Open Collections

UBC Theses and Dissertations

UBC Theses Logo

UBC Theses and Dissertations

Aqueous chemistry of vanadium complexes of therapeutic interest Glover, Nicholas 1993

Your browser doesn't seem to have a PDF viewer, please download the PDF to view this item.

Item Metadata


831-ubc_1994-0035.pdf [ 2.42MB ]
JSON: 831-1.0059504.json
JSON-LD: 831-1.0059504-ld.json
RDF/XML (Pretty): 831-1.0059504-rdf.xml
RDF/JSON: 831-1.0059504-rdf.json
Turtle: 831-1.0059504-turtle.txt
N-Triples: 831-1.0059504-rdf-ntriples.txt
Original Record: 831-1.0059504-source.json
Full Text

Full Text

AQUEOUS CHEMISTRY OF VANADIUM COMPLEXES OF THERAPEUTIC INTEREST by NICHOLAS GLOVER B.Sc. (Honours), University of East Anglia, Norwich, Norfolk, U.K., 1990 A THESIS SUBMITTED IN PARTIAL FULFILLMENT OF THE REQUIREMENTS FOR THE DEGREE OF MASTER OF SCIENCE in THE FACULTY OF GRADUATE STUDIES (Department of Chemistry) We accept this thesis as conforming to the required standard  The University of British Columbia November 1993 © Nicholas Glover, 1993  In presenting this thesis  in  partial fulfilment of the  requirements for an advanced degree at the University of British Columbia, I agree that the Library shall make it freely available for reference and study. I further agree that permission for extensive copying of this thesis for scholarly purposes may be granted by the head of my department or by his or her representatives. It is understood that copying or publication of this thesis for financial gain shall not be allowed without my written permission.  (Signature)  Department of The University of British Columbia Vancouver, Canada Date  DE-6 (2/88)  Abstract , 2 The aqueous chemistry of bis(maltolato)oxovanadium(IV), VO(ma) (ma) 2 [V0 , and bis(maltolato)dioxovanadate(V), ] 2 bis(kojato)oxovanadium(1V), VO(ka) were studied as part of a research interest regarding their potential as insulin mimetic therapeutic agents. UV-Visible spectroscopy provided information regarding changes in the immediate 2 and vanadium environment with variable pH, and determined that oxidation of VO(ma) 2 occurs upon exposure to air. Variable pH cyclic voltammetric electrochemical VO(ka) 2 displayed reversible investigation of the compounds of interest revealed that only VO(ma) electrochemical behaviour.  Room temperature and frozen solution ESR spectral  2 in - to VO(ma) (ma) [V0 ] parameters for the complexes were determined. Reduction of 2 V NMR spectroscopy was used to acidic solution was determined by ESR spectroscopy. ’5 probe the type, and stoichiometry, of complexes present in aqueous solution as a function ’V 2 to vanadium(V) complexes was proven by 5 2 and VO(ka) of pH. Oxidation of VO(ma) NMR spectra.  Potentiometric titrations were used to determine ligand protonation  constants, and metal-ligand stability constants.  11  TABLE OF CONTENTS page Abstract  ii  Table of Contents  iii  List of Tables  v  List of Figures  vi  List of Abbreviations  viii  Acknowledgements  x  Dedication  xi  Chapter 1  Introduction  1  Chapter 2  Experimental  10  1. Solution Preparation and Chemicals  10  2. Spectroscopic Studies  11  2.2.1. 51 V NMR Spectroscopy 2.2.2. ESR Spectroscopy  11  2.2.3. UV-Visible Spectroscopy  13  12  2.3. Electrochemical Studies  13  2.4. Potentiometric Equilibrium Measurements  14  2.4.1. Instrumentation  14  2.4.2. Calibration  14  2.4.3. Potentiometric Tiirations  15  2.4.4. Computations  16  2.4.5. Data Collection  17  111  Chapter 3  ResuJts and Dscuion 3.1. 3.2. 3.3. 3.4. 3.5.  Chapter 4  Absorption Spectroscopy Electrochemistry ESR Spectroscopy V NMR Speceroscopy 51 Stability Constant Determinations  Suggestions for Future Study  .  19 19 28 35 46 61 73 75  References  iv  LIST OF TABLES Table 3.1.  2 2 and VO(ka) Visible and IR spectra] parameters of VO(ma)  25  Table 3.2.  2 Variable pH electrochemistry of VO(rna)  32  Table 3.3.  , 2 , VO(ka) 2 Isotropic ESR spectral parameters of VO(ma) V0 and 5 [ma) 4 NH ( 2 ] O) VO(H 2  Table 3.4.  41  , 2 , VO(ka) 2 Anisotropic ESR spectral parameters of VO(ma) ) 5 0 VO(H V0 and 2 [ma) 4 NH. ( 2 ]  42  Table 3.5.  5V NMR spectroscopic properties  46  Table 3.6.  Variable pH 5 W NMR chemical shifts of vanadate  49  Table 3.7.  51  Table 3.8.  Log $3 values for the vanadyl-maltol system  Table 3.9.  Log  V NMR chemical shifts of vanadate-maltol complexes  13 values for the vanadyl-kojic acid system  V  56 66 67  LIST OF FIGURES Figure 1.1.  O) VO(H 2 Representation of 5  3  Figure 1.2.  (3-Hydroxy-4-pyrone)-vanadium compounds investigated in this study  7  Figure 3.1.  2 Variable pH visible spectra of VO(ma)  21  Figure 3.2.  Graph oft versus pH for VO(ma) at X. = 620 nm  22  Figure 3.3.  2 with time following Variation in the visible spectra of VO(ma) exposure to air  23  Figure 3.4.  V0 [ma) 4 NH. ( ] Variable pH UV spectra of 2  27  Figure 3.5.  Reduction potentials for vanadium species as a function of pH.. 28  Figure 3.6.  2 Variable pH electrochemistry of VO(ma)  31  Figure 3.7.  2 Variable pH electrochemistry of VO(ka)  33  Figure 3.8.  V0 [ma) 4 NH ( ] Variable pH electrochemistry of 2  34  Figure 3.9.  2 Variable pH isotropic ESR spectra of VO(ma)  37  Figure 3.10.  , 2 , VO(ka) 2 Isotropic ESR spectra of VO(ma) O) VO(H 2 V0 and 5 [ma) 4 NH ( ] 2  39  Figure 3.11.  (ma) K[V0 ] Representation of the structure of 2  40  Figure 3.12.  Experimental and simulated anisotropic ESR spectra 43  of VO(ma)2 Figure 3.13.  Possible ESR spectral lines due to vanadyl-rat transferrin. 2 treated STZ-diabetic rat spectrum non-treated (VO(ma) -  STZ-diabetic rat spectrum)  45  Figure 3.14.  V NMR spectra of vanadate Variable pH 51  48  Figure 3.15.  Variable maltol concentration 51V NMR spectra  50  Figure 3.16.  Variable vanadate concentration 51V NMR spectra  52  Figure 3.17.  V0 [ma) 4 NH ( ] Variable pH ‘v NMR spectra of 2  55  Figure 3.18.  2 ’V NMR spectra of VO(ma) Variable pH 5  57  vi  Figure 3.17.  (ma) NH.jV0 ] Variable pH 51 V NMR spectra of 2  55  Figure 3.18.  Variable pH 51V NMR specua of VO(ma)2  57  Figure 3.19.  2 Variable pH V NMR spectra of VO(ka)  59  Figure 3.20.  Vanadyl-maltol potentiomethc titration curves  65  Figure 3.21.  Vanadyl-kojic acid potentiometic titration curves  67  Figure 3.22.  V0 potentiometric titration curve [ma) 4 NH. ( 2 ]  68  Figure 3.23.  Speciation diagrams for solutions containing a 1:2 ratio of 72  pyrone (mM) :VO 2  Scheme I  V NMR chemical shifts and possible routes of Equilibria, 51  formation for vanadium(V) and vanadium(IV) complexes with the maltolate anion  vfl  60  LIST OF ABBREVIATIONS Meaning  Abbreviation a  (moles base added moles acid present)/moles ligand  A  angstrom  A  cm-’) 4 electron-nuclear spin coupling constant (lO  Ar  argon  -  overall stability constant degrees Celsius 1 cm-  wave number  6  chemical shift  £  cm1 (Mmolar absorptivity )  E  electrochemical potential (volts) e.m.f. constant of cell  ESR  electron spin resonance  g  Landé effective electron g-factor  G  Gauss  H  hydrogen ion  Hka (ka)  kojic acid, 5-hydroxy-2-hydroxymethyl-4-pyrone  Hma (ma)  maltol, 3-hydroxy-2-methyl-4-pyrone  Hz  Hertz  ic/a  anodic/cathodic current  I  ionic strength  JR  infrared  K  degrees Kelvin absorbance wavelength (nm)  yin  L  ligand  mL  millilitre  mM  millimolar  mV  millivolt  M  metal ion  M  ) 3 (after number) molarity (mol din-  nm  nanometre  NMR  nuclear magnetic resonance  uv=o  vibrational stretching mode  pH  -Iog[H]  pKa  ligand protonation constant (log 3 o,ij)  pK  ionic product constant of water  ppm  parts per million  a  standard deviation  s  seconds  S.H.E.  standard hydrogen electrode  STZ  streptozotocin  UV-Vis  ultraviolet-visible vanadium-51  V  volts  V02+  vanadyl ion  2 VO(ka)  bis(kojato)oxovanadium(W)  2 VO(ma)  bis(maltolato)oxovanadium(IV)  (ma) 2 [V0 ]  bis(maltolato)dioxovanadate(V) anion  X-band  Ca.  9.5 GHz  ix  ACKNOWLEDGEMENTS Acknowledgements are due to those who have helped make this Thesis possible. Firstly, I would like to thank Dr. Chris Orvig, for guidance, understanding and support throughout my time at U.B.C.. and in the preparation and proofreading of this thesis. My sincere appreciation is extended to Dr. Lucio Gelmini for experimental assistance, tedious proofreading and invaluable collaborative efforts. I would also like to express my gratitude to Dr. Geoff Herring for many hours of assistance with the ESR measurements. The help given by support staff in the Chemistry Department, in particular the NMR spectroscopy staff, was greatly appreciated. Many past and present members of the Orvig group, especially Dr. Martha Kline, have offered invaluable assistance, and I thank them. Special thanks to Violet Yeun for the rat plasma samples. Financial support from U.B.C. in the form of a Teaching Assistantship is gratefully acknowledged. A simple thanks is not enough to express my appreciation of the support, encouragement and love that Mum, Dad, Steve, Elaine and the rest of my family have given me, and their pride in my achievements. I would dearly like to thank Jim and Colette, who have welcomed me so warmly into their family, and have helped Barbara and I in so many ways. Finally, the biggest thanks of all are for my wife Barbara, for her continuing wholehearted love and support, and much more.  x  This work is dedicated to my Dad, Ronald Glover. The qualities ofperseverance, appreciation and pride in doing your best are gifts I will always treasure.  xi  Chapter 1 Introduction Vanadium was first recognized by Nils Gabriel Sefstrom in 1831, and named for its many multicoloured solutions after Vanadis, the Norse goddess of beauty. Vanadium is found widely distributed in the world, and has a natural abundance greater than that of copper, zinc or lead. It is accumulated in above average concentrations in coal, oil and in certain rock formations, such as sandstone and shale. In addition it is found diversely in the plant and animal kingdoms. In mammals it is an ultratrace element, found particularly in blood plasma. 1,2 The discovery by Henze in 1911 of high concentrations of vanadium in the blood cells of tunicates, a form of marine ascidian, did much to promote early interest in the 3 Since then, the essentiality of vanadium in certain biological chemistry of vanadium. organisms has been confirmed, and increasing evidence suggests vanadium is an essential 4 Amavadin, a low molecular weight protein found in the trace element in mammals. 5 A toadstool genus Amanira, was the first vanadium-containing protein to be isolated. bromoperoxidase from the marine algae Ascophyllum nodosum was the first enzyme found 7 The discovery by Cantley et al in 1977 that ’ 6 to have a vanadium-containing active site. vanadate is a potent inhibitor of Na, K-ATPases led to numerous studies on the inhibitory 8 Recent findings have established that vanadium and stimulatory effects of vanadium. compounds can exhibit insulin-like properties in biological systems, and this has led to an 9 There have been several intensive research effort in this area during the past 10 years. 06 extensive reviews on the chemistry of vanadium in biological systems.’ The coordination chemistry of V’ species, generally referred to as vanadates, is 2 moieties; the latter is often manifested in a complex dominated by the VO and cis-VO of distorted octahedral geometry. The molecular geometries of vanadium(V) complexes 1  can be more  varied than this, however, as is evidenced in recent studies showing some  -, in both size 3 4 - and P0 3 4 17 The similarity of V0 unusual stereochemistries for vanadates. and geometry, has led many investigators to speculate that this may be one of the factors 11 The inhibitory effects of vanadium contributing to the biological relevance of the metal.  on many phosphohydrolases (vide infra) have been attributed to the vanadium(V) oxidation V NMR state; Gresser and Tracey have published extensively in this area, utilizing 51 8 These and other findings have prompted numerous spectroscopy as a structural probe.’ investigations of the biological chemistry of vanadates, for example vanadium(V) 23 and the role of vanadium(V) in halide promoted catalysis of NADH to NADPH, 24 Under physiological conditions, it is believed that a oxidation by hydrogen peroxide. considerable proportion of vanadate is reduced to vanadium(IV); glutathione and catechol ’ 4 are among the cellular constituents that have been shown to do this. The aqueous chemistry of both vanadium(V) and vanadium(IV) species is characterized by a series of complex equilibria and hydrolyses, still the subject of much 2627 The formation of oligomers is highly favoured in these systems, with study and debate. many pH-dependent species formed. The speciation of vanadates is discussed further in V NMR is the Chapter 3.4. Owing to the favourable nuclear properties of vanadium, 51 spectroscopic method of choice for studying vanadate systems. oiety V0 m By far the most common vanadium(IV) compounds contain the 2 (oxovanadium(IV) ion or, more commonly, vanadyl ion), which is the most stable oxocation of the first row transition metals. The vanadyl ion is analogous to the divalent transition and alkaline earth metals, and reacts to form cationic, anionic and neutral square 7 with the oxo-oxygen occupying an axial pyramidal complexes with a variety of ligands,1’ position (Figure l.l.). The relatively short vanadyl to oxo-oxygen bond length, —1.6 A, is I Although strictly referring to a bound substrate, ligand as used here refers to both bound and unbound substrates. § A less common vanadyl complex coordination geometry that is occasionally observed is trigonal 11 bipyramidal, as adopted by bis(2-methyl-8-quinolinato)oxovanadium(1V).  2  indicative of considerable multiple bond character, in most studies considered to be a double bond. These complexes are reactive, substitution reactions occur fairly rapidly in the apical position trans to the vanadyl oxygen (i.e. for a solvent molecule), and are often followed by rearrangement in the plane of the ligands. Vanadyl to trans -ligand bond lengths are typically some 0.2  -  0.3  A longer than those of the corresponding equatorial  ligands.  A  2.1  A 2.4A  Figure 1.1.  ) 9 O VO(H 2 . 2 Representation of the penta(aquo)oxovanadium(W) ion, 5  In biology, vanadium(IV) is found in the blood of ascidians (and ubiquitously throughout nature) and is thought to form strong complexes with proteins, for example the 0 In addition, vanadium(IV) also forms stable 1 16 transferrin. 3 iron-carrier proteins such as ’ 12 It complexes with small in vivo chelating ligands such as citrate, ATP and amino acids. appears unlikely therefore that free VO exists in biological systems. Dissolved oxygen in plasma may cause oxidation to vanadium(V), and it is probable that a partition exists ). Oxovanadium(IV) has a 3032 physiologically between vanadium(V) and vanadium(IV pararnagnetic, 3d’, electronic ground state, and favourable nuclear properties that make it a powerful ESR spin probe (Chapter 3.3.). 3  O [VO(H ] 2 , ) but above pH 2 In highly acidic solution oxovanadium(IV) exists as 5 27 However, the degree of oxidation can be air oxidation to form vanadates occurs. effectively retarded by using anionic, chelating ligands, with hard donor atoms such as 0 or N. The speciation of vanadyl as a function of pH is discussed further in Chapters 3.3. and 3.4. Only the species found in the acidic regime can be assigned with any certainty. Inherent difficulties have made the elucidation of all equilibria and species existing in neutral and basic solution thus far impossible. 33 ’ 26 Pharmacological effects of vanadium previously noted include the lowering of cholesterol and glucose levels, a diuretic effect, the contraction of blood vessels, and the enhancement of the oxygen affinity of hemoglobin and myoglobin. In addition, vanadium has a direct cardiac effect, similar to digitalis, possibly caused by the inhibition of the Na, K+ATPase.  Several vanadium compounds have been studied with respect to their  cytostatic activity, with moderate success; vanadyl species seem to be most active. These 16 Vanadium is used extensively in various findings have been summarized elsewhere. types of industry (i.e. in alloys, as a catalyst, etc), and exposure to high vanadium concentration levels is not uncommon. Thus, the potential for vanadium toxicity is a viable concern. The current consensus is that there is little risk associated with routine exposure 1 to the metal, provided caution is observed. Diabetes mellitus encompasses a heterogeneous group of diseases with various etiologies. These diseases are all characterized by variable and chronic hyperglycemia and other disturbances in carbohydrate and lipid metabolism, as well as a variety of vascular and neurological complications. The prevalence of the disease in the North American general population is approximately seven million. Diabetes mellitus may be directly or  34 indirectly responsible for some 400,000 deaths per year. The treatment for diabetes mellitus is dominated by the regular use of insulin, a hormone secreted in the pancreas to decrease blood glucose levels, or stimulate glucose  4  oxidation. In diabetes sufferers the production of insulin is compromised, leading to a characteristically elevated blood glucose level. Insulin treatment therefore involves the introduction of the hormone to attain normal concentration levels of insulin or of glucose. However, the digestive juices of the intestinal tract render insulin inactive. This means that oral administration is not possible, and therefore the hormone has to be injected directly into the patient. It is this inconvenient method of administration, and its incumbent complications, that have promoted the search for an orally active therapeutic agent to be  35 used in conjunction with, or to replace, insulin treatment. The action of vanadium compounds on glucose metabolism was first noted by Shechter et al, who reported that vanadates displayed insulin-like properties in vitro in rat 37 However, the 36 results that were subsequently confirmed by other groups. adipocytes; levels of vanadate needed to stimulate insulin-like response in diabetic rats has been 38 Vanadyl (as the sulphate salt) has associated with severe toxic side effects, even death. been shown to be just as effective as vanadate at lowering blood glucose levels in vivo, as 39 These results have prompted the synthesis, and well as being considerably less toxic. subsequent testing, of vanadyl based compounds as potential insulin mimetic agents. One such compound developed in this laboratory is bis(maltolato)oxovanadium(IV) (Figure 1.2(a)), abbreviated as BMOV in the pharmacological literature and as VO(ma)2 in this study. ° The rationale behind this formulation is the low toxicity of the ligand (the 4 common food additive maltol) and the belief that the complex formed would be less toxic. The use of an organic framework around the active vanadyl centre was exploited so as to provide greater lipophilicity and gastrointestinal absorption, thereby potentially decreasing , its 2 the dosage needed to produce a positive effect. Since the development of VO(ma) insulin mimetic properties have been shown to be highly encouraging. As well, in comparison to other vanadium treatments, it displays considerable improvement in cardiac 413 dysfunction in diabetic rats and in the prevention of diabetes induced pathoIogy.  5  There have been studies that question the assumption that vanadate is inherently toxic. It is also noted that vanadyl can be oxidized to vanadate both in vitro and in This has led to the formulation of a related vanadate-based maltol compound , in order to compare the chemistry and V0 bis(maltolato)dioxovanadate(V), 2 [ma) 4 NH ( ] physiological effects of two similar vanadyl and vanadate complexes (Figure 1.2(c)). This compound has been isolated, fully characterized, and a crystal structure determination has recently  been  45 performed.  In addition, a further vanadyl compound,  , has been synthesized using kojic acid as a maltol 2 bis(kojato)oxovanadium(IV), VO(ka) analogue (Figure 1.2(b)). Kojic acid is a biologically relevant ligand precursor, obtained 46 These three compounds from the action of the Aspergillus mold family on carbohydrates. were actively tested for insulin mimetic properties at the onset of this study, and are the foci 47 of this work. This investigation follows a previous M.Sc. thesis from this laboratory, concerned with the synthesis and characterization of further potential insulin mimetic vanadyl and vanadate compounds, and is a contribution to a number of ongoing studies.  6  3 CH 0  II  —  2  Ih  Bis(malto)ato)oxovanadium(IV)  I  0 v,I  2 [(b) Bis(kojato)oxovanadium(IV)  I  [3  1(c) Bis(rnaltolato)dioxovanadate(V)  Figure 1.2.  I  (3-Hydroxy-4-pyrone)-vanadium compounds investigated in this study. 2 and (c) 2 V0 ma)2]• [ 4 NH ( (a) VO(ma)2, (b) VO(ka)  7  Recently, there have been a significant number of studies performed concerning the biological chemistry of vanadium; the number of papers published in this field has risen almost exponentially over the last 25 years. Despite this burgeoning study, there have been few examples of concerted spectroscopic and potentiometric investigations in the literature. The majority of papers published within this field thus far deal with the preparation of various vanadium-ligand complexes, whereas relatively little work has been performed on the solution chemistry. It is well known that oxovanadium(IV) and vanadium(V) form stable complexes with hard 0-donor ligands, such as carboxylates, hydroxycarboxylates, 164854 However, there are relatively few hydroxamates, salicylates, catecholates, etc. examples of comprehensive investigations of the solution chemistry of such complexes. Relevant previous studies will be introduced as needed to corroborate and/or illustrate pertinent points in this work. Apparently, comprehensive synthetic, spectroscopic and in vivo studies are necessary in order to better understand the coordination and physiological chemistry of these compounds. The work presented here represents preliminary data concerning their aqueous chemistry. In this study UV-Visible spectroscopy has been used to monitor changes in the immediate vanadium environment as a function of pH. Air oxidation of the vanadyl compounds has been postulated from time dependent spectra (Chapter 3.1.). Electrochemistry provides information regarding the redox behaviour of the compounds. The variation in cyclic voltammetric waves with pH is presented, and possible redox reactions are discussed (Chapter 3.2.). ESR is a particularly powerful technique for the study of pararnagnetic vanadyl compounds. Changes in the ESR signals with pH have been monitored, and values for isotropic and anisotropic constants determined. These were compared to results previously determined for similar compounds. The appearance of ESR  signals in spectra obtained from solutions of vanadium(V) complexes proves that reduction V NMR spectra were used to monitor the pH to vanadium(IV) occurs (Chapter 3.3.). 51  8  , and to provide stoichiomethc information regarding complex V0 [ma) 4 NH ( ] titration of 2 formation and speciation in solution (Chapter 3.4.). Potentiomethc titrations were used to determine ligand protonation constants, and metal-ligand stability constants. Speciation diagrams were constructed for the metal:ligand systems (Chapter 3.5.). Suggestions for future studies are proposed (Chapter 4). It is hoped that this information provides a basis from which to study further these potentially therapeutic agents, and to design better analogues.  9  Chapter 2 Experimental 2.1. Solution Preparation and Chemicals Water was deionized (Barnstead D8902 and 1)8904 cartridges), then distilled (Corning MP-1 Megapure Still). Solutions of NaOH (—0.15 M) were prepared from dilutions of a 50% aqueous solution of sodium hydroxide (reagent grade, less than 0.1% O with freshly boiled, distilled and deionized wateT, and standardized 3 C 2 Na ) potentiometrically against potassium hydrogen phthalate (KHP, BDH certified, dried overnight at 180 °C and used without further purification). Hydrochloric acid solutions were prepared from dilutions of 12 M HCI (reagent grade) with distilled, deionized water,  and standardized potentiometrically against the standard NaOH solution. A Gran plot was used to calculate the concentration of the standard solutions.l Maltol (3-hydroxy-2-methyl-4-pyrone) and kojic acid (2-hydroxymethyl-5-hydroxyy-pyrone) were obtained from Aldrich and used without further purification. Vanadyl solutions used for potentiometric determinations were obtained from dilutions of a vanadium atomic absorption standard (Sigma). The amount of excess acid present in the solution was determined by a Gran plot.  Vanadium sulphate trihydrate, for  spectrophotometric titrations, was obtained from Aldrich, analyzed for water content, and used as supplied. ° 4 , was prepared as previously described, 2 Bis(maltolato)oxovanadium(IV), VO(ma) by the addition of two equivalents of maltol to one equivalent of vanadium sulphate. The complex was recrystallized prior to use, and fully characterized by IR, NMR and elemental I The Gran plot. (1) Before the endpoint of the titration, plot Vt vs. Vt*eT (2) After the titiation  endpoint plot Vt vs. Vt*eT. The x-intercept yields the equivalence volume, Ve (Vt is the volume of 56 ’ 55 titrant, E is the emf of the cell, F is the Faraday constant).  10  , was prepared and purified under an inert 2 analysis. Bis(kojato)oxovanadium(IV), VO(ka) 47 atmosphere and characterized in a similar fashion. Ammonium bis(maltolato)dioxovanadate(V), 2 [ma) 4 NH ( 1 V0 , was prepared by the addition of maltol to a solution of 3 VO and stirring overnight in air.” The compound 1 NIL , was recrystallized and fully characterized by IR, NMR and elemental analysis prior to use. The dioxo- anion can be obtained using a number of countercations: NH, K+ etc. Samples of VO(ma) 2 and 2 [ 4 NH ( ma)2] V0 were kindly supplied by Dr. L. Gelmini; 2 was kindly supplied by Y. Zhou, both from this laboratory. VO(ka) Concentrations of solutions were determined by dilution of accurately weighed samples in analytical standard, graduated glassware.  Where appropriate, vanadium  . 2 containing solutions were prepared and kept under an inert atmosphere, i.e. Ar or N 2.2. Spectroscopic Studies V NMR Spectroscopy 2.2.1. 51  Vanadium-51 NMR spectra (78.864 MHz) were recorded on solutions in 5mm o.d. 3 sealed in a tubes at room temperature, using a Varian XL-300 Spectrometer. Neat VOC1 5mm o.d. tube was used as an external standard, and assigned a chemical shift (6) of 0.0 57 The spectra were obtained by judicious choice of number of transients (typically ppm. -.5,000), pulse width (8  -  12  spectral window (1 x iO Hz) and line broadening (1  -  3  Hz), depending on the system under observation. The error associated with each chemical shift is estimated to be ± 0.2 ppm.  § To retard air oxidation of vanadyl containing solutions. 11  2.2.2. ESR Spectroscopy X-band ESR spectra, recorded as the first derivative of absorption, were obtained at Ca. 9.60 GHz using a Bruker ECS-180 ESR spectrophotometer. A Eurotherm B-Vt 2000 flowthrough variable temperature controller was used to provide temperatures of 100-120 K at the sample position in the cavity for anisotropic spectra. Calibration of microwave frequency and magnetic field were performed automatically with Bruker ECS-180 software, using an E.LP. 625A CW microwave frequency counter. Samples were contained in either 4 mm o.d. glass tubes, or 20 tL microcapillaries. Typically, instrument settings were: field set 1000  -  =  3200 35000; scan range -  10000; time constant  =  I  -  =  1200  -  1600 0; power =2 20 mW; gain  2 s; modulation amplitude  -  =  1  -  5 0; scan time  =  1  -  =  3  mm. Frozen solutions were prepared by slow pre-freezing of the sample from the bottom up in a Dewar of liquid nitrogen, before admittance to the spectrophotometer cavity. This lowers sample tube mortality rates due to freeze fracturing, thus protecting the cavity, and lessens solute aggregation during freezing, and facilitates glass formation. Analyses of the spectral parameters were performed using software routines developed in this department, run on a 486-PC computer. Isotropic parameters were determined by software analysis directly from the spectral data. Anisotropic parameters were determined by spectral simulation in the following manner: (1) Estimates of likely g-factors and A-values were determined from the experimental spectra. (2) A spectral simulation was obtained using 58 (3) The experimental and simulated spectra were compared, by software these values. routine. (4) Estimates of simulated parameters were refined until the comparison could not be further improved. Estimated errors in the spectral parameters were determined by the minimum change in a particular parameter necessary to cause an observable change in the simulated spectrum.  12  2.2.3. UV-Visible Spectroscopy UV-Visible spectra of aqueous solutions were recorded in 1 cm quartz cells at 25 C under Ar. Spectra were measured between 200  -  450 nm and 500  -  900 nm using a  Shimadzu UV-2100 Spectrometer linked to a data terminal. A cell containing 0.15 M NaCI was employed as a solvent blank.  2.3 Electrochemica) Studies Cyclic voltammetric data were obtained using a Princeton Applied Research Model  264 polarographic analyzer/stripping voltammeter, and a P.A.R. Model RE0089 X-Y recorder. Electrochemical measurements were carried out under an argon atmosphere at room temperature. Solution concentrations were i0 M in complex and 0.15 M in supporting electrolyte (sodium chloride for aqueous solutions). Voltammograms were recorded using a platinum working electrode, a platinum wire counter electrode and a Ag/AgCI reference electrode checked periodically relative to a 1.0 x i0 M solution of ferrocene in acetonitrile containing 0.1 M TEAP for which the ferrocenium/ferrocene 59 reduction potential was 400 mV and E =72 mV at a scan rate of 100 mV/s.  13  2.4. Potentiometric Equilibrium Measurements 2.4.1. Instrumentation Potentiometric titrations were performed using an automated potentiometric titration apparatus developed in this laboratory. The emf measurements were made with a Fisher Accumet 925 pH meter equipped with Orion Ross research grade glass and reference electrodes. The solutions were maintained at a constant temperature of 25 ± 0.1C in water jacketed beakers by means of a Julabo UC circulating water bath. The beakers were securely sealed by means of a rubber bung, with entries for the electrodes, argon gas line,  and automatic burette tip. The solutions were gently stirred with a magnetic stir bar on an insulated magnetic stir plate, and were kept constantly degassed under an argon atmosphere. The ionic strength was adjusted to 0.15 M (isotonic) by the addition of NaCL A Metrohm Dosimat 665 automatic burette was used to deliver the standardized acid and base solutions. The automated potentiometric titration system was controlled by a dedicated Acer 710 PC-XT computer running computer programs specifically written for this application.  2.4.2. Calibration The electrodes were calibrated before, and occasionally after,l each titration run by the potentiometric titration of an acidic solution of known concentration with standard A1. The 60 NaOH. The titrations were performed using the computer program CTITR results of these titrations were analyzed by a non-linear least squares fitting program (CAL ). The program reads calibration titration data (mL titrant, observed mV readings) 6 ANAL °  and refines values for electrode E and slope. Only values in the ranges pH = 2.3 2.9 and -  1 To determine if significant electrode drift had occwTcd during the course of the titration. 14  pH  =  10.8  -  11.3 are refined in the calibration, as these yield the most linear electrode  ’ Outside these regions it was sometimes necessary to calculate both F and the 6 response. slope directly from calibration data within a specific range, to give best fit in buffer regions and areas of steeply changing slope, as well as in extremes of acid and base. 2.4.3. Potentiometric Titrations Potentiometric titration apparatus was controlled by use of the computer program ° The values of electrode F and the slope were refined in the calibration and 6 TITRA1. were used to determine pH valuest from observed millivolt readings using equation 2.1. mV =  -  log [H]  .  slope  (2.1)  The critical factors that determine the accuracy of automated titrations are: (1) allowing equilibrium to be reached, and (2) waiting for pH electrode response to stabilize. The program allows for this by measuring and comparing data sets of observed emf values (and hence pH). If the variance in a measured data set is not acceptable, then the program continues to take further readings until a suitable value for the variance is obtained. Only  then will the computer instruct a further aliquot of titrant to be added. The program then waits for a pre-determined period of time, whilst stabilization of emf readings occurs, before resuming data set measurement. The program yields output of mL titrant added and observed pH. It is prudent to choose a number of experimental points that is not too low (< 20) nor high (> 120).63 Typically, —50 80 data points are obtained in each titration run. -  Titration runs were repeated with different concentrations, and on different days to ensure reproducibility. t Because of the widespread use of the term pH in potentiometric studies, it is perhaps necessary to comment on its definition and measurement. The quantity pH, defined as -logai+. is not explicitly measurable; however, under conditions of constant ionic strength maintained by an inert supporting electrolyte, activity coefficents are essentially constant. Thus the potential of the hydrogen electrode 62 In varies linearly with hydrogen ion concentration (-log[H]), as well as with hydrogen ion activity. this study pH refers to -log[Hj, as defined above.  15  2.4.4. Computations The computer program BEST” was used to determine the successive overall stability constants, (3, for each of the metal ligand systems M H L (Metal Hydrogen Ligand). The program sets up simultaneous mass-balance equations for all components present at each point in the titration curve, and calculates pH values for these points according to the set of stability constants judiciously chosen by the user, and the concentrations of each component. The program then minimizes the sum of the squared differences between the observed and calculated pH values from variations of the selected stability constants. A standard deviation, a, is given by the equation 2.2, where N is the number of titration points. 2 / (N-i) a = L (pHk. pH)  (2.2)  This represents the fit in log units for the reproducibility of the titration curve. Adjustment is continued until there is no further minimization of fit. The program incorporates a weighting factor (numerically calculated square reciprocal slope) to increase sensitivity of the calculations to buffer regions, and to offset points encountered in the less accurate regions caused by a steeply changing slope. The treatment of data using BEST allows for the variation in selected equilibrium constants to minimize the least squares fit of observed to calculated data. Of primary importance in obtaining accurate results with this computational method is the judicious choice of viable complexes to be included in the calculation. 65 utilizes a simplified version of the algorithm found The computer program PKAS in BEST and was used to calculate protonation constants in the simpler H L system.  16  2.4.5. Data Collection The equilibrium constants described in this study  are defined according to  equation 2.3:  2 xVO  +  yIP +zL  -  ,(.)z O +Y•Z )x(H+) (V + 2 X  (2.3)  Calibration Data collected in the pH ranges 2.3 2.9 and 10.8 11.3 were used in the calibration. -  -  Typically some 20-30 data points were obtained in each run. Vanadyl-maltol system (i) Protonation Constant of Maltol. Data collected from 15 titrations (—800 data points), in the concentration range 0.3 mM  [HL]  0.4 mM, pH 9 were used to determine the protonation constant of the  hydroxyl proton of maltol. (ii) Stability Constants of Vanadyl Maltol Complexes. Data were collected from 15 titrations (—800 data points), in the concentration range 0.8 mM  [HL]  5.0 mM, 0.4 mM  VO 2.0 mM, pH 7. The metal:ligand ratio was  varied in the order 1:1, 1:2, 1:3, 1:10. Vanadyl-kojate system (i) Protonation Constant of Kojic Acid. Data were collected from 6 titrations (—250 data points), in the concentration range 0.3mM [HL]0.4mM,pH8.  17  (ii) Stability Constants of Vanadyl-Kojic Acid complexes.  Data were derived from 11 titrations (—700 data points), in the concentration range 0.4 mM  [HL]  4.0 mM, 0.4 mM  ] 0.8 mM, pH 8. The metal:ligand ratio was 2 [VO  varied in the order 1:1, 1:2, 1:3, 1:10.  [VOi(ma)i] 4 N11 Data were derived from 5 titrations (—250 data points). (Conditions: 5 mM, 0.15 M NaCI, 25 C, under Ar). Data analysis was performed using the Henderson-Hassleblach equation, 2.4.  log [HA]/[A] = log 3 + log [H]  (2.4)  A plot of log [HA]I[Ai vs log [H] yields an intercept corresponding to log 3, with a slope of 1.56  18  Chapter 3 Results and Discussion 3.1. Absorption Spectroscopy Oxovanadium(IV) complexes have visible spectra that are quite distinct from other vanadium(IV) species.” The strong axial perturbation of the vanadium-oxo group has a  69 Absorption spectra have been used to significant effect in determining the spectrum. probe the variation in the ligand field environment around the metal centre in vanadyl 7074 These spectra have been used complexes with small chelating ligands and in proteins. to assign electronic transitions, and to help elucidate structural changes in solution. A convenient method for predicting the stoichiometry of a dominant species in 75 The technique has been used, among solution is the method of continuous variation. others, to study the stoichiometry of the interaction of vanadate with monosaccharides and 76 and vanadyl interaction with phytic acid and other phosphorus ligands of nucleosides, 77 The method is simple, and provides a first approximation for biological significance. further studies. A series of solutions with the mole fraction of the ligand X ranging between 0 and 1 were prepared and absorbance measurements performed at a selected wavelength (typically between 625  765 nm). As both the metal ion and the complex  -  absorb appreciably at this wavelength, and the ligand does not, the quantity determined, defined as:  =  can be  A Ath (Ath, is the theoretical absorbance due to the -  metal alone). A plot of Ai versus mole fraction X (Job plot) should have a maximum corresponding to the composition of the complex. Three maxima at XL = 0.6 were observed in the Job plot for the VO  -  maltol system, under argon, at three wavelengths in  the visible region (X= 765 nm, 700 nm, 625 nm). These correspond to a solution consisting  19  of a major component of 1:2 metal:ligand stoichiometry, and a minor component of 1:1 75. stoichiometry The visible spectra of a typical pH titration of VO(ma)2 under argon are shown in Figure 3.1. For pH <3 the spectra show a maxima at —760 nm with a shoulder at higher energy. These are similar to the spectrum of the aqueous vanadyl ion (see Figure 3.1. 71 For 3 <pH < 8 a band appears at Inset), which has absorbances at 760 nm and 625 nm. —620 nm, and becomes more defined with increasing basicity. The spectra generally indicate that there is a change in the ligand field environment around V0 in this region, 78 The band at and the spectra appear significantly different from that of the vanadyl ion.  —760 nm shifts to lower energy with increasing pH. Above pH >9, the general increase in absorbance and the observed shift to higher energies in the region 500 nm  < <  600 nm  78 A plot of £ (620 nm) vs pH (Figure 3.2.) suggest hydrolysis and/or oxidation of vanadyl. shows that for the region 3 <pH  <  8  £  remains relatively constant, indicating that in this  region the ligand environment around the VO centre remains essentially unchanged, and is presumed to be an 04 or an 05 ligand donor set. Similar inferences regarding such changes in ligand environment have been drawn from observations reported in the 70 In very acidic or very basic solutions, a different environment is observed, literature. corresponding to aquo- and hydrolyzed/oxidized species. A study by WUthrich clearly demonstrates the difference in the appearance of spectra 79 The spectra observed in this study are of 1:1 and 1:2 vanadyl:ligand complexes. ° From the 8 consistent with those previously assigned for vanadyl:ligand 1:2 complexesY’ work of Kozlowski et al, significant dimerization of vanadyl complexes is manifested by ’ hence it is also possible to 8 the appearance of four distinct bands in the absorption spectra; discount the significant formation of dimers as major species in these solutions. The predominance of a 1:2 complex is indicated from these data.  20  pH 10.5 b  f  : <  500  700 (rim)  600  800  900  .Avtg?N  Figure 3.1.  2 (1 mM, 0.15 M NaCI, 25 Variable pH visible spectra of VO(rna) Ar) Ordinate represents variation in £ (arbitrary units).  21  under  30 25.  w  20. 15. 10.  I  1  Figure 3.2.  3  I  I  7 5 •)og[H+]  I  I  9  11  2 at Plot oft (molar absorbance) versus pH for VO(ma)  =  620 nm. Data  and conditions from Figure 3.1.  2 at pH 6 was exposed to air, and In a related experiment, a solution of VO(ma) monitored at successive time intervals over a number of hours. The variations in spectra, Figure 3.3., show that there is a change in the ligand environment around the vanadyl centre with time, as the solution changed colour from green to orange. The bands at 620  nm and 870 nm decrease in intensity with increasing time, and eventually disappear, while the intense band below 550 nm remains. If oxidation of the complex occurs, then it is the d-d bands which are expected to disappear upon removing the lone 3d eIectron. The 82 and therefore the intense band below 550 nm is indicative of a charge transfer band, spectra are strongly suggestive of an oxidation to a vanadium(V) species. The oxidant is  presumed to be dissolved oxygen, as evidenced by the lack of oxidation of solutions that are stored under argon. Saito has observed a similar oxidation for the conversion of  l nta)] 83 ( [VVO . [VWO(nta)] to 2  § Note, however, that the onset of oxidation in vanadyl complexes is both concentration and pH dependent. I ma refers to nitriloacetic acid.  22  \ime  (Hours) .24 10  1i3 I  I  500  600 700  1  800  —1 900  )(nm)  Figure 3.3.  Visible spectra at various times after the exposure to air of an aqueous solution of VO(ma)2 (1 mM, pH 6.0,0.15 M NaC1, 25 C).  23  t who proposed ’ 8 Similar observations were made in a study by Stewart and Porte, that bisligand-vanadyl complexes, such as those investigated, undergo reversible oxidation in aqueous solution to form dioxobis(ligand)vanadate(V) anions. The mechanism by which this one electron oxidation with molecular oxygen proceeds is not yet understood. A 2 proceeds is given in an unbalanced possible pathway by which the oxidation of VO(ma)  form in equation 3.1.  2 VO(ma)  OII{O -  ma) (3.1) 7 ( 2 [V0 J  In Porte’s study, the sixth coordinated ligand was postulated as being trans to the oxo - is (ma) 2 [V0 2 is oxidized and that anionic ] group. It appears likely that VO(ma) subsequently formed in these solutions. A crystal structure determination performed at 45 Hence, initial oxygen (ma) is a cis complex. 2 K[V0 U.B.C. recently proves that ] substitution might occur in the trans position to V=O, followed by apical to basal rearrangement to form a cis product. Stewart and Porte also note that no evidence was 85 found for dimer formation. In the analogous kojic acid complex similar observations were made. In the variable 2 the appearance of the spectra are also strongly indicative of pH visible spectra of VO(ka) the dominance of a 1:2 species throughout the neutral pH region. A solution exposed to air,  and monitored by visible spectroscopy with time, shows that gradual oxidation also occurs in this instance. The assignment of bands in the absorption spectra of vanadyl complexes has been a ’ Typically three bands are observed in the visible region: 67 matter of some controversy. bands I (between 910 625 nm), II (between 690 520 nm), and III (between 475 -  -  nm). Molar absorbances of 5  -  -  330  1 are typical. The energy level scheme 150 M-’ cm-  developed by Ballhausen and Gray is perhaps most often used to interpret  86 However, some studies have questioned the universal oxovanadium(IV) visible spectra. 24  87 and there is no real agreement on the ordering validity of the Bailhausen-Gray scheme,’ of the vanadyl energy levels. The V=O bond stretching frequency, UV, is an important characteristic of 67 The magnitude oxovanadium(IV) complexes, and is generally observed at 985±50 cm. of the stretching frequency is particularly sensitive to the a-donor qualities of the other ligands bound to the vanadyl centre. Donors that increase the electron density of the vanadium centre will reduce its acceptor properties towards the multiply bonded oxygen  (decreased pit  -+  dit donation). A decrease in the vanadium-oxygen bond order is reflected  in a decrease in uvo.  4 are It has been found in general that values of uvo —950 cm  indicative of octahedral geometry, and values of —980 - 1000 cm-’ of square-pyramidal  68 ’ 67 geometry with the V=O axial, and having no ligand trans. 2 are given in Table 3.1. 2 and VO(ka) Visible and IR spectral parameters for VO(ma) The data reveal that the spectral parameters for the ligand complexes of interest are 89 ’ 88 comparable to those previously obtained for similar vanadyl species.  Table 3.1.  . 2 2 and VO(ka) Visible and IR spectral parameters of VO(ma)  (M-’c&)  uvo (cur’)  Complex  ) (cm-’)  £  2 VO(ma)  625, 86O  15, 30  995  2 VO(ka)  610, 8401  17,32  980  pH 5.0  I  pH5.5  In all these visible spectra the tail end of a high intensity charge-transfer band is seen below 500 nm in the UV region of the spectrum. Vanadium(V) species however, lacking d 0 in ma)2] N}Lt[V ( electrons, show no transitions in the visible region. A pH titration of 2 the UV region is shown in Figure 3.4. A peak at  25  =  273 nm has a maximum absorption  corresponding to lowest pH. Between pH 1.5 and 4.0 the gradual decrease in intensity of this absorption is observed, with a single isosbestic point noted at 292 run. In the pH range 4.5  -  7.5 the spectra are almost identical, indicating that only one major species exists in  this region. Above pH 8.0 the increase in intensity of an absorbance at ) = 318 nm is observed, with a concomitant reduction in the intensity of the band at 273 nm. A single pseudo-isosbestic point is noted at 295 nm. The bands correlate to those of protonated ( = 273 nm) and deprotonated Q. = 318 nm) maltol ligand, as determined in this study, and it is likely that the bands are due to it -÷ it transitions within the ligand. In acidic solution the bands possibly denote species involving protonated maltol ligand and vanadyl and vanadate-maltol complexes. In the neutral pH region it appears likely that anionic and/or ’V NMR data reported in Section 3.4., and the protonated [V02(ma) ] is prevalent, from 5 2 spectra are assigned thus. In basic solution vanadates predominate, hence the band observed at 318 nm presumably corresponds to uncomplexed free ligand, L-.  26  .  0  g .  ‘P  C C  0  C C C 0 ‘P  0 C 0 0 0 0  Figure 3.4.  25C  300  260  320  350  ) (2mM, 0.15 M NaC1, 25 C, 0 H ia V n N [ 4 ( ] Variable p1-I UV specta of 2 ction of peak intensity with increasing under Ar). Arrows indicate the dire pH.  27  3.2. Electrochemistry Elecuochemical data are useful in obtaining information about the redox properties Vm, V”, and VV of metal complexes. Of the common oxidation states of vanadium, only reducing are of any significance when considering physiological systems, V being far too oxidation to exist under normal conditions. In acid, the predominant species for these . As the p14 is raised, hydrolysis occurs and a number 2 2 and cis -VO states are V, VO e pH reduction of monomeric and oligomeric vanadium species are formed. The variabl  potentials for various vanadium species are shown in Figure 3.5.  w  vs z  w  Figure 3.5.  Reduction potentials (versus S.H.E.) for various vanadium species as a in function of p11. Boundary lines coirespond to values where species indicate adjacent regions are in equal concenuations. Short dashed lines uncertainty in the location of the boundary.  28  In the systems under observation V’’  -  V” is most likely to be the only important  , have shown that the H 0 redox couple. Previous electrochemical studies, performed in 2 potential for this couple in five coordinate vanadyl species generally lies between +0.3 V 1 (vs S.H.E.), for example in 89 ((3,5-Dimethylpyrazolyl)borato)dichlorooxovanadium(IV), 9 ’ and +0.9 V. as observed for (N,N-Disalicylidenecthylenediamine)oxovanadium(IV), and is dependent upon the ligands bound to the vanadyl centre. In cyclic voltammetry, the potential of a small, stationary, working electrodeê is changed linearly with time starting from a potential where no electrode reaction occurs and moving to potentials where reduction or oxidation occurs. After traversing the potential region of interest, the direction of the linear sweep is reversed. The reversibilty of a cyclic voltammetric wave is usually compared to that of a known system. One such standard is the ferrocenium/ferrocene redox couple, a one electron reversible process with tE =72 mV (versus S.H.E.), Ic/ia =1, at 100 mV/s.l To investigate the redox behaviour of these systems the variable pH electrochemistry of VO(ma) 2 was studied using cyclic voltammetry (Figure 3.6.). The measurements were referenced to the ferrocenium/ferrocene reduction potential. For 5 <pH < 8, an irreversible 1 >1 (Table 3.2.). As the acidity of the solution is wave is observed; E >>72 mV, i /i increased, the wave becomes increasingly quasi-reversible (jc/ial). The wave obtained at pH 2.5 closest resembles reversibility as defined by the ferrocenium/ferrocene redox couple (zE —72 mV,  =l), and from ESR measurements reported in Section 3.3., is likely to 1 IC! i  . In the extremes of acid and 2 correspond to a solution containing predominantly VO(ma) base the waves obtained appear similar to those of oxovanadium(IV) and vanadium(V)  2 value of 0.77 V versus the standard f 1 respectively, as determined in this study. The E hydrogen electrode at pH 2.5 lies within the range of the expected values for the V” -+ VW  i.e.Ptasusedinthisstudy. I Where: E  Ecathodic -Edic, ic = cathodic current, 1 a = anodic current.  29  redox couple.t The process occurring can be summarized as in equation 3.2. The marked dependence of (quasi-) reversibility with pH can be explained by the increasing extent of hydrolysis of the complex with increasing basicity, and by the steady conversion of 2 to various vanadium(V) species, as implicated from ESR and 51V NMR studies VO(ma) presented herein. Hence, depending on the pH of the solution, a mixture of species is obtained. The redox couple corresponding to vanadium TV ++ III frequently has potentials ranging from 0.2 V to 1.0 V; in this study no waves were observed in the region 1.4 V -  -  to 0.0 V or above 0.75 V, indicating that there is only one redox couple of importance involved in these systems.  O(ma) 2 ’ 1 [V ]  [VVO(ma)z]4  —  +  e (3.2)  2 the In the variable pH electrochemistry of a corresponding solution of VO(ka) observed waves are irreversible in all pH regions; àE >>72 mV,  /i >1 in each case  (Figure 3.7.). In addition to the reasons stated for solutions of VO(ma)2, these observations may be attributed to the effect of the hydroxyl group of the kojate ligand. This may react with neighbouring vanadium centres to form kojato chain complexes, hence causing the . Further studies are necessary to 2 irreversible redox behaviour observed with VO(ka) establish this.  t Potentials reported versus the standard hydrogen electrode are derived from those experimentally obtained using AgfAgCl electrodes as follows: E (S.H.E.) = E (AgIAgCI) + O.2223.  30  25  55  3D  60  3  65  40  775  2PjsAt  4.57w O•075  .075  Figure 3.6.  o  VO(ma) (0.15 M NaCI ,25 C., under Ar, Variable pH electrochemistry of 2 scan rate 100 mV/s. electrode: Ag/AgCI)  31  Table 3.2.  Redox potentials (Volts) obtained from a cyclic voltammetric investigation of VO(ma) 2 at variable pH (25 C, 0.15 M NaC1, under Ar).  pH  Ei (V)  liE (mV)b  ic/ia  2.5  0.77  71  1.00  3.0  0.73  68  1.03  3.5  0.70  65  1.03  4.0  0.69  64  1.01  5.0  0.67  55  >>1  a Potentials reported vs standard hydrogen electrode (E Ag!AgC1 b =72mV,ic/ia=1.00forFc/Fc.  +  0.222); ± 0.01 V.  V0 gave irreversible waves for pH [ma) 4 NH ( ] The variable pH electrochemistry of 2 >4 (Figure 3.8.). As the pH is lowered, the waves become increasingly quasi-reversible. 1 =0.70 V, E =67 mV and i/i =—1. This corresponds to the V”/V” redox At pH 2.5: E . ESR observations 2 couple, and quite reasonably, appears similar to that for VO(ma) V0 [ma) 4 NH ( ] presented in Section 3.3. indicate that as acidity is lowered in a solution of 2 2 are observed. In addition, the first ESR signal detected increasing amounts of VO(ma) occurs at pH <4, at which point quasi-reversibility is also manifested. Hence, the findings 2 exhibits reversible electrochemical behaviour in aqueous appear to indicate that VO(ma) 1 solution, whilst its hydrolysis and oxidation products do not. Differences between E V0 determined here presumably are due to [ma) 4 NH ( 2 2 and reduced ] values of VO(ma) 2 present in either solution, the “concentration” of the differing amounts of VO(ma) 2 being both pH and concentration dependent. VO(ma)  32  ===  70  30  7.5  80  40’ 4.5  90  i 6  50 .075  Figure 3.7.  0.075  0  VO(ka) (0.15 M NaCI ,25 C., under Ar, Variable pH electrochemistry of 2 scan rate = 100 mV/s, electrode: AgIAgC1)  33  pH60 5.5 50’ 4.5 40  30 2OA{  .25  .0175  Figure 3.8.  0•  (ma) (0.15 M NaCI ,25 C., NH4VO ] Variable pH electrochemistry of 2 under Ar, scan rate = 100 mV/s, electrode: Ag/AgCI)  34  3.3. ESR Spectroscopy Vanaclyl complexes have a low lying singly degenerate ground state due to their near  1 electronic idealized C4v geometries, axial distortion due to the oxo group, and 3d 0 configuration. This makes vanadyl complexes valuable as spin probes in ESR. The d vanadium(V) species do not exhibit ESR spectra. In general, the VW oxidation state requires a noncubic field for ESR observability, fulfilled in vanadyl complexes. Another simplifying feature of these spectra is the absence of interelectronic effects, since there is 94 ESR is a valuable speciroscopic tool, being several orders of only one d electron. magnitude more sensitive than most other spectroscopic techniques. The ESR spectra of vanadyl complexes at room temperature exhibit characteristic widely-spaced eight line patterns due to the coupling of the unpaired electron with the large moment of the —100% 1V nucleus (I = f2). abundant 5  In addition to the intensity and line shape of the absorption, two fundamental parameters can be derived from the ESR spectral measurement. These are the Landé effective electron g-factor and the electron-nuclear spin coupling constant A. The isotropic  0 and g 0 can be determined from the position and spacing of the eight ESR parameters A resonant lines for the room temperature spectrum. Frozen solutions of axial d’ systems (e.g. vanadyl complexes) show two sets of overlapping eight line resonance components, one set due to the parallel (N) features and the other due to the perpendicular (J..) features. Ideally, ESR spectra can yield much useful information to the coordination chemist: (1) identification of the metal, oxidation state, and spin state; (2) identification of the binding site (ligands) and local symmetry; (3) determination of the concentration of the , are highly 1 95 The hyperfine coupling constants A, given in lO cmparamagnetic ion. sensitive to the ligand field around the vanadyl centre; linewidths to the size of the complex 7 This has been demonstrated in the fine structure solution. 9 ’ and to the viscosity of the 96  35  observed in the ESR spectra of vanadyl albumin, which clearly shows two distinct binding s. These large protein complexes exhibit near rigid limiting spectra due to 98 environment 72 It is also possible to obtain spectra of samples containing slow molecular tumbling. minute concentrations of vanadium, thus allowing for the study of tissue and sub-cellular fractions. Studies of this kind have been performed on organs from vanadium treated 101 00 and as a probe of pH in intact vanadocytes from Ascidia cerawdes. diabetic rats,” The use of ESR is probably one of the most powerful tools in the study of vanadyl 102 containing species. ) in aqueous solution is of 5 O [VO(H 2 The hydrolysis of the vanadyl ion ] considerable importance when considering the ESR spectra of vanadyl complexes. The  2 spectrum loses intensity as the pH is increased from 2 to 4 due to the formation of an VO . A sharp decline in the intensity of the spectra 2 ESR-silent hydrolysis species, [VO(OH)] after pH —4 marks the onset of the formation of a grey precipitate of [VO(OH)]. By pH —6 precipitation is complete, and the solution is ESR-silent. No further signals are - forms. 3 observable until pH >11, when significant amounts of ESR active VO(OH) These observations have an important ramification; between pH —5  -  10, ESR signals  observed in solution must be solely due to vanadyl-ligand complexes, and have no attributable contribution from vanadyl hydrolysis products. 2 are depicted in Figure 3.9. Below pH 1.5 The variable pH ESR spectra of VO(ma) O) VO(H , 2 the spectra (not shown) appear similar to those observed in this study for 5  2 reported in observations that concur with results obtained in the visible spectra of VO(ma) Section 3.1  Analysis and comparison of the spectral parameters confirm that in highly  O) In the region 2 <pH <9, a VO(J4 . 2 acidic solution VO(ma)2 is converted to 5 O) is observed, having greater intensity in more VO(H 2 spectrum distinct from that of 5 acidic solution. Analysis of the spectral parameters, given in Table 3.3., reveals no 0 throughout the pH range. Hence, it may be concluded that there is difference in g 0 or A  36  only one major vanadium(IV) species throughout this region, and based on a 1:2 . 2 metal:Iigand dominant species determined in the visible spectroscopic study, it is VO(ma) Spectral integration reveals that between p144-7 the spectra arc all of equal intensity, and hence “concentration”. Assuming that in the neutral pH range ESR signals are solely due to vanadyl-ligand complexes, then in neutral pH 65% of the total vanadium in solution is 0) VO(H , 2 present as VO(ma) . This is detennined from spectral integrations relative to 5 2 which is assumed to correspond to 300% of the total vanadium in solution under highly 103 As basicity is raised spectral intensity wanes, and by pH 10 the acidic conditions. spectra has all but disappeared into the baseline. This is consistent with conversion of  V NMR spectra reported in Section vanadium(IV) to vanadium(V), as indicated by the 5 2 (not shown) revealed similar 3.4. The corresponding variable pH ESR spectra of VO(ka) trends; spectral parameters are reported in Table 3.3.  3000  Figure 3.9.  3200  3400  pE 9.0 8.0 7.0 6.0 5.0 4.0 3.5 3.0 2.5 2.0  3600  3800  Variable pH ESR spectra of VO(ma)2 (0.15 M NaC1, 25 C, under Ar, spectra depicted on an equal intensity scale).  37  V0 (not shown) no [ma) 4 NH ( 2 In the variable pH ESR spectra of ESR-silent d° ] spectra are observed above pH 4. As the acidity of the solution is increased a typical vanadyl ESR spectra is observed in increasing intensity. Analysis of the isotropic spectral parameters indicates that this species is VO(ma)2, as the parameters are identical (Table ) forms, as evidenced by the appearance of the spectra, and 5 O VO(H 3.3.). Below pH 1.5 2 the spectral parameters. It is apparent therefore that the vanadium(V) species has been reduced and that there exists a facile route for the interconversion between these vanadium(IV) and coffesponding vanadium(V) ligand complexes in aqueous solution. 2 is given in A closer examination of the room temperature spectrum of VO(ma) Figure 3.10., and shows the presence of a minor species in the spectrum, indicated in the figure by an arrow. The nature of this minor species may be attributed to an isomer of VO(ma) with trans ketonic oxygens. The presence of suchcis and trans isomers of 2 .’° An alternative 5 vanadyl complexes has previously been resolved by ESR spectroscopy ) Further studies are needed to 5 O VO(H . explanation is that the distortion is due to 2 confirm its identity.  38  o k%%II 0 1 % 0 3 H  3 H  srans.bis(maholato)oxovanadjumav)  th.bis(makolato)oxovanadjumay)  X.Band ESR 298 K  eb•  (ma) 2 VO  (ka) 2 VO  0 ma) V ( ] !2 4 NH  m=  “‘2  I  3000  3200  3800  3400  I  I  I  3800  000  1G3  ) VO(H O . 2 0 and 5 a) R V m N [ 4 ( ] VO(ka) 2 , VO(ma) 2 , Figure 3.10. 1soopic specua of 2  39  [o..v.o=1o3.7’  Figure 3.11.  I  . (ma) K[V0 ] 2 Chem 3.DTM representation of the structure of 45  , Figure 3.10., it was not possible to 2 In the corresponding ESR spectrum of VO(ka) unequivocally determine if such a minor species exists. However, comparison of the O) reveals a spectral distortion, VO(H 2 spectra of these pyrone complexes with that of 5 particularly evident in the ni  =  -  /2, /2 and /2 lines.  40  . V0 [ma) 4 NH ( 1 ,2 2 Isotropic ESR spectral parameters of VO(ma)2, VO(ka)  Table 3.3.  Complex  &) (g 0  b) (10 cmA ( 0 ) 1  2 VO(ma)  1.963  103.7  2 VO(ka)  1.963  104.5  V0 [ma) 4 NH ( 2 ]  1.963  103.6  ) 5 O VO(H 2  1.968  106.7  a±OO(j1 b±01 (10 cm-’). The anisotropic spectral parameters for the complexes of interest were determined by spectral simulation of the frozen glass spectra obtained at —110 K. The anisotropic 2 are shown in Figure 3.12. spectrum, and corresponding simulated spectrum, of VO(ma) Often X-band spectra can be described by axially symmetric tensors, in which case g = gzz, g.L = g,  =  11 = gyy, A  and A.L =  =  Ayy. Simulation of the anisotropic spectum  2 using these parameters failed to adequately describe the experimental of VO(ma) spectrum. In a study of the analogous anisotropic ESR spectral parameters of VO(acac)2, 6° Rhombicity is the authors determined that the complex exhibited slight rhombicity.t’ manifested in cases where xx  -  g>  0.006 and A 1  =  4 x10- cm1 (for axes, see  28 Upon closer examination of the frozen glass spectrum of VO(ma)2, Figure 3.12. inset).  the manifestation of rhombicity is evident from the position and spacings of spectral lines, 2 isomer further particularly in the central region of major overlap (the minor VO(ma) complicates the spectrum). The spectrum is satisfactorily simulated with a slight rhombic distortion with parameters g = 1.979, g = 1.973, g = 1.938 (g g = 0.006), A = -  t VO(acac) 2 refers to bis(acetylacetonato)oxovanadium(IV). 41  21 , =60 and A 3 55, A,  =  169 (A  -  =5 (10-4 cm’)). By this method the frozen glass  2 was similarly simulated with slight rhombic distortion, whilst the spectrum of VO(ka) , quite reasonably, is coincident with that of V0 [ma) 4 NH ( 2 frozen glass spectrum of ] , and has identical parameters (Table 3.4.). 2 VO(ma) Table 3.4.  , 2 Anisotropic (simulated) spectral parameters for VO(ma)2, VO(ka) ) 0. VO(H ’ 2 (ma) and 5 NHJVO ] 2  Complex  g,(a)  g(a)  g(1)  Acj) cm4 (10) 1  AyyO’) cm’) 4 (1O-  Azz ) 0 cml) t (1O-  2 VO(ma)  1.979  1.973  1.938  55  60  169  2 VO(ka)  1.979  1.973  1.938  58  62  170  4 NH (ma) 2 [V0 ]  1.979  1.973  1.938  55  60  169  ) 5 O VO(H 2  1.986  1.986  1.934  70  70  180  a± 0.001. b± 1 (l0- cm ). 1  42  X.Band ESR 110K  2 VO(m a)  Experimental  Simulated  -4  250G  Figure 3.12. Experimental and simulated anisotopic ESR specua of VO(ma)2.  43  In general, the ESR spectra of V02+ complexes are strongly dependent on the donor atoms bound to the metal; e.g. they are different for nitrogen and oxygen donors. There are also observable, but smaller, differences with different types of ligands that contain the same donor atom. An analysis of the spectral parameters for these vanadium-pyrone complexes, Tables 3.3. and 3.4., reveals a distinct similarity between parameters. It is first noted that the g values remain invariant between the maltolato and kojato complexes, even though the latter has an electron-withdrawing hydroxymethyl group. In general, this has been found for other ligand types in oxovanadium complexes. For example, fluorine or phenyl substitution does not affect the g values of -diketone, -ketimine and -diketonate ° The hyperfine splittings 8 ” 7 ’° 95 complexes. substituent effects, as evidenced in the 1  -  ,  A and  are more sensitive to  1 change upon hydroxymethyl 4 cnr 3 x10  . It is noted that, except for the most powerful 2 2 relative to VO(ma) substitution for VO(ka) electron-withdrawing or donating groups, the changes in Ac, in oxovanadium(IV) 95 Chasteen has correlated ESR spectral complexes upon substitution are usually small. data with ligand type for a variety of comp]exes. The determined spectral parameters are similar to those previously determined for 04 ligand coordinated oxovanadium(W). 2 treated S1’Z-diabetic rat In a preliminary experiment, the blood plasma of a V0(ma) was investigated at -110 K by ESR spectroscopy. As blood plasma contains many 2 whose ESR signal appears in the region of paramagnetic impurities, particularly Mn interest, it was necessary to improve the spectrum to determine if any vanadyl signal could be observed. This was achieved by subtraction of the ESR spectrum of the blood plasma  2 untreated) rat from the experimental from an STZ-control (i.e. diabetic, but V0(ma) spectrum. The final spectrum (experimental minus control) is given in Figure 3.13., and . The signals are 2 clearly shows prominent V0 lines, in addition to those caused by Mn likely to correspond to protein-bound vanadyl; the spectrum is annotated according to x. 0 comple 1 ” 1 spectral lines expected for a vanadyl-rat transferrin 72  44  2600  3000  3400  3800  4200  4600 (6)  Figure 3.13. Possible ESR spectral lines due to vanadyl-rat transferrin. (Subtraction of non-treated STZ-diabetic rat plasma spectrum from VO(ma)2 treated STZ-diabetic rat plasma spectrum)  45  3.4.  S1V  NMR Spectroscopy  Vanadium has a particularly attractive nucleus for NMR spectroscopy. However, it t vanadium(IV) signals is only the d° (VV) oxidation state that is observed. Paramagnetic d do not appear in the experiment; the large electron-nucleus hyperfine splitting constant for 111 paramagnetic species causes large chemical shifts, and extreme line broadening. Nuclear properties of 51y are given in Table 3.5. Table 3.5.  112 V nucleus. NMR properties of the 51 99.76 %  Natural abundance Nuclear spin H = 1.00) 1 Relative receptivity ( 3 Chemical shift range relative to VOCI  0.38 +200 to -1000 ppm  Vanadium has a relatively large shielding range that manifests itself in the sensitivity of the 51V nucleus to minor changes in its chemical environment. The chemical shift (6) is V NMR spectroscopy, and is probably the most important parameter discussed in 5 strongly dependent on the electronegativity of the ligands surrounding the nucleus. In general, the electronegativity criterion can be used to distinguish between a nucleus surrounded by thio-ligands (6 -300 to +200 ppm) or oxo- or nitrogen ligands (6 -700 to 113 There is also a marked increase in shielding, and hence 6, observed when -300 ppm).  introducing steric strain via 5- or 6- membered chelate rings, with bulky ligands, and to a 114 lesser extent with increase in coordination number.  Line widths for the vanadium  complexes encountered in these spectra are relatively broad, No correlation of line width to ligand environment is noted in the literature, except the general observation that the magnitude of line width parallels ligand symmetry about the V’ centre, narrower in more s. Line widths are not noted in this study, apart from such 57 symmetric environment generalizations.  46  ‘v  NMR spectroscopy has been employed in many investigations of biological  interest. For example, the ‘V nucleus has been utilized as a metalloenzyme probe of the 20422 and to investigate ” 7 1 12,1 13.1 ‘ 6 to probe vanadate-protein interactions, transferrins, t ” 115  V NMR is the speciroscopic probe of choice 51  vanadate as a phosphate analogue. for investigating vanadate systems.  As a reference for further studies, the hydrolysis behaviour of vanadate was studied by pH  titration.  In acidic solution the dominant vanadium(V) species is “free” VO, and  appears as a broad line at 8= -544 ppm (Figure 3.14.), indicating the low symmetry of this species.l At pH 4 the dominant peaks appear at 8= 422, -497 and -514 ppm in a 1:2:2 ratio and are all attributable to decavanadate. The structure of decavanadate consists of ten 6 octahedra sharing edges, with two oxygen atoms in the interior of the molecule (VO) 126 surrounded by six vanadium atoms.  The result is three different vanadium  114 Minor resonances are also seen at 6 = environments, and hence three chemical shifts. -559, -569 and -580 ppm, and are due to the following vanadate species: monomeric  4” ) respectively. 5 , tetrameric vanadate (Vt) and pentameric vanadate (V (V ) vanadate 1 These are observed with greater intensity at pH 7, but are shifted slightly. The dominant , 8= -576 ppm; while a new peak corresponding to dimeric 4 species at neutral pH is V , appears at 8 2 vanadate, V  =  -572 ppm. All these peaks are relatively narrow and sharp,  indicating high symmetry around vanadium in these complexes. At pH 9, the major species , whilst V, dominates at pH 11. The deprotonation of these species accounts 4 1 and V are V 2 (8= -572 to -560 for the variation in chemical shift for V, (6= -559 to -536 ppm) and V ppm). Chemical shifts are summarized in Table 3.6., and are in good agreement (±1 ppm) with those obtained by Pettersson er a!. in an extensive study using a combination of 127. V NMR spectroscopy potentiometric titrations and 51  I  spectroscopic utrations in the acidic regime for this complex cearIy show that reduction to a vanadium(TV) species also occurs, and indeed V(IV) is the dominant oxidation state.  ESR  47  vi pE  JI:  2.1.0  9.0 7.0  _10  _AaJL  4.0  2.0 •  —400  Figure 3.14.  •  -440  -480 -520  —560  -600 6 (ppm)  25 ‘V NMR specta of a pH titradon of vanadate (10 mM, 0.15 M NaC1, •C under Ar).  48  Chemical shifts (6) of vanadates present in aqueous solution (Data obtained  Table 3.6.  in this study). Vanadate  V 1 2 V 4 V 5 V 10 V 2 VO  Proposed formula  HV0 2 4 3 7 0 2 HV 12 O 4 V 15 0 V 5 oO 1 V 6  8 (ppm)  pH range  -559to -536 -572 to -560 -576  3-12 5 10 5-9 5-9 3-6 1-4  -584 -422 -497 -514 -544  -  The reaction of vanadate with increasing concentrations of maltol was studied in order to examine the nature and stoichiometry of the species formed in solution. To a 10 mM solution of vanadate at pH 7,* were added increasing equivalents of maltol, and the  ‘v NMR  spectra observed (Figure 3.15.). With no maltol added, the only observable  4 (8= -576 2 (8= -572 ppm), V species are the vanadate oligomers, Vi (6= -559 ppm), V 5 (6= -584 ppm). For M:L ratios of 1:’,2 to 1:10, the gradual appearance of ppm) and V peaks due to increasing complexation, nominally assigned as peak a (6= -496 ppm) and peak b (8= -509 ppm), are observed with higher ligand concentrations. There is no change  in the chemical shifts of the vanadate oligomers, and no other resonances are detected. As the ligand concentration is increased, and the metal-ligand complexes are formed in increasing amounts, one can observe the concomitant reduction in the intensity of the peak ). These observations indicate that peaks a and b 5 4 (and to a lesser extent V assigned to V are due to vanadate complexation with maltol.  § From ESR observations, reduction to vanadium(IV) does not occur above p114.  49  a  Ji1 1:2  1:1  —  1:0.5  —  jV4 3:0 .iIipI..1 ,,WJ1i .j piiI IpIIIpI(IpIIIJIIIIjIIIIeIIoIIjI IlIIIipiI 1 i.I.pfl I I flIl1j14lII4IiI II 1  -400  440  -480 -520  -560  -600 ‘6 (ppm)  increasing Figure 3.15. ‘V NMR spectra of the reaction of vanadate (10 mM) with mM. concentrations of maltol (0 mM, 5 mM, 10 mM, 20 mM, 50 mM, 100 Conditions: 0.15 M NaC1, pH 7,25 ‘C, under Ar). 50  V NMR spectra The concentrations of the various vanadate species observed in the 51 were determined by integration. The NMR sample contained a known total amount of  vanadate; hence, having first calculated the relative ratio of each integrated resonance, the concentration of each species could be determined, and is equal to the product of total vanadate concentration and the relative ratio of that integrated peak. Following the method 124 the signal corresponding to a, 8= -496 ppm, was analyzed for maltol ’ 2t of Tracey stoichiometry; the integrations of peak b were too small to determine stoichiometry with any degree of accuracy. Using this method, a 1:2 vanadate:ligand stoichiometry is determined for resonance a. The reaction of vanadate and maltol at increasing vanadate concentrations was studied by  ‘v  NMR, the spectra obtained shown in Figure 3.16. The total maltol  concentration was 10 mM and the pH was 5.5, chosen to study the interaction of the complexes with decavanadate. With a large excess of ligand (metal:ligand ratios of 1:20 to 1:5) the only observable resonance is peak a (6  =  -496 ppm), which further supports  evidence that a 1:2 complex is the major vanadate-maltol species. For a metal:ligand ratio 1 (8= -560 ppm) are observed. For a 1:1 ratio, of 1:2, peaks a, b (6= -508 ppm) and V 0 signals are seen in their usual V 5 and 1 ,V 4 ,V 1 appreciable vanadate formation is noted; V positions. In excess vanadate (5:1 ratio) peak b is not observed, and peak a appears as a shoulder on the 8  =  -497 ppm decavanadate peak. A concentrated aqueous solution of  V0 yields a lone resonance at 8= -495 ppm. Studies in non-aqueous solvents [ma) 4 NH ( 2 ] (ma) anion persists as the major species of a 2 [V0 further support evidence that the ] V0 Peak a is therefore assigned to anionic [V02(ma)2]. ma)2]; solution of 2 [ 4 NH (  51  a 0 V  lvi  1  Ab  N:Z. 5:1  JL  1:1  1:2  -—--  1:5  1. —  1:20 ,.ie.J..  In Ij I IjIII IIIII4IIIIIIIIIjII  -400  Figure 3.16.  -440  -480 -520 -360  -600 5 (pp)  ‘V NMR specua of the reaction of maltol (10 mM) with increasing vanadate concentations (0.5 mM, 1 mM, 2 mM, 5 mM, 10 mM, 50 mM. Conditions: 0.15 M NaCI, pH 5.5,25 C, under Ar). 52  The metal-centered oxygen atom transfer reactions of vanadium have been likened to OV-O2 -OR (R =aikyl) to an organic esteT, L 2 those of organic functional groups: VOL ” Thus, an “anhydricle” can be esterified, acylated or base 1 2 to an acid anhydride. VOL V0 alkylated derivatives have been synthesized and [ma) 4 NH ( ] hydrolyzed. A number of 2 45 In addition, the formation of a characterized (i.e. , 2) R= Et, Me, iPr). VO(OR)(ma VO(OH)(ma) have a) ) and a protonated complex, 2 , dinuclear species, ] V-O-VO(ma [(m 0 2 V NMR chemical 45 51 been proposed in organic solvent solutions of these derivatives. shifts of these complexes are summarized in Table 3.7. O was studied by variable pH 51V NMR; ma) NH4V ( ] The aqueous chemistry of 2 the spectra are shown in Figure 3.17. In pH 2 solution two very broad resonances are observed at 8 = -460 ppm and 6= -510 ppm. As the pH is raised to 3.5, the two peaks have  increased in intensity, and are located at 8= -483 and -509 ppm. The integrated peak areas indicate that the two species are in approximately equal proportion at this point. The position and nature of these two signals indicate that they are due to vanadate-maltol complexes, as they lie outside the ranges of vanadate hydrolysis products. The chemical shifts of the two peaks continue to coalesce until the pH is raised to 4.5, upon which the -, (ma) 2 [V0 signals are seen at 8= -496 and -509 ppm, corresponding to peaks a, anionic ]  and b in the previous investigations. A minor peak at 8= -519 ppm is also observed at this point, which is assigned to decavanadate (8 presumed that the 6  =  =  —-520 ppm). From Figure 3.16. it is  -497 ppm decavanadate peak is located under, and contributes to,  resonance a. In the neutral pH region, vanadate hydrolysis products appear in the spectra, 4 (6= -576 2 (6= -570 to -560 ppm) and V and are assigned to Vi (8= -560 to -537 ppm), V ppm). For pH >3.5, resonance a grows in intensity, whilst peak b decreases. Above pH 8, 1 is the 1 begins to dominate to the detriment of the other signals. In basic solution, V V major species.  53  The partial protonation of the maltol ligands in the coordination sphere of vanadium 0 as the solution is acidified result in the formation of 2 and their displacement by H )] (x = 1 or 2) which, in the acidic regime, can be assigned as a contributor (ma)(H 2 [VO O to resonance b (6= -509 to -529 ppm). This is implicated by the relative intensities of the  peaks as the pH is reduced. These observations are in agreement with those previously A1 NMR spectra of AP’ complexes with pyrone and observed in the variable pH 27 ’ In addition, from chemical shifts noted in Table 3.7, it is also 3 ” 30 pyridinone Iigands.’ possible that in acidic solution the resonance has some contribution from the protonated .) 2 [VO(OH)(ma species ]  In acidic solution the resonance is very broad, and may  correspond to overlap from several peaks. It is also noted that in neutral to basic solution resonances a and b do not change in intensity, under conditions in which neither a mono ligand or protonated complex would be expected to exist. Therefore, it is postulated that peak b is due to another species in this pH region; a possibility being another isomer of cis bis(maltolato)diooxovanadate(V), with one maltol ligand inverted by 180 so that a hydroxy oxygen is trans to a V=O (see Scheme I). In effect, several distinct vanadate species have a similar chemical shift, 8  =  -509 ppm, that manifest themselves under  conditions in which formation of that species is favoured. Further studies are necessary to unequivocally assign this resonance with increasing pH. Chemical shifts of these vanadate maltol complexes are given in Table 3.7. V0 is [ma) 4 NH ( 2 It has been shown using ESR that an appreciable amount of ] reduced in acidic solution to VO(ma)2 (Section 3.3.). Therefore there will be an increase in paramagnetic vanadium(IV) complexes at low pH. The variation in chemical shifts of peaks a and b observed with increasing acidity in the pH study is attributed to paramagnetic shifting. The lower the pH, the greater the degree of reduction to V(JV) complexes, hence the greater concentration of paramagnetic species in solution and the more exaggerated the variation in the chemical shifts.  54  pE 12.0  —400  —440  —480 —520  —560  —600  2.0 (ppm)  [rna) 4 NH ( ] V0 (5 mM, 0.15 M NaC1, Figure 3.17. Variable pH V NMR spectra of 2 25 ‘C, under Ar). 55  Table 3.7.  Chemical shifts and assignments for vanadate-maltol complexes formed in aqueous solution. (0.15 M NaC1, 25’ C).  Chemical Shift (6)  45 Assignment  -509 ppm  0 -f ma) anothercis-[V ( 2 ] ) (H [V0 0 2 (ma)) and/or 13 [VO(OH)(ma) ] and/or 2  -496 ppm  (ma) 2 cis-[V0 ]  a) * 2 [VO(OMe)(m ]  -420 ppm -470 ppm  0V-O-VO(ma) 2 [(ma) )  (ma)2]- isomer, with a hydroxy oxygen trans to a V=O. 2 t i.e. a cis-[V0 § i.e. in MeOH V NMR spectra of VO(ma)2 were studied. In a similar experiment the variable pH 51 1 (Figure 3.18.). As In basic solution a peak is observed at 8= -537 ppm corresponding to V the pH is lowered, peaks at 8= -496 and -510 ppm appear. These clearly correspond to 1 is peaks a and b in the previous study. As the acidity of the solution is increased, V shifted to -560 ppm, and disappears below pH 5.5. In acidic solution, peaks a and b separate and are paramagnetically shifted. No spectra are observed below pH 2. These results prove that the same vanadium(V) complexes form in a solution of V0(ma)2, even under relatively acidic conditions. The most likely mechanism for the formation of these 2 with 02 to form [V02(ma)2]-, and vanadate complexes is by the oxidation of VO(ma) V subsequently vanadate oligomers, as implicated in Section 3.1. Comparison of the 51 V0 (Figures [ma) 4 NH. ( ] NMR spectra obtained in the variable pH studies of VO(ma)2 and 2 3.17. and 3.18.) reveals that the spectra are nearly coincident at identical pH values. The  2 1 vanadate oligomer is somewhat more favoured in the case of V0(ma) formation of the V however.  56  Iv’ I  ab  :*t  10.0  S .0  7•0 -  5.5  45  3•5  2•0  .Li1Il I’  -400 Figure 3.18.  I.II..lIiII1.nhpI  -440  I4IIIiIIl uIj!lIrj.I 41141141  I’14t’’I’J’’4’.,.II’’’’l..n.j  -480 -520 -560  -600 8 (ppm)  VO(ma) (10 mM, 0.15 M NaCI, 25 C). Variable pH V NMR specta of 2 57  2 V NMR spectra obtained in the corresponding variable pH study of VO(ka) The 5  are very similar to those observed for VO(ma)2 (Figure 3.19.). In this case however signal intensity is appreciably less, as the complex is less water soluble. A broad peak centred around 8= -495 ppm (&) and a shoulder at 6 = -510 ppm (W) were detected, in addition to 1 peaks. The peaks corresponding to vanadate-kojate complexes occur in a similar V position to those observed in the mahol systems. This is not entirely unexpected, as the immediate 51-vanadium nuclear environment is very similar in both complexes (i.e. an 05,6 donor set). Corresponding ESR spectral parameters for the two ligands are also very V similar. Tracey et al has observed similar chemical shifts in similar positions in the 51 NMR spectra of a number of carboxylate ligands.’ In addition, the signals observed in the experiment are broad, and have a low signal/noise ratio, and hence are more difficult to (ma) NH4[V0 ] 2 and 2 V NMR spectra of VO(ka) assign. Comparison of the variable pH 51 (Figures 3.17. and 3.19.) reveals that the spectra are nearly identical at corresponding pH values. 2 In a preliminary investigation a blood plasma sample obtained from a V0(ma) V NMR spectroscopy. The spectrum (not shown) shows a treated STZ-rat was studied by 51 . Interpretation of this result must be 1 lone resonance at 8= -555 ppm, corresponding to V treated with caution, as the number of scans necessary to obtain a signal (100,000 scans) necessitates a lengthy time period in which the sample was exposed to air. Hence, the signal could be due to vanadyl to vanadate conversion in vivo, or may be due to air oxidation of vanadyl in the sample. Further studies are necessary to help elucidate the mechanism and cause of oxidation.  It is significant, however, that signals due to  vanadium(V)-transferrin, which typically has resonances at 8  =  -529 ppm and -531  16 were not observed. ppm, 1 ’ 115 The various equilibria descibed in this section are summarized in Scheme I.  58  0  pB  10 . 0  8.0  —  7.0  5.5  3.5  I  -400  Figure 3.19.  ‘440  -480 -520 -560  4. 4 .I  2.0 -600 6 (ppm)  2 (10 mM, 0.15 M NaC1, 25 C). Variable pH V NMR spectra of VO(ka) 59  3 CH  oJOHl +  oJ.OH 0 Z  0H H’?  .5O9ppj  7-  3 CH oJo  4 O_LO%11  I  -  \  ,,f =  ° 3 -$O9ppm CH  CH 3 o’L0s%II  7/”  /  [6= 496ppm) w  0 2 .H  OZ0\/0  %..dLy0’ OH2  CH  2 OH .5O9ppm  rCH3  ROH  OJ,OR  oil 3 CH  o  4CH3 \0  3 H  [=42Opj  o0,i0b [ESR Sign  Scheme 1.  Equilibria, V NMR chemical shifts and possible routes of fomadon for vanadium(V) and vanadium(IV) complexes with the maltolate anion. (0.15 M NaCI, 25 C). 60  3.5. Stability Constant Determinations Stability constants are essentially equilibrium constants for metal complex formation, and have long been employed as a quantitative indication of the affinity of a 63 With the relatively recent introduction of powerful ligand for a metal ion in solution. 132 the accuracy of calculated ’ 75 computational methods for the analysis of titration data,  constants have considerably improved. Although there is still some concern regarding the 133 provided a careful and thorough experimental technique validity of all published work, is employed, such as that described herein, reliable and reproducible data can be obtained. Stability constants are frequently quoted in terms of stepwise, or successive, protonation constants for the ligand and either overall or successive formation constants for the metal ligand complex. The relationship between successive (K) and overall  (f)  formation  constants, is that the latter are a cumulative product of the former. Due to the computer programs employed for data reduction (BEST, PKAS (Chapter 2.4.4.)) stability constants quoted in this study are generally given in terms of the overall stability constants, J3. Fundamentally, a stability constant describes an equilibrium quotient involving the activities of reacting species in solution at equilibrium. This is often defined as the ratio of products of the activities of the reaction products, each raised to the appropriate power, to the product of the activities of the reactants. The determination of the activities of these complex ionic species is both complicated and time consuming. However, it has been observed that concentration parallels activities when the ionic strength is regulated by a 56. non-reacting electrolyte in great excess  Although some physical chemists have  occasionally questioned the validity of the “concentration constant” approximation, it has become common practice to measure equilibrium constants as a quotient of concentrations 63 The two most common at constant ionic strength maintained by a supporting electrolyte. techniques employed in stability constant determination are potentiometry and  61  spectrophotometry.  Spectrophotometry is applicable to studies performed at low  concentration or pH, making it useful in cases where ligands are sparingly soluble, and can be used with almost all systems, even proteins and macromolecules. However, the relative ease of use and applicability of data obtained for computation make potentiometry the most 132 attractive, and popular, technique. The ligands maltol and kojic acid contain the 3-hydroxy-4-pyronate moiety as the coordinating group to the metal centre, and have been investigated previously as ligand  precursors for Lewis acids. When deprotonated both of these ligand precursors behave as bidentate Lewis bases, which form relatively stable, neutral, tris-ligand complexes with l). Therefore they are expected to form relatively 1 aluminium(ll group 13 metals such as ’ strong complexes with the hard Lewis acid VO. The positioning of a hydroxy group in an a-position relative to a keto group is found widespread in many naturally occurring 46 compounds, and makes these ligands attractive as potential physiological ligands. Comparison with the analogous 3-hydroxy-2-methyl-pyridinones shows that the pyrones  are stronger acids (pyridinone pKa’s —9- 10; pyrone piCa’s ‘7 8).135 The remarkable ease of hydrolysis of vanadyl chelates, and the potential for oxidation and/or hydrolysis of uncoordinated vanadium(IV), has limited the number of studies on vanadyl chelate formation constants. Indeed, a survey of the literature reveals a veritable dearth of information, and the accuracy of the results obtained in earlier studies must be questioned (due to outmoded computational methodology, poor experimental technique, and lack of consideration for the complexity of the system). Vanadyl chelate complexes hydrolyze as, or more, easily than any of the other metal chelates of the first transition series, and often prefer to combine with hydroxyl ions rather than a second mole of ligand. However, in comparison, vanadyl chelates are often found to be more stable than other transition metal chelates. This may be a consequence of the influence of the highly electronegative oxygen bound to vanadium; the resulting higher electropositive character of  62  the vanadium would make it a stronger Lewis acid, and lead to more stable chelates.  Thermodynamic studies have been scarce, but indicate a high +AS’ value as a major contributor to chelate stability in these comp1exes.’ Vanadyl complexes with catechol as the ligand have been thoroughly investigated, and values of log  =  37 These represent determined. 1 ’ 16.8 and log 1,O,2 = 31.5 81  particularly high vanadyl complex stability constants, and bis(catecholate) species are found to be stable to hydrolysis even at pH >12. The formation of vanadyl chelates with phenolic and/or carboxylate containing ligands yield typical values of log i,oj = —5-12 and log  =  _1220.70,138.140 These figures are representative of the range of determined  stability constants. In a series of papers, Gilard et al 78,1411U have determined stability constants of various vanadyl-amino acid complexes, that fall into the range given. These studies are noteworthy owing to the experimental scrutiny and thoroughness of the authors. Similarly, the work of the Umeâ group on various vanadate systems represents a high .’ However, when comparing reported 46 45 computation 1 standard of data collection and ’ results it is prudent to remember that the type of supporting ionic media, the ionic strength and the solution temperature all contribute to the calculated equilibrium constants; comparison between values determined under different conditions is only indicative, not precise. Smith and Martell have collected a library of stability constant data, for a variety 147 of metal-ligand complexes.  As previously mentioned, it is the formation of  oxovanadium(IV) hydrolysis products in aqueous solution that makes assignments of the complexes formed difficult. The nature of species present in neutral or basic solutions is still somewhat controversial, but it is apparent that hydrolysis products predominate with increasing basicity, and should be included in any equilibrium model. In acidic solution ) predominates. With increasing basicity the species 5 O VO(H “free”, unhydrolyzed, 2 (OH) log (VO) , VO(OH), log f i,.i.o = 5.791, and 2 3 -  ¶  =  -  6.70, form in neutral pH.  log ‘s of the vanadyl hydrolysis species are quoted re-calculated at experimental conditions 33 NaCI, 25 ‘C) by the author according to the method of Bacs and Mesmer.  63  (0.15 M  :OH stoichiometries have been proposed to exist in the pH range 5 11; (i) 2:5 2 Two VO -  predominates, its formation constant is j - 27 (OH) 2 [(VO) ) and (ii) 2:6. Assuming 5 s,o = 3  -  ] predominates, log 1,.3,O = 3 22.06. In the basic regime, the monomeric species [VO(OH) —18 (not included in BEST calculations; this species is certainly important for pH >1 1).26.273  (i) Vanadyl-maltol system The protonation constant of maltol was determined by PKAS analysis: log oj,i  =  8.42 ± 0.01 (0.15 M NaC1, 25 C). Literature values determined for the pICa of maltol are , 25 ‘C).’ 3 47 A representative titration 34 and 8.38 (0.1M KNO 8.62 (0.6 M NaC1, 25 ‘C)’ curve for the vanadyl-maltol system is given in Figure 3.20. The curve corresponding to the vanadyl-maltol titration is shifted from the maltol ligand curve by about 2 units, indicating two moles of ligand consumed in the vanadyl titration. The sharp pH increase at a  =  2 in the vanadyl-maltol curve also suggests formation of a 1:2 complex. A buffer  region appears between pH 8  -  10, which is indicative of the formation of hydrolysis  32 The thermodynamic reversibility of the system was investigated by titration products.” with standard HC1 solution, as depicted in Figure 3.20.; the results of the forward and reverse titrations are in good agreement, with the two curves almost superimposable.  64  1210-  fr  8+  C III;’  4. 2 0 -2-101  2345  a  —0-— Maltol V02+ Maltol Reverse  -G  Figure 3.20.  -  Titration curves obtained from the potentiometric study of the system: V02+ maltol. Conditions: 0.15 M NaC1, 25 C, metal:ligand =1:2. (a =(moles base added-moles acid present)/moles ligand present). -  The vanadyl-pyrone system is quite complicated, and calculations involving refinement of  values of all possible species that may form (i.e. x, y and z quite  general) gave results with no physical meaning. The following strategy was employed. The species log 3 i,o,i and log D1,o.2 were refined initially, with the appropriate vanadyl 33 Other plausible species were hydrolysis constants, ligand protonation constant and pK.t gradually introduced into the refinement, and were discarded if the fit worsened. Combinations of plausible species were refined, until finally no improvement in fit could t pKw refers to the water dissociation constant. pKw =13.765 ,as calculated by Baes for conditions of 33 O.15M NaC1, 25C.  65  be realized. The results of the refinement using the model generating the lowest a, and hence the best fit, are given in Table 3.8. Table 3.8.  Composition and log  f3 values for species in the vanadyl-maltol system, as  refined in the model of best fit as determined by BEST data reduction.& Composition  x, y, z  log  [VO(ma)] ] 2 [VO(ma) [VO(OH)(ma)] -) [VO(OH)(ma ] 2 H) ma)] [VO(O ( 2  1 0 1 1 0 2 1 1 1 1 1 2 1 -2 1 2-2 2  8.49(2) 15.10(4) 1.36 (5) 5.51 (7) -4.57 (4) 5.29 (1)  ) OH)2(ma)2] 2 ([(VO  -  -  ,  a The number in parentheses refers to the standard deviation between successive runs. Log 1 quoted as the mean value. (ii) Vanadyl-kojic acid system Analysis of the kojic acid pKa titration data gave log 3, =7.77 ± 0.01 (0.15 M NaC1, , 25 3 25 C). Literature values determined for the pKa of kojic acid are 7.75 (0.1 M KNO 47 and 7.61 (0.6 M NaC1, 25 .C).1M ‘C)’ A representative titration curve of the vanadyl-kojate system is given in Figure 3.21. The figure shows a sharp inflection in pH at a =2 for the vanadyl-kojic acid titration curve, and a two unit difference between that and the curve for kojic acid alone. This indicates the formation of a 1:2 complex. The thermodynamic reversibility of the vanadyl-kojate titration was checked by reverse titration, with good agreement between forward and reverse curves. Data analysis was performed using the same procedure as for the maltol system, and the results of the refinement are given in Table 3.9.  66  1:.  QI  o  I  I  22  4  6  —0-— Kojic Acid --.0-- V02+ Kojic Acid Reverse -  Figure 3.21. Titration curves obtained from the potentiometric study of the system: -kojic acid. Conditions: 0.15 M NaC1, 25 C, metai:ligand 1:2. (a 2 VO =(moles base added moles acid present)/moles ligand present). -  Table 3.9.  Composition and log f3 values for species refined using BEST for the vanadyl-kojic acid system.  Composition  x, y, z  [VO(ka)] ] 2 [VO(ka) [VO(OH)(ka)] H) ka)][VO(O ( 2  1 1 1 1  67  0 1 0 2 1 1 -2 1  -  log f3  ,,  9.88 (1) 16.37 (5) 1.36 (5) 4.57 (4)  V0 [ma) 4 NH ( 2 (iii) J Giacomelli and co-workers have determined that it is possible to protonate the oxo tes. To this end, 2 V0 was investigated [ma) 4 NH ( ] 9 )dioxovanada group in cisbis(ligand 12 by potentiometric titration, a representative curve shown in Figure 3.22. The protonation constant for the complex was determined by the Henderson-Hassleblach equation: log 1.2  =  ni,..  6.83 (6). This value is comparable to that obtained by Giacomelli et al, log  6.3, for bis(8-quinolinato)dioxovanadate(V). A second protonation constant for their anion was also determined (by acid solubility measurements) as log  ,  =  3.4.  It was not  possible to determine a value for a second protonation constant in this study, either through potentiometric or spectrophotometric means. However, this does not entirely discount the possible formation of such a species at pH <3.  987. —  +  I.;’  5 4-  3-3  —  I  I  -2  .1  —o-—-  0  1  2  3  a Forward  —a-—  Reverse  Figure 3.22. Titration curve obtained from the potentiometric determination of the (first) V0 (0.15 M NaC1, 25 C, under Ar). [ma) 4 NH ( ] protonation constant for 2  68  The refined stability constants in Tables 3.8. and 3.9. offer a plausible hypothesis for several of the species present in each solution. The 1:1 and 1:2 complexes can be confidently defined, and show low standard deviations. Species important above pH >6, where appreciable hydrolysis occurs, show higher standard deviations, and are harder to ascertain. Of considerable importance to the accuracy of the model is the nature of the vanadyl hydrolysis species in neutral to basic conditions. The uncertainty in the value 27 however, refinements attributed to Pz-5,o is certainly a concern for investigators; attempted without this constant gave a very poor fit. It is apparent that hydrolysis, of both free and bound vanadium, occurs in these regions, as proposed by the visible spectroscopy,  ‘v  NMR and ESR studies presented herein. The inclusion of the  f2.-s.o  hydrolysis  constant in vanadyl-ligand stability constant determinations is a practice observed in other 48 laboratories.’ The metal-ligand equilibria for the non-hydrolyzed species in these systems may adequately be described by equations 3.3 and 3.4.  V02+  +  VO(L) + L-  —  —  f3  (3.3)  2 Pi VO(L)  (3.4)  VO(L)  L-.  i 1(VO)L,( ) 2 ][Lj 2 [VO  These pyrone ligands are found to have a reasonably high affinity for vanadyl, evidenced by the determined log 3 values. In comparison, the results show that these pyrones have similar stabilities to those found for other phenolic and/or carboxylate containing ligands, i.e. log  =  5.59 for (malonato)oxovanadium(IV), log Di,o, = 9.48  47 A comparison between the two ligands reveals that for bis(malonato)oxovanadium( V).’  69  the acidity of maltol (pKa = 8.42) is somewhat lower than that of kojic acid (pICa = 7.77). This difference may be attributed to the inductive effect of the methyl group in maltol. Vanadyl binds most effectively to hard electronegative atoms, i.e. F, Cl, 0 and N. Fluoro and oxygen donor containing compounds are especially stable. In general, complexes with oxygen as the donor atoms are found to follow the expanded Irving49 The presence of one or two nitrogen donors places VO lower in the Williams series.’ 27 When the stability constants of these vanadyl-y-pyrones are compared to the series. stabilities of other divalent metal-y-pyrone complexes, the resulting series is indeed found . in addition, it is also noted 2 2 >Co >Mn 2 >2+ >Ni to follow the order: V02+ >C’j that the magnitude of the stability constants previously determined for divalent metal kojato complexes are in general larger than those for the analogous maltolato complexes, 453 Presumably, it is the effect that observations that are also noted in this study.t 136,147,150  the hydroxyl group in kojic acid has upon the stability of the metal-ligand complexes, in comparison to the methyl group of maltol, that is responsible for the slight difference in the stability constants for the two ligands. Speciation diagrams are an effective indicator of the nature of the complexes formed in solution, although they are by no means a substitute for spectroscopic investigation. Information garnered from such diagrams must be treated purely as a gauge for solution speciation. The quality of information is directly proportional to the quality of the stability constant data from which they are constructed. In this example, the poorly understood vanadyl hydrolysis, and the nature of minor hydrolysis species, questions the complete validity of the proposed diagrams in the basic regime. Speciation diagrams for complex formation in the vanadyl-maltol system are shown calculated at mM vanadyl concentration for a 1:2 metal-ligand system from stability constants obtained in this study (Figure 3.231). t However, it must be noted that several of these studies were performed before the introduction of new computational methods, or contain results obtained under less than rigorous experimental conditions. Hence, these trends are indicative at best.  70  In highly acidic pH, free V02+ is observed. The log pH 2  -  i,o,i  complex dominates in the region  6; the log Di,o,z complex begins to form by pH 4, and has a maximum  “concentration” at pH —7. This has obvious ramifications for the chemist interested in physiological modeling. In basic solution hydrolyzed complexes, and vanadyl hydrolysis products, dominate. The diagrams succinctly demonstrate the effects of concentration on complex formation. At low concentrations in neutral solution, hydrolysis products are able to effectively compete with the ligand for vanadyl. The result is that very little of the 1:2 product is obtained at this concentration level. In the basic regime, hydrolysis products dominate, as proposed from results presented herein. In the corresponding vanadyl-kojic acid systems, Figure 3.23.ii, similar observations  are made. However, in this case the higher overall stability of the kojato complexes is demonstrated by the dominance of the log ( 1,o,i and log (31,0,2 curves. The difference in 3 stability of the maltol and kojate complexes is manifested in a less dramatic concentration dependence for complex formation in the latter case. An obvious flaw in the computer software currently available for determining stability constants is the lack of a provision to include an electron as a fourth variable in the data computation and curve fitting procedures. Until this is possible, and a complete understanding of vanadyl hydrolysis is reached, the absolute dependibilty of data obtained for species in the neutral to basic regime will remain questionable.  71  100-  --  a  -Iog[H+] (i) 100C  St..  50  S  e  25-  /  ‘. S.  ‘  S  1  2  3  5 4 .tog[H+]  _,  —  -.  0-  ——  ,  6  7  8  (ii) Figure 3.23.  -Ligand (1:2, 10-2 mM VO). (I) Maltol 2 Speciation diagrams for VO (ii) Kojic Acid. Conditions: 1=0.15 M NaCI, 25 C, under Ar. . (OH) 2 [(VO) ] (OH)], e. 5 2 ], d. [VO(L) 2 a. VO, b. [VO(L)], c. [VO(L)  72  Chapter 4 Suggestions for Future Work The results presented herein clearly show that facile processes occur for the (ma) in aqueous solution. This is 2 [V0 interconversion between VO(ma) 2 and anionic 1 (ma) and V 2 NH.jV0 evidenced by ESR signals observed in solutions of ESR silent ] . Dissolved oxygen is the presumed 2 NMR signals in solutions of NMR silent VO(ma) , as shown by a retardation of the (ma) 2 [V0 2 to ] oxidant in the conversion of VO(ma) process in solutions maintained under an inert atmosphere. An investigation of the mechanism and kinetics of this oxidation are appropriate subjects for future study This oxidation process has obvious implications in vivo as dissolved oxygen is prevalent physiologically, particularly in plasma. Therefore it is assumed that a similar process will occur in the mammalian system. The similarities in the chemistry of the (bismaltolato)oxovanadium(IV) and (biskojato)oxovanadium(IV) complexes is demonstrated in their similar behaviour and visible, ESR and ‘V NMR spectral properties. Recent studies in collaboration with Dr. 3. H. McNeill of the Faculty of 2 display 2 and VO(ka) Pharmaceutical Sciences at UBC have concluded that while VO(ma) favourable insulin mimicking properties when administered to STZ-treated diabetic rats, V0 does not, and can indeed be severely toxic. This [ma) 4 NH ( 2 the vanadate analogue ] appears surprising in view of the similar, pH dependent, chemistry reported herein. As the pH of the mammalian system changes dramatically from stomach to intestines, it is apparent that in vivo investigations are necessary to determine the physiological site of absorption, and the oxidation state during absorption. ESR, and less likely 51v NMR, as demonstrated in this study, are suitable probes for studying the redox behaviour, ligand environment, and oxidation state of the compounds in the blood and organs of treated 73  animals, and could be effectively employed in this case. As both ESR and NMR can be 154 they used to monitor the diffusion of molecules acmss membranes and red blood cells, could be the basis for model studies into absorption of vanadyl-complex molecules from the gastrointestinal tract into the blood.  74  References (1)  Jandhyala, B. S.; Hom,G. 3. Life Sciences 1983,33, 13254340.  (2)  Chasteen, N. D.; Lord, E. M.; Thompson, H. 3. In Frontiers in Bioinorganic  Chemistry; A. Xavier, Ed.; VCH Publishers: 1986;133-141. (3)  Smith, M. 3. Experienda 1989,45,452.  (4)  Byrne, A. R.; Kosta, L. Sci. Total Environ. 1978,10, 17-30.  (5)  Bayer, W. F.; Kneifel, H. Z. Natwforsch. 1972,27,207.  (6)  Vilter, H. Phytochemisrry 1984,23, 1387.  (7)  Cornman, C. R.; Kampf, 3.; Soo Lah, M.; Pecoraro, V. L. bzorg. Chem. 1992,31, 2035-2043  (8)  Cantley, L. C.; Josephson, L.; Warner, R.; Yanaqisawa, M.; Lechene, C.; Guidotti,  G. JBiol. Chem. 1977,252,7421. (9)  Shechter, Y. In Vanadium in Biological Systems; N. D. Chasteen, Ed.; Kiuwer Academic: Dordrecht, 1990;100-128.  (10)  Biggs, W. R.; Swinehart, 3. H. Metal Ions Biol. Syst. 1976,6, 142-196.  (11)  Kustin, K.; Macara, I. G. Comments Inorg. Chem. 1982,2, 1-22.  (12)  Chasteen, N. D. Struc:. Bond. 1983,53, 105-138.  (13)  Boyd, D. W.; Kustin, K. Adv. Inorg. Biochem. 1984,6,311-365.  (14)  Nechay, B. R.; Nanninga, L. B.; Nechay, P. S. E.; Post, R. L; Grantham, J. 3.;  Macara, L 0.; Kubena, L. F.; Phillips, T. D.; Nielsen, F. H. Fed. Proc. 1986,45, 123-132. (15)  Wever, R.; Kustin, K.Adv. lnorg. Chem. 1990,35, 81-115.  (16)  Rehder, D. Angew. Chem. mt. Ed. Engl. 1991,30, 148-167.  (17)  Butler, A.; Canano, C. 3. Coord. Chem. Rev. 1991,109,61-105.  (18)  Tracey, A. S.; Gresser, M. 3. Proc. Nat!. Acad. Sci. USA. 1986,83, 609-613.  75  (19)  Gresser, M. 3.; Tracey, A. S.; Stankiewicz, P.3. Adv. Protein Phosphatases 1987,4, 35.  (20)  Tracey, A. S.; Gresser, M. 3.; Liu, S. I. Am. Chem. Soc. 1988,110,5869-5874.  (21)  Tracey, A. S.; Galeffi, B. Inorg. Chem. 1989,28, 1726-1734.  (22)  Gresser, M. 3.; Tracey, A. S. In Vanadiumin Biological Systems; N. D. Chasteen, Ed.; Kiuwer Academic: Dordrecht, 1990;63-80.  (23)  Liochev, S.; Ivancheva, E.; Fridovich, L Arch. Biochem. Biophys. 1989,269, 188193.  (24)  Secco, F. Inorg. Chem. 1980,19,2722-2725.  (25)  A1-Bayati, M.; Raabe, 0. G.; Gin, S. N.; Knaak, 3. B. .1. Am. Coil. Tox. 1991,10, 233-241.  (26)  Henry, R. P.; Mitchell, P. C. H.; Prue, 3. E. I. Chem. Soc. Dalton Trans. 1973, 1156.  (27)  Vilas Boas, L. F.; Costa Pessoa, 3. In Comprehensive Coordination Chemistry; S. 0. Wilkinson, R. D. Gillard and 3. A. McCleverty, Ed.; Pergamon Press: Oxford, 1987; Vol. 3;453-584.  (28)  Chasteen, N. D. Biol. Magn. Res. 1981,3,53-119.  (29)  Hon, P. K.; Belford, R. L.; Pfluger, C. E. J. Chem. Phys. 1965,58, 1195.  (30)  Harris, W. R.; Carrano, C. J. I. Inorg. Biochem. 1984,22,201-218.  (31)  Harris, W. R.; Friedman, S. B.; Silberman, D. J. Jnorg. Biochem. 1984,20, 157169.  (32)  Sabbioni, E.; Marafante, E. Bioinorg. Chem. 1978,9,389-407.  (33)  Baes, C. F.; Mesmer, R. E. The Hydrolysis of Cwions; John Wiley & Sons: New York, 1986.  (34)  Atkinson, M. A.; Maclaren, N. K. Sci. Am. 1990,263,62-71.  (35)  Shechter, Y. Diabetes 1990,39, 1-5.  (36)  Shechter, Y.; Karlish, S. 3. D. Nature 1980,284,556-558.  76  (37)  Degani, A.; Shechter, Y.; Gochin, M.; Karlish, S. 3. D. Biochemistry 1981,10,  2433-2438. (38)  Heyliger, C. E.; Tahiliani, A. 0.; McNeil, 3. H. Science 1985,227, 1474-1477.  (39)  Ramanadham, S.; Mongold, 3.3.; Brownsey, R. W.; eros, 0.; McNeill, 3.11. American Journal ofPhysiology 1989,257, H904-H91 1.  (40)  McNeil, 3. H.; Yeun, V.0.; Hoveyda, H. R.; Orvig, C. J. Med. Chem. 1992,35, 1489-1491.  (41)  Dai, S.; Yeun, V. 0.; Orvig, C.; McNeill, 3. H. Pharmacol. Comm. 1993,3,311321.  (42)  Yuen, V. G.; Orvig, C.; McNeil, 3. H. Can. I. Physiol. Pharmacol. 1993,71,263269.  (43)  Yuen, V. 0.; Orvig, C.; McNeill, 3. H. Can. I. Physiol. Pharmacol. 1993,71,270276.  (44)  Paulson, D. J.; Kopp, S. 3.; Tow, 3. P.; Peace, D. G. I. Pharmacol. Exp. Ther. 1987, 240,529-534.  (45)  Gelmini, L.; Glover, N.; HerringF.G.H.;Rettig,S.J.;Orvig, C. Manuscript in Preparation 1993.  (46)  Beelik, A.; Purves, C. B. Can. I. Chem. 1955,33, 1361-1374.  (47)  Zhou, Y. M.Sc. Thesis, University of British Columbia, 1993.  (48)  Kustin, K.; Pizer, P. Inorg. Chem. 1970,9, 1536.  (49)  Pettit, L. D.; Swash, 3. L. M. I. Chem. Soc. Dalton Trans. 1978,286.  (50)  Che, T. M.; Kustin, K. Inorg. Chem. 1980,19,2275.  (51)  Holland, R. M.; Tapscott, R. E. I. Coord. Chem. 1981, 17.  (52)  Fisher, D. C.; Barclay-Peet, S. 3.; Balfe, C. A.; Raymond, K. N. lnorg. Chem. 1989, 28,4399.  (53)  Branca, M.; Micera, G.; Dessi, A.; Sanna, D. I. Inorg. Biochem. 1992,45, 169-177.  77  (54)  Buglyó, P.; Kiss, T. I. Coord. Chem. 1991,22,259.  (55)  Rossotti, F. 3. C.; Rossotti, H. I. Chem. Ed. 1965,42,375-378.  (56)  Rossotti, H. The Study ofIonic Equilibria; Longman: New York, 1978.  (57)  Rehder, D. Magn. Reson. Rev. 1984,9, 125-233.  (58)  White, L. K.; Belford, R. L I. Am. Chem. Soc. 1976,98,4428  (59)  Gagne, R. R.; Koval, C. A.; Lisensky, G. C. Inorg. Chem. 1980,19,2854-2855.  (60)  Cacheris, W. P.;Franczyk, T. S. In Raymond Research Group: University of California, Berkeley, 1990.  (61)  May, P. M.; Williams, D. R.Talanza 1982,29,249.  (62)  Bates, R. 0. Determination ofpH; John Wiley: New York, 1973.  (63)  Martell, A. E.; Motekaitis, R. Determination and Use ofStability Constajus; VCH Publishers: New York, 1988.  (64)  Motekaitis, R. 3.; Martell, A. E. Can. I. Chem. 1982,60,2403-2409.  (65)  Motekaitis, R. 3.; Martell, A. E. Can. I. Chem. 1982,60, 168-173.  (66)  Ortolano, T. R.; Selbin, 3.; McGlynn, S. P. J. Chem. Phys. 1964,41,262-269.  (67)  Selbin, 3. Chem. Rev. 1965,65, 153-175.  (68)  Selbin, 3. Coord. Chem. Rev. 1966, 1, 293-314.  (69)  Bafihausen, C. 3.; Gray, H. B. Inorg. Chem. 1962, 1, 111-122.  (70)  Tomiyasu, H.; Gordon, G. I. Coord. Chem. 1973,3,47-56.  (71)  Lever, A. B. P. Inorganic Electronic Spectroscopy Elsevier: Amsterdam,  1986, p. 260 (72)  Chasteen, N. D.; Grday, 3. K.; Holloway, C. E. Inorg. Chem. 1986,25,2754-2760.  (73)  Branca, M.; Micera, 0.; Dessi, A.; Sanna, D. I. lnorg. Biochem. 1992,45, 169-177.  (74)  Ferrer, E. G.; Williams, P. A. M.; Baran, E. J. I. Inorg. Biochem. 1993,50,253262.  78  (75)  Meloun. M.; Havel, 3.; Hogfeldt, E. Computation ofSolution Equilibria: A Guide to Methods in Potentiometiy, Extraction, and Spectrophotometry; Ellis Horwood Ltd:  Chichester, 1988. (76)  Geraldes, C. F. 0.; Casim, M. M. C. A. I. Inorg. Biochem. 1989,35,79-93.  (77)  Williams, P. A. M.; Baran, E. 3. Biol. Trace Eleme,u Res. 1993,36, 143-150.  (78)  Costa Pessoa, 3.; Vilas Boas, L F.; Gillard, R. D.; Lancashire, R. 3. Polyhedron 1988,7, 1245-1262.  (79)  Wuthrich, K. Helv. Chim. Acca. 1965,48, 1012-1017.  (80)  Cooper, S. R.; Koh, Y. B., Raymond, KJ. Am. Chem. Soc. 1982,104,5092-5102.  (81)  Jezowska-Bojczuk, M.; Kozlowski, H.; Zubor, A.; Kiss, T.; Branca, M.; Micera, 0.; Dessi, A. I. Chem. Soc. Dalton Trans. 1990,2903-2907.  (82)  Cornman, C. R.; Colpas, G. 3.; Hoeschele, 3. D.; Kampf, 3., Pecararo, V.LJ. Am. Chem. Soc. 1993, 114, 9925-9933.  (83)  Saito, K. In Coordination Chemistry; Volume 20; D. Baherjea, Ed. 1980.  (84)  Porte, A. L.; Stewart, C. P.1. Chem. Soc. Dalton Trans. 1972, 1661-1666.  (85)  Ortolano, T.R., Selbin, 3., McGlynn, S.P. 1. Chem. Phys. 1964,41(1), 262-268.  (86)  Baithausen, C. 3.; Gray, H. B. Inorg. Chem. 1962,1, 111-122.  (87)  Selbin, 3. Can. 1. Chem. 1966,293-315.  (88)  Borovik, A. S., Dewey, T. M., Raymond, K. N.Inorg. Chem. 1993,32,413-421.  (89)  Bonadies, 3. A., Carrano, C. JJ. Am. Chem. Soc. 1986,108,4088-4095.  (90)  Israel, Y.; Meites, L. In Standard Potentials in Aqueous Solution; A. 3. Bard, R. Parsons and 3. Jordan, Ed.; Dekker: New York, 1985.  (91)  Bonadies, 3. A.; Butler, W. M.; Pecoraro, V. L.; Carrano, C. 3. Inorg. Chem. 1987, 26, 1218-1222.  (92)  Collison, D.; Powell, A. K.; Eardley, D. R.; Mabbs, F. E.; Turner, S. S. Inorg. Chem. 1993,32,664-671.  79  (93)  Bard, A. 3.; Faulkener, L. R. Electrochemical Methods  -  Fundamentals and  Applications; John Wiley & Sons: New York, 1980. (94)  Wertz, 3. E.; Bolton, 3. R. Electron Spin Resonance  -  Elementary Theory and  Practical Applications; Chapman and Hall: New York, 1986. (95)  Boucher, L. J.; Tynan, E. C.; Yen, T. F. In Electron Spin Resonance ofMetal Complexes; T. F. Yen, Ed.; Plenum: New York, 1969; 111-130.  (96)  Kivelson, D.; Lee, S.-K. 1. Chem. Phys. 1964,41, 1896-1903.  (97)  Wilson, R.; Kivelson, D. .1. Chem. Phys. 1965,44, 154-166.  (98)  Casey, 3. D.; Chasteen, N. D. I. Inorg. Biochem. 1980,13, 111-126.  (99)  Sakurai, H.; Hiram, 3.; Michibata, H. Biochem. Biophys. Res. Commun. 1987,149, 411-416.  (100) Sakurai, H.; Tsuchiya, K.; Nukatsuka, M.; Kawada, 3.; Ishikawa, S.; Yoshida, H.;  Komatsu, M. J. Clin. Biochem. Nutr. 1990,8, 193-200. (101) Frank, P.; Carlson, R. M. K.; Hodgson, K. 0. Inorg. Chem. 1986,25,470-478. (102) Eaton, S. S.; Eaton, G. R. In Vanadium in Biological Systems; N. D. Chasteen, Ed.; Kiuwer Academic: Dordrecht, 1990;199-222. (103) Mustafi, D.; Telser, 3.; Makinen, M. W. I. Am. Chem. Soc. 1992, 114,6219-6226. (104) Ehde, P. M.; Andersson, I.; Petersson, L. Ada. Chem. Scand. A 1986,40,489-499. (105) Sakurai, H.; Shimomura, S.; Ishizu, K. Inog. Chim. Acta 1981,55, L65-L69. (106) Yordanov, N. D.; Zdravkova, M. Polyhedron 1993,12,635-639. (107) Kuska, H. A.; Rogers, M. T. Inorg. Chem. 1966,5,3113. (108) Beddoes, R. L.; Collison, D.; Mabbs, F. E.; Passand, M. A. Polyhedron 1990,9, 2483-2489. (109) Chasteen, N. D.; Lord, E. M.; Thompson, H. 3.; Grady, 3. K. Biochim. Biophys. Acta 1986,884,84-92.  80  (110) Vanadiwn in Biological Systems  -  Physiology and Biochemistry; Chasteen, N. D.,  Ed.; Kiuwer Academic: Dordecht, 1990. (111) Harris, R. K. Nuclear Magnetic Resonance : A Physicochemical View Pitman: London, 1983. (112) Rehder, D. In Vanadiunz in Biological Systems; N. D. Chasteen, Ed.; Kiuwer Academic: Dordrecht, 1990;173-198. (113) Rebder, D.; Weidemann, C.; Duch, A.; Priebsch, W. Inorg. Chem. 1988,27,584587. (114) O’Donnell, S.; Pope, M. T. I. Chem. Soc. Dalton Trans. 1976,2290-2297. (115) Butler, A.; Danzitz, M. 3.; Eckert, H. J. Am. Chem. Soc. 1987,109, 1864-1865. (116) Butler, A.; Eckert, H. I. Am. Chem. Soc. 1989, 111, 2802-2809. (117) Crans, D. C.; Shin, P. K. Inorg. Chem. 1988,27, 1791-1806. (118) Crans, D. C.; Bunch, R. L.; Theisen, L. A. I. Am. Chem. Soc. 1989, 111, 75977607.  (119) Crans, D. C.; Wong, C.-H.; Daniels, L.; Pederson, R. L.; Durrwachter, 3. R.; Drueckhammer, D. 0. J. Org. Chem. 1989,54,70-77. (120) Crans, D. C.; Rithner, C. D.; Theisen, L. A. I. Am. Chem. Soc. 1990,112,29012908. (121) Rehder, D. Inorg. Chem. 1988,27,4312-4316. (122) Rehder, D.; Hoist, H.; Quaas, R.; Hinrichs, W.; Hahn, U.; Saeñger, W. J. Inorg. Biochem. 1989, 37, 141-150. (123) Tracey, A. S.; Jaswai, 3. S.; Gresser, M.; Rehder, D. Inorg. Chem. 1990,29,42834288. (124) Tracey, A. S.; Gresser, M. 3. Inorg. Chem. 1990,29,2267-2271. (125) Tracey, A. S.; Gresser, M. 3.; Liu, S. Biochemistry 1992,31,2677-2685. (126) Habayeb, M. A.; Hileman, 0. E. Can. I. Chem. 1980,58,2255-2261.  81  (127) Pettersson, L.; Hedman, B.; Andersson, 1.; Ingri, N. ChemicaScripta. 1983,22, 254-264. (128) Blair, A. 3.; Pantony, D. A.; Minkoff, G. I. I. bzorg. Nuc. Chem. 1958,5,316-331. (129) Giacomelli, A.; Floriani, C.; Duane, A.; Chiesi-Villa, A. Inorg. Chem. 1982,21, 3310-3316. (130) Finnegan, M. M.; Lutz, T. G.; Nelson, W. 0.; Smith, A.; Orvig, C. Jnorg. Chem. 1987,26,2171-2176. (131) Nelson, W. 0.; Karpishin, T. B.; Rettig, S. 3.; Orvig, C. Inorg. Chem. 1988,27, 1045-1051. (132) Zékány, L.; Nagyp1, I. Computational Methodsfor the Determination of Stability Constants; Plenum Press: New York, 1985. (133) Rossotti, H. S. Talanw 1974,21,809-829. (134) Hedlund, T.; Ohman, L-O. Ac:. Chem. A 1988,42,702-709. (135) Clevette, D. J.; Orvig, C. Polyhedron 1990, 9, 151-161. (136) Gerard, C.; Hugel, R. P.1. Chem. Research 1978,404-405,4875-4889. (137) Lal, K.; Agarwal, R. P. Bull. Chem. Soc. Japan 1967,40, 1148-1152. (138) Nechay, B. R.; Nanninga, L. B.; Nechay, P. S. E. Arch. Biochem. Biophys. 1986, 251, 128-138. (139) Lorenzotti, A.; Leonesi, D.; Cingolani, A.; Bernardo, P. D. 1. Inorg. Nuc. Chem. 1981,43,737-738. (140) Mont, G. E.; Martell, A. E. 1. Am. Chem. Soc. 1966,88, 1387-1393. (141) Costa Pessoa, 3.; Vilas Boas, L F.; Gillard, R. D. Polyhedron 1989,8, 1173-1199. (142) Costa Pessoa, 3.; Marques, R. L.; Vilas Boas, L. F.; Gillard, R. D. Polyhedron 1990, 9, 81-98. 2125. (143) Costa Pessoa, 3.; Vilas Boas, L. F.; Gillard, R. D. Polyhedron 1990,9,2101-  82  (144) Costa Pessoa, 3.; Antunes, 3. L.; Vilas Boas, L. F.; Gilard, R. D. Polyhedron 1992,  11, 1149-1461. (145) Pettersson, L; Hedman, B.; Nenner, A. M.; Andersson, I. Acm. Chem. Scand. A  1985,39,499-506. (146) Ehde, P. M.; Peuersson, L.; Glaser, 3. Acm. Chem. Scand. A1991, 45,998-1005. (147) Smith, R. M.; Martell, A. E. Critical Stability Constants; Plenum Press: New York, 1974-1989. (148) Costa Pessoa, 3., Personal Communication. (149) Huheey, 3. E. Inorganic Chemistry; Third edition; Harper & Row: New York, 1983, 276-277. (150) Bryant, B. E.; Fernelius, W. C. I. Am. Chem. Soc. 1954,76,5351-5352. (151) Okac, A.; Kolarik, Z. Collect. Czechoslov. Chem. Commun. 1959,266-272. (152) Gerard, C. Bull. Chem. Soc. France 1979,1,451-456. (153) Gerard, C.; Hugel, R. P. C. R. Acad. Sc. Paris 1982,295, 175-177.  (154) Herring, G. Unpublished Results  83  


Citation Scheme:


Citations by CSL (citeproc-js)

Usage Statistics



Customize your widget with the following options, then copy and paste the code below into the HTML of your page to embed this item in your website.
                            <div id="ubcOpenCollectionsWidgetDisplay">
                            <script id="ubcOpenCollectionsWidget"
                            async >
IIIF logo Our image viewer uses the IIIF 2.0 standard. To load this item in other compatible viewers, use this url:


Related Items