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Ligand ChemistrySandeep Badarla

ContentsArticles

Ligand 1Crystal field theory 9Denticity 14Chelation 16Hapticity 20Trans-spanning ligand 23Linkage isomerism 24Bridging ligand 25Metal–ligand multiple bond 27Non-innocent ligand 29Chiral ligand 34Ligand dependent pathway 37Ligand field theory 37

ReferencesArticle Sources and Contributors 40Image Sources, Licenses and Contributors 41

Article LicensesLicense 42

Ligand 1

Ligand

Cobalt complex [HCo(CO)4] with five ligands

In coordination chemistry, a ligand is an ion or molecule (see also:functional group) that binds to a central metal atom to form acoordination complex. The bonding between metal and ligandgenerally involves formal donation of one or more of the ligand'selectron deficient pairs. The nature of metal-ligand bonding can rangefrom covalent to ionic. Furthermore, the metal-ligand bond order canrange from one to three. Ligands are viewed as Lewis bases, althoughrare cases are known involving Lewis acidic "ligands."[1][2]

Metals and metalloids are bound to ligands in virtually allcircumstances, although gaseous "naked" metal ions can be generatedin high vacuum. Ligands in a complex dictate the reactivity of thecentral atom, including ligand substitution rates, the reactivity of theligands themselves, and redox. Ligand selection is a critical consideration in many practical areas, includingbioinorganic and medicinal chemistry, homogeneous catalysis, and environmental chemistry.

Ligands are classified in many ways: their charge, their size (bulk), the identity of the coordinating atom(s), and thenumber of electrons donated to the metal (denticity or hapticity). The size of a ligand is indicated by its cone angle.

HistoryThe composition of coordination complexes have been known since the early 1800s, e.g. Prussian blue and coppervitriol. The key breakthrough occurred when Alfred Werner reconciled formulas and isomers. He showed, amongother things, that the formulas of many cobalt(III) and chromium(III) compounds can be understood if the metal hassix ligands in an octahedral geometry. The first to use the term "ligand" were Alfred Stock and Carl Somiesky, inrelation to silicon chemistry. The theory allows one to understand the difference between coordinated and ionicchloride in the cobalt ammine chlorides and to explain many of the previously inexplicable isomers. He resolved thefirst coordination complex called hexol into optical isomers, overthrowing the theory that chirality was necessarilyassociated with carbon compounds.[3][4]

Strong field and weak field ligandsIn general, ligands are viewed as electron donors and the metals as electron acceptors(citation. metals are electrondonors). Bonding is often described using the formalisms of molecular orbital theory. The HOMO (HighestOccupied Molecular Orbital) can be mainly of ligands or metal character.Ligands and metal ions can be ordered in many ways; one ranking system focuses on ligand 'hardness' (see alsohard/soft acid/base theory). Metal ions preferentially bind certain ligands. In general, 'soft' metal ions prefer weakfield ligands, whereas 'hard' metal ions prefer strong field ligands. According to the molecular orbital theory, theHOMO of the ligand should have an energy that overlaps with the LUMO (Lowest Unoccupied Molecular Orbital)of the metal preferential. Metal ions bound to strong-field ligands follow the Aufbau principle, whereas complexesbound 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 anew HOMO and LUMO (the orbitals defining the properties and reactivity of the resulting complex) and a certainordering of the 5 d-orbitals (which may be filled, or partially filled with electrons). In an octahedral environment, the5 otherwise degenerate d-orbitals split in sets of 2 and 3 orbitals (for a more in depth explanation, see crystal fieldtheory).

Ligand 2

3 orbitals of low energy: dxy, dxz and dyz2 of high energy: dz2 and dx2−y2

The energy difference between these 2 sets of d-orbitals is called the splitting parameter, Δo. The magnitude of Δo isdetermined by the field-strength of the ligand: strong field ligands, by definition, increase Δo more than weak fieldligands. Ligands can now be sorted according to the magnitude of Δo (see the table below). This ordering of ligandsis 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: dz2 and dx2−y23 orbitals of high energy: dxy, dxz and dyz

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 thed-orbitals are influenced by 6 ligands. When the coordination number is neither octahedral nor tetrahedral, thesplitting becomes correspondingly more complex. For the purposes of ranking ligands, however, the properties of theoctahedral 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 strongeffect on virtually all the properties of the resulting complexes. E.g. the energy differences in the d-orbitals has astrong effect in the optical absorption spectra of metal complexes. It turns out that valence electrons occupyingorbitals 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 oneorbital to another orbital under influence of light) can be correlated to the ground state of the metal complex, whichreflects the bonding properties of the ligands. The relative change in (relative) energy of the d-orbitals as a functionof 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-ligandbond can be further stabilised by a formal donation of electron density back to the ligand in a process known asback-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.

Metal-EDTA complex, whereinthe aminocarboxylate is a

hexadentate chelating ligand.

Classification of ligands as L and X

Especially in the area of organometallic chemistry, ligands are classified as L and X(or combinations of the two). The classification scheme - the "CBC Method" forCovalent Bond Classification - was popularized by M.L.H. Green and "is based onthe notion that there are three basic types [of ligands]... represented by the symbolsL, X, and Z, which correspond respectively to 2-electron, 1-electron and 0-electronneutral ligands."[5][6] L ligands are derived from charge-neutral precursors and arerepresented by amines, phosphines, CO, N2, and alkenes. X ligands typically arederived from anionic precursors such as chloride but includes ligands where salts ofanion do not really exist such as hydride and alkyl. Thus, the complexIrCl(CO)(PPh3)2 is classified as an MXL3 complex, since CO and the two PPh3ligands are classified as L's. The oxidative addition of H2 to IrCl(CO)(PPh3)2 givesan 18e- ML3X3 product, IrClH2(CO)(PPh3)2. EDTA4- is classified as an L2X4ligand, as it features four anions and two neutral donor sites. Cp is classified as anL2X ligand.[7]

Ligand 3

Cobalt(III) complex containingsix ammonia ligands, which aremonodentate. The chloride is not

a ligand.

Polydentate and polyhapto ligand motifs andnomenclature

Denticity

Denticity (represented by κ) refers to the number of times a ligand bonds to a metalthrough non-contiguous donor sites. Many ligands are capable of binding metal ionsthrough multiple sites, usually because the ligands have lone pairs on more than oneatom. Ligands that bind via more than one atom are often termed chelating. A ligandthat 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 commonlyformed by linking donor groups via organic linkers. A classic bidentate ligand is ethylenediamine, which is derivedby the linking of two ammonia groups with an ethylene (-CH2CH2-) linker. A classic example of a polydentateligand is the hexadentate chelating agent EDTA, which is able to bond through six sites, completely surroundingsome 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 hexidentate, binds as aκ6-ligand, the amines and the carboxylate oxygen atoms are not contiguous. In practice, the n value of a ligand is notindicated explicitly but rather assumed. The binding affinity of a chelating system depends on the chelating angle orbite angle.

Complexes of polydentate ligands are called chelate complexes. They tend to be more stable than complexes derivedfrom 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 largering that at least partially surrounds the central atom and bonds to it, leaving the central atom at the centre of a largering. The more rigid and the higher its denticity, the more inert will be the macrocyclic complex. Heme is a goodexample: the iron atom is at the centre of a porphyrin macrocycle, being bound to four nitrogen atoms of thetetrapyrrole macrocycle. The very stable dimethylglyoximate complex of nickel is a synthetic macrocycle derivedfrom the anion of dimethylglyoxime.

HapticityHapticity (represented by η) refers to the number of contiguous atoms that comprise a donor site and attach to ametal center. Butadiene forms both η2 and η4 complexes depending on the number of carbon atoms that are bondedto the metal.[7]

Ligand motifs

Outer-sphere ligandsIn coordination chemistry, the ligands that are directly bonded to the metal (that is, share electrons), form part of thefirst coordination sphere and are sometimes called "inner sphere" ligands. "Outer-sphere" ligands are not directlyattached to the metal, but are bonded, generally weakly, to the first coordination shell, affecting the inner sphere insubtle ways. The complex of the metal with the inner sphere ligands is then called a coordination complex, whichcan be neutral, cationic, or anionic. The complex, along with its counterions (if required), is called a coordinationcompound.

Ligand 4

Trans-spanning ligandsTrans-spanning ligands are bidentate ligands that can span coordination positions on opposite sides of a coordinationcomplex.[8]

Ambidentate ligandUnlike polydentate ligands, ambidentate ligands can attach to the central atom in two places but not both. A goodexample of this is thiocyanate, SCN−, which can attach at either the sulfur atom or the nitrogen atom. Suchcompounds give rise to linkage isomerism. Polyfunctional ligands, see especially proteins, can bond to a metal centerthrough different ligand atoms to form various isomers.

Bridging ligandA bridging ligand links two or more metal center. Virtually all inorganic solids with simple formulas arecoordination polymers, consisting of metal centres linked by bridging ligands. This group of materials includes allanhydrous binary metal halides and pseudohalides. Bridging ligands also persist in solution. Polyatomic ligands suchas carbonate are ambidentate and thus are found to often bind to two or three metals simultaneously. Atoms thatbridge metals are sometimes indicated with the prefix "μ" (mu). Most inorganic solids, are polymers by virtue of thepresence of multiple bridging ligands.

Metal–ligand multiple bondMetal ligand multiple bonds some ligands can bond to a metal center through the same atom but with a differentnumber of lone pairs. The bond order of the metal ligand bond can be in part distinguished through the metal ligandbond angle (M-X-R). This bond angle is often referred to as being linear or bent with further discussion concerningthe degree to which the angle is bent. For example, an imido ligand in the ionic form has three lone pairs. One lonepair is used as a sigma X donor, the other two lone pairs are available as L type pi donors. If both lone pairs are usedin pi bonds then the M-N-R geometry is linear. However, if one or both these lone pairs is non-bonding then theM-N-R bond is bent and the extent of the bend speaks to how much pi bonding there may be. η1-Nitric oxide cancoordinate to a metal center in linear or bent manner.

Specialized ligand types

Non-innocent ligandNon-innocent ligands bond with metals in such a manner that the distribution of electron density between the metalcenter and ligand is unclear. Describing the bonding of noninnocent ligands often involves writing multipleresonance forms that have partial contributions to the overall state.

Bulky ligandsBulky ligands are used to control the steric properties of a metal center. They are used for many reasons, bothpractical and academic. On the practical side, they influence the selectivity of metal catalysts, e.g. inhydroformylation. Of academic interest, bulky ligands stabilize unusual coordination sites, e.g. reactive coligands orlow coordination numbers. Often bulky ligands are employed to simulate the steric protection afforded by proteins tometal-containing active sites. Of course excessive steric bulk can prevent the coordination of certain ligands.

Ligand 5

Chiral ligandsChiral ligands are useful for inducing asymmetry within the coordination sphere. Often the ligand is employed as anoptically pure group. In some cases, e.g. secondary amines, the asymmetry arises upon coordination. Chiral ligandsare essential components of asymmetric homogeneous catalysis.

Common ligandsSee nomenclature.

Virtually every molecule and every ion can serve as a ligand for (or "coordinate to") metals. Monodentate ligandsinclude virtually all anions and all simple Lewis bases. Thus, the halides and pseudohalides are important anionicligands whereas ammonia, carbon monoxide, and water are particularly common charge-neutral ligands. Simpleorganic species are also very common, be they anionic (RO− and RCO2

−) or neutral (R2O, R2S, R3−xNHx, and 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 informing the coordinate bond. Also, metals can bind to the σ bonds in for example silanes, hydrocarbons, anddihydrogen (see also: agostic interaction).In complexes of non-innocent ligands, the ligand is bonded to metals via conventional bonds, but the ligand is alsoredox-active.

Examples of common ligands (by field strength)In the following table the ligands are sorted by field strength (weak field ligands first):

Ligand formula (bonding atom(s)in bold)

Charge Most common denticity Remark(s)

Iodide (iodo) I− monoanionic monodentate

Bromide (bromido) Br− monoanionic monodentate

Sulfide (thio or less commonly"bridging thiolate")

S2− dianionic monodentate (M=S), orbidentate bridging(M-S-M')

Thiocyanate (S-thiocyanato) S-CN− monoanionic monodentate ambidentate (see also isothiocyanate,below)

Chloride (chlorido) Cl− monoanionic monodentate also found bridging

Nitrate (nitrato) O-NO2− monoanionic monodentate

Azide (azido) N-N2− monoanionic monodentate

Fluoride (fluoro) F− monoanionic monodentate

Hydroxide (hydroxo) O-H− monoanionic monodentate often found as a bridging ligand

Oxalate (oxalato) [O-C(=O)-C(=O)-O]2− dianionic bidentate

Water (aqua) H-O-H neutral monodentate monodentate

Nitrite (nitrito) O-N-O− monoanionic monodentate ambidentate (see also nitro)

Isothiocyanate (isothiocyanato) N=C=S− monoanionic monodentate ambidentate (see also thiocyanate,above)

Acetonitrile (acetonitrilo) CH3CN neutral monodentate

Ligand 6

Pyridine C5H5N neutral monodentate

Ammonia (ammine or lesscommonly "ammino")

NH3 neutral monodentate

Ethylenediamine en neutral bidentate

2,2'-Bipyridine bipy neutral bidentate easily reduced to its (radical) anion oreven to its dianion

1,10-Phenanthroline phen neutral bidentate

Nitrite (nitro) N-O2− monoanionic monodentate ambidentate (see also nitrito)

Triphenylphosphine PPh3 neutral monodentate

Cyanide (cyano) CN− monoanionic monodentate can bridge between metals (bothmetals bound to C, or one to C and oneto N)

Carbon monoxide (carbonyl) CO neutral monodentate can bridge between metals (bothmetals bound to C)

Note: The entries in the table are sorted by field strength, binding through the stated atom (i.e. as a terminal ligand),the 'strength' of the ligand changes when the ligand binds in an alternative binding mode (e.g. when it bridgesbetween metals) or when the conformation of the ligand gets distorted (e.g. a linear ligand that is forced throughsteric interactions to bind in a non-linear fashion).

Other general encountered ligands (alphabetical)In this table other common ligands are listed in alphabetical order.

Ligand formula (bonding atom(s) in bold) Charge Mostcommondenticity

Remark(s)

Acetylacetonate (Acac) CH3-C(O)-CH2-C(O)-CH3 monoanionic bidentate In general bidentate,bound through bothoxygens, butsometimes boundthrough the centralcarbon only,see also analogousketimine analogues

Alkenes R2C=CR2 neutral compounds with aC-C double bond

Benzene C6H6 neutral and other arenes

1,2-Bis(diphenylphosphino)ethane (dppe) Ph2PC2H4PPh2 neutral bidentate

1,1-Bis(diphenylphosphino)methane (dppm) C25H22P2 neutral Can bond to 2 metalatoms at once,forming dimers

Corroles tetradentate

Crown ethers neutral primarily for alkaliand alkaline earthmetal cations

2,2,2-crypt hexadentate primarily for alkaliand alkaline earthmetal cations

Ligand 7

Cryptates neutral

Cyclopentadienyl (Cp) [C5H5]− monoanionic Althoughmonoanionic, by thenature of itsoccupied MO's, it iscapable of acting asa tridentate ligand.

Diethylenetriamine (dien) C4H13N3 neutral tridentate related to TACN,but not constrainedto facialcomplexation

Dimethylglyoximate (dmgH−) monoanionic

Ethylenediaminetetraacetate (EDTA) (HOOC-CH2)2N-(CH2)2-N(CH2-COOH)2 tetra-anionic hexadentate actual ligand is thetetra-anion

Ethylenediaminetriacetate trianionic pentadentate actual ligand is thetrianion

Ethyleneglycol-bis(oxyethylenenitrilo)-tetraacetate(EGTA)

(HOOC-CH2)2N-(CH2)2-O-(CH2)2-O-(CH2)2-N(CH2-COOH)2 tetra-anionic octodentate

glycinate (Glycinato) monoanionic bidentate other α-amino acidanions arecomparable (butchiral)

Heme dianionic tetradentate macrocyclic ligand

Nitrosyl NO+ cationic bent (1e) and linear(3e) bonding mode

Oxo O dianion monodentate sometimes bridging

Pyrazine N2C4H4 neutral ditopic sometimes bridging

Scorpionate ligand tridentate

Sulfite monoanionic monodentate ambidentate

2,2',5',2-Terpyridine (terpy) neutral tridentate meridional bondingonly

Triazacyclononane (tacn) (C2H4)3(NR)3 neutral tridentate macrocyclic ligandsee also theN,N',N"-trimethylatedanalogue

Tricyclohexylphosphine (C6H11)3P or (PCy3) neutral monodentate

Triethylenetetramine (trien) neutral tetradentate

Trimethylphosphine PMe3 neutral monodentate

Tri(o-tolyl)phosphine P(o-tolyl)3 neutral monodentate

Tris(2-aminoethyl)amine (tren) (NH2CH2CH2)3N neutral tetradentate

Tris(2-diphenylphosphineethyl)amine (np3) neutral tetradentate

Terpyridine C15H11N3 neutral tridentate

Tropylium C7H7+ cationic

Carbon dioxide CO2 see transition metalcarbon dioxidecomplex

Ligand 8

Ligand exchangeLigand exchange (also ligand substitution) is a type of chemical reaction in which one ligand in a chemicalcompound is replaced by another ligand. One type of pathway for substitution is the Ligand dependent pathway. Inorganometallic chemistry this can take place by associative substitution or by dissociative substitution. Another formof ligand exchange is seen in the nucleophilic abstraction reaction.

PronunciationPronounced ˈlɪgənd with the first syllable sounding like the word "lithium" or 'laɪgənd with the first syllablesounding like the word "lie". [9]

References[1] Cotton, Frank Albert; Geoffrey Wilkinson, Carlos A. Murillo (1999). Advanced Inorganic Chemistry. pp. 1355. ISBN 0-471-19957-5,

9780471199571.[2] Miessler, Gary L.; Donald Arthur Tarr (1999). Inorganic Chemistry. pp. 642. ISBN 0-13-841891-8, 9780138418915.[3] Jackson, W. Gregory; Josephine A. McKeon, Silvia Cortez (1 October 2004). "Alfred Werner's Inorganic Counterparts of Racemic and

Mesomeric Tartaric Acid: A Milestone Revisited". Inorganic Chemistry 43 (20): 6249–6254. doi:10.1021/ic040042e. PMID 15446870.[4] Bowman-James, Kristin (2005). "Alfred Werner Revisited: The Coordination Chemistry of Anions". Accounts of Chemical Research 38 (8):

671–678. doi:10.1021/ar040071t. PMID 16104690.[5] Green, M. L. H. (20 September 1995). "A new approach to the formal classification of covalent compounds of the elements". Journal of

Organometallic Chemistry 500 (1–2): 127–148. doi:10.1016/0022-328X(95)00508-N. ISSN 0022-328X.[6] http:/ / www. columbia. edu/ cu/ chemistry/ groups/ parkin/ mlxz. htm[7][7] Hartwig, J. F. Organotransition Metal Chemistry, from Bonding to Catalysis; University Science Books: New York, 2010. ISBN

1-891389-53-X[8][8] von Zelewsky, A. "Stereochemistry of Coordination Compounds" John Wiley: Chichester, 1995. ISBN 047195599.[9] "Ligand - Definition and more from the Free Merriam-Webster Dictionary" (http:/ / www. merriam-webster. com/ dictionary/

ligand?show=0& t=1303577063). Merriam Webster. . Retrieved 23 April 2011.

Crystal field theory 9

Crystal field theoryCrystal field theory (CFT) is a model that describes the breaking of degeneracies of electronic orbital states, usuallyd or f orbitals, due to a static electric field produced by a surrounding charge distribution (anion neighbors). Thistheory has been used to describe various spectroscopies of transition metal coordination complexes, in particularoptical spectra (colours). CFT successfully accounts for some magnetic properties, colours, hydration enthalpies, andspinel structures of transition metal complexes, but it does not attempt to describe bonding. CFT was developed byphysicists Hans Bethe and John Hasbrouck van Vleck[1] in the 1930s. CFT was subsequently combined withmolecular orbital theory to form the more realistic and complex ligand field theory (LFT), which delivers insight intothe process of chemical bonding in transition metal complexes.

Overview of crystal field theory analysisAccording to CFT, the interaction between a transition metal and ligands arises from the attraction between thepositively charged metal cation and negative charge on the non-bonding electrons of the ligand. The theory isdeveloped by considering energy changes of the five degenerate d-orbitals upon being surrounded by an array ofpoint charges consisting of the ligands. As a ligand approaches the metal ion, the electrons from the ligand will becloser to some of the d-orbitals and farther away from others causing a loss of degeneracy. The electrons in thed-orbitals and those in the ligand repel each other due to repulsion between like charges. Thus the d-electrons closerto the ligands will have a higher energy than those further away which results in the d-orbitals splitting in energy.This splitting is affected by the following factors:•• the nature of the metal ion.•• the metal's oxidation state. A higher oxidation state leads to a larger splitting.•• the arrangement of the ligands around the metal ion.• the nature of the ligands surrounding the metal ion. The stronger the effect of the ligands then the greater the

difference between the high and low energy d groups.The most common type of complex is octahedral; here six ligands form an octahedron around the metal ion. Inoctahedral symmetry the d-orbitals split into two sets with an energy difference, Δoct (the crystal-field splittingparameter) where the dxy, dxz and dyz orbitals will be lower in energy than the dz2 and dx2-y2, which will have higherenergy, because the former group is farther from the ligands than the latter and therefore experience less repulsion.The three lower-energy orbitals are collectively referred to as t

2g, and the two higher-energy orbitals as e

g. (These

labels are based on the theory of molecular symmetry). Typical orbital energy diagrams are given below in thesection High-spin and low-spin.Tetrahedral complexes are the second most common type; here four ligands form a tetrahedron around the metal ion.In a tetrahedral crystal field splitting the d-orbitals again split into two groups, with an energy difference of Δtetwhere the lower energy orbitals will be dz2 and dx2-y2, and the higher energy orbitals will be dxy, dxz and dyz -opposite to the octahedral case. Furthermore, since the ligand electrons in tetrahedral symmetry are not orienteddirectly towards the d-orbitals, the energy splitting will be lower than in the octahedral case. Square planar and othercomplex geometries can also be described by CFT.The size of the gap Δ between the two or more sets of orbitals depends on several factors, including the ligands andgeometry of the complex. Some ligands always produce a small value of Δ, while others always give a largesplitting. The reasons behind this can be explained by ligand field theory. The spectrochemical series is anempirically-derived list of ligands ordered by the size of the splitting Δ that they produce (small Δ to large Δ; seealso this table):I− < Br− < S2− < SCN− < Cl− < NO3

− < N3− < F− < OH− < C2O4

2− < H2O < NCS− < CH3CN < py < NH3 < en <2,2'-bipyridine < phen < NO2

− < PPh3 < CN− < CO

Crystal field theory 10

It is useful to note that the ligands producing the most splitting are those that can engage in metal to ligandback-bonding.The oxidation state of the metal also contributes to the size of Δ between the high and low energy levels. As theoxidation state increases for a given metal, the magnitude of Δ increases. A V3+ complex will have a larger Δ than aV2+ complex for a given set of ligands, as the difference in charge density allows the ligands to be closer to a V3+ ionthan to a V2+ ion. The smaller distance between the ligand and the metal ion results in a larger Δ, because the ligandand metal electrons are closer together and therefore repel more.

High-spin and low-spin

Low Spin [Fe(NO2)6]3− crystal field diagram

Ligands which cause a large splitting Δ of the d-orbitalsare referred to as strong-field ligands, such as CN− andCO from the spectrochemical series. In complexes withthese ligands, it is unfavourable to put electrons into thehigh energy orbitals. Therefore, the lower energyorbitals are completely filled before population of theupper sets starts according to the Aufbau principle.Complexes such as this are called "low spin". Forexample, NO2

− is a strong-field ligand and produces alarge Δ. The octahedral ion [Fe(NO2)6]3−, which has 5d-electrons, would have the octahedral splitting diagram shown at right with all five electrons in the t2g level.

High Spin [FeBr6]3− crystal field diagram

Conversely, ligands (like I− and Br−) which cause asmall splitting Δ of the d-orbitals are referred to asweak-field ligands. In this case, it is easier to putelectrons into the higher energy set of orbitals than it isto put two into the same low-energy orbital, becausetwo electrons in the same orbital repel each other. So,one electron is put into each of the five d-orbitalsbefore any pairing occurs in accord with Hund's ruleand "high spin" complexes are formed. For example,Br− is a weak-field ligand and produces a small Δoct. So, the ion [FeBr6]3−, again with five d-electrons, would havean octahedral splitting diagram where all five orbitals are singly occupied.

In order for low spin splitting to occur, the energy cost of placing an electron into an already singly occupied orbitalmust be less than the cost of placing the additional electron into an eg orbital at an energy cost of Δ. As noted above,eg refers to the dz2 and dx2-y2 which are higher in energy than the t2g in octahedral complexes. If the energy requiredto pair two electrons is greater than the energy cost of placing an electron in an eg, Δ, high spin splitting occurs.The crystal field splitting energy for tetrahedral metal complexes (four ligands) is referred to as Δtet, and is roughlyequal to 4/9Δoct (for the same metal and same ligands). Therefore, the energy required to pair two electrons istypically higher than the energy required for placing electrons in the higher energy orbitals. Thus, tetrahedralcomplexes are usually high-spin.The use of these splitting diagrams can aid in the prediction of the magnetic properties of coordination compounds.A compound that has unpaired electrons in its splitting diagram will be paramagnetic and will be attracted bymagnetic fields, while a compound that lacks unpaired electrons in its splitting diagram will be diamagnetic and willbe weakly repelled by a magnetic field.

Crystal field theory 11

Crystal field stabilization energyThe crystal field stabilization energy (CFSE) is the stability that results from placing a transition metal ion in thecrystal field generated by a set of ligands. It arises due to the fact that when the d-orbitals are split in a ligand field(as described above), some of them become lower in energy than before with respect to a spherical field known asthe barycenter in which all five d-orbitals are degenerate. For example, in an octahedral case, the t2g set becomeslower in energy than the orbitals in the barycenter. As a result of this, if there are any electrons occupying theseorbitals, the metal ion is more stable in the ligand field relative to the barycenter by an amount known as the CFSE.Conversely, the eg orbitals (in the octahedral case) are higher in energy than in the barycenter, so putting electrons inthese reduces the amount of CFSE.

Octahedral crystal field stabilization energy

If the splitting of the d-orbitals in anoctahedral field is Δoct, the three t2gorbitals are stabilized relative to thebarycenter by 2/5 Δoct, and the egorbitals are destabilized by 3/5 Δoct. Asexamples, consider the two d5

configurations shown further up thepage. The low-spin (top) example hasfive electrons in the t2g orbitals, so the total CFSE is 5 x 2/5 Δoct = 2Δoct. In the high-spin (lower) example, the CFSEis (3 x 2/5 Δoct) - (2 x 3/5 Δoct) = 0 - in this case, the stabilization generated by the electrons in the lower orbitals iscanceled out by the destabilizing effect of the electrons in the upper orbitals.

Crystal Field stabilization is applicable to transition-metal complexes of all geometries. Indeed, the reason that manyd8 complexes are square-planar is the very large amount of crystal field stabilization that this geometry produceswith this number of electrons.

Explaining the colors of transition metal complexesThe bright colors exhibited by many coordination compounds can be explained by Crystal Field Theory. If thed-orbitals of such a complex have been split into two sets as described above, when the molecule absorbs a photon ofvisible light one or more electrons may momentarily jump from the lower energy d-orbitals to the higher energy onesto transiently create an excited state atom. The difference in energy between the atom in the ground state and in theexcited state is equal to the energy of the absorbed photon, and related inversely to the wavelength of the light.Because only certain wavelengths (λ) of light are absorbed - those matching exactly the energy difference - thecompounds appears the appropriate complementary color.As explained above, because different ligands generate crystal fields of different strengths, different colors can beseen. For a given metal ion, weaker field ligands create a complex with a smaller Δ, which will absorb light of longerλ and thus lower frequency ν. Conversely, stronger field ligands create a larger Δ, absorb light of shorter λ, and thushigher ν. It is, though, rarely the case that the energy of the photon absorbed corresponds exactly to the size of thegap Δ; there are other things (such as electron-electron repulsion and Jahn-Teller effects) that also affect the energydifference between the ground and excited states.

Crystal field theory 12

Which colors are exhibited?

Color wheel

This color wheel demonstrates which color a compound will appear if it onlyhas one absorption in the visible spectrum. For example, if the compoundabsorbs red light, it will appear green.λ absorbed versus color observed400 nm Violet absorbed, Green-yellow observed (λ 560 nm)450 nm Blue absorbed, Yellow observed (λ 600 nm)490 nm Blue-green absorbed, Red observed (λ 620 nm)570 nm Yellow-green absorbed, Violet observed (λ 410 nm)580 nm Yellow absorbed, Dark blue observed (λ 430 nm)600 nm Orange absorbed, Blue observed (λ 450 nm)650 nm Red absorbed, Green observed (λ 520 nm)

Crystal field splitting diagrams

Crystal field splitting diagrams (π-acceptor ligands)

Octahedral Pentagonal bipyramidal Square antiprismatic

Square planar Square pyramidal Tetrahedral

Trigonal bipyramidal

Crystal field theory 13

Gallery

Octahedral Pentagonal bipyramidal Square planar Square pyramidal

Tetrahedral Trigonal bipyramidal Pentagonal pyramidal

References[1] Van Vleck, J. (1932). "Theory of the Variations in Paramagnetic Anisotropy Among Different Salts of the Iron Group". Physical Review 41:

208. Bibcode 1932PhRv...41..208V. doi:10.1103/PhysRev.41.208.

• Zumdahl, Steven S (2005). Chemical Principles (5th ed.). Houghton Mifflin Company. pp. 550–551, 957–964.ISBN 0-669-39321-5.

• Silberberg, Martin S (2006). Chemistry: The Molecular Nature of Matter and Change (4th ed.). New York:McGraw Hill Company. pp. 1028–1034. ISBN 0-8151-8505-7.

• Shriver, D. F.; Atkins, P. W. (2001). Inorganic Chemistry (4th ed.). Oxford University Press. pp. 227–236.ISBN 0-8412-3849-9.

• Housecroft, C. E.; Sharpe, A. G. (2004). Inorganic Chemistry (2nd ed.). Prentice Hall. ISBN 978-0130399137.• Miessler, G. L.; Tarr, D. A. (2003). Inorganic Chemistry (3rd ed.). Pearson Prentice Hall. ISBN 0-13-035471-6.

Denticity 14

Denticity

Atom withmonodentate ligands

Denticity refers to the number of atoms in a single ligand thatbind to a central atom in a coordination complex.[1][2] In manycases, only one atom in the ligand binds to the metal, so thedenticity equals one, and the ligand is said to be monodentate(sometimes called unidentate). Ligands with more than onebonded atom are called polydentate or multidentate.

The word denticity is derived from dentis, the Latin word fortooth. The ligand is thought of as biting the metal at one ormore linkage points. Denticity is distinguished from hapticity,in which electrons of a bond or conjugated series of bonds arelinked to the central metal without the metal-ligand bondbeing localized to a single ligand atom.

Classes of denticity

Polydentate ligands are chelating agents[3] and classified bytheir denticity. Some atoms cannot form the maximum possible number of bonds a ligand could make. In that caseone or more binding sites of the ligand are unused. Such sites can be used to form a bond with another chemicalspecies.

• Bidentate (also called didentate) ligands bind with two atoms, an example being ethylenediamine.

Structure of the pharmaceutical Oxaliplatin, which featurestwo different bidentate ligands.

• Tridentate ligands bind with three atoms, an example beingterpyridine. Tridentate ligands usually bind via two kindsof connectivity, called "mer" and "fac." Cyclic tridentateligands such as TACN and 9-ane-S3 bind in a facialmanner.

• Tetradentate ligands bind with four atoms, an examplebeing triethylenetetramine (abbreviated trien). Tetradentateligands bind via three connectivities depending on theirtopology and the geometry of the metal center. Foroctahedral metals, the linear tetradentate trien can bind viathree geometries. Tripodal tetradentate ligands, e.g.tris(2-aminoethyl)amine, are more constrained, and onoctahedra leave two cis sites. Many naturally occurring macrocyclic ligands are tetradentative, an example beingthe porphyrin in heme.

•• Pentadentate ligands bind with five atoms, an example being ethylenediaminetriacetic acid.• Hexadentate ligands bind with six atoms, an example being EDTA (although it can bind in a tetradentate manner).• Ligands of denticity greater than 6 are well known. The ligands

1,4,7,10-tetraazacyclododecane-1,4,7,10-tetraacetate (DOTA) and diethylene triamine pentaacetate (DTPA) areoctadentate. They are particularly useful for binding lanthanide ions, which typically have coordination numbersgreater than 6.

Denticity 15

Relationship between "linear" bi-, tri- and tetradentate ligands (red) bound to an octahedral metal center. The structures markedwith * are chiral owing to the backbone of the tetradentate ligand.

Stability constantsIn general, the stability of a metal complex correlates with the denticity of the ligands. The stability of a complex isrepresented quantitatively in the form of Stability constants. Hexadentate ligands tend to bind metal ions morestrongly than ligands of lower denticity.

External links• EDTA chelation lecture notes. [4] 2.4MB PDF - Slide 3 on denticity

References[1] IUPAC Gold Book denticity (http:/ / goldbook. iupac. org/ D01594. html)[2][2] von Zelewsky, A. "Stereochemistry of Coordination Compounds" John Wiley: Chichester, 1995. ISBN 047195599.[3] IUPAC Gold Book chelation (http:/ / goldbook. iupac. org/ C01012. html)[4] http:/ / people. chem. byu. edu/ dvd/ chem223/ lecture_notes/ 11_EDTA_Chelation. pdf/ at_download/ file

Chelation 16

Chelation

Metal-EDTA chelate

Chelation is the formation or presence of two or more separatecoordinate bonds between a polydentate (multiple bonded) ligand and asingle central atom.[1] Usually these ligands are organic compounds,and are called chelants, chelators, chelating agents, or sequesteringagents.

The ligand forms a chelate complex with the substrate. Chelatecomplexes are contrasted with coordination complexes composed ofmonodentate ligands, which form only one bond with the central atom.

Chelants, according to ASTM-A-380, are "chemicals that form soluble,complex molecules with certain metal ions, inactivating the ions so thatthey cannot normally react with other elements or ions to produceprecipitates or scale."

The word chelation is derived from Greek χηλή, chelè, meaning claw;the ligands lie around the central atom like the claws of a lobster.[2]

The chelate effect

Ethylenediamine ligand, binding to a centralmetal ion with two bonds

Cu2+ complexes with methylamine (left) andethylenediamine (right)

The chelate effect describes the enhanced affinity of chelating ligandsfor a metal ion compared to the affinity of a collection of similarnonchelating (monodentate) ligands for the same metal.Consider the two equilibria, in aqueous solution, between thecopper(II) ion, Cu2+ and ethylenediamine (en) on the one hand andmethylamine, MeNH2 on the other.

Cu2+ + en [Cu(en)]2+ (1)Cu2+ + 2 MeNH2 [Cu(MeNH2)2]2+ (2)

In (1) the bidentate ligand ethylene diamine forms a chelate complexwith the copper ion. Chelation results in the formation of afive–membered ring. In (2) the bidentate ligand is replaced by twomonodentate methylamine ligands of approximately the same donorpower, meaning that the enthalpy of formation of Cu—N bonds isapproximately the same in the two reactions. Under conditions of equalcopper concentrations and when the concentration of methylamine istwice the concentration of ethylenediamine, the concentration of thecomplex (1) will be greater than the concentration of the complex (2).The effect increases with the number of chelate rings so theconcentration of the EDTA complex, which has six chelate rings, ismuch much higher than a corresponding complex with two monodentate nitrogen donor ligands and fourmonodentate carboxylate ligands. Thus, the phenomenon of the chelate effect is a firmly established empirical fact.

Chelation 17

The thermodynamic approach to explaining the chelate effect considers the equilibrium constant for the reaction: thelarger the equilibrium constant, the higher the concentration of the complex.

[Cu(en)] =β11[Cu][en][Cu(MeNH2)2]= β12[Cu][MeNH2]2

Electrical charges have been omitted for simplicity of notation. The square brackets indicate concentration, and thesubscripts to the stability constants, β, indicate the stoichiometry of the complex. When the analytical concentrationof methylamine is twice that of ethylenediamine and the concentration of copper is the same in both reactions, theconcentration [Cu(en)] is much higher than the concentration [Cu(MeNH2)2] because β11 >> β12.An equilibrium constant, K, is related to the standard Gibbs free energy, ΔG by

ΔG = −RT ln K = ΔH − TΔS where R is the gas constant and T is the temperature in kelvins. ΔH is the standard enthalpy change of the reactionand ΔS is the standard entropy change.It has already been posited that the enthalpy term should be approximately the same for the two reactions. Thereforethe difference between the two stability constants is due to the entropy term. In equation (1) there are two particleson the left and one on the right, whereas in equation (2) there are three particles on the left and one on the right. Thismeans that less entropy of disorder is lost when the chelate complex is formed than when the complex withmonodentate ligands is formed. This is one of the factors contributing to the entropy difference. Other factors includesolvation changes and ring formation. Some experimental data to illustrate the effect are shown in the followingtable.[3]

Equilibrium log β ΔG ΔH /kJ mol−1 −TΔS /kJ mol−1

Cd2+ + 4 MeNH2 Cd(MeNH2)42+ 6.55 -37.4 -57.3 19.9

Cd2+ + 2 en Cd(en)22+ 10.62 -60.67 -56.48 -4.19

These data show that the standard enthalpy changes are indeed approximately equal for the two reactions and that themain reason for the greater stability of the chelate complex is due to the entropy term, which is much lessunfavourable, indeed, it is favourable in this instance. In general it is difficult to account precisely forthermodynamic values in terms of changes in solution at the molecular level, but it is clear that the chelate effect ispredominantly an effect of entropy.Other explanations, including that of Schwarzenbach,[4] are discussed in Greenwood and Earnshaw (loc.cit).

In natureVirtually all biochemicals exhibit the ability to dissolve certain metal cations. Thus, proteins, polysaccharides, andpolynucleic acids are excellent polydentate ligands for many metal ions. Organic compounds such as the amino acidsglutamic acid and histidine, organic diacids such as malate, and polypeptides such as phytochelatin are also typicalchelators. In addition to these adventitious chelators, several biomolecules are specifically produced to bind certainmetals (see next section).[5][6][7][8]

In biochemistry and microbiologyVirtually all metalloenzymes feature metals that are chelated, usually to peptides or cofactors and prostheticgroups.[8] Such chelating agents include the porphyrin rings in hemoglobin and chlorophyll. Many microbial speciesproduce water-soluble pigments that serve as chelating agents, termed siderophores. For example, species ofPseudomonas are known to secrete pyocyanin and pyoverdin that bind iron. Enterobactin, produced by E. coli, is thestrongest chelating agent known.

Chelation 18

In geologyIn earth science, chemical weathering is attributed to organic chelating agents, e.g. peptides and sugars, that extractmetal ions from minerals and rocks.[9] Most metal complexes in the environment and in nature are bound in someform of chelate ring, e.g. with a humic acid or a protein. Thus, metal chelates are relevant to the mobilization ofmetals in the soil, the uptake and the accumulation of metals into plants and microorganisms. Selective chelation ofheavy metals is relevant to bioremediation, e.g. removal of 137Cs from radioactive waste.[10]

ApplicationsChelators are used in producing nutritional supplements, fertilizers, chemical analysis, as water softeners,commercial products such as shampoos and food preservatives, medicine, heavy metal detox, and industrialapplications.In 2010, the Asia-Pacific region was the largest outlet, generating about 45% of worldwide demand for chelatingagents. The region was followed by Western Europe and North America. The global chelating agents market isexpected to reach more than 5 million tonnes in 2018.[11]

Nutritional supplementsIn the 1960's, scientists developed the concept of chelating a metal ion prior to feeding the element to the animal.They believed that this would create a neutral compound, protecting the mineral from being complexed withinsoluble salts within the stomach, rendering the metal unavailable for absorption. Amino acids, being effectivemetal binders, were chosen as the prospective ligands, and research was conducted on the metal-amino acidcombinations. The research supported that the metal-amino acid chelates were able to enhance mineral absorption.During this period, synthetic chelates were also being developed. An example of such synthetics isethylenediaminetetraacetic acid (EDTA). These synthetics applied the same concept of chelation and did createchelated compounds; however, these synthetics were too stable and not nutritionally viable. If the mineral was takenfrom the EDTA ligand, the ligand could not be used by the body and would be expelled. During the expulsionprocess the EDTA ligand will randomly chelate and strip another mineral from the body.[12]

According to the Association of American Feed Control Officials (AAFCO), a metal amino acid chelate is defined asthe product resulting from the reaction of a metal ion from a soluble metal salt with a mole ratio of one to three(preferably two) moles of amino acids. The average weight of the hydrolyzed amino acids must be approximately150 and the resulting molecular weight of the chelate must not exceed 800 Da.Since the early development of these compounds, much more research has been conducted, and has been applied tohuman nutrition products in a similar manner to the animal nutrition experiments that pioneered the technology.Ferrous bis-glycinate is an example of one of these compounds that has been developed for human nutrition. [13]

FertilizersMany mineral deficiencies can occur in plants, such as iron chlorosis, which can reduce the nutritional benefits ofcrops and eventually result in plant death. Mineral chelates have been used to alleviate the mineral deficiencies ofaffected crops through liquid foliar applications. These fertilizers are also used to prevent deficiencies fromoccurring and improving the overall health of the plants. [14]

Heavy metal detoxificationChelation therapy is the use of chelating agents to detoxify poisonous metal agents such as mercury, arsenic, and lead by converting them to a chemically inert form that can be excreted without further interaction with the body, and was approved by the U.S. Food and Drug Administration in 1991. In alternative medicine, chelation is used as a treatment for autism, although this practice is controversial due to the absence of scientific plausibility, lack of FDA

Chelation 19

approval, and its potentially deadly side-effects.[15]

Although they can be beneficial in cases of heavy metal poisoning, chelating agents can also be dangerous. Use ofdisodium EDTA instead of calcium EDTA has resulted in fatalities due to hypocalcemia.[16]

Other medical applicationsAntibiotic drugs of the tetracycline family are chelators of Ca2+ and Mg2+ ions.EDTA is also used in root canal treatment as a way to irrigate the canal. EDTA softens the dentin facilitating accessto the entire canal length and to remove the smear layer formed during instrumentation.Chelate complexes of gadolinium are often used as contrast agents in MRI scans.

Chemical applicationsHomogeneous catalysts are often chelated complexes. A typical example is the ruthenium(II) chloride chelated withBINAP (a bidentate phosphine) used in e.g. Noyori asymmetric hydrogenation and asymmetric isomerization. Thelatter has the practical use of manufacture of synthetic (–)-menthol.Citric acid is used to soften water in soaps and laundry detergents. A common synthetic chelator is EDTA.Phosphonates are also well known chelating agents. Chelators are used in water treatment programs and specificallyin steam engineering, e.g., boiler water treatment system: Chelant Water Treatment system.

Products such as Bio-Rust and Evapo-Rust are chelating agents sold for the removal of rust from iron and steel.

References[1] IUPAC definition of chelation. (http:/ / goldbook. iupac. org/ C01012. html)[2] The term chelate was first applied in 1920 by Sir Gilbert T. Morgan and H. D. K. Drew, who stated: "The adjective chelate, derived from the

great claw or chele (Greek) of the lobster or other crustaceans, is suggested for the caliperlike groups which function as two associating unitsand fasten to the central atom so as to produce heterocyclic rings."Morgan, Gilbert T.; Drew, Harry D. K. (1920). "CLXII.—Researches on residual affinity and co-ordination. Part II. Acetylacetones ofselenium and tellurium". J. Chem. Soc., Trans. 117: 1456. doi:10.1039/CT9201701456. (nonfree access)

[3] Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (http:/ / www. amazon. com/Chemistry-Elements-Second-Edition-Earnshaw/ dp/ 0750633654) (2nd ed.). Butterworth–Heinemann. ISBN 0080379419. . p 910

[4] Schwarzenbach, G (1952). "Der Chelateffekt". Helv. Chim. Acta 35 (7): 2344–2359. doi:10.1002/hlca.19520350721.[5] U Krämer, J D Cotter-Howells, J M Charnock, A H J M Baker, J A C Smith (1996). "Free histidine as a metal chelator in plants that

accumulate nickel". Nature 379 (6566): 635–638. doi:10.1038/379635a0.[6] Jurandir Vieira Magalhaes (2006). "Aluminum tolerance genes are conserved between monocots and dicots". Proc Natl Acad Sci USA 103

(26): 9749–9750. doi:10.1073/pnas.0603957103. PMC 1502523. PMID 16785425.[7] Suk-Bong Ha, Aaron P. Smith, Ross Howden, Wendy M. Dietrich, Sarah Bugg, Matthew J. O'Connell, Peter B. Goldsbrough, and

Christopher S. Cobbett (1999). "Phytochelatin synthase genes from Arabidopsis and the yeast Schizosaccharomyces pombe" (http:/ / www.plantcell. org/ cgi/ content/ full/ 11/ 6/ 1153?ck=nck). Plant Cell 11 (6): 1153–1164. doi:10.1105/tpc.11.6.1153. PMC 144235.PMID 10368185. .

[8] S. J. Lippard, J. M. Berg “Principles of Bioinorganic Chemistry” University Science Books: Mill Valley, CA; 1994. ISBN 0-935702-73-3.[9] Dr. Michael Pidwirny, University of British Columbia Okanagan, http:/ / www. physicalgeography. net/ fundamentals/ 10r. html[10][10] Prasad (ed). Metals in the Environment. University of Hyderabad. Dekker, New York, 2001[11] Market Study on Chelating Agents by Ceresana Research (http:/ / www. ceresana. com/ en/ market-studies/ additives/ chelating-agents).[12] Ashmead, H. DeWayne (1993). The Roles of Amino Acid Chelates in Animal Nutrition. Westwood: Noyes Publications.[13] Albion Laboratories, Inc. "Albion Ferrochel Website" (http:/ / www. albionferrochel. com). . Retrieved July 12, 2011.[14] Ashmead, et al., [ed], H. DeWayne (1986). Foliar Feeding of Plants with Amino Acid Chelates. Park Ridge: Noyes Publications.[15] Doja A, Roberts W (2006). "Immunizations and autism: a review of the literature". Can J Neurol Sci 33 (4): 341–46. PMID 17168158.[16] U.S. Centers for Disease Control, "Deaths Associated with Hypocalcemia from Chelation Therapy" (March 3, 2006), http:/ / www. cdc. gov/

mmwr/ preview/ mmwrhtml/ mm5508a3. htm

Hapticity 20

Hapticity

Ferrocene contains twoη5-cyclopentadienyl

ligands

The term hapticity is used to describe how a group of contiguous atoms of a ligand arecoordinated to a central atom. Hapticity of a ligand is indicated by the Greek character'eta', η. A superscripted number following the η denotes the number of contiguous atomsof the ligand that are bound to the metal. In general the η-notation is only used whenthere is more than one atom coordinated (otherwise the κ-notation is used, see alsohapticity vs. denticity).

History

The need for additional nomenclature for organometallic compounds became apparent inthe mid-1950s when Dunitz, Orgel, and Rich described the structure of the "sandwichcomplex" ferrocene by X-ray crystallography[1] where an iron atom is "sandwiched"between two parallel cyclopentadienyl rings. Cotton later proposed the term hapticityderived from the adjectival prefix hapto (from the Greek haptein, to fasten, denoting contact or combination) placedbefore the name of the olefin,[2] where the Greek letter η (eta) is used to denote the number of contiguous atoms of aligand that bind to a metal center. The term is usually employed to describe ligands containing extended π-systemsor where agostic bonding is not obvious from the formula.

Historically important compounds where the ligands are described with hapticity• Ferrocene - bis(η5-cyclopentadienyl)iron• Uranocene - bis(η8-1,3,5,7-cyclooctatetraene)uranium• W(CO)3(PPri

3)2(η2-H2 ) - the first compound to be synthesized with a dihydrogen ligand.[3][4]

• IrCl(CO)[P(C6H5)3]2(η2-O2) - the dioxygen derivative which forms reversibly upon oxygenation of Vaska'scomplex.

ExamplesThe η-notation is encountered in many coordination compounds:• Side-on bonding of molecules containing σ-bonds like H2:

• W(CO)3(PiPr3)2(η2-H2)[3][4]

• Side-on bonded ligands containing multiple bonded atoms, e.g. ethylene in Zeise's salt or with fullerene, which isbonded through donation of the π-bonding electrons:• K[PtCl3(η2-C2H4)].H2O

•• Related complexes containing bridging π-ligands:• (μ-η2:η2-C2H2)Co2(CO)6 and (Cp*2Sm)2(μ-η2:η2- N2)[5]

• Dioxygen in bis{(trispyrazolylborato)copper(II)}(μ-η2:η2-O2),Note that with some bridging ligands, an alternative bridging mode is observed, e.g. κ1,κ1, like in(Me3SiCH2)3V(μ-N2-κ1(N),κ1(N'))V(CH2SiMe3)3 contains a bridging dinitrogen molecule, where themolecule is end-on coordinated to the two metal centers (see hapticity vs. denticity).

• The bonding of π-bonded species can be extended over several atoms, e.g. in allyl, butadiene ligands, but also incyclopentadienyl or benzene rings can share their electrons.

• Apparent violations of the 18-electron rule sometimes are explicable in compounds with unusual hapticities:

Hapticity 21

• The 18-VE complex (η5-C5H5)Fe(η1-C5H5)(CO)2 contains one η5 bonded cyclopentadienyl, and one η1

bonded cyclopentadienyl.• Reduction of the 18-VE compound [Ru(η6-C6Me6)2]2+ (where both aromatic rings are bonded in an

η6-coordination), results in another 18VE compound: [Ru(η6-C6Me6)(η4-C6Me6)].• Examples of polyhapto coordinated heterocyclic and inorganic rings: Cr(η5-C4H4S)(CO)3 contains the sulfur

heterocycle thiophene and Cr(η6-B3N3Me6)(CO)3 contains a coordinated inorganic ring (B3N3 ring).

Electrons donated by "π- ligands" vs. hapticity

Ligand Electronscontributed

(neutral counting)

Electronscontributed

(ionic counting)

η1 Allyl 1 2

η3-Allylcyclopropenyl

3 4

η3-Allenyl 3 4

η2-Butadiene 2 2

η4-Butadiene 4 4

η1-cyclopentadienyl 1 2

η3-cyclopentadienyl 3 4

η5-cyclopentadienylpentadienylcyclohexadienyl

5 6

η2-Benzene 2 2

η4-Benzene 4 4

η6-Benzene 6 6

η7-Cycloheptatrienyl 7 6

η8-Cyclooctatetraenyl 8 10

Changes in hapticityThe hapticity of a ligand can change in the course of a reaction.[6] E.g. in a redox reaction:

Hapticity 22

Here one of the η6-benzene rings changes to a η4-benzene.Similarly hapticity can change during a substitution reaction:

Here the η5-cyclopentadienyl changes to an η3-cyclopentadienyl, giving room on the metal for an extra 2-electrondonating ligand 'L'. Removal of one molecule of CO and again donation of two more electrons by thecyclopentadienyl ligand restores the η5-cyclopentadienyl. The so-called indenyl effect also describes changes inhapticity in a substitution reaction.

Hapticity vs. denticityHapticity must be distinguished from denticity. Polydentate ligands coordinate via multiple coordination sites withinthe ligand. In this case the coordinating atoms are identified using the κ-notation, as for example seen in coordinationof 1,2-bis(diphenylphosphino)ethane (Ph2PCH2CH2PPh2), to NiCl2 asdichloro[ethane-1,2-diylbis(diphenylphosphane)-κ2P]nickel(II). If the coordinating atoms are contiguous (connectedto each other), the η-notation is used, as e.g. in titanocene dichloride:dichlorobis(η5-2,4-cyclopentadien-1-yl)titanium.[7]

Hapticity and fluxionalityMolecules with polyhapto ligands are often "fluxional", also known as stereochemically non-rigid. Two classes offluxionality are prevalent for organometallic complexes of polyhapto ligands:• Case 1, typically: when the hapticity value is less than the number of sp2 carbon atoms. In such situations, the

metal will often migrate from carbon to carbon, maintaining the same net hapticity. The η1-C5H5 ligand in(η5-C5H5)Fe( η1-C5H5)(CO)2 rearranges rapidly in solution such that Fe binds alternatingly to each carbon atomin the η1-C5H5 ligand. This reaction is degenerate and, in the jargon of organic chemistry, it is an example of asigmatropic rearrangement.

• Case 2, typically: complexes containing cyclic polyhapto ligands with maximized hapticity. Such ligands tend torotate. A famous example is ferrocene, Fe(η5-C5H5)2, wherein the Cp rings rotate with a low energy barrier aboutthe principal axis of the molecule that "skewers" each ring (see rotational symmetry). This "ring whizzing"explains, inter alia, why only one isomer can be isolated for Fe(η5-C5H4Br)2. In this case, the rotamers are notnecessarily degenerate, but the rotational barriers have low energies of activation.

Hapticity 23

References[1] J. Dunitz, L. Orgel, A. Rich (1956). "The crystal structure of ferrocene". Acta Crystallographica 9 (4): 373–5.

doi:10.1107/S0365110X56001091.[2] F. A. Cotton (1968). "Proposed nomenclature for olefin-metal and other organometallic complexes". J. Am. Chem. Soc. 90 (22): 6230–6232.

doi:10.1021/ja01024a059.[3] Kubas, Gregory J. (March 1988). "Molecular hydrogen complexes: coordination of a σ bond to transition metals" (http:/ / pubs. acs. org/ doi/

abs/ 10. 1021/ ar00147a005). Accounts of Chemical Research 21 (3): 120–128. doi:10.1021/ar00147a005. .[4] Kubas, Gregory J. (2001). Metal Dihydrogen and σ-Bond Complexes - Structure, Theory, and Reactivity (http:/ / books. google. com/

?id=SSn_OQAACAAJ) (1 ed.). New York: Kluwer Academic/Plenum Publishers. ISBN 978-0-306-46465-2. LCCN 00059283. .[5] D. Sutton (1993). "Organometallic diazo compounds". Chem. Rev. 93 (3): 995–1022. doi:10.1021/cr00019a008.[6] Huttner, Gottfried; Lange, Siegfried; Fischer, Ernst O. (1971). "Molecular Structure of Bis(Hexamethylbenzene)-Ruthenium(0)". Angewandte

Chemie, International Edition in English 10 (8): 556–557. doi:10.1002/anie.197105561.[7] "IR-9.2.4.1 Coordination Compounds: Describing the Constitution of Coordination Compounds: Specifying donor atoms: General" (http:/ /

www. iupac. org/ reports/ provisional/ abstract04/ RB-prs310804/ Chap9-3. 04. pdf). Nomenclature of Inorganic Chemistry –Recommendations 1990 (the 'Red Book') (http:/ / old. iupac. org/ reports/ provisional/ abstract04/ connelly_310804. html) (Draft March 2004ed.). IUPAC. 2004. p. 16. .

Trans-spanning ligandTrans-spanning ligands are bidentate ligands that can span opposite sites of a complex with square-planargeometry. A wide variety of ligands that chelate in the cis fashion already exist, but very few can link oppositevertices on a coordination polyhedron. Early attempts to generate trans-spanning bidentate ligands relied onpolymethylene chains to link the donor functionalities, but such ligands often lead to coordination polymers.

HistoryA diphosphane linked with pentamethylene was claimed to span across a square planar complex. This early attemptwas followed by ligands with more rigid backbones. "TRANSPHOS" was the first trans-spanning diphosphaneligand that usually coordinates to palladium(II) and platinum(II) in a trans manner. TRANSPHOS featuresbenzo[c]phenanthrene substituted by diphenylphosphinomethyl (Ph2PCH2) groups at the 1 and 11 positions.[1][2] Thepolycyclic framework suffers sterically clashing hydrogen centers.

Trans-spanning ligand 24

Xantphos, SPANphos, TRANSDIP and related ligandsXantphos is a trans-spanning ligand, without the steric problems associated with TRANSPHOS. SPANphos iscomparable to XANTPHOS but more reliably trans-spanning. TRANSDIP, based on a α-cyclodextrin, is the firstligand to give exclusively trans-spanned complexes, even with d8 metal ion halides.[3]

References[1] Destefano, N. J.; Johnson, D. K.; Lane, R. M.; Venanzi, L. M. (1976). "Transition-Metal Complexes with Bidentate Ligands Spanning

Trans-Positions. I. The Synthesis of 2,11-Bis(Diphenylphosphinomethyl)Benzo[C]-Phenanthrene, A Ligand Promoting the Formation ofSquare Planar Complexes". Helvetica Chimica Acta 59 (8): 2674–2682. doi:10.1002/hlca.19760590806.

[2] Mochida, J. A.; Mattern, J. C.; Bailar, Jr., J. (1975). "Stereochemistry of Complex Inorganic Compounds. XXXV. A Complex Containing aLigand That Spans Trans Positions". J. Am. Chem. Soc. 97 (11): 3021–3026. doi:10.1021/ja00844a017.

[3] Poorters, L.; Armspach, D.; Dominique, M.; Toupet, L.; Choua, S.; Turek, P. (2007). "Synthesis and Properties of Transdip, A Rigid ChelatorBuilt Upon a Cyclodextrin Cavity: is Transdip an Authentic Trans - Spanning Ligand?". Chemistry: A European Journal 13 (34): 9448–9461.doi:10.1002/chem.200700831. PMID 17943701.

Linkage isomerismLinkage isomerism is the existence of co-ordination compounds that have the same composition differing with theconnectivity of the metal to a ligand.Typical ligands that give rise to linkage isomers are:• thiocyanate, SCN-

• selenocyanate, SeCN-

• nitrite, NO2-

• sulfite, SO32-

Examples of linkage isomers are violet-colored [(NH3)5Co-SCN]2+ and orange-colored [(NH3)5Co-NCS]2+. Theisomerization of the S-bonded isomer to the N-bonded isomer occurs intramolecularly.[1] In the complex,dichlorotetrakis(dimethyl sulfoxide)ruthenium(II), linkage isomerism of dimethyl sulfoxide ligands can be observedin the NMR spectrum due to the effect of S vs. O bonding on the methyl groups of DMSO. The proper notation forlinkage isomerism is the kappa notation where the atom directly bonding to the metal is proceeded by the lowercaseGreek letter kappa; κ. For example, NO2

- is represented as nitrito-κ-N and nitrito-κ-O, replacing the old system oftrivial names such as nitro and nitroso.[2]

HistoryThe first reported example of linkage isomerism had the formula [Co(NH3)5(NO2)]Cl2. The cationic cobalt complexexists in two separable linkage isomers. In the yellow-coloured isomer, the nitro ligand is bound through nitrogen. Inthe red linkage isomer, the nitrito is bound through one oxygen atom. The O-bonded isomer is often written as[Co(NH3)5(ONO)]2+. Although the existence of the isomers had been known since the late 1800s, only in 1907 wasthe structural difference explained.[3] It was later shown that the red isomer converted to the yellow isomer uponUV-irradiation. In this particular example, the formation of the nitro isomer (Co-NO2) from the nitrito isomer(Co-ONO) occurs through the rearrangement of the molecular structure. Thus, no bonds are broken duringisomerization.

Linkage isomerism 25

Structures of the two linkage isomers of [Co(NH3)5(NO2)]2+.

References[1] Buckingham, D. A.; Creaser, I. I.; Sargeson, A. M. (1970). "Mechanism of Base Hydrolysis for CoIII(NH3)5X2+ Ions. Hydrolysis and

Rearrangement for the Sulfur-Bonded Co(NH3)5SCN2+ Ion". Inorg. Chem. 9 (3): 655–661. doi:10.1021/ic50085a044.<[2] IUPAC. Compendium of Chemical Terminology, 2nd ed. (the "Gold Book"). Compiled by A. D. McNaught and A. Wilkinson. Blackwell

Scientific Publications, Oxford (1997). XML on-line corrected version: http:/ / goldbook. iupac. org (2006-) created by M. Nic, J. Jirat, B.Kosata; updates compiled by A. Jenkins. ISBN 0-9678550-9-8.

[3] Werner, A. (1907). "Über strukturisomere Salze der Rhodanwasserstoffsäure und der salpetrigen Säure". Ber. 40 (1): 765–788.doi:10.1002/cber.190704001117.

Bridging ligandA bridging ligand is a ligand that connects two or more atoms, usually metal ions.[1] The ligand may be atomic orpolyatomic. Virtually all complex organic compounds can serve as bridging ligands, so the term is usually restrictedto small ligands such as pseudohalides or to ligands that are specifically designed to link two metals.In naming a complex wherein a single atom bridges two metals, the bridging ligand is preceded by the Greekcharacter 'mu', μ[2], with a superscript number denoting the number of metals bound to the bridging ligand is bound.μ2 is often denoted simply as μ.

An example of a μ2 bridging ligand

Illustrative bridging ligands

Virtually all ligands are known to bridge, with the exception ofamines and ammonia.[3] Particularly common inorganic bridgingligands include

• OH−,• O2−,

• S2−,• SH−,• NH2

−

• NH2− (imido)• N3− (nitrido)•• CO• Halides•• Hydride•• Cyanide

Bridging ligand 26

Cyanide usually bridges via M-NC-M' linkages, unlike the other entries on this list.Many organic ligands form strong bridges between metal centers. Many common examples include derivatives ofthe above inorganic ligands (R = alkyl, aryl):• OR−,• SR−,• NR2

−

• NR2− (imido)• P3− (phosphido)• PR2

− (phosphido, note the ambiguity with the preceding entry)• PR2− (phosphinidino)

Polyfunctional ligandsPolyfunctional ligands can attach to metals in many ways and thus can bridge metals in diverse ways, includingsharing of one atom or using several atoms. Examples of such polyatomic ligands are the oxoanions CO3

2− and therelated Carboxylate, PO4

3−, and the polyoxometallates. Several organophosphorus ligands have been developed thatbridge pairs of metals, a well-known example being Ph2PCH2PPh2.

Examples

Compound Formula Description

{(Fe(III)(OH2)4)2(µ-OH)2}4+ In this example hydroxide plays the role of a μ2 bridgingligand. Notice in the name of the compound μ2 has beensimplified to μ.[4]

(η6-C6H6)2Ru2Cl2(μ-Cl)2 In this particular complex, two chloride ligands are terminaland two are μ2 bridging. The η in the beginning of theformula denotes the hapticity of the benzene ligands.

B2H6 This classic borane compound, diborane features two μ2

bridging hydrides.

Bridging ligand 27

(Co(CO)3)3(μ3-(C-tBu)) This compound contains a μ3 bridging carbyne ligand(C-tBu).

See Also•• Bridging carbonyl

References[1] Nic, M.; Jirat, J.; Kosata, B., eds. (2006–). "bridging ligand" (http:/ / goldbook. iupac. org/ B00741. html). IUPAC Compendium of Chemical

Terminology (Online ed.). doi:10.1351/goldbook.{{{file}}}. ISBN 0-9678550-9-8. .[2] Nic, M.; Jirat, J.; Kosata, B., eds. (2006–). "µ- (mu)" (http:/ / goldbook. iupac. org/ M03659. html). IUPAC Compendium of Chemical

Terminology (Online ed.). doi:10.1351/goldbook.{{{file}}}. ISBN 0-9678550-9-8. .[3] Werner, H. (2004). "The Way into the Bridge: A New Bonding Mode of Tertiary Phosphanes, Arsanes, and Stibanes". Angew. Chem. Int. Ed.

43 (8): 938–954. doi:10.1002/anie.200300627. PMID 14966876.[4] Classifications of ligands. http:/ / chimge. unil. ch/ En/ complexes/ 1cpx7. htm

Metal–ligand multiple bondIn Chemistry, a metal–ligand multiple bond describes the interaction of certain ligands with a metal with a bondorder greater than one.[1] Coordination complexes featuring multiply bonded ligands are of both scholarly andpractical interest. Transition metal carbene complexes catalyze the olefin metathesis reaction. Metal oxointermediates are pervasive in oxidation catalysis. oxygen evolving complex.

Most common classes of complexes showingmetal ligand multiple bonds

As a cautionary note, the classification of a metal ligand bond as being"multiple" bond order is ambiguous and even arbitrary because bondorder is a formalism. Furthermore, the usage of multiple bonding is notuniform. Symmetry arguments suggest that most ligands engage metalsvia multiple bonds. The term 'metal ligand multiple bond" is oftenreserved for ligands of the type CRn and NRn (n = 0, 1, 2) and ORn (n =0, 1) where R is H or an organic substituent, or pseudohalide.Historically, CO and NO+ are not included in this classification, nor arehalides.

Metalligand multiple bond 28

Pi-donor ligandsIn coordination chemistry, a pi-donor ligand is a kind of ligand endowed with filled non-bonding orbitals thatoverlap with metal-based orbitals. Their interaction is complementary to the behavior of pi-acceptor ligands. Theexistence of terminal oxo ligands for the early transition metals is one consequence of this kind of bonding. Classicpi-donor ligands are oxide (O2-), nitride (N3-), imide (RN2-), alkoxide (RO-), amide (R2N-), and fluoride. For latetransition metals, strong pi-donors form anti-bonding interactions with the filled d-levels, with consequences for spinstate, redox potentials, and ligand exchange rates. Pi-donor ligands are low in the spectrochemical series.[1]

Multiple bond stabilizationMetals bound to so-called triply bonded carbyne, imide, nitride (nitrido), and oxide (oxo) ligands are generallyassigned to high oxidation states with low d electron counts. The high oxidation state stabilizes the highly reducedligands. The low d electron count allow for many bonds between ligands and the metal center. A d0 metal center canaccommodate up to 9 bonds without violating the 18 electron rule, whereas a d6 species can only accommodate 6bonds.

Reactivity explained through ligand hybridizationA ligand described in ionic terms can bond to a metal through however many lone pairs it has available. For examplemany alkoxides use one of their three lone pairs to make a single bond to a metal center. In this situation the oxygenis sp3 hybridized according to valence bond theory. Increasing the bond order to two by involving another lone pairchanges the hybridization at the oxygen to an sp2 center with an expected expansion in the M-O-R bond angle andcontraction in the M-O bond length. If all three lone pairs are included for a bond order of three than the M-O bonddistance contracts further and since the oxygen is a sp center the M-O-R bond angle is 180˚ or linear. Similarly withthe imidos are commonly referred to as either bent (sp2) or linear (sp). Even the oxo can be sp2 or sp hybridized. Thetriply bonded oxo, similar to carbon monoxide, is partially positive at the oxygen atom and unreactive towardsbronsted acids at the oxygen atom. When such a complex is reduced, the triple bond can be converted to a doublebond at which point the oxygen no longer bears a partial positive charge and is reactive towards acid.

Conventions

Bonding representationsImido ligands, also known as imides or nitrenes, most commonly form "linear six electron bonds" with metal centers.Bent imidos are a rarity limited by complexes electron count, orbital bonding availability, or some similarphenomenon. It is common to draw only two lines of bonding for all imidos, including the most common linearimidos with a six electron bonding interaction to the metal center. Similarly amido complexes are usually drawn witha single line even though most amido bonds involve four electrons. Alkoxides are generally drawn with a singlebond although both two and four electron bonds are common. Oxo can be drawn with two lines regardless ofwhether four electrons or six are involved in the bond, although it is not uncommon to see six electron oxo bondsrepresented with three lines.

Representing oxidation statesThere are two motifs to indicate a metal oxidation state based around the actual charge separation of the metalcenter. Oxidation states up to +3 are believed to be an accurate representation of the charge separation experiencedby the metal center. For oxidation states of +4 and larger, the oxidation state becomes more of a formalism withmuch of the positive charge distributed between the ligands. This distinction can be expressed by using a Romannumeral for the lower oxidation states in the upper right of the metal atomic symbol and an Arabic number with a

Metalligand multiple bond 29

plus sign for the higher oxidation states (see the example below). This formalism is not rigorously followed and theuse of Roman numerals to represent higher oxidation states is common.

[MIIILn]3+ vs. [O=M5+Ln]3+

References[1] "Metal–Ligand Multiple Bonds: The Chemistry of Transition Metal Complexes Containing Oxo, Nitrido, Imido, Alkylidene, or Alkylidyne

Ligands" W. A. Nugent and J. M. Mayer; Wiley-Interscience, New York, 1988.

Further reading (specialized literature)• Heidt, L.J.; Koster, G.F.; Johnson, A.H. "Experimental and Crystal Field Study of the Absorption Spectrum at

2000 to 8000 A of to Manganous Perchlorate in Aqueous Perchloric Acid" J. Am. Chem. Soc. 1959, 80,6471–6477.

• Rohde,J; In,J.; Lim, M.H.; Brennessel, W.W.; Bukowski, M.R.; Stubna, A.; Muonck, E.; Nam, W.; Que L."Crystallographic and Spectroscopic Characterization of a Nonheme Fe(IV)O Complex" Science VOL 2991037–1039.

• Decker, A.; Rohde,J.; Que, L.; Solomon, E.I. "Spectroscopic and Quantum Chemical Characterization of theElectronic Structure and Bonding in a Non-Heme FeIVO Complex" J. Am. Chem. Soc. 2004, 126, 5378–5379.

• Aliaga-Alcalde, N.; George, S.D.; Mienert, B.; Bill, E.; Wieghardt, K.; Neese, F. "The Geometric and ElectronicStructure of [(cyclam-acetato)Fe(N)]+: A Genuine Iron(V) Species with a Ground-State Spin S=1/2" Angew.Chem. Int. Ed. 2005, 44, 2908–2912.

Non-innocent ligandIn chemistry, a (redox) non-innocent ligand is a ligand in a metal complex where the oxidation state is unclear.Typically, complexes containing non-innocent ligands are redox active at mild potentials. The concept assumes thatredox reactions in metal complexes are either metal or ligand localized, which is a simplification, albeit a useful one.

Redox Reactions of Complexes Containing Innocent vs. Non-Innocent LigandsConventionally, redox reactions of coordination complexes are assumed to be metal-centered. The reduction ofMnO4

- to MnO42- is described by the change in oxidation state of manganese from 7+ to 6+. The oxide ligands do

not change in oxidation state, remaining 2- (a more careful examination of the electronic structure of the redoxpartners reveals however that the oxide ligands are affected by the redox change). Oxide is an innocent ligand.Another example of conventional metal-centered redox couple is [Co(NH3)6]3+/[Co(NH3)6]2+. Ammonia is innocentin this transformation.

A clear example of redox non-innocent behavior of ligands is observed for [Ni(S2C2Ph2)2]z, which exists in three oxidation states: z = 2-, 1-, and 0. If the ligands are always considered to be dianionic (as is done in formal oxidation state counting), then z = 0 requires that that nickel has a formal oxidation state of +IV. The formal oxidation state of the central nickel atom therefore ranges from +II to +IV in the above transformations (see Figure). However, the formal oxidation state is different from the real (spectroscopic) oxidation state based on the (spectroscopic) metal

Non-innocent ligand 30

d-electron configuration. The stilbene-1,2-dithiolate behaves as a redox non-innocent ligand, and the oxidationprocesses actually take place at the ligands rather than the metal. This leads to the formation of ligand radicalcomplexes. The charge-neutral complex (z =0) is therefore best described as a Ni2+ derivative of S2C2Ph2

-. Thediamagnetism of this complex arises from anti-ferromagnetic coupling between the unpaired electrons of the twoligand radicals.The complex Cr(2,2'-bipyridine)3 is a derivative of Cr(III) bound to three radical anions of 2,2'bipyridine, which is inthis case also behaving as a redox non-innocent ligand. On the other hand, one-electron oxidation of[Ru(2,2'-bipyridine)3]2+ is localized on Ru and the bipyridine is behaving as a normal, innocent ligand in this case.

HistoryC.K. Jørgenson (Cologny-Geneva) described ligands as "innocent" and "suspect": "Ligands are innocent when theyallow oxidation states of the central atoms to be defined. The simplest case of a suspect ligand is NO..."[1]

Redox non-innocent ligands have been intensively investigated spectroscopically by the groups of K. Wieghardt(MPI Mülheim a/d Ruhr) and W. Kaim (Stuttgart) over the past years. Quite recently it became obvious that redoxnon-innocent ligands are not just a spectroscopic curiosity, as the radical reactivity of redox non-innocent ligandswas demonstrated to play a crucial role in the mechanism of bio-catalytic processes mediated by severalmetallo-enzymes (e.g. Gallactose Oxidase, Cytochrome P450, methane mono-oxygenase). More recently, somesynthetic research groups have started to systematically investigate the (catalytic) reactivity of transition metalcomplexes with redox non-innocent ligands in organometallic chemistry.

Typical Ligands that often behave as Redox Non-Innocent Ligands• O2 and NO.[2]

Ligands with extended pi-delocalization such as porphyrins and phthalocyanines, ligands with the generalisedformulas [D-CR=CR-D]2- or D=CR-CR=D (D = O, S, NR’ and R, R' = alkyl or aryl), and similar related systems areoften non-innocent. For example:• dioxalenes, such as catecholates.[3]

• dithiolenes, such as 1,2-maleonitriledithiolate• diimines such as derivatives of 1,2-diaminobenzene, α-diimines, and dimethylglyoxime.•• pyridine-2,6-diimine ligands (relevant in polymerisation and hydrogenation catalysis).

The pyridine-2,6-diimine ligand can be easily reduced by one or two electrons.[4][5][6]

Non-innocent ligand 31

Redox Non-Innocent Ligands in Organometallic Chemistry and CatalysisIn paramagnetic organometallic complexes of Rh and Ir (metallo-radicals),[7] ethene ligands, amido ligands, and(reactive) carbene ligands are sometimes also behaving as 'redox non-innocent' ligands:

• Solvent coordination to some metallo-radical IrII(ethene) species transfers the spin-density from the metal to theredox non-innocent ethene ligand, after which direct radical coupling reactions with the olefinic ligand radicalbecome possible.[8][9][10]

• Oxidation of certain RhI-amido and IrI-amido complexes does not lead to the expected MII-amido species. Insteadthe unexpected MI-aminyl radical complexes are formed.[11]

• Carbene formation from diazo compounds at metallo-radical IrII species unexpectedly leads to formation of'carbene radicals'. This is a result of the redox non-innocent character of Fischer-type carbenes, whereone-electron reduction of the carbene ligand by IrII leads to formation of carbon centered 'carbene radicals'coordinated to IrIII. These 'carbene radicals' reveal interesting radical-type reactivities.[12]

Non-innocent ligand 32

Redox Non-Innocent Ligands in BiologyMetalloenzymes often feature non-innocent ligands. A common non-innocent ligand is found in metalloporphyrins.In the enzyme cytochrome P450, the porphyrin ligand sustains oxidation during the catalytic cycle. In other hemeproteins, such as myoglobin, ligand-centered redox does not occur and the porphyrin is innocent.

Galactose Oxidase (GOase) provides a seminal example for the involvement of reactive non-innocent ligands inbio-catalytic turnover.[13][14] GOase converts chemo-selectively primary alcohols with O2 into aldehydes and H2O2,with impressive turnover frequencies. The active site of the enzyme GOase contains a tyrosinyl radical which iscoordinated to a CuII ion. In the key steps of the catalytic cycle, a cooperative Brønsted-basic ligand-sitedeprotonates the alcohol, and subsequently the oxygen atom of the tyrosinyl radical abstracts a hydrogen atom fromthe alpha-CH functionality of the coordinated alcoholate substrate. Thus, the tyrosinyl radical is a reactive fragmentin the catalytic cycle which cooperates with the Cu site. This is essential for the function of the enzyme, because theCu-ion is only capable of one-electron transformations. It is the interplay of the 1e reactivity of the ligand radical andthe 1e reactivity of the metal which makes the overall process possible. The radical abstraction nature of the processmakes the process extremely fast. Anti-ferromagnetic coupling between the unpaired spins of the tyrosine radicalligand and the d9 CuII ion (open-shell singlet ground state) explains the observed diamagnetic nature of the restingstate of the enzyme, as was confirmed by synthetic model studies.[15]

The oxygen molecule in oxyhemoglobin (or oxymyoglobin) would appear to satisfy the definition of a non-innocentligand. Deoxyhemoglobin is ferrous and pentacoordinate, the (innocent) ligands being four N of the porphyrin andNe of the proximal histidine. An O2 molecule binds the sixth coordination position. There is evidence from anumber of lines that partial electron transfer from the Fe to O2 occurs, so that the complex is better described assuperoxide anion bound to ferric heme [16], although spin coupling makes the complex diamagnetic. The change inoxidation and spin state of the Fe results in a change of bond length to the five innocent ligands which results in theheme switching from a "domed" to a planar conformation, which in turn drives conformational changes in theprotein responsible for the cooperativity of O2 binding. This cooperativity is essential for efficient oxygen transport,so in a way we all owe our lives to the suspect nature of the O2 ligand!

Non-innocent ligand 33

References[1] Jørgensen, Chr. K. (1966). "Differences between the four halide ligands, and discussion remarks on trigonal-bipyramidal complexes, on

oxidation states, and on diagonal elements of one-electron energy". Coordination Chemistry Reviews 1 (1-2): 164–178.doi:10.1016/S0010-8545(00)80170-8.

[2] Kaim, W.; Schwederski, B. (2010). "Non-innocent ligands in bioinorganic chemistry—An overview". Coordination Chemistry Reviews. 254(13-14) (13-14): 1580–1588. doi:10.1016/j.ccr.2010.01.009.

[3] Piero Zanello, P.; Corsini, M. (2006). "Homoleptic, mononuclear transition metal complexes of 1,2-dioxolenes: Updating theirelectrochemical-to-structural (X-ray) properties". Coordination Chemistry Reviews 250 (15-16): 2000–2022. doi:10.1016/j.ccr.2005.12.017.

[4] de Bruin, B.; Bill, E.; Bothe, E.; Weyhermüller, T.; Wieghardt, K. (2000). "Molecular and Electronic Structures ofBis(pyridine-2,6-diimine)metal Complexes [ML2](PF6)n(n = 0, 1, 2, 3; M = Mn, Fe, Co, Ni, Cu, Zn)". Inorganic Chemistry 39 (13):2936–2947. doi:10.1021/ic000113j.

[5] Budzelaar, P.H.M.; de Bruin, B.; Gal, A.W.; Wieghardt, K.; van Lenthe, J.H. (2001). "Metal-to-Ligand Electron Transfer in DiiminopyridineComplexes of Mn−Zn. A Theoretical Study". Inorganic Chemistry 40 (18): 4649–4655. doi:10.1021/ic001457c.

[6] Chirik, P.J.; Wieghardt, K. (2010). "Radical Ligands Confer Nobility on Base-Metal Catalysts". Science. 327 (5967) (5967): 794–795.doi:10.1126/science.1183281. PMID 20150476.

[7] de Bruin, B.; Hetterscheid, D.G.H.; Koekkoek, A.J.J.; Grützmacher, H. (2007). 5. In Karlin, Kenneth D.. "The Organometallic Chemistry ofRh-, Ir-, Pd-, and Pt-Based Radicals: Higher Valent Species". Progress in Inorganic Chemistry 55: 247–354. doi:10.1002/9780470144428.

[8] Hetterscheid, D.G.H.; Kaiser, J.; Reijerse, E.; Peters, T.P.J.; Thewissen, S.; Blok, A.N.J.; Smits, J.M.M.; de Gelder, R.; de Bruin, B. (2005)."IrII(ethene): Metal or Carbon Radical?". Journal of the American Chemical Society 127 (6): 1895–1905. doi:10.1021/ja0439470.PMID 15701024.

[9] Hetterscheid, D.G.H.; Bens, M.; de Bruin, B. (2005). "IrII(ethene): Metal or Carbon Radical? Part II: Oxygenation via iridium or directoxygenation at ethene?". Dalton Transactions (5): 979–984. doi:10.1039/b417766e. PMID 15726153.

[10] de Bruin, B.; Hetterscheid, D.G.H. (2007). "Paramagnetic (Alkene)Rh and (Alkene)Ir Complexes: Metal or Ligand Radicals?". EuropeanJournal of Inorganic Chemistry 2 (2): 211–230. doi:10.1002/ejic.200600923.

[11] Büttner, T.; Geier, J.; Frison, G.; Harmer, J.; Calle, C.; Schweiger, A.; Schönberg, H.; Grützmacher, H. (2005). "A Stable Aminyl RadicalMetal Complex". Science. 307 (5707) (5707): 235–238. doi:10.1126/science.1106070. PMID 15653498.

[12] Dzik, W.I.; Reek, J.N.H.; de Bruin, B. (2008). "Selective C-C Coupling of Ir-Ethene and Ir-Carbenoid Radicals". Chemistry: A EuropeanJournal 14 (25): 7594–7599. doi:10.1002/chem.200800262. PMID 18523935.

[13] Whittaker, M.M.; Whittaker, J.W. (1993). "Ligand interactions with galactose oxidase: mechanistic insights". Biophysical Journal. 64 (3):762–772.

[14] Wang, Y.; DuBois, J. L.; Hedman, B.; Hodgson, K. O.; Stack, T. D. P. (1998). "Catalytic Galactose Oxidase Models: BiomimeticCu(II)-Phenoxyl-Radical Reactivity". Science. 279 (5350) (5350): 537–540. doi:10.1126/science.279.5350.537.

[15] Müller, J.; Weyhermüller, T. Bill, E.; Hildebrandt, P.; Ould-Moussa, L.; Glaser, T.; Wieghardt, K. (1998). "Why Does the Active Form ofGalactose Oxidase Possess a Diamagnetic Ground State?". Angewandte Chemie International Edition. 37 (5) (5): 616–619.doi:10.1002/(SICI)1521-3773(19980316)37:5<616::AID-ANIE616>3.0.CO;2-4.

[16] Hui Chen†, Masao Ikeda-Saito and Sason Shaik (2008). "Nature of the Fe−O2 Bonding in Oxy-Myoglobin: Effect of the Protein". J. Am.Chem. Soc. (130): 14778–14790. doi:10.1021/ja805434m.

Chiral ligand 34

Chiral ligandIn chemistry a chiral ligand is a specially adapted ligand used for asymmetric synthesis. This ligand is anenantiopure organic compound which combines with a metal center by chelation to form an asymmetric catalyst.This catalyst engages in a chemical reaction and transfers its chirality to the reaction product which as a result alsobecomes chiral. In an ideal reaction one equivalent of catalyst can turn over many more equivalents of reactantwhich enables the synthesis of a large amount of a chiral compound from achiral precursors with the aid of a verysmall (often expensive) chiral ligand.

First discoveryThe first such ligand, the diphosphine DiPAMP was developed in 1968 by William S. Knowles of MonsantoCompany, who won the 2001 Nobel Prize in Chemistry,[1] and ultimately used in the industrial production ofL-DOPA.

Privileged ligandsMany thousands of chiral ligands have been prepared and tested since then but only several compound classes havebeen found to have a general scope. These ligands are therefore called privileged ligands.[2][3] Important membersdepicted below are BINOL, BINAP, TADDOL, DIOP, BOX and DuPhos (a phosphine ligand), all available asenantiomeric pairs.

Other members are Salen, cinchona alkaloids and phosphoramidites. Many of these ligands possess C2 symmetrywhich limits the number of possible reaction pathways and thereby increasing enantioselectivity.

Chiral ligand 35

Chiral fence

Chiral ligands work asymmetricinduction somewhere along thereaction coordinate. The imagedepicted on the right gives a generalidea how a chiral ligand may induce anenantioselective reaction. The ligand(in green) has C2 symmetry with itsnitrogen, oxygen or phosphorus atomshugging a central metal atom (in red).In this particular ligand the right side issticking out and its left side pointsaway. The substrate in this reduction isacetophenone and the reagent (in blue)a hydride ion. In absence of the metaland the ligand the re face approach ofthe hydride ion gives the(S)-enantiomer and the si faceapproach the (R)-enantiomer in equal amounts (a racemic mixture like expected). The ligand/metal presence changesall that. The carbonyl group will coordinate with the metal and due to the steric bulk of the phenyl group it will onlybe able to do so with its si face exposed to the hydride ion with in the ideal situation exclusive formation of the (R)enantiomer. The re face will simple hit the chiral fence.[4] Note that when the ligand is replaced by its mirror imagethe other enantiomer will form and that a racemic mixture of ligand will once again yield a racemic product. Alsonote that if the steric bulk of both carbonyl substituents is very similar the strategy will fail.

Chiral ligand 36

Chiral counterionsIn a novel concept, so-called chiral ions team up with traditional cationic catalysts in asymmetric synthesis asdemonstrated in this allene hydroxyalkoxylation in which the active catalyst is a salt of gold(I) and a phosphate of achiral binaphthol:[5][6]

References[1] Nobel prize 2001 www.nobelprize.org Link (http:/ / nobelprize. org/ nobel_prizes/ chemistry/ laureates/ 2001/ public. html)[2] Design of chiral ligands for asymmetric catalysis: From C2-symmetric P,P- and N,N-ligands to sterically and electronically nonsymmetrical

P,N-ligands Andreas Pfaltz and William J. Drury III PNAS, April 20, 2004 vol. 101 no. 16 5723-5726 doi:10.1073/pnas.0307152101[3] Privileged Chiral Catalysts Tehshik P. Yoon, Eric N. Jacobsen Science 14 March 2003: Vol. 299. no. 5613, pp. 1691 - 1693

doi:10.1126/science.1083622 PMID 12637734[4] Chiral and C2-symmetrical bis(oxazolinylpyridine)rhodium(III) complexes: effective catalysts for asymmetric hydrosilylation of ketones

Hisao Nishiyama, Hisao Sakaguchi, Takashi Nakamura, Mihoko Horihata, Manabu Kondo, and Kenji Itoh Organometallics; 1989; 8(3) pp846 - 848; doi: 10.1021/om00105a047

[5] A Powerful Chiral Counterion Strategy for Asymmetric Transition Metal Catalysis Gregory L. Hamilton, Eun Joo Kang, Miriam Mba, F.Dean Toste. Science 317, 496 (2007) doi: 10.1126/science.1145229

[6] Starting catalyst: 1,2-bis(diphenylphosphino)ethane (dppm) gold(I) chloride complex

Ligand dependent pathway 37

Ligand dependent pathwayThere are two types of pathway for substitution of ligands in a complex. The ligand dependent pathway is the onewhereby the chemical properties of the ligand affect the rate of substitution. Alternatively, there is the ligandindependent pathway, which is where the ligand does not have an effect.This is of vital importance in the world of inorganic chemistry and complex ions.

Ligand field theoryLigand field theory (LFT) describes the bonding, orbital arrangement, and other characteristics of coordinationcomplexes.[1] It represents an application of molecular orbital theory to transition metal complexes. A transitionmetal ion has nine valence atomic orbitals, five (n)d, one (n+1)s, and three (n+1)p orbitals. These orbitals are ofappropriate energy to form bonding interaction with ligands. The LFT analysis is highly dependent on the geometryof the complex, but most explanations begin by describing octahedral complexes, where six ligands coordinate to themetal.[2]

σ-BondingThe molecular orbitals created by coordination can be seen as resulting from the donation of two electrons by each ofsix σ-donor ligands to the d-orbitals on the metal. In octahedral complexes, ligands approach along the x-, y- andz-axes, so their σ-symmetry orbitals form bonding and anti-bonding combinations with the dz2 and dx2−y2 orbitals.The dxy, dxz and dyz orbitals remain non-bonding orbitals. Some weak bonding (and anti-bonding) interactions withthe s and p orbitals of the metal also occur, to make a total of 6 bonding (and 6 anti-bonding) molecular orbitals.

Ligand-Field scheme summarizing σ-bonding in the octahedral complex [Ti(H2O)6]3+.

Ligand field theory 38

In molecular symmetry terms, the six lone pair orbitals from the ligands (one from each ligand) form six symmetryadapted linear combinations (SALCs) of orbitals, also sometimes called ligand group orbitals (LGOs). Theirreducible representations that these span are a1g, t1u and eg. The metal also has six valence orbitals that span theseirreducible representations - the s orbital is labeled a1g, a set of three p-orbitals is labeled t1u, and the dz2 and dx2−y2orbitals are labeled eg. The six σ-bonding molecular orbitals result from the combinations of ligand SALC's withmetal orbitals of the same symmetry.

π-bondingπ bonding in octahedral complexes occurs in two ways: via any ligand p-orbitals that are not being used in σbonding, and via any π or π* molecular orbitals present on the ligand.The p-orbitals of the metal are used for σ bonding (and are the wrong symmetry to overlap with the ligand p or π orπ* orbitals anyway), so the π interactions take place with the appropriate metal d-orbitals, i.e. dxy, dxz and dyz. Theseare the orbitals that are non-bonding when only σ bonding takes place.One important π bonding in coordination complexes is metal-to-ligand π bonding, also called π backbonding. Itoccurs when the LUMOs of the ligand are anti-bonding π* orbitals. These orbitals are close in energy to the dxy, dxzand dyz orbitals, with which they combine to form bonding orbitals (i.e. orbitals of lower energy than theaforementioned set of d-orbitals). The corresponding anti-bonding orbitals are higher in energy than the anti-bondingorbitals from σ bonding so, after the new π bonding orbitals are filled with electrons from the metal d-orbitals, ΔOhas increased and the bond between the ligand and the metal strengthens. The ligands end up with electrons in theirπ* molecular orbital, so the corresponding π bond within the ligand weakens.The other form of coordination π bonding is ligand-to-metal bonding. This situation arises when the π-symmetry por π orbitals on the ligands are filled. They combine with the dxy, dxz and dyz orbitals on the metal and donateelectrons to the resulting π-symmetry bonding orbital between them and the metal. The metal-ligand bond issomewhat strengthened by this interaction, but the complementary anti-bonding molecular orbital fromligand-to-metal bonding is not higher in energy than the anti-bonding molecular orbital from the σ bonding. It isfilled with electrons from the metal d-orbitals, however, becoming the HOMO of the complex. For that reason, ΔOdecreases when ligand-to-metal bonding occurs.The greater stabilisation that results from metal-to-ligand bonding is caused by the donation of negative charge awayfrom the metal ion, towards the ligands. This allows the metal to accept the σ bonds more easily. The combination ofligand-to-metal σ-bonding and metal-to-ligand π-bonding is a synergic effect, as each enhances the other.As each of the six ligands has two orbitals of π-symmetry, there are twelve in total. The symmetry adapted linearcombinations of these fall into four triply degenerate irreducible representations, one of which is of t2g symmetry.The dxy, dxz and dyz orbitals on the metal also have this symmetry, and so the π-bonds formed between a centralmetal and six ligands also have it (as these π-bonds are just formed by the overlap of two sets of orbitals with t2gsymmetry.)

High and low spin and the spectrochemical seriesThe six bonding molecular orbitals that are formed are "filled" with the electrons from the ligands, and electronsfrom the d-orbitals of the metal ion occupy the non-bonding and, in some cases, anti-bonding MO's. The energydifference between the latter two types of MO's is called ΔO (O stands for octahedral) and is determined by thenature of the π-interaction between the ligand orbitals with the d-orbitals on the central atom. As described above,π-donor ligands lead to a small ΔO and are called weak- or low-field ligands, whereas π-acceptor ligands lead to alarge value of ΔO and are called strong- or high-field ligands. Ligands that are neither π-donor nor π-acceptor give avalue of ΔO somewhere in-between.

Ligand field theory 39

The size of ΔO determines the electronic structure of the d4 - d7 ions. In complexes of metals with these d-electronconfigurations, the non-bonding and anti-bonding molecular orbitals can be filled in two ways: one in which as manyelectrons as possible are put in the non-bonding orbitals before filling the anti-bonding orbitals, and one in which asmany unpaired electrons as possible are put in. The former case is called low-spin, while the latter is calledhigh-spin. A small ΔO can be overcome by the energetic gain from not pairing the electrons, leading to high-spin.When ΔO is large, however, the spin-pairing energy becomes negligible by comparison and a low-spin state arises.The spectrochemical series is an empirically-derived list of ligands ordered by the size of the splitting Δ that theyproduce. It can be seen that the low-field ligands are all π-donors (such as I-), the high field ligands are π-acceptors(such as CN- and CO), and ligands such as H2O and NH3, which are neither, are in the middle.I− < Br− < S2− < SCN− < Cl− < NO3

− < N3− < F− < OH− < C2O4

2− < H2O < NCS− < CH3CN < py (pyridine) < NH3< en (ethylenediamine) < bipy (2,2'-bipyridine) < phen (1,10-phenanthroline) < NO2

− < PPh3 < CN− < CO

HistoryLigand field theory was developed during the 1930s and 1940s as an alternative to crystal field theory (CFT). CFTdescribes certain properties of coordination complexes but is based on a model that emphasizes electrostaticinteractions between ligand electrons with the d-electrons on the metal. CFT does not describe bonding. Ligand FieldTheory, in a sense, combined CFT and the then-emerging molecular orbital theory.

References[1][1] Schläfer, H. L.; Gliemann, G. "Basic Principles of Ligand Field Theory" Wiley Interscience: New York; 1969[2] G. L. Miessler and D. A. Tarr “Inorganic Chemistry” 3rd Ed, Pearson/Prentice Hall publisher, ISBN 0-13-035471-6.

Article Sources and Contributors 40

Article Sources and ContributorsLigand  Source: http://en.wikipedia.org/w/index.php?oldid=502666113  Contributors: 194.200.130.xxx, 1exec1, 4lex, Ace Frahm, Achemgeekf, Ahoerstemeier, Akhilchirravuri,Alexnullnullsieben, AlphaEta, Altenmann, Amangill, Amateria1121, AnnaFrance, Apple2, Ausinha, Ave matthew, Axiosaurus, Basicdesign, Beetstra, Benbest, Benjah-bmm27, Bhellis,Biscuittin, Bissinger, Brandon, Brichcja, Bryan Derksen, CMSherwood, Cadmium, Centrx, Ceyockey, Chem8240bv, ChemMater, Christian75, Cmrufo, Conrad.Irwin, Conversion script, Crystalwhacker, Cstmoore, Dcirovic, Dilbert08, Dmb000006, Dormroomchemist, Drphilharmonic, Dualus, Dwmyers, Erajda, Fnielsen, Gadfium, Gaius Cornelius, Garethhegarty, Gentgeen, Giftlite,Gilliam, Gioto, Habj, Hugo-cs, Itub, Jrissman, Jrtayloriv, Junior Brian, Jynto, K. Hiippari, Kenster85hero, Ktsquare, Lfh, Lifeformnoho, Lkinkade, Luís Felipe Braga, Magwich77,Manderson198, Martarius, Mattbr, Mayooranathan, Mets501, Michael Hardy, Mj455972007, Mrnatural, Nihiltres, Nina Gerlach, Nlu, Nono64, OMCV, Organometallics, Patrick Lutz, Pentalis,Petrb, Pgholder, Pixeltoo, Postglock, Pwdent, Quadalpha, Quadro, R'n'B, Red, Rej8, Rifleman 82, Sadads, Serein (renamed because of SUL), SilverShatter, SilveryWolf, Slightsmile, Smokefoot,Sodium, SrKAawa, Srnec, Stewartadcock, T.c.w7468, T.vanschaik, TDogg310, TheAMmollusc, Thegeneralguy, Thiseye, ThomasYun, Tlabshier, Tmangray, V8rik, Valenciano, Vinci cool,Walkerma, Waltpohl, Wangi, Wayne Slam, Whoop whoop pull up, Wickey-nl, Woohookitty, Yaluen, YitzHandel, Zenohockey, Zoom321, ТимофейЛееСуда, 152 ,زرشک anonymous edits

Crystal field theory  Source: http://en.wikipedia.org/w/index.php?oldid=505556984  Contributors: 1727quimico, 17Drew, Alansohn, Anton Khorev, Aponar Kestrel, ArgentiumOutlaw, Beetstra,Benjah-bmm27, Bit Lordy, Blethering Scot, Brichcja, CarbonX, Carsrac, Chemicalengineer03, Claviere, Dirac66, DragonflySixtyseven, Freestyle-69, Gaius Cornelius, Gmanley, Hoffmeier,J991, Jaganath, Jaraalbe, Jh51681, Kazkaskazkasako, Kingpin13, LinDrug, Manuelkuhs, Materialscientist, Michael Hardy, Moorfrogger, MrBland, OMCV, Okedem, Out of Phase User, PSUPHYS514 F06, Passw0rd, Petergans, Pettythug, PleaseStand, Prateek khanna, Puppy8800, Quadell, Quantockgoblin, RaseaC, Red Thrush, Reedy, Rifleman 82, SWpens11, Sfan00 IMG, SilentNemesis2710, Smokefoot, Snehashistamlukwb, T.vanschaik, Tabatharose, V8rik, Walkerma, Welsh, Xiglofre, YanA, Zaiken, 78 anonymous edits

Denticity  Source: http://en.wikipedia.org/w/index.php?oldid=503385743  Contributors: Anypodetos, Axiosaurus, CMSherwood, CommonsDelinker, Crystal whacker, Dr. F.C. Turner, DrPhen,Imareaver, Lamro, MrBell, Rod57, Roux, Smokefoot, Wickey-nl, 6 anonymous edits

Chelation  Source: http://en.wikipedia.org/w/index.php?oldid=501870390  Contributors: A876, AdamRoach, Albmont, Alexei Kouprianov, Alphachimp, Anders.Warga, Andonic, Anupam,Arcadian, Archfool, AxelBoldt, Axewiki, Beetstra, Benbest, Bensaccount, Bill.albing, Bug42, Bunnyhop11, Burschik, C S, Cacycle, CaptinJohn, Cburnett, CecilWard, Centrx, Chowbok, Crystalwhacker, Cuaxdon, Cutefidgety, Dcirovic, Dfranke, Dieter Simon, Dj Capricorn, Dmcarey1, Dmn, Doodle77, Drilnoth, Drsibia, Dwmyers, Elinor McCartney, Erianna, Eubulides, Filelakeshoe,Finemann, FrozenMan, Gaius Cornelius, Gene Nygaard, Gonzonoir, Gpunketesh, Grandpa Larsen, Ground Zero, HappyCamper, Herald83, Hughcharlesparker, JaGa, Jkbrown, JoJan, JohnNevard, Johner, JonRichfield, Jpbrenna, Jtyndall02, Julesd, Kauczuk, Keenan Pepper, Keira Vaughn, Kelson, Kjmoran, Kristenq, Ktsquare, Kwamikagami, Langhorner, Lorenzarius, Maelli,Mboverload, Mellery, MrBell, NaOH, Narayanese, NawlinWiki, Nbarth, Neparis, Nihiltres, Ohnoitsjamie, Ojigiri, Okedem, Organometallics, Outsidelookin, Passw0rd, PaulNovitski, Persian PoetGal, Petergans, Pgan002, Physchim62, Pinethicket, Pjacobi, Pquijal, Rajah, Ransu, Recognizance, Rifleman 82, Rjwilmsi, Sam Hocevar, Scharks, Sdel, Shaddack, Shalom Yechiel, Shunnosuke,SlamDiego, Smokefoot, Srnec, StevenDH, Sticky Parkin, StradivariusTV, TJRC, TMorris13, Talgalili, Tomas e, Tortoise0308, Trusilver, Unyoyega, Varlaam, Vuo, WLU, Westerness,Wickey-nl, Yaris678, Zymatik, زرشک, มือใหม่, 119 anonymous edits

Hapticity  Source: http://en.wikipedia.org/w/index.php?oldid=491347596  Contributors: Beetstra, Beetstra public, Brichcja, CMSherwood, Dr. Sunglasses, Edgar181, Eno-ja, Euchiasmus, GaiusCornelius, Haeleth, HappyApple, HappyCamper, Hooperbloob, Lhynard, Nergaal, PedroDaGr8, R'n'B, RandomP, Rifleman 82, Sasuke Sarutobi, Shalom Yechiel, Shinryuu, Smokefoot,T.vanschaik, Tabletop, Tetracube, V8rik, YanA, 8 anonymous edits

Trans-spanning ligand  Source: http://en.wikipedia.org/w/index.php?oldid=467580083  Contributors: Christian75, Dualus, Firebat08, Itub, Lamro, OMCV, Rifleman 82, Smokefoot, Tassedethe,1 anonymous edits

Linkage isomerism  Source: http://en.wikipedia.org/w/index.php?oldid=479104667  Contributors: CWenger, Charles Matthews, Christian75, Eggilicious, GeeJo, Goldenrowley, JamesBWatson,Physchim62, Puppy8800, Retropunk, Rifleman 82, Shalom Yechiel, Smokefoot, Srnec, Tiddly Tom, V8rik, 10 anonymous edits

Bridging ligand  Source: http://en.wikipedia.org/w/index.php?oldid=495206398  Contributors: Alan Au, Axiosaurus, Beetstra, Benjah-bmm27, CMSherwood, Chem-awb, Dslate2123, Itub,Mets501, Mikespedia, Patrick Lutz, PedroDaGr8, Physchim62, Quaxmonster, Rifleman 82, Smokefoot, Srnec, Tomásdearg92, Useight, 5 anonymous edits

Metal–ligand multiple bond  Source: http://en.wikipedia.org/w/index.php?oldid=504420935  Contributors: Chem8240bv, Dicklyon, EdChem, ErikHaugen, Hydrogen Iodide, Nergaal, OMCV,Rifleman 82, Smokefoot, SunCreator, TenPoundHammer, V8rik, Wickey-nl, 8 anonymous edits

Non-innocent ligand  Source: http://en.wikipedia.org/w/index.php?oldid=503504714  Contributors: Bruintje71, Eaberry, Headbomb, Nick Y., Rifleman 82, Rjwilmsi, Smokefoot, V8rik, 42anonymous edits

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Ligand field theory  Source: http://en.wikipedia.org/w/index.php?oldid=465494912  Contributors: Axiosaurus, Borgx, Brichcja, Calvero JP, Centrx, Dirac66, Jaganath, Kelix, Michael Hardy,Mladjowie, OMCV, Omegakent, Ophiucusthesorceror, Out of Phase User, Passw0rd, Petergans, Physchim62, Puppy8800, Sadads, Smokefoot, T.vanschaik, Themusicking, Tijmz, V8rik, YanA,33 anonymous edits

Image Sources, Licenses and Contributors 41

Image Sources, Licenses and ContributorsFile:HCo(CO)4-3D-balls.png  Source: http://en.wikipedia.org/w/index.php?title=File:HCo(CO)4-3D-balls.png  License: Public Domain  Contributors: Benjah-bmm27File:Metal-EDTA.svg  Source: http://en.wikipedia.org/w/index.php?title=File:Metal-EDTA.svg  License: Public Domain  Contributors: Smokefoot derivative work: Chamberlain2007 (talk)File:CoA6Cl3.png  Source: http://en.wikipedia.org/w/index.php?title=File:CoA6Cl3.png  License: Public Domain  Contributors: Cwbm (commons), SmokefootFile:CFT - Low Spin Splitting Diagram 2.png  Source: http://en.wikipedia.org/w/index.php?title=File:CFT_-_Low_Spin_Splitting_Diagram_2.png  License: GNU Free Documentation License Contributors: YanAImage:CFT - High Spin Splitting Diagram 2.png  Source: http://en.wikipedia.org/w/index.php?title=File:CFT_-_High_Spin_Splitting_Diagram_2.png  License: GNU Free DocumentationLicense  Contributors: YanAFile:Crystal Field Splitting 4.png  Source: http://en.wikipedia.org/w/index.php?title=File:Crystal_Field_Splitting_4.png  License: GNU Free Documentation License  Contributors: YanAFile:Colorwheel.jpg  Source: http://en.wikipedia.org/w/index.php?title=File:Colorwheel.jpg  License: Public Domain  Contributors: Original uploader was Tabatharose at en.wikipediaImage:Octahedral crystal-field splitting.png  Source: http://en.wikipedia.org/w/index.php?title=File:Octahedral_crystal-field_splitting.png  License: GNU Free Documentation License Contributors: en:YanAFile:Pentagonal bipyramidal.png  Source: http://en.wikipedia.org/w/index.php?title=File:Pentagonal_bipyramidal.png  License: GNU Free Documentation License  Contributors: Originaluploader was YanA at en.wikipedia (Original text : YanA)Image:Square antiprismatic.png  Source: http://en.wikipedia.org/w/index.php?title=File:Square_antiprismatic.png  License: GNU Free Documentation License  Contributors: Original uploaderwas YanA at en.wikipedia (Original text : YanA)Image:Square planar.png  Source: http://en.wikipedia.org/w/index.php?title=File:Square_planar.png  License: GNU Free Documentation License  Contributors: Original uploader was YanA aten.wikipedia (Original text : YanA)Image:Square pyramidal.png  Source: http://en.wikipedia.org/w/index.php?title=File:Square_pyramidal.png  License: GNU Free Documentation License  Contributors: Original uploader wasYanA at en.wikipedia (Original text : YanA)Image:Tetrahedral.png  Source: http://en.wikipedia.org/w/index.php?title=File:Tetrahedral.png  License: GNU Free Documentation License  Contributors: Original uploader was YanA aten.wikipedia (Original text : YanA)Image:Trigonal bipyramidal.png  Source: http://en.wikipedia.org/w/index.php?title=File:Trigonal_bipyramidal.png  License: GNU Free Documentation License  Contributors: Originaluploader was YanA at en.wikipedia (Original text : YanA)Image:Octahedral-3D-balls.png  Source: http://en.wikipedia.org/w/index.php?title=File:Octahedral-3D-balls.png  License: Public Domain  Contributors: Benjah-bmm27Image:AX7E0-3D-balls.png  Source: http://en.wikipedia.org/w/index.php?title=File:AX7E0-3D-balls.png  License: Public Domain  Contributors: Benjah-bmm27Image:Square-planar-3D-balls.png  Source: http://en.wikipedia.org/w/index.php?title=File:Square-planar-3D-balls.png  License: Public Domain  Contributors: Benjah-bmm27, Zzyzx11Image:Square-pyramidal-3D-balls.png  Source: http://en.wikipedia.org/w/index.php?title=File:Square-pyramidal-3D-balls.png  License: Public Domain  Contributors: Benjah-bmm27Image:Tetrahedral-3D-balls.png  Source: http://en.wikipedia.org/w/index.php?title=File:Tetrahedral-3D-balls.png  License: Public Domain  Contributors: Benjah-bmm27Image:Trigonal-bipyramidal-3D-balls.png  Source: http://en.wikipedia.org/w/index.php?title=File:Trigonal-bipyramidal-3D-balls.png  License: Public Domain  Contributors: Benjah-bmm27Image:Pentagonal-pyramidal-3D-balls.png  Source: http://en.wikipedia.org/w/index.php?title=File:Pentagonal-pyramidal-3D-balls.png  License: Public Domain  Contributors: Benjah-bmm27File:Hexaaquasodium-3D-balls.png  Source: http://en.wikipedia.org/w/index.php?title=File:Hexaaquasodium-3D-balls.png  License: Public Domain  Contributors: Ben MillsImage:Oxaliplatin.svg  Source: http://en.wikipedia.org/w/index.php?title=File:Oxaliplatin.svg  License: Public Domain  Contributors: Calvero.Image:Linear2-4Chelate.png  Source: http://en.wikipedia.org/w/index.php?title=File:Linear2-4Chelate.png  License: Public Domain  Contributors: SmokefootImage:Metal-EDTA.png  Source: http://en.wikipedia.org/w/index.php?title=File:Metal-EDTA.png  License: Public Domain  Contributors: User:YikrazuulFile:M-en1.png  Source: http://en.wikipedia.org/w/index.php?title=File:M-en1.png  License: Public Domain  Contributors: Amada44, Ba10r, DMacks, Pieter Kuiper, Wickey-nlImage:Cu chelate.png  Source: http://en.wikipedia.org/w/index.php?title=File:Cu_chelate.png  License: Public Domain  Contributors: . Original uploader was Petergans at en.wikipediaImage:Equilibrium.svg  Source: http://en.wikipedia.org/w/index.php?title=File:Equilibrium.svg  License: Public Domain  Contributors: L'AquatiqueImage:StrikeO.png  Source: http://en.wikipedia.org/w/index.php?title=File:StrikeO.png  License: Public Domain  Contributors: . 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