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Aqueous coordination chemistry of aluminum and gallium with 3-hydroxy-4-pyridinone ligands Nelson, William Otto 1988

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AQUEOUS COORDINATION CHEMISTRY OF ALUMINUM AND GALL WITH 3-HYDROXY-4-PYRIDINONE LIGANDS by WILLIAM OTTO NELSON B.Sc, University of Wisconsin-Superior, 1984 A THESIS SUBMITTED IN PARTIAL FULFILLMENT OF THE REQUIREMENTS FOR THE DEGREE OF DOCTOR OF PHILOSOPHY in THE FACULTY OF GRADUATE STUDIES (Department of Chemistry) We accept this thesis as conforming to the required standard THE UNIVERSITY OF BRITISH COLUMBIA December 1988 © William Otto Nelson, 1988  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.  Department The University of British Columbia Vancouver, Canada  DE-6 (2/88)  i A B S T R A C T Aluminum(III) and gallium(III) complexes with the following synthesized: 3-hydroxy-2-methyl-4(lH)-pyridinone (Hmpp), the 1-methyl (Hdp (Hmepp), and 1-hexyl (Hmhpp) derivatives of Hmpp, and pyridinone)]-a-aminopropionic acid (mimosine). The 3-hydroxy-4-pyridinone employed in this study (except mimosine) were prepared by the heterocycle, 3-hydroxy-2-methyl-4-pyrone, to the corresponding nitrogen he reaction with a primary arnine. These bidentate ligands contain an and their conjugate bases form neutral complexes with trivalent metal metal complexes were fully characterized by mass spectrometry and and ultraviolet spectroscopy. The structures of several ligands and metal complexes were diffraction. Hmpp, Hdpp, and Hmepp crystallize as centrosym hydrogen bonded dimeric units. The geometric isomers of crystallize as the dodecahydrate in which the water molecules are rings similar in structure to that of ice Ih. The oxygen atoms hydrogen bonded to bridging waters so that the water rings a interconnected in a three-dimensional array. An analogous water netw structures of Al- and Ga(mepp3). The proton NMR spectra in and indicate the met fluxional above -30 °C. Variable-temperature proton NMR experime exchange process as to geometric isomerization. Liga experiments using proton NMR indicated the isomerization follows an than intramolecular pathway in Variable-pHAl NMR experiments tris-ligand aluminum complexes to be resistant to hydrolysis from pH facial  CD3OD  facial  D2O  meridional  27  CD3OD.  ii constants of the metal-ligand complexes were determined by potentiome this study indicates the gallium complexes have a similar pH regi The overall formation constants for the tris-ligand aluminum and galli all greater than 10, indicating that these ligands could compete for in blood plasma models. Water solubilities and octanol/water partition ligands and metal complexes were measured and they indicate th complexes foG ra animal biodistribution experiments. The results of t study show that under conditions of ligand excessGa is redirecte however, thGea-ligand complexes do not localize in any organs. greatly enhance the removal of the radionuclidefromthe body. 03  76  76  76  iv TABLE OF CONTENTS  Abstract  ii  Table of Contents  iv  List  viii  of  Tables  List of Figures  x  List of Abbreviations  xiii  Acknowledgements  xvi  Chapter I  1  General Introduction  Chapter n Synthesis and Characterization  10  A. 3-Hydroxy-4-Pyridinone Ligands  10  2.1 Introduction  10  2.2 Materials and Methods  15 Method A  2.2.1  Preparation of B zmpp  15  2.2.2  Preparation of Hmpp  16  2.2.3  Preparation of Hdpp  16  2.2.4 2.2.5  Preparation of Hmhpp Preparation of H2exn  16 16  Method B  2.3  2.2.6  Preparation of Hmpp  17  2.2.7  Preparation of Hdpp  17  2.2.8  Preparation of Hmepp  18  Discussion of the Synthetic Procedure  2.4 Characterization of the 3-Hydroxy-4-Pyridinone Ligands 2.4.1 2.4.2  Elemental Analysis Infrared Spectroscopy  18 22 23 24  V  2.4.3 Proton NMR Spectroscopy 29 2.4.4 Electron Impact Mass Spectrometry 33 B. Tris(3-Hydroxy-4-Pyridinonato) Metal Complexes 36 2.5 Introduction 36 2.6 Material and Methods 41 2.6.1 Preparation of Al(mpp3) 42 2.6.2 Preparation of Ga(mpp3) 42 2.6.3 Preparation of Al(dpp3) 42 2.6.4 Preparation of Ga(dpp3) 42 2.6.5 Preparation of Al(mhpp>3 42 2.6.6 Preparation of Ga(mhpp3) 43 2.6.7 Preparation of Al(mepp3) 43 2.6.8 Preparation of Ga(mepp3) 43 2.6.9 Preparation of Al(mimo3) 43 2.6.10 Preparation of Ga(mimo3)»H20 44 2 . 7 D i s c u s s i o n o f t h e S y n t h e t i c P r o c e d u r e 2.8CompClheax raecsterization of the Tris(3-Hydroxy-4-Pyri4 d8 inonato) Metal 2.8.1 Elemental Analysis 49 2.8.2 Infrared Spectroscopy 50 2.8.3 Mass Spectrometry 57 Chapter HI Solid State Studies 58 A. 3-Hydroxy-4-Pyridinone Crystal Structures 58 3.1 Introduction 58 3.2 Results and Discussion 62 B. M(dpp3) Crystal Structures 68 3.3 Introduction 68  vi 3.4 Results and Discussion 70 3.4.1 The/ac-M(dpp)3 Unit 70 3.4.2 The Hydrogen Bonded Water Network C. M(mepp)3 Crystal Structures 79 3.5 Introduction 79 3.6 Results and Discussion 81 Chapter IV NMR Studies 93 A. Aluminum-27 NMR Spectroscopy 93 4.1 Introduction 93 4.2 Materials and Methods 96 4.3 Results and Discussion 97 B. Variable-Temperature Proton NMR Spectroscopy 109 4 . 4 Introduction 109 4.5 Materials and Methods 112 4.6 Results and Discussion 112 Chapter V Solution Studies 129 A. Aqueous Solubility 129 5.1 Introduction 129 5.2 Materials and Methods 130 5.2.1 Determination of Molar Absorptivity 131 5.2.2 Determination of Aqueous Solubility 13 5.3 Results and Discussion 133 5.3.1 Ultraviolet Spectroscopy 133 5.3.2 Aqueous Solubility at 25 and 37 °C B. Octanol/Water Partition Coefficients 140 5.4 Introduction 140 5.5 Materials and Methods 142  vii 5.6 Results and Discussion 144 C. Potentiometric Equilibrium Measurements 150 5.7 Introduction 150 5.8 Results and Discussion 151 Chapter VI Gallium-67 Biodistribution Study 161 6.1 Introduction 161 6 . 2 Materials and Methods 162 6 . 3 Results and Discussion 163 Chapter VII Conclusion and Suggestions for Future Work References 171 Appendix 180 Procedure A.l. Potentiometric Equilibrium Measurements 185  vii LIST OF TABLES Table 1.1. Effective ionic radii and transferrin bindingPage constants Table 2.1. Results of thefreeligand elemental analyses Table 2.2. Infrared bands for Hmpp and substituted analogues Table 2.3. Selected infrared absorption bands in the free ligands Table 2.4. Proton NMR chemical shift data for thefreeligands Table 2.5. Mass spectral data for the free ligands Table 2.6. Comparison of T| values and C/TR for selected biva Table 2.7. Results of elemental analyses of the metal complexes Table 2.8. Assignment of VM-0 the Al and Ga complexes Table 2.9. Characteristic IR absorptions of the Al and Ga com Table 2.10. *H NMR data for mimosine and M(mimo)3 complexes Table 2.11. *H NMR data for the tris-ligand metal complexes Table 2.12. Data from FAB-MS spectra of the tris-ligand metal Table 3.1. Bond lengths and angles for thefreeligands Table 3.2. Comparison of bond lengths between 2- and 4-pyridi Table 3.3. Free ligand H-bond parameters and IR stretchingfreq Table 3.4. Bonding parameters for the M(dpp)3»12H20 complexes Table 3.5. Intra-annular torsion angles of the chelate rings in M Table 3.6. Bond angles for the metal-ligand interactions in M(dpp)3 Table 3.7. Metal ionic radii and M-O bond lengths in M(dpp)3 Table 3.8. Hydrogen bond distances and angles for M(dpp)3»12H20 Table 3.9. Bond parameters for M(mepp)3»12H20 complexes Table 3.10. Unit cell dimensions for M(dpp)3 and M(mepp)3 com Table 4.1. NMR properties of several quadrupolar nuclei Table 4.2.A1 NMR data for the tris(3-hydroxy-4-pyridinonato) Al com m  72  ix Table 4.3. Kinetic data for rris(3-hyclroxy-4-pyridinonato) metal complexe Table 5.1. ^ and E for the B-band in the free ligands Table 5.2. and £ for the B-band in the metal-ligand comple Table 5.3. Water solubility of free ligands at 25 °C Table 5.4. Water solubility of the metal-ligand complexes at 25 Table 5.5. Log P values of some common drugs Table 5.6. Log P values for the 3-hydroxy-4-pyrone complexes Table 5.7. Log P values for the 3-hydroxy-4-pyridinone complexes Table 5.8. Log K and log (3n for 3-hydroxy-4-pyridinone metal Table A. 1. Molecular weights for ligands and metal-ligand comple Table A.2. Crystallographic data for Hmpp, Hdpp, and Hmepp Table A.3. Crystallographic data for M(dpp)3 complexes Table A.4. Crystallographic data for M(mepp)3 complexes Table A.5. Hydrogen bond distances and angles for In(dpp)3«12H20 Table A.6. Hydrogen bond distances and angles for M(mepp)3»12H20 max  X  max  X  LIST OF FIGURES Page 3  Figure 1.1. Speciation diagram for ImM Ga-*+ at 25 "C Figure 1.2. Catechol to o-semiquinone to o-benzoquinone redox cycle  6  Figure 1.3. Heterocyclic ligands comparable to o-semiquinone  7  Figure 2.1. Mechanism for the conversion of a 4-pyrone to a 4-pyridinone  10  Figure 2.2. Precursor 3-hydroxy-4-pyrones  11  Figure 2.3. Ligand structures and abbreviations  14  Figure 2.4. Schematic representation of Method A  18  Figure 2.5. Infrared spectrum of Hmpp from 1700 to 300 cm  28  -1  Figure 2.6. Proton NMR spectrum of Hmpp in ( C D ) S O at 80 M H z  32  Figure 2.7. Mass spectrum of H2exn  35  Figure 2.8. Five-membered chelate ring  36  Figure 2.9. Tris(N-substituted-3-hydroxy-4-pyridinonato) metal complexes  41  Figure 2.10. Equation for the formation of the tris-ligand metal complexes  44  Figure 2.11. IR spectra of Al(mpp)3 from 1450 to 1200 cm-1  46  3  2  Figure 2.12. IR spectra of Hmpp and M(mpp)3 complexes from 900 to 300 cm  -1  52  Figure 2.13. IR spectra of Al(dpp)3 and Hdpp from 1700 to 1400 cm'  53  Figure 3.1. 4-Pyridinone resonance forms  58  Figure 3.2. O R T E P view of Hmpp, Hdpp, and Hmepp  60  Figure 3.3. Resonance forms for 2-pyridinone  63  Figure 3.4 Hdpp hydrogen bonded dimeric unit  65  Figure 3.5. Stereoview of the packing arrangement in Hmpp  66  Figure 3.6. Stereoview of the packing arrangement in Hdpp  67  Figure 3.7. Stereoview of the packing arrangement in Hmepp  68  Figure 3.8. ORTEP view down the c axis of the M(dpp)3 unit cell packing  71  Figure 3.9. Stereoview of the M(dpp)3 unit cell packing and stereoview of ice Ih  72  1  xi Figure 3.11 0. ORTEP view oftahepatrtis(liogfand)theportiHo-nbondoifngM(ndeptpw)o3rk Figure 3.12. ORTEP view of the tris0igand) portion of M(mepp)3 Figure 3.13. ORTEP view down the axis of the unit cell p Figure 3.14. ORTEP view down the c axis of the unit cell Figure 4.1.A1 NMR spectra of Al(mpp)3 at pH 1.8 Figure 4.2. Variable pHAl NMR spectra of Al(mpp)3 Figure 4.3.A1 NMR spectra of Al(mpp)at pH 8.9 Figure 4.4. Variable-pHAl NMR spectra of Al(dpp)3 Figure 4.5. Variable-pHAl NMR spectra of Al(mimo)3 Figure 4.6. Variable-pHA1 NMR spectra of 3:1 Hmepp to A1( Figure 4.7. Isomers of trivalent metals with asymmetric bidentate l Figure 4.8. 300 MHz *H NMR spectrum of Ga(dpp)3 in CD O Figure 4.9. 300 MHz *H NMR spectrum of Al(dpp>3 in Figure 4.10. 300 MHzH NMR spectrum of Al(dpp)in (CD)SO Figure 4.11. 300 MHz *H NMR spectrum of Al(dpp)3 in D2 Figure 4.12. CHspectra for Al(dpp)3 and Ga(dpp)3 in at Figure 4.13. Experimental and simulated CH3spectra in Al(dpp)3 Figure 4.14. Variable-temperature *H NMR spectra of CH3in Al( Figure 4.15. Variable-temperature *H NMR spectra of Ga(dpp)3 in Figure 4.16. NMR spectrum of 3:1 Al(ma)3 to Al(dpp)in Figure 4.17. *H NMR spectra of H and Hb in Al(dpp)3, Al Figure 4.18. Variable-temperatureH NMR spectra of Hb in mixed-ligan Figure 5.1. Equilibrium equations and constants for M with bidentat Figure 5.2. Speciation diagrams for Al-Hdpp at 3 mM and 3 Figure 5.3. Comparison ofAl NMR spectra and Al-Hdpp speciation c  72  27  72  27  3  27  72  l  C  3  32  CD3OD  c3  C  C  H  l  3  l  a  3 +  27  C  x i i Figure 5.4. Speciation diagram for Ga-Hdpp 1 Figure 5.5. Protonation equations for 3-hydroxy-4-pyridinone ligands and Figure 5.6. Plot of Al complexation vs. log of total ligand co Figure 5.7. Plot of Hmpp, Hdpp, and Hmepp logP3 values Figure5.8. Plot of Hmpp, Hdpp, and Hmepp pK values Figure 6.1. Liver biodistribution plotted asGa uptake vs. ligand Figure 6.2. Graph of %Ga complexed vs. ligand injection concentra 3+  76  a  xii L I S T O F A B B R E V I A T I O N S Abbreviation Meaning acac acetylacetonate a-RT a-substituted tropolonate AI13 [A104Ali2(OH)24(H0)i2] Al(dpp)3 tris(3-hydroxy-1,2-m^ethyl-4-pyriaUnonato)aluminvim(III) Al(ma)3 tris(maltolato)alummum(IJI) Al(mepp)3 tris(3-hydroxy-2-methyl-1 -ethyl-4-pyridinonato)aluminum(nT) Al(mhpp)3 tris(3-hydroxy-2-methyl-1 -hexyl-4-pyrioUnonato)alurninurn(lII) Al(mimo)3 rrisn(Tmi osniaot)aluminum(III) Al(mpp)3 rris(3-hyo!roxy-2-methyl-4-pyridinonato)aluminurnn(i) Pn overall formation constant B-band benzenoid ultraviolet absorbance band BBB blood-brain barier Bzma 3-benzyloxy-2-methyl-4-pyrone Bzmpp 3-benzyloxy-2-methyl-4(lH)-pyridinone Catechol 1,2-dihydroxybenzene Chlorokojic acid 5-hydroxy-2-(chloromethyl)-4-pyrone C/IR charge to ionic radius ratio AV frequency shift in Hertz between peaks in absence Av experimentally observed frequency shift S. vibrational in-plane bending mode £ molar absorptivity E D T A e t h y l e n e d i a m i n e t e t r a a c e t i c a c i d FAAS flame atomic absorption spectroscopy +7  2  e  x y  fac  facial geometrxiicv isomer Ga(ma3) tris(maltolato)gallium(JTJ) Ga(dpp3) tris(3-hydroxy-1,2-cumemyl-4-pyricUnonato)gaUiijmi(III) Ga(mepp3) tris(3-hydroxy-2-methyl-1 -emyl-4-pyricUnonato)galliumn(i) Ga(mhpp3) Ga(mimo3) Ga(mpp3)  rj H-bond Hdpp Hmepp Hmhpp Hmpp IR K Kojic acid L-dopa Maltol mer mQ Mimosine NMR V. P H2exn  n  Xmax  tris(3-hyclroxy-2-methyl-l-hexyl-^pyricUnonato)galliumn(i) tris(mimosinato)gam'umn (i)  ms(3-hyo!roxy-2-memyl-4-pyridinonato)gallium(III) absolute hardness hydrogen bond 3-hydroxy-1,2-dimethyl-4-pyridinone 1,6-oU(3-hyd^oxy-2-methyl-4-pyridinone)hexane 3-hydroxy-2-methyl-1 -ethyl-4-pyridinone 3-hydroxy-2-methyl-1 -hexyl-4-pyridinone 3-hydroxy-2-methyl-4(lH)-pyridinone infrared stepwise formation constant 5-hydroxy-2-(hydroxymethyl)-4-pyrone wavelength of maximum absorbance L-3,4-dihydroxyphenylalanine 3-hydroxy-2-methyl-4-pyrone meridional geometric isomer milUCurrie; equals 3.7 x 10disintegrations/second P-[N-(3-hydroxy-4-pyridinone)]-a-aminopropionic acid nuclear magnetic resonance vHi-borcatiaonoal/watesrtretcphairntgitionmodce oefficient 7  XV  7C -y  vibrational out-of-plane bending mode  pyromeconic acid  3-hydroxy-4( lH)-pyridinone  tfac  trifluroacetylacetonate  T  coalescence temperature  X  c  Ti  spin-lattice (longitudinal) relaxation  T  spin-spin (transverse) relaxation  2  TQ  nuclear quadrupole relaxation  UV  correlation time  W1/2  peak width at half height  ultraviolet  xvi ACKNOWLEDGEMENTS I w o u l d l i k e t o t h a n k t h e t e c h n i c a l s t a f f o f t h e M a s s S p e CS h. eO mi. stryCha Dn epartamnedntMasruipeptotratAusstatfrfiafoorfaltlhetheN irMRhelpl.ab Iforwouthledireshpeelc S t e v e t h e e l d e r a n d S t e v e t h e y o u n g e r i n t h e g l a s s b l o w i n g s h o p . the lTehnegthyfincaonncviearlsatiaosnssistanacbeoutin the fan d t o O v e f o r l i s t e n i n g t o o r m o f a U n i v e r s i t y G r a d u a t e TSeiagcmhaingXiA.ssistantship is gratefully acknowledged, as is a research g T h o s e w h o s e w o r k a p p e a r s i n t h i s t h e s i s a r e d e s e r v i n g o Thhiserefp orreomptI dweotuelrdminatliiokne toof tthheankelemSetnetvaelRaentatliygsest,oDw h o m I o w e a v e f o r a l l h i qhuaavnitnigtiestoofsmesplelciatiaosn dmiuacghrams,octanToilmasK.Iadnidd,CaarnedyDfoorn,theiTrerrisynt the hot lab. T h i s w o u l d n o t b e c o m p l e t e w i t h o u t a c k n o w l e d g i n g t h e s u p p o r t otof aAlllexithsefm e m b e r s o f t h e " O r v i g T e a m " p a s t a n d p r e s e n t . A o r c h a n g i n g O r v i g ' s s t u d e n t i n t o O r v i g ' s s t u d e n t s , a n d hniomtselfonlyforoldbeeringthawnillinhgistoreseatrackhedo n a g r a d u a t e s t u d e n t w h o , irector but also had more hair F i n a l l y , t o M a r i a n w h o w o u l d m a k e m y l i f e a m i s e r y i f SleisrsIslaoancely."Newt wherever he might be, "thank's for making tho DNMR3  xvii  To Marianw,anted andatodocmty p a r e n t s , B u d a n d D o r i s , w h o a l or in the family.  IofleavYeogimyBerrcaolleianguestheonsevetnhtehOg ra vim geteoafmtwhieth19t6 h0 e iWmomr orl "It ain't over 'til it's over."  Chapter I  1  General Introduction  Aluminum is the most abundant metal and the third most earth's crust. It is bound as oxides, primarily alumina, and c Surface water concentrations have remained minimal due to the low Al minerals; however, acidification due to acid precipitation has facil Al from soil to surface waters.Elevated concentrations of Al have and rivers in regions throughout the northern temperate zone that acidic substances, including eastern Canadaand the northeastern United The increasing concentrations of Al in surface waters can unforeseen consequence of human activity. By design, municipal water high levels of Al (up to 1 mg/L) because Ai2(SO,4)3 is com during water purification.Al is also found in relatively high concentra products used for human consumptionincluding drugs (e.g., 10-15 buffered aspirin), processed foods (cheese can be up to 0.7% A (about 5% Al). In spite of this exposure, Al is largely excluded low gastrointestinal absorption and efficient renal excretion. On average the range of 20-50 mg/day and the total body burden in normal As recendy as 1974 Al was regarded as a generally benig last twenty years a large body of evidence has accumulated t neurological dysfunctions that appear to have an environmental etio encephalopathy (DE) and amyotrophic lateral sclerosis (ALS) are espe because of their association with Al in drinking water. Patients on 1  2  4  5  2 the victims of DE and a correlation has been established betwee dialysis solutions and this often fatal syndrome.Natives of the is inordinately high incidence of ALS (commonly referred to as Lou G has been hypothesized that this could be related to the high lev and surface waters of the regions where the disease is endemic The association of Al with Alzheimer's disease was first recogni experiments where it was found that aluminum-induced encephalopathy course and histopathology similar but not identical to that found disease.-It has since been established that victims of Alzheimer' levels (10 to 30 times above normal) in the nuclei of the a has not been proven that Al is the cause of Alzheimer's or disorder, the circumstantial evidence is compelling when coupled with neurotoxicity of Al (first documentedin 1886). The case against several studies have been undertaken to assess chelating agents fo "benign" element from the body.''It was in the context of consequences of the increasing environmental exposure to aluminum that interested in the aqueous coordination chemistry of this group 13 me Our interest in gallium is due to its utilization as a radio agent. Ga has two isotopes that lend themselves to the det medicine:Ga (ti/2 = 78.1 h; y= 93.3,185, 300 KeV; accelerat = 68.3m;y=511 KeV from (3+ annihilation; generator product). T 1960's thatGa (administered as the citrate) localized in soft tissue considerable clinical research and the citrate (the only commercial radiopharmaceutical ofGa) is now widely used in oncological nuclea Despite its availability and easy detection, the use ofGa has been of tumors and inflammatory lesions.' 9  121  41  48 51  76  76  76  76  012  3 The aqueous coordination chemistry of Ga with multidentate ligan hydroxy, amino, carboxylate, and catechol moieties has been ex medicine in mind;however, there has not been much work of this trivalent metal with simple bidentate ligands. One reason f instability of the group 13 metals; Al and Ga undergo extensive that is pH, concentration, and time dependent.-Figure 1.1 is dependent speciation of Ga(the speciation diagram for Al is ve diagram is simplified as no polymeric species are included and it i nature of the Ga-hydroxide complexes. Al and Ga form insoluble pH 4-8 region and the logarithm of the overall formation constants gelatinous precipitates are 12.7 and 23.2, respectively (based on d chelate these metals in an aqueous environment, a ligand must b hydroxide ion. 42-22  572  +3  %tootafl . Ga 40 An  3+  Ga hexaaqua ion Ga(OH) Ga(OH)2 Ga(OH)3 Ga(OH)4  H  10 PH Figure 1.1. Speciation diagram for ImM Gaat 25 °C (base +3  4 We are interested in ligands and metal-ligand complexes that relevant". For the purposes of our research, biologically relevant com the following properties: molecular weight (MW) < 600, water lipophilicity to cross cell membranes, neutral charge at physiological p metal-ligand complexes) significant thermodynamic stability. In addition, t to be reasonably nontoxic as they are to be used in animal Ga. This is of use both to determine the efficacy of the l this radionuclide and to gain some idea of the biodistribution of t are no Al isotopes suitable for this type of study. InGa imaging experiments, water solubility at the |J.M leve sufficient; for facile spectral characterization of the metal-ligand complexes mM is desirable. The constraints on size, lipophilicity, and charge of our interest in ligands and metal-ligand complexes that could concei The brain is separated from the circulatory system by the most se the blood-brain barrier (BBB). It has been shown that the per directly related to the size and lipophilicity (and therefore charge) o upper limit on size has yet to be determined, but it is thought not inhibited by size for compounds of MW < 500. For in vivo experiments, the need for thermodynamic stability by the afinity of the blood serum protein transferrin for trivalent m is an iron transport protein that has two binding sites for metal (The ionic radii and transferrin binding constants (Ki and K) for listed in Table 1.1.) The large binding constants and the high transferrin binding sites in human blood (-50 p^M ) make transferr of trivalent metal ions. The size specificity of transferrin binding difference in the stability of the Al and In complexes. WhenGa 76  76  72  2  92  76  5 imaging, the observed biodistribution is that oGfa-transferrin as citrate for the metal ion. A ligand must be able to maintain the radi the high levels of transferrin in the circulatory system even to b imaging agent. Table 1.1. Effective ionic radii (six coordinate)and transferrin bin Al Ga In Fe IonicLogRadiKuis (A) 12.0 . 5 3 5 0 . 6 2 0 0 . 8 0 0 0 . 6 4 5 9 2 0 . 3 3 0 . 5 2 2 . 7 Log K.2 12.3 19.3 25.5 22.1 From ref. 30. In from ref. 28, Al and Fe from ref. 29, and Ga fro The group 13 metals have an affinity for oxygen containin research in our laboratory has concentrated on bidentate ligands wi a-hydroxyketone binding groups. My initial project was the synthe complexes with several 1,2-dihydroxybenzene derivatives (catechols). Variou have been used as ligands for a number of metals,including In addition, the catecholamines (dopamine, noradrenaline, and adrenali medullary hormones responsible for the coordination of the sympathetic and the adrenal medulla. The catechol ligand of greatest interest dihydroxyphenylalanine (L-dopa). L-Dopa is not a neurotransmitter, but precursor that is converted to dopamine by L-dopa decarboxylase.LBBB and it is the principal therapeutic for the treatment of Parkin neurological disorder considered to have an environmental etiology. 76  3  a  b  23  53  63  6 The ability of L-dopa to localize in the brain made it a research. However, it has two properties that are liabilities within study: the tris-ligand metal complexes are trianionic and free (and readily oxidized in water. The first property makes it more difficu complexes to cross cell membranes. The second property is one s they can be oxidized to the free radical o-semiquinone and to o The redox sensitivity of the catechol ligands complicated the tris(catecholato) Al and Ga complexes, and we were unable to with these metals.  Figure 1.2. The catechol to o-semiquinone to o-benzoquinone r Our research group enjoyed greater success with two classes of containing the a-hydroxyketone moiety. The binding group is the o-semiquinone (Fig. 1.3) and a number of transition metal comple ligands have been isolated.These heterocyclic ligands do not hav properties of the catechols and their tris-ligand metal complexes are pH. Contemporaneous with my catechol research, other members synthesizing Al and Ga complexes with several 3-hydroxy-4-pyrones. proved the most interesting was 3-hydroxy-2-methyl-4-pyrone (maltol), occurring compound that is used as a food additive. In animal collaborators, tris(maltolato)aluminum(III) (Al(ma)3) was found to b 23  73  7 neurotoxic as Al-lactate, the agent commonly used to induce alumin The work in our laboratory established the afinity of the oc-hydrox for Al and Ga; however, maltol was not able to redirectG preliminary imaging experiments. -4-pyridinone xy-4-pyrone 3-hydroxyO o-semO iquinone 3-hydroO O" O' O*  76  I  H Figure 1.3. Heterocyclic ligands comparable to o-semiquinone. The 3-hydroxy-4-pyridinone* nitrogen heterocycles offered several advant the 3-hydroxy-4-pyrones. The primary advantage was synthetic versatilit in 4-pyrone is essentially inert whereas a variety of substituents can nitrogen in 4-pyridinone. The 3-hydroxy-4-pyridinones are also strong indicated by the pKof the 3-hydroxyl group) and this should complexes that are more stable than those formed by the 3-hydro my research was diverted from the catechols and was directed tris(3-hydroxy-4-pyridinonato) Al and Ga complexes. One of the 3-hyd ligands employed in this study has been used in clinical trials as iron overload diseases'.The ligand was shown to be nontoxic class of nitrogen heterocycles would be suitable for animal biodistribu a  0943  * These compounds are also commonly referred to as 4-pyridones or 3,4-dihydroxypyridines.  8  felt the 3-hydroxy-4-pyridinones were similar enough to maltol to attributes (sufficient water solubility and lipophilicity) of the Al- and while having the potential to form tris-ligand metal complexes of variability. My project involved the synthesis of a homologous series of hydroxy-4-pyridinones. This was done to provide bidentate ligand lipophilicity that could possibly have different biodistributions. A potenti ligand was also synthesized by linking together two 4-pyridinones vi Once isolated, the ligands were used to synthesize tris3-(hydroxy-4-py Ga complexes that (by design) were all near the upper size into the brain.* The ligands and metal-ligand complexes were f elemental analysis, mass spectrometry, and infrared, proton NMR spectroscopy. Several studies were then undertaken to determine compounds fit our criteria for biological relevance. The water solub and, to indicate the degree of lipophilicity, the octanol-water part determined. The hydrolytic stability of the Al complexes was examin potentiometrictitrationswere used to determine the thermodynamic stabilit Ga complexes. The proton NMR spectra of the metal complexes indicated an occurring and this process was identified by a variable-temperature crystals of several ligands and metal complexes were grown and were determined by X-ray diffraction. A concerted effort was made state structures to the aqueous solution behavior of these compounds * See Table A.1 in the Appendix for a listing of molecular weights.  9 was placed on examining the hydrogen bonding interactions that were the solid state and were also found to persist in solution for bo tris-ligand metal complexes. Based on the results of this research project, several of the Ga biodistribution study. This study is stil underway and the prel briefly discussed in Chapter VI. Under conditions of ligand excess pyridinones can redirectGa from transferrin; however, the resulting biodi significantly different from that of Ga-citrate and it is doubtful these any applicability as radiopharmaceutical imaging agents. Because the tran constants for Ga are seven orders of magnitude greater than those compete forGa has the potential to form stable Al complexes biodistribution study does indicate the 3-hydroxy-4-pyridinone ligands hav therapeutic Al chelators and this application may be pursued in the f This research project had another goal separate from the ex coordination chemistry of group 13 metals and the appraisal of the class of ligands. This goal was the development of a methodology applicable for future projects undertaken in this laboratory. We es identify and assess the techniques that would be the most usef biological potential of a ligand prior to the transition from in vitro 76  76  76  Chapter II A.  10  Synthesis and Characterization  3-Hydroxy-4-Pyridinone Ligands  2.1 Introduction The first synthesisof a 4-pyridinone was reported in 1884 considerable literature has accumulated on these nitrogen heterocycles for their synthesis can be grouped into three categories: ring closur conversion of other heterocyclic ring systems, and substitution and pyridine or its derivatives. A review of the literature indicated that simplest method for the synthesis of 3-hydroxy-4-pyridinones. 14  42-44  Figure 2.1. Mechanism for the conversion of a 4-pyrone to a  1 1 One of the oldest ring conversions involves the ammonolysis cyclic ether 4-pyrone. The accepted mechanism (Fig. 2.1) is n primary amine, followed by ring opening, loss of water, and corresponding 4-pyridinone.There is no direct proof offered for th molecular orbital calculations predict an enhanced probability of nucl occurring at position 2.Further indirect evidence is the effect ring the conversion reaction: by induction (when in the 2 position) electron-withdrawing groups enhance reactivity while electron-donating group opposite effect Although the electron-donating hydroxyl substituent can reduce the the reaction, there are many examples of the conversion of 3corresponding 3-hydroxy-4-pyridinones. The 3-hydroxy-4-pyrones that were as possible synthetic precursors in this study are listed below. 54  64  44  Ri,R, R3 = H Pyromeconic a ,;RR==CC 0H lO | 11 R!R=RH;=RH HOH KoMjeicco * ° R R = H; R = CH Maltol Figure 2.2. Precursor 3-hydroxy-4-pyrones. The conversion reaction was first used for the structural determ occurring compounds. The structures of maltol(in 1906) and k were confirmed by their reactivity with primary amines; conversely 4-pyridinones were also verified with the conversion reaction. Meco R  2  JLoH 2 Rs  l v  f l  ? 1  2  3  2  3 2  3  2  3  74  2 3  12 pyrolysis product pyromeconic acid were employed in the structure hydroxy-l-alanine-4-pyridinone (mimosine*); this was a difficult problem t some controversy prior to its definitive resolution in 1947 The the conversion reaction followed from its investigative role: kojic acid a in the attempted synthesis of 4-piperidinols,pyromeconic acid was the the first total synthesis of mimosine,and a series of N-alkyl s was made from pyromeconic and meconic acid. The conversion reaction with the 3-hydroxy-4-pyrones, however, somewhat inconsistent. For example, meconic acid gave the expecte methyl-, ethyl-, and propylamine but not with n-butylamine or a-ph reported yields were further evidence of inconsistency: a 40% yield 15% yield with ethylamine. The enhanced reactivity imparted by c shown by the even lower yield (10%) for the reaction of methylamine. An extensive study on the conversion reaction concluded and less hindered amines would give the greatest yields of 4-py even the results with ammonia were subject to variation: Heyns 30% yield for the ammonolysis of kojic acid but others found unproductive. It is possible to improve the reactivity of the 3-hydroxy-4-pyrone hydroxyl group. The utility of including a blocking step has long it was found that maltol and pyromeconic acid would not undergo 3-methoxy derivatives would react to give the expected 4-pyridin 49,50  15  25  35  34  Earlier authors refer to this compound as leucaenol or leucaenine, a natural product isolated from the seeds of Leucaena glauca in 1937. It was subsequendy shown to be identical to mimosine,firstisolated in 1936 from the sap of the tropical shrub Mimosa pudica; since mimosine was coinedfirst,proof of identity made the other names redundant  13 blocking-deblocking sequence used dimethyl sulfate to form the methoxy deblocking by acid hydrolysis in HI or HBr. The blocking reaction but therigorousconditions for deblocking could be a major experiment Spenser and Notation attributed the poor yield and irreproducibili mimosine synthesis to the six hours of refluxing in aqueous HI methyl blocking group. His mimosine synthesis (from pyromeconic a blocking group because it could be removed under less strenuo continued interest in mimosine* led Harris to attempt the improveme ancillary to this goal, he reported a preparation for the ammonolys conditions that also employed the benzyl blocking group.His prepar starting point for our synthesis of N-alkyl substituted-3-hydroxy-4-pyridinones. Maltol was chosen as the synthetic precursor despite the electr group in the 2 position. By induction the methyl substituent re conversion reaction, but by the ortho effectit reduces the acidity (pKof 8.36 for maltol vs. 7.69 for pyromeconic acid and 7.66 therefore a stronger base and a better ligand for Lewis acid m found that maltol formed Al complexes of greater stability than did from our laboratory had shown maltol to be a good ligand for importantly, the metal complexes had the desired degree of w lipophilicity.We felt the 2-methyl substituent would play a similar chemistry of the 3-hydroxy-4-pyridinones synthesized from maltol thereby for the reduced effectiveness of the conversion reaction. 65  75  a  73  Animals fed the seeds or leaves of Leucaena glauca suffered from hair loss, and mimosine was found to be the active ingredient responsible for this. This depilitory activity was the reason for the renewed interest in its synthesis as it was thought to show promise as a chemical de-fleecing agent to improve wool harvesting.  14 3-Hydroxy-2-methyl-4-pyridinone ligands were synthesized from maltol following primary amines: methyl-, ethyl- and /i-hexylamine; 1,6-diaminohexane. The diamine was used to synthesize a bispyridi potentially tetradentate ligand. The monopyridinone ligands were named letters of the substituents, e.g., 3-Hydroxy-2-methyl-l-hexyl-4(para)-pyridinone Hmhpp. The substituents were ordered to emphasize the acidic pro base of the ligand could be readily identified, e.g., mhpp". T named from its hexane progenitor again stressing the protons. The synthesized in this study are in Figure 2.3. 3-Hydroxy-2-mO ethyl-4-pyridinone R=H CH (CH)CH R l,6-Di(3-hydroxy-2-memyl-4-pyridinone)hexane 3  CH2CH3  25 3  Hmpp Hdpp H m e p p Hmhpp Hexn 2  Figure 2.3. Ligand structures and abbreviations.  15 2.2 Materials and Methods Of the preparations in Method A, that of Hmpp was others were reported previously by us, with ours being the fi Hmpp has been synthesized by other methods.The synthesis of and of Hmhpp from maltol D-glucosidehad been reported prior experimentation was undertaken to simplify Method A by eliminating the results are reported as Method B. The precedents and ra synthetic effort will be presented in Section 2.3. All chemicals were reagent grade or better and were us purification. The progress of the reactions was monitored by thin (TLC) on silica gel plates with 5% methanol in as the material 3-benzyloxy-2-methyl-4-pyrone (Bzma) was made from the available maltol (Aldrich) by the method of Harris (> 95% yie yellow liguid that was used without further purification (TLC pure) were measured with a Mel-Temp apparatus and are uncorrected. U the quoted yields were for purified compound and they were calcu characterization data, including the elemental analyses, are found in S Method A 2.2.1 3-Benzvloxv-2-methvl-4(lHVpyridinone. Bzmpp. This intermediate wa by the treatment of Bzma with a solution of ammonia in ethano Harris.Crystallization from hot ethanol gave 4.70 g (60% yield) Mp 166 °C. 16  46  CH2CI2  65  16 22. 3-Hvdroxy2--methvl-4(lH)-pvridinone. Hmpp. Bzmpp was deblocked hydrolysis with 40% HBr in acetic acid.The crude product was was added to adjust to pH 8. The solution was c gave 2.63 g (47% yield) of pink crystals. Mp 265 °C dec. 22..3 3 -Hvdroxy-21-.dimethvl-4-pvridinone. Hdpp. A solution of Bzma mmol) in 50 mL ethanol and 100 mL 40% methylamine in w a sealed flask at 20 °C for 72 hours. The excess amine wa residue taken up in mL water. The organic phase was e mL) and evaporation gave a yellow oil. The benzyl ether (Bzd THF and hydrogenated under ambient conditions over 5% Pd/C hydrogen uptake ceased. The solution was filtered and the solid water (3 x 100 mL). The aqueous solution was concentrate commenced and then chilled at 10 °C overnight. Recrystallization fro 3.46 g (52% yield) of white crystals. Mp 260 "C dec. 22..4 3 -Hvdroxv2--methvl-1 -hexvl-4-pyridinone. H mhpp. Bzma (71.0 g, and hexylamine (16.6 g, 164 mmol) were dissolved in 100 sealed flask at 20 °C for 72 hours. The solution was concentr 12M HC1 was added to the oily residue. After heating on excess acid was removed in vacuo. To the residue was added adjusted to 8 with 2 M NaOH at which time a precipitate f solid gave 3.40 g (48% yield) of a yellow powder. Mp 123 22..5 1 .6-Di(3-hydroxy2--methyl-4-pyridinone)hexane. Bzma (30. g, mmol) and 1,6-diaminohexane (80. g, 690. mmol) were stired hours at 20 °C. Concentration in vacuo gave an oil that concentrated HC1 and heated for 30 minutes on a steambath. 65  NH4OH  100  H2exn.  1 7 vacuo and the residue was dissolved in 50 mL water. The pH KOH and the resulting precipitate was collected. The filtrate was evaporation left a solid residue that was added to the earlier prec yield) of a yellow powder. Sublimation gave analytically pure samp Method B The reactions were done under positive N2 pressure. The condensing columns but the reactions were only heated to 50 aqueous recrystallization, the pH was adjusted to take advantage of between the 3-hydroxy-4-pyridinones and maltol (pKof -9.8 and Purification at pH 8 ensured disproportionately greater ionization of the material thus aiding in the removal of the primary contaminant i solution pH was measured with a Fisher Accumet model 805 pH 2.2.6 3-Hydroxv-2-methvl-4-(lH)pyridinone. Hmpp. To a chilled solu (2.53 g, 20.1 mmol) in 30 mL of water was added 4 followed by the slow addition of 5 mL 6 N HC1. The p solution was heated for 36 hours. Concentration in vacuo and hours gave 1.42 g of crystalline product. Recrystallization from hot (48% yield). 2.2.7 3-Hvdroxv-1.2-dimethvl-4-pvridinone. Hdpp. Maltol (5.05 g, 40 mL water was placed in a 3-necked flask to which was connecte 15 mL 40% methylamine in water (193 mmol) and 50 mL w for 12 hours with the amine being slowly added during the fir concentrated in vacuo and cooling at 10 °C overnight gave microcrystalline solid. The filtrate was extracted with 150 mL liquid-liquid extractor. Evaporation of the solvent left a residue tha a  C  1 8 acetone and added to the filtrant. Purification by sublimation gave white powder. 2.2.8 3-Hvdroxv-2-methvl-1 -ethy1-4-pvridinone. Hmepp. A solution 20.0 mmol), 70% ethylamine in water (13.3 g, 206 mmol) in an ice bath. 30 mL of 6 N HC1 was slowly added solution was heated for 24 hours, transferred to a continuous li the product was extracted into mL for 6 hours. Evap a brown solid that was washed with cold acetone. Sublimation ga a white powder. Mp 205 °C dec. 2.3 Discussion of the Synthetic Procedure Initially, the ammonolysis of maltol was attempted under the reported herein for its benzyl ether derivative Bzma. The resulting conversion to Hmpp confirmed the utility of employing a blocking (Fig.2.4) can be divided into three steps: blocking of the ring conversion reaction, and deblocking to give the 3-hydroxy-4-pyridinone. reaction is a Williamson ether synthesis with maltol going to Bzma o O O OH CH Figure 2.4. The schematic representation of Method A. 100  R  CH2CI2  R  3  19 The conversion reaction was done at ambient temperature and was obtained after three days. A similar procedure allowed seven Hdpp was the same as in Method A (~50%).Attempts we reaction by refluxing, but product isolation was complicated by byproducts and yields were significantly lower. The reaction was mixtures because of the hydrophobicity of Bzma and this may hav earlier study had concluded that the conversion reaction was more alcohol.In the synthesis of Hmpp, the 4-pyridinone benzyl ethe as an analytically pure solid when triturated by acetone and recrys The other benzyloxy intermediates only formed as oils (probably d N-alkyl substituents) and the preparations proceeded direcdy to the deb For Hmpp, deblocking was done by catalytic hydrogenation o in 40% HBr in acetic acid, with the latter giving the greater reversed with Hdpp, and hydrogenation gave better results. Acid hy was used for Hmhpp, and H2exn required concentrated HC1 to group. Restricted access to the hydrogenation apparatus and the ne deblocking left acid hydrolysis as the method most often employ reaction was done with a five-fold excess of amine, and the by rotary evaporation. For 1,2-diaminohexane this was insufficient an gave the dihydrogenchloride salt of the precursor amine as a byp product was difficult and undoubtedly contributed to the poorer yield to that for Hmhpp (18 vs. 48%). The problem of salt formati in all the preparations albeit to a lesser extent. The least desirable aspect of Method A was thetimerequired up to five days to complete. A possibility for shortening the pro reaction in a sealed glass tube: maltol and its 3-methoxy derivative 36  15  2 0 respectively when heated with aniline in sealed tubes (40 hours at with sealed vessels were unsuccessful as both maltol and Bzma fa ammonia or n-hexylamine. The decision to forego further experime technique was facilitated by the physical restrictions of the glass tub recent attempts to circumvent sealed tube reactions in a study arylamines.Another way to speed up the procedure would be blocking step. This step was included to avoid deprotonation of the the increased electron-donating capacity of the O" anion would imped 4-pyrone resonance hybrid most susceptible to nucleophilic attack. Inste proton, the ionic state of maltol could be controlled by buffering pKof the hydroxyl proton. Hdpp was obtained in a 55% yield from a procedure employi acetate buffer. The reaction was done at slightly acidic, neutral, the later giving the best results.The pH was not specified but gave an initial pH of 9; however, we were unable to achi Hmepp with this procedure. It was felt the problem lay with t the chromatographic separation step this necessitated) and the results experimentation with buffered systems. Recently the conversion reaction without blocking or buffering.This simple preparation gave a 50% of maltol with methylamine, but the results were poorer with eth (24% and 21% yield, respectively) showing once more the inconsist reaction. This limitation made it worthwhile to continue to look for have wider applicability than the above preparations and that would A. But "method" is the wrong word; rather what is reported an experimental strategy that recognizes the primary importance of the reactants (and consequently the solution pH) to the success of the 66  a  26  76  2 1 The utility of this strategy can be demonstrated by examining Hdpp without blocking or buffering.The ratio of reactants was maltol; given the concentrations used, the initial pH was 1.5 an ionized. At the start of the reaction there was essentially no u present. As the reaction was heated, the concentration of the reduced and after the prescribed 5 hours of refluxing the pH methylamine would be 99% ionized and unreactive. Somewhere in extremes there is a pH region where the nucleophilic attack mechan preparation does not control the time spent at this optimal pH. is a nucleophile able to cleave 4-pyronesso the avoidance of s reason for controlling the solution pH. A simple way to control the pH without buffering would base and this was done in the preparation of Hdpp as reported was added to an aqueous maltol solution over 4 hours and b amine (Bp -6.3 °C), the reaction was heated at 50 °C over simple preparation gave higher yields (70 vs. 52%) and was fast Method A. It was a cleaner preparation than the other synthes repeats of literature preparations) and gave microcrystalline precipitate d reaction mixture. The reaction with ammonia provided a better trial for this ammonolysis of 3-hydroxy-4-pyrones of which we are aware employed a sealed reaction vessel, or both. This is probably due to the -33.4 °C) and its reduced nucleophilicity compared to alkylamines. choice of ammonolysis rather than aminolysis to explore the necessi blocking group was an unfortunate one. Consistent with the diffi attempts to prepare Hmpp simply by the slow addition of ammoni 76  86  2 2 results were obtained with an ammonium chloride buffer. Buffering produced no reaction (monitored by TLC) but a significant decrease occurred upon overnight refluxing at pH 9.3. These results coupled unbuffered reaction (pH > 10) indicated a narrow range of optim For the Hmpp preparation in Method B, the solution was initial concentrations such that the ratio of unionized reactants would maltol. This was a one pot synthesis that gave the same y was again a simpler and faster preparation. The N-ethyl derivative an analogous manner to Hmpp. The buffered preparation gave a either Method A (35%) or the unbuffered preparation cited previously By controlling the ionic state of the reactants it is possible 3-hydroxy-4-pyridinone analogues without using a blocking group. The same or better and the preparations are simpler and faster for Meth Method A. The routine synthesis of 3-hydroxy-4-pyridinones in our done without a blocking group, as are the exploratory efforts to ligands. 2.4 Characterization of the 3-Hvdroxy-4-Fyridinone Ligands The ligands were characterized by elemental analysis, infrared NMR spectroscopy, and electron impact mass spectrometry (EI-MS). T data were completely consistent with the structures as shown in F attention was paid to the hydrogen bonding (H-bonding) in these spectra indicated that the compounds were H-bonded polymers in the the intermolecular H-bonds involved the hydroxyl and the carbonyl mo  2 3 results for Hmpp showed the cyclic secondary amine was also NMR experiment indicated the H-bonding persisted in solution. The elemental analyses (C, H, N) were performed by Microanalytical Laboratory of this Department. The IR spectra were r Elmer PE 783 in the range 4000-200 cm. All samples were spectra were referenced to polystyrene film. The proton NMR spect a Bruker WP-80 and a WP-400 instrument. The 80 MHz author, and spectra at 400 MHz were supplied by the U.B.C was performed on a Kratos MS 50 spectrometer and all mass U.B.C. mass spectrometry service. 2.4.1 Elemental Analysis Prior to submission for analysis, all samples were puri recrystallization or by sublimation in vacuo (0.03 torr) with heating. sublimation were relatively mild (<140 °C) except with H2exn. insoluble for recrystallization and the analyzed sample was sublimed Wood's metal bath. Table 2.1. Results of the ligand elemental analyses (Found/[Calculated] 1-  %G  %N  57.58 [57.57]  60.45 [60.41]  62.56 [62.72]  68.56 [68.87]  64.81 [65.03]  5.70 [5.65]  6.49 [6.53]  7.28 [7.25]  9.01 [9.15]  7.19 [7.29]  11.23 [11.24]  10.00 [10.07]  8.98 [9.14]  6.61 [6.69]  8.40 [8.43]  24  2.4.2 Infrared Spectroscopy Of the spectroscopic techniques employed in the study of the the most useful as it affords a facile confirmation of the outcom The ring skeleton vibrational modes of the pyridinones and the py found in benzene.This typically results in four ring-stretching modes 1400 cm" , and the pattern changes predictably when the cycli nitrogen. An examination of the IR spectrum readily confirms the to the corresponding 3-hydroxy-4-pyridinone. In the 1960's there were a number of IR spectroscopic s motivated primarily by the difficulty of assigning the carbonyl stretch compounds".Although none of these studies included any 3-hydro was possible to make tentative spectral assignments based on this in Tables 2.2 and 2.3 were supported by reference to Bella correlations in the infrared.Hmpp is of special interest because o donor in this compound; the assignments for the modes in Hmp are included with those of the benzyl ether intermediate (Bzmpp), analogue, d2-Hmpp in Table 2.2. Deuteration was by repeated and was undertaken to determine the relative strength of the assignments for the other compounds are in Table 2.3. The I 1700 to 300 cmis reproduced as Figure 2.5. The abbreviati modes are: V, stretching; 8, in-plane bend; n, out-of-plane ben symmetric. The assignment of VOH in Hmpp is confirmed by Bzmpp band disappears and by d2-mpp, where it shifts to 2420 cm 24  1  1976  27  1-  25 1.35, indicative of a relatively weak H-bond. The theoretical reduced masses) and for free OH it is 1.355.This ratio strength of the H-bond increases and can reach unity for ver combination of a lower zero point energy for H than D, O — H — O interaction explains this low ratio.In the N-alkyl deriva below 3200 cm"which suggests stronger H-bonding than in Hmpp the lower these compounds, e.g., 1645 cmin Hmpp and In all cases the bands are broad and within the range for n H-bonding. The assignment of is confirmed by Bzmpp, d2-Hmpp, a at 2800 cm"in Hmpp, shifts to 2650 cm"on benzylation, mean = 2180 cm) on deuteration, and finally disappears on is 1-28 and splitting of is not without precedent, deuterio-2-pyridinone.The authors list a 3100 cnr band (in K compound which supports our assignment, and also state that the 2-thiapyridinone is the strongest NH hydrogen bonding yet recorde zwitterionic structures. The deuteration ratio for the thiapyridinone is N—H—S stretch appears at 3160 cm. It is impossible to separate the and the higher energy the compounds. The bands are extensively coupled and there localized solely in the carbonyl bond.The highest wave number character and the relatively low energy of this band (below 16 indicates it is acting as a H-bond acceptor. In the SCH regi bands are observed with additional bands for the compounds consistent with the presence of additional methyl and methylene extensive mechanical coupling of the 6OH and Vco modes in the 37  47  1  m  1-  Vc=o  27  1  1  V N H  1-  (VNH/VND)  57  V N D  1  1-  Vc=o  67  V  26  so that an unambiguous assignment is not p o s s i b l e . ' 74  77  The 8Ri g and T C R j n  ng  assignments correlate well with the literature and the sharp out-of-plane bending mode (ca. 820 cm ) is the most distinctive feature of the lower energy region in all of the spectra. -1  Table 2.2. Infrared bands for Hmpp and substituted analogues (cm ). -1  Hmpp  d2-Hmpp  V O H (D)  3270 b  2420  VNH(D)  2800 b  2180  2650 b  vc=o  1645  1630  1630 sh  and  1620  1550  1620  VRing  1540  1540 sh  1535  1500  1490  1500  1420 w  1420 w  1410  VcOand  1300,1270  1325  1190(Vcoc)  50H(D)  1245  900 b  8Ring  1225  1245 sp  1215 sp  1110  1150  1110  1045  1090  1050  Assignment  Bzmpp  27 Table 2.3. Selected infrared absorption bands (cm). All bands are Assignment Hmpp Hdpp Hmepp Hmhpp H2exn (b) 3270 3150 3180 3195 3195 VCHring (w)3100 3010 3040 3060 3060 2920 w2940 w2980 w296 0 , 2 9 3 5 2 9 6 0 2860 2890 1645 1630 1630 1630 1630 aVnRding 1 6 2 0 1 5 6 5 1 5 7 5 1 5 8 0 1 5 8 0 1 5 4 0 1 5 3 0 1 5 3 0 1 5 3 0 1 5 3 0 1 5 0 0 1 5 1 5 1 5 1 0 1 5 1 0 1 5 0 5 1420 w1400 w1405 w1405 w1400 m 1450 w1460 1450 w1465 w1460 w 1380 11 33 83 05m1365 11 33 85 00w1 3 8 0 w 1360 m Vcoand 1300, 12701280 1260 1270 sh 1300,1280 1245 1250 b 1230 1240 b 1235 b SRing (w) 1225 1230 1250 1220 1220 1 1 1 0 1 1 2 5 1 1 3 5 1 1 7 0 1 1 7 0 1045 1 0 5 5 1 0 4 0 1 1 3 0 1 0 4 0 1030 1040 TCRing 830 820 830 850 850 m, medium absorption; w, weak absorption; sp, split band; 1-  V O H  VCH3(2)  vc=o  8asCH  8syCH  80H  28  Figure 2.5. Infrared spectrum of Hmpp from 1700 to 300  29 2 . 4 . 3  Proton NMR Spectroscopy The spectra of Hmpp, Hdpp, and Hmepp were recorded in D2O; the lipophilic  Hmhpp and the insoluble H2exn required 20%  CD3COOD/D2O. The acidic N H and O H  protons were detectable if the samples were rigorously dried and the spectra were recorded in aprotic solvents: (CD3)2SO for Hmpp and Hdpp, and CDCI3 for Hmepp and Hmhpp. H2exn was only soluble in acidic solution so it was not possible to detect the hydroxyl proton signals. The spectra in D 0 were referenced to an external (CH3)3Si(CH2)3S03Na 2  (DSS) signal and the spectra in aprotic solvents were internally referenced to the solvent signal. The chemical shifts listed in Table 2 . 4 were recorded at 4 0 0 M H z except those of the O H protons that were at 80 MHz.  The spectra have A B doublets for the ring protons (J b of 7 to 8 Hz), a singlet for a>  the ring methyl group, and a series of signals for the methyl and methylene protons of the R substituent. The signal from the protons on the carbon direcdy attached to the ring nitrogen is shifted downfield from that of the ring methyl group. Deshielding by the electronegative nitrogen is the reason for this and for assigning the lowfield doublet to H . a  78  For Hdpp and Hmepp, the signals from the alkyl protons are easily assigned. With  the exception of the CH2d,e methylene protons in Hmhpp, we did not attempt to resolve the second order spectra of the alkyl protons in Hmhpp or H2exn. In all cases, peak integrations are consistent with the assignments as given in Table  2 . 4 .  In the solid state, Hmpp forms intermolecular H-bonds and this interaction persists in solution as evidenced by the spectra recorded in (CD3)2SO. As the sample concentration is increased, the O H resonance shifts downfield from 5 . 6 4 0 . 0 9  ppm at 0 . 0 3  M to 6 . 9 0  ppm at  M. The direction of the shift is indicative of H-bonding and a dependence on  concentration is typical of intermolecular H-bonding. The O H chemical shift is also  30 affected by the concentration of water in the solvent. The spectr (0.09 M) prepared in a drybox under N2 is reproduced as Fi is assigned to the broad peak at 6.90 ppm; leaving the N atmosphere for several days caused the OH resonance to shift to were obtained with a 0.03 M solution of Hmpp and with H probably due to the diffusion of atmospheric water into the previou upfield direction of the shift suggests the intermolecular H-bonds are between the 3-hydroxy-4-pyridinone and water. The NH resonance in Hmpp is assigned to the broad signa a peak in Bzmpp at 11.30 ppm (both in at 80 MH D2O to NMR samples of Hmpp results in the loss of both a of the second order coupling to H which appears as a broad only with this compound (see Fig. 2.6). The presence of an ppm in 2-pyridinone has been taken to suggest that the compou dimers in solutionjust as it does in the solid state.Our resu the existence of intermolecular H-bonded arrays in solution with both acting as proton donors. The IR spectra and the crystallographic clearly show that this is the case in the solid state. (CD3)2SO  a  97  08  31  Table 2.4. Proton NMR chemical shift (8) data for the free ligands at 400 MHz (ppm).  o  R= H  bH^JL/OH H  Xx N"  a  C H 3d , l  CH  |  3  CH . C H  2d 3e  3 r C  CR.CH  R  2d  C H . CH  2e  2f  2g  C H  U  2h  CH . 31 H  CH  2 d  CH  2 e  CH  2 f  CH  2 g  CH  2 h  3  CH  \  C C  2 i  -N  Hmpp  Hdpp  Hmepp  Hmhpp  OH  6.90  6.35  6.43  5.96  Ha(d)  7.58  7.58  7.56  7.57  7.55  H  6.51  6.46  6.43  6.71  6.69  2.34  2.36  2.33  2.16  2.13  *3.73 (s)  4.00 (q)  3.87 (t)  *1.26 (t)  1.37 (tt)  yi.37 (m)  CH f,g,h 2  0.85 (m)  20.93 (m)  CH  3 i  0.38 (m)  Assignment  b  (d)  CH  3 c  CH  2 d  CH  2 e  (s)  H exn** 2  x  3.87 (t)  abbreviations: s = singlet; d = doublet; q = quartet; t = triplet; tt = triplet of triplets; m = multiplet. * ReadCH3. ** For H exn, the two rings are apparently equivalent, e.g. read: H to two protons. 2  x  Read C H ^ i with integral equal to four protons,  y ReadCH  2e>h  z Read C H f , 2  g  "  a>a  « with integral equal  32  9 03  O  n rt  Figure 2.6. Proton NMR spectrum of Hmpp in (CD3)2SO at 80 MHz.  33 2.4.4 Electron Impact Mass Spectrometry The molecular ions, base peaks, and selected fragment ions listed in Table 2.5 and the mass spectrum of F^exn is reprod assignments were made by comparison with the literature on 2-pyr scant work available on the 4-pyridinones.The ring fragments excl CO and HCO* (M-29) and not by the loss of HCN is common in the 2-pyridinones. 3-Hydroxy-2-pyridinonefragments w H2O (M-18) and this did not occur in any of the 3-hydroxy-4 The molecular radical cation (M*) is present in all the spectr result of the ability of the ring nitrogen to localize positive cha Hmpp, the M* at mz/ 125 is the base peak and mz/ 97 of HCO respectively. The mz/ 125 occurs in four of the five three; the prominence of the ring radical cation is due t dihydroxypyridinium moiety which is known to be favored in the CO and HCO* results in the mz/ 111 and 110 peaks in peaks in Hmepp. Hmepp can also fragment by cleavage of t hydrogen migration to give the ring molecular cation at mz/ 125; from this ion is the best explanation for the prominent peaks at this spectrum). This fragmentation mode is not operable with the Hdpp spectrum was without a mz/ 125 peak. The loss of carbonyl fragments from the M* is not observ Hmhpp or The N-hexyl group in Hmhpp fragments by methylene groups and again (by hydrogen migration) the ring rad peak. The bispyridinone fragments through loss of one ring 18  (M-28)  28  H2exn.  H2exn  34  followed by the same methylene cascade as in Hmhpp. The base peak is m/z 125 and both spectra have peaks at m/z 97 and 96. The least ambiguous spectroscopic evidence for the structure of H2exn is provided by its mass spectrum as the IR and proton N M R spectra afford little distinction between the mono- and bispyridinone.  Table 2.5. Mass spectral data (m/z) with the percent relative intensity in parenthesis.  Ligand  M*  MOO  M-HCO  Hmpp  125*  97 (9)  96(47)  Hdpp  139 (98)  110*  111 (25)  Hmepp  153 (74)  125 (91)  124 (28)  Hmhpp H2exn  Fragment Peaks  97 (58)  96*  209 (37)  125*  97 (14)  96 (20)  332 (15)  208 (64)  125*  97 (10)  * Base peak (relative intensity = 100%).  96(31)  36  B. Tris(3-Hydroxy4--Pyridinonato) Metal Complexes  2.5 Introduction With the exception of the conjugate bases of the 3-h are anionic bidentate ligands that can chelate a metal ion via t and the carbonyl oxygens. Each ligand would form a five-membered metal center and the tris-ligand metal complexes would be of neutr H2exn,  R  CH  3  F i g u r e 2 . 8 . F i v e m e m b e r e d c h e l a t e ring. W e refer to as a "potentially" tetradentate ligand becau bidentate moieties; however, the position of the oxygen atoms on the hexyl bridge make it virtually impossible for both rings to could chelate two metals to give a dimer as was similar diprotic tetradentate ligand containing two l-hydroxy-2-pyridinone rin L i g a n d s w i t h o x y g e n d o n o r a t o m s a r e u s u a l l y h a r d b a s e s . easily fit the definition of hard acids as acceptor atoms of high size. Therefore, the reaction of the 3-hydroxy-4-pyridinones with conform to the Hard and Soft Acid Base (HSAB) principle, H2exn  H2exn  M2L3  3 7 coordinate to hard bases.The HSAB principle is a generalization facts; however, a theoretical derivation has been developed for t inorganic chemistry.The derivation is based in part on the concept (I - A) (JI) w hich has an operational (and approximate) definition of ionization potential and A is electron affinity. This work has als definition of hardness as "the resistance of the chemical potential to of electrons". Using the TJ values, it is possible to rank metal cations by To demonstrate the utility of this, the rj values and the charge (C/IR) ratios (based on data from ref. 30) are compared for s metal ions in Table 2.6. (The left hand column is included f highest T| value, and Ca is a hard acid, Ag is a soft borderline in the earlier classification system). The C/IR ratio has the relative hardness of metal ions and the T| values allow a property. The first row transition metals (al high spin) have sim is significantly harder than its neighbors, as expected for its d el as indicated by the rj values. By either the TJ values or the c order of hardness for the group 13 metals is Al» Ga > In. Table 2.6. ComparisonCo/fIRr\ values(eV) aC n/ dIRC/IR ratio (A ) for B+ 111 Cr 4.88 9.1 Al 5.61 45.8 Ca+19.5 Mn 4.65 8.8 Ga 4.84 17 Cu 8.27 Fe 4.65 12.1 I+ n 3.75 13 Ag 6.96 Co 4.92 8.9 3.34 10.4 58  68  T|  +2  =-  +1  58  5  Tl  3  2  +2  Tl  +3  3+  3+  3+  +3  3+  1+  68  3  -1  38 Another measure of the relative acidity of a metal ion is constant of the hexaaquo species defined as: [M(H 0) ]  [M(H 0) OH]  3+  2  2+  6  2  pKa  5  +  H  +  Based on pKvalues of 2.6, 4.0, and 5.0,the order of The hardness of Al is exaggerated by both of the scales in Tab greatest in the absolute hardness value. Aluminum's large T] value fourth ionization potential which requires the removal of an electron f core. The ionization potentials of Gaand Inare based on the electron, respectively. Not surprisingly, the ionization potential for Al that of Gaor In (11,563 vs. 6150 and 5571 kJ mol").The relevance of the fourth ionization potential of Al was raised by t absolute hardness ranking,and it is clear that the T) values do relative acidity of the group 13 metals. This is in contrast to t transition metals where the ionization potentials are due to the removal from the same shell. The reason for the order of the hexaaquo pKvalues can b ionization potential of these metals. Instead of decreasing on going increases from 2744 to 2962 kJ mol", and the value for In of Al (2704 kJ mol"). The reason given for this is d-block co higher effective nuclear charge for Ga and, to a lesser degree, of the filed d shells do not completely shield the 10 added posi When this is contrasted to the shielding afforded by the Ne core of Al (or the enhanced acidity of Gaand In) is quite understa 52  a  +3  +3  3+  +3  +3  1 78  68  a  1  1  3+  +3  +3  39  Regardless of the scale that is used, however, the group 13 metals are hard Lewis acids and any aqueous synthetic chemistry with A l or Ga requires the consideration of their pH dependent speciation. The complexity of their aqueous behavior can be attributed to the affinity of the hard base water (r\ = 9.5) for these hard metal ions. The following equations stress the p H dependence of hydroxide formation.  Al(OH)  3  + 3 H+  "  [A1(H 0) ] 6  A l 3+ + 3 H 0 2  ^ = 1 0 . 0 ,  Kso, =  2  89  3+  = 10 10.7  From the final equation, 3 X 1 0 "  12  x  IQ - 3pH  M is the highest concentration of free A l  allowed by amorphous Al(OH)3 at p H 7.4. The practical consequence is that the typical synthetic concentrations (mM) are supersaturated in Al(OH)3 and a similar situation pertains with Ga. Unless a ligand has a high affinity for these metals, a gelatinous precipitate will form. Work in our laboratory with 3-hydroxy-4-pyrones had indicated that the a-hydroxyketone group was a strong enough binder of A l and G a t o preclude hydroxide formation at neutral p H and this was confirmed by the large overall formation constant P3 reported for Al(ma)3.  59  (For a definition of P3 refer to Section 5.7.)  The  3-hydroxy-4-pyrones are stronger acids than their pyridinone analogues because of the heterocyclic oxygen's greater capacity (compared to that of nitrogen) for stabilizing the negative charge of the deprotonated hydroxyl anion. The 3-hydroxy-4-pyridinone ligands used in this study have a hydroxyl p K ~1 p H unit higher than maltol. This increased a  basicity ensured there would be no hydroxide formation in the synthesis of the tris-ligand A l and Ga complexes.  4 0 There have been a number of aqueous solution studies of bidentate oxygen containing ligands, e.g., Al with salicylate ions carboxylic acidsand Ga with hydroxyaromatic ligands.Recent in biological and environmental roles of Al produced several studies on with the potentially tridentate citric acid.'Potentiometric studies indicate are good binders of these metals but do not afford any synthetic coordination chemistry is probably the best place to look for relevan high spin Feis close in size and electronic properties to Ga(s pertinent work was Raymond's synthesis and structural characterization o with l-hydroxy-2- and 3-hydroxy-2-pyridinones.Raymond has also synthesis of Ga complexes with catecholate and benzohydroxamate ligan Coordination chemistry with the 3-hydroxy-4-pyridinone ligands provid less fertile ground as regards synthetic antecedents. In the mid were touted as extractants for the separation ofGa and ^Zn rad same time the stability constants of mimosine (and several related number of divalent metal ions were determined'.Another study 29.2 for tris(mimosinato)aluminum(III), although their methodology was qu Several N-substituted-3-hydroxy-4-pyridinones have been evaluated as chelating the treatment of iron overload diseases and a large formation constan determined for Fe(dpp)3.This result was supported by the log P3 Fecomplex with 3,4-dihydroxypyridine. Based on our results with the 3-hydroxy-4-pyrones and the ab synthesis of Al and Ga complexes was attempted with the 3-hydroxyprepared in our laboratory and with L-mimosine (abbreviated as mimo included to ascertain what effect the amino acid substituent would h the tris-ligand metal complexes. It was also felt the M(mimo)3 co  09  19  29  3949  +3  +3  59  76  8799  001  +3  59  99  4 1 interesting biodistribution as mimosine is the 4-pyridinone analogue of t a compound that readily enters the brain (see Chapter I). Figure 2.9 is a diagrammatic representation of the tris(3-hydroxy metal complexes that were successfully isolated and characterized in thi M(dpp)3, and M(mhpp)3 have been reported by this laboratoryan of the synthesis of M(mepp)3 and Ga(mimo)3 of which we are 101  M ( m p p ) = CH, = H M(dpp) CH M ( m e p p ) M N CHCH M(mhpp) ( C H ) C H =H CHCH(NH)C0H M(mimo) Figure 2.9. Tris(N-substituted-3-hydroxy-4-pyridinonato) metal complexes. R i  3 3  R  3  J  3  3  3  2 3  R i  2  25 3  2  2  3  2.6 Material and Methods The preparations were the same for Al and Ga with any given for the Al preparations and only the reactant concentrations an the corresponding Ga preparations. The only difference of synthetic reduced aqueous solubility of the Ga complexes. For a given lig ~50% less soluble than its Al analogue and this necessitated only preparations.  42 Unless stated otherwise, the reported yields were for analyticall Mimosine was available commercially (Sigma, approx. 99%) and w purification. The other ligands were prepared and purified as describ 2.6.1 Tris(3-hyo^oxy-2-methyl-4-pyridinonato)aluminum(in). Al(mpp)^. Hmpp g, 10.0 mmol) and A1(NC«3)3»9H20 (1.24 g, 3.30 mmol) w The pH was raised to 8 with 2 M NaOH and the volum 70 °C. While the suspension was stil hot, the pink solid wa gave 1.06 g (83% yield). Mp 310 °C dec. 2.6.2 Tris(3-hydroxv-2-methvl-4-pvridinonato)gallium(III). Ga(mpp). 3 To Hm (1.09 g, 8.70 mmol) in 20 mL H0 was added 1.45 M Yield 1.13 g, 88%. Mp 290 °C dec. 2.6.3 Tris(3-hydroxv-1.2-dimethyl-4-pyridinonato)aluminum(III). Al(dpp)3. Hdpp (1.69 g, 12.2 mmol) and A1(N03)3«90 H (1.52 g, 4.05 mmo water. The pH was raised to 8 with 12 M and t heating at 70° C. The pale pink product was collected after gave 1.39 g (79% yield). Mp 300 °C dec. 2.6.4 Tris(3-hvdroxv-1.2-dimethvl-4-pvridinonato)gallium(III). Ga(dpp)3. To H (1.22 g, 8.80 mmol) in 20 mL water was added 1.45 Yield 1.10 g, 79%. Mp 280 °C dec. 2.6.5 Tris(3-hydroxv-2-methyl-l-hexyl-4-pvridinonato)aluminum(III). Al(mhpp)3. Hmhpp (0.86 g, 4.1 mmol) and AlCl3»6H20 (0.33g, 1.3 m (1:1) methanol/water. Deprotonation by the addition of 12 M by heating at 80 °C for one hour to remove excess ammonia ;  2  2  NH4OH  NH  43  was extracted into CH2CI2 (2 x 50 mL) and the solution was dried over MgS04. Vacuum distillation left a brown oil that was triturated with hexanes. Drying in vacuo (18 hours) gave 590 mg of a yellow powder (67% yield). Purification was by sublimation (300 °C, IO" torr). Mp 305 °C dec. 2  2.6.6  Tris(3-hvdroxv-2-methvl-l-hexvl—4-pvridinonato)gallium(III). Ga(mhpp)3.  To Hmhpp (0.63 g, 3.0 mmol) in 40 mL (1:1) methanol/water was added 1.45 M GaCl3 (0.69 mL, 1.0 mmol). Yield 0.50 g, 71% . Mp 285 °C dec. 2.6.7  Tris(3-hv(iroxv-2-memvl-l-emvl-4-pyridinonato)aluminum(III). Al(mepp)3.  Hmepp (422 mg, 2.89 mmol) and Al(NO3)3«9H 0 (360 mg, 0.96 mmol) were dissolved in 2  25 mL of water forming a pale pink solution of pH 2.2. The pH was raised to 7 by the slow addition of 2N NaOH and then the solution was heated at 70 °C for 30 minutes. The cloudy yellow solution was transferred to a liquid-liquid extractor and the product extracted into 75 mL CH2CI2 over 12 hours. The organic layer was removed in vacuo. The pink solid was washed with diethyl ether and gave 425 mg (92% yield). Mp 250 °C dec. 2.6.8  Tris(3-hvdroxv-2-methvl-l-ethvl-4-pvridinonato)gallium(III). Ga(mepp)3. To  Hmepp (669 mg, 4.37 mmol) dissolved in 15 mL of water was added 1 mL of 1.45 M Ga(Cl3) aqueous solution. Yield 652 mg, 85%. Mp 240 "C dec. 2.6.9  Tris(mimosinato)aluminum(III). Al(mimo)3. L-mimosine (591 mg, 2.98 mmol)  and A1(N03)3»9H20 (371 mg, 0.99 mmol) were dissolved in 25 mL of water forming an orange colored solution of pH 1.8. Concentrated N H 4 O H was slowly added to bring to pH 6.5 and the yellow solution was left stirring at 20 °C for 1 hour. Methanol was added until cloudy and the solution was cooled at 5 °C for 16 hours. The pink precipitate was collected, the filtrate was concentrated to 15 mL in vacuo, and the precipitation procedure was repeated a second time. The two precipitates were combined giving 510 mg (83% yield). M p 2 5 0 ° C d e c .  4 4 2.6.10 Tris(rrumosinato)gallium(IID hydrate. Ga(mimo)H 320. To L-mimosi mg, 2.99 mmol) dissolved in 25 mL of water was added 1. (0.69 mL, 1.00 mmol). Yield 571 mg, 86%. Mp 240 °C 2.7 Discussion of the Synthetic Procedure ,  The preparation of the tris(3-hydroxy-4-pyridinonato) Al and Ga relatively straightforward (Fig. 2.10). The 3:1 stoichiometry was closely formation of the tris-ligand metal complexes was thermodynamically favo necessary to use an excess of ligand to push the reaction to were in water except for M(mhpp)3; because of the lipophilic Nmethanol to water solution was used. To overcome purification pro the preparations, other solvents (methanol, ethanol, acetone) were t were obtained in water. The heating step was ostensibly include reaction but its primary utility became that of volume reduction to tris-ligand metal complexes. rawiistehtB oOHpH 8 A HL = 3-hydroxy-4-pyridinone ligand M = Al or Ga X = NO Figure 2.10. Equation for the formation of the tris-ligand met 3 HL  +  MX  3  •  ML  3  +  3BX  +  3  H 0 2  4 5 All of the complexes were prepared in good yield (70 to 9 any side reactions observed. The only synthetic complication was products from the salt formed in the neutralization process. This s either by solvent extraction or by fractional crystallization in those solubility properties of the complexes precluded simple extraction. M(mp M(mimo)3 were only soluble in protic solvents and they were crystallization. The initial preparations of Al(mpp)3 used AlCl3«6H20 with NaOH reaction mixture was concentrated (at 70 °C) until Al(mpp)3 prec advantage of the large difference in solubility between the product a (Al(mpp)3 was predicted to be soluble at the mM level). The filtered; the IR and proton NMR spectra of the precipitate wer Al(mpp)3 formulation and showed no proton containing impurities. analytical results were indicative of a non-carbon containing contaminan 54.18, found 46.93) and the likely cause was coprecipitation of was repeated using A1(N03)3 since NaN03 is twice as water s has a strong sharp IR band at 1380 cm"(a region free of is invisible in the IR. A simple experiment showed that NO3" was present in the easy purification method: the IR spectrum of the initial precipitate h cnrband but if the solution was concentrated further and filtered this band disappeared only to reappear when the solution was a before filtering. The hot (on the left) and the cold (on the spectra are reproduced as Figure 2.11 (the NO3" band at 1380 asterisk). The simple expedient of hot filtration gave analytically p yield. The Al(mpp)3 preparation set the pattern for the rest of 1  1  46 Al(NO>3 was used as the starting material to take advantage of the IR visibility of NO3". As the most convenient source of Ga was the chloride salt (obtained by dissolving the metal in HO), the reaction conditions were optimized for the Al preparation before being repeated with G a C l 3 .  » 1400  I  I  I  1200  I  1  1400  1200  Figure 2.11. The IR spectra of Al(mpp)3 from 1450 to 1200 cm' . 1  Al(dpp)3 did not precipitate from a hot solution (due to its high solubility temperature coefficient) but overnight cooling at 20 °C gave the pure product in good yield. To avoid problems with coprecipitation, the reaction mixture was more dilute than with Al(mpp)3 and N H 4 O H was used as the base ( N H 4 salts are more water soluble than Na +  +  salts). In the preparation of Al(mimo)3, fractional crystallization from hot water did not work nor was it possible to induce precipitation by cooling as the complex was too water soluble. Better results were obtained by using methanol as a second solvent; when precipitation was allowed to proceed slowly (16 hours), an analytically pure product was obtained.  47 Fractional crystallization gave low yields (50%) of Al(mepp)3. complex is not soluble to a measurable extent in any aprotic isolate pure product in yields > 85% by continuous liquid-liquid Al(mhpp)3 is quite lipophilic and forms an intractable g ummy preparation from was attempted using anhydrous organic solven line. The product was isolated in low yield (10%) but the complications inherent in a highly reactive starting material. Further ex A1(N03)3, alcohol-water mixtures, a nd product extraction with reasonable preparation. The yields were somewhat low (70%) and 0.08 torr) was needed to give samples suitable for analysis. Despite considerable effort, we were unable to isolate any m with the ligand. Reactions using a 3:2 stoichiometry invariab that were so insoluble as to defy characterization or purification. sample from an Al reaction did give an molecular ion but structural evidence to emerge from this work. The insolubility of an insurmountable problem. itself is only soluble in acidic solution; this is rather su is readily soluble in organic solvents and the other ligands all solubility. The anomalous behavior of the bispyridinone may be H-bonded dimeric units. The hydrophilic sites in the molecule (t moieties) would be involved in intermolecular H-bonds and th hydrophobic: the result would be water insolubility. The polar Hcomplex lipophobic: the result would be complete insolubility. A explain the insolubility of the metal complexation reaction precipitate, would have hydrophobic hexyl bridges and lipophobic metal centers. A1(CH3)3  CH2CI  H2exn  AI2L3  H.2exn  4 8 It is also possible for to polymerize by forming H-bonded dimeric units. The monopyridinones would be much less likely t may offer a better explanation for the higher melting point and (both near 300 °C), and for the decreased solubility of wh ligands. In the metal complexation reaction, oligimers could be for Whether due to dimerization or polymerization, the established insolubi the conjectured insolubility of its dimer is reasonable. 2.8 Characterization of the Tris(3-Hvdroxv-4-Pvridinonato') Metal Complexe H2exn  H2exn  M2L3  The IR and proton NMR spectra of the tris-ligand metal com changes from those of the free ligands. The changes are consis the deprotonated hydroxyl and the carbonyl oxygen atoms. The el molecular ions in the mass spectra are consistent with an ML3 the free ligands were analyzed in some detail; therefore, only t free and complexed ligand need be addressed in this section. Unlike the free ligands, the tris-ligand metal complexes are n EI-MS but rather require positive ion fast atom bombardment (FAB spectra were recorded with an AEI MS 9 and the samples tipped probe in either a glycerol or a thioglycerol matrix dependent exception, the instrumentation and conditions are the same as characterization of the free ligands (see Section 2.4).  49  2.8.1 Elemental Analysis The samples were submitted for analysis under N because complexes were hygroscopic. Rigorous drying was required to rem recrystallized products: heating in vacuo (0.05 torr) at > 65 °C Those samples that were sublimed were not subjected to further d 2  Table 2.7. Results of elemental analyses (Found [Calculated]) Compound Formula %H %N Al(mpp)3 [CiHAilN06] 5 3.89[54.13]4.69[4.55]10.60[10.52] Ga(mpp)3 [Ci8Hi8GaN 06]48.65[48.90]4.29[4.11] 9.53[9.51] Al(dpp)3 57.40[57.13]5.54[5.49] 9.67[9.52] Ga(dpp)3 [C2iH24GaN06] 5 1.98[52.09]5.10[5.01] 8.54[8.68] Al(mepp)3 59.72[59.61]6.25[6.27] 8.90[8.69] Ga(mepp)3»H20 [C24H32GaN0] 5 3.52[53.85]6.00[5.85] 7.72[7.85] Al(mimo)3 [C4HAlNOi2] 46.61[46.60]4 .60[4.41]13.52[13.59] Ga(mimo)3»H20 [C24H9GaNOi3] 42.18[42.43]4 .08[4.31]12.20[12.37] Al(mhpp)3 66.50[66.34]8.19[8.35] 6.34[6.45] Ga(mhpp) [C36H54GaN 0 ]62.42[62.25]7.94[7.85] 5.93[6.05] %C  8 8  3  3  [C21H24AIN3O6]  3 [C24H30AIN3O6]  3 7  2  7 2  6  6  2 [C36H54ALN3O6]  3  36  50  2.8.2 Infrared Spectroscopy IR studies on pyridine derivatives show that the substituen independently of the ring. The free ligand IR spectra were conclusion and predictably, the spectra of the tris-ligand metal corresponding free ligand are quite similar. The bands are broade general features of the spectra are the same. Also, the spectra with any given ligand are virtually identical above 800 cm(a cm"). There are, however, three diagnostic differences between the and complexed ligand: the appearance of three new bands belo and the bathochromic shift of Vc=0. The three new tentatively assigned as VM-0 but they are probably also coupled to modes.This is the only region of the spectrum where distincti in mass between the two metal ions are observable as at consistently at lower energy in the Ga complexes. Figure 2.12 spectra of Hmpp (top), Al(mpp3) (middle), and Ga(mpp3) (bottom) (the VMO - bands are marked with asterisks). The zwitterionic ami the JR spectrum of mimosine but the low energy region is cle utility in assigning the of the M(mimo3) complexes. The loss of the is the most noticeable change in M(mepp3 ), a nd M(mhpp) . The region above 3100 cnrhas onl 3450 cm"due to water. For M(mpp3), the band at 3280 c hypsochromic shift of this mode (at 2800 cm"in Hmpp) as 24  1-  1  VoHi  201  VM-O  1  V O H  1  3  1  5 1 intermolecular H-bonds. The assignment is not unusual for this m much sharper in the tris-ligand metal complex than in the free l TablMe(nW2 . 8 . A s s i g n m e n t o f V M O ( c m , K B r d i s k s ) M ( m r M ( d M ( m < M ( m i m o ) 3 ) 3 3 > P ) 3 1PP)3 5 P ) 3 Al Ga Al Al Ga Al Ga Al Ga 735 730 710 7 1 0 705 705 720 720 660655 b G a 640 625 575 570 580 565 580 570 620600 b 460365 b440 370 460 410 475 b340 b440<300 1-  The characteristic four band pattern of mixed Vc=o and formation of the tris(3-hydroxy-4-pyridinonato) metal complexes (Table most pronounced bathochromic shift occurring for the highest energy The superimposed spectra of Al(dpp(on top) and Hdpp (on ilustrate the changes typically observed. Table 2.9. Characteristic infrared absorptions (cnr). All are sharp an Assignment M(mpp)3 M(dpp)3 M(mepp)3 M(mhpp)3  V  3  1  Vc=0  and  1620  1605  1600  1605  1595  1560  1555  1560  1515  1515  1520  1495  1490  1490  VRing  1535 1505  sbh  52  54 2.8.2 Proton NMR Spectroscopy The spectral changes on formation of the tris(3-hydroxy-4-pyridi complexes are minor, but there is a diagnostic shift in the reso The resonances are closer together and the H signal is shifted sligh 15 Hz). The alkyl proton resonances are also shifted but in a Under the same conditions as described previously for the free li are absent in all the tris-ligand metal complexes. The NH signal i respectively for Al- and Ga(mpp)3 (in (CI>3)2SO at 80 MHz). An exchange process is occurring in the tris-ligand Al comp temperature; however, the chemical shift changes are small (from and are only observable on higherfieldstrength instruments. This where a difference between the metals is observed as only th observed for the Ga complexes at the ambient probe temperature (t not unexpected as Ga is known to be more labile than Al. NMR of the tris(3-hydroxy-4-pyridinonato) Al and Ga complexes Chapter IV.) The fluxionality of the Al complexes results in over chemical shift of the central signal is reported. The spectral assi M(mimo)3 complexes are listed in Table 2.11 and the peak integr the assignments. It is possible for mimosine to chelate via the amino acid ni although this was considered to be unlikely given the relative softn spectra are indicative of chelation by the a-hydroxyketone moiety, but amino acid bands in both the free and complexed ligand makes a  301  5 5 entirely binding at this site. The compounds are not soluble in not possible to look for the acidic protons by NMR; however, shifts of free and complexed mimosine in D2O (Table 2.10) sho the H and Hb resonances that were seen with the other li (mimosine is the only ligand in this study without a 2-methyl su so that it is between the AB doublets, rather than downfield complexes. The exocyclic proton signals, found as three multiplets that each integrate as one proton, are broadened but they did chelation. The similarity in the movement of the ring proton mimosine is chelating the metals in an fashion analogous to t pyridinone ligands employed in this study. Table 2.1R0e.corded*HatNM R c h e m i c a l s h i f t s ( 8 ) f o r m i m o s i n e a n d 400 MHz. a  M  I  Cr  ^CH CH(NH )C0 H 2  2  2  H  Assignment H(d) H(s) H (d) a  c  b  Mimosine 7.78 7.82 6.85  Al(mimo)3 Ga(mimo)3  7.73 7.46 6.72  7.72 7.54 6.77  56 Table 2.11. *H NMR chemical shifts (8) for the tris-ligand metal complexes (ppm). Recorded at 400 MHz.  n  M(mpp)  R = H  3  CH  M(dpp)  3  3 d  CH j CH  3 e  CH  2 e  2<  M(mpp)3  2 d  CH  M(mepp)  3  CH  M(dpp)3  2 f  CH  2 g  CH  2 h  CH  3i  M(mepp)3  Al  Ga  Al  Ga  Al  Ga  Ha(d)  7.50  7.48  7.56  7.49  7.53  H (d)  6.56  6.61  6.50  6.47  2.31  2.34  2.35  2.43  M(mhpp)  3  M(mhpp)3 Al  Ga  7.52  7.02  7.24  6.48  6.54  6.48  6.70  2.32  2.36  2.38  2.56  4.08  4.08 q  3.82  3.96 t  1.30*  1.30*  1.66  1.73 m  CH f,g,|»  1.27  1.30 m  CH  0.85  0.87 m  b  CH  3 c  CH  2 d  CH  2 e  (s  3.81*  3.80 s*  2  3i  * Read C H . 3  All recorded in D2O except: @ = CD3OD; # = CDCI3. D2O spectra externally referenced to DSS; spectra in other solvents are internally referenced. Abbreviations: s = singlet; d = doublet; q = quartet; t = triplet; m = multiplet.  57  2.8.3 Mass Spectrometry  The FAB-MS spectral data are listed in Table 2.12. In every case the molecular ion is found as the HML3 (HM ) peak, and the M L 2 base peak is formed by the loss of one +  +  +  ligand. For the Ga complexes, the peaks are in the natural isotopic ratio of 3:2 G a to 69  71  Ga. The presence of peaks due to the matrix make the assignment of the spectra beyond  this point profidess.  Table 2.12. Data from FAB-MS spectra of the tris-ligand metal complexes (m/z). M(rr PP)3  HML3  ML  + 2  4  M(n* 3>P)3  M(m W)3  M(m mo)3  M(di>P)3 Al Ga  Al  Ga  Al  Ga  Al  Ga  Al  Ga  400  444 442  442  486 484  484  528 526  652  695 693  619  662 660  275  319 317  303  347 345  331  375 373  443  487 485  421  465 463  58  Chapter III A.  Solid State Studies  3-Hydroxy-4-Pyridinone Crystal Structures  3.1 Introduction The structures of Hmpp, Hdpp, and Hmepp (Fig. 3.2) were established by single crystal X-ray diffraction. Apart from the obvious motive of structure confirmation, the crystallographic study was undertaken to establish the extent of double bond delocalization in these compounds, to determine if N-alkyl substitution increased the delocalization, and to discover what effect intermolecular H-bonding had on the crystal packing arrangement.  o  o'  o"  o"  R  R  R  R  o  Figure 3.1. 4-pyridinone resonance forms.  There are several possible resonance forms for 4-pyridinone  104  (Fig.3.1) and this is  of some importance when considering the potential stability of the tris(4-pyridinonato) metal complexes.  A significant contribution from the resonance forms with a partial  negative charge on the carbonyl oxygen might lead to complexes of increased stability as the bonding interactions with the hard acids A l and Ga are primarily electrostatic in nature.  105  Delocalization of the ring double bonds and a lengthening of the carbonyl bond  59  would be evidence of resonance hybridization in the 4-pyridinone ligands. It was thought that alkyl substitution at the ring nitrogen might enhance its ability to accommodate a positive charge thereby increasing the contribution from the pyridinium resonance form; the veracity of that supposition was examined by comparing the structures of Hmpp and its N-alkylated analogues. Intermolecular H-bonding had been indicated by the IR and proton N M R spectra and not surprisingly, the 3-hydroxy-4-pyridinones crystallized as dimeric H-bonded units. Our interest in this interaction is due to the potential connection between H-bonding and the solubility properties of the ligands and the metal-ligand complexes.  In addition, our  research group would eventually like to enter the field of inclusion chemistry and extensively H-bonded structures have potential for the synthesis of clathrate compounds. Knowledge of the packing arrangement could be useful in determining the suitability of Hmpp, Hdpp, and Hmepp as host molecules. The structures of Hmpp and Hdpp have been reported previously by this laboratory and this is the first report of the structure of Hmepp. The single crystals of 60  Hmpp were grown from a supersaturated water solution: a twice recrystallized sample of the ligand was dissolved in dilute acetic acid (pH 2), the p H was adjusted to 8 with 2 M NaOH, and overnight cooling at 10 °C gave crystals suitable for X-ray diffraction. Crystals of Hdpp could not be grown from water and were obtained by liquid-liquid diffusion from methanol with diethyl ether as the second solvent. Hmepp crystals were grown by slow evaporation from a dilute water solution. A l l of the crystal structures in this thesis were determined by Dr. Steven J. Rettig of the U.B.C. structural chemistry laboratory. In all cases, the crystals were stable under ambient atmosphere and were mounted on glass fibers for X-ray diffraction.  Figure 3.2. ORTEP view of Hmpp, Hdpp, and Hmepp.  61 Table 3.1. Bond lengths (A) and angles (deg) for the free ligands (bond lengths in the upper and bond angles in the lower portion of the table). In this and all subsequent listings of bond parameters, the estimated standard deviation (esd) is in parenthesis after the data entry. Compound  Atoms Hmpp  Hdpp  Hmepp  0(1)-C(3)  1.358 (2)  1.360 (3)  1.357 (3)  0(2)-C(4)  1.280 (2)  1.272 (3)  1.258 (3)  N - C(2)  1.356 (2)  1.369 (4)  1.378 (3)  C(2) - C(3)  1.371 (2)  1.376 (4)  1.370 (3)  C(3) - C(4)  1.431 (2)  1.430 (4)  1.433 (3)  C(4)-C(5)  1.411 (2)  1.407 (4)  1.410(3)  C(5) - C(6)  1.365 (2)  1.364 (4)  1.350 (4)  C(6) - N  1.345 (2)  1.352 (4)  1.351 (3)  C(2) - C ( l )  1.493 (2)  1.489 (4)  1.490 (4)  1.482 (4)  1.485 (3)  N - C(7)  1.484 (4)  C(7) - C(8) 0(1)-C(3)-C(2)  117.8 (2)  118.2 (3)  118.3 (2)  O(l) - C(3) - C(4)  120.4 (2)  119.2(2)  118.8 (2)  0(2)-C(4)-C(3)  120.0 (2)  120.5 (3)  121.3 (2)  0(2) - C(4) - C(5)  124,4 (2)  124.3 (3)  124.3 (2)  C(6) - N - C(2)  121.8 (2)  120.8 (2)  120.5 (2)  N - C(2) - C(3)  119.0(2)  118.6 (3)  118.6 (2)  C(2) - C(3) - C(4)  121.8 (2)  122.5 (3)  122.9 (2)  C(3) - C(4) - C(5)  115.5 (2)  115.2 (3)  114.3 (2)  C(4) - C(5) - C(6)  120.7 (2)  121.1 (3)  121.9 (2)  C(5) - C(6) - N  121.1 (2)  121.7 (3)  121.8 (2)  N - C(2) - C ( l )  118.3 (2)  119.0 (3)  120.0 (2)  C ( l ) - C(2) - C(3)  122.7 (2)  122.3 (3)  121.5 (2)  C(6) - N - C(7)  118.4(3)  117.3(2)  C(2) - N - C(7)  120.7 (3)  122.0 (2)  N - C(7) - C(8)  111.8 (2)  62  3.2 Results and Discussion  The six-membered pyridinone rings are slightly non-planar and the maximum deviations from the mean planes are 0.019(2) A in Hmpp, 0.009(3) A in Hdpp, and 0.013(3) A in Hmepp; the distortions from planarity are toward a C(4) envelope, a C(2)C(5) boat, and a N-C(4) boat respectively. A n examination of the bond angles (Table 3.1) shows only minor differences between the three structures. N-alkylation results in a smaller intra-annular bond angle at nitrogen (C(2)-N-C(6)), and the C(3)-C(4)-C(5) angle is compressed in all three compounds to a minimum of 114.3° in Hmepp.  This  compression is indicative of the strength of the C=0 bond (shortest in Hmepp) and is due in part to the lone pairs of electrons on the oxygen.  The bulk of the N-alkyl groups  predictably affects the position of the ring methyl group (C(l)) and this increases the N-C(2)-C(l) angle from 118.3° in Hmpp to 120° in Hmepp. The small changes in the bond angles involving the oxygen atoms are due to the packing of the H-bonded dimeric units. The nitrogen coordination is planar within experimental error in all three compounds and the bond lengths indicate partial delocalization of the formal double bonds.  By  comparing the observed bond lengths with the values calculated (using the empirical relationship between bond order and bond length) for the double bond character consonant with the probable resonance forms, it is possible to relate double bond delocalization (as determined crystallographically) to the approximate contributions from each resonance hybrid.  106  Calculations of this nature were done for 2-pyridinone, a structural isomer of  4-pyridinone that is capable of similar resonance interactions involving the nitrogen lone pair of electrons.  80  The formula used was  R =•—^  1  f ) ^ * ^*  w n e r e  ^  1 S  t n e  observed bond length, R i and R are the standard carbon single- and double bond lengths 2  63  and x is the bond order expressed as the fraction of double bond character.  The  combination of 2-pyridinone resonance forms that gave the best agreement with the observed bond lengths are shown in Figure 3.3.  O + N  50  15  o  o  20  0  0  • A 10  5  Figure 3.3. Resonance forms for 2-pyridinone (percent contribution is in bold type).  Analogous calculations were not done for the 3-hydroxy-4-pyridinones as a comparison of the bond lengths listed in Table 3.2 clearly indicates the similarity in the degree of double bond delocalization between Hmpp and 2-pyridinone. (Compound C was included to show that N-alkyl and hydroxy substitution did not appreciably alter bond delocalization in the 2-pyridinone ring.) The carbonyl bond is significantly longer in Hmpp and the similarity in the ring bond lengths is readily apparent; additionally, the mean of the six ring bonds is the same and Hmpp even has a slightly smaller standard deviation (a). This indicates there is as much delocalization in Hmpp and suggests there may be as significant a contribution from the pyridinium resonance forms for Hmpp as was seen in 2-pyridinone (35%).  Contrary to our expectations, N-alkylation did not increase the  amount of double bond delocahzation in the 4-pyridinones. Compared to Hmpp, there was no significant change in Hdpp bond lengths and the changes in Hmepp were indicative of a decrease rather than an increase as the C(5)=C(6) and C=0 bonds were shorter and the C - N bonds were longer than in Hmpp (see Table 3.1).  64 Table 3.2. Comparison of bond lengths (A). (The ring carbons are numbered as in Hmpp.) O 3  .OH  ^OH A  2  "CH  J  x N ^ O  3  Hmpp  2-Pyridinone  Bond  #  I  n- Bu 3-Hydroxy-1 -butyl-2-pyridinone  Compound A  B#  c@  c=o  • 1.28  1.24  1.24  C(2)-C(3)  1.37*  1.44  1.43  C(3)-C(4)  1.43  1.33*  1.35*  C(4)-C(5)  1.41  1.42  1.41  C(5)-C(6)  1.37*  1.37*  1.33*  N-C(2)  1.36  1.40  1.37  N-C(6)  1.34  1.34  1.38  Mean (a)  1.38 (0.033)  1.38 (0.044)  1.38 (0.038)  Data from ref. 80.  @ Data from ref. 107. The 3-hydroxy-4-pyridinones crystallized as centrosymmetric  0-H-0=C  hydrogen bonded dimeric units. (The dimeric unit from Hdpp is included as Figure 3.4 and the packing arrangements of Hmpp, Hdpp, and Hmepp are given as Figures 3.5, 3.6, and 3.7 respectively.)  The dimeric units of Hdpp and Hmepp are separated from one  another by normal van der Waals distances but in Hmpp each dimer is linked to four others by N-H—0=C hydrogen bonds to form a three-dimensional network. The strength of the  65  H-bonds can best be judged by examining both the H-bond parameters and the IR stretching frequencies (Table 3.3). In all three structures, the O - H - O bond lengths are intermediate (<2.70 A  1 0 8  ) and the angles are within the range typical for H-bonds (140-  180*109). in Hmpp, the N - H - O distance and VNH indicate relatively strong nitrogen H-bonds;  108  the constraints imposed by the N H hydrogen bonding causes a weakening of  the O - H - O bonds as indicated by the bond angle of 144* and by the VOH of 3270 cm-  1  (VOH > 3200 c m is classified as weak H-bonding ). A comparison of VOH in the three -1  108  ligands supports the proposition that the weakening of the bonds between the dimeric units in Hmpp is due to the formation of the three-dimensional network.  Figure 3.4 The Hdpp hydrogen bonded dimeric unit.  Table 3.3. A comparison of the free ligand H-bond parameters and the IR stretching frequencies. Values involving the nitrogen in Hmpp are in italics. Compound  Atoms Hmpp  Hdpp  Hmepp  H • O (A)  1.90(3) 1.90(2)  1.94 (5)  1.83(4)  (N)O-O (A)  2.670 (2) 2.796(2)  2.692 (3)  2.659 (2)  154 (3)  150 (2)  3150  3180  ( J V j O - H - O (deg)  Vw)OH (cm- ) 1  144(3) 166(2)  3270  2800  66  In the packing of the dimeric units, the alkyl groups are staggered to allow the closest spacing of the pyridinone ring;, and this is most readily seen in the unit cell of Hdpp (see Fig. 3.6). Comparing Hdpp to Hmepp shows that the N-ethyl group disrupts the stacking of the dimeric units; this is reflected in an increase in the volume of the unit cell (both compounds crystallized in the orthorhombic space group* Pbca ) from 1312 A in 3  Hdpp to 1628 A in Hmepp. The bulkier ethyl group also results in an appreciably lower 3  calculated density for Hmepp (1.14 g/cm ) than either Hmpp or Hdpp (1.42 and 1.41 3  g/cm respectively). Despite the similarities in the strength of the H-bonds, the Hmepp 3  crystal lattice appears to have been weakened due to the steric requirements of the N-ethyl group. This conclusion is supported by the lower melting point (205 vs. 260 *C) and the greater water solubility (five-fold increase at 25 *C) of Hmepp when compared to Hdpp.  Figure 3.5. Stereoview of the packing arrangement in Hmpp.  * The crystallographic data (including the unit cell dimensions) for Hmpp, Hdpp, and Hmepp are in the Appendix as Table A.2.  67  Figure 3.6. Stereoview of the packing arrangement in Hdpp.  68  Figure 3.7. Stereoview of the packing arrangement in Hmepp.  69  B.  M(dpp)3 Crystal Structures  3.3 Introduction The structures of A1-, Ga-, and In(dpp) were determined; the facial (fac) isomers 3  crystallize in the trigonal space group* P3 and the empirical formula is M(dpp)3«12H20. The threefold symmetry of the fac isomers results in an asymmetric unit consisting of 1/3 of a metal ion, one ligand, and four water molecules. The water molecules form hexagonal rings that are H-bonded to the chelating oxygen atoms of the complexes by bridging waters. The water molecules do not coordinate the metal; rather they are structural waters  110  which form a three dimensional framework that is an important factor in  determining the structure of the inorganic complex. The unit cell diagram (Figs. 3.8 and 3.9) accentuates the spatial relationship of the water channels to the tris(pyridinonato)metal portion of the structure (/ac-M(dpp)3 unit). The view perpendicular to the symmetry axes of the water channels and the/ac-M(dpp)3 units in Figure 3.10 affords the best perspective of the entire water network. Because of the uniqueness and the structural importance of the water network, the Results and Discussion will be divided into two parts: the /ac-M(dpp)3 unit in Section 3.4.1 and the hexagonal water network in Section 3.4.2. The A l and Ga complexes are isostructural but due to a slight variation in the packing arrangement, the In complex is not crystallographically equivalent; however, the/ac-M(dpp)3 units of the three complexes are isostructural (Fig. 3.11) and the bonding parameters will be examined together. Despite  * The crystallographic data for the M(dpp)3 complexes are in the Appendix as Table A.3.  70  the differences in packing, the water network is the same in all three compounds and will be addressed as such. The single crystals were grown from supersaturated water solutions: an aqueous suspension was heated to 80 °C, filtered, and crystals formed after several days at 20 °C. The crystals were stable under ambient atmosphere and the integrity of the water network was exemplified by Al(dpp) which successfully analyzed for 12 waters. The crystals 3  were collected by filtration, dried overnight in a desiccator, and submitted for analysis under nitrogen: expected (found) C 38.35 (38.15); H 7.37 (7.39); N 6.39 (6.31). The strength of crystal lattice accounted for the rigorous conditions necessary to isolate the anhydrous compounds (see Section 2.8.1). To determine the importance of the lattice waters to the geometry of the/ac-M(dpp) units, numerous attempts were made to grow 3  crystals from other solvents, either by liquid-liquid diffusion or by slow evaporation. The latter technique gave In(dpp) crystals from 95% ethanol, but the complex again formed as 3  the dodecahydrate. This was the only system that produced crystals suitable for X-ray diffraction. The A l - and Ga(dpp) structures were reported previously by 3  laboratory. 101  111  this  The synthesis of In(dpp) was analogous to that of its congeners and 3  the structure was reported in a study of In complexes.  Although my doctoral research  112  did not involve the synthesis of indium complexes, the synthetic procedures were based on my work with the tris(3-hydroxy-4-pyridinone) A l and Ga complexes. I was directly involved in the growing of single crystals of In(dpp) and the similarity of the three 3  structures motivated the inclusion of the In complex in this thesis. A number of In complexes were also part of the octanol/water partition coefficient study that is included in Chapter V .  71  Figure 3.8. O R T E P view down the c axis of the unit cell packing of the M(dpp)3 complexes ( M = A l and Ga).  72  73  i  *  >4  Figure 3.10. ORTEP view of a part of the H-bonding network of waters inM(dpp)3. This view down the a axis shows all of the independent O atoms; all of the atoms in the M(dpp)3 units are omitted except for the M 0 6 octahedral coordination sphere.  74 3.4.1  The fac-M(dpph Unit  The ligand rings are planar within experimental error. There are no significant differences in the ring bond lengths between the A l and Ga complexes and only minor changes are seen in comparison to the In complex (Table 3.4). For In(dpp)3, one C - N bond is longer but the other is shorter and the same relationship holds for the formal double bonds (C(2)-C(3) and C(5)-C(6)); this produces no net change in double bond delocalization. Comparison to the free ligand reveals only minor changes and the extent of ring double bond delocalization is essentially unaffected by metal chelation. The delocalization in the C - 0 bonds is greater in the metal complexes than in the free ligand: the difference between 0(1)-C(3) and 0(2)-C(4) decreases from 0.088 A in Hdpp to 0.028 A, 0.038 A, and 0.054 A in A1-, Ga-, and In(dpp) respectively. The 3  extent of delocalization is clearly seen by comparing bond lengths with the related dihydroxy ligands, the catechols. The average C-O distance (calculated from a number of different metal complexes) is 1.36(1) A for the catecholate anion and 1.29(1) A for the delocalized semiquinone radical anion.  32  For the dpp anions in our metal complexes, the  0(1)-C(3) bonds are signifcantly shorter (mean = 1.337(8) A) than the catecholate bond and the 0(2)-C(4) bonds are the same (mean = 1.297(8) A) as the semiquinone bond (within the esd). The only significant variations in ligand bond angles occur in the carbon atoms (C(3) and C(4)) that are part of the chelate ring. Chelation causes a compression in the interior angles of the chelate ring: compared to the free ligand values, the 0(1)-C(3)-C(4) and 0(2)-C(4)-C(3) bond angles are decreased in the A l and Ga complexes and are the same in the In complex. The smaller the metal the greater the compression. Further evidence of this is seen in (C6Hs)2B(dpp) where these angles for the dpp" anion are further compressed to 111.8(4)° and 113.6(3)° respectively.  113  Figure 3.11. ORTEP view of the tris(ligand) portion of the M(dpp>3 complexes ( M = Al, Ga, and In).  76 Table 3.4. Bonding parameters for the M(dpp)3 12H20 complexes. The bond lengths (A) are in the upper portion and the bond angles (deg) are in the lower portion of the table. For comparison, the parameters for the Hdpp ligand have been included in the last column. #  M(dpp)3  Atoms M=  Al  Ga  In  Hdpp  0(1)-C(3)  1.327(3)  1.342 (5)  1.343 (3)  1.360 (3)  CK2)-C(4)  1299 (3)  1.304 (5)  1.289 (3)  1.272 (3)  N - C(2)  1.369 (3)  1.372 (5)  1.355 (3)  1.369 (4)  C(2) - C(3)  1.385 (3)  1.382 (6)  1.393 (3)  1.376 (4)  C(3)-C(4)  1.423 (3)  1.409 (6)  1.403 (3)  1.430 (4)  C(4)-C(5)  1.398 (3)  1.403(6)  1.410 (3)  1.407 (4)  C(5) ~ C(6)  1.360 (3)  1.359 (6)  1.344 (4)  1.364 (4)  C(6) - N  1.360 (3)  1.349 (5)  1.371 (3)  1.352 (4)  C(2)-C(l) N - C(7)  1.493 (3)  1.486 (7)  1.487 (4)  1.489 (4)  1.476 (3)  1.470 (6)  1.463 (3)  1.482 (4)  0(1)-C(3)-C(2)  124.5 (2)  122.4 (4)  120.1 (2)  118.2 (3)  CXI) - C(3) - C(4)  115.3 (2)  116.9(4)  119.1 (2)  119.2 (2)  0(2)-C(4)-C(3)  116.0(2)  117.9(4)  120.0 (2)  120.5 (3)  0(2)-C(4)-C(5)  125.7 (2)  124.2 (4)  123.0 (2)  124.3 (3)  C(6) - N - C(2)  121.2 (2)  121.1 (4)  120.2(2)  120.8 (2)  N - C(2) - C(3)  119.0(2)  118.9(4)  119.7(2)  118.6 (3)  C(2)-C(3)-C(4)  120.2 (2)  120.7 (4)  120.8(2)  122.5 (3)  C(3)-C(4)-C(5)  118.4(2)  117.9 (4)  116.9(2)  115.2(3)  C(4)-C(5)-C(6)  119.5(2)  119.7(4)  120.8 (2)  121.1 (3)  C(5) - C(6) - N  121.7 (2)  121.7 (4)  121.6 (2)  121.7 (3)  C(6)-N-C(7) C(2) - N - C(7)  117.7(2)  117.7(4)  117.0(2)  118.4 (3)  121.1 (2)  121.1 (4)  122.8 (2)  120.7 (3)  N - C(2) - C(l)  119.2 (2)  119.4(4)  118.8 (2)  119.0(3)  C(l) - C(2) - C(3)  121.9 (2)  121.7 (4)  121.4 (2)  122.3 (3)  77  Table 3.5. Intra-annular torsion angles of the chelate rings in the M(dpp)3 complexes (deg). Atoms M = Al  Ga  In  0(2)-M-0(l)-C(3)  - 5.05 (14)  - 5.4 (3)  7.47 (14)  M-0(l)-C(3)-C(4)  4.7 (2)  5.0 (4)  - 7.2 (2)  0(l)-C(3)-C(4)-0(2)  - 1.0 (3)  - 0.8 (6)  0.8 (3)  M-0(2)-C(4)-C(3)  - 3.1 (2)  - 3.8 (5)  6.0 (3)  0(l)-M-0(2)-C(4)  4.50 (15)  5.0 (3)  - 7.21 (15)  The increasing size of the metal ion is accompanied by a reduction in ligand strain as evinced by compression of the O-C-C bond angles. By this criterion, In best fits the dpp" anion since it caused the least deviation from the bond angles in the free ligand. This is a tentative conclusion as the changes are small and crystal packing forces could be a factor in the observed variations. The planarity of the chelate rings (Table 3.5) varies in the opposite direction. The metal ion is shifted out of the plane and the deviation from planarity is the greatest for In(dpp)3. There is a compression of the M(dpp)3 units from ideality along the C3 axis leading to a decrease in the 0(l)-M-0(2) angles and an increase in the exocyclic 0(l)-M-0(2') angles in all the structures (Table 3.6). The exocyclic 0(1)-M-0(1') and 0(2)-M-0(2') angles are 90°(±1°) in the A l and Ga complexes, but they increase to 93.58 and 92.40 A respectively in the In complex The buckling of the chelate ring and the compression along the C 3 axis increase directly with the ionic radius of the metal center.  A n examination of the packing  arrangement (see Figs. 3.8 and 3.10) shows that along the c axis (parallel to C3) the  78  distance between fac-M(dpp)  units is governed largely by a single bridging 0(3) water.  3  The length of the H-bonds from 0(3) to the ligand oxygens does not vary significantly among the three structures. If the rigid water network will not expand, an increase in the size of the metal center must result in compression along the C3 axis and a twisting of the chelate plane.  Table 3.6. Bond angles (deg) for the metal-ligand interactions in M(dpp) complexes. 3  Atoms  Al  Ga  In  O(l) - M - 0(2)  84.23 (6)  83.22 (12)  77.87 (6)  0(1)-M-0(1)'  90.81 (8)  90.90 (12)  93.58 (6)  0(l)-M-0(2)'  95.71 (7)  96.65 (12)  97.74 (7)  171.85 (7)  170.48 (12)  166.18 (6)  0(2) - M - 0(2)'  89.83 (8)  90.01 (12)  92.40 (6)  M - O(l) - C(3)  112.20 (13)  110.7 (2)  110.82 (13)  M - 0(2) - C(4)  112.00 (14)  110.9 (3)  111.39 (14)  0(2)-M-0(l)'  M=  t r a  n  S  The M - O bond lengths are listed in Table 3.7; the differences in the M - O bond lengths and the differences in metal ionic radii are included in italics. Comparison of the differences shows that the M - O bond lengths are reasonably close to the values predicted by the differences in the size of the metal. The Al-0(2) bond shows the most deviation and this slight increase from the predicted length suggests the metal-oxygen bonding is weakest for Al(dpp) . The strength of the Ga-0 bonds is seen by comparison to the 1.986(6) A 3  average Ga-O distance in K [tris(catecholato)Ga(JJi)]. The dative bond (Ga-0(2)) is the 3  34  same length as the catecholate bond (within esds) and Ga-O(l) is significantly shorter. The  79  relative shortness of Ga-0(2) is further evidence of a significant contribution from the pyridinium resonance forms with a partial negative charge on 0(2). The length of the Ga-0 bonds also compares well to the bonding in tris(3-hydroxy-l-butyl-2-pyridinonato)iron (IH);  107  when adjusted for the differences in metal ionic radii (a difference of 0.025 A), the  M - 0 distances are 2.024 and 1.955 A.  Table 3.7. Metal ionic r a d i i and M-O bond lengths in M(dpp)3 complexes (A). 30  3  M= Al  a  In  Ga  Ionic Radius  0.535  0.085  0.620  0.180  M - 0(1)  1.893 (2)  0.074  1.967 (3)  0.167  2.134 (2)  M - 0(2)  1.923 (2)  0.067  1.990 (3)  0.175  2.165 (2)  0.800  Going along a row, the differences between entries in adjoining columns are in italics.  The structural differences between the free and the complexed ligand are minor and the changes are indicative of increased double bond delocalization in the metal-ligand complexes.  The/ac-M(dpp)3 units are isostructural and the variations between the  structures are readily rationalized by consideration of the metal ionic radii and the rigidity of the packing arrangement. The metal-oxygen interactions are strong and compare favorably to those of other oxygen containing bidentate ligands known to be good chelators of trivalent metals.  80 3.4.2 The Hydrogen Bonded Water Network  The water molecules form a three dimensional array: half (H20(3) and H2CX6)) form a bridge from the/ac-M(dpp)3 units to the hexagonal channels formed by the other half of the water molecules (H 0(4) and H20(5)) in the corners of the unit cell (see Figure 2  3.10 for numbering). According to the classification of Falk and K n o p ,  1 1 4  in which  waters are designated by the number and type of hydrogen bonding water neighbors (a is a proton acceptor and d is a proton donor), H20(3) is a, H2CX6) dda, and the waters in the hexagonal channels (H20(4) and H20(5)) are ddaa.  Table 3.8. H-bond distances (A) and angles for M(dpp)3»12H20.  Interaction M=  a  o-  H- O Al  O - H • O (deg)  •O Ga  Al  Ga  Al  Ga  0(3)-H(a> 0(1)  1.95(6)  1.98(9)  2.861(3)  2.859(5)  161(4)  165(8)  0(3)-H(b)-0(2)  2.08(4)  2.19(6)  2.849(3)  2.842(6)  164(4)  159(7)  0(4)-H(a)-0(6)  2.12(5)  2.19(6)  2.772(3)  2.765(6)  159(5)  164(10)  0(4)-H(b)-0(4)  1.62(8)  1.58(12)  2.747(3)  2.746(5)  168(5)  154(8)  0(5)-H(a)-0(4)  2.06(10)  2.44(12)  2.802(4)  2.807(7)  150(9)  145(22)  0(5)-H(c)-0(5)  a  1.50(13)  1.53(24)  2.793(4)  2.779(7)  174(6)  160(10)  0(5)-H(d)-0(5)  2.17(13)  0(6)-H(a)-0(3)  1.72(4)  1.99(8)  2.729(4)  2.734(7)  174(3)  167(8)  0(6)-H(b)-0(5)  2.00(4)  1.91(7)  2.791(4)  2.778(7)  174(3)  171(6)  149(12)  2.793(4)  This interaction involves H(05b) for the Ga compound.  81  The H-bonds* (Table 3.8) between the ligand O atoms (O(l) and 0(2)) and H20(3) are relatively strong considering these oxygen atoms are chelating the metal atoms. A l l six of the chelating O atoms are hydrogen bonded to 0(3) water molecules; the latter form an infinite chain down the c axis bridging from the hydroxy 0(1) in one ligand to the keto 0(2) of a ligand rotated by 120° and translated by one unit cell. (This is why Figure 3.8 shows the contents of two unit cells). The O - H and O - O distances for the chelating oxygens vary from 1.95(6) to 2.19(6) A and from 2.842(6) to 2.861(3) A respectively. Not surprisingly, the chelating hydroxy O forms shorter O - H bonds than the chelating keto O. The hexagonal channels of water molecules in the corners of the unit cell are the most unique feature in these structures.  (The water 0(5) was found to be twofold  disordered in the A l complex and is shown with four half protons bound to it in Figure 3.10.) Each of the water rings has crystallographically imposed 3 or S6 symmetry and the rings essentially adopt the structure of i c e  1 1 5  in its stable low pressure form, ice-I . (Refer n  to Fig.3.9 for a comparison of the water rings and the structure ice.) Every water molecule in the ring is hydrogen bonded to four nearest neighbors with the added distinction that the overall structure is predominantly proton-ordered, unlike i c e - I which is completely n  disordered. The O - O distances in the channels (those involving 0(4), 0(5), and 0(6)) vary from 2.75 to 2.81 A compared with the average value of 2.75 A for O - O in ice-I at n  100K.U6 Within each water ring the hydrogen bonding is h o m o d r o m i c ' 117  119  because of the  unidirectional circular bonding pattern. A l l the O - H - O bonds run in a counterclockwise direction when viewed down the hexagonal axis (see Figure 3.9). The arrangement is crystallographically imposed by the space group (P3), however, and does not occur * There were no significant differences in the H-bond parameters for In(dpp)3 so they have been placed in the Appendix as Table A.5  82  independent of symmetry constraints as do the water networks in the structures of some nucleosides  120  and P-cyclodextrins. ' 121  122  Homodromic hydrogen bonding arrangements  are favored (and more frequently observed despite symmetry constraints) over heterodromic or antidromic arrays because of an inherent lower dipole moment.  118  There  is a considerable cooperative effect which results in increased hydrogen bonding activity for a hydroxyl group when it is already the donor or acceptor in a hydrogen bond. Quantum mechanical calculations have confirmed that chain structures (particularly cyclic) of hydrogen bonds are energetically favored over individual interactions.  123  This is the first example where this arrangement of water rings occurs in hydrates containing (relatively) large metal complexes; there has recently been reported a tris(2-pyridinonato) iron complex that contains hexagonal water rings, but the rings are discrete and there is no three-dimensional water network. well known in the crystal structures of i c e  1 1 5  -  1 1 6  107  Water rings are, however,  and the clathrate h y d r a t e s  124  '  125  The  clathrate hydrates are crystalline compounds which consist of a hydrogen bonded water host network (often a H40O20 pentagonal dodecahedron) within which a guest is held by an interaction which varies from weakly hydrogen bonding, to ionically bonding where one or more ion is associated with, or incorporated in, the water framework.  126  We note  similarities here with hexamethylenetetramine hexahydrate ((CH2)6N4*6H 0), an unusual 2  hydrate in which the host lattice is not based on a regular polyhedron, so there are no well defined polyhedral cavities.  127  It shares with the structures reported herein the hexagonal  water rings; however, the water rings are staggered around the cage-like amine molecules in a spiral (instead of a linear chain) to which the latter are bound. The above observations suggest that the M(dpp) complexes present (when 3  crystallized from water) appropriate conditions for the formation of the water channels in what is a previously unobserved, but energetically favored, hydrogen bonding arrangement. This probably results by virtue of both the complex size and the hydrophobic  83  core of the unit cell formed by the pairs of methyl groups on the ligands from the two /ac-M(dpp)3 units. There is a clear alternation of hydrophilic and hydrophobic regions along the ab diagonals of the unit cell; no doubt this feature also contributes to the unique hydrogen bonding arrangement. It may be that a driving force for the formation of the dodecahydrates is a variation on hydrophobic c l a t h r a t i o n » 124  cooperativity  109  125  as well as the  of H-bonding networks.  C.  M(mepp)3 Crystal Structures  3.5 Introduction  The importance of the H-bonded water network to the M(dpp) structure was 3  stressed in the preceding sections; because of the constraints inherent in such a rigid structure, it was felt that even a small variation in the ligand might be sufficient to alter the nature of the water network or to preclude its formation altogether.  In the packing  arrangement of tris(3-hydroxy-l-butyl-2-pyridinonato)iron(III) there are hexagonal water rings H-bonded to the chelating oxygens.  107  There are no bridging waters and no water  channels: one water ring is sandwiched between two metal complexes (forming discrete units) and alternate waters in the ring are H-bonded to the chelating hydroxyl oxygens (the carbonyl oxygens do not act as H-bond acceptors) of one of the metal complexes. As soon as the first M(dpp)3 structure was determined, we wondered if the water network was  84  unique to the Hdpp ligand; this very interesting Fe structure piqu water arrangements that might be possible with the other 3-hydroxy4 Efforts to grow more crystals culminated in the determination Al- and Ga(mepp)3. The crystal growing attempts with the other with litle success and the M(mepp)3 complexes were synthesized structural study. The crystals were grown by slow evaporation (o from dilute water solutions. Crystal growth was much slower t complexes and the single crystals were considerably smaller. The with an ethyl group is, of course, a very minor structural c change was sufficient to alter significantly the packing arrangement of it was thought a similar outcome might be possible with the tri When the Al(mepp)3 crystals analyzed as a decahydrate (the prepared analogously to the previously mentioned Al(dpp)3 sample), were dealing with a similar structure. The detennination of the cry the water network was able to accommodate the N-ethyl group: Al(mepp)3«12H20, the space group* was trigonal P3, and the unchanged.** The Ga(mepp)3 crystal structure showed that a change conceit with N-ethyl substitution was also insufficient to alter the and Ga complexes were essentially isostructural; the modifier "essentiall by the same crystallographic inequality that distinguished the In(dpp)3 and Ga analogues. For this reason, the In(dpp)3 packing arrangem previously and the nature of the inequality will be examined by a packing in the Al- and Ga(mepp)3 complexes. * The crystallographic data for the M(mepp)3 complexes are in the Appendix as Table A.4. ** Because there was no significant differences from the water network in the M(dpp)3 structures, the M(mepp)3 H-bond parameters are in the Appendix as Table A.6.  85  3.6 Results and Discussion  The/ac-M(mepp)3 units of the A l and Ga complexes are isostructural (Fig. 3.12), as is readily apparent by a comparison of the bond lengths and angles in Table 3.9. The C(5)-C(6) bond is significandy longer in the A l complex (by 0.02 A) and that is the only difference in ring bond lengths. There appears to be an increase in delocalization compared to the free ligand but this is primarily due to change in only one bond, C(5)-C(6), that was the shortest for any of the structures solved (free ligands and metal complexes) in Hmepp and was the longest in Al(mepp)3~ 1.350 vs. 1.380 A. There is more delocalization in the C - 0 bonds than in either Hmepp or the M(dpp)3 complexes. The difference between the two C-O bonds is 0.099 A in Hmepp, and 0.020 and 0.024 A in A l - and Ga(mepp)3 respectively.  This leads to the same conclusion as was reached with the M(dpp)3  complexes; i.e., an increase in double bond delocalization occurs upon chelation. Aside from the obvious difference in the N-alkyl groups, the /<ac-M(dpp)3 and /ac-M(mepp)3 units are isostructural. The ligand bond lengths and angles are virtually the same (compare Tables 3.4 and 3.9). The chelate ring angles show a similar compression along the C3 axis and the M - O bond lengths are comparable. In the water network, some of the H-bonds are slightly longer but, as was the case with the increased size of the metal ion, there is no evidence from the H-bond parameters that the water network is significantly affected by the additional methylene group in the M(mepp)3 complexes.  86  Figure 3.12. ORTEP view of the tris(ligand) portion of the M(mepp)3 complexes.  87 Table 3.9. Bond parameters for M(mepp)3«12H20 complexes. The bond lengths (A) are in the upper portion and bond angles (deg) are in the lower portion. For comparison, the parameters for the Hmepp ligand are in the last column. Atoms Al  Ga  M-O(l) M-0(2)  1.894 (1) 1.930 (1)  1.962 (1) 2.00 (1)  0(1) - C(3) 0(2) -C(4)  1.317 (2) 1.297 (2)  1.327 (2) 1.303 (2)  1.357 (3) 1.258 (3)  N - C(2) C(2)-C(3) C(3)-C(4) C(4)-C(5) C(5)-C(6) C(6) - N  1.373 1.388 1.424 1.399 1.380 1.342  1.376 1.387 1.422 1.396 1.360 1.349  1.378 1.370 1.433 1.410 1.350 1.351  C(2)-C(l) N - C(7) C(7) - C(8)  1.494 (2) 1.485 (2) 1.504 (3)  1.485 (3) 1.488 (3) 1.500 (4)  O(l) - M - 0(2) o(i) - M - o(iy O(l) - M - 0(2)' OO-M-OdW, 0(2) - M - 0(2)' M - O(l) - C(3) M - 0(2) - C(4)  84.23 (5) 90.46 (6) 94.87 (5) 172.50 (5) 90.93 (6) 112.1 (1) 111.3 (1)  83.03 (6) 90. 52 (6) 95.29 (6) 171.34 (5) 91.79 (6) 111.2 (1) 110.7 (1)  0(1)-C(3)-C(2) 0(1)-C(3)-C(4) 0(2)-C(4)-C(3) 0(2)-C(4)-C(5)  124.0 115.5 116.2 125.6  (2) (1) (1) (1)  122.4 117.0 117.5 124.9  (2) (2) (2) (2)  118.3 118.8 121.3 124.3  (2) (2) (2) (2)  C(6) - N - C(2) N - C(2) - C(3) C(2)-C(3)-C(4) C(3)-C(4)-C(5) C(4)-C(5)-C(6) C(5) - C(6) - N  121.6 118.8 120.5 118.1 119.2 121.8  (1) (2) (1) (1) (2) (2)  120.9 118.8 120.6 117.6 120.1 122.0  (2) (2) (2) (2) (2) (2)  120.5 118.6 122.9 114.3 121.9 121.8  (2) (2) (2) (2) (2) (2)  N - C(2) - C(l) C(l)-C(2)-C(3) C(6) - N - C(7) C(2) - N - C(7) N - C(7) - C(8)  120.5 (2) 120.7 (2) 117.4 (2) 121.0 (2) 112.6(2)  120.2 121.0 118.0 121.1 112.5  (2) (2) (2) (2) (2)  120.0 121.5 117.3 122.0 111.8  (2) (2) (2) (2) (2)  M=  (2) (2) (2) (2) (3) (3)  Hmepp  (3) (3) (3) (3) (3) (3)  (3) (3) (3) (3) (4) (3)  1.490 (4) 1.485 (3) 1.484 (4)  88  Table 3.10. Unit cell dimensions for M(dpp) and M(mepp)3 complexes. 3  Dimension  M(dpp)3 M =  Al  Ga  In  Al  Ga  16.600 (2)  16.6549 (6)  16.842 (1)  17.1734 (8)  17.247 (1)  6.877 (1)  6.8691 (4)  6.8078 (7)  6.827 (1)  6.830 (2)  Volume (A )  1641.3 (3)  1650.1 (1)  1672.3 (2)  1743.7 (3)  1759.4 (1)  D @ (g/cm )  1.33  1.47  1.48  1.33  1.40  a (A)  c(A) 3  3  c  @ D  M(mepp)3  c  = calculated density.  There are some differences in the unit cell dimensions (Table 3.10) of the structures that were determined. (In the trigonal space group P3, a =b * c, a = P = 90° and y = 120°.) The rigidity of the water network that was responsible for the compression of the M-O-C bond angles along the C3 axis (parallel to the c- axis) can be seen in the length of c. For the M(dpp)3 complexes, this dimension decreases (slighdy as the metal radius increases) to a minimum for In; the increase in the size of the N-alkyl substituent likewise causes a decrease in this parameter, e.g. 0.050 A between Al(dpp)3 and Al(mepp)3- It is along the ab diagonal that the changes in the metal and ligand are accommodated by the water network. The result is an increase in a:, e.g., a 0.242 A  increase for In- vs.  Al(dpp) , and a 0.573 A increase for Al(mepp)3 vs. Al(dpp)3. 3  The packing diagram of Al(mepp) (Fig. 3.13) shows how the N-ethyl group fits 3  into the hydrophobic core that is made up of one ligand from each of the two/ac-M(mepp)3 units in the cell. The flexibility of the structure in the ab diagonal is due to the large ring that encircles the core and consists of bridging waters, one side of the hexagonal water  89  channels, and the M - O bonds in the /ac-M(mepp)3 units (the oxygen atoms have been darkened to highlight this ring). The ligands in the core are separated by 3.5±1 A; i.e., normal van der Waal's contacts. The points where the ligand carbon atoms approach the water rings, C(8)-0(5) and C(6)-0(6), are likewise separated by ~3.5 A. To maintain this distance from the water network, the N-ethyl group must twist out of the ligand plane. The length of c (> 6.8 A) is considerably greater than van der Waal's contact distances so there is room to accommodate the N-ethyl group without an increase in c and, therefore, without an increase in the length of the H-bonds that determine this dimension. Focussing on this inner core and its encircling ring, it is possible to see why the increase in metal radius did not disrupt the water network. Increasing the length of the O - M - 0 portion of the inner ring pushes the water channels away from each other (increases a) but no strain is put on the water channels because they are connected via the /ac-M(mepp)3 unit rather than directly connected by H-bonds. The structure can be thought of as four rigid water columns held together by the more flexible /ac-M(mepp)3 units. A n increase in the size of the metal simply pushes the water columns apart in the a and b directions, and in the one direction where direct strain could be placed on water H bonds, the octahedral metal complexes are compressed before the H-bonds are stretched. The crystallographic inequality that distinguishes In(dpp)3 from its A l and Ga analogues also separates the M(mepp) complexes from each other. A comparison of the 3  packing diagrams of A l - and Ga(mepp)3 clearly indicates that the unit cells are not equivalent. There is a rotation of the/ac-M(mepp)3 units by ca. 60° about the C axis with 3  respect to the water network. Each unit cell contains one A and one A stereoisomer; this was necessary for the ligands in the hydrophobic core to lie parallel to each other and, therefore, to define the dimensions of the core. In order for the ligands to be as close as possible, the rings are staggered so the C(l) methyl groups point away from each other. In the unit cell of Al(mepp)3 (see Fig. 3.13), the A isomer is in the upper left of the cell and  90  the methyl groups project into the plane of the paper, the A isomer is in the lower right and the methyl groups are toward the viewer. Figure 3.14 shows the packing in the unit cell of Ga(mepp)3 and here the position of the stereoisomers is reversed. This was referred to as a crystallographic inequality simply because it only involves the unit cells and does not alter the/ac-M(mepp) units or the interaction of the water rings with these units. It is possible 3  this alteration of the packing arrangement is caused by the size of the metal as for both ligands, the change occurs in the tris-ligand complex with the largest metal; i.e., In(dpp)3 and Ga(mepp>3. However, it is not readily apparent why this rotation of the ligand orientation occurs and it must be considered simply a crystallographic oddity.  91  Figure 3.13. ORTEP view down the c axis of the unit cell packing of Al(mepp)3.  92  93  Chapter IV  A.  N M R Studies  Aluminum-27 NMR Spectroscopy  4.1 Introduction The  2 7  A l nucleus has a nuclear spin of 5/2 and, therefore, it has a nuclear  quadrupolar moment (Q). When a quadrupolar nucleus is placed in a magnetic field, the N M R energy levels are perturbed by quadrupolar effects. This quadrupolar interaction is affected by the direction of the electric field gradient that is fixed by the molecular framework. Thus the quadrupolar energy can be modulated by the Brownian motion of the molecule and if this occurs at the proper rate, spin-lattice relaxation (Ti) will be induced. Because the molecular motion is random, it has random phase and this leads to the loss of phase coherence between nuclei, i.e., spin-spin relaxation (T2). The result is an efficient  magnetic  relaxation mechanism dominated by nuclear  quadrupole  relaxation. 128,129 The electric field gradients are generated by ligand field asymmetry so there is a direct connection between the line width of the A 1 N M R signal and the geometry of the 2 7  coordinated ligands or coordinated solvent molecules.  Octahedrally solvated A l  3 +  ,  tetrahedral AIX4", and AI2X6 dimers exhibit high symmetry and have relatively narrow line widths. For trigonal species, the line widths (measured as peak widths at half height, W1/2) become much larger.  The overall range of line widths is from 3 H z (for  [A1(H20)6] ) to over 6000 H z . 1 3+  30  There have been a number of studies using A l N M R 2 7  line widths as a probe to determine the coordination number of organoaluminum complexes  94  in organic solvents.  A far lesser amount of data has been collected in aqueous solution  131  for A l complexes with biologically active l i g a n d s . " 132  134  The A l isotope is 100% naturally abundant and has a relative receptivity one-fifth 2 7  that of the proton. Comparison with the 1.11% receptivity of  1 3  natural abundance and 0.016 relative  C indicates why A 1 N M R does not require the instrument time typical of 27  C N M R . The practicality of A l N M R can best be illustrated by comparing its nuclear  1 3  2 7  properties to those of its congener Ga and of two other quadrupolar nuclei,  1 7  0 and  1 5  N,  that are used as N M R nuclei. The data in the last two columns of Table 4.1 were calculated from the following equation for nuclear spin quadrupolar relaxation (TQ) that is valid in the Umit of fast motion.  128  1  _ 1 _ 1 _ 3 TQ " T I " T " 40 2  r  21 + 3 /eQ\ ' ~(21 - 1) I R J ( d z ) 2  2  a 7  2  2 1  I is the nuclear spin quantum number, eQ is the electric quadrupole moment, h is Planck's dv constant divided by 2%, -j-^ is the maximum electric field gradient at the nucleus, and T is 2  c  the correlation time for Brownian motion. The line width factor, defined as (21 + 3) Q I (21-1)  2  2  can be used to compare the line widths of quadrupolar nuclei. Normalized to the value for 2 7  A l , the relative line widths of the Ga isotopes are the largest and the other two  quadrupolar nuclei have peaks < 10% that of A l . However, when the differences in 2 7  abundance and receptivity are used to determine the relative peak heights, defined as abundance * receptivity line width factor 2 7  A l N M R has a hypothetical sensitivity an order of magnitude (or more) greater than the  other quadrupolar nuclei in Table 4.1. As a result of this, high resolution A l N M R has 2 7  been used to quantify A l  3 +  to 1 ppm (37 p M ) .  1 3 5  95  Table 4.1. The N M R properties of several quadrupolar nuclei (refs. 128 and 129-Appendix 2) NMR frequency (MHz)  Relative receptivity  Isotope  I  27A1  5/2  100  26.08  0.206  0.149  1.00  54.6  69Ga  3/2  60.4  24.04  0.042  0.19  6.7  1.00  Ga  3/2  39.6  30.55  0.057  0.12  2.6  2.24  0  5/2  0.037  13.56  0.00011  -0.026  0.03  0.003  1  99.63  7.23  0.001  0.01  0.07  3.8  1/2  99.98  100.0  1  —  —  —  7 1  1 7  14  N  lH  1  Quadrupole Relative moment line (10- m ) width  Relative peak height  Isotopic abundance (%)  28  2  1  Relative (to H) receptivity at a constant field with equal numbers of nuclei.  2  Relative (to A l ) line width.  3  Relative (to G a ) peak height.  2  l  2 7  69  We did attempt some studies with G a N M R , but its lack of sensitivity (see Table 7 1  4.1) coupled with the low solubility of our Ga complexes made it impossible to detect any signals in a practical time span.  2 7  A l N M R served as a characterization technique and it  was also used to readily determine the success of our tris-ligand A l complex preparations. The constraints on instrument access unfortunately restricted this latter application. Our primary utilization of this N M R technique was the determination of the hydrolytic stability of the tris(3-hydroxy-4-pyridinonato) A l complexes. 27  Al,  1 3 6  A 1 N M R has been used extensively to examine the p H dependent speciation of  and the chemical shifts and narrow line widths of the predominant A l species at pH  < 4, [A1(H20)6] , and at pH > 9, [Al(OH)4]", are quite different from those of the 3+  tris-ligand complexes in this study. When a solution of an A l complex was sampled at a  3  96  series of p H values from < 2 to > 10, the A l N M R spectra readily indicated the pH region 2 7  where the hydrolysis products first appeared. The broad peaks of our complexes made it difficult to quantify by this technique, but it did paint a clear picture of what we referred to as the "window" of hydrolytic stability for Al(ma)3. using  2 7  37  Similar A l hydrolysis experiments  A 1 N M R have been carried out with the tris-ligand A l complexes of  acetylhydroxamate, ligands.  134  137  oxalate,  138  lactate,  139  and a number of hydroxycarboxylate  The closest comparison to the spectra of the tris(3-hydroxy-4-pyridinonato) A l  complexes were the spectra of Al(ma)3.  4.2 Materials and Methods  The A l N M R spectra were recorded at ambient probe temperature (ca. 18 °C) on a 2 7  Varian XL-300 N M R spectrometer at 78.16 M H z and a pulse width of 15 p.s. The first experiments used a sweep width of 20 K H z and an acquisition time ( T ) of 0.20 s. 6500 aq  Transients were needed to achieve a reasonable signal to noise ratio (S/N) and the recording of one spectrum required 22 minutes. Because of the broad lines, we felt the resolution could be reduced in order to complete one variable-pH experiment (involving ten or more samples) in the two hour time blocks available on this instrument. Combinations of sweep width (20 to 60 KHz) and T  a q  (0.05 to 0.40 s) were tried before settling on 50 K H z and  0.12 s. These acquisition parameters allowed the resolution of both the narrow signals from the hydrolysis products and the broad signals from the Al-ligand complexes. If the sample concentrations were > 10 m M , 3500 transients gave a good S/N and the 7 minute run time per sample was compatible with the limitations on instrument availability. The spectra were referenced to the [A1(H20)6] signal (set as zero) from 0.20M A1(C1C>4)3 in 3+  97  0.10M HCIO4. D2O was added for locking and downfield chemical shifts were positive. The spectra were recorded by the author. Purified compounds, distilled and deionized water, and a Fisher Accumet model 805 pH meter (calibrated with p H 4 and 10 buffer solutions) were used in the variable-pH experiments. Solutions of the A l complexes were made in 15 mL H2O and 2 m L D2O (for a lock signal) and the initial p H was adjusted to < 2 by the addition of 8 M HC1. The pH was raised by the addition of 8 M NaOH and the solution was equilibrated between pH changes for > 10 minutes. Aliquots were withdrawn and filtered through glass fibers into the N M R tube.  The sampling was initiated at the lowest p H and was repeated at  approximately one pH unit intervals up to p H > 10. The hydrolysis was completely reversible within the 2-3 hour time frame of the experiment.  4.3 Results and Discussion  The chemical shift range of the A l nucleus is approximately 450 ppm and, like the 2 7  line widths, the shift is characteristic of the ligand symmetry about the A l atom. The signals for hexacoordinate A l nuclei are quoted as occurring from 20 to -46 ppm and A1C>6 species usually appear close to 0 p p m .  1 3 0  There are exceptions to this rule: the  tris(hydroxamato) A l complexes and the alumichrome trihydroxamato peptides resonate at 36-42 p p m  1 3 7  and a dimeric acetate species gives a signal at 38 p p m .  1 3 6  These exceptions  have led to the proposal that downfield shifts near 38 ppm could be characteristic of chelation by small-ring-forming bidentate ligands.  136  The tris(3-hydroxy-4-pyronato) A l  complexes have chemical shifts of ca. 39 ppm and the tris(3-hydroxy-4-pyridinonato) A l complexes resonate at essentially the same frequency. The downfield shifts are due to the  98  inequivalence of the chelating hydroxy and carbonyl oxygen atoms. The A l nuclei are not subject to a rigorously octahedral field and this is also reflected in the relatively broad peaks of these tris-ligand A l complexes. The chemical shifts and the W1/2 values of the tris(3-hydroxy-4-pyridinonato) A l complexes are listed in Table 4.2; the data are from the variable-pH experiment spectra of the p H 7-8 samples. Al(mhpp)3 is not sufficiently water soluble for the variable-pH experiment and its spectrum was recorded in CD3OD. The values for Al(ma)3 are included for comparison and to put these data in perspective, Al(acac)3 (D3 point group) resonates at 0 ppm with W1/2 = 100 H z .  Table 4.2.  2 7  1 3 0  A l N M R data for the tris(3-hydroxy-4-pyridinonato) A l complexes. > Al(ma)3  Al(mpp)3  Al(dpp)3  Al(mepp)3  Al(mimo)3  Al(mhpp)3  ppm  38  38  37  37  37  38  Wi/ (Hz)  900  600  700  780  1600  1400  2  Our other data (spectroscopic and crystallographic) indicate the A106 coordination sphere is very similar for all the 3-hydroxy-4-pyridinone ligands; it is not likely that differences in ligand symmetry would account for the variations in line widths that were observed. N M R line widths are also affected by temperature, solvent viscosity, exchange processes, and the mass of the solute.  129  The proton N M R indicate an exchange process is  occurring at room temperature in the A l complexes and this could contribute to the variations in line widths. In comparison to Al(mpp) , the larger R groups in Al(mimo)3 3  and Al(mhpp)3 may be decreasing the molecular tumbling rate; this would increase the T  c  for Brownian motion and, as indicated in the above equation, the relaxation rate (and therefore the W1/2) for the larger tris-ligand A l complexes would increase.  99  After Subtraction  Before Subtraction  Solvent Blank  I i i I i I i i I i I i I I I I i M '200  150  Figure 4.1.  2 7  100  50  I I I i N I M 0 -50  i i | I I i I | i I I I | -100  -150  PPM  -200  A l N M R spectra of Al(mpp)3 at pH 1.8 showing the background correction.  100  The minimum concentration for the variable-pH studies was 10 m M ; this cutoff point is not due to the sensitivity of the technique but rather is due to the background signal from the A l in the ceramics of the N M R probe. This is a well documented phenomenon 140  141  and the  probe A l has a broad (~6000 Hz) signal centered at ca. 60 ppm. The background signal from the probe overlaps the signals from the tris-ligand A l complexes (near 40 ppm). The probe signal is also out of phase with the signals from the solution A l nuclei which makes it difficult to properly phase the spectrum. It is possible to do a background correction by subtracting the free induction decay (FID) of a solvent blank from the sample FID and Fourier transforming the resulting signal. Figure 4.1 depicts the before and after spectra of a p H 1.8 Al(mpp)3 sample with an initial A l concentration of 15 m M . (The sharp signal at 0 is from [A1(H20)6] and the smaller broad peak at 17 ppm is from [Al(dpp)(H20)4] .) 2+  3+  This background  correction works well but the experimental procedure presented an additional complication. As the p H was raised, the composition of the solution changed due to the addition of base and this was hard to duplicate in the solvent blank. If the conductivity of the solvent blank was sufficiently different from that of the sample, it effectively detuned the probe so the phase shifts were altered and the subtraction did not work. This problem was handled by making several solvent blanks and using the one that gave the best results for a particular sample.  The interference from the background signal can be avoided by using A l concentrations > 20 m M . For Al(mepp)3 and Al(mimo)3,40 m M solutions were used; Al(mpp)3 is soluble to 16 m M and background subtraction was necessary in a number of the spectra. Al(dpp)3 is only soluble to 1.2 m M at neutral pH, but.we were able to make a supersaturated solution with an initial A l concentration of 30 m M at pH 1.8. The insolubility of the tris-ligand A l complex is due to the incorporation of water molecules (see Chapter HI) and this is not an instantaneous process. There was no appreciable precipitation in the base solution over the two to three hours required for the variable-pH experiment. The N M R spectra were recorded immediately after sampling to avoid precipitation in the N M R tube. (Crystals did form in the tubes  101  overnight and it was this experiment that first indicated X-ray diffraction grade crystals of this compound could be grown from water.) The aqueous solubility of Al(mhpp)3 is similar to that of Al(dpp)3 and a concentrated solution can be made in acidic solution. However, its low solubility is due to the lipophilic N-hexyl groups and the variable-pH experiment was not feasible because a pervasive precipitate formed at ca. p H 2.5. (A detailed summary of the solubility properties of these complexes is in Chapter V.) The variable-pH A l N M R spectra for Al(mpp)3 are shown as Figure 4.2. The acidic 2 7  hydrolysis is evinced by a shoulder at p H 3.4 on the upfield side of the Al(mpp)3 peak at 37 ppm. This results from the partial protonation of ligands in the coordination sphere of the A l and their replacement with waters. At p H 3.1 this shoulder is resolved into two shoulders resulting from the [Al(mpp)2(H20)2] (28 ppm) and [Al(mpp)(H20)4] (17 ppm) species. +  2+  When the solution is acidified further, the completely hydrated species [A1(H20)6]  3+  is  observed at 0 ppm. As the p H drops the signals from the hydrolysed species gain in intensity at the expense of the Al(mpp)3 peak. Even at the lowest p H of 1.8, there is still a signal from the monoligand species although most of the A l is in [A1(H20)6] . When the p H is raised, 3+  the ligands are replaced by the hydroxide ion to form ultimately [Al(OH)4]" at p H 11 (80 ppm, W1/2 = 60 Hz). The basic hydrolysis of A l  3 +  is time dependent and this explains why no  intermediate species analogous to those formed at acidic p H were found in the time frame of the experiment. The Al(mpp)3 pH 8.9 sample was left (in an N M R tube) at room temperature for two weeks and the spectrum recorded at that time is contrasted to the initial spectrum in Figure 4.3. After two weeks, the intensity of the Al(mpp)3 signal is reduced by one-third and the broad peak at 61 ppm is probably due to four coordinate mixed aquo/hydroxo/mpp A l complexes. In the basic hydrolysis of tris(lactato)aluminum(III), mixed hydroxo/lactato complexes are ascribed to a broad peak at 60 p p m .  139  102  i i i | i i i i | i i i I | i i i i | i i i i | i i i i | i i i i | i i i i | i  200  0  100  -100  ppm  Figure 4.2. The variable-pH A 1 N M R spectra of Al(mpp)3. 27  103  —I  I . | I  200 Figure 4.3.  2 7  JI  I  j  I  Ii  ! | 1 i I I | 1M  I | I . I I  100  0  j  i I II  |  i  .100  I I I | 1  ppm  A l N M R spectra ofAl(mpp)3 at p H 8.9. The bottom spectrum is 2 hours  after sampling and the top spectrum is the same N M R sample 2 weeks later.  A comparison of the variable-pH spectra of Al(mpp)3 with those for Al(ma)3 indicates that the tris(3-hydroxy-4-pyridinone) complex is more stable to both acidic and basic hydrolysis. In the spectra of Al(ma)3, the signal from [A1(H20)6] first appears at 3+  p H 3.2 and Al(OH)4" is evident at p H 9; the spectra of Al(mpp)3 show no signals from either of these species at similar pH values. The difference is particularly evident at higher p H where the [Al(OH)4]" signal does not occur until p H 10.8 in Al(mpp)3. The Al(dpp)3 spectra (Fig. 4.4) are similar to those of Al(mpp)3: the intermediate species occur at 26 ppm ([Al(dpp) (H 0)2j ) and at 14 ppm ([Al(dpp)(H20)4] ). At p H 2  2  2+  +  2.3 there appears to be significantly more of the bisligand species as the two peaks at 26 and 14 ppm are resolvable, unlike in the Al(mpp) spectra. The only significant change in 3  104  the Al(mimo)3 spectra (aside from broader peaks) is a shift in the window of hydrolytic stability toward lower p H (Fig. 4.5). It must be stated that it is somewhat unfair to compare the spectra of the A l complexes in this study. We are observing a competition for the A l ion among water, hydroxide ions, and the 3-hydroxy-4-pyridinone ligands; the concentrations of the first two are fixed at any given p H but the ligand concentrations were different for each complex studied. It is possible to conclude that the tris(3-hydroxy4-pyridinonato) A l complexes are more stable to hydrolysis than Al(ma)3 on the basis of this experiment, however, since the initial Al(ma)3 concentration was larger (50 mM) than any A l concentration used in this study. The first variable-pH experiment with Al(mepp)3 (Fig. 4.6) gave spectra that showed a second hydrolysis product at p H > 9.0 (62 ppm, W1/2 = 50 Hz). This narrow peak was likely due to the polymeric [A104Ali2(OH)24(H20)i2] (commonly referred to 7+  as  " A I 1 3 " )  species whose symmetric tetrahedral AIO4 core resonates at 62.5 p p m .  1 3 1  A I 1 3  is the principal Al N M R detectable base hydrolysis product that is formed when solutions 27  containing only [A1(H20)6] are neutralized by the addition of 2.5 equivalents of base. 3+  The polymer is unstable at the acidic p H where it forms and it can undergo slow transformations to give other less symmetric species that have much broader signals. The appearance of the spectrum is very different depending on the speed of the hydrolysis and this added to the difficulty in determining the precise nature of the A l hydrolysis reactions.  136  Given the similarity of the 3-hydroxy-4-pyridinone ligands, it seemed unlikely that Al(mepp)3 was undergoing hydrolysis in a significantly different fashion from the other complexes where no A I 1 3 signals were observed. In the other variable-pH experiments, the isolated tris-ligand A l complexes were used as the starting material, but in this experiment A1(NC>3)3 and the free ligand (in stoichiometric ratio) were used. It appears there was a slight excess of A l that was hydrolysed to the A l 13 polymer, as the length of the  105  experiment was short enough to ensure this would be the preferred hydrolysis product. This was confirmed by repeating the experiment with purified Al(mepp)3 as the starting material: there was no AI13 signal in the spectra at pH 9.6 and the addition of A1(NC»3)3 resulted in the appearance of the 62 ppm peak that subsequently disappeared when an excess of the free ligand was introduced. The p H region where the AI13 signal was found is not as readily explained. The AI13 species forms at ca. pH 4 and at p H > 8 it should disappear as, in the absence of other ligands, [Al(OH)4]" is the dominant A l hydrolysis product at higher p H .  1 3 6  The lack of a  signal at lower p H could be ascribed to an exchange with the Al(mepp)3 complex, but the persistence of the signal (there was no significant change in the spectrum after one hour) at higher p H is something of a mystery.  The AI13 species was seen in hydrolysis  experiments with tris(lactato)aluminum; there was a weak signal at 62 ppm that appeared at p H 7 and was completely gone by p H 10 where only [Al(OH)4]" was present.  139  No  signal due to AI13 was found in a p H controlled study of a series of hydroxy carboxylic acids even at 1:1 ligand to A l ratios.  134  The robustness of these A l complexes is evident from their wide window of hydrolytie stability, from p H < 4 to > 9. They appear to be more stable to hydrolysis than Al(ma)3 and this is corroborated by the formation constants (see Chapter V) as the log P3 for the tris(3-hydroxy-4-pyridinonato) A l complexes is ~32 compared to a log P3 of 22.5 for Al(ma)3.  59  The variable-pH experiment with Al(mepp)3 directly illustrates the ability of  the 3-hydroxy-4-pyridinones to compete with the hydroxide ion. The results of this variable-pH A 1 N M R study indicate the tris(3-hydroxy-4-pyridinonato) A l complexes 27  should resist in vivo hydrolysis except in the highly acidic conditions of the stomach.  106  Figure 4.4. The variable-pH A 1 N M R spectra of Al(dpp>3. 27  107  M  I j  I I I I ) I I I I J I I I I ) I I  wo  ido  |  |  I  i  |  ||  i  I |  11  | • | | ,  -lboppm  Figure 4.5. The variable-pH ?A1 N M R spectra of Al(rrumo) . 2  i . ,  3  108  1 I 11I I 1I | 1I I M I I 1I I I I I 1I I I I I I I 11I I I I I M I  2<Jo  ldo  o  -iboppm  Figure 4.6. Variable-pH A 1 N M R spectra of a 3:1 Hmepp to A1(N03)3 rnixture. 27  109  B.  Variable-Temperature Proton NMR Spectroscopy  4.4 Introduction In tris-ligand metal complexes, asymmetric bidentate ligands may assume a facial (fac) or a meridional (mer) geometry. The fac and mer geometric isomers* are each enantiomeric pairs of A and A stereoisomers.  (Figure 4.7 is a generalized diagram  illustrating the four isomers.) The rearrangement reactions are geometric isomerization and racemization; these reactions can occur separately or simultaneously depending on the mechanism of the rearrangement. Three types of mechanisms have been proposed: (1) the complete dissociation of one ligand to give a four-coordinate intermediate, (2) the rupture of one metal-ligand bond to give a five-coordinate intermediate, and (3) twisting processes (Bailar or RSy-Dutt twists) that do not require the cleavage of any metal-ligand bonds.  142  The majority of research on the rearrangement reactions of tris-ligand metal complexes has been directed towards differentiating between the two mtramolecular mechanisms.  fac-A  fac-A  mer-A  mer-A  Figure 4.7. Isomers of tris-ligand metal complexes with asymmetric bidentate ligands.  *They are frequently referred to as cis and trans isomers respectively.  110  Gordon and H o l m  1 4 3  define two limiting types of tris-ligand systems based on the  rates of the rearrangement reactions. The kinetically "slow" systems are those in which the geometric isomers can be completely separated and partially resolved. The reaction rates are ? 10" s" and complexes containing inert metal ions such as C r 2  1  3 +  and C o  3 +  144  are in this  category. The second type is designated "fast"; this means that the intramolecular rearrangements are rapid enough to prevent separation or resolution of the isomers, but they are not so rapid as to disallow isomer detection by N M R spectroscopy or other techniques such as low temperature H P L C .  1 4 5  Tris-ligand A l and Ga complexes fall into  the fast category. N M R is particularily suited to the examination of the rearrangement reactions because the fac isomer has a threefold symmetry axis and the mer isomer is asymmetric; the three ligands of xhefac isomer are magnetically equivalent and the chemical shifts of the nuclei on these ligands will be different from that of their inequivalent counterparts in the mer isomer. If the chemical shift differences are large enough, the isomers may be identified and it is possible to measure the rates of isomerization and racemization. In the early 1960s, Fay and Piper used variable temperature  1 9  F N M R to examine the fluxional  behavior of tris(trifluroacetylacetonate) (tfac) complexes with a variety of metals including A l , Ga, and I n .  1 4 6  This classic work established the utility of variable-temperature N M R  for the study of stereochemically nonrigid inorganic complexes. Proton N M R has been used extensively to study the rearrangement reactions of tris(P-diketonato) metal complexes and the group 13 metals have figured prominently in this w o r k .  144  -  147  tris-ligand A l and Ga complexes with other ligands such as E D T A tropolonates ( a - R T )  1 4 9  -  1 5 0  1 4 8  The fluxionality of and a-substituted  has also been examined by N M R techniques. The tropolone  complexes are of particular interest as these bidentate ligands chelate the metal center with the same binding group (the a-hydroxyketone moiety) as the 3-hydroxy-4-pyridinones.  Ill The tris(3-hydroxy-4-pyridinonato) A l complexes exhibit complex proton N M R spectra at the ambient probe temperature of the 300 M H z N M R instrument. Under the same conditions, the Ga analogues give the expected first order spectra. The spectra of the A l complexes can be rationalized on the basis of fac-mer isomerization that is slow enough to result in overlapped signals from the geometric isomers. A variable-temperature N M R study of the M(dpp) complexes was conducted to verify that ligand rearrangement is the 3  source of the difference between spectra of the analogous tris-ligand A l and Ga complexes. The results of this study indicate that two exchange processes are occurring and the coalescence temperature (T ) for the higher temperature process is near room temperature c  for Al(dpp) and is at -9 °C for Ga(dpp) . 3  3  The solubility properties of the M(dpp) complexes necessitate the use of protic 3  solvents. The majority of this work was done in C D O D due to the solubility of the 3  M(dpp) complexes and the large temperature range available with this solvent. Ga(dpp) 3  3  is too insoluble in D 0 to allow the acquisition of a spectrum in a reasonable length of time. 2  Ga(dpp) is quite soluble in ( C D ) S O but it does not reach the region of slow exchange 3  3  2  above the freezing point (18.5 °C) of this solvent. Therefore, the difference in lability between the two metals was examined by variable-temperature N M R in C D O D . Because 3  of the importance of water interactions to the solid state structure of the M(dpp)  3  complexes, we were also interested in determining if the rate of the rearrangement reaction for Al(dpp) is significantly different in water. The exchange process for Al(dpp) was 3  3  examined in D 0 , ( C D ) S O , and C D O D ; because of the low water solubility of 2  3  2  3  Al(dpp) , the fluxionality of the more water soluble Al(mpp) was also investigated in 3  D Q. 2  3  112 4.5 Materials and Methods  The spectra were recorded on a Varian XL-300 N M R spectrometer equipped with a variable temperature probe. The thermocouple was calibrated using a methanol calibration standard and was accurate within ± 1 ° over a range of -70 to 60 °C. The spectra in CD3OD and (CD3)2SO were referenced to the solvent peak arid those in D 0 were 2  referenced internally to (CD3)2CO. The spectra were recorded by the author. A l - and Ga(dpp)3 and Al(mpp)3 were prepared and purified as reported herein (Section 2.6). In(dpp)3 and Al(ma)3 were synthesized by others in this laboratory (C. Matsuba and M . Finnegan) and were purified by recrystallization. The equilibrium distribution of the Al(dpp)3 geometric isomers in the absence of exchange was determined by computer simulation using the N M R line-shape program D N M R 3 .  151  4.6 Results and Discussion  The spectra of Ga- and Al(dpp)3 in C D 3 O D at 18 °C clearly show the differences that prompted this variable-temperature N M R study (Figs. 4.8 and 4.9~the spectral regions without signals have been omitted for clarity). The singlet for the C H 3 group has a W1/2 C  of 1.6 H z in Ga(dpp)3 compared to an exchange broadened 7.0 H z in Al(dpp)3. The downfield doublets in the Al(dpp)3 spectrum are also broadened and a second signal is just starting to appear as shoulders on the Hb doublet. The Al(dpp)3 spectra in (CD3)2SO (Fig. 4.10) and D 2 O (Fig. 4.11), also at 18 °C, afford better resolution of the signals from the fac and mer isomers due to the slightly larger chemical shift difference and the higher coalescence temperature in these solvents. In both spectra, there are four distinct signals from the three inequivalent CH3 groups of the mer isomer and the one unique methyl of C  the fac isomer.  113  Hdd  e a f B1  9  s:2  8 B  a I B 8 • »t > «a a I I  B;Z t I 11 t t I I  O;E  I i • • »I »  t I I  II  l f  2;E • I l t 11 I I  re I t t  Y  II  • 11  II  g;e I I I  II  B;E  II  I 11  I I  11  I I 11  It  o;tI I  11  I I 11  II  1I  CIO(H*CD)  o 3  HDN  H3  o o  £  LO  oo o  o  Ndd 2-9  fr'9  9*9  B-9  07  27  f 7  97  97  I I t I 1 1 I 1 I I I I I I I I I I I 11 I I I I I I I I 1 I I I I 1 1 I M I I I I I 11 I t i I I I I 1 1 1 1 i t 1 I I I I 11 I 1 1 1 1 i N 11 1 1 1 1 I I 1 1 1 1 I 1 1 1 I I 1 1 1 1  11  1  115  Figure 4.10. 300 M H z proton N M R spectrum of Al(dpp)3 in (CD3) SO at 18 °C. 2  Figure 4.11. 300 M H z proton N M R spectrum of Al(dpp)3 in D 0 at 18 C . 2  e  117  The chemical shift differences in the downfield doublets are quite small so the H  a  and H5 ring protons are of limited utility for identifying the geometric isomers in any of the solvents. The same is true of the N-CH3 signal. For this reason, the kinetic parameters were determined from the temperature-dependent C E ^ spectra. Among the four signals, the one that was different in intensity was assigned to the /ac-isomer, this assignment rationale was used by Piper and Fay for the fluorine resonances in the M(tfac)3 complexes.  146  Figure 4.12 shows the C H ^ signals for A l - and Ga(dpp)3 in the absence of  exchange (-30 °C) and the signal assigned to the fac isomer is marked with an asterisk. (The scale in Hertz indicates the magnitude of the peak separation and not the chemical shift) Two of the mer isomer signals in Ga(dpp)3 appear as shoulders, but the assignment of the peak with the odd intensity to the fac isomer has been made with other tris-ligand metal complexes where only two of the four peaks are resolvable.  149  ji 1111111 iji 111 j 111\j\11111111 y 11111111j JM11111111) 111 ^111111111 y 111111 n Figure 4.12. CH3c spectra for Al(dpp)3 (left) and Ga(dpp)3 (right) in CD3OD at -30 °C..  A statistical distribution of the geometric isomers would give four peaks of equal intensity, i.e., the fac isomer would be 25% of the total concentration. It has been stated that the mer isomer is the more stable isomer due to its lower dipole moment and this was  118  used to explain the smaller than statistical (18%) equilibrium distribution for the kinetically fast Al(tfac)3 complex.  146  Gordon and Holm maintained that unless the complexes were  sterically constrained, a statistical or nearly statistical distribution of isomers would usually be formed in solution.  143  ./ac-isomer at equilibrium  The A l - and Ga(cc-RT)3 complexes had a slight excess of the 150  and this was also the situation for Al(dpp) . D N M R 3 was 3  used to simulate the equilibrium spectrum of Al(dpp)3 in the absence of exchange and the best fit (Fig. 4.13) was obtained with a 32% distribution for the /ac-isomer. The enhanced stability of the fac isomer was interesting when considered in concert with the solid state structures (see Chapter III). It was thought that the water networks imposed a facial geometry on the M(dpp)3»12H20 complexes, but these N M R results in methanol suggest that the fac isomer enjoys at least a small thermodynamic advantage independent of the H-bonded water network. The close match of the computer simulation also supports the assignment of the larger signal to the/ac-isomer.  Figure 4.13. Experimental (left) and simulated (right) Al(dpp)3 CH3c spectra in CD3OD (-30 *C).  For the exchange of nuclei between two inequivalent sites, the rate of exchange at Tt Av the temperature of coalescence (kj ) is given by: k x =—p— where Av is the frequency c  c  V2  separation (in Hertz) between the resonance components in the absence of exchange.  152  119 For the CH3 spectra, A v (the experimentally observed frequency separation) was taken as C  e  the difference between the signal for the/ac-isomer and the furthest downfield signal from the mer-isomer. A source of error in using this equation is the temperature dependence of the chemical shifts that is presumably due to solvent-solute interactions.  153  To correct for  this temperature dependence, the A v was plotted against temperature at several points in e  the region of slow exchange, and the line was extrapolated to the region of fast exchange. Then Av was read from the extrapolated line at the T . In our systems, the values of A v c  e  were small (5-8 Hz) and the variations due to temperature dependence were barely within the resolution error of the instrument; therefore, the required adjustments in Av were minor. The value of the free energy of activation (AG+Tc) was calculated from the Eyring equation  154  (assuming the transmission coefficient to be unity): k r = - j p T e -AGf/RTc c  c  where h is Planck's constant, K B is Boltzman's constant, and R is the gas constant. The simplified Gutowsky-Holm equation  152  is used to calculate the rate constants  for an exchange process from the peak separations in the slow exchange region; the rate constants can then be used to determine the E , AH^, and AS^ for the exchange reaction. a  The size of the error in using this approximation is related to the length of T and the 2  magnitude of Av. When Av is small and T is short, the errors become comparable to the 2  calculated rate constants and the peak separation method should not be used.  155  A value of  less than 0.33 for the ratio of T (expressed as W1/2) to Av was cited as the minimum 2  requirement for using the simplified equation.  156  For the CH3 spectrum in Al(dpp)3, the C  T determined by computer simulation is 0.135 s (W1/2 = 2.36 Hz) and the Av is 5.1 Hz. 2  The ratio of W1/2 to Av is 0.46 and the simplified equation could not be used for this complex, or for Ga(dpp)3 either as the CH3 signal from the Ga complex is even more C  poorly resolved.  120  The preferred method for estimating rate constants is line-shape simulation. The D N M R 3 computer program allows mutual and non-mutual exchange (referred to in the program documentation as "otherwise"). The non-mutual exchange routine has multiple chemical configurations that can have unequal populations and it is possible to simulate the equilibrium spectrum of Al(dpp) in the absence of exchange with this routine; however, 3  the non-mutual routine cannot directly accommodate exchange between chemical configurations of different symmetry. With the mutual exchange routine, the population difference between the fac- and mer-isomers cannot be incorporated so the larger peak due to the fac isomer is not modelled. Because of these limitations, the D N M R 3 simulated spectra did not correspond well enough to the experimental spectra to merit inclusion in this thesis. Qualitatively, the M(dpp) complexes appear to undergo two exchange processes in 3  a fashion similar to that of the M(cc-RT) complexes. 149  3  150  This is illustrated by the C H  3 c  spectra for Al(dpp) over a temperature range of -90 to 38 °C which are included as Figure 3  4.14. The spectra for Ga(dpp) are similar and Figure 4.15 is a view of the exchange 3  process for all of the protons in the complex (the spectral regions without signals have been omitted in this reproduction). In both M(dpp) complexes there is a low temperature 3  exchange process (LTP) with a T near the -98 °C low temperature limit of our C D O D c  3  study. There is a second exchange process with a T of 21 and -9 °C in A l - and Ga(dpp) , c  3  respectively. The kinetic parameters at the T for this higher temperature process (HTP) are c  listed in Table 4.3. The L T P for the M(oc-RT) complexes was identified as racemization 3  by means of a trigonal twist and the HTP was fac-mer isomerization. By comparison to the spectra of the M(cc-RT) complexes, the variable-temperature N M R spectra of the M(dpp) 3  3  complexes support the conclusion that the exchange process near ambient temperature is due to fac-mer isomerization.  122  Figure 4.15. Variable-temperature proton N M R spectra of Ga(dpp)3 in CD3OD.  123  Table 4.3. Kinetic data for tris(3-hydroxy-4-pyridinonato) metal complexes at the T . c  Complex  Solvent [e ]  Av (Hz)  T (°C)  Ga(dpp)3  CD3OD  4.8  -9  10.7  14.2 ± 1.2  5.1  21  11.3  15.7 ± 1.3  ( C D ) S O [40.7]  7.2  38  16.0  16.5 ± 1.4  D 2 O [78.5]  8.0  34  17.8  16.0 ± 1.4  D2O  6.0  36  13.3  16.5 + 1.4  Al(dpp)3  c  CD3OD  3  Al(mpp)  3  [32.7]  2  c  a  kr^s' ) 1  AG^Tc (kcal/mol)  a  Errors were estimated to be ± 2 °.  b  Errors were estimated assuming an order of magnitude error in the rate constants.  c  Dielectric constants at 25 °C.  The M(dpp)3 complexes are more labile than A l - and Ga(tfac)3 that have a AG+TC of 21.4 and 18.7 kcal/mol respectively (in C D C I 3 ) . the CF3 groups is well documented,  158  1 5 7  Since the rate-accelerating influence of  it would appear that the 3-hydroxy-4-pyridinones  constitute a relatively labile tris-ligand metal system. This is only speculation, however, as the contribution from solvent effects has not been taken into account. The rearrangement rates for d° and d  1 0  metals are dependent on the ionic radius of the m e t a l  144  and the  M(dpp)3 complexes exhibit the predicted kinetic order (Al < Ga < In). Ga(dpp)3 has a lower AG+Tc than Al(dpp)3 and the In analogue gives only an averaged spectrum down to -80 "C in C D 3 O D .  For a series of aprotic solvents, it was found that the T and AG^Tc of Al(tfac)3 c  decreased as the dielectric constant of the solvent increased.  159  There was no indication of  a similar trend in this study and the kinetic parameters for the Al(dpp)3 exchange process were essentially the same in the three solvents that were used (see Table 4.3). The kinetic  b  124  data for Ni(phenanthroline)  3  showed no simple correlation between the racemization rate  and the dielectric constant of the solvent (this study included water and methanol).  158  The  results for Al(dpp) are in agreement with this conclusion. 3  In order to obtain mechanistic information on the rearrangement reaction, it is necessary to ascertain that one is dealing with an intramolecular process. This can be accomplished by examining the N M R spectrum of a mixture of the tris-ligand metal complex of interest and another tris-ligand metal complex. If it takes an appreciable amount of time to form the mixed-ligand species, it can be concluded that ligand exchange is slower than the rearrangement reaction, e.g., no N M R signals attributable to mixed-ligand species were detected in one hour at room temperature for a mixture of Al(acac) and tris(l-phenyl3  5-methylhexane-2,4-dionato)aluminum(III).  160  When ligand exchange is completed in a  shorter period of time (in minutes at room temperature), the N M R spectrum at a temperature above the T for the rearrangement reaction can be used to determine the c  relative rates of the two processes.  If multiple signals assignable to the mixed-ligand  species are observed at a temperature where the rearrangement process gives an averaged spectrum, it can again be concluded that the intermolecular process of ligand exchange occurs at a slower rate. This is considered sufficient proof that the rearrangement reaction is an intramolecular process. Our interest in the possibility of solvent effects on the exchange process made it worthwhile to qualitatively compare the rates of ligand exchange and ligand rearrangement for Al(dpp) in C D O D . The other tris-ligand aluminum complex used in the ligand 3  3  exchange experiment was Al(ma) . The Al-maltol complex was chosen because, unlike the 3  tris(3-hydroxy-4-pyridinonato) A l complexes used in this study, the chemical shift of the ring protons was sufficiently removed from that of Al(dpp) to permit easy differentiation 3  of the signals from the two ligands.  125  The ratio of the N M R sample was 3:1 Al(ma)3 to Al(dpp)3 and Figure 4.16 shows the proton N M R spectrum fifteen minutes after mixing (at a probe temperature of 18 °C). It is evident that mixed-ligand species have already formed and the 2:1 ratio of the signals from the N-methyl group of the dpp" ligand (at 3.82 ppm) suggests the formation of Al(dpp) ma and Al(ma) dpp species. Figure 4.17 highlights the spectra of the H and Hb 2  2  a  ring protons: in the 3:1 mixture, the signals from the maltol in the mixed-ligand species are shifted slightly upfield from those of Al(ma)3 and the signals at the chemical shift of Al(dpp)3 indicate that all of the dpp* ligands are on mixed-ligand species. The T for Al(dpp)3 in C D 3 O D is 21 °C and as the top spectrum in Figure 4.17 c  illustrates, the ligand rearrangement rate is slow at 18 °C. However, the signals from the mixed-ligand species are sharp and well resolved at this temperature. This is evidence that the ligand exchange rate is comparable to the rate of ligand rearrangement. To verify this, the 3:1 mixture was cooled (in 10° increments) to -50 °C and the variable-temperature N M R spectra of the Hb proton on the dpp' ligand are reproduced as Figure 4.18. The sharp overlapped doublets of the 18 °C spectrum are significantly broadened at -10 °C (a similar effect is observed for the other signals in the spectrum). At -20 °C the ligand exchange rate is slow enough to permit the observation of signals from the inequivalent dpp ligands on the Al(dpp)3. (ma) species. The ligand exchange rate is slower at -30 °C and this can be n  n  considered the spectrum in the absence of ligand exchange as it is unchanged at -40 and -50 °C. It is apparent that ligand exchange is occurring at a rate somewhat faster than that of the rearrangement reaction. It is necessary to cool the sample to a temperature comparable to the T of the more labile Ga(dpp)3 (-9 °C) in order to slow the ligand exchange process c  enough to observe the signals from all of the mixed-ligand species.  126  Figure 4.16. Proton N M R spectrum of 3:1 Al(ma)3 to Al(dpp)3 in CD3OD at 18 "C.  Figure 4.17. Proton N M R spectra of Ha and Hb in Al(dpp)3 (top), Al(ma)3 (middle) and the 3:1 Al(ma) to Al(dpp)3 mixture (bottom). A l l spectra in CD3OD at 18 *C. 3  128  Figure 4.18. Variable-temperature proton N M R spectra of H5 in the dpp" ligand on the mixed-ligand species (in CD3OD and temperatures are °C).  The ligand exchange experiment indicates the rearrangement reaction for Al(dpp)3 in CD3OD is an intermolecular process. The racemization of Ni(phenanthroline)3 in water 2+  proceeds through a dissociative mechanism and the similarity of the activation energies for the process in methanol (and a number of other solvents) is used to support the proposition that the same mechanism is occurring in all of the solvents studied.  158  By analogy, the  similarity of AG+Tc in D 0 and (CD3)2SO to the value in CD3OD suggests there is an 2  intermolecular exchange occurring in all of these solvents. It can be concluded on the basis of the variable-temperature N M R experiments that fac-mer isomerization is the reason for the fluxional N M R spectra of the tris(3-hydroxy4-pyridinonato) A l and Ga complexes . It appears the exchange process is intermolecular and, therefore, the isomerization mechanism likely involves ligand dissociation. In addition, the solubility properties of these A l and Ga complexes make it very unlikely that the mechanistic details of their ligand rearrangement processes could be determined.  129  Chapter V A.  5.1  Solution Studies  Aqueous Solubility  Introduction During the synthesis of the ligands and the tris-ligand metal complexes, their  interesting and at times unexpected aqueous solubilities were duly noted. This led to solid state studies on single crystals of several compounds grown from, and in the case of the metal complexes grown with, water. The crystal packing arrangements were dictated by H-bonding and we wanted to discover if this form of molecular association also affected the solution behavior of these compounds.  The accurate measurement of aqueous  solubility was one method used to establish the correlation between the solid state and solution properties. Accurate knowledge of the aqueous solubility was also of practical value as regards synthetic and biological applications. The unusual aqueous solubility properties of nitrogen heterocycles were noted as early as 1899 when it was found that the insertion of hydroxyl groups progressively decreased the water solubility of purine. This unexpected result was also observed for 7C-deficient N-heterocycles as exemplified by 4-hydroxy-2-pyridinone which is 160 times less water soluble than either 2- or 4-pyridinone. To explain this, Albert suggested that intermolecular H-bonding from the ring nitrogen to the exocyclic oxygen atom was preferred over H-bonding to water molecules.  104  The reduced aqueous solubility of the  3-hydroxy-4-pyridinones can be attributed to strong intermolecular H-bonds, but with the exception of Hmpp, the ring nitrogen was not directly involved. In comparison to the free  130  ligands, the interesting feature of the metal-ligand complexes was that their solubilities were governed by solvent rather than intermolecular H-bonding and still the result was a lowered aqueous solubility. The solubility was determined at 25 °C and in an attempt to gauge the relative strength of the H-bonds, the change in solubility upon heating to 37 °C was also measured for several compounds. In keeping with our interest in biological applicability, the measurements were done in isotonic p H 7.4 buffer. The 3-hydroxy-4-pyridinones have a TC->7t*  transition (~280 nm) and the solute concentrations were measured by monitoring  this band. Since it was employed primarily as an analytical technique (it was also used in the determination of the «-octanol/water partition coefficients), ultraviolet ( U V ) spectroscopy was included in this chapter rather than with the other spectroscopic techniques in Chapter II.  5.2 Materials and Methods  Water was deionized with Barnstead D8902 and D8904 cartridges and distilled in a Corning MP-1 Megapure still. Trizma-7.4 HC1 (tris(hydroxymethyl)aminomethane) is available from Sigma. The isotonic buffer solution was 0.15 M NaCl and 0.05 M Trizma which gives p H 7.42 (±0.05) at 25°C. Solution temperatures (+0.1 °C) were maintained with water-jacketed beakers and a Julabo circulating waterbath. Samples were centrifuged at 3300 rpm in a Centrific Centrifuge model 228. A l l aliquots were withdrawn with Eppendorf digital pipettes, either 10-100 U.L or 100-1000 p L (1-1.5% error). The electronic spectra were measured from 370 to 210 nm with a Perkin Elmer Coleman 124 U V - V i s spectrophotometer (1 cm cell). Trizma absorbs in the U V near the 210 nm instrument cut-off point so it was necessary to use buffer as the reference solution.  131  L-Mimosine was recrystallized from water; the other ligands and metal complexes were prepared and purified as previously described in Chapter JJ. A l l samples were rigorously dried prior to weighing.  5.2.1  Procedure for the Determination of Molar Absorptivity (£)  By comparison to other 3-hydroxy-4-pyridinones, a reasonable estimate of £ for 161  the free ligands was 12,000 L moHcnr for the selected band (A, ax)- Based on this, 1 1  m  m M stock solutions (in 100 m L volumetric flasks) were made from accurately weighed compound. Five 10 mL standard solutions were made (5 to 80 p.M for the free ligands, 3 to 40 p M for the metal complexes) that covered an absorbance range of 0.070 to 1.200. Plots of concentration (c) versus absorbance (A) were made and all the compounds obeyed Beer's Law (A = £bc with b = 1 cm) over these concentration ranges. Each plot had at least five points and in some cases a second set of measurements was made and all the points were included in the graph. The values of £ were calculated by linear regression analysis (correlation coefficients > 0.999). The £ values in Tables 5.1 and 5.2 have an estimated error of ± 500 L moHcm ; this is consistent with the propagation of dilution _1  and instrument errors and is < the standard deviations (a) calculated from several of the Beer's Law plots. The instrument error in  X  max  is ± 1 nm.  5.2.2 Procedure for the Determination of Aqueous Solubility at 25 and 37 °C  Suspensions of each compound in buffer were placed in the 25 °C waterbath, stirred for 30 minutes, and equilibrated for an additional hour to allow the finely suspended  132  solids to settle. Aliquots were withdrawn from the saturated solutions, filtered through glass wool into centrifuge tubes, and centrifuged for five rninutes. This step was necessary because of the large dilutions required to obtain concentrations suitable for spectroscopy; any solids in the final aliquot were redissolved upon dilution and resulted in errors in the calculated solubility. Accurate aliquots were withdrawn from the centrifuged samples and diluted with buffer to the concentrations necessary to give absorbances between 0.500 and 1.500 (the dilution factors ranged from 10 for Ga(mhpp)3 to 16,667 for Hmepp). Four measurements were made for each solution and the average absorbance (A25) and the previously determined £ values were used to calculate the concentrations. Propagation of the weighing, dilution, and instrument errors would allow the reporting of concentrations to three significant figures. The additional systematic errors inherent in solubility studies (such as the effects of small amounts of impurities) and in this procedure (such as the lack of temperature control once the samples were removed from the waterbath) limited our confidence so the concentrations were reported with two significant figures. In Tables 5.3 and 5.4, the concentrations have an estimated error of ±5 in the second digit. The concentrations of M(mpp)3, M(dpp)3, and M(mepp)3 were measured on three separate occasions and the reported values for a (r=3) supported this estimated error. For the absorbance measurements at 37 °C, the suspensions were equilibrated at the higher temperature for 30 minutes and treated as above. The absorbances (A37) were corrected for changes in the dilution factor and the percent change on heating was calculated  133 5.3 Results and Discussion 5.3.1  Ultraviolet Spectroscopy  The ultraviolet spectra of pyridine and its derivatives are similar to that of benzene; they have an E (ethylinic) and a B (benzenoid) band above 210 nm that originate from Tt—>TC*  transitions.  n->7t* transition  162  Due to the lone pair of electrons on the nitrogen, there is a weak  (R-band) that is generally only observable in the vapor phase. The B-band  in pyridine is at 257 nm (in water with £ = 2750 L moHcnr ); conjugated and/or electron1  donating substituents cause this band to shift to lower energy. 162  cause bathochromic shifts of 18 and 5 nm respectively.  104  163  O H and CH3 groups  In good agreement with this,  the B-band in the N-substituted ligands is at 280 ± 2 nm and is at a slightly shorter wavelength in Hmpp (Table 5.1). Compared to their respective ligands, the B-band in the metal-ligand complexes show a small bathochromic shift (9-16 nm) and the Ga complexes are consistendy at the longest wavelength (Table 5.2).  Table 5.1.  X  max  (nm) and £ x 10 (L moHcnr ) for the B-band in the free ligands. 1  -3  Hmpp  Hdpp  Hmepp  Hmhpp  Mimosine  H2exn  A-max  273  278  279  279  282  278  £  12.4  13.6  14.2  12.4  15.0  26.1  134  Table 5.2. Xmax (nm) and £ x IO" (L moHcnr ) for the B-band in the metal-ligand complexes. 1  3  M(m ?P)3  M(d 'P)3 Al Ga  Al  Ga  X  285  286  290  £  26.4  27.1  28.4  Al  Ga  M(m 1PP>3 Al Ga  291  292  293  293  29.4  30.7  30.8  28.8  M(mepp)3  M(mimo)3 Al  Ga  294  291  294  30.0  30.2  30.5  For pyridine derivatives, the intensity of the B-band can be greatly enhanced by auxochromic substituents (saturated groups with nonbonded electrons).  163  This  enhancement accounts for the large increase in the £ between the 3-hydroxy-4-pyridinones and pyridine. Increasing solvent polarity also has a marked hyperchromic effect on the B-band attributable to solvent H-bonding with the nitrogen lone pair of electrons and this increase in intensity is at a maximum for pyridinium salts.  162  The smaller £ values for  Hmpp and the M(mpp)3 complexes may be due to a decreased electron density on the ring nitrogen when compared to the N-substituted compounds. This could result in somewhat weaker solvent H-bonds and, therefore, a reduced hyperchromic effect. The differences are small (from 1200 to 4300 L moHcm ) but consistent; the one exception is Hmhpp and 1  it is possible that the hydrophobic N-hexyl group could disrupt the H-bonding sufficiently to account for the lower £ of this ligand.  135  5.3.2 Aqueous Solubility at 25 and 37 °C  At 25 °C, Hmpp is less than half as water soluble as Hdpp (Table 5.3).  The  replacement of a hydrophilic amino with a hydrophobic methyl group increases water solubility and intermolecular H-bonding is the likely explanation for this unusual behavior. The evidence from IR spectroscopy and the crystal structure indicated that the O - H - O hydrogen bonds in Hmpp were weaker than in Hdpp so the decrease in water solubility must be due to the N-H—0=C hydrogen bonds. The strength of the Hmpp H-bonds in solution was shown by the proton N M R experiment in which the intermolecular interactions persisted in (CD3)2SO, a good H-bond acceptor solvent. It is interesting to compare these compounds to 4-pyridinone and 4-hydroxy-2pyridinone whose water solubility (at 20 °C) is 105 and 0.65 m M respectively.  104  Despite  the presence of two methyl groups, Hdpp has a solubility comparable to 4-pyridinone. This could be due to the relative position of the H-bonding sites and, therefore, the extent of the H-bonded network. When the H-bond sites are ortho, as in Hdpp, it is possible to form dimeric units. When they are para, as in 4-pyridinone, the most likely interaction would be the formation of H-bonded chains that could reduce the relative water solubility. The addition of a second H-bond donor in 4-hydroxy-2-pyridinone presents the possibility of a three-dimensional H-bonded structure, and as with Hmpp, the result is a further reduction in water solubility. The situation is not so straightforward when Hmepp is compared to Hdpp; the substitution of an ethyl for a methyl group increases the water solubility almost sixfold. The increased volume of the unit cell and the appreciably lower melting point for Hmepp compared to that of Hdpp led us to expect the water solubility to be greater, but the magnitude of the difference is surprising. The compounds both form dimeric units and on  136  the basis of IR stretching frequencies, the H-bonds are of similar strength. The large increase in solubility must be due to a weaker crystal lattice in Hmepp; although intermolecular H-bonds are important, they are by no means the only factor in determining aqueous solubility.  Table 5.3. Water solubility (mM) at 25 °C.  Solubility  Hmpp  Hdpp  Hmepp  Hmhpp  Mimosine  38  95  570  5.6  19  H2exn 0.56  Hmhpp is minimally water soluble and, unlike the other 3-hydroxy-4-pyridinones studied, Hmhpp is soluble in a number of aprotic solvents (including diethyl ether, CH2CI2, and acetonitrile). It is apparent that the N-hexyl group, and not the potential for H-bonding, is the deciding factor in the solubility properties of Hmhpp. The results for H2exn indicate why this compound was difficult to characterize and support the conclusion that the lower water solubility is due to H-bonded polymerization. The low water solubility of L-mimosine, a compound with four H-bonding sites, strongly suggests that intermolecular H-bonding is forming a crystal lattice capable of resisting solvation by water molecules (or by any other solvents short of dilute acids and bases). The increased size of the tris-ligand metal complexes should result in a decrease in water solubility compared to that of the free ligand. However, the water solubility of the free ligands is strongly influenced by their ability to form intermolecular H-bonds via the cc-hydroxy ketone moiety and formation of the metal complexes removes this H-bonding site. It is possible that diminished intermolecular association could offset the increase in  137  size and result in greater not lesser aqueous solubility for the metal complexes. This appears to be the case with mimosine as Al(mimo)3 is five times as soluble as the free ligand (Table 5.4). The substituent zwitterion becomes the dominant factor in determining the water solubility and, given the acidity of the ammonio proton (pKa of 7 in the free ligand ), one would predict appreciable water solubility. Ga(mimo)3 is approximately 97  one-half as soluble as its Al analogue and the other ligands show a similar contrast between their Al and Ga complexes. The reduction in solubility is consistent with the larger size of the metal center, but the magnitude of the difference is not easily rationalized considering that in no instance does the metal account for more than 16% of the mass of the tris-ligand metal complex.  Table 5.4. Water solubility (mM) at 25 *C and % change f ^ m ^ S f 1 " ^ x 100  M(m 'P)3 Al Ga Sol. %  M(d 'P)3 Al Ga  16(1) 7.5(5) 42  Al  Ga  1.2(1) 0.70(2) 19(1) 11(1) 1.2  20  M(mepp)3  0.74  3.3  1.9  M(m 1PP)3 Al Ga  M(mimo)3 Al  Ga  0.69  0.20  95  55  12  3.6  500  289  The M(mimo) complexes are the exception rather than the rule and the rest of the 3  tris-ligand metal complexes have a much lower water solubility than that of the free ligand. The M(dpp)3 complexes have the maximum change as the solubility of Ga(dpp)3 is < 1% that of Hdpp (see bottom row in Table 5.4). The decrease is not as large in the M(mpp)  3  complexes (e.g., 42% for Al(mpp)3); this is probably due to the lower solubility of the free ligand and the availability of the NH group for solvent H-bonding in the metal-ligand complexes. That the M(mhpp)3 complexes are less soluble than Hmhpp is not surprising  138  since there are three lipophilic hexyl groups per molecule instead of one. The large decrease in solubility for the M(dpp)3 complexes results in solubilities on the order of those for the hydrophobic M(mhpp)3 complexes and this is rather surprising. The solid state packing arrangement presents a possible explanation for the low water solubility of the M(dpp)3 complexes. Solid water consists of hexagonal water rings interconnected in a honeycomb arrangement to give a rigid quasi-infinite network (refer to Fig. 3.9). On melting, some of the H-bonds are disrupted and the network of water molecules becomes irregular and less open than in the ice structure. It is generally accepted that there would still be an infinite network linking together many finite discrete networks of varying sizes but the exact structure of the finite units is still very much open to debate.  In the solid state, the  164  M(dpp)3 complexes are able to fit into and in a sense hold together separate water structures. A similar situation could be occurring in aqueous solution; rather than being pulled apart by the solvent molecules, the metal-ligand complexes could be incorporated into the disordered fabric of liquid water to form insoluble polymers of the inorganic complexes and the water columns. The M(dpp)3 and M(mepp)3 complexes have the same solid state packing arrangement and both show similar large decreases in water solubility when compared to the free ligands. But the M(mepp)3 complexes are ca.16 times as soluble as the M(dpp)3 complexes and they are even more water soluble than the M(mpp) complexes despite the 3  obvious difference in hydrophilicity between the N-ethyl and N - H groups. Hmepp is substantially more water soluble than Hdpp for reasons that are not readily apparent and the M(mepp)3 complexes could be forming water polymers that are simply more soluble than their M(dpp)3 analogues.  139  The difference between the M(dpp) and M(mepp)3 complexes is further shown by 3  the change in absorbance found on heating to 37 °C. The experiment was motivated by the large aqueous solubility temperature coefficient that was observed in the synthesis of the M(dpp)3 complexes and was exploited in crystal growing to make the supersaturated solutions from which the single crystals were obtained. Neither the M(mpp)3 nor the M(mepp)3 complexes appear to exhibit this property and the results of the quantitative study confirm the experimental observations. On heating to 37 °C, the absorbance of the M(dpp)3 solutions increases by nearly 200% while the M(mpp)3 and M(mepp)3 solutions show no significant change (A37 is within ± 10% of A25). (For perspective, the absorbance of all the free ligands increased by ca. 30%.) This simple experiment serves to illustrate the unique aqueous solution behavior of the M(dpp)3 complexes. The 3-hydroxy-4-pyridinone ligands synthesized in this study varied comparatively little in size or electronic configuration yet they exhibited a considerable range of water solubilities. The order of water solubility expected for the free ligands was Hmhpp < Hmepp < Hdpp < Hmpp and the order found was Hmhpp < Hmpp < Hdpp « Hmepp. This illustrates the interesting, albeit somewhat unexpected, results that can be obtained by simple N-alkyl substitution. The determination of the aqueous solubilities also made it possible to correlate the spectroscopic and crystallographic studies directly to a property of concern to our goal of developing biologically relevant ligands.  140  B.  Octanol/Water Partition Coefficients  5.4 Introduction The distribution of a solute between two immiscible phases in which it is soluble has long been a subject of investigation.  The early workers established that the  concentrations of the solute in the two phases was constant and independent of the relative solution v o l u m e s .  165  While the physical chemists strove to develop a theoretical  description of the partition phenomena, applications for this physicochemical property developed from the discovery (near the turn of the century) that the relative narcotic activities of drugs was often direcdy related to their oil/water partition coefficient.  166  In the  late 1960's, Hansen and co-workers proposed that the rc-octanol/water system was a good model for the lipoidal biophase in living organisms and, therefore, could be useful in studying the distribution of solutes between blood and lipid in living organisms.  167  In  pharmacological research, the n-octanol/water partition coefficient (P)* has been accepted as the operational definition of lipophilicity and is now widely employed in the design and development of new bioactive compounds. The measurement of log P for the Ga and In complexes was useful in establishing their potential as radiopharmaceutical imaging agents (In has two suitable radioisotopes). The increasing interest in A l chelators for the treatment of various neurological disorders  This may be represented as K nol/water (ow)> Pow or D - Log P is used in the computational methods for the estimation of P and in the numerous quantitative structure-activity relationship (QSAR) studies involving this property; therefore, log P is the most frequently (but by no means exclusivley) encountered symbol for this property. octa  o w  141  linked to a l u m i n u m ' 15  168  made it worthwhile to determine the log P values for the free  ligands and the tris-ligand A l complexes. A log P of 2 has been proposed as ideal for the design of barbituates, but the complexity of brain uptake makes it impossible to set a lower log P limit in neurological drug design.  27  Levin reported a structure-activity relationship  between the permeability of the blood-brain barrier, and the log P and molecular weight of a substance.  26  He found that drugs with log P values as low as -3.0 could penetrate the  brain in detectable amounts. The following table is intended merely to give an idea of the range in lipophilicity found for small molecules that readily enter the brain.  Table 5.5. The log P values of some common drugs.  27  ethanol  -0.31  morphine  0.07  nicotine  0.45  cocaine  1.05  caffeine  0.08  phenobarbital  1.42  codeine  0.23  amphetamine  -0.84  The log P can be measured by shaking a solute with the immiscible solvents and measuring the solute concentration in one or both of the phases. Conceptually simple, the consensus opinion is that the classic "shake-flask" method is experimentally difficult and very time c o n s u m i n g . ' 169  170  The difficulties in measuring the partition coefficient have at  times caused wide variations in the reported log P values; in a recent study, an author went so far as to eliminate from consideration all log P values that were in conflict with those reported by Hansch and his co-workers.  171  (This reflects both the high regard that is  afforded Hansch and the difficulties of measuring P.) This has led to the development of numerous alternative methods for determining log P v a l u e s .  1 7 0  '  1 7 2  '  1 7 3  The most widely  used alternative is reverse phase high performance liquid chromatography (HPLC) with  142  methanol-water as the eluent. * 170  174  The H P L C retention times are converted to log P by  means of a standard set of solutes for which the shake-flask log P values have been measured; however, this method has been criticized as being unreliable, especially for hydrophobic compounds that require high proportions of methanol (> 50%) in the eluent.  169  The log P is an additive-constitutive property of a substance and as such it can be estimated using a substituent constant, JC, defined in an analogous manner to the Hammet a constant.  165  A manual algorithm was developed that has since given way to a  computerized method of estimating log P: for a homologous series, the log P values are measured for the compounds with key structural features capable of mutual interaction and then estimated for the other members of the series. The computations are non-trivial (and expensive) and the estimated log P values are often quite different from the experimentally derived results.  Despite the problems in its execution, the shake-flask method is  considered the most accurate and reliable method for measuring log P values.  169  5.5 Materials and Methods  Reagent grade rc-octanol was distilled and the first and last quarters were discarded. The p H 7.4 isotonic buffer was prepared as described in Section 4.2. The two solvents were mutually saturated by stirring a 1:1 mixture overnight and the saturated solvents were used for all measurements. The ligands and metal-ligand complexes were purified by recrystallization or sublimation. The solute concentration in the aqueous phase was determined by monitoring the absorbance of the B-band; the spectra were recorded on the U V spectrophotometer used for the aqueous solubility study. Trial solutions made with  143  n-octanol saturated buffer indicated the small amount of n-octanol was not affecting the position or intensity of the B-band. The reference solution was n-octanol saturated buffer that had been centrifuged for the same length of time as the sample solutions. This was necessary because /i-octanol absorbs near 210 nm and a small difference in its concentration between the reference and sample solutions was found to cause serious baseline drift. The initial absorbance was kept between 0.5 and 1.0 which required concentrations of -30 p M for the metal complexes and ~60 p M for the free ligands. 1 m M stock solutions were diluted (using graduated cylinders) to make 25 mL solutions at the desired initial concentrations. A 10 mL aliquot was withdrawn and placed in a 15 mL centrifuge tube labelled as the initial solution. Two 6 mL aliquots were withdrawn and placed in centrifuge tubes labelled as extraction tubes. 6 mL of /i-octanol was added to the extraction tubes and they were inverted 100 times (>2 minutes contact time). After >15 minutes equilibration, the tubes were centrifuged until the two layers were visibly clear (typically at least 15 minutes). The n-octanol layer was removed with a pasteur pipette, and the aqueous layer was used to rinse the cuvettes and to make one absorbance reading per tube. The initial solution was also centrifuged and two absorbance readings were made. This was repeated three times per compound thus producing six measurements of the initial absorbance and six of the post-extraction value. The values of P were calculated from the absorbance of the aqueous phase as f o l l o w s  175  and the mean log P (r > 6) and a were reported.  (initial absorbance)-(post-extraction absorbance) post-extraction absorbance  x  volume of buffer volume of w-octanol  From 0 to 25 °C, the log P can vary from 0.005 to 0.01 log units/degree;  165  this is  within the error for the shake-flask method and no attempt was made to control the  144  partitioning temperature. Ideally, samples of both phases should be analyzed to check for material balance as a guard against unforeseen losses. This requires an analytical procedure for both phases which further increases the time required for the measurements. It has been shown that if care is taken to ensure that no special solute interactions are occurring, reliable results can be obtained by analyzing only one phase.  176  5.6 Results and Discussion  Once we decided on using the shake-flask method, it remained to choose an analytical technique. One option was flame atomic absorption spectroscopy (FAAS) and this technique has been used to determine the log P for a variety of Hg complexes.  177  The  viscosity of n-octanol did present problems for sample aspiration, but a procedure was developed and several A l and Ga complexes were studied (analyses done by T. Karpishin). The compound that was used to verify our F A A S procedure was Al(ma)3. Shake-flask extraction with U V spectroscopy as the analytical technique gave a log P of -0.17 for 57 p M Al(ma)3  37  (this value was reproduced by our collaborators at another laboratory ). 178  Using the F A A S procedure, the Al(ma)3 log P was -1.22 ± 0.10 for a 2.5 m M solute concentration and a similar large deviation from the value determined with U V spectroscopy was seen for Ga(ma)3. The log P was also found to be concentration dependent as the Al(ma)3 log P values ranged from -1.70 at 15 m M to -1.10 at 1.5 m M . The partitioning should be done at the lowest solute concentrations possible since P is concentration dependent and only theoretically valid at infinite dilution; however, concentrations of 10 M are considered sufficiendy dilute for neutral molecules that have _1  little tendency to associate in s o l u t i o n .  176  The Al(ma)3 log P is not constant at m M  145  concentrations and this suggests intermolecular association. The fact that the log P is an order of magnitude higher at 57 | i M than at 2.5 m M certainly supports this conclusion. To avoid errors from solute association, the partition concentrations can be decreased until no further change in log P is f o u n d .  176  This is not possible with F A A S because of the  detection limits for A l and Ga (the bottom end of the working range is 1.5 and 0.72 m M , respectively ). 179  We experimented with F A A S because of problems with the procedure we used to measure the log P values of tris(3-hydroxy-4-pyronato) A l and Ga complexes (analyses done by T.Lutz). This procedure was prone to give large fluctuations in log P values and 37  changing the analytical technique offered a possible improvement. F A A S did solve the problem of precision, but the large discrepancy from the "true" log P for Al(ma) made the 3  accuracy of the results suspect. The experiments with F A A S did alert us to the potential complication that intermolecular association presented. The large molar absorptivities of the 3-hydroxy-4-pyridinones made U V spectroscopy the analytical technique of choice for this study. (pM detection limits). The problems with our first procedure for determining log P were not rectified by changing the analytical technique so adjustments were made to the extraction steps and the revised procedure is reported herein. Ga(ma) was used to test the revised procedure and 3  the log P (in brackets in Table 5.6) was in agreement with our previously published value. The error of ±0.03 log units was calculated on the basis of two separate experiments (r=12) and it was the same as the error considered acceptable for the shake-flask method.  176  The  lower limit that can be measured with acceptable precision by this procedure is a log p of -1.75. For log P values lower than this, the difference between the initial and postextraction absorbances (initial absorbance of 1 and 1:1 extraction) is < three times the instrument error (± 0.005 absorbance units).  146  The accepted way to extend the log P range is to increase the ratio of n-octanol to buffer, thereby increasing the difference between the two absorbance readings (this also decreases the error in log P). However, extractions at these ratios formed emulsions that were very difficult to remove and the failure to get completely clear solutions is known to produce large errors in log P values.  176  We found that letting the solutions stand for > 24  hours did not remove the emulsions and even centrifuging for several hours did not always give clear solutions.  Based on the log P values of the 3-hydroxy-4-pyrone metal  complexes (Table 5.6), the lower limit on log P imposed by 1:1 extractions was acceptable as it was thought the 3-hydroxy-4-pyridinone metal complexes would have log P values within an order of magnitude of these results. We also felt that a log P of -1.75 made a reasonable cut-off point as compounds with values below this are quite hydrophilic and this would significantly limit their ability to cross cell membranes. The log P values for a number of In complexes were determined using the revised procedure.  112  When the results for the In complexes of maltol and kojic acid are included  with the log P values determined previously for the A l and Ga analogues (Table 5.6), a reasonable pattern emerges (for the structures of the 3-hydroxy-4-pyrones see Fig.2.2). With its additional hydroxyl group, kojic acid is considerably more hydrophilic than maltol. For both ligands, the metal complexes have smaller log P values than the free ligands, and the relative order of the lipophilicity of the metal-ligand complexes is the same; i.e. In > A l > Ga. The results for the tris-ligand metal complexes with pyromeconic acid and cholorokojic acid (5-hydroxy-2-(chloromethyl)-4-pyrone) were predictable: removing the ring methyl group from maltol (to give pyromeconic acid) produced a decrease in lipophilicity and replacing the methyl O H group in kojic acid with a C l atom (to give chlorokojic acid) produced an increase in lipophilicity.  147  Table 5.6. Log P values for the 3-hydroxy-4-pyrone complexes (a values are in parenthesis).  Al  ML Ga  In  - 0.17  - 0.29 [-0.22(3)]  -0.009(8)  -1.12  --  -0.62(6)  - 1.06  - 1.10  -0.82(6)  --  --  -0.31(4)  3  Free Ligand Maltol  3  0.090  Pyromeconic acid Kojic acid  - 0.64  Chlorokojic acid  a  These values are from ref. 180 and were determined by the shake-flask method.  The log P values for the 3-hydroxy-4-pyrones are presented as an example of a well behaved metal-ligand system. The log P values are reasonable for metal complexes and changes in structure produce predictable changes in the log P. By comparison, the 3-hydroxy-4-pyridinones could be called a "truculent" metal-ligand system based on their log P values (Table 5.7). With the exception of A l - and Ga(mhpp)3, the metal complexes have much lower than expected log P values. Instead of being similar to the values of the 3-hydroxy-4-pyrone metal complexes, they are at or near the lower log P limit for the shake-flask procedure employed in this study. Professor Hansch graciously agreed to measure the log P of a Ga(dpp) sample and the result reported to us (italicized entry in 3  Table 5.7) is in agreement with the value that was determined in our laboratory . The error in log P (±0.10 log units) is significantly larger than the ± 0.02 log units typical for what Hansch referred to as "well behaved compounds."  181  The tris(3-hydroxy-4-pyridinonato)  metal complexes have log P values that are lower and more difficult to measure than was predicted on the basis of the results for the 3-hydroxy-4-pyrone system.  148  Table 5.7. Log P values for the 3-hydroxy-4-pyridinone complexes. ML  3  Free Ligand  Al  Ga  In  Hmpp  - 0.52 (8)  < - 1.75  - 1.51 (6)  < - 1.75  Hdpp  - 0.74 (8)  < - 1.75  - 1.55 (20) -1.59 (10)  < - 1.75  Hmepp  - 0.37 (3)  - 1.68 (6)  - 1.64 (20)  Hmhpp  0.95 (10)  1.32 (20)  1.38 (12)  It is possible that the interaction of water molecules with the tris(3-hydroxy-4pyridinonato) metal complexes is affecting the partitioning process. The partitioning in this laboratory was done at 30 p M solute. concentrations, and Hansch used Ga(dpp)  3  concentrations from 27 to 3 uM without observing any concentration dependence. This indicates that intermolecular association is not responsible for the low log P values. The solid state and the water solubility studies demonstrated the strength of the interaction between water molecules and the metal complexes. If the metal-ligand complexes are associating with liquid water in a fashion comparable to the M(dpp) and M(mepp) solid 3  3  state structures, the reduced log P values would be reasonable as the most lipophilic portion of the complex would be shielded from the n-octanol molecules (see Chapter III). It is difficult to account for the effects of solvent-solute association on partitioning, but water 165  association comparable to that found in the solid state structures could explain both the lower water solubility and the lower log P values of these metai-ligand complexes. The free ligands were certainly better behaved than the metal complexes, but the order of lipophilicity was unexpected; instead of Hmpp < Hdpp < Hmepp < Hmhpp as predicted on the basis of the N-substituents, Hmpp was more lipophilic than Hdpp. If one  149  equates lipophilicity to a lower water solubility, Hmpp was significantly more lipophilic than Hdpp and this was thought to be the result of intermolecular H-bonding (see Table 5.3). The log P of Hmpp was measured at 52 p M and H-bonded dimerization is not considered to be a significant factor in log P measurements at concentrations below the m M level.  165  Also, if there is intermolecular association, the log P should be smaller not larger  than expected. Once again the physical properties of the homologous series of ligands did not follow the predicted order. The M(mhpp)3 complexes have log P values larger than that of Hmhpp; obviously the lipophilicity imparted by three hexyl groups dominates the partitioning process just as it governed the aqueous solubility. The M(mpp)3, M(dpp)3, and M(mepp)3 complexes have smaller log P values than their respective ligands and this indicates they would have potential as therapeutic chelating agents. The ideal chelating agent would be a ligand that is lipophilic enough to be distributed to the sites of metal accumulation and that forms metalligand complexes more hydrophilic than the free l i g a n d .  168  This would diminish  redistribution of the metal to other tissues and would expedite the elimination of the metal from the body. Taking the removal of A l from the brain as a hypothetical clinical situation, the ligands—Hmpp, Hdpp, and Hmepp—are small enough and are sufficiently lipophilic to cross the blood-brain barrier (based on Levin's structure-activity relationship ). The 26  tris-ligand A l complexes are at least an order of magnitude more hydrophilic than the free ligands and this fits the above criteria for a therapeutic chelating agent.  150  C.  Potentiometric Equilibrium Measurements  5.7 Introduction In the reaction of a trivalent metal with a bidentate ligand, a series of equilibria are established involving the metal, the ligand, and metal-ligand complexes. These equilibria are illustrated in Figure 5.1 along with the definitions of the stepwise (K ) and overall (Pn) n  formation constants. The generalized formula for the overall formation constants is p = [MLn] [M] [ L P * n  M  +  L  ML + L  ML  2  +  L  [ML]  ^2  ML  ^  ML  -p:  ML  K i = jjfffo  2  3  pSSfip,-  K  2  =  K  3  =" M ^ L T  p! = K ,  P2 = K i K  2  & = K-K2K3  Figure 5.1. Equilibrium equations and constants for the reaction of a trivalent metal with bidentate ligands.  The above equations describe the equilibria for A l and Ga in aqueous solution (neglecting mixed hydroxo complexes). For the 3-hydroxy-4-pyridinone ligands, L is the conjugate base formed by the dissociation of the hydroxyl proton, and the neutral ML3 complex was predicted to be the dominant metal-ligand species at physiological p H . Variable-pH A l N M R experiments provided a qualitative picture of the pH-dependent 2 7  151  speciation of the A l complexes that clearly supported this prediction. To obtain a quantitative evaluation of the A l speciation, and to confirm our assumption of an analogous speciation for the Ga complexes, potentiometric titrations were performed to determine the formation constants. The preparation and standardization for all the solutions and the potentiometric titration procedure are included in the Appendix as Procedure A . l . This procedure was developed by Dr. David Clevette of this laboratory and it has been reported previously.  182  The computational methods used to calculate the formation constants were developed elsewhere and are referenced as such in the procedure.  The data for Hmpp, Hdpp,  Hmepp, and Hmhpp (Table 5.8) were reported to the author by Dr. Clevette.  5.8 Results and Discussion  The 3-hydroxy-4-pyridinone ligands form A l and Ga complexes of great stability (Table 5.8); the overall stability constants P3 are all > 10  30  at 25 °C and an ionic strength of  0.15 M NaCl (isotonic). In the equilibrium calculations a hydrolysis model consisting of the species [ M ( O H ) ] ( n  3_n  ) (n = 1, 2, 3, 4) with formation constants according to ref. 25 +  was applied. The polynuclear species [Al2(OH)2] and [Al3(OH)4] were also included 4+  5+  in the calculations. In conjunction with the studies in this laboratory, Professor Staffan Sjoberg examined the equilibrium reactions of A l with Hdpp and Hmpp. He found that effects due to possible mixed A l - OH" - L" species were negligible at total ligand to 3 +  aluminum ratios of 1, 2, 3, and 5 ;  1 8 3  therefore, a total ligand to metal ratio of just greater  than three was used in all the potentiometric titrations performed in our laboratory.  152 Table 5.8. Log protonation constants (K) and log p for the equilibrium reactions of A l , n  Ga, and In with the 3-hydroxy-4-pyridinone ligands synthesized in this study (25 °C, 0.15 M NaCl).  Hmpp  Hdpp  Hmepp  Hmhpp  logKi  9.80 (1)  9.86 (3)  9.81 (2)  9.92 (2)  logK2  3.65 (1)  3.70 (1)  3.64 (2)  3.59 (1)  Al  11.87 (3)  11.91 (2)  11.75 (4)  11.51 (1)  Ga  13.34 (3)  13.17 (15)  13.15 (9)  In  13.51 (1)  13.60 (2)  13.53 (1)  Al  22.54 (3)  22.83 (2)  22.52 (5)  Ga  24.41 (1)  25.16 (15)  24.98 (11)  In  23.70 (1)  23.93 (3)  23.78 (1)  Al  32.05 (3)  32.25 (5)  32.17 (6)  Ga  32.85 (6)  34.32 (14)  33.88 (13)  In  32.76 (2)  32.93 (3)  32.80 (1)  Constant  log Pi  Metal  a  logp2  a  log P3  a  a  b  0  ~  22.49 (1) —  31.71 (3) —  The In values were included in reference to Figure 5.7. Data for the Ga and In complexes were not available due to their low water solubility.  153  The affinity of the hydroxide ion for A l and Ga has been alluded to at several points in this thesis and one of the primary objectives of this research project was the synthesis of ligands that would prevent the formation of metal hydroxides at neutral p H . The competition between the ligands and the hydroxide ion can be represented graphically by using the data in Table 5.8 to produce plots of metal speciation as a function of solution pH. Since the variation in log p values among the four ligands is minor, the Al-Hdpp n  speciation diagrams are representative of the ligands studied. In Figure 5.2, the speciation diagram for 3 m M Hdpp (top figure) shows no hydrolysis at neutral p H and [Al(OH)4]" does not occur in significant amounts (> 1% of the total Al) below p H 8.5. Even with a thousand-fold dilution in ligand concentration (bottom figure), the metal-ligand species still predominate from p H 4 to 8 and 60% of the initial A l is in Al(dpp) at physiological pH. 3  The qualitative observations of an absence of hydrolysis products in the synthesis of the tris(3-hydroxy-4-pyridinonato) A l complexes are verified by these quantitative results. The p H 4 to 9 window of hydrolytic stability for the tris-ligand A l complexes that was determined by the variable-pH A l N M R study (see Section 4.3) is confirmed by the 2 7  potentiometric titration data. Using the same concentrations of A l and Hdpp as in the N M R experiment ([Al] = 35 m M and 3:1 A l to Hdpp) it is possible to "simulate" the A l N M R 2 7  spectral results in a qualitative fashion. In Figure 5.3, several spectra from the variable-pH 2 7  A l N M R study are shown next to bar graphs of the equilibrium speciation profile  calculated at the same pH. The agreement between the two methods is very good and this comparison also serves as a verification of the assignment of the upfield A l N M R signals 2 7  to the mono- and bisligand A l species.  Figure 5.2. Speciation diagrams for 3:1 Hdpp to A l at 3 m M (top) and 3 p M (bottom) ligand concentrations (25 °C, 0.15 M NaCl).  155  Figure 5.3. Left: variable-pH A 1 N M R spectra for 0.035 M Al(dpp)3. Chemical shifts 27  [Al(OH) r (80),  (ppm) are:  4  [Al(dpp) ] (38), [Al(dpp)2(H 0)2] (26), 3  [Al(dpp)(H 0)4] (14), and [ A 1 ( H 0 ) ] 2+  2  2  6  2  3+  +  (0).  Right: A l speciation profiles calculated for 0.035 M A l using data at 25 °C and 0.15 M NaCl (L = dpp).  156  It was not possible to perform variable-pH G a N M R experiments because of the 7 1  low sensitivity of this quadrupolar nucleus and the low solubilities of the Ga complexes. Their hydrolytic stability could be inferred from the similarity of their synthesis, characterization, and physical studies data to that of the corresponding A l complexes. The determination of the formation constants establishes that the Ga-ligand complexes exhibit the same resistance to hydrolysis as their A l analogues and this is readily apparent from the speciation diagram for 3 m M Hdpp and 1 m M Ga (Fig. 5.4). The greater stability of the Ga-hydroxides and the larger log p* values of the Ga-Hdpp complexes results in a slight n  shift in the region of hydrolytic stability towards lower pH. A comparison of the A l and Ga speciation diagrams shows that the cross-over point at which the [M(H20)6] species 3+  is reduced to 50% of the total metal concentration is shifted from a p H of 2.2 to a p H of 1.4 and the [M(OH)4]" species appears at a lower p H and in greater concentration in the Ga diagram.  Figure 5.4. Speciation diagram for 3 m M Hdpp and 1 m M Ga at 25 °C and 0.15 M NaCl.  157  The 3-hydroxy-4-pyridinone ligands and their precursor maltol are amphoteric (Figure 5.5). The two stepwise protonation constants for the ligands (Ki and K ) are 2  given in Table 5.8. At the conception of this research project, it was thought that the N-alkyl groups would enhance the ability of the ring nitrogen to stabilize a positive charge. This should have resulted in more ring double bond delocalization and larger K values by 2  increasing the stability of the pyridinium resonance hybrid and the pyridinium cation. The data from the solid state studies (refer to Section 3.2) and these protonation constants both show that this anticipated effect of N-alkylation did not occur. The log (J3 values for the metal complexes also indicate the metal binding efficacy of the 3-hydroxy-4-pyridinone ligands was not significantly affected by N-alkylation.  O  O  OH  Figure 5.5. Protonation equations for 4-pyridinone ligands (X = N - R ) and maltol (X = O)  Hmpp is the 4-pyridinone ligand closest in structure to maltol and comparison of the protonation constants shows that the 4-pyrones are stronger acids. (The data for maltol were collected in 0.6 M NaCl; because the magnitude of the formation constants is 59  dependent on ionic strength, the following values for H m p p  1 8 3  are also at 0.6 M NaCl and  are slightly different from the values at 0.15 M NaCl reported in Table 5.8.)  Maltol has a  hydroxyl log K i of 8.38 while that for Hmpp is 9.58. In both heterocyclic rings there is an  158 additional protonation constant with a log K2 of -0.71 in maltol (p. = 0.5)  58  and 3.74 in  Hmpp. The 4.5 log unit difference in K2 is an effect of the ring nitrogen atom, which is better able to delocalize a positive charge than a ring oxygen. This would stabilize the dihydroxypyridinium cation in acidic solution and increase K2 for Hmpp. The difference in K i is due to the greater ability of the ring oxygen to delocalize the negative charge of the deprotonated hydroxyl oxygen. This would stabilize the conjugate base and decrease K i for maltol. The net effect is an increase in the basicity of the chelating oxygen atoms in the 3-hydroxy-4-pyridinone ligands. This results in a large enhancement of the formation constant for the metal-ligand complexes: log P3 for Al(ma)3 is 22.48 compared to a value of 30.41 for Al(mpp) . 3  100 m  • A • A  • •  A  x 50  lpMAl 50 u M Tf sites 0.1 m M citrate pH7.4; 25°C  •  ^ 8  A"  •  Hdpp edta maltol catechol  A  0 -7  -6  -3  -5 -4 log [total ligand]  -2  Figure 5.6. Plot of Al " " complexation (%) vs. log of total ligand concentration (conditions 3  1  and ligands are indicated in the legend).  Figure 5.6 is a plot that compares the metal binding affinities of several ligands regardless of denticities. Using formation constants taken from the literature for A l complexes with citrate  184  (100 pM) and transferrin (50 p M vacant sites), a simple model 29  of the metal binding capacity of human blood can be made (pH 7.4, 25 "C, 0.15 M NaCl).  159  For 1 | i M A l , the plot shows that Hdpp is more efficient (lower ligand concentration required) at complexing 100% of the A l than is E D T A tetraprotic), pyrocatechol  186  1 8 5  (which is hexadentate and  (bidentate and diprotic), and m a l t o l (bidentate and 59  monoprotic). It must be emphasized that this model is limited in an absolute sense, but it is valid in a comparative sense and it certainly illustrates the enhanced metal binding capacity of the 3-hydroxy-4-pyridinone ligands. The solid state structures of A1-, Ga-, and In(dpp)3 showed little variation in bonding parameters, although the M - O bond lengths did indicate the bonds were somewhat weaker in Al(dpp) (see Section 3.4.1). However, the Ga and In complexes have 3  significantly larger log {3 values than the A l analogues and this is illustrated by a plot of the n  log p3 values for Hdpp, Hmpp, and Hmhpp (Figure 5.7).  This graph shows that the  order of stability for the tris-ligand metal complexes is the same as the order of the first dissociation constants for the hexxaquo ions; i.e., Ga > In > A l (see Section 2.5). Using the p K of the 3-hydroxyl proton (log K i ) as a measure of basicity, a plot of p K vs. ligand a  a  (Figure 5.8) correlates exactly with the relative stabilities of the metal complexes. Hdpp is the strongest base and it forms the most stable metal complexes. Taken together, these two graphs nicely illustrate that this ligand-metal system does comply to the H S A B principle paraphrased as: "the harder acid prefers the harder base." The formation constants for the metal-ligand complexes show that the 3-hydroxy-4pyridinones are very good bidentate chelators for the group 13 metals. These quantitative results are in complete agreement with our experimental observations and they also verify the hydrolytic stability of the tris-ligand complexes as established by A 1 N M R . The 2 7  predictive capability that these data affords is particularly useful in the context of our stated goal of using in vitro techniques to assess the suitability of a ligand for in vivo studies.  160  34 34.033.533.0 32 .51 32.0 31.5  Hmpp  Hdpp  Hmepp  Figure 5.7. Plot of log P3 for the Ga, In, and A l complexes (as indicated in the legend) vs. ligand.  9.90  i  T  9.851  9.801  9.75  Hmpp  Hdpp  Hmepp  Figure 5.8. Plot of pKa of the hydroxyl proton vs. ligand (25 *C, 0.15 M NaCl).  161  Chapter VI  GaIIium-67 Biodistribution Study  6.1 Introduction Our goal of developing radiopharmaceutical imaging agents and our interest in aluminum, which has no suitable isotopes for radiolabelling studies, explain why G a is 6 7  the radionuclide that was chosen for these imaging experiments. The favorable nuclear decay properties and the relative paucity of work with this radionuclide (see Chapter I) were also factors that prompted our interest in G a . Despite the differences in electronic 6 7  configuration and ionic radius, the aqueous coordination chemistry of A l and Ga is very similar. They are only found in the +3 oxidation state in water and their aqueous chemistry is dominated by their shared property of strong Lewis acidity. These G a biodistribution 6 7  experiments with the 3-hydroxy-4-pyridinones can be considered first order approximations of the biological fate of the A l analogues. The similarity of in vitro aqueous behavior makes G a the best model available for A l . 6 7  The ligands—Hmpp, Hdpp, Hmhpp, and L-mimosine—form tris-ligand Ga complexes that are sufficiently water soluble for a biodistribution study; the solubilities range from 0.74 m M for Ga(dpp)3 to 55 m M for the mimosine complex. Ga(mhpp)3 has a log P of 1.4 and the other complexes have log P values near -1.6.  None of these  complexes has the the same combination of these properties seen in Ga-maltol (water solubility of 31 m M and log P of -0.22), but they do provide a range of lipophilicity that could result in variations in the biodistribution of G a . The formation constants for the 6 7  Ga-ligand complexes are large enough to ensure hydrolytic stability in vivo.  162  Once the ligands and tris-ligand Ga complexes were completely characterized, a preliminary screening experiment with G a was performed to determine if the 3-hydroxy6 7  4-pyridinones showed promise as imaging agents.  It was at this point that the  developmental work with maltol ceased as the image obtained for Ga-maltol was identical 67  to that of Ga-citrate, that is to say, the same as Ga-transferrin. 67  o7  187  However, the results  with Hdpp indicated that this ligand was affecting the G a biodistribution. Based on this 6 7  screening and the in vitro studies, a biodistribution study using Hdpp, Hmpp, Hmhpp, and L-mimosine was initiated. This study is still in progress; therefore, only a brief description of the methodology and a summary of results to date will be presented herein. This work was done in collaboration with Dr. Donald Lyster and a complete report of the results is forthcoming.  188  6.2 Materials and Methods  The biodistribution study was done in the Radiopharmacy at Vancouver General Hospital.  The animal work-up and the data analysis were performed by Dr. D. Lyster,  T. Rihela, and G. Webb. The ligands were synthesized and purified in our laboratory. Solutions at the proper concentrations were prepared in isotonic Trizma pH 7.4 buffer and delivered to the Radiopharmacy for radiolabelling and animal injection. The Ga-citrate 67  starting material was available by purchase. For the preliminary screening experiment, anesthetized rabbits were injected with 0.43 m M Hdpp solution containing 0.5 mCi of Ga-citrate (33 nM) and the imaging was 67  done with a Siemans Large Field of View Gamma Camera. For a control, the same dosage  163  of Ga-citrate without Hdpp was injected in a second rabbit and the images were 67  compared at various time intervals over a period of hours. The biodistribution experiments were performed in B A L B / c mice which were sacrificed at varying time intervals up to 2 days post-injection. The percent injected dose per organ (i.d./organ) values were obtained for blood, liver, spleen, stomach, kidneys, lungs, heart, and brain. The organs were placed in vials and the radiation was measured in a well gamma counter. In all the experiments, the ligand to Ga molar ratio was kept constant (~10 to 1) and each mouse was given a 0.1 m L injection containing 1 p C i of 4  6 7  G a . For example, the injection concentrations were 3.4 p M Hdpp and 0.25 n M G a . 6 7  The study was then repeated using the same amount of  6 7  G a and concentrated (near  saturation) ligand solutions. Again using Hdpp as the example, the injection concentrations were 80 m M Hdpp and 0.25 n M  6 7  G a for a molar ratio of 10 to 1. 8  The injection  concentrations of the other ligands were 1, 20, and 40 m M for Hmhpp, L-mimosine, and Hmpp respectively. 67  As a control in both the dilute and concentrated ligand studies,  Ga-citrate solutions without added ligands were also administered to a separate test  population.  6.3 Results and Discussion  The  6 7  G a biodistribution experiments with dilute ligand solutions were done to  assess the potential of these ligands as G a imaging agents. The injection concentration 6 7  of G a necessary to produce the optimal radiation is on the n M level; therefore, formation 6 7  of the tris-ligand G a complex is ensured by using p M ligand solutions. To assess the 6 7  164  efficacy of an imaging agent, it is desirable to keep the ligand concentration at a minimum because of the influence that a large excess of ligand can have on the biodistribution. Metal ions are thought to cross membranes by one of two generalized processes: active transport requiring energy dependent ion pumps or passive transport; i.e., adsorption onto the cell membrane followed by diffusion into the c e l l .  189  The latter process may be  facilitated by ligands that could aid the transport of the hydrophilic metal ion across the hydrophobic membrane. This can be envisaged as passage from the aqueous medium to the interior of the cell along a ligand cascade of increasing binding strength.  190  Therefore,  the abundance and strength of extracellular metal-ligand binding affects metal uptake which is clearly demonstrated by the effect organic ligands can have on metal toxicity.  189  The  objective of cellular accumulation could be adversely affected by a large excess of the imaging agent itself. This was of particular importance with these ligands since, with the exception of Hmhpp, the tris-ligand Ga complexes were significantly less lipophilic than was considered ideal. The preliminary screening had indicated that the ligands were affecting the G a biodistribution with p M ligand injection concentrations. To minimize 6 7  potential interference with cellular uptake, the same ligand to metal ratio was maintained in the dilute ligand experiments. The assessment of these ligands as chelating agents was done with concentrated ligand solutions. The objective of a chelating agent is the removal of a metal from the body and high concentrations of ligand in the blood can only facilitate this goal. In long term experiments, high levels can be maintained by repeatedly administering the ligand. We were conducting experiments of a shorter duration and were interested in determining if injection of saturated ligand solutions would be sufficient to produce a change in biodistribution. If this were to occur, it would be good evidence of the potential of these ligands as chelating agents.  165  The rjreUminary results of this study can be summarized be examining the uptake of 6 7  G a (as percent of total G a ) in the liver at 24 hours post-injection. The bar graph (Fig. 67  6.1) contains the data for the four ligands, both the dilute and concentrated experiments, and for the Ga-citrate control. The liver is the principle organ where transferrin 67  accumulates; therefore, any alteration from the biodistribution of Ga-transferrin can be 67  most readily observed in this organ. The uptake of Ga-citrate (taken to be that of Ga67  transferrin) is shown in the last column. It is obvious that these 3-hydroxy-4-pyridinones significantly reduce the uptake of  6 7  G a in both the dilute and concentrated ligand  experiments. This confirms in vivo what the in vitro experiments had shown—the 3hydroxy-4-pyridinones are very good bidentate chelators of Ga..  10  LIGAND • I  Dilute  Figure 6.1. Liver biodistribution plotted as  ES3 Concentrated 6 7  G a uptake vs. ligand for the dilute and  concentrated ligand experiments. These results indicate that these 3-hydroxy-4-pyridiones will not be useful as imaging agents; although they redirect G a from transferrin at concentrations that would 6 7  166  be viable for further development, they are not localizing the radionuclide in any particular organ. The G a uptake in the other organs examined is similar to that of the liver, and the 6 7  levels of G a are uniformly reduced from that of Ga-citrate. It appears that the ligands 6 7  67  enhance the removal of the metal from the body, presumably via the urine; but this cannot be definitely proven until this study is completed. It was thought that the lipophilic Hmhpp complexes might have a different biodistribution. This did not occur, although it is possible that Ga-Hmhpp may be removed through the hepatobiliary tract. This also has 67  yet to be determined. We intend to exploit the ability of this class of ligands to redirect 6 7  G a from transferrin. This will be done by synthesizing additional 3-hydroxy-4-  pyridinone ligands with other substituents that could alter their biodistribution. The further decrease in G a uptake seen for the concentrated ligand experiments is 6 7  especially encouraging when viewed in concert with the positive results from the extensive studies of Hdpp as an iron chelating agent.  39  The fact that the two trivalent metal ions  closest in size to A l can both be mobilized by Hdpp definitely indicates the potential of the 3-hydroxy-4-pyridinones as A l chelators. The results are even more impressive when the differences in transferrin binding constants are considered since the affinity of transferrin for A l is ca. eight and ten orders of magnitude less than for Ga and Fe, respectively. This in vivo study also allows us to assess the relative merits of the methodology that we have been employing in our laboratory. The log P values for these ligands were significantly lower than ideal and this is undoubtedly a factor in the rapid elimination of the metal-ligand complexes from the body. Knowledge of the difference in log P values between the ligand and the metal complexes is very useful since the biodistribution results further indicate their potential as chelating agents. The aqueous solubility data were used in the design of this biodistribution experiment and this simple determination merits inclusion in future studies.  167  The formation constants can be employed in a manner complementary to experiments such as these biodistribution studies. Figure 6.2 is a plot that accounts for the dilution of the ligand and the G a in the circulatory system. It indicates what ligand 6 7  injection concentration (in this case, Hdpp) is necessary to compete effectively with transferrin, using the transferrin binding constants and concentrations of human blood serum. This graph demonstrates both the utility of this model and the difficulties encountered in applying any simple model to the complex interactions in vivo. The model predicts that 80 m M Hdpp used in the concentrated ligand experiment is sufficient to complex almost all of the G a and this is supported by the results of the liver uptake of 6 7  6 7  G a . The model also predicts that p.M Hdpp should not alter the biodistribution, but the  animal study indicates some redirection from transferrin. It is necessary to improve the model by including other factors such as Fe concentrations and concentrations of metal binding proteins such as albumin. This is not an easy task but the predictive capacity shown by the simple model definitely indicates this is a goal worth pursuing.  100 -i  80 -  % o f Ga complexed with Hdpp  1 picoM Ga 0.1 m M citrate 50 p;M Tf sites pH7.4; 25°C  60 40 20 H  • • • •  —r20  —r~ 40  60  80  100  120  m M Hdpp injection concentration Figure 6.2. Graph of %Ga complexed as a function of Hdpp injection concentration (conditions as indicated in legend).  168  Chapter VII  Conclusion and Suggestions for Future Work  The primary objective of our research is the investigation of the aqueous coordination chemistry of A l and Ga. We are also interested in biological applications of this chemistry with the dual goals of developing G a radiopharmaceuticals and using G a 6 7  6 7  as a model for the biodistribution of A l . The preceding chapter addresses the biological aspects of our work and it serves as both the conclusion for that part of our research and as an indicator of the direction our future research in that area will take. The core of my research was the synthesis and characterization of several 3-hydroxy-4-pyridinone ligands and their A l and Ga complexes. The ligand synthesis initially involved the use of a benzyl blocking group, but because of the length of the procedure, a preparation was developed using buffered solutions to control the ionic state of the starting materials, maltol and the primary amines. The buffered preparations pointed us towards the next logical step in this synthesis and this work is now being carried out in our laboratory. The formation of Al(ma)3 enhances the reactivity of the complexed maltol ligands to nucleophilic attack. The electropositive metal acts to remove electron density from the 4-pyrone ring and it is well established that electron-withdrawing groups facilitate the conversion of 4-pyrones to 4-pyridinones. By the reaction of Al(ma)3 with an excess of methylamine, Al(dpp)3 can be produced almost quantitatively. This metal template reaction is now being used in the synthesis of other 3-hydroxy-4-pyridinone ligands. We attempted the synthesis of multidentate ligands and the bispyridinone H2exn was the initial product of this work. However the solubility properties of this compound restricted its use as a ligand. Multidentate ligands with the cc-hydroxyketone moiety could  169  be made by joining two 4-pyridinone rings via an amide linkage and this should afford increased aqueous solubility.  3-Hydroxy-2-methyl-l-(P-ethylamino)-4-pyridinone,  formed by the conversion of maltol with ethylenediamine, was prepared and isolated in our laboratory. Using a dicarboxylic acid (such as butanedioic or pentanedioic acids), it would be possible to link two of these ligands together via their N-ethylamine substituents. The coupling agent dicyclohexylcarbodiimide (DCC) could be used to form the peptide bonds and the length of the bridge between the rings could easily be varied as aliphatic dicarboxylic acids and D C C are readily available commercially. One of the most interesting aspects of our work was the unusual solid state structures that were determined for the M(dpp)3»12H20 complexes.  Because of the  extensive water network in these structures, attempts were made to grow crystals from water of the other metal complexes.  This culminated in the determination of the  M(mepp)3»12H20 structures. The stability of the water network was somewhat unexpected and a close analysis of the structures suggests that a slighdy larger N-alkyl group would disrupt its formation. The synthesis of 3-hydroxy-2-methyl-l-propyl-4-pyridinone is now being carried out; if the metal complexes of this ligand are sufficiently water soluble, attempts will be made to grow crystals to determine whether the water networks can still be formed with this larger N-substituent. Also, variable temperature potentiometric titrations will be performed to establish the enthalpic and entropic contributions of this water network to the overall thermodynamic stability of the M(dpp)3 complexes. The variable-temperature N M R experiments with A l - and Ga(dpp)3 established that these metal complexes undergo fac-mer isomerization. Ligand exchange experiments with Al(dpp)3 and Al(ma)3 indicated that the ligand rearrangement process likely involved a dissociative mechanism. Further ligand exchange studies are planned to determine the lability of several of the tris-ligand metal complexes in aqueous solution. If the exchange  170  rates permit N M R detection, the ligand exchange between the tris(3-hydroxy-4-pyridinato) A l complexes and Al-transferrin is the ultimate goal of this research. The affinity of the 3-hydroxy-4-pyridinone ligands for A l and Ga made the synthesis of the metal complexes relatively straightforward. The stability of the metal complexes is due to the basicity of both the oxygen atoms of the bidentate ligand. The ring nitrogen is primarily responsible for this base strength because of the stability of the pyridinium resonance hybrid and pyridinium cation. The hydrolytic stability of the A l complexes was established by variable-pH A 1 N M R experiments, and potentiometric 27  equilibrium measurements showed the thermodynamic stability of the tris ligand A l and Ga complexes. The favorable properties of these ligands reported in this thesis ensure this ligand system will be the basis of further research in our group.  171  References  1.  Driscoll, C. T. Environ. Health Perspect. 1985, 63, 93.  2.  Aluminum in the Canadian Environment; Havas, M . , Jaworski, J . F., Eds.;  3.  Cronan, C.S.; Schofield, C. L . Science 1979,204, 304.  4.  Alfrey, A . C. Adv. Cline. Chem. 1983,23, 69.  5.  Krishnan, S. Can. Res. 1988, 32.  6. 7. 8.  National Research Council of Canada, Assoc. Committee Sci. Criteria Environ. Quality, Rep. No. N R C C 24759,1986; p 331.  Sorenson, J. R. J.; Campbell, I. R.; Tepper, L. B.; Lingg, L . B. Environ. Health  Perspect. 1974, 8, 3.  Macdonald, T. L.; Martin, R. B. TIBS 1988,13, 15. Krueger, G . L.; Morris, T. K.; Suskind, R. R.; Widnek, E. M . CRC Crit. Rev. 1984,13, 1.  Toxicol.  9.  Alfrey, A . C ; Legendre, G . R.; Kaehny, W. D. N. Eng. J. Med. 1976,294, 184.  10.  Perl, D. P.; Gajdusek, D. C ; Garutto, R. M . ; Yanagihara, R. T.; Gibbs, C. J .  11.  Crapper, D. R.; Krishnan, S. S.; Quittkat, S. Brain 1976, 99, 67.  12.  Crapper, D. R.; Krishnan, S. S.; Dalton, A . J . Science 1973,180, 511.  13.  Crapper, D. R.; DeBoni, U . in Aluminum Neurotoxicity; Liss, L., Ed.; Pathotox Publishers: Illinois, 1980; pp 3-15.  14.  Dolken, A . Arch. Exp. Pathol. Pharmakol. 1897,40, 58.  15.  Llobet, J. M . ; Domingo, J. L.; Gdmez, M . ; Tomds, J. M . ; Corbella, Pharmacol.  16.  Edwards, C. L.; Hayes, R. L. J. Nucl Med. 1969,10, 103.  17.  Hoffer, P. /. Nucl. Med. 1980,21, 394.  18.  Hayes, R. L.; Huber, K. F. Metal Ions Biol. Syst. 1983,16, 279.  19.  Tsan, M . F.; Scheffel, U . /. Nucl. Med. 1986, 27, 1215.  20.  Hoffer, P. /. Nucl. Med. 1980,21, 484; Int. J. Nucl. Med. Biol.  21.  Tsan, M . B. /. Nucl. Med. 1985,26, 88.  Science 1982, 217, 1053.  Toxicol. 1987, 60, 280.  1981, 5, 243.  172 22.  Green, M . A . ; Welch, M . J.; Mathias, C. J.; Fox, K. A . A . ; Knabb, R. M . ; Huffman, J . C. J. Nucl Med. 1985, 26, 170.  23.  Taliaferro, C. H . ; Martell, A . E . Inorg. Chem. 1985,24, 2408 and references therein.  24.  Taliaferro, C. H . ; Martell, A . E . Inorg. Chim. Acta 1985,85, 9 and references therein.  25.  Baes, C. F.; Mesmer, R. E. The Hydrolysis of Cations; Wiley: New York, 1976; pp 112-123, 313-327.  26.  Levin, V . A . /. Med. Chem. 1980, 23, 682.  27.  Hansch, C ; Bjorkroth, J . P.; Leo, A . /. Pharm. Sci. 1987, 76, 663.  28.  Kulprathipanja, S.; Hnatowich, D. J.; Beh, R.; Elmaleh, D. Int. J. Nucl. Med. Biol. 1979, 6, 138.  29.  Martin, R. B.; Savory, J.; Brown, S; Bertholf, R.L.; Wills, M . R. Clin. Chem. 1987, 33, 405.  30.  Shannon, R.D. Acta. Crystallogr. 1976, A32, 75.  31.  Harris, W. R.; Pecoraro, V . L. Biochemistry 1983, 22, 292.  32.  Pierpont, C. G.; Buchanan, R. M . Coord. Chem. Rev. 1981,38, 45.  33.  Hancock, R. A . ; Orszulik, S. T. Polyhedron 1982,1, 313.  34.  Borgias, B. A . ; Barclay, S. J.; Raymond, K. N . /. Coord. Chem . 1986,15, 109.  35.  Veca, A . ; Dreisbach, J . H . /. Chem. Ed. 1988, 65, 108.  36.  Rajan, K . S.; Mainer, S.; Davis, J . M . Bioinorg. Chem. 1978, 9, 187.  37.  Finnegan, M . M . ; Lutz, T. G.; Nelson, W . O.; Smith, A . ; Orvig, C. Inorg. Chem. 1987,26,2171.  38.  McLachlan, D. R. Neurobiol. Aging 1986, 7, 525.  39.  Kontoghiorghes, G. J.; Aldouri, M . A . ; Hoffbrand, A . V.; Barr, J.; Wonke, B.; Kourouclaris, T.; Sheppard, L . Br. Med. J. 1987,295, 1509.  40.  Albert, A . Selective Toxicity, 7th. ed.; Chapman and Hall: New York, 1985; pp 430-489.  41.  Lerch, J . U . Monatshefte 1884, J , 407.  42.  Boulton, A.J.; McKillop, A . in Comprehensive Heterocyclic Chemistry Volume 2; Katritzky, A . R., Rees, C.W., Eds.; Pergamon Press: New York, 1984; pp 1-27.  173 43.  Brody, F.; Ruby, P. R. in The Chemistry of Heterocyclic Compounds Volume 14: Pyridine and its Derivatives Part One; Klingsberg, E.,Ed.; Interscience: New York, 1960; pp 99-589.  44.  Meislich, H . in The Chemistry of Heterocyclic Compounds Volume 14: Pyridine and its Derivatives Part Three; Klingsberg, E., Ed.; Interscience: New York, 1962; pp 509-890.  45.  Elkaschef, M . A - F . ; Nosseir, M . H . /. Am. Soc. Chem. 1960, 82, 4344.  46.  Brown, R. D. /. Chem. Soc. 1951, 2670.  47.  Peratoner, A . ; Tamburello, A . Gazz. Chim. Ital. 1906,36, 33.  48.  Yabuta, J . /. Chem. Soc. 924, 575.  49.  Bickel, A . F. J. Am. Chem. Soc. 1947, 69, 1801.  50.  Adams, R.; Jones, V . V . /. Am. Chem Soc. 1947, 69, 1803.  51.  Campbell, K. N . ; Ackerman, J. F.; Campbell, B. K . /. Org. Chem. 1950,15, 337.  52.  Adams, R.; Johnson, J . L . /. Am. Chem. Soc. 1949, 71, 705.  53.  Kleipool, R. J . C ; Wibaut, J . P. Rec. Trav. Chim. Pays Bos 1950, 69, 1041.  54.  Heyns, K.; Vogelsang, G . Chem. Ber. 1954, 87, 1377.  55.  Spenser, I. D.; Notation, A . D . Can. J. Chem. 1962,40, 1374.  56.  Harris, R. L . N . Aust. J. Chem. 1976,29, 1329.  57.  March, J. Advanced Organic Chemistry; McGraw H i l l : New Y o r k , 1977; p 258.  58.  Chous, G.; Benoit, R. L. /. Org. Chem. 1967, 32, 3974.  59.  Hedlund, T.; Ohman, L.-O. Acta Chem. Scand. in press.  60.  Nelson, W . O.; Karpishin, T. B.; Rettig, S. J.; Orvig, C. Can. J. Chem. 1988, 66, 123.  61.  Chem. Abstr. 1985,102, 113315Z. (patent)  62.  Severin, Th.; Loidl, A . Z. Lebensm. Unter.-Forsch, 1976,161, 119.  63.  Kontoghiorghes, G.J. Ph.D. Thesis, University of Essex, U . K. , 1982.  64.  Yasue, M . ; Kawamura, N . ; Sakakibara, J. Yakugaku Zasshi, 1970, 90, 1222.  65.  Jakopcic, K.; Tamhina, B.; Zorko, F.; Herak, M.J. /. Inorg. Nucl. Chem. 1977, 39, 1201.  66.  Looker, J . H.; Cliffton, M . D. /. Heterocycl. Chem. 1986,23, 5.  174 67.  Kontoghiorghes, G . J . ; Sheppard, L . Inorg. Chim. Acta 1987,136, L l l .  68.  Cavalier, L . F. Chem Rev. 1947, 41, 525.  69.  Katritzky, A . R.; Jones, R. A . /. Chem Soc. 1960, 2947.  70.  Cook, D . Can. J. Chem. 1963, 41, 515,.2575.  71.  Batts, B. D.; Spinner, E. Aust J. Chem. 1969,22, 2581.  72.  Bellamy, L . J. The Infrared Spectra of Complex Molecules Volume 1,3rd. ed.;  73.  Chapman and Hall: New York, 1975; p 433.  Odinokov, S. E.; Nabiullin, A . A . ; Mashkovsky, A . A . ; Glazunov, V . P.  Spectrochim. Acta 1983, 39A, 1055.  74.  Bellamy, L . J. The Infrared Spectra of Complex Molecules Volume 2, 2nd. ed.;  75.  Bellamy, L . J . ; Rogasch, P. E. Proc. Roy. Soc. (A) 1960,257, 98.  76.  Katritzkky, A . R.; Taylor, P. J. in Physical Methods in Heterocyclic Chemistry  77.  Hadzi, D.; Sheppard, N . Proc. Roy. Soc. (A) 1953, 216, 247.  78.  White, R. F. M . ; Williams, H . in Physical Methods in Heterocyclic Chemistry  79.  Cox, R. H . ; Bothner-By, A . A . /. Phys. Chem. 1969, 73, 2465.  80.  Penfold, B. R. Acta Crystallogra. 1953,6, 591.  81.  Porter, Q. N . Mass Spectrometry of Heterocyclic Compounds, 2nd ed.; Wiley:  82.  Undheim, K.; Hurum, T. Acta Chem.Scand. 1972, 26, 2075.  83.  Maquestiau, A . ; van Haverbeke, Y.; de Meyer, C ; Katritzky, A . R.; Cook, M . J . ; Page, A . D. Can. J. Chem. 1975, 53, 490.  84.  Scarrow, R. C ; White, D. L.; Raymond, K . N . /. Am. Chem. Soc. 1985,107, 6540.  85.  Pearson, R. G . /. Am. Chem. Soc. 1963, 85, 3533.  86.  Chapman and Hall: New York, 1980; pp 240-280.  Volume TV', Katritzky, A . R., Ed.; Academic Press: New York, 1971; pp 266-432.  Volume TV; Katritzky, A . R., Ed.; Academic Press: New York, 1971; pp 177-216.  New York, 1985; pp 603-663..  Pearson, R. G . /. Chem. Ed. 1987, 7, 561.; Pearson, R. G . Inorg. Chem. 1988,  27, 734.  87.  Cotton, F. A.; Wilkinson, G. Advanced Inorganic Chemistry, 5th ed.; Wiley: New York, 1988; pp 209.  88.  Greenwood, N . N . ; Earnshaw, A . Chemistry of the Elements; Pergamon Press: Oxford, 1984; pp 250-256.  175 89.  Martin, R. B. Clin. Chem. 1986,32, 1797.  90.  Secco, R ; Venturini, M . Inorg. Chem. 1975,14, 1978.  91.  Motekaitis, R. J.; Martell, A . E. Inorg. Chem. 1984,23, 18.  92.  Letkeman, P.; Martell, A . E.; Motekaitis, R. J. /. Coord. Chem. 1980,10, 47.  93.  Martin, R. B. J. Inorg. Biochem. 1986,28, 181.  94.  Gregor, J . E.; Powell, H . K. J. Aust. J. Chem. 1986,39, 1851.  95.  Scarrow, R. C.; Riley, P. E.; Abu-Dari, K.; White, D.L.; Raymond, K. N . Inorg. Chem. 1985,24, 954.  96.  Tamhina, B.; Herak, M . J.; Jakopcic, K. /. Less-Common Met. 1973,33, 289.; Herak, M . J.; Tamhina, B.; Jakopcic, K. /. Inorg. Nucl. Chem. 1973,35, 1665.  97.  Stunzi, H.; Perrin, D. D.; Teitei, T.; Harris, R. L . N . Aust. J. Chem. 1979,32, 21.  98.  Stunzi, H.; Harris, R. L . N . ; Perrin, D. D.; Teitei, T. Aust. J. Chem. 1980,33, 2207.  99.  Tsai, W. C.; Ling, K . - H . /. Chin. Biochem. Soc. 1973,2, 70.  100.  Kontoghiorghes, G. J. Inorg. Chim. Acta 1987,135, 145 and references contained therein.  101.  Nelson, W. O.; Karpishin, T. B.; Rettig, S. J.; Orvig, C. Inorg. Chem. 1988,27, 1045.  102.  Nakamoto, K. Infrared and Raman Spectra of Inorganic and Coordination  103.  Kido, H . ; Saito, K. /. Am. Chem. Soc. 1988,110, 3187.  104.  Albert, A . Heterocyclic Chemistry an Introduction; Athlone Press: London, 1959; p 424.  105.  Klopman, G. Chemical Reactivity and Reaction Paths; Klopman, G., Ed.; Wiley:  106.  Pauling, L. The Nature of the Chemical Bond, 3rd ed.; Cornell University Press: Ithaca, N.Y., 1960; pp 221-264. '  107.  Scarrow, R. C ; Raymond, K. N . Manuscript submitted to Inorg. Chem.  108.  Novak, A . Struct. Bond. (Berlin) 1974,18, 111.  109.  Jeffrey, G. A . ; Takagi, S. Acc. Chem. Res. 1978,11, 264.  110.  Evans, R. C. An Introduction to Crystal Chemistry, 2nd ed.; Cambridge Univ. Press: Cambridge, 1964; pp 283-301.  111.  Nelson, W . O.; Rettig, S. J.; Orvig, C. /. Am. Chem. Soc. 1987,109, 4121.  Compounds, 3rd. ed.; Wiley-Interscience: New Y o r k , 1978; pp 249-311.  New York, 1974; pp 55-166.  176 112.  Matsuba, C. A . ; Nelson, W . O.; Rettig, S. J.; Orvig, C. Inorg. Chem. 1987,27, 3935.  113..  Nelson, W . O.; Orvig, C ; Rettig, S. J.; Trotter, J. Can. J. Chem. 1988, 66, 132.  114.  Falk, M . ; Knop, O. in Water: A Comprehensive Treatise Volume 2; Franks, F., Ed.; Plenum Press: New York - London, 1973; pp 55-113.  115.  Peterson, D. W.; Levy, H . A . Acta Crystallogr. 1957,10, 70.  116.  Kamb, B. in Structural Chemistry and Molecular Biology, Rich, A . , Davidson, N . , Eds.; W. H . Freeman: San Francisco, 1968; pp 507-542.  117.  Saenger, W . Nature (London) 1979,279, 343.  118.  Saenger, W . Nature (London) 1979,280, 848.  119.  Saenger, W.; Lindner, K. Angew. Chem. Int. Ed. Engl. 1980,19, 398.  120.  Neidle, S.; Berman, H . M . ; Shieh, H . S. Nature (London) 1980, 288, 129.  121.  Betzel, C ; Saenger, W.; Hingerty, B.; Brown, G . M . /. Am. Chem. Soc. 1984, 106, 7545.  122.  Zabel, V . ; Saenger, W.; Mason, S.A. /. Am. Chem. Soc. 1986,108, 3664.  123.  Del Bene, J.; Pople, J . A . /. Chem. Phys. 1970,52, 4858; 1973,58, 3605.  124.  Jeffrey, G . A . ; McMullan, R. K . Prog. Inorg. Chem. 1967, 8, 43.  125.  Jeffrey, G . A . Acc. Chem. Res. 1969, 2, 344.  126.  Wells, A . F. Structural Inorganic Chemistry, 5th ed.; Clarendon Press: Oxford, 1984; p 660.  127.  Mak, T. C. W . /. Chem. Phys. 1965,43, 2799.  128.  Akitt, J . W . Ann. Rep. NMR Spectrosc. 1972,5A, 465.  129.  Harris, R. K. Nuclear Magnetic Resonance Spectroscopy; Pitman: London , 1983; pp 66-94, 118-142, Appendix 2.  130.  Hinton, J. F.; Briggs, R. W . in NMR and the Periodic Table; Harris, R. K., Mann, B.E., Eds.; Academic Press: London, 1978; pp 279-308.  131.  Delpuech, J. J. in NMR of Newly Accessible Nuclei Volume 2; Laszlo, P., Ed.; Academic Press: London, 1983; pp 153-195.  132.  Karlik, S. J.; Elgavish, G. A . ; Pillai, R. P.; Eichhorn, G. L . /. Magn. Reson. 1982, 49, 164.  133.  Karlik, S. J.; Elgavish, G . A . ; Eichhorn, G . L . /. Am. Chem. Soc. 1983,105, 602.  134.  Greenaway, F. T. Inorg. Chim. Acta 1986, 116, L21.  177 135.  Bertsch, P. M . ; Barnhisel, R. I.; Thomas, G.W.; Layton, W . J.; Smith, S. L. Anal. Chem. 1986,58, 2583.  136.  Akitt, J. W.; Milic, N . B. J. Chem. Soc, Dalton Trans. 1984, 981 and references therein.  137.  Llinas, M . ; De Marco, A . /. Am. Chem. Soc, 1980,102, 2226.  138.  Jaber, M . ; Bertin, R ; Thomas-David, G. Can. J. Chem. 1977, 55, 3689.  139.  Karlik, S. J.; Elgavish, G. A . ; Eichhorn, G. L. Inorg. Chem. 1983,22, 525.  140.  Akitt, J . W.; Mann, B. E. /. Magn. Reson. 1981,44, 584.  141.  Benn, R.; Ruflinska, A . ; Janssen, E.; Lehmkuhl, H . Organometallics 1986,5, 825.  142.  Girgis, A . Y.; Fay, R. C. /. Am. Chem. Soc. 1979, 92, 7061.  143.  Gordon, J . G.; Holm, R. H . /. Am. Chem. Soc. 1970, 92, 5319.  144.  Pignolet, L . H . Top. Curr. Chem. 1975,56, 91.  145.  Henderson, D. E.; Saltzman, J . J.; Uden, P. C ; Cheng, Z. Polyhedron 1988, 7, 369.  146.  Fay, R. C ; Piper, T. S. /. Am. Chem. Soc. 1963, 85, 500.  147.  Hatakeyama, Y.; Kido, H.; Harada, M . ; Tomiyasu, H.; Fukutomi, H . Inorg. Chem. 1988, 27, 992.  148.  Gennaro, M . C ; Mirti, P.; Casalino, C. Polyhedron 1983,2, 13.  149.  Eaton, S. S.; Hutchinson, J . R.; Holm, R. H . ; Muetterties, E. L. /. Am. Chem. Soc. 1972, 94, 6411.  150.  Eaton, S. S.; Eaton, G. R.; Holm, R. H . ; Muetterties, E. L. /. Am. Chem. Soc. 1973,95, 1116.  151.  Binsch, B. G.; Kleier, D. A . , Quantum Chemistry Program Exchange QPCE, Program No. 165; D N M R 3 , Chemistry Department, Indiana University.  152.  Gutowsky, H . S.; Holm, C. H . /. Chem Phys. 1956,25, 1228.  153.  Fay, R. C ; Piper, T. S. Inorg. Chem. 1964,3, 348.  154.  Glasstone, S.; Laidler, K.; Eyring, H . The Theory of Rate Processes; McGrawH i l l : New York, 1941; p 190.  155.  Allerhand, A . ; Gutowsky, H . S.; Jonas, J.; Meinzer, R. A . /. Am. Chem. Soc. 1966, 88, 3185.  156.  Pinnavaia, T. J.; Sebeson, J . M . ; Case, D. A . Inorg. Chem. 1969,8, 644.  157.  Grossman, D. L.; Haworth, D. T. Inorg. Chim. Acta 1984, 84, L17.  178 158.  Basolo, F.; Pearson, R. G. Mechanisms of Inorganic Reactions, 2nd ed.; Wiley: New York, 1967; pp 300-334.  159.  Fay, R. C.; Piper, T. S. /. Am. Chem.Soc. 1962, 84, 2303.  160.  Hutchinson, J. R.; Gordon, J . G.; Holm, R. H . Inorg. Chem. 1971,10, 1004.  161.  Imafuku, K.; Takahashi, K.; Matsumura, H . Bull. Chem. Soc. Jpn. 1979,52, 111.  162.  Silverstein, R. M . ; Bassler, G. C.; Morrill, T. C. Spectrometric Identification of Organic Compounds, 4th ed.; Wiley: New York, 1981; pp 305-331.  163.  Mason, S. F. in Physical Methods in Heterocyclic Chemistry, Volume II;  164.  Geiger, A . ; Mausbach, P.; Schnitker, J. in Water and Aqueous Solutions; Neilson, G.W., Enderby, J. E.; Eds.; Adam Hilgef: Bristol, 1986; pp 15-30.  165.  Leo, A . ; Hansch, C.; Elkins, D. Chem. Rev. 1971, 71, 525.  166.  Meyer, H . Arch. Exptl. Pathol. Pharmakol. 1899,42, 110.  167.  Helmer, F.; Hansch, C.; Kiehs, K. Biochemistry 1968, 7, 2858.  168.  Yokel, R. A . ; Kostenbauder, H . B. Toxicol. Appl. Pharmacol. 1987, 91, 281.  169.  Leo, A . J . /. Pharm. Sci. 1987, 76, 166.  170.  Partition Coefficient Determination and Estimation; Dunn IV, W . J., Block, J. H . ,  171.  Dunn III, W . J.; Koehler, M . G.; Grigoras, S. /. Med. Chem. 1987, 30, 1121.  172.  Clarke, F. H . ; Cahoon, N . M . ; /. Pharm. Sci. 1987, 76, 611.  173.  Katritsky, A . R.; Ed., Academic Press: New York & London, 1963; pp 1-84.  Pearlman, R.S., Eds.; Pergamon Press: Oxford, 1986; p 154.  Gobas, F. A . P. C ; Lahittete, J. M . ; Garofalo, G.; Shiu, W. Y . ; Mackay, D .  /. Pharm. Sci. 1988, 77, 265.  174.  De Kock, A . C ; Lord, D. A . Chemosphere 1987,16, 133.  175.  Herz, W . "Der Verteilungssatz", Ferdinand Enke, Stuttgart 1909, p 5.  176.  Purcell, W . P.; Bass, G . G.; Clayton, J . M.Strategy of Drug Design: a guide to  177.  Arnold, A . P.; Canty, A . J.; Moors, D. W.; Deacon, G . B. /. Inorg. Biochem.  biological activity; Wiley: New York; Appendix I pp 126-143. 1983,19, 319.  178.  McLachlan, D. R. Personal communication.  179.  Bassett, J.; Denney, R. C ; Jeffery, G . H . ; Mendham, J . Vogel's Textbook of Quantitative Inorganic Analysis, 4th ed.; Longman: London and New York, 1985; pp 834-835.  179 180.  Kontoghiorghes, G . Inorg. Chim. Acta 1988, 257, 101.  181.  Hansch, C. Personal communication.  182.  Lutz, T. G.; Clevette, D. J.; Rettig, S. J.; Orvig, C. Submitted for publication.  183.  Clevette, D. J.; Nelson, W. O.; Nordin, A . ; Orvig, C ; Sjoberg, S. Submitted for publication.  184.  Ohman, L . -O. Inorg. Chem. 1988,27, 2565.  185.  Martell, A . E.; Smith, R. M . Critical Stability Constants; Plenum: New York, 1974-1982; Vols. 1-5.  186.  Ohman, L. -O.; Sjoberg, S. Polyhedron 1983,2, 1329.  187.  Larson, S. M . Seminars in Nuclear Medicine 1978, 8, 193.  188.  Lyster, D. M . ; Clevette, D. J.; Nelson, W. O.; Rihela, T.; Webb, G. A . ; Orvig, C. Manuscript in preparation.  189.  Langston, W . J.; Bryan, G. W . in Complexation of Trace Metals in Natural  190.  Williams, R. J. P. Proc. R. Soc. (B) 1981,213, 361.  191.  Gran, G. The Analyst 1952, 77, 661; Anal. Chim. Acta 1988,206, 111.  192.  Motekaitis, R. J.; Martell, A . E. Can. J. Chem. 1982,60, 168.  193.  Motekaitis, R. J.; Martell, A . E. Can. J. Chem. 1982,60, 2403.  Waters; Kramer, C. J. M . , Duinker, J. C , Eds.; Martinus Nijhoff/Dr W. Junk: The Hague, 1984; pp 375-392.  180  Appendix  Table A . l . Molecular weights (MW) of the 3-hydroxy-4-pyridinone ligands and their tris-ligand metal complexes. (The entries in italics are for the 3-hydroxy-4-pyrone ligands)  M W of Tris-Ligand Metal Complexes Ligand  MW  AIL3  GaL3  InL  Hmpp  125.1  399.4  442.1  487.2  Hdpp  139.1  441.2  484.0  529.1  Hmepp  153.1  483.2  526.0  571.1  Mimosine  198.2  618.6  661.3  706.4  Hmhpp  209.3  651.8  694.6  739.7  H2exn*  332.4  1048  1134  1224  Maltol  126.2  402.3  445.0  490.1  Kojic acid  142.1  450.2  493.0  538.1  •Molecular weights are for the M2L3 dimer.  3  181  Table A.2. Crystallographic data for the 3-hydroxy-4-pyridinone ligands.  compound  Hmpp  Hdpp  Hmepp  formula  C6H7NO2  C7H9NO2  C HnN0  formula weight  125.1  139.2  153.2  crystal system  monoclinic  orthorhombic  othorhombic  space group  Pl\ln  Pbca  Pbca  a (A)  6.8351 (4)  7.3036 (4)  12.5907 (8)  b(A) c(A)  10.2249 (4)  13.0490 (6)  11.7477 (6)  8.6525 (4)  13.7681 (7)  11.0040(6)  P(deg)  105.215 (4)  V(A3)  583.51 (6)  1312.2 (1)  1627.6 (3)  Z  4  8  8  D (g/cm3)  1.424  1.409  1.14  F(000)  264  592  592  diffractometer  Enraf-Nonius CAD4-F  Enraf-Nonius  Rigaku AFC6  u (Cu-K ) (cm-l)  8.65  8.21  6.62  radiation  Cu-Ka 1.540562  Cu-K  1.540562  1.544390  1.544390  N i filter  N i filter  temperature  23°C  23°C  23°C  Mmax (deg)  150  150  150.1  reflections with I > 3a (I)  914  857  1228  number of variables  111  128  145  R; Rw  0.037; 0.046  0.044; 0.046  0.053; 0.085  max A/a (final cycle)  0.05  0.05  0.02  goodness of fit indicator  2.349  1.014  3.50  0.20  0.32  0.45  c  a  WA)  residual density  (e/A ) 3  a  8  Cu-K  2  a  1.54178 graphite-monochromated  182 Table A.3. Crystallographic data for the M(dpp) complexes (recorded with a Enraf-Nonius 3  CAD4-F diffractometer). compound  Al(dpp) -12H 0  Ga(dpp)3-12H 0  In(dpp) -12H 0  formula  C21H48AIN3O18  C iH48GaN30i8  C2iH48lnN Oi  formula weight  657.6  700.3  745.44  crystal system  Trigonal  Trigonal  Trigonal  space group  P3  P3  P3  (A)  16.600 (2)  16.6549 (6)  16.842 (1)  c(A)  6.877 (1)  6.8691 (4)  6.8078 (7)  V(A3)  1641.3 (3)  1650.1 (1)  1672.3 (2)  Z  2  2  2  D (g/cm.3)  1.331  1.470  1.480  F(000)  704  740  776  radiation  Cu-K  Cu-K«  Mo-K  1.540562 1.54439  1.540562 1.54439  0.70930 0.71359  nickel filter  nickel filter  p (cm )  11.97  17.89  7.67  temperature  22°C  22°C  21°C  150  150  60  reflections with I > 3a (I)  1662  1653  2496  number of variables  202  195  190  R; Rw  0.045; 0.051  0.047; 0.055  0.033; 0.037  max A/a (final cycle)  0.17  0.023  0.027  goodness of fit indicator  1.020  1.023  1.492  0.23  0.48  -0.55 to +0.75 (near In)  a  3  c  W(A)  -1  2Q x ma  (deg)  residual density  (e/A ) 3  2  a  2  2  3  2  3  8  a  graphite monochromator  183  Table A.4. Crystallographic data for the M(mepp)3 complexes (recorded with a Rigaku AFC6).  compound  Al(mepp)3-12H20  Ga(mepp)3«12H20  formula  C24H54AIN3O18  C24H54GaN Oi8  formula weight  699.7  742.4  crystal system  Trigonal  Trigonal  space group  P3  P3  a (A)  17.1734 (8)  17.247 (1)  c(A)  6.827 (1)  6.830 (2)  V(A3)  1743.7 (3)  1759.4 (1)  z  2  2  D (g/cm )  1.33  1.40  F(000)  752  788  radiation  Cu-K  3  c  a  3  Mo-K  a  1.54178 graphite-monochromated  0.71069 graphite-monochromated  p (cm- )  11.56  8.50  temperature  21°C  21°C  2Qmax (deg)  150.3  55.0  reflections with I > 3a (I)  1157  1918  number of variables  207  215  R; Rw  0.032; 0.038  0.029; 0.036  max A/a (final cycle)  0.23  0.02  goodness of fit indicator  1.63  1.46  residual density (e/A )  0.10  0.28  tea  (A) 1  3  184  Table A.5. Hydrogen bond distances (A) and angles for In(dpp)3«12H20. Interaction  Atoms O-H  (A)  H O  (A)  o-o (A)  O - H - O (deg)  0(3)-H(03a)-0(l)  1.00  1.92  2.900(3)  165  0(3)-H(03b)-0(2)  0.87(3)  1.99(3)  2.839(3)  168(3)  0(4)-H(04a)-0(6)  0.70(4)  2.07(5)  2.744(4)  160(5)  0(4)-H(04b)-0(4)  0.87(9)  1.90(9)  2.755(3)  169(6)  0(4)-H(04c)-0(4)  0.72(8)  2.04(8)  2.755(3)  170(6)  »0(5)-H(05a)-O(4)  0.82  1.99  2.789(3)  166  0(5)-H(05b)-0(5)  0.77(5)  2.01(5)  2.782(3)  179(5)  0(6)-H(06a)~0(3)  0.66(4)  2.10(4)  2.759(4)  177(5)  0(6)-H(06b)-0(5)  0.74(4)  2.06(4)  2.780(4)  164(4)  Table A.6. Hydrogen bond distances (A) and angles for the M(mepp)3«12 H2O complexes. Interaction  O-H Al  H- 0 Ga  Al  0 - •O Ga  Al  O - H - 0 (deg) Ga  Al  Ga  0(3)-H(l)-0(l)  0.81(5) 0.75(4)  2.10(5) 2.14(4)  2.877(2) 2.883(3)  161(4) 173(4)  0(3)-H(2) 0(2)  0.92(7) 0.77(3)  1.95(7) 2.08(3)  2.838(2) 2.834(3)  163(5) 168(3)  0(4)-H(3)-0(6)  0.83(7) 0.85(4)  1.99(7) 1.95(7)  2.795(3) 2.784(3)  166(4) 167(4)  0(4)-H(4)-0(4)  0.85(8) 0.71(7)  1.97(7) 2.11(7)  2.811(2) 2.810(3)  170(6) 172(5)  0(5)-H(6)-0(4)  0.89(1) 0.81(3)  2.04(6) 2.03(4)  2.828(3) 2.821(3)  145.83 164(3)  0(5)-H(5)-0(5)  0.82(6) 0.76(3)  2.02(6) 2.08(4)  2.833(2) 2.835(3)  171(4) 175(4)  0(6)-H(7)-0(3)  0.91(6) 0.79(4)  1.86(6) 2.00(4)  2.771(3) 2.763(3)  173(4) 165(4)  0(6)-H(8)-0(5)  0.76(8) 0.77(4)  2.11(8) 2.12(4)  2.859(3) 2.849(4)  168(5) 159(4)  185  Appendix Procedure A . l  Potentiometric Equilibrium Measurements  Potentiometric measurements of the ligands in the absence, and presence, of metal ions were performed with an Orion Research E A 920 p H meter equipped with Orion Ross research grade glass and reference electrodes. A Metrohm automatic buret (Dosimat 665) was used to add the standard NaOH. The temperature was maintained at 25.0 ±0.1 °C throughout with waterjacketed beakers and a Julabo circulating bath, and the ionic strength was adjusted to 0.15 M (isotonic) by the addition of NaCl. A l l solutions were continuously degassed with prepurified A r during the course of a titration. The ligands were twice recrystallized or sublimed; concentrations were obtained by weighing. A l l metal-containing solutions were obtained from appropriate dilution of atomic absorption standard solutions of A l and Ga (Sigma or Aldrich). The exact amount of excess acid present in the metal ion solutions was determined by a Gran's p l o t  191  of ( V + V ) x 10 P _  G  t  H  vs.  V , where V = the initial volume of 1:1 metal-Na2EDTA solution, and V is the volume of added t  G  t  standard NaOH. The base consumed is equal to the excess acid plus the N a E D T A protons. The 2  metal-ligand titrations were performed at a total ligand (C) to total metal ion (B) concentration ratio of just greater than three. NaOH solutions (0.1 M) were prepared from dilutions of 50% NaOH (less than 0.1% N a C 0 3 ) with freshly boiled, distilled, deionized water and standardized 2  potentiometrically against potassium hydrogen phthalate (KHP). The electrodes were calibrated with standard aqueous HC1 and NaOH solutions to read -log[H ] directly. The range of -log[H ] available was limited from 1.5 to 12 in which the +  +  electrode behavior was reversible and linear. Protonation and deprotonation reactions of the ligands were studied within the range 2 < -log[H ] < 10 and the metal-ligand titrations were +  studied within the range 1.5 < -log[H ] < 4. A l l titrations were performed in sets of 3 or 4 runs. +  186  The proton dissociation constant of the ligands were determined by using the Fortran computer program P K A S .  1 9 2  In the M(IU) systems, the computations allowed for the presence of  M(OH)2+, M ( O H ) , M ( O H ) and M ( O H ) \ In addition, A l ( O H ) 2  +  3  4  2  2  4 +  and A l ( O H ) 3  4  5 +  were  included. Formation constants for these various metal species were taken from ref. 25. The stability constants for the main species M L Fortran computer program B E S T .  1 9 3  2 +  , M L , and ML3 were determined by using the 2  +  This program sets up simultaneous mass-balance equations  for all the components present at each addition of base, and calculates the pH at each data point according to the current set of stability constants and total concentrations of each component. Stability constants judiciously chosen by the user are automatically adjusted in order to minimize the sum of squares of differences between the calculated and observed values of - l o g [ H j . +  Adjustment is continued until there is no further improvement in the fit. The constants are reported to the second decimal place, which is representative of the reproducibility of the potentiometric equipment employed.  


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