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Ligand
In chemistry, a '''ligand''' is either an atom, ion, or molecule (see also: functional group) that bonds to a central metal, generally involving formal donation of one or more of its electrons. The metal-ligand bonding ranges from covalent bond|covalent to more ionic. Furthermore, the metal-ligand bond order can range from one to three. Ligands are viewed as Lewis bases, although rare cases are known involving Lewis acidic "ligands."[Cotton, F. A. and Wilkinson, G., ''Advanced Inorganic Chemistry'', John Wiley and Sons: New York, 1988. ISBN 9780471199575]
Metal and metalloids are bound to ligands in virtually all circumstances, although gaseous "naked" metal ions can be generated in high vacuum. Ligands in a complex dictate the reactivity of the central atom, including ligand substitution rates, the reactivity of the ligands themselves, and redox. Ligand selection is a critical consideration in many practical areas, including bioinorganic chemistry|bioinorganic and medicinal chemistry, homogeneous catalysis, and environmental chemistry.
Ligands are classified in many ways: their charge, size (bulk), the identity of the coordinating atom(s), and the number of electrons donated to the metal (denticity or hapticity). The size of a ligand is indicated by its Ligand cone angle|cone angle.
Inner- vs out-sphere ligands
In coordination chemistry, the ligands that are directly bonded to the metal (that is, share electrons), are sometimes called "inner sphere" ligands. "Outer-sphere" ligands are not directly attached to the metal, but are bonded, generally weakly, to the first coordination shell, affecting the inner sphere in subtle ways. The complex of the metal with the inner sphere ligands is then called a coordination complex, which can be neutral, cationic, or anionic. The complex, along with its counter ions (if required), is called a coordination compound.
Strong field and weak field ligands
In general, ligands are viewed as donating electrons to the central atom. Bonding is often described using the formalisms of molecular orbital theory. In general, Lone electron pair|electron pairs occupy the HOMO of the ligands.
Ligands and metal ions can be ordered in many ways, one ranking system focuses on ligand 'hardness' (see also HSAB theory|hard soft acid base theory). Metal ions preferentially bind certain ligands. In general, 'hard' metal ions prefer weak field ligands, whereas 'soft' metal ions prefer strong field ligands. From a Molecular orbital theory|MO point of view, the HOMO/LUMO|HOMO of the ligand should have an energy that overlaps with the LUMO of the metal preferential. Metal ions bound to strong-field ligands follow the Aufbau principle, whereas complexes bound to weak-field ligands follow Hund's rule.
Binding of the metal with the ligands results in a set of molecular orbitals, where the metal can be identified with a new HOMO and LUMO (the orbitals defining the properties and reactivity of the resulting complex) and a certain ordering of the 5 d-orbitals (which may be filled, or partially filled with electrons). In an octahedral environment, the 5 otherwise degenerate d-orbitals split in sets of 2 and 3 orbitals (for a more in depth explanation, see crystal field theory).
- :3 orbitals of low energy: ''dxy'', ''dxz'' and ''dyz''
- :2 of high energy: ''d''''z''2 and ''d''''x''2−''y''2
The energy difference between these 2 sets of d-orbitals is called the splitting parameter, Δo. The magnitude of Δo is determined by the field-strength of the ligand: strong field ligands, by definition, increase Δo more than weak field ligands. Ligands can now be sorted according to the magnitude of Δo (see the table Ligand#Examples of common ligands (by field strength)|below). This ordering of ligands is almost invariable for all metal ions and is called spectrochemical series.
For complexes with a tetrahedral surrounding, the d-orbitals again split into two sets, but this time in reverse order:
- :2 orbitals of low energy: ''d''''z''2 and ''d''''x''2−''y''2
- :3 orbitals of high energy: ''d''''xy'', ''d''''xz'' and ''d''''yz''
The energy difference between these 2 sets of d-orbitals is now called Δt. The magnitude of Δt is smaller than for Δo, because in a tetrahedral complex only 4 ligands influence the d-orbitals, whereas in an octahedral complex the d-orbitals are influenced by 6 ligands. When the coordination number is neither octahedral nor tetrahedral, the splitting becomes correspondingly more complex. For the purposes of ranking ligands, however, the properties of the octahedral complexes and the resulting Δo has been of primary interest.
The arrangement of the d-orbitals on the central atom (as determined by the 'strength' of the ligand), has a strong effect on virtually all the properties of the resulting complexes. E.g. the energy differences in the d-orbitals has a strong effect in the optical absorption spectra of metal complexes. It turns out that valence electrons occupying orbitals with significant 3d-orbital character absorb in the 400-800 nm region of the spectrum (UV-visible range). The absorption of light (what we perceive as the color) by these electrons (that is, excitation of electrons from one orbital to another orbital under influence of light) can be correlated to the ground state of the metal complex, which reflects the bonding properties of the ligands. The relative change in (relative) energy of the d-orbitals as a function of the field-strength of the ligands is described in Tanabe-Sugano diagrams.
In cases where the ligand has low energy LUMO, such orbitals also participate in the bonding. The metal-ligand bond can be further stabilised by a formal donation of electron density back to the ligand in a process known as ''back-bonding.'' In this case a filled, central-atom-based orbital donates density into the LUMO of the (coordinated) ligand. Carbon monoxide is the preeminent example a ligand that engages metals via back-donation. Complementarily, ligands with low-energy filled orbitals of pi-symmetry can serve as pi-donor.
Polydentate and polyhapto ligand motifs and nomenclature
Denticity
Denticity (represented by κ) refers to the number of times a ligand bonds to a metal through non-contiguous donor sites. Many ligands are capable of binding metal ions through multiple sites, usually because the ligands have lone pairs on more than one atom. Ligands that bind via more than one atom are often termed ''Chelation|chelating''. A ligand that binds through two sites is classified as ''bidentate,'' and three sites as ''tridentate''. The "bite angle" refers to the angle between the two bonds of a bidentate chelate. Chelating ligands are commonly formed by linking donor groups via organic linkers. A classic ''bi''dentate ligand is ethylenediamine, which is derived by the linking of two ammonia groups with an ethylene (-CH2CH2-) linker. A classic example of a ''poly''dentate ligand is the hexadentate chelating agent EDTA, which is able to bond through six sites, completely surrounding some metals. The number of times a polydentate ligand bind to a metal centre is symbolized with "κn", where "n" indicates the number sites by which a ligand attaches to a metal. EDTA4−, when it is sexidentate, binds as a κ6-ligand, the amines and the carboxylate oxygen atoms are not contiguous. In practice, the n value of a ligand is not indicated explicitly but rather assumed. The binding affinity of a chelating system depends on the chelating angle or bite angle.
Complexes of polydentate ligands are called ''chelate'' complexes. They tend to be more stable than complexes derived from monodentate ligands. This enhanced stability, the chelate effect, is usually attributed to effects of entropy, which favors the displacement of many ligands by one polydentate ligand. When the chelating ligand forms a large ring that at least partially surrounds the central atom and bonds to it, leaving the central atom at the centre of a large ring. The more rigid and the higher its denticity, the more inert will be the macrocyclic complex. Heme is a good example: the iron atom is at the centre of a porphyrin macrocycle, being bound to four nitrogen atoms of the tetrapyrrole macrocycle. The very stable dimethylglyoximate complex of nickel is a synthetic macrocycle derived from the anion of dimethylglyoxime.
Hapticity
Hapticity (represented by η) refers to the number contiguous atoms that comprise a donor site and attach to a metal center. Butadiene forms both η2 and η4 complexes depending on the number of carbon atoms that are bonded to the metal.
Single atom bonding motifs
Ambidentate ligand
Unlike polydentate ligands, ambidentate ligands can attach to the central atom in two places but not both. A good example of this is thiocyanate, SCN−, which can attach at either the sulfur atom or the nitrogen atom. Such compounds give rise to linkage isomerism. Polyfunctional ligands, see especially proteins, can bond to a metal center through different ligand atoms to form various isomers.
Bridging ligand
Bridging ligand link two or more metal centers. Polyatomic ligands such as Carbonate|CO22- are especially prone to bridge. The bonding is complicated because polyatomic ligands are ambidentate and thus the capacity for many different linkage isomers. Atoms that bridge metals are sometimes indicated with prefix of "μ" (mu). Most inorganic solids, e.g. FeCl2, are polymers by virtue of the presence of multiple bridging ligands.
Metal ligand multiple bond
Metal ligand multiple bonds some ligands can bond to a metal center through the same atom but with a different number of lone pairs. The bond order of the metal ligand bond can be in part distinguished through the metal ligand bond angle (M-X-R). This bond angle is often referred to as being linear or bent with further discussion concerning the degree to which the angle is bent. For example, an imido ligand in the ionic form has three lone pairs. One lone pair is used as a sigma X donor, the other two lone pairs are available as L type pi donors. If both lone pairs are used in pi bonds then the M-N-R geometry is linear. However, if one or both these lone pairs is non-bonding then the M-N-R bond is bent and the extent of the bend speaks to how much pi bonding there may be. η1-Nitric oxide can coordinate to a metal center in linear or bent manner.
Specialized ligand types
Non-innocent ligand
Non-innocent ligands bond with metals in such a manner that the distribution of electron density between the metal center and ligand is unclear. Describing the bonding of noninnocent ligands often involves writing multiple Resonance (chemistry)|resonance forms which have partial contributions to the overall state.
Trans-spanning ligand
Trans-spanning ligands are bidentate ligands that can span coordination positions on opposite sides of a coordination complex.
Common ligands
- ''See Complex (chemistry)#Naming complexes|nomenclature.''
Virtually every molecule and every ion can serve as a ligand for (or "coordinate to") metals. Monodentate ligands include virtually all anions and all simple Lewis bases. Thus, the halides and pseudohalides are important anionic ligands whereas ammonia, carbon monoxide, and water are particularly common charge-neutral ligands. Simple organic species are also very common, be they anionic (alkoxide|RO− and Carboxylate|RCO2−) or neutral (Ether|R2O, Thioether|R2S, amine|R3−xNHx, and phosphine|R3P). The steric properties of some ligands are evaluated in terms of their cone angles.
Beyond the classical Lewis bases and anions, all unsaturated molecules are also ligands, utilizing their π-electrons in forming the coordinate bond. Also, metals can bind to the σ bonds in for example silanes, hydrocarbons, and dihydrogen (see also: agostic interaction).
In complexes of non-innocent ligands, the ligand is bonded to metals via conventional bonds, but the ligand is also redox-active.
Examples of common ligands (by field strength)
In the following table the ligands are sorted by field strength (weak field ligands first):
See also
- 2-Mercaptoindole
- Crystal field theory
- Ligand field theory
- Coordination chemistry
- Inorganic chemistry
- Radioligand
- Tanabe-Sugano diagram
- Spectrochemical series
- Scatchard equation
References
Category:Coordination chemistry
Category:Chemical bonding
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