate energy to form bonding interaction with ligands. The LFT analysis is highly dependent on the geometry of the complex, but most explanations begin by describing octahedral complexes, where six ligands coordinate to the metal. Other complexes can be described by reference to crystal field theory.
π backbonding, also called π backdonation, is a concept from chemistry in which electrons move from an atomic orbital on one atom to an appropriate symmetry antibonding orbital on a π-acceptor ligand. It is especially common in the organometallic chemistry of transition metals with multi-atomic ligands such as carbon monoxide, ethylene or the nitrosonium cation. Electrons from the metal are used to bond to the ligand, in the process relieving the metal of excess negative charge. Compounds where π backbonding occurs include Ni(CO)4 and Zeise's salt. IUPAC offers the following definition for backbonding:
A description of the bonding of π-conjugated ligands to a transition metal which involves a synergic process with donation of electrons from the filled π-orbital or lone electron pair orbital of the ligand into an empty orbital of the metal (donor–acceptor bond), together with release (back donation) of electrons from an nd orbital of the metal (which is of π-symmetry with respect to the metal–ligand axis) into the empty π*-antibonding orbital of the ligand.
The isolobal principle is a strategy used in organometallic chemistry to relate the structure of organic and inorganic molecular fragments in order to predict bonding properties of organometallic compounds. Roald Hoffmann described molecular fragments as isolobal "if the number, symmetry properties, approximate energy and shape of the frontier orbitals and the number of electrons in them are similar – not identical, but similar." One can predict the bonding and reactivity of a lesser-known species from that of a better-known species if the two molecular fragments have similar frontier orbitals, the highest occupied molecular orbital (HOMO) and the lowest unoccupied molecular orbital (LUMO). Isolobal compounds are analogues to isoelectronic compounds that share the same number of valence electrons and structure. A graphic representation of isolobal structures, with the isolobal pairs connected through a double-headed arrow with half an orbital below, is found in Figure 1.
Reductive elimination is an elementary step in organometallic chemistry in which the oxidation state of the metal center decreases while forming a new covalent bond between two ligands. It is the microscopic reverse of oxidative addition, and is often the product-forming step in many catalytic processes. Since oxidative addition and reductive elimination are reverse reactions, the same mechanisms apply for both processes, and the product equilibrium depends on the thermodynamics of both directions.
The Jahn–Teller effect is an important mechanism of spontaneous symmetry breaking in molecular and solid-state systems which has far-reaching consequences in different fields, and is responsible for a variety of phenomena in spectroscopy, stereochemistry, crystal chemistry, molecular and solid-state physics, and materials science. The effect is named for Hermann Arthur Jahn and Edward Teller, who first reported studies about it in 1937. The Jahn–Teller effect, and the related Renner–Teller effect, are discussed in Section 13.4 of the spectroscopy textbook by Bunker and Jensen.
In chemistry, octahedral molecular geometry describes the shape of compounds with six atoms or groups of atoms or ligands symmetrically arranged around a central atom, defining the vertices of an octahedron. The octahedron has eight faces, hence the prefix octa. The octahedron is one of the Platonic solids, although octahedral molecules typically have an atom in their centre and no bonds between the ligand atoms. A perfect octahedron belongs to the point group Oh. Examples of octahedral compounds are sulfur hexafluoride SF6 and molybdenum hexacarbonyl Mo(CO)6. The term "octahedral" is used somewhat loosely by chemists, focusing on the geometry of the bonds to the central atom and not considering differences among the ligands themselves. For example, [Co(NH3)6]3+, which is not octahedral in the mathematical sense due to the orientation of the N-H bonds, is referred to as octahedral.
A quintuple bond in chemistry is an unusual type of chemical bond, first reported in 2005 for a dichromium compound. Single bonds, double bonds, and triple bonds are commonplace in chemistry. Quadruple bonds are rarer but are currently known only among the transition metals, especially for Cr, Mo, W, and Re, e.g. [Mo2Cl8]4− and [Re2Cl8]2−. In a quintuple bond, ten electrons participate in bonding between the two metal centers, allocated as σ2π4δ4.
Copper proteins are proteins that contain one or more copper ions as prosthetic groups. Copper proteins are found in all forms of air-breathing life. These proteins are usually associated with electron-transfer with or without the involvement of oxygen (O2). Some organisms even use copper proteins to carry oxygen instead of iron proteins. A prominent copper proteins in humans is in cytochrome c oxidase (cco). The enzyme cco mediates the controlled combustion that produces ATP.
The 18-electron rule is a chemical rule of thumb used primarily for predicting and rationalizing formulas for stable transition metal complexes, especially organometallic compounds. The rule is based on the fact that the valence orbitals of transition metals consist of five d orbitals, one s orbital and three p orbitals which can collectively accommodate 18 electrons as either bonding or nonbonding electron pairs. This means that the combination of these nine atomic orbitals with ligand orbitals creates nine molecular orbitals that are either metal-ligand bonding or non-bonding. When a metal complex has 18 valence electrons, it is said to have achieved the same electron configuration as the noble gas in the period. The rule is not helpful for complexes of metals that are not transition metals, and interesting or useful transition metal complexes will violate the rule because of the consequences deviating from the rule bears on reactivity. The rule was first proposed by American chemist Irving Langmuir in 1921.
Tanabe–Sugano diagrams are used in coordination chemistry to predict absorptions in the UV, visible and IR electromagnetic spectrum of coordination compounds. The results from a Tanabe–Sugano diagram analysis of a metal complex can also be compared to experimental spectroscopic data. They are qualitatively useful and can be used to approximate the value of 10Dq, the ligand field splitting energy. Tanabe–Sugano diagrams can be used for both high spin and low spin complexes, unlike Orgel diagrams, which apply only to high spin complexes. Tanabe–Sugano diagrams can also be used to predict the size of the ligand field necessary to cause high-spin to low-spin transitions.
In coordination chemistry the bite angle is the ligand–metal–ligand bond angle of coordination complex containing a bidentate ligand. This geometric parameter is used to classify chelating ligands, including those in organometallic complexes. It is most often discussed in terms of catalysis, as changes in bite angle can affect not just the activity and selectivity of a catalytic reaction but even allow alternative reaction pathways to become accessible.
An electronic effect influences the structure, reactivity, or properties of molecule but is neither a traditional bond nor a steric effect. In organic chemistry, the term stereoelectronic effect is also used to emphasize the relation between the electronic structure and the geometry (stereochemistry) of a molecule.
The d electron count is a chemistry formalism used to describe the electron configuration of the valence electrons of a transition metal center in a coordination complex. The d electron count is an effective way to understand the geometry and reactivity of transition metal complexes. The formalism has been incorporated into the two major models used to describe coordination complexes; crystal field theory and ligand field theory, which is a more advanced version based on molecular orbital theory.
Spin states when describing transition metal coordination complexes refers to the potential spin configurations of the central metal's d electrons. In many these spin states vary between high-spin and low-spin configurations. These configurations can be understood through the two major models used to describe coordination complexes; crystal field theory and ligand field theory, which is a more advanced version based on molecular orbital theory.
Magnetochemistry is concerned with the magnetic properties of chemical compounds. Magnetic properties arise from the spin and orbital angular momentum of the electrons contained in a compound. Compounds are diamagnetic when they contain no unpaired electrons. Molecular compounds that contain one or more unpaired electrons are paramagnetic. The magnitude of the paramagnetism is expressed as an effective magnetic moment, μeff. For first-row transition metals the magnitude of μeff is, to a first approximation, a simple function of the number of unpaired electrons, the spin-only formula. In general, spin-orbit coupling causes μeff to deviate from the spin-only formula. For the heavier transition metals, lanthanides and actinides, spin-orbit coupling cannot be ignored. Exchange interaction can occur in clusters and infinite lattices, resulting in ferromagnetism, antiferromagnetism or ferrimagnetism depending on the relative orientations of the individual spins.
In inorganic chemistry, the cis effect is defined as the labilization of CO ligands that are cis to other ligands. CO is a well-known strong pi-accepting ligand in organometallic chemistry that will labilize in the cis position when adjacent to ligands due to steric and electronic effects. The system most often studied for the cis effect is an octahedral complex M(CO)
5X where X is the ligand that will labilize a CO ligand cis to it. Unlike trans effect, where this property is most often observed in 4-coordinate square planar complexes, the cis effect is observed in 6-coordinate octahedral transition metal complexes. It has been determined that ligands that are weak sigma donors and non-pi acceptors seem to have the strongest cis-labilizing effects. Therefore, the cis effect has the opposite trend of the trans-effect, which effectively labilizes ligands that are trans to strong pi accepting and sigma donating ligands.
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