The Chelate Effect

Effect entropy effect common to all chelate systems, but often additional stabilization results 

A more sophisticated approach would be to consider the reaction  

99 Ethylenediamine and ammonia are almost identical in field strengths (their / factors differ by less than 3%, Chapter I IK and so the LFSE for complexes with the same metal ion will be almost identical. However, ethylenediamine is also a stronger base than ammonia from the inductive effect of the methylene groups. In addition there may be enthalpy differences in the form of ring strain or other sieric effects. See Table 12.6 and the following discussion. 

Thermodynamic contributions to the chelate affect in complexes of nidteipi) and copper ll)a 

Entropy changes associated with solvation and rotational differences are also important. The driving force for this reaction (AG = -50 kJ mol- ) comes predominantly from Ihe 7AS term. 

4 Entropies oTcheUtionahauid be compared with 33v4n(n = number of chelale rings) based on = tilMn 35.3. 

Finally, chelating ligands such as acelylacetone enjoy re.sonance stabilization as a result of forming six-membered rings having some aromatic character. Acelylacetone (2,4-pentanedione) coordinates as an anionic enolale ligand  

Eniropy changes associaled with sotvalion and raiaiional dWerenecs are also imponaoi. The driving force for this reaction idC -50 kJ comes prcdominaniiy froni ihe 7 lemu 

With trivalent metals, acetylacetone thus forms neutral tris complexes such as [Al(acac),], [TiCacacly], [Crlacac),], and [Co(acac),]. As a result of resonance, the two M—O bonds in each of these complexes are equal in length, as are the two C—O and the two ring C—C bonds, giving a symmetric structure (only one ring shown)  

Tlie chelate effect may be defined as the unusual stability of a coordination compound involving a chelating, multidentate ligand as compared with equivalent compounds involving monodentate ligands. To see the magnitude of this effect, consider the following two reactions represented in Equations (6.9) and (6.10)  

Note that the overall stability constant (i of the tris(ethylenediamine) complex is about 10 orders of magnitude greater than that of the equivalent hexaammine complex. 

What could be the cause of such a large difference in thermodynamic stability After all, the number of Ni -N coordinate-covalent bonds is six in both the products of these two reactions, so the enthalpy changes (Ai/) involved when these bonds are formed should be fairly similar. That seems to leave entropy as the major explanation for the effect. Indeed, the rationale for the chelate effect can be understood in two ways, both related to the relative probabilities that the two reactions will occur. First, consider the number of reactants and products in the two cases. As written more explicitly in Equations (6.11) and (6.12), it is apparent that the number of ions and molecules scattered throughout the water structure in the first reaction stays the same (seven in both the reactants and the products). In the second reaction, however, three ethylenediamine molecules replace six water molecules in the coordination sphere, and the number of particles scattered at random throughout the aqueous solution increases from four to seven. The larger number of particles distributed randomly in the solution represents a state of higher probability or higher entropy for the products of the second reaction. Therefore, the second reaction is favored over the first due to this entropy effect. 

Recall that the relative importance of the enthalpy and entropy changes in a reaction is given by the expression for the free-energy change shown in Equation (6.13)  

A large increase in the entropy of a reaction is reflected in a more negative value for AG. The equilibrium constant is in turn related to the change in free energy by the expression given in Equation (6.14)  

Chugaev conduded that five- or six-membered chelate rings are the most favored in coordination compoundsl and this conclusion is supported by quantitative data. Compounds with seven-membered rings are close in thermodynamic stability to analogous coordination compounds without 

The number of liberated water molecules is typically twice the number of incoming ligands when bidentate ligands like ethylenediamine are employed, Eq. 1.10  

As a result of this difference, the entropy component of the standard free energy change is affected, Eq. 1.11  

the standard free energy change associated with Eq. 1.10 is more negative than that associated with Eq. 1.9 and so the tris(ethylenediamine) complex is thermodynamically more stable than the hexa(ammino) complex. The enthalpy component of the free energy change may also be a contributing factor in the chelate effect. 5,16 

Observations on the chelate effect and, in particular, the importance of ring size, enable predictions to be made concerning the formation of stable coordination compounds with polydentate ligands that contain a variety of individual functional groups. Work on the synthesis of organic ligands 

K1K2K3K4 for complexing by four NH3 molecules must be compared to K3K2 for complexing by two en molecules because two of the latter occupy four coordination sites. 

Thermodynamically, there is little difference between enthalpies of formation of bonds to NH3 or en for a given metal. Both involve the formation of coordinate bonds between nitrogen atoms and the metal ions. However, the larger formation constants for the en complexes must reflect a more negative value of AG. Recalling that 

Metal ions dissolved in water are effectively complexed to water molecules. Displacing the set of water ligands, partially or entirely by another set in such aqua metal ions, results in forming what is more conventionally known as complexes. Displacement of water molecules by multidentate ligands results in more stable complexes than similar systems with none or fewer chelates. 

This has been defined by Ken Raymond as the negative logarithm of the concentration of the free or uncomplexed metal ion (pM — —log [M ]) and it is calculated from the formation constant for a total ligand concentration of 10 M and [M ]tot 10 M under standard conditions (i.e., pH 7.4 and 25 °C). 

The blood of the patients with thalassaemia contains an abnormal form of haemoglobin. 

FIGURE 2.3 (a) The metal chelator desferrioxamine (DFO) and (b) its complex with iron. 

FIGURE 2.4 The structures of corrin (left) and porphyrin (right). 

The number of monodentate groups such as hydroxide, halide, water molecule, etc., bonded to a central metal ion, usually equates to its coordination number. However, a remarkable feature occurs if two or more of these ligand donor groups are united, possibly by a short chain of hydrocarbons there will be a markedly enhanced preference for forming such bidentate complexes with the transition metal ion, and indeed in the ultimate, it is possible to go through tri- to tetra- to penta-to hexadentate chelate ligands such as EDTA (ethylenediaminetetra-acetic acid) which is widely described in this book. 

Most chelate rings involving octahedral transition metal ions tend to be five- or six-membered since these best satisfy the 90 ° angle between 

For completeness, it must be noted that the main group metal ions, because they prefer ionic bonding, are normally found in association with monodentate ligands such as the water molecule, chloride, carbonate, etc., in human systems, but even they can be complexed with a well chosen polydentate ligand. 

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Page, M. L., Jencks, W. P. Entropic contributions to rate accelerations in enzymic and intramolecular interactions and the chelate effect. Proc. Natl. Acad. Sci. USA 68 (1971) 1678-1683... 

The coordination of bidentate ligands is generally more efficient than expected on the basis of the binding affinity of monodentate analogues. This is referred to as the chelate effect. For reviews, see (a) Schwarzenbach, G. Helv. Chim. Acta, 1952, 35, 2344 (b) reference 75. 

On the basis of the values of AS° derived in this way it appears that the chelate effect is usually due to more favourable entropy changes associated with ring formation. However, the objection can be made that and /3l-l as just defined have different dimensions and so are not directly comparable. It has been suggested that to surmount this objection concentrations should be expressed in the dimensionless unit mole fraction instead of the more usual mol dm. Since the concentration of pure water at 25°C is approximately 55.5 moldm , the value of concentration expressed in mole fractions = cone in moldm /55.5 Thus, while is thereby increased by the factor (55.5), /3l-l is increased by the factor (55.5) so that the derived values of AG° and AS° will be quite different. The effect of this change in units is shown in Table 19.1 for the Cd complexes of L = methylamine and L-L = ethylenediamine. It appears that the entropy advantage of the chelate, and with it the chelate effect itself, virtually disappears when mole fractions replace moldm . ... 

Probably the most satisfactory model with which to explain the chelate effect is that proposed by G. Schwarzenbach If L and L-L are present in similar concentrations and are competing for two coordination sites on the metal, the probability of either of them coordinating to the first site may be taken as equal. However, once one end of L-L has become attached it is much more likely that the second site will be won by its other end than by L, simply because its other end must be held close to the second site and its effective concentration where it matters is therefore much... 

The formation of a single complex species rather than the stepwise production of such species will clearly simplify complexometric titrations and facilitate the detection of end points. Schwarzenbach2 realised that the acetate ion is able to form acetato complexes of low stability with nearly all polyvalent cations, and that if this property could be reinforced by the chelate effect, then much stronger complexes would be formed by most metal cations. He found that the aminopolycarboxylic acids are excellent complexing agents the most important of these is 1,2-diaminoethanetetra-aceticacid (ethylenediaminetetra-acetic acid). The formula (I) is preferred to (II), since it has been shown from measurements of the dissociation constants that two hydrogen atoms are probably held in the form of zwitterions. The values of pK are respectively pK, = 2.0, pK2 = 2.7,... 

The factors which influence the stability of metal ion complexes have been discussed in Section 2.23, but it is appropriate to emphasise here the significance of the chelate effect and to list the features of the ligand which affect chelate formation ... 

An investigation of the chelate effect the binding of bidentate phosphine and arsine chelates in square-planar transition metal complexes. D. M. A. Minahan, W. E. Hill and C. A. McAuliffe, Coord. Chem. Rev., 1984, 55, 31-54 (153). 

The Chelate Effect and Polydentate Ligands 147 Table 8-1. Stability constants for some nickel(ii) complexes of ammonia and 1,2-diaminoethane. 

Ligand-Field Stabilization Energies 8.2.3 Contributions to the Chelate Effect - The Entropy... 

Two examples are noteworthy of the chelate effect in the formation of acyl complexes. The reaction of NaMn(CO)5 with MeSCH2CH2Cl at —78°C... 

The complex Cu(II)2(0-BISTREN) is much more acidic than the free Cu2+ ion, by a factor of more than three log units. This is primarily due to the presence of two Cu(II) ions, because the formation constant of the Cu2(OH)+ complex is not much less than that for the Cu2(0-BISTREN) complex with hydroxide. This is not a good indication of how well two free Cu2+ ions would bind hydroxide compared to the Cu2(0-BISTREN) complex, however, since one must take into account the dilution effect operative in the chelate effect to make the comparison more realistic (90). Thus, the formation constant for the Cu2OH+ complex above applies for the standard reference state of 1 M Cu2 +. In contrast, in 10 6 M Cu2+, for example, the pH at which Cu2(OH) + would form is raised from pH 5.6 to 11.6, ignoring the fact that Cu(OH)2(s) would precipitate out long before this pH as reached. By comparison, the acidity of the Cu2(0-BISTREN) complex is not affected by dilution and would still form the hydroxide complex at pH 3.9 if present at a 10"6 M concentration. 

The complexes of zinc with pyridylphenylketone and pyridylphenylmethanol are stabilized by the chelate effect. The complexes formed were dependent on anions present and the ratio of the... 

These compounds have been prepared via oxidative addition reactions between the appropriate phosphate or phosphine and either a quinone or via displacement reactions with a suitable diol. Compounds 81 and 83 were prepared by such a displacement reaction between monocyclic pentaoxyphosphorane 317 and 3-fluorocatechol or catechol in toluene, respectively. This reaction takes advantage of the chelation effect of forming a bicyclic system from a monocyclic one <1997IC5730>. Compound 82 and compound 84 were synthesized via oxidation addition between tetrachloroquinone and the respective sulfur-containing cyclic phosphate or phosphine <1997IC5730>. Compound 93 was prepared from the phosphine 318 and the diol 319 in the presence of iV-chlorodiisopropylamine in an ether solution <1998IC93>. 

Because the ethylenediamine forms chelate rings, the increased stability compared to Nff3 complexes is called the chelate effect. For both ligands, the atoms donating the electron pairs are nitrogen atoms. The difference in stability of the complexes is not related to the strength of the bonds between the metal ion and nitrogen atoms. 

Because of the chelate effect, ligands that can displace two or more water molecules from the coordination sphere of the metal generally form stable complexes. One ligand that forms very stable complexes is the anion ethylenediaminetetraacetate (EDTA4-),... 

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