Metal Complexes in Aqueous Solutions

The idea of coordination of ligands to metal ions goes back to the theories of Alfred Werner, and his painstaking accumulation of evidence^1,2 of the coordination numbers and geometries of complexes of metal ions such as Co(III) and Pd(II). The achievement of Werner was the realization that metal ions have typical numbers of ligands (see Appendix for Glossary of terms) coordinated to them, such as the Co(III) and Pd(II) ions with their coordination numbers of six and four respectively, and that these numbers were not necessarily the same as the oxidation state of the metal ion. Werner further realized that ligands were arranged in space in definite geometries, such as the octahedral coordination geometry of Co(III) complexes, or square-planar geometry of Pd(II) complexes. Specific coordination geometries accounted for the occurrence of optical and geometrical isomers in these complexes, as summarized in Figure 1.1.

When a unidentate ligand such as NH₃ or SCN⁻ replaces a water molecule on a metal ion to form a complex in aqueous solution, several factors contribute to the final free energy of complex formation. This section summarizes these factors as they are presently understood.

The Chelate Effect¹ produces increased stability for the complexes of chelating ligands as compared with those of open-chain analogues. This is seen for the formation constants of complexes of Ni(II) with n-dentate polyamines, as compared with the analogous complexes with ammonia.²

Macrocyclic ligands have traditionally been divided into two classes, those with oxygen donors, such as the crown ethers and cryptands, discovered by Pedersen¹ and Lehn,² and the nitrogen donor macrocycles, first investigated by workers such as Curtis³ and Busch.⁴ The sulfur-donor macrocycles^5–7 resemble the nitrogen donor macrocycles more closely in their coordinating properties, and so tend to be grouped with them. The division into mainly oxygen-donor macrocycles on the one hand, and nitrogen donor macrocycles on the other, derives from a tendency on the part of investigators to study one or the other type of ligand, which in turn rests on the very different types of metal ions which are strongly complexed by the two classes of ligand. The oxygen donor macrocycles tend to complex well with metal ions such as the larger alkali and alkaline earth metal ions, and the larger post-transition metal ions such as Pb(II), Tl(I), or Hg(II). The nitrogen donor macrocycles complex well with transition metal ions, as well as the post-transition metal ions. The sulfur-donor macrocycles are weakly complexing with all metal ions, but complex best with the same group as the nitrogen donor macrocycles. This chapter is not an attempt to cover all aspects of macrocyclic chemistry, but rather an attempt to provide insight into the factors that govern the complexation of metal ions by macrocyclic ligands. For other aspects of macrocyclic chemistry, a selection of excellent books and reviews is available.^8–17

It is the purpose of this chapter to use the principles that have been described in previous chapters on metal complexes, chelates, and macrocyclic compounds to develop an understanding of the use of metal complexes or ligands for medical purposes. Previous chapters have emphasized the factors involved in designing metal complexes with maximum stability, and with a large degree of selectivity. Both of these factors are important in biological systems because of the fact that all metal ions are subject to interaction with the natural ligands that are present in body fluids for the purpose of storage, transport and the regulation of the activity of natural metal ions that are needed for various metabolic purposes. Therefore, in order for a ligand to be effective in biological systems, or a metal complex to retain its integrity in competition with natural carriers, the thermodynamic stability of the complex in question must be maximized. When a ligand is designed to remove a certain metal ion because it has reached toxic levels, it must be as selective as possible for that metal ion so as not to disturb the metals that are naturally present. However, high thermodynamic stability and selectivity are not the sole requirements of the ligand or of the metal complex being considered, because there are many other factors in biological systems which must be taken into consideration. Such factors include the method of administration (oral, sub-cutaneous, intravenous injection, etc.), bioavailability, membrane permeability, toxicity, and rapid elimination of the ligand and its metal chelate without spreading the undesired metal to other organs throughout the body. It will be the philosophy of this chapter that the thermodynamic stability of a metal chelate should be maximized, and the principles described in this book will be used for this purpose.

In Chapters 1 to 4, the factors that control selective complexation of metal ions were discussed. Here factors that control strength of metal-ligand interactions as these might relate to metal ions in biology are examined. It is not possible here to give an exhaustive account of selective complexation of metal ions in biology. Rather, factors that control selectivity are examined as they relate to selected examples in biology, and it is hoped that readers will be able to identify these same principles acting in their own area of interest in metal ions in biology.

The equilibrium constant involving the formation of a metal complex from the aquo metal ion and the most basic form of the ligand is a standard measure of the effectiveness of the ligand in coordinating metal ions. The constants involved are called stability constants or formation constants. Most complex formation reactions are measured in aqueous medium under controlled conditions, and the formation constants generally apply to that medium. However, ligands that are not soluble in water but are soluble in organic solvents are frequently employed and their formation constants with metal ions are often determined in mixed solvents such as dioxane-water (up to 70% dioxane by volume); ethanol- and methanol-water systems are also quite common. For completely organic systems such as acetonitrile or tetrahydrofuran, metal complexes can be formed quite readily but their formation constants are generally not known. Approximate values have been used occasionally for such systems but there is no mathematical expression that can relate the equilibrium constants in such systems to the formation constants or stability constants in water or water/organic mixtures.