Advances in Inorganic Chemistry (eBook)

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1995 | 1. Auflage
408 Seiten
Elsevier Science (Verlag)
978-0-08-057891-0 (ISBN)

Infos

Advances in Inorganic Chemistry presents timely and informative summaries of the current progress in a variety of subject areas within inorganic chemistry ranging from bio-inorganic to solid state studies. Thisacclaimed serial features reviews written by experts in the area and is an indispensable reference to advanced researchers. Each volume of Advances in Inorganic Chemistry contains an index, and each chapter is fully referenced.

Inhalt

Front Cover 1
Advances in Inorganic Chemistry, Volume 42 4
Copyright Page 5
Contents 6
Chapter 1. Substitution Reactions of Solvated Metal Ions 10
I. Introduction 11
II. Solvent Exchange and Ligand Substitution on Main Group Metal Ions 23
III. Solvent Exchange and Ligand Substitution on Transition Metal Ions 34
IV. Solvent Exchange and Ligand Substitution on the Trivalent Lanthanides and Diosouranium(VI) 65
V. Solvent Exchange and Ligand Substitution on Some Molybdenum and Tungsten Cluster Complexes 75
VI. Concluding Comments 81
VII. Appendix: Ligand Abbreviations, Formulae, and Structures 83
References 87
Chapter 2. Lewis Acid–Base Behavior in Aqueous Solution: Some Implications for Metal Ions in Biology 98
I. Introduction 98
II. Lewis Acid–Base Interactions in the Gas Phase 99
III. Comparison of Behavior of Lewis Acids and Bases in Aqueous Solution and in the Gas Phase 104
IV. Significance of HSAB Ideas for Zinc-Containing Metalloenzymes 112
V. Chelate Ring Size and Metal Ion Selectivity 118
VI. The Neutral Oxygen Donor 128
VII. The Negative Oxygen Donor 140
VIII. The Nitrogen Donor 142
IX. Sulfur Donors 146
X. Systems Containing More Than One Metal Ion—The Reverse Chelate 149
References 152
Chapter 3. The Synthesis and Structure of Organosilanols 156
I. General Introduction 156
II. General Synthetic Methods 160
III. The Acidity and Basicity of the Silanol Group 185
IV. Structural Studies of Silanols 196
V. Summary and Conclusions 257
References 261
Chapter 4. Studies of the Soluble Methane Monooxygenase Protein System: Structure, Component Interactions, and Hydroxylation Mechanism 272
I. Introduction 272
II. Structural Studies of sMMO Hydroxylase 274
III. Component Interactions 278
IV. Mechanism of Hydroxylation 284
V. Conclusions 294
References 295
Chapter 5. Alkyl, Hydride, and Hydroxide Derivatives of the s- and p-Block Elements Supported by Poly(pyrazolyl)borato Ligation: Models for Carbonic Anhydrase, Receptors for Anions, and the Study of Controlled Crystallographic Disorder 300
I. Introduction 301
II. Syntheses, Structures, and Steric Properties of Poly(pyrazoly1)hydroborato Ligands 303
III. Terminal Alkyl Derivatives of the s- and p-Block Metals Supported by Poly(pyrazoly1)borato Ligation 317
IV. Terminal Hydride Derivatives of the s- and p-Block Metals Supported by Poly(pyrazoly1)borato Ligation 349
V. Terminal Hydroxide Derivatives of the s- and p-Block Metals Supported by Poly(pyrazo1yl)borato Ligation 357
VI. Anion Coordination by Protonated Tris(pyrazoly1)hydroborato Derivatives 379
VII. Controlled Crystallographic Disorder' in [TpRR']MX Complexes: Bond Length Artifacts as Determined by Single Crystal X-Ray Diffraction 381
VIII. Summary 390
References 390
Index 404
Contents of Previous Volumes 414

Textauszug

Substitution Reactions of Solvated Metal Ions

Stephen F. Lincoln; André E. Merbach    Department of Chemistry, University of Adelaide, Australia; and Institut de Chimie Minérale et Analytique, Université de Lausanne, Switzerland

I Introduction

A GENERAL ASPECTS

The most common ligand substitution on a solvated metal ion is the exchange of a water molecule in the first coordination sphere of a metal ion with a water molecule from the second coordination sphere. Water exchange on metal ions may also be viewed as being particularly fundamental, as it is through the intervention of another type of ligand in this process that the stoichiometry of the first coordination sphere is changed and a vast range of metal complexes is formed. Accordingly, it is appropriate to review the water exchange process and the 18 orders of magnitude variation in the water exchange rate constant, H2O, and lability exhibited by metal ions towards this process. This is most effectively commenced through an inspection of Fig. 1, which in its evolving forms has been one of the most intellectually focusing compilations of information about ligand substitution since it was published by Eigen in 1963 (1). At one end of the lability scale it is seen that the mean lifetime of a water molecule in the first coordination sphere of [Rh(H2O)6]3 +, H2O=1/kH2O, is 14.4 years, whereas at the other extreme, that of a water molecule bound to Cs+ is ~2 × 10− 10 s, during which time light travels ~6 cm. Clearly the microscopic interpretation of the macroscopic kinetic and related observations of water exchange and its substitution in the first coordination sphere by other ligands, the ligand substitution mechanism, must account for this great variation in lability reflected in the range of H2O characterizing the metal ions in Fig. 1.

Fig. 1 Mean lifetimes of a single water molecule in the first coordination sphere of a given metal ion, H2O, and the corresponding water exchange rate constants, H2O.The tall bars indicate directly determined values, and the short bars indicate values deduced from ligand substitution studies. References to the plotted values appear in the text.

The metal ions may be conveniently considered in three categories, the first of which is the main group metal ions which, for a given ionic charge, exhibit an increase in H2O with an increase in ionic radius, rM(2). Thus, the lability of the alkali metal ions increases by an order of magnitude as rM increases from Li+ to Cs+, and in general terms these ions are particularly labile coincident with their low ionic charge-to-rM ratio. In contrast, the lability of the alkaline earth metal ions spans approximately seven orders of magnitude, coincident with their higher charge and larger variation in rM. At this point it is appropriate to note that the more labile of the metal ions in these two sets are particularly difficult to characterize in terms of their numbers of coordinated water molecules. Thus, the coordination number of Li+ is variously quoted as 4 and 6, and that of Cs+ as 9, while intermediate values are quoted for Na+, K+, and Rb+. For Be2 + and Mg2 +, coordination numbers of 4 and 6, respectively, are well established, but a range from 6 to 10 is quoted for Ca2 + and similar or greater values are anticipated for Sr2 + and Ba2 +(3–5). As rM increases with coordination number, so the metal ion–water dipole interaction weakens and lability increases. The labilities exhibited by the six-coordinate sets—Zn2 +, Cd2 +, and Hg2 +(1) and Al3 +, Ga3 +, and In3 +(6–8)—vary over two and six orders of magnitude, respectively, and it is not surprising that the smallest of these ions, Al3 +, is the least labile in the light of the preceding discussion.

The second category is the transition metal ions, all of which in Fig. 1 are six-coordinate with the exception of Pt2 + and Pd2 +, which are square-planar four-coordinate (6–9). Their labilities are strongly influenced by the electronic occupancy of their d orbitals. This is illustrated by the divalent first-row transition metal ions, which should exhibit similar labilities to Zn2 + on the basis of their rM; instead, however, their labilities encompass seven orders of magnitude. On a similar basis, the trivalent first-row transition metal ions might be expected to be of similar lability to Ga3 +, but instead they exhibit a lability variation of 11 orders of magnitude, with Cr3 + being at the lower end of this lability scale. Second-row Ru3 + and Rh3 + are slightly more labile and three orders of magnitude less labile than Cr3 +, respectively.

The third category is the heavy eight-coordinate trivalent lanthanides, whose lability decreases with the progressive filling of the 4f orbitals and the resulting lanthanide contraction, and which are very labile as a consequence of their large rM(7, 10, 11).

B THE FORMATION OF METAL COMPLEXES

It is generally considered that some orientation of water in the second coordination sphere of the archetypal aqua ion, [M(H2O)n]m +, occurs (4, 12), but exchange of this water with bulk water appears to occur at close to diffusion-controlled rates. Hence, it is only in the case of the most labile metal ions that the rate of entry of a substituting ligand, Lx−, into the first coordination sphere is within an order of magnitude or so of its rate of entry into the second coordination sphere. Usually, the entry of Lx− into the first coordination sphere is preceded by the formation of an outer-sphere complex (13), in which Lx− resides in the second coordination sphere of [M(H2O)n]m+. An impressive armory of kinetic techniques has facilitated the study of ligand substitution processes in solution, ranging from those occurring at close to diffusioncontrolled rates to those taking place over extended times (14–16). However, it is seldom the case that more than one stage of the movement of a monodentate ligand from the bulk water environment into the first coordination sphere is observed, except in ultrasonic studies where as many as three steps have been detected. It is such studies which have provided the basis for most current discussion of the mechanisms of metal complexation. Thus, the formation of metal complexes has been conveniently formalized in a mechanism proposed by Eigen and Tamm (17), and illustrated by Eq. (1), where Mm+ is the metal ion, Lx− is the substituting ligand, and only those water molecules interposed between Mm+ and Lx−are shown. Initially, in a diffusioncontrolled step, Mm+ and Lx− form a species (1) in which they are separated by two layers of water molecules. This is followed by a fast step in which Lx− enters the second solvation sphere or second coordination sphere of Mm+ to form an outer-sphere complex (2), and the inner-sphere complex (3) is formed in the final and slowest step characterized by k23, which is frequently denoted ki because it indicates the step in which Lx− interchanges between the second and first coordination spheres. By analogy, k32 is often denoted k− i. The formation of (1) is seldom detected (18), and as a consequence, the simpler mechanism shown in Eq. (2) is often discussed instead and is sometimes referred to as the Eigen–Wilkins mechanism (19).

12Mm++Lx−⇌k10k01M⋅OH2⋅OH2⋅Lm−x+⇌k21k12M⋅OH2⋅Lm−x+MLm−x+k23⇌k323

  (1)

m++Lx−⇌k21k12M⋅OH2⋅Lm−x+⇌k32k23MLm−x+

  (2)

12=K0=4πNR3/3000exp−zMzLe02/εRkT

  (3)

m++Lx−⇌kbkfMLm−x+

  (4)

obs=k1K0[Lx−]1+K0[Lx−]+k−i

  (5)

The successive equilibria are characterized by K12 and K23, respectively, and when K12 (often denoted K0) cannot be directly determined, it may be estimated from the Fuoss equation (3), where R is the distance of closest approach of Mz+ and Lx− (considered as spherical species) in M · OH2 · L(m − x)+, ε is the solvent dielectric constant, and zM and zL are the...

Produkt-Details

• Erscheint lt. Verlag: 10.10.1995
• Mitarbeit: Herausgeber (Serie): AG Sykes
• Sprache: englisch
• Themenwelt: Sachbuch/Ratgeber
• Themenwelt: Naturwissenschaften ► Chemie ► Anorganische Chemie
• Themenwelt: Technik
• ISBN-10: 0-08-057891-8 / 0080578918
• ISBN-13: 978-0-08-057891-0 / 9780080578910