Ligand chemistry

19,727 views
19,329 views

Published on

This document provides an overview about Ligand chemistry.

Published in: Technology, Education
1 Comment
7 Likes
Statistics
Notes
No Downloads
Views
Total views
19,727
On SlideShare
0
From Embeds
0
Number of Embeds
10
Actions
Shares
0
Downloads
378
Comments
1
Likes
7
Embeds 0
No embeds

No notes for slide

Ligand chemistry

  1. 1. Ligand ChemistrySandeep Badarla PDF generated using the open source mwlib toolkit. See http://code.pediapress.com/ for more information. PDF generated at: Sun, 05 Aug 2012 04:56:58 UTC
  2. 2. ContentsArticles Ligand 1 Crystal field theory 9 Denticity 14 Chelation 16 Hapticity 20 Trans-spanning ligand 23 Linkage isomerism 24 Bridging ligand 25 Metal–ligand multiple bond 27 Non-innocent ligand 29 Chiral ligand 34 Ligand dependent pathway 37 Ligand field theory 37References Article Sources and Contributors 40 Image Sources, Licenses and Contributors 41Article Licenses License 42
  3. 3. Ligand 1 Ligand In coordination chemistry, a ligand is an ion or molecule (see also: functional group) that binds to a central metal atom to form a coordination complex. The bonding between metal and ligand generally involves formal donation of one or more of the ligands electron deficient pairs. The nature of metal-ligand bonding can range from covalent to ionic. Furthermore, the metal-ligand bond order can range from one to three. Ligands are viewed as Lewis bases, although rare cases are known involving Lewis acidic "ligands."[1][2] Metals and metalloids are bound to ligands in virtually all circumstances, although gaseous "naked" metal ions can be generated in high vacuum. Ligands in a complex dictate the reactivity of the Cobalt complex [HCo(CO)4] with five ligands central atom, including ligand substitution rates, the reactivity of the ligands themselves, and redox. Ligand selection is a critical consideration in many practical areas, including bioinorganic 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 the number of electrons donated to the metal (denticity or hapticity). The size of a ligand is indicated by its cone angle. History The composition of coordination complexes have been known since the early 1800s, e.g. Prussian blue and copper vitriol. The key breakthrough occurred when Alfred Werner reconciled formulas and isomers. He showed, among other things, that the formulas of many cobalt(III) and chromium(III) compounds can be understood if the metal has six ligands in an octahedral geometry. The first to use the term "ligand" were Alfred Stock and Carl Somiesky, in relation to silicon chemistry. The theory allows one to understand the difference between coordinated and ionic chloride in the cobalt ammine chlorides and to explain many of the previously inexplicable isomers. He resolved the first coordination complex called hexol into optical isomers, overthrowing the theory that chirality was necessarily associated with carbon compounds.[3][4] Strong field and weak field ligands In general, ligands are viewed as electron donors and the metals as electron acceptors(citation. metals are electron donors). Bonding is often described using the formalisms of molecular orbital theory. The HOMO (Highest Occupied 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 also hard/soft acid/base theory). Metal ions preferentially bind certain ligands. In general, soft metal ions prefer weak field ligands, whereas hard metal ions prefer strong field ligands. According to the molecular orbital theory, the HOMO 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 complexes bound to weak-field ligands follow Hunds rule. Binding of the metal with the ligands results in a set of molecular orbitals, where the metal can be identified with a new HOMO and LUMO (the orbitals defining the properties and reactivity of the resulting complex) and a certain ordering of the 5 d-orbitals (which may be filled, or partially filled with electrons). In an octahedral environment, the 5 otherwise degenerate d-orbitals split in sets of 2 and 3 orbitals (for a more in depth explanation, see crystal field theory).
  4. 4. Ligand 2 3 orbitals of low energy: dxy, dxz and dyz 2 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 is determined by the field-strength of the ligand: strong field ligands, by definition, increase Δo more than weak field ligands. Ligands can now be sorted according to the magnitude of Δo (see the table below). This ordering of ligands is almost invariable for all metal ions and is called spectrochemical series. For complexes with a tetrahedral surrounding, the d-orbitals again split into two sets, but this time in reverse order. 2 orbitals of low energy: dz2 and dx2−y2 3 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 the d-orbitals are influenced by 6 ligands. When the coordination number is neither octahedral nor tetrahedral, the splitting becomes correspondingly more complex. For the purposes of ranking ligands, however, the properties of the octahedral complexes and the resulting Δo has been of primary interest. The arrangement of the d-orbitals on the central atom (as determined by the strength of the ligand), has a strong effect on virtually all the properties of the resulting complexes. E.g. the energy differences in the d-orbitals has a strong effect in the optical absorption spectra of metal complexes. It turns out that valence electrons occupying orbitals with significant 3d-orbital character absorb in the 400-800 nm region of the spectrum (UV-visible range). The absorption of light (what we perceive as the color) by these electrons (that is, excitation of electrons from one orbital to another orbital under influence of light) can be correlated to the ground state of the metal complex, which reflects the bonding properties of the ligands. The relative change in (relative) energy of the d-orbitals as a function of the field-strength of the ligands is described in Tanabe-Sugano diagrams. In cases where the ligand has low energy LUMO, such orbitals also participate in the bonding. The metal-ligand bond can be further stabilised by a formal donation of electron density back to the ligand in a process known as back-bonding. In this case a filled, central-atom-based orbital donates density into the LUMO of the (coordinated) ligand. Carbon monoxide is the preeminent example a ligand that engages metals via back-donation. Complementarily, ligands with low-energy filled orbitals of pi-symmetry can serve as pi-donor. 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" for Covalent Bond Classification - was popularized by M.L.H. Green and "is based on the notion that there are three basic types [of ligands]... represented by the symbols L, X, and Z, which correspond respectively to 2-electron, 1-electron and 0-electron neutral ligands."[5][6] L ligands are derived from charge-neutral precursors and are represented by amines, phosphines, CO, N2, and alkenes. X ligands typically are derived from anionic precursors such as chloride but includes ligands where salts of anion do not really exist such as hydride and alkyl. Thus, the complex IrCl(CO)(PPh3)2 is classified as an MXL3 complex, since CO and the two PPh3 Metal-EDTA complex, wherein ligands are classified as Ls. The oxidative addition of H2 to IrCl(CO)(PPh3)2 gives the aminocarboxylate is a hexadentate chelating ligand. an 18e- ML3X3 product, IrClH2(CO)(PPh3)2. EDTA4- is classified as an L2X4 ligand, as it features four anions and two neutral donor sites. Cp is classified as an L2X ligand.[7]
  5. 5. Ligand 3 Polydentate and polyhapto ligand motifs and nomenclature Denticity Denticity (represented by κ) refers to the number of times a ligand bonds to a metal Cobalt(III) complex containing through non-contiguous donor sites. Many ligands are capable of binding metal ions six ammonia ligands, which are through multiple sites, usually because the ligands have lone pairs on more than one monodentate. The chloride is not a ligand. atom. Ligands that bind via more than one atom are often termed chelating. A ligand that binds through two sites is classified as bidentate, and three sites as tridentate. The "bite angle" refers to the angle between the two bonds of a bidentate chelate. Chelating ligands are commonly formed by linking donor groups via organic linkers. A classic bidentate ligand is ethylenediamine, which is derived by the linking of two ammonia groups with an ethylene (-CH2CH2-) linker. A classic example of a polydentate ligand is the hexadentate chelating agent EDTA, which is able to bond through six sites, completely surrounding some metals. The number of times a polydentate ligand bind to a metal centre is symbolized with "κn", where "n" indicates the number sites by which a ligand attaches to a metal. EDTA4−, when it is hexidentate, binds as a κ6-ligand, the amines and the carboxylate oxygen atoms are not contiguous. In practice, the n value of a ligand is not indicated explicitly but rather assumed. The binding affinity of a chelating system depends on the chelating angle or bite angle. Complexes of polydentate ligands are called chelate complexes. They tend to be more stable than complexes derived from monodentate ligands. This enhanced stability, the chelate effect, is usually attributed to effects of entropy, which favors the displacement of many ligands by one polydentate ligand. When the chelating ligand forms a large ring that at least partially surrounds the central atom and bonds to it, leaving the central atom at the centre of a large ring. The more rigid and the higher its denticity, the more inert will be the macrocyclic complex. Heme is a good example: the iron atom is at the centre of a porphyrin macrocycle, being bound to four nitrogen atoms of the tetrapyrrole macrocycle. The very stable dimethylglyoximate complex of nickel is a synthetic macrocycle derived from the anion of dimethylglyoxime. Hapticity Hapticity (represented by η) refers to the number of contiguous atoms that comprise a donor site and attach to a metal center. Butadiene forms both η2 and η4 complexes depending on the number of carbon atoms that are bonded to the metal.[7] Ligand motifs Outer-sphere ligands In coordination chemistry, the ligands that are directly bonded to the metal (that is, share electrons), form part of the first coordination sphere and are sometimes called "inner sphere" ligands. "Outer-sphere" ligands are not directly attached to the metal, but are bonded, generally weakly, to the first coordination shell, affecting the inner sphere in subtle ways. The complex of the metal with the inner sphere ligands is then called a coordination complex, which can be neutral, cationic, or anionic. The complex, along with its counterions (if required), is called a coordination compound.
  6. 6. Ligand 4 Trans-spanning ligands Trans-spanning ligands are bidentate ligands that can span coordination positions on opposite sides of a coordination complex.[8] Ambidentate ligand Unlike polydentate ligands, ambidentate ligands can attach to the central atom in two places but not both. A good example of this is thiocyanate, SCN−, which can attach at either the sulfur atom or the nitrogen atom. Such compounds give rise to linkage isomerism. Polyfunctional ligands, see especially proteins, can bond to a metal center through different ligand atoms to form various isomers. Bridging ligand A bridging ligand links two or more metal center. Virtually all inorganic solids with simple formulas are coordination polymers, consisting of metal centres linked by bridging ligands. This group of materials includes all anhydrous binary metal halides and pseudohalides. Bridging ligands also persist in solution. Polyatomic ligands such as carbonate are ambidentate and thus are found to often bind to two or three metals simultaneously. Atoms that bridge metals are sometimes indicated with the prefix "μ" (mu). Most inorganic solids, are polymers by virtue of the presence of multiple bridging ligands. Metal–ligand multiple bond Metal ligand multiple bonds some ligands can bond to a metal center through the same atom but with a different number of lone pairs. The bond order of the metal ligand bond can be in part distinguished through the metal ligand bond angle (M-X-R). This bond angle is often referred to as being linear or bent with further discussion concerning the degree to which the angle is bent. For example, an imido ligand in the ionic form has three lone pairs. One lone pair is used as a sigma X donor, the other two lone pairs are available as L type pi donors. If both lone pairs are used in pi bonds then the M-N-R geometry is linear. However, if one or both these lone pairs is non-bonding then the M-N-R bond is bent and the extent of the bend speaks to how much pi bonding there may be. η1-Nitric oxide can coordinate to a metal center in linear or bent manner. Specialized ligand types Non-innocent ligand Non-innocent ligands bond with metals in such a manner that the distribution of electron density between the metal center and ligand is unclear. Describing the bonding of noninnocent ligands often involves writing multiple resonance forms that have partial contributions to the overall state. Bulky ligands Bulky ligands are used to control the steric properties of a metal center. They are used for many reasons, both practical and academic. On the practical side, they influence the selectivity of metal catalysts, e.g. in hydroformylation. Of academic interest, bulky ligands stabilize unusual coordination sites, e.g. reactive coligands or low coordination numbers. Often bulky ligands are employed to simulate the steric protection afforded by proteins to metal-containing active sites. Of course excessive steric bulk can prevent the coordination of certain ligands.
  7. 7. Ligand 5 Chiral ligands Chiral ligands are useful for inducing asymmetry within the coordination sphere. Often the ligand is employed as an optically pure group. In some cases, e.g. secondary amines, the asymmetry arises upon coordination. Chiral ligands are essential components of asymmetric homogeneous catalysis. Common ligands See nomenclature. Virtually every molecule and every ion can serve as a ligand for (or "coordinate to") metals. Monodentate ligands include virtually all anions and all simple Lewis bases. Thus, the halides and pseudohalides are important anionic ligands whereas ammonia, carbon monoxide, and water are particularly common charge-neutral ligands. Simple organic species are also very common, be they anionic (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 in forming the coordinate bond. Also, metals can bind to the σ bonds in for example silanes, hydrocarbons, and dihydrogen (see also: agostic interaction). In complexes of non-innocent ligands, the ligand is bonded to metals via conventional bonds, but the ligand is also redox-active. Examples of common ligands (by field strength) In the following table the ligands are sorted by field strength (weak field ligands first): Ligand formula (bonding atom(s) Charge Most common denticity Remark(s) in bold) Iodide (iodo) monoanionic monodentate I− Bromide (bromido) monoanionic monodentate Br− Sulfide (thio or less commonly dianionic monodentate (M=S), or S2− "bridging thiolate") bidentate bridging (M-S-M) Thiocyanate (S-thiocyanato) monoanionic monodentate ambidentate (see also isothiocyanate, S-CN− below) Chloride (chlorido) monoanionic monodentate also found bridging Cl− Nitrate (nitrato) monoanionic monodentate O-NO2− Azide (azido) monoanionic monodentate N-N2− Fluoride (fluoro) monoanionic monodentate F− Hydroxide (hydroxo) monoanionic monodentate often found as a bridging ligand O-H− Oxalate (oxalato) dianionic bidentate [O-C(=O)-C(=O)-O]2− Water (aqua) H-O-H neutral monodentate monodentate Nitrite (nitrito) monoanionic monodentate ambidentate (see also nitro) O-N-O− Isothiocyanate (isothiocyanato) monoanionic monodentate ambidentate (see also thiocyanate, N=C=S− above) Acetonitrile (acetonitrilo) CH3CN neutral monodentate
  8. 8. Ligand 6 Pyridine C5H5N neutral monodentate Ammonia (ammine or less NH3 neutral monodentate commonly "ammino") Ethylenediamine en neutral bidentate 2,2-Bipyridine bipy neutral bidentate easily reduced to its (radical) anion or even to its dianion 1,10-Phenanthroline phen neutral bidentate Nitrite (nitro) monoanionic monodentate ambidentate (see also nitrito) N-O2− Triphenylphosphine PPh3 neutral monodentate Cyanide (cyano) monoanionic monodentate can bridge between metals (both CN− metals bound to C, or one to C and one to N) Carbon monoxide (carbonyl) CO neutral monodentate can bridge between metals (both metals 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 bridges between metals) or when the conformation of the ligand gets distorted (e.g. a linear ligand that is forced through steric 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 Most Remark(s) common denticityAcetylacetonate (Acac) CH3-C(O)-CH2-C(O)-CH3 monoanionic bidentate In general bidentate, bound through both oxygens, but sometimes bound through the central carbon only, see also analogous ketimine analoguesAlkenes R2C=CR2 neutral compounds with a C-C double bondBenzene C6H6 neutral and other arenes1,2-Bis(diphenylphosphino)ethane (dppe) Ph2PC2H4PPh2 neutral bidentate1,1-Bis(diphenylphosphino)methane (dppm) C25H22P2 neutral Can bond to 2 metal atoms at once, forming dimersCorroles tetradentateCrown ethers neutral primarily for alkali and alkaline earth metal cations2,2,2-crypt hexadentate primarily for alkali and alkaline earth metal cations
  9. 9. Ligand 7Cryptates neutralCyclopentadienyl (Cp) monoanionic Although [C5H5]− monoanionic, by the nature of its occupied MOs, it is capable of acting as a tridentate ligand.Diethylenetriamine (dien) C4H13N3 neutral tridentate related to TACN, but not constrained to facial complexation monoanionicDimethylglyoximate (dmgH−)Ethylenediaminetetraacetate (EDTA) (HOOC-CH2)2N-(CH2)2-N(CH2-COOH)2 tetra-anionic hexadentate actual ligand is the tetra-anionEthylenediaminetriacetate trianionic pentadentate actual ligand is the trianionEthyleneglycol-bis(oxyethylenenitrilo)-tetraacetate (HOOC-CH2)2N-(CH2)2-O-(CH2)2-O-(CH2)2-N(CH2-COOH)2 tetra-anionic octodentate(EGTA)glycinate (Glycinato) monoanionic bidentate other α-amino acid anions are comparable (but chiral)Heme dianionic tetradentate macrocyclic ligandNitrosyl cationic bent (1e) and linear NO+ (3e) bonding modeOxo O dianion monodentate sometimes bridgingPyrazine N2C4H4 neutral ditopic sometimes bridgingScorpionate ligand tridentateSulfite monoanionic monodentate ambidentate2,2,5,2-Terpyridine (terpy) neutral tridentate meridional bonding onlyTriazacyclononane (tacn) (C2H4)3(NR)3 neutral tridentate macrocyclic ligand see also the N,N,N"-trimethylated analogueTricyclohexylphosphine (C6H11)3P or (PCy3) neutral monodentateTriethylenetetramine (trien) neutral tetradentateTrimethylphosphine PMe3 neutral monodentateTri(o-tolyl)phosphine P(o-tolyl)3 neutral monodentateTris(2-aminoethyl)amine (tren) (NH2CH2CH2)3N neutral tetradentateTris(2-diphenylphosphineethyl)amine (np3) neutral tetradentateTerpyridine C15H11N3 neutral tridentateTropylium cationic C7H7+Carbon dioxide CO2 see transition metal carbon dioxide complex
  10. 10. Ligand 8 Ligand exchange Ligand exchange (also ligand substitution) is a type of chemical reaction in which one ligand in a chemical compound is replaced by another ligand. One type of pathway for substitution is the Ligand dependent pathway. In organometallic chemistry this can take place by associative substitution or by dissociative substitution. Another form of ligand exchange is seen in the nucleophilic abstraction reaction. Pronunciation Pronounced ˈlɪgənd with the first syllable sounding like the word "lithium" or laɪgənd with the first syllable sounding 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 Werners 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] Hartwig, J. F. Organotransition Metal Chemistry, from Bonding to Catalysis; University Science Books: New York, 2010. ISBN 1-891389-53-X [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.
  11. 11. Crystal field theory 9 Crystal field theory Crystal field theory (CFT) is a model that describes the breaking of degeneracies of electronic orbital states, usually d or f orbitals, due to a static electric field produced by a surrounding charge distribution (anion neighbors). This theory has been used to describe various spectroscopies of transition metal coordination complexes, in particular optical spectra (colours). CFT successfully accounts for some magnetic properties, colours, hydration enthalpies, and spinel structures of transition metal complexes, but it does not attempt to describe bonding. CFT was developed by physicists Hans Bethe and John Hasbrouck van Vleck[1] in the 1930s. CFT was subsequently combined with molecular orbital theory to form the more realistic and complex ligand field theory (LFT), which delivers insight into the process of chemical bonding in transition metal complexes. Overview of crystal field theory analysis According to CFT, the interaction between a transition metal and ligands arises from the attraction between the positively charged metal cation and negative charge on the non-bonding electrons of the ligand. The theory is developed by considering energy changes of the five degenerate d-orbitals upon being surrounded by an array of point charges consisting of the ligands. As a ligand approaches the metal ion, the electrons from the ligand will be closer to some of the d-orbitals and farther away from others causing a loss of degeneracy. The electrons in the d-orbitals and those in the ligand repel each other due to repulsion between like charges. Thus the d-electrons closer to 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 metals 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. In octahedral symmetry the d-orbitals split into two sets with an energy difference, Δoct (the crystal-field splitting parameter) where the dxy, dxz and dyz orbitals will be lower in energy than the dz2 and dx2-y2, which will have higher energy, 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 t2g, and the two higher-energy orbitals as eg. (These labels are based on the theory of molecular symmetry). Typical orbital energy diagrams are given below in the section 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 Δtet where 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 oriented directly towards the d-orbitals, the energy splitting will be lower than in the octahedral case. Square planar and other complex 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 and geometry of the complex. Some ligands always produce a small value of Δ, while others always give a large splitting. The reasons behind this can be explained by ligand field theory. The spectrochemical series is an empirically-derived list of ligands ordered by the size of the splitting Δ that they produce (small Δ to large Δ; see also this table): I− < Br− < S2− < SCN− < Cl− < NO3− < N3− < F− < OH− < C2O42− < H2O < NCS− < CH3CN < py < NH3 < en < 2,2-bipyridine < phen < NO2− < PPh3 < CN− < CO
  12. 12. Crystal field theory 10 It is useful to note that the ligands producing the most splitting are those that can engage in metal to ligand back-bonding. The oxidation state of the metal also contributes to the size of Δ between the high and low energy levels. As the oxidation state increases for a given metal, the magnitude of Δ increases. A V3+ complex will have a larger Δ than a V2+ complex for a given set of ligands, as the difference in charge density allows the ligands to be closer to a V3+ ion than to a V2+ ion. The smaller distance between the ligand and the metal ion results in a larger Δ, because the ligand and metal electrons are closer together and therefore repel more. High-spin and low-spin Ligands which cause a large splitting Δ of the d-orbitals are referred to as strong-field ligands, such as CN− and CO from the spectrochemical series. In complexes with these ligands, it is unfavourable to put electrons into the high energy orbitals. Therefore, the lower energy orbitals are completely filled before population of the upper sets starts according to the Aufbau principle. Complexes such as this are called "low spin". For Low Spin [Fe(NO2)6]3− crystal field diagram example, NO2− is a strong-field ligand and produces a large Δ. The octahedral ion [Fe(NO2)6]3−, which has 5 d-electrons, would have the octahedral splitting diagram shown at right with all five electrons in the t2g level. Conversely, ligands (like I− and Br−) which cause a small splitting Δ of the d-orbitals are referred to as weak-field ligands. In this case, it is easier to put electrons into the higher energy set of orbitals than it is to put two into the same low-energy orbital, because two electrons in the same orbital repel each other. So, one electron is put into each of the five d-orbitals High Spin [FeBr6]3− crystal field diagram before any pairing occurs in accord with Hunds rule and "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 have an 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 orbital must 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 required to 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 roughly equal to 4/9Δoct (for the same metal and same ligands). Therefore, the energy required to pair two electrons is typically higher than the energy required for placing electrons in the higher energy orbitals. Thus, tetrahedral complexes 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 by magnetic fields, while a compound that lacks unpaired electrons in its splitting diagram will be diamagnetic and will be weakly repelled by a magnetic field.
  13. 13. Crystal field theory 11 Crystal field stabilization energy The crystal field stabilization energy (CFSE) is the stability that results from placing a transition metal ion in the crystal 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 as the barycenter in which all five d-orbitals are degenerate. For example, in an octahedral case, the t2g set becomes lower in energy than the orbitals in the barycenter. As a result of this, if there are any electrons occupying these orbitals, 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 in these reduces the amount of CFSE. If the splitting of the d-orbitals in an octahedral field is Δoct, the three t2g orbitals are stabilized relative to the barycenter by 2/5 Δoct, and the eg orbitals are destabilized by 3/5 Δoct. As examples, consider the two d5 Octahedral crystal field stabilization energy configurations shown further up the page. The low-spin (top) example has five electrons in the t2g orbitals, so the total CFSE is 5 x 2/5 Δoct = 2Δoct. In the high-spin (lower) example, the CFSE is (3 x 2/5 Δoct) - (2 x 3/5 Δoct) = 0 - in this case, the stabilization generated by the electrons in the lower orbitals is canceled 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 many d8 complexes are square-planar is the very large amount of crystal field stabilization that this geometry produces with this number of electrons. Explaining the colors of transition metal complexes The bright colors exhibited by many coordination compounds can be explained by Crystal Field Theory. If the d-orbitals of such a complex have been split into two sets as described above, when the molecule absorbs a photon of visible light one or more electrons may momentarily jump from the lower energy d-orbitals to the higher energy ones to transiently create an excited state atom. The difference in energy between the atom in the ground state and in the excited 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 - the compounds appears the appropriate complementary color. As explained above, because different ligands generate crystal fields of different strengths, different colors can be seen. For a given metal ion, weaker field ligands create a complex with a smaller Δ, which will absorb light of longer λ and thus lower frequency ν. Conversely, stronger field ligands create a larger Δ, absorb light of shorter λ, and thus higher ν. It is, though, rarely the case that the energy of the photon absorbed corresponds exactly to the size of the gap Δ; there are other things (such as electron-electron repulsion and Jahn-Teller effects) that also affect the energy difference between the ground and excited states.
  14. 14. Crystal field theory 12 Which colors are exhibited? This color wheel demonstrates which color a compound will appear if it only has one absorption in the visible spectrum. For example, if the compound absorbs red light, it will appear green. λ absorbed versus color observed 400 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) Color wheel 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
  15. 15. 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.
  16. 16. Denticity 14 Denticity Denticity refers to the number of atoms in a single ligand that bind to a central atom in a coordination complex.[1][2] In many cases, only one atom in the ligand binds to the metal, so the denticity equals one, and the ligand is said to be monodentate (sometimes called unidentate). Ligands with more than one bonded atom are called polydentate or multidentate. The word denticity is derived from dentis, the Latin word for tooth. The ligand is thought of as biting the metal at one or more linkage points. Denticity is distinguished from hapticity, in which electrons of a bond or conjugated series of bonds are linked to the central metal without the metal-ligand bond being localized to a single ligand atom. Classes of denticity Atom with monodentate ligands Polydentate ligands are chelating agents[3] and classified by their denticity. Some atoms cannot form the maximum possible number of bonds a ligand could make. In that case one or more binding sites of the ligand are unused. Such sites can be used to form a bond with another chemical species. • Bidentate (also called didentate) ligands bind with two atoms, an example being ethylenediamine. • Tridentate ligands bind with three atoms, an example being terpyridine. Tridentate ligands usually bind via two kinds of connectivity, called "mer" and "fac." Cyclic tridentate ligands such as TACN and 9-ane-S3 bind in a facial manner. • Tetradentate ligands bind with four atoms, an example being triethylenetetramine (abbreviated trien). Tetradentate ligands bind via three connectivities depending on their topology and the geometry of the metal center. For octahedral metals, the linear tetradentate trien can bind via Structure of the pharmaceutical Oxaliplatin, which features two different bidentate ligands. three geometries. Tripodal tetradentate ligands, e.g. tris(2-aminoethyl)amine, are more constrained, and on octahedra leave two cis sites. Many naturally occurring macrocyclic ligands are tetradentative, an example being the 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) are octadentate. They are particularly useful for binding lanthanide ions, which typically have coordination numbers greater than 6.
  17. 17. Denticity 15 Relationship between "linear" bi-, tri- and tetradentate ligands (red) bound to an octahedral metal center. The structures marked with * are chiral owing to the backbone of the tetradentate ligand. Stability constants In general, the stability of a metal complex correlates with the denticity of the ligands. The stability of a complex is represented quantitatively in the form of Stability constants. Hexadentate ligands tend to bind metal ions more strongly 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] 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
  18. 18. Chelation 16 Chelation Chelation is the formation or presence of two or more separate coordinate bonds between a polydentate (multiple bonded) ligand and a single central atom.[1] Usually these ligands are organic compounds, and are called chelants, chelators, chelating agents, or sequestering agents. The ligand forms a chelate complex with the substrate. Chelate complexes are contrasted with coordination complexes composed of monodentate 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 that they cannot normally react with other elements or ions to produce precipitates 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] Metal-EDTA chelate The chelate effect The chelate effect describes the enhanced affinity of chelating ligands for a metal ion compared to the affinity of a collection of similar nonchelating (monodentate) ligands for the same metal. Consider the two equilibria, in aqueous solution, between the copper(II) ion, Cu2+ and ethylenediamine (en) on the one hand and methylamine, MeNH2 on the other. Cu2+ + en [Cu(en)]2+ (1) Cu2+ + 2 MeNH2 [Cu(MeNH2)2]2+ (2) Ethylenediamine ligand, binding to a central In (1) the bidentate ligand ethylene diamine forms a chelate complex metal ion with two bonds with the copper ion. Chelation results in the formation of a five–membered ring. In (2) the bidentate ligand is replaced by two monodentate methylamine ligands of approximately the same donor power, meaning that the enthalpy of formation of Cu—N bonds is approximately the same in the two reactions. Under conditions of equal copper concentrations and when the concentration of methylamine is twice the concentration of ethylenediamine, the concentration of the complex (1) will be greater than the concentration of the complex (2). Cu2+ complexes with methylamine (left) and The effect increases with the number of chelate rings so the ethylenediamine (right) concentration of the EDTA complex, which has six chelate rings, is much much higher than a corresponding complex with two monodentate nitrogen donor ligands and four monodentate carboxylate ligands. Thus, the phenomenon of the chelate effect is a firmly established empirical fact.
  19. 19. Chelation 17 The thermodynamic approach to explaining the chelate effect considers the equilibrium constant for the reaction: the larger 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 the subscripts to the stability constants, β, indicate the stoichiometry of the complex. When the analytical concentration of methylamine is twice that of ethylenediamine and the concentration of copper is the same in both reactions, the concentration [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 reaction and ΔS is the standard entropy change. It has already been posited that the enthalpy term should be approximately the same for the two reactions. Therefore the difference between the two stability constants is due to the entropy term. In equation (1) there are two particles on the left and one on the right, whereas in equation (2) there are three particles on the left and one on the right. This means that less entropy of disorder is lost when the chelate complex is formed than when the complex with monodentate ligands is formed. This is one of the factors contributing to the entropy difference. Other factors include solvation changes and ring formation. Some experimental data to illustrate the effect are shown in the following table.[3] Equilibrium log β ΔG ΔH /kJ mol−1 −TΔS /kJ mol−1 6.55 -37.4 -57.3 19.9 Cd2+ + 4 MeNH2 Cd(MeNH2)42+ 10.62 -60.67 -56.48 -4.19 Cd2+ + 2 en Cd(en)22+ These data show that the standard enthalpy changes are indeed approximately equal for the two reactions and that the main reason for the greater stability of the chelate complex is due to the entropy term, which is much less unfavourable, indeed, it is favourable in this instance. In general it is difficult to account precisely for thermodynamic values in terms of changes in solution at the molecular level, but it is clear that the chelate effect is predominantly an effect of entropy. Other explanations, including that of Schwarzenbach,[4] are discussed in Greenwood and Earnshaw (loc.cit). In nature Virtually all biochemicals exhibit the ability to dissolve certain metal cations. Thus, proteins, polysaccharides, and polynucleic acids are excellent polydentate ligands for many metal ions. Organic compounds such as the amino acids glutamic acid and histidine, organic diacids such as malate, and polypeptides such as phytochelatin are also typical chelators. In addition to these adventitious chelators, several biomolecules are specifically produced to bind certain metals (see next section).[5][6][7][8] In biochemistry and microbiology Virtually all metalloenzymes feature metals that are chelated, usually to peptides or cofactors and prosthetic groups.[8] Such chelating agents include the porphyrin rings in hemoglobin and chlorophyll. Many microbial species produce water-soluble pigments that serve as chelating agents, termed siderophores. For example, species of Pseudomonas are known to secrete pyocyanin and pyoverdin that bind iron. Enterobactin, produced by E. coli, is the strongest chelating agent known.
  20. 20. Chelation 18 In geology In earth science, chemical weathering is attributed to organic chelating agents, e.g. peptides and sugars, that extract metal ions from minerals and rocks.[9] Most metal complexes in the environment and in nature are bound in some form of chelate ring, e.g. with a humic acid or a protein. Thus, metal chelates are relevant to the mobilization of metals in the soil, the uptake and the accumulation of metals into plants and microorganisms. Selective chelation of heavy metals is relevant to bioremediation, e.g. removal of 137Cs from radioactive waste.[10] Applications Chelators 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 industrial applications. In 2010, the Asia-Pacific region was the largest outlet, generating about 45% of worldwide demand for chelating agents. The region was followed by Western Europe and North America. The global chelating agents market is expected to reach more than 5 million tonnes in 2018.[11] Nutritional supplements In the 1960s, 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 with insoluble salts within the stomach, rendering the metal unavailable for absorption. Amino acids, being effective metal binders, were chosen as the prospective ligands, and research was conducted on the metal-amino acid combinations. 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 is ethylenediaminetetraacetic acid (EDTA). These synthetics applied the same concept of chelation and did create chelated compounds; however, these synthetics were too stable and not nutritionally viable. If the mineral was taken from the EDTA ligand, the ligand could not be used by the body and would be expelled. During the expulsion process 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 as the 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 approximately 150 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 to human 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] Fertilizers Many mineral deficiencies can occur in plants, such as iron chlorosis, which can reduce the nutritional benefits of crops and eventually result in plant death. Mineral chelates have been used to alleviate the mineral deficiencies of affected crops through liquid foliar applications. These fertilizers are also used to prevent deficiencies from occurring and improving the overall health of the plants. [14] Heavy metal detoxification Chelation 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
  21. 21. 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 of disodium EDTA instead of calcium EDTA has resulted in fatalities due to hypocalcemia.[16] Other medical applications Antibiotic 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 access to 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 applications Homogeneous catalysts are often chelated complexes. A typical example is the ruthenium(II) chloride chelated with BINAP (a bidentate phosphine) used in e.g. Noyori asymmetric hydrogenation and asymmetric isomerization. The latter 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 specifically in 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 units and 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 of selenium 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. OConnell, 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] 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
  22. 22. Hapticity 20 Hapticity The term hapticity is used to describe how a group of contiguous atoms of a ligand are coordinated 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 atoms of the ligand that are bound to the metal. In general the η-notation is only used when there is more than one atom coordinated (otherwise the κ-notation is used, see also hapticity vs. denticity). History The need for additional nomenclature for organometallic compounds became apparent in Ferrocene contains two η5-cyclopentadienyl the mid-1950s when Dunitz, Orgel, and Rich described the structure of the "sandwich ligands complex" ferrocene by X-ray crystallography[1] where an iron atom is "sandwiched" between two parallel cyclopentadienyl rings. Cotton later proposed the term hapticity derived from the adjectival prefix hapto (from the Greek haptein, to fasten, denoting contact or combination) placed before the name of the olefin,[2] where the Greek letter η (eta) is used to denote the number of contiguous atoms of a ligand that bind to a metal center. The term is usually employed to describe ligands containing extended π-systems or 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(PPri3)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 Vaskas complex. Examples The 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 Zeises salt or with fullerene, which is bonded 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 the molecule 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 in cyclopentadienyl or benzene rings can share their electrons. • Apparent violations of the 18-electron rule sometimes are explicable in compounds with unusual hapticities:
  23. 23. 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 Electrons Electrons contributed contributed (neutral counting) (ionic counting) 1 2 η1 Allyl 3 4 η3-Allyl cyclopropenyl 3 4 η3-Allenyl 2 2 η2-Butadiene 4 4 η4-Butadiene 1 2 η1-cyclopentadienyl 3 4 η3-cyclopentadienyl 5 6 η5-cyclopentadienyl pentadienyl cyclohexadienyl 2 2 η2-Benzene 4 4 η4-Benzene 6 6 η6-Benzene 7 6 η7-Cycloheptatrienyl η8-Cyclooctatetraenyl 8 10 Changes in hapticity The hapticity of a ligand can change in the course of a reaction.[6] E.g. in a redox reaction:
  24. 24. 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-electron donating ligand L. Removal of one molecule of CO and again donation of two more electrons by the cyclopentadienyl ligand restores the η5-cyclopentadienyl. The so-called indenyl effect also describes changes in hapticity in a substitution reaction. Hapticity vs. denticity Hapticity must be distinguished from denticity. Polydentate ligands coordinate via multiple coordination sites within the ligand. In this case the coordinating atoms are identified using the κ-notation, as for example seen in coordination of 1,2-bis(diphenylphosphino)ethane (Ph2PCH2CH2PPh2), to NiCl2 as dichloro[ethane-1,2-diylbis(diphenylphosphane)-κ2P]nickel(II). If the coordinating atoms are contiguous (connected to each other), the η-notation is used, as e.g. in titanocene dichloride: 5 [7] dichlorobis(η -2,4-cyclopentadien-1-yl)titanium. Hapticity and fluxionality Molecules with polyhapto ligands are often "fluxional", also known as stereochemically non-rigid. Two classes of fluxionality 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 atom in the η1-C5H5 ligand. This reaction is degenerate and, in the jargon of organic chemistry, it is an example of a sigmatropic rearrangement. • Case 2, typically: complexes containing cyclic polyhapto ligands with maximized hapticity. Such ligands tend to rotate. A famous example is ferrocene, Fe(η5-C5H5)2, wherein the Cp rings rotate with a low energy barrier about the 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 not necessarily degenerate, but the rotational barriers have low energies of activation.
  25. 25. 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 2004 ed.). IUPAC. 2004. p. 16. . Trans-spanning ligand Trans-spanning ligands are bidentate ligands that can span opposite sites of a complex with square-planar geometry. A wide variety of ligands that chelate in the cis fashion already exist, but very few can link opposite vertices on a coordination polyhedron. Early attempts to generate trans-spanning bidentate ligands relied on polymethylene chains to link the donor functionalities, but such ligands often lead to coordination polymers. History A diphosphane linked with pentamethylene was claimed to span across a square planar complex. This early attempt was followed by ligands with more rigid backbones. "TRANSPHOS" was the first trans-spanning diphosphane ligand that usually coordinates to palladium(II) and platinum(II) in a trans manner. TRANSPHOS features benzo[c]phenanthrene substituted by diphenylphosphinomethyl (Ph2PCH2) groups at the 1 and 11 positions.[1][2] The polycyclic framework suffers sterically clashing hydrogen centers.
  26. 26. Trans-spanning ligand 24 Xantphos, SPANphos, TRANSDIP and related ligands Xantphos is a trans-spanning ligand, without the steric problems associated with TRANSPHOS. SPANphos is comparable to XANTPHOS but more reliably trans-spanning. TRANSDIP, based on a α-cyclodextrin, is the first ligand 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 of Square 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 a Ligand 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 Chelator Built 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 isomerism Linkage isomerism is the existence of co-ordination compounds that have the same composition differing with the connectivity 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+. The isomerization 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 observed in the NMR spectrum due to the effect of S vs. O bonding on the methyl groups of DMSO. The proper notation for linkage isomerism is the kappa notation where the atom directly bonding to the metal is proceeded by the lowercase Greek letter kappa; κ. For example, NO2- is represented as nitrito-κ-N and nitrito-κ-O, replacing the old system of trivial names such as nitro and nitroso.[2] History The first reported example of linkage isomerism had the formula [Co(NH3)5(NO2)]Cl2. The cationic cobalt complex exists in two separable linkage isomers. In the yellow-coloured isomer, the nitro ligand is bound through nitrogen. In the 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 was the structural difference explained.[3] It was later shown that the red isomer converted to the yellow isomer upon UV-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 during isomerization.
  27. 27. 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 ligand A bridging ligand is a ligand that connects two or more atoms, usually metal ions.[1] The ligand may be atomic or polyatomic. Virtually all complex organic compounds can serve as bridging ligands, so the term is usually restricted to 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 Greek character 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. Illustrative bridging ligands Virtually all ligands are known to bridge, with the exception of amines and ammonia.[3] Particularly common inorganic bridging ligands include • OH−, • O2−, An example of a μ2 bridging ligand • S2−, • SH−, • NH2− • NH2− (imido) • N3− (nitrido) • CO • Halides • Hydride • Cyanide
  28. 28. 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 of the 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 ligands Polyfunctional ligands can attach to metals in many ways and thus can bridge metals in diverse ways, including sharing of one atom or using several atoms. Examples of such polyatomic ligands are the oxoanions CO32− and the related Carboxylate, PO43−, and the polyoxometallates. Several organophosphorus ligands have been developed that bridge 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 bridging ligand. Notice in the name of the compound μ2 has been [4] simplified to μ. (η6-C6H6)2Ru2Cl2(μ-Cl)2 In this particular complex, two chloride ligands are terminal and two are μ2 bridging. The η in the beginning of the formula denotes the hapticity of the benzene ligands. B2H6 This classic borane compound, diborane features two μ2 bridging hydrides.
  29. 29. 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 bond In Chemistry, a metal–ligand multiple bond describes the interaction of certain ligands with a metal with a bond order greater than one.[1] Coordination complexes featuring multiply bonded ligands are of both scholarly and practical interest. Transition metal carbene complexes catalyze the olefin metathesis reaction. Metal oxo intermediates are pervasive in oxidation catalysis. oxygen evolving complex. As a cautionary note, the classification of a metal ligand bond as being "multiple" bond order is ambiguous and even arbitrary because bond order is a formalism. Furthermore, the usage of multiple bonding is not uniform. Symmetry arguments suggest that most ligands engage metals via multiple bonds. The term metal ligand multiple bond" is often reserved 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 are halides. Most common classes of complexes showing metal ligand multiple bonds
  30. 30. Metalligand multiple bond 28 Pi-donor ligands In coordination chemistry, a pi-donor ligand is a kind of ligand endowed with filled non-bonding orbitals that overlap with metal-based orbitals. Their interaction is complementary to the behavior of pi-acceptor ligands. The existence of terminal oxo ligands for the early transition metals is one consequence of this kind of bonding. Classic pi-donor ligands are oxide (O2-), nitride (N3-), imide (RN2-), alkoxide (RO-), amide (R2N-), and fluoride. For late transition metals, strong pi-donors form anti-bonding interactions with the filled d-levels, with consequences for spin state, redox potentials, and ligand exchange rates. Pi-donor ligands are low in the spectrochemical series.[1] Multiple bond stabilization Metals bound to so-called triply bonded carbyne, imide, nitride (nitrido), and oxide (oxo) ligands are generally assigned to high oxidation states with low d electron counts. The high oxidation state stabilizes the highly reduced ligands. The low d electron count allow for many bonds between ligands and the metal center. A d0 metal center can accommodate up to 9 bonds without violating the 18 electron rule, whereas a d6 species can only accommodate 6 bonds. Reactivity explained through ligand hybridization A ligand described in ionic terms can bond to a metal through however many lone pairs it has available. For example many alkoxides use one of their three lone pairs to make a single bond to a metal center. In this situation the oxygen is sp3 hybridized according to valence bond theory. Increasing the bond order to two by involving another lone pair changes the hybridization at the oxygen to an sp2 center with an expected expansion in the M-O-R bond angle and contraction in the M-O bond length. If all three lone pairs are included for a bond order of three than the M-O bond distance contracts further and since the oxygen is a sp center the M-O-R bond angle is 180˚ or linear. Similarly with the imidos are commonly referred to as either bent (sp2) or linear (sp). Even the oxo can be sp2 or sp hybridized. The triply bonded oxo, similar to carbon monoxide, is partially positive at the oxygen atom and unreactive towards bronsted acids at the oxygen atom. When such a complex is reduced, the triple bond can be converted to a double bond at which point the oxygen no longer bears a partial positive charge and is reactive towards acid. Conventions Bonding representations Imido 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 similar phenomenon. It is common to draw only two lines of bonding for all imidos, including the most common linear imidos with a six electron bonding interaction to the metal center. Similarly amido complexes are usually drawn with a single line even though most amido bonds involve four electrons. Alkoxides are generally drawn with a single bond although both two and four electron bonds are common. Oxo can be drawn with two lines regardless of whether four electrons or six are involved in the bond, although it is not uncommon to see six electron oxo bonds represented with three lines. Representing oxidation states There are two motifs to indicate a metal oxidation state based around the actual charge separation of the metal center. Oxidation states up to +3 are believed to be an accurate representation of the charge separation experienced by the metal center. For oxidation states of +4 and larger, the oxidation state becomes more of a formalism with much of the positive charge distributed between the ligands. This distinction can be expressed by using a Roman numeral for the lower oxidation states in the upper right of the metal atomic symbol and an Arabic number with a

×