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METAL ION COMPLEXING PROPERTIES OF THE HIGHLY PREORGANIZED LIGANDS 8-QUINOLYL-1,10-PHENANTHROLINE AND 1,10-PHENANTHROLINE-2,9-DICARBOXYALDEHYDE Adam Lawson Brenneman A Thesis Submitted to the University of North Carolina Wilmington in Partial Fulfillment of the Requirements for the Degree of Master of Science Department of Chemistry and Biochemistry University of North Carolina Wilmington 2010 Approved by Advisory Committee Dr. Sridhar Varadarajan . Dr. S. Bart Jones Dr. Robert D. Hancock Chair . Accepted by DN: cn=Robert D. Roer, o=UNCW, ou=Dean of the Graduate School & Research, [email protected], c=US Date: 2010.07.08 09:52:52 -04'00' _________________________________ Dean, Graduate School . TABLE OF CONTENTS ABSTRACT ....................................................................................................................... iii ACKNOWLEDGMENTS ................................................................................................. iv LIST OF TABLES ...............................................................................................................v LIST OF FIGURES ......................................................................................................... viii INTRODUCTION ...............................................................................................................1 METHODS ........................................................................................................................15 Synthesis of PDALD.....................................................................................................17 UV-Vis spectrophotometric titrations involving 8QP ..................................................19 Fluorescence studies involving 8QP .............................................................................25 UV-Vis spectrophotometric titrations involving PDALD ............................................26 RESULTS AND DISCUSSION ........................................................................................29 Synthesis of PDALD.....................................................................................................29 8QP Protonation Constants ...........................................................................................32 PDALD Protonation Constants .....................................................................................38 HyperChem MM Calculations ......................................................................................45 Titrations Involving 8QP ..............................................................................................52 Fluorescence studies involving 8QP .............................................................................75 Titrations Involving PDALD ........................................................................................77 Crystal Structure Results.............................................................................................108 CONCLUSIONS..............................................................................................................113 LITERATURE CITED ....................................................................................................118 ii ABSTRACT Highly preorganized ligands have shown greater stability constants as well as increased metal-ion selectivities over their straight-chain analogs. The preorganized ligand 8-quinolyl1,10-phenanthroline (8QP) was studied to determine its formation constants with various aqueous metal-ions as well as its metal-ligand complex crystal structure. Formation constants were determined from UV/Vis spectrophotometry detection methods using the absorbance spectra as a function of pH. Formation constants for the metal ions Cu2+, Ni2+, Zn2+, Ag2+, and Cd2+ are reported and the crystal structure for the [Cu(8QP)] complex is also reported. Fluorescence properties of the free ligand 8QP and three of its most stable metal complexes ([Cu(8QP)], [Zn(8QP)], and [Ni(8QP)]) were also examined. 1,10-phenanthroline-2,9-dicarboxyaldehyde (PDALD) was synthesized by a literature method and was subjected to purity verification for studies into its formation constants with various aqueous metal-ions. Formation constants were determined from UV/Vis spectrophotometry detection methods using the absorbance spectra as a function of pH. Formation constants for the metal ions Cd2+, Cu2+, Pb2+, Zn2+, and Th4+ are reported amongst others. iii ACKNOWLEDGEMENTS I would like to sincerely thank Dr. Hancock for all of his help and support over the past few years. His patience dealing with my work schedule and more than a couple, spur of the moment, extended overseas disappearances is something that I have been extremely grateful for. Not many people would have been as flexible or allowed me the freedom that he has, and it has really made my time in this program unique and enjoyable. I would also like to thank my other committee members, Dr. Jones and Dr. Varadarajan, for their help and guidance throughout my career. The faculty and staff of the Biology and Chemistry Departments, as well as the Athletic Department, have really made my time here at UNCW fantastic. Thanks to all those who have supported me in my efforts to continue my education; Dr. Ballard, Dr. Dillaman, Coach Dave Allen, and Dr. Beard. Special thanks to my parents, Lynn and Diane, my brother Jeremy, my girlfriend Abbey, and the Murray family for their support and encouragement. iv LIST OF TABLES Table Page 2+ 1. Stability constants of Ni complexes in a series of polyamine ligands to show the effect of increasing ligand denticity .......................................11 2. EXCEL spreadsheet for free ligand used to calculate protonation equilibria. .......36 3. Solutions for absorbance values produced by „SOLVER‟ module in determining pKa of 8QP ........................................................................................37 4. Protonation equilibria and constants for 8QP free ligand ......................................37 5. Solutions for absorbance values produced by „SOLVER‟ module in determining pKa of PDALD ..................................................................................41 6. Protonation equilibria and constants for PDALD free ligand ................................41 7. Protonation constants and formation constants with a selection of metal ions with 8QP in 1.0 M NaClO4 at 25°C. ....................................................49 8. Protonation constants and formation constants with a selection of metal ions with PDALD in 1.0 M NaClO4 at 25°C ...............................................51 9. Solutions for the pKa and absorbance parameters in determining logK1 of 8QP-Cu complex .....................................................................................57 10. Summary of equations at each protonation equilibrium of Cu(8QP). ...................57 11. Solutions for the pKa and absorbance parameters in determining logK1 of 8QP-Ni complex ......................................................................................62 12. Summary of equations at each protonation equilibrium of Ni(8QP) .....................62 13. Solutions for the pKa and absorbance parameters in determining logK1 of 8QP-Zn complex .....................................................................................67 14. Summary of equations at each protonation equilibrium of Zn(8QP) ....................67 v 15. Solutions for the pKa and absorbance parameters in determining logK1 of 8QP-Cd complex .....................................................................................72 16. Summary of equations at each protonation equilibrium of Cd(8QP) ....................72 17. Solutions for the pKa and absorbance parameters in determining logK1 of PDALD-Cu complex ...............................................................................80 18. Summary of equations at each protonation equilibrium of Cu(PDALD) ..............80 19. Solutions for the pKa and absorbance parameters in determining logK1 of PDALD-Cd complex ...............................................................................83 20. Summary of equations at each protonation equilibrium of Cd(PDALD) ..............83 21. Solutions for the pKa and absorbance parameters in determining logK1 of PDALD-Ca complex ...............................................................................87 22. Summary of equations at each protonation equilibrium of Ca(PDALD) ..............87 23. Solutions for the pKa and absorbance parameters in determining logK1 of PDALD-Gd complex...............................................................................91 24. Summary of equations at each protonation equilibrium of Gd(PDALD) ..............91 25. Solutions for the pKa and absorbance parameters in determining logK1 of PDALD-Pb complex ...............................................................................95 26. Summary of equations at each protonation equilibrium of Pb(PDALD) ..............95 27. Solutions for the pKa and absorbance parameters in determining logK1 of PDALD-Th complex ...............................................................................99 28. Summary of equations at each protonation equilibrium of Th(PDALD) ..............99 29. Solutions for the pKa and absorbance parameters in determining logK1 of PDALD-Zn complex .............................................................................103 30. Summary of equations at each protonation equilibrium of Zn(PDALD) ............103 vi 31. Solutions for the pKa and absorbance parameters in determining logK1 of PDALD-UO2 complex ..........................................................................107 32. Summary of equations at each protonation equilibrium of UO2(PDALD)..........107 33. Crystal data and structure refinement for [Cu(8QP)2](ClO4)2 .............................111 34. Bond lengths and angles of interest in the complex cation [Cu(8QP)2]2+ ...........112 35. Comparison of protonation constants and logK1 values with a selection of metal ions between 8QP and terpy ..................................................................114 36. Comparison of logK1 and ΔlogK1 values with a selection of metal ions with PDALD, phen, and PDALC. .......................................................................116 vii LIST OF FIGURES Figure Page 1. Structure of the [Co(NH3)6]Cl3 complex indicating the octahedral coordination geometry.. ...........................................................................................1 2. Structure of [Gd(DTPA)OH2] complex used as a MRI contrast agent ....................2 3. Chemical equation showing the Fe2+ + H2O2 making ROS .....................................5 4. CHEF Effect diagram ..............................................................................................7 5. Ethylenediamine v. 1,10-phenanthroline .................................................................8 6. Effect of metal ion size on relative affinity for the ligand DPyA, which forms a 6-membered chelate ring, compared to bipy, which forms a 5-membered chelate ring. ...........................................................................9 7. Illustration of the effect of the formation of chelate rings on metal ion selectivity .........................................................................................................10 8. Metal ion classification chart according to Pearson‟s HASB theory. ....................12 9. Plot of logK1 DIEN versus logK1 NH3 for various metal ions to illustrate which metal ions form strong bonds with nitrogen donors.....................13 10. Ligands discussed in this paper..............................................................................14 11. Schematic of flow cell apparatus used in all UV-vis titration experiments ...........15 12. Schematic of PDALD synthesis.............................................................................17 13. FT-IR spectra of 1,10-phenanthroline-2,9-dicarboxyaldehyde. ............................29 14. 1 15. Overview of the absorbance v. wavelength spectra for the 8QP free ligand at 4 different pH‟s to show the progression of the spectral curve ..............34 16. Absorbance v. wavelength spectra for the 8QP free ligand for pH ≈ 0.5 – 8.5.. ......................................................................................................35 17. Plot of measured and theoretical absorbance vs. pH for 8QP at 223 nm...............35 H-NMR spectrum of a) 1,10-phenanthroline-2,9-dicarboxyaldehyde (PDALD) in DMSO-d6 and b) neocuprine in DMSO-d6 ......................................30 viii 18. Variation of absorption at the wavelengths indicated, as a function of pH for 4.0 x 10-6 M 8QP in 1.0 M NaClO4 at 25 °C. The points are experimental values of absorbance, and the solid lines are theoretical curves of absorbance vs pH calculated on the basis of two protonation constants, of 6.82 and 2.40 for 8QP .......................................................................38 19. Overview of the absorbance v. wavelength spectra for the PDALD free ligand at 4 different pH‟s to show the progression of the spectral curve ..............39 20. Absorbance v. wavelength spectra for the PDALD free ligand from pH ≈ 2.0 – 7.5 ........................................................................................................40 21. Plot of measured and theoretical absorbance vs. pH for PDALD at 270 nm ........40 22. Variation of absorption at the wavelengths indicated, as a function of pH for 2.0 x 10-5 M PDALD in 0.1 M NaClO4 at 25 °C. The points are experimental values of absorbance, and the solid lines are theoretical curves of absorbance vs pH calculated on the basis of two protonation constants, of 7.05 and 2.12 for PDALD.................................................................42 23. Calculated strain energy (kcal/mol) versus metal-nitrogen bond length (Å) curve for 8QP. The arrows indicate points on the curve where specific metal ions fit the curve ..........................................................................................46 24. [Cd(8QP)(H2O)2] MM ...........................................................................................47 25. [Ca(8QP)(H2O)3] MM ...........................................................................................48 26. [Zn(8QP)(H2O)2] PM3 ...........................................................................................48 27. [Pb(PDALD)(OH)2] complex ................................................................................50 28. UV spectra of 3.9 x 10-6 M 8QP and 4 x 10-5 M Cu(ClO4)2 in 0.01 M HClO4 and 0.09 M NaClO4. Initial spectrum of 8QP-Cu complex at pH = 2.15. Final spectrum of 8QP- Cu hydroxide complex at pH = 7.18. ............54 29. Plot of corrected absorbance (data points) and theoretical absorbance versus pH for titration of 8QP and Cu(ClO4)2 for varying wavelengths ..............55 30. UV spectra of 4.0 x 10-6 M 8QP and 4.0 x 10-6 M Cu(ClO4)2 in a 1.0 M (0.1 M HClO4 and 0.9 M NaClO4) solution ...........................................................56 31. Plot of corrected absorbance (data points) and theoretical absorbance versus pH for titration of 8QP and Ni(ClO4)2 for varying wavelengths ...............60 32. UV spectra of 4.0 x 10-6 M 8QP and 4.0 x 10-3 M Ni(ClO4)2 in a 1.0 M (0.1 M HClO4 and 0.9 M NaClO4) solution ...........................................................61 ix 33. Plot of corrected absorbance (data points) and theoretical absorbance versus pH for titration of 8QP and Zn(ClO4)2 for varying wavelengths...............65 34. UV spectra of 4.0 x 10-6 M 8QP and 4.0 x 10-3 M Zn(ClO4)2 in a 1.0 M (0.1 M HClO4 and 0.9 M NaClO4) solution ...........................................................66 35. Plot of corrected absorbance (data points) and theoretical absorbance versus pH for titration of 8QP and Cd(ClO4)2 for varying wavelengths ..............70 36. UV spectra of 4.0 x 10-6 M 8QP and 4.0 x 10-3 M Cd(ClO4)2 in a 1.0 M (0.1 M HClO4 and 0.9 M NaClO4) solution ...........................................................71 37. Fluorescence - Plot of emission v. wavelength for the 8QP free ligand, Zn(8QP) complex, Ni(8QP) complex, and Cu(8QP) complex ..............................76 38. Plot of corrected absorbance (data points) and theoretical absorbance versus pH for titration of PDALD and Cu(ClO4)2 for varying wavelengths ........78 39. Major differences between the spectra of the free ligand and the spectra where the Cu2+ complex is present ........................................................................79 40. Plot of corrected absorbance and theoretical absorbance versus pH for titration of PDALD and Cd(ClO4)2 for varying wavelengths .........................81 41. UV spectra of 2.0 x 10-5 M PDALD and 2.0 x 10-5 M Cd(ClO4)2 in a 0.1 M (0.01 M HClO4 and 0.09 M NaClO4) solution .......................................................82 42. Plot of corrected absorbance and theoretical absorbance versus pH for titration of PDALD and Ca(ClO4)2 for varying wavelengths .........................85 43. UV spectra of 2.0 x 10-5 M PDALD and 2.0 x 10-5 M Ca(ClO4)2 in a 0.1 M (0.01 M HClO4 and 0.09 M NaClO4) solution .......................................................86 44. Plot of corrected absorbance and theoretical absorbance versus pH for titration of PDALD and Gd(ClO4)3 for varying wavelengths .........................89 45. UV spectra of 2.0 x 10-5 M PDALD and 2.0 x 10-5 M Gd(ClO4)3 in a 0.1 M (0.01 M HClO4 and 0.09 M NaClO4) solution ......................................90 46. Plot of corrected absorbance and theoretical absorbance versus pH for titration of PDALD and Pb(ClO4)2 for varying wavelengths ..........................93 47. UV spectra of 2.0 x 10-5 M PDALD and 2.0 x 10-5 M Pb(ClO4)2 in a 0.1 M (0.01 M HClO4 and 0.09 M NaClO4) solution ......................................94 x 48. Plot of corrected absorbance and theoretical absorbance versus pH for titration of PDALD and Th(NO3)4 for varying wavelengths ..........................97 49. UV spectra of 2.0 x 10-5 M PDALD and 2.0 x 10-5 M Th(NO3)4 in a 0.1 M (0.01 M HClO4 and 0.09 M NaClO4) solution ......................................98 50. Plot of corrected absorbance and theoretical absorbance versus pH for titration of PDALD and Zn(ClO4)2 for varying wavelengths .......................101 51. UV spectra of 2.0 x 10-5 M PDALD and 2.0 x 10-5 M Zn(ClO4)2 in a 0.1 M (0.01 M HClO4 and 0.09 M NaClO4) solution ....................................102 52. Plot of corrected absorbance and theoretical absorbance versus pH for titration of PDALD and UO2(NO3)2 for varying wavelengths ......................105 53. UV spectra of 2.0 x 10-5 M PDALD and 2.0 x 10-5 M UO2(NO3)2 in a 0.1 M (0.01 M HClO4 and 0.09 M NaClO4) solution ....................................106 54. Structure of the complex cation [Cu(8QP)2]2+, showing the numbering scheme for atoms relevant to discussion of the coordination sphere around the copper. ................................................................................................109 55. Structure of [Pt(8QP)Cl]+ illustrating the non-planarity due to angle strain characteristic of 8QP complexes.. ..............................................................110 56. Comparison of terpy and 8QP binding site accessibility with small metal ions .............................................................................................................114 xi INTRODUCTION Understanding of coordination chemistry dates back to Alfred Werner‟s research with transition metal-amine complexes in the late 19th and early 20th centuries. Werner defined coordination numbers in 1893, along with his research determining the octahedral configuration of [Co(NH3)6Cl3], that led him to propose that neutral or anionic ligand molecules coordinate in geometrical arrangements (coordination geometry) around a central transition metal atom. He found that the geometry of a complex differed according to the coordination number, or number of atoms which would coordinate to the metal ion.1 Figure 1 : Structure of the [Co(NH3)6]Cl3 complex as postulated by Alfred Werner, indicating the octahedral coordination geometry. Subsequent research built off the foundation laid by Werner and his contemporaries, and led to the development of ligands with very specific characteristics; among them: metal ion selectivity, relative strength or weakness of the metal-ligand complex, and specialty functions such as fluorescence upon complexation. The concept of preorganization2 has allowed chemists to design and synthesize ligands which form very strong complexes with the intended target and has allowed metal-ligand complexes to find application in medical, industrial, and technological fields, previously untouched by inorganic chemistry. One of the most commonly used applications is the medical imaging technique known as magnetic resonance imaging (MRI). MRI uses contrast agents to provide detailed images of the body which offer differentiation between soft tissues, making it especially useful in neurological, musculoskeletal, and oncological inquiries.3 Image contrast is derived from the polarization of coordinated water molecules onto Gd3+ ions. The gadolinium complex with diethylenetriaminepentaacetic acid (DTPA) was the first paramagnetic chelate used clinically as an MRI contrast agent and is still the most common.4 A good contrast agent such as [Gd(DTPA)H2O]2- must fulfill three basic requirements. The ligand complex must be highly stable to mask the toxic effects of injection of gram amounts of gadolinium. High water solubility is also important due to the need for a relatively concentrated, small-volume, injectable solution. The third requirement is that it must have the ability to enhance the relaxation rate between protons of water molecules located in the inner sphere of the Gd3+ ion and surrounding water molecules.3 The amount of image contrast that can be attained is dependant upon the rate of relaxation, as well as the amount of water molecules which may be found within the inner shell. Figure 2 : Structure of [Gd(DTPA)OH2] complex used as a MRI contrast agent 2 As seen in Figure 2, Gd3+ has a coordination number of nine5, with eight of its coordination sites occupied by the octadentate DTPA ligand. This leaves only one site for a water molecule to coordinate and limits the amount of contrast resolution which can be attained. Future contrast agents may remedy this situation by preorganization of hemicyclic ligands with fewer bonds to the metal ion, leaving more sites available for water coordination 6 and potentially increasing contrast. Recent research with gadofullerene complexes has also shown great promise in the development of safer, more soluble, and more effective contrast agents.7 A variety of metal ions can be found in the human body, such as copper, zinc, and calcium are essential at low concentrations and play critical roles in maintaining life. Toxic metals such as cadmium, mercury, and lead may have relatively high concentrations in the body although they serve no known biological function.8 Humans are exposed to metals on a daily basis through their diets and from the environment.9 Whether deemed essential or not, when metal ion concentrations exceed normal human homeostatic levels, there will be adverse effects. These effects may not be very noticeable as metal ions accumulate throughout the body; but over time, they begin to disrupt enzyme pathways, interfere with protein synthesis, cause cancer and eventually death. Chelation therapy is the administration of a chelating agent in order to remove heavy metals from the body. Some practitioners of alternative medicine claim that it is very effective in removing metastatic calcium deposits in patients with arteriosclerosis10; however, the term chelation therapy is generally applied to treatment in cases of acute heavy metal exposure or chronic poisoning. Of all heavy metals, lead poisoning is the most commonly reported and is especially common among children six years and younger because of their susceptibility to the lead content 3 in paints, soil in many areas, and in toys. Lead exposure in children leads to stunted growth, learning disabilities, decreased motor skills, and lowered performance on intelligence tests.11 The stunted growth may be due to a detoxification mechanism in humans which deposits up to 90% of ingested lead into bone matrix, inhibiting proper bone growth.12 The toxic effects are brought on by the remaining 10% which bind to sulfhydral groups of cysteine residues and deform and inactivate proteins. This effect on protein function is widely considered to be the most likely mode of toxicity for lead as well as other heavy metals including mercury and cadmium.8 Ethylenediamine tetraacetic acid (EDTA) was introduced in the 1940s in the treatment of patients with lead poisoning and continues to be used today only in patients with extremely high blood lead levels (BLL).10 This is due to the fact that it removes iron as well as the intended heavy metals from the body and requires either musculoskeletal injection or intravenous administration. Depending on the BLL, patients may require multiple chelation treatments to completely clear the metal from the blood as well as multiple intravenous iron replacement treatments. Dimercaptosuccinic acid (DMSA) is another chelation therapy medication which is commonly used in less severe cases. It is useful in removing both lead and mercury, and is often the preferred course of action because it may be administered orally.13 However they are administered, chelation therapy medications bind strongly with the unwanted metal ions in the blood, prevent them from causing any further damage to the body, and allow the metal to be harmlessly excreted through the urine. The most common form of dementia in older people is Alzheimer‟s disease. It is an incurable, degenerative, and terminal disease which affects as many as 4.5 million people in the United States.14 The progression of the disease is marked by the aggregation of β-amyloid (Aβ) 4 peptides that trigger neurodegeneration. Iron, copper, and zinc concentrations in these Aβ deposits are substantial and are thought to play major roles in the development of Alzheimer‟s.15 Zn2+ has been shown to trigger the Aβ aggregation and plaque formation. The neurotoxicity of the β-amyloid deposits is thought to be due to the generation of damaging reactive oxygen species. This is commonly attributed to reactions similar to the Fenton reaction seen in Fig.3, between redox active transition metals and hydrogen peroxide.16 Figure 3 : An example of a common reaction which occurs within Aβ plaques, this reaction shows the oxidation of Fe2+ to Fe3+ and resulting hydroxyl radicals. Cherney et al. have shown evidence that Cu2+/Zn2+ chelation is effective in inhibiting β-amyloid accumulation in vivo.15 Future research is necessary, but the development of a ligand which could bind selectively with Cu2+ and Zn2+ to prevent and decrease Aβ aggregate deposition may be a viable method for the prevention and treatment of Alzheimer‟s disease. Early diagnosis is essential to impeding the progression of Alzheimer‟s disease. Physicians rely on symptoms such as cognitive impairment and dementia to show before they are able to make a probable diagnosis and begin treatment, which in many cases is already too late. Currently, the only way to get a definitive diagnosis of Alzheimer‟s disease is from a postmortem histological analysis. Development of fluorescent markers capable of safely crossing the blood brain barrier and quantifying zinc and copper levels in vivo would allow physicians to detect the presence of conditions conducive of Aβ aggregation before the onset of symptoms. 5 A molecular sensor is a molecule which has the ability to signal the presence of a specific substrate. Seen in Figure 4, 9,10-bis(TMEDA)anthracene is a fluorescent marker which can signal the presence of zinc in a biological sample, and utilizes common principles of molecular sensor design. Upon complexation with the target, the sensor undergoes a change which can be observed and quantified by investigators, most often changes in fluorescence or absorbance. A significant increase in the magnitude of fluorescent emission upon chelation of a metal ion is a phenomenon known as chelation-enhance fluorescence (CHEF).17 This effect is often exploited in development of new chemosensors by using extended aromatic systems with amine donor groups attached. While the free ligand is unbound in solution, the inherent fluorescence of the extended aromatic system will be quenched due to the proximity of the lone pairs of amine groups. As seen in Figure 4, once the metal binds, the amine lone pair is involved in the complexation of the metal and is unable to donate an electron to the excited aromatic, allowing the system to fluoresce and signal the chelation equilibrium.17 Whether designing ligands for medical application as previously discussed, or for application in industrial, manufacturing, or environmental fields, certain concepts and characteristics must be considered. Preorganization, chelate ring size18, and the number and type of donor atoms19 all play roles in the behavior and specificity of the ligand. The term preorganization describes the design of a ligand structure with a conformation conducive to chelation of the target metal ion. This is often achieved by the slight modification of ligands with desired binding properties by adding structural elements which sterically hinder any deviation from the optimal conformation.2,20 6 Molecular recognition is prevalent in biological systems and occurs naturally in DNAprotein, RNA-ribosome, and receptor-ligand interactions.21 Figure 4 : CHEF Effect -- Structures of 9,10-bis(TMEDA)anthracene (1) and its fluorescent bis(ZnCl2) complex (2). The bottom diagram shows the relative fluorescent intensity of 1 vs. 2 and 9,10-dimethylanthracene (DMA) in acetonitrile.17 (All solutions at 10-4 M) 7 Figure 5 : This diagram shows how ethylenediamine can freely rotate along each single bond. This means it has to overcome more energy to bind a metal ion, lowering the logK1. With the addition of an ethylene bridge, forming 1,10-phenanthraline, the nitrogens in the 1,10-position are fixed and unable to rotate. Studies have shown that making it unnecessary for the ligand to rotate to form a complex leads to an increase in logK1. 18 This concept was first applied to ligand design after 1967 when Charles Pedersen published evidence that crown ethers, two-dimensional organic compounds, were able to recognize and selectively bind with alkali metal ions to form highly structured complexes.22 This was followed by research findings published in 1969 on the design, synthesis, and binding properties of cryptands by the French scientist J. M. Lehn.23 Then, in 1974, Donald J. and Jane M. Cram published findings and introduced a whole new field of study called “Host-Guest Chemistry”. Cram went on to expand on Pedersen‟s work by synthesizing three-dimensional molecules that could mimic the way natural molecules functioned. In 1987 when Pedersen, Lehn, and Cram were jointly awarded the Nobel Prize for Chemistry, Cram emphasized two major points: preorganization is the central determinant of binding power and complementarity is the central determinant of structural recognition. Cram used preorganization techniques to make heterocyclic molecules which coordinate specific metals with incredible efficacy due to their structural rigidity and shape.20 Hemicyclic 8 ligands such as 8QP and PDALD are of particular interest due to their ability to form complexes of similar strength as macrocycles, while being considerably easier and cheaper to synthesize. In addition to pre-organization and complementarity, chelate ring size and the number and type of donor atoms all play roles in behavior and specificity when designing ligands. Chelate ring size is important to consider when designing a ligand for a specific function. Figure 6 shows that there is a direct correlation between the size of the ring that metal chelation forms and the atomic radius of coordinated metal ions with the highest logK1 values.24 The chelate ring size rule states that five-membered chelate rings, such as the one present in 1,10-phenanthroline from Figure 7, coordinate with the least steric strain to large metal ions, while a 6-membered chelate ring such as that of DPN (dipyridonaphthalene) coordinates with the least strain to a very small metal ion.25 Figure 6 : Effect of metal ion size on relative affinity for the ligand DPyA, which forms a 6membered chelate ring, compared to bipy, which forms a 5-membered chelate ring. Formation constant data from NIST.24 9 Figure 7 : Illustration of the effect of the formation of chelate rings on metal ion selectivity. (Left side) Ligands that form of 5-membered chelate rings upon complexing a metal ion are selective for large metals (preferred ionic radius ~ 1.0 Å and M-L bond length 2.5 Å) , where 6membered rings are selective for small metal ions (preferred ionic radius ~ 0.3 Å and M-L bond length 1.6 Å) (right side). This effect can be attributed to difference in the M-L bond length for the 5- and 6-membered chelate rings and the corresponding steric hinderence which larger metal ions have to overcome to form a bond in a 6-member ring configuration. The importance of the number of donor atoms, or denticity, was discussed earlier as it applied to the Gd3+-DTPA complex. DTPA was said to be octodentate meaning that the ligand had eight points of attachment to the gadolinium ion. Since the coordination geometry of different metal ions is usually known, the necessary ligand shape and denticity are considered when designing a target specific molecule. Whenever possible, it is important to use multidentate ligands as opposed to multiple mono- or bidentate ligands, due to the entropic advantages of the former explained by the chelate effect.26 This can be seen in Table 1 by comparing the values for adding two monodentate ligands (NH3) and one bidentate ligand like ethylenediamine (en) with Ni2+.27 10 Table 1. Stability constants of Ni2+ complexes in a series of polyamine ligands to show the effect of increasing ligand denticity. Polyamine denticity, n EN 2 DIEN 3 TRIEN 4 TETREN 5 PENTEN 6 log βn (NH3) 5.08 6.85 8.12 8.93 9.08 logK1 (polyamine) 7.47 10.7 14.4 17.4 19.1 Ionic Strength = 0.5 M EN DIEN TRIEN TETREN PENTEN NH3CH2CH2NH2 NH3(CH2CH2NH2)2H NH3(CH2CH2NH2)3H NH3(CH2CH2NH2)4H NH3(CH2CH2NH2)5H log βn (NH3) = log(K1 x K2 ---- x Kn) When looking for a ligand which simply binds the metal ion very tightly, a ligand may be designed to coordinate at all possible sites on the metal ion. In the case of MRI contrast agents, it is important to leave sites available for water molecules to coordinate, so the denticity of the ligand should be less than the coordination number of the corresponding metal. 11 In addition to the number of donor atoms, the type of donor atoms can be altered to take advantage of the properties of the target metal. Pearson‟s HASB Principle states that hard acids form more thermodynamically stable complexes with hard bases, and soft acids complex better with soft bases. Donor atom and metal ion hardness trends according to Pearson can be seen in Figure 8 below.28 Figure 8 : Metal ion classification chart according to Pearson‟s HASB theory. Nitrogen donors are often used in coordination chemistry because it is a synthetically convenient point of attachment, and displays stronger coordination properties with many metal ions than neutral oxygen donors.18 The affinity of individual metal ions for neutral nitrogen donors can be characterized by the linear free energy relationship shown in Figure 9. logK1 values for ammonia are used as indicators of the affinity for nitrogen donors. 12 Figure 9 : Plot of logK1 DIEN versus logK1 NH3 for various metal ions to illustrate which metal ions form strong bonds with nitrogen donors. PDALD and 8QP have two and three unsaturated nitrogens, respectively. Unsaturated nitrogen donors have low basicity, which is important because strong complexes not only rely on formation constants but also the relative difficulty to remove the protons from donor groups in order to permit complex formation.18 Neutral and negative oxygen donors can also be added to a ligand to increase affinity for large metal ions and metal ions classified as hard acids, respectively. All these factors were considered in the selection of the ligands investigated in this study: 8-quinolyl-1,10-phenanthroline (8QP) and 1,10-phenanthroline-2,9-dicarboxyaldehyde (PDALD). 13 Figure 10 : Structure, nomenclature, and abbreviation of ligands discussed in this paper. With the chelate ring size rule in mind, investigation of the ligand 8QP is of special interest to determine whether the inclusion of both a five-membered and six-membered ring will make it target mid-sized metal ions such as Zn2+ and Cu2+. Comparison of 8QP with terpyradine (terpy), a similar three nitrogen donor ligand with two five-membered rings, should provide evidence that the inclusion of a six-membered ring increases affinity for smaller metal ions. Investigation of PDALD will seek to compare its binding properties with the similar ligands 2,9-bis(hydroxymethyl)-1,10-phenanthroline (PDALC) and 1,10-phenanthroline (phen), seen in Figure 10. Comparisons will be made with phen based on the addition of neutral oxygen donors (PDALD) versus negative oxygen donors (PDALC). The purpose of this study is to thoroughly characterize the stability constants, logK1, of metal complexes formed with the ligands 8QP and PDALD as well as investigate the fluorescent properties. 14 METHODS Equipment Specifications UV/Vis absorbance spectra were recorded for aqueous metal-ligand titration experiments using a double beam Cary 1E UV/Vis spectrophotometer (Varian, Inc.) with WinUV Version 2.00(25) software. A 1.0 cm quartz flow cell, fitted with a variable flow peristaltic pump, was used to circulate the metal-ligand solution continuously throughout the series of titration experiments. Titrant mixing was enhanced using a magnetic stir bar and stir plate under the titration vessel. Sample temperature remained constant at 25.0±0.1°C, stabilized using a temperature controlled flow cell. All pH values for titration experiments were recorded using a SympHony SR60IC pH meter (VWR Scientific, Inc.). Calibration occurred before each titration and consisted of either titrating a 25 mL solution of 0.010 M HClO4 in 0.090 M NaClO4 with 50 mL of 0.010 M NaOH in 0.090 M NaClO4 and calculating E0 to determine correlation between mV readings and calculated pH, or calibration using pH 4.00, 7.00, and 10.00 buffer solutions prior to each titration. Figure 11 : Schematic of flow cell apparatus used in all UV-vis titration experiments. 15 All aqueous metal and free ligand stock solutions were prepared using deionized (DI) water. Aqueous metal-ligand solutions were prepared at ionic strengths of either 0.10 M (0.010 M HClO4 / 0.090 M NaClO4) or 1.0 M (1.0 M HClO4), with the distinction noted in each case and the titrant used being 0.10 M or 1.0 M NaOH, respectively. The titrant solution was allowed to equilibrate for 7 minutes between each addition. Absorbance scan ranges were taken from 200 to 350 nm at a rate of 600.00 nm/min. Absorbance spectra were referenced using DI H2O and a 1.0 cm quartz cell filled with DI H2O was placed in the path of the reference beam. A Bruker 400 MHz NMR spectrometer was used to obtain 1H-NMR spectra for analysis of PDALD organic synthesis. 1 H-NMR samples were prepared in DMSO-d6 and spectra were referenced to the DMSO peak at 2.49 ppm. A Thermo Scientific Nicolet 6700 FT-IR instrument (Thermo Nicolet Corp.) with WinFirst software was used to obtain infrared absorption spectra. The samples for FT-IR analysis were prepared as KBr Pellets with a 7 mm die press (Pike Technologies). Fluorescence spectra were obtained using a HORIBA Jobin Yvon Fluorolog-3 scanning spectrofluorometer equipped with a 450 W Xe short-arc lamp and a R928P emission detector (high sensitivity 240-850 nm). Excitation wavelengths were scanned from 250-500 nm at 5 nm increments. Emission wavelengths were scanned from 365-800 nm at 5 nm increments. The spectra obtained were 3D excitation and emission spectra and were reported in S1/R1 mode, processed by the HORIBA Jobin Yvon software package, FluorEssence (v 2.1). 16 Acquisition of Materials 8QP was synthesized29 by a research group supervised by Dr. Randolph Thummel at the University of Houston. All other chemicals and reagents used in this study were of analytical grade and were purchased commercially. Synthesis of PDALD was carried out for the purposes of this study using the oxidation method shown and described below.30 Figure 12 : Schematic of PDALD synthesis Synthesis of PDALD In aromatic heterocycles, methyl groups located in the α-position to a nitrogen atom oxidize relatively easily into their corresponding aldehyde with the addition of selenium dioxide, SeO2.30 With this in mind, 3.00 g of neocuprine (2,9-dimethyl-1,10-phenanthroline monohydrate) and 7.50 g of SeO2 were combined in the bottom of a 250 mL round bottom flask. 200 mL of a 5% DI H20 / 95% p-dioxane solution were added to dissolve the compounds. A magnetic stir bar was placed into the flask before the mixture was immersed in a paraffin bath and heated to 100 °C. The solution was stirred and refluxed for three hours. The hot solution was quickly filtered through a layer of Celite, which had been compacted on the filter paper by 17 seating with p-dioxane, by vacuum filtration. The solution was allowed to cool slowly to room temperature before being refrigerated overnight. The following morning, the solution was warmed to room temperature before the yellow precipitate was isolated by vacuum filtration. This was placed in a 500 mL round bottom flask with 300 mL of tetrahydrofuran, recrystallized, and refiltered. The product was characterized by 1H-NMR spectroscopy and FT-IR analysis and found to be pure PDALD. This process resulted in 0.68 g of product for a yield of 22.9 %. Solutions for and methods of pH electrode calibration Two stock solutions were prepared for mV corrected pH electrode calibration. The first solution (0.010 M HClO4 / 0.090 NaClO4) was prepared in a 25 mL volumetric flask using 21.5 µL HClO4 (11.6 M) and 0.275 g NaClO4, filled to the line with DI H2O. The second solution (0.010 M NaOH / 0.090 NaClO4) was prepared in a 50 mL volumetric flask using 50 µL NaOH (10.0 M) and 0.551 g NaClO4, filled to the line with DI H2O. The 25 mL 0.010 M HClO4 / 0.090 NaClO4 solution was placed into the temperature controlled UV-vis sample container with a magnetic stir bar. The 50 mL of the 0.010 M NaOH / 0.090 NaClO4 solution was added in 1.00 mL increments, recording the pH and mV reading after each titration. Using the Microsoft Excel and a variation of the Nernst equation, one can create a calibration curve which gives corrected pH values based on slight changes in the electric potential, measured in mV, of a system. This is done using the graphing function of Excel to plot the calculated pH vs. mV over the course of the titration and extracting a linear equation for the line of best fit. The Nernst Equation (1) is useful in determining ion concentration of one unknown if all other variables are known. The ionic strength of the 25 mL solution was known to be 0.1 M and the voltage (in mV) was recorded for each titration. 18 Ecell E 0 RT ln H zF (1) A graph was made plotting potential (mV) against volume (mL) of base added throughout the titration. The [H+] is calculated by plugging the x-intercept value into Eq. (2). The pHcalc is found using Eq. (3). [H+]calc = (mLi x Mi) – (mLadd x Madd x 25/x-int)/( mLi + mLadd) pHcalc = -log[H+]calc (2) (3) Following these calculations, a graph is plotted of mV against the pHcalc from Eq. (3). The rest of the variables in the Nernst Equation are constants, and Eq. (1) can be modified to Eq. (4) after plotting mV vs. pHcalc rather than log[H+]. The slope of the linear equation found using the Excel graphing function is the Nernst number (N), and the y-intercept value is E0. Eq. (5) solves for the corrected pH value of any Ecell value, measured experimentally in mV. Ecell = E0 – N pHcorr (4) pHcorr = ( E0 - Ecell ) / N (5) 19 8QP Solution Preparations UV/Vis spectrophotometric titrations involving 8QP The following stock solutions were prepared for acid-base titrations of aqueous 8QP and metal-8QP solutions. A 3.9 x 10-6 M (0.0012 g in 1000 mL DI H2O) solution of 8QP was prepared at a pH ≈ 2 (0.01 M HClO4), with an ionic strength of 0.1 M (0.09 M NaClO4). A 4.5 x 10-6 M (0.0014 g in 1000 mL DI H2O) solution of 8QP was prepared at pH ≈ 0 (1.0 M HClO4), with an ionic strength of 1.0 M. NaOH solutions were prepared at corresponding ionic strength for each titration experiment. A 0.1 M NaOH solution was prepared in a 500 mL volumetric flask adding 5.0 mL of 10.0 M NaOH and filling to volume. A 1.0 M NaOH solution was prepared in a 250 mL volumetric flask adding 25 mL of 10.0 M NaOH and filling to volume. Protonation constants were determined for the free ligand by preparing a 50.00 ± 0.05 mL aliquot of 8QP stock solution to be placed into the flow cell apparatus. The solution was then titrated with the NaOH stock of appropriate molarity. Absorbance values were noted for 204 nm, 223 nm, 243 nm, 291 nm, and 315 nm, and pH was recorded following each titrant addition. Solution for titration of 8QP with copper (II) Three separate titration experiments were performed. A stock copper solution of 0.10 M Cu(ClO4)26H2O (1.8525 g, Aldrich, 99%, in 50 mL of DI H2O) was prepared. For the first titration the concentrations for both the copper and 8QP were 3.9 × 10-6 M. A 50.00 ± 0.05 mL solution was prepared of 3.9 × 10-6 M 8QP stock solution at 0.1 M ionic strength containing 19.5 L of a 0.01 M dilution of the stock copper solution. This solution was placed into the flow cell apparatus as previously described and was titrated with the 0.1 M NaOH stock solution. The second titration was run at a 100:1 copper:8QP concentration. A 50.00 ± 0.05 mL solution was 20 prepared of 3.9 × 10-6 M 8QP stock solution at 0.1 M ionic strength containing 195.0 L of the 0.1 M stock copper solution. This solution was placed into the flow cell apparatus as previously described and was titrated with the 0.1 M NaOH stock solution. The third titration was run at a 1:1 copper:8QP concentration in a 1.0 M ionic strength solution. A 50.00 ± 0.05 mL solution was prepared of 4.0 × 10-6 M 8QP stock solution at 1.0 M ionic strength containing 20.0 L of a 0.01 M dilution of the stock copper solution. This solution was placed into the flow cell apparatus as previously described and was titrated with the 1.0 M NaOH stock solution. Solution for titration of 8QP with zinc (II) Five separate titration experiments were performed. A stock copper solution of 0.0333 M Zn(ClO4)26H2O (0.6333 g, Aldrich, 99%, in 50 mL of DI H2O) was prepared. For the first titration the concentrations for both the copper and 8QP were 3.9 × 10-6 M. A 50.00 ± 0.05 mL solution was prepared of 3.9 × 10-6 M 8QP stock solution at 0.1 M ionic strength containing 6.0 L of the 0.0333 M stock zinc solution. This solution was placed into the flow cell apparatus as previously described and was titrated with the 0.1 M NaOH stock solution. The second titration was run at a 10:1 zinc:8QP concentration. A 50.00 ± 0.05 mL solution was prepared of 3.9 × 10-6 M 8QP stock solution at 0.1 M ionic strength containing 58.5 L of the 0.0333 M stock zinc solution. This solution was placed into the flow cell apparatus as previously described and was titrated with the 0.1 M NaOH stock solution. The third titration was run at a 100:1 zinc:8QP concentration. A 50.00 ± 0.05 mL solution was prepared of 3.9 × 10-6 M 8QP stock solution at 0.1 M ionic strength containing 585.0 L of the 0.0333 M stock zinc solution. This solution was placed into the flow cell apparatus as previously described and was titrated with the 0.1 M NaOH stock solution. The fourth titration was run at a 1:1 zinc:8QP concentration in a 1.0 M ionic 21 strength solution. A 50.00 ± 0.05 mL solution was prepared of 4.0 × 10-6 M 8QP stock solution at 1.0 M ionic strength containing 20.0 L of a 0.01 M dilution of the stock zinc solution. This solution was placed into the flow cell apparatus as previously described and was titrated with the 1.0 M NaOH stock solution. The fifth titration was run at a 1000:1 zinc:8QP concentration in a 1.0 M ionic strength solution. A 50.00 ± 0.05 mL solution was prepared of 4.0 × 10-6 M 8QP stock solution at 1.0 M ionic strength containing 6.00 mL of the 0.0333 M stock zinc solution. This solution was placed into the flow cell apparatus as previously described and was titrated with the 1.0 M NaOH stock solution. Solution for titration of 8QP with nickel (II) Four separate titration experiments were performed. A stock nickel solution of 0.0333 M Ni(ClO4)26H2O (0.6095 g, Aldrich, 99%, in 50 mL of DI H2O) was prepared. The first titration was run at a 10:1 nickel:8QP concentration. A 50.00 ± 0.05 mL solution was prepared of 3.9 × 10-6 M 8QP stock solution at 0.1 M ionic strength containing 58.5 L of the 0.0333 M stock zinc solution. This solution was placed into the flow cell apparatus as previously described and was titrated with the 0.1 M NaOH stock solution. The second titration was run at a 1:1 nickel:8QP concentration in a 1.0 M ionic strength solution. A 50.00 ± 0.05 mL solution was prepared of 4.5 × 10-6 M 8QP stock solution at 1.0 M ionic strength containing 7.0 L of the 0.0333 M stock nickel solution. This solution was placed into the flow cell apparatus as previously described and was titrated with the 1.0 M NaOH stock solution. The third titration was run at a 100:1 nickel:8QP concentration in a 1.0 M ionic strength solution. A 50.00 ± 0.05 mL solution was prepared of 4.5 × 10-6 M 8QP stock solution at 1.0 M ionic strength containing 675.0 L of the 0.0333 M stock nickel solution. This solution was placed into the flow cell apparatus as 22 previously described and was titrated with the 1.0 M NaOH stock solution. The fourth titration was run at a 1000:1 nickel:8QP concentration in a 1.0 M ionic strength solution. A 50.00 ± 0.05 mL solution was prepared of 4.0 × 10-6 M 8QP stock solution at 1.0 M ionic strength containing 6.00 mL of the 0.0333 M stock nickel solution. This solution was placed into the flow cell apparatus as previously described and was titrated with the 1.0 M NaOH stock solution. Solution for titration of 8QP with cadmium (II) Three separate titration experiments were performed. A stock cadmium solution of 0.001 M Cd(ClO4)26H2O (0.0210 g, Aldrich, 99%, in 50 mL of DI H2O) was prepared. For the first titration the concentrations for both the cadmium and 8QP were 3.9 × 10-6 M. A 50.00 ± 0.05 mL solution was prepared of 3.9 × 10-6 M 8QP stock solution at 0.1 M ionic strength containing 195.0 L of the 0.001 M stock cadmium solution. This solution was placed into the flow cell apparatus as previously described and was titrated with the 0.1 M NaOH stock solution. The second titration was run at a 10:1 cadmium:8QP concentration in a 1.0 M ionic strength solution. A 50.00 ± 0.05 mL solution was prepared of 4.5 × 10-6 M 8QP stock solution at 1.0 M ionic strength containing 2.20 mL of the 0.001 M stock cadmium solution. This solution was placed into the flow cell apparatus as previously described and was titrated with the 1.0 M NaOH stock solution. The third titration was run at a 1000:1 cadmium:8QP concentration in a 1.0 M ionic strength solution. A new 0.0963 M Cd(ClO4)26H2O (1.050 g, Aldrich, 99%, in 25 mL of DI H2O) stock solution was prepared. A 50.00 ± 0.05 mL solution was prepared of 4.5 × 10-6 M 8QP stock solution at 1.0 M ionic strength containing 2336.0 L of the 0.0963 M stock cadmium solution. This solution was placed into the flow cell apparatus as previously described and was titrated with the 1.0 M NaOH stock solution. 23 Solution for titration of 8QP with lead (II) A stock lead solution of 0.0033 M Pb(ClO4)26H2O (0.0762 g, Aldrich, 97%, in 50 mL of DI H2O) was prepared. The concentrations for both the lead and 8QP were 3.9 × 10-6 M. A 50.00 ± 0.05 mL solution was prepared of 3.9 × 10-6 M 8QP stock solution at 0.1 M ionic strength containing 58.5 L of the 0.0033 M stock lead solution. This solution was placed into the flow cell apparatus as previously described and was titrated with the 0.1 M NaOH stock solution. Solution for titration of 8QP with iron (III) A stock iron solution of 0.0033 M Fe(ClO4)3.H20 (Aldrich, 99%, in 50 mL of DI H2O) was prepared. The concentrations for both the iron and 8QP were 3.9 × 10-6 M. A 50.00 ± 0.05 mL solution was prepared of 3.9 × 10-6 M 8QP stock solution at 0.1 M ionic strength containing 58.5 L of the 0.0033 M stock iron solution. This solution was placed into the flow cell apparatus as previously described and was titrated with the 0.1 M NaOH stock solution. Solution for titration of 8QP with calcium (II) A stock calcium solution of 0.0033 M Ca(ClO4)26H2O (Aldrich, 99%, in 50 mL of DI H2O) was prepared. The concentrations for both the calcium and 8QP were 3.9 × 10-6 M. A 50.00 ± 0.05 mL solution was prepared of 3.9 × 10-6 M 8QP stock solution at 0.1 M ionic strength containing 58.5 L of the 0.0033 M stock calcium solution. This solution was placed into the flow cell apparatus as previously described and was titrated with the 0.1 M NaOH stock solution. 24 Solution for titration of 8QP with thorium (IV) A stock thorium solution of 0.01 M Th(NO3)4H20 (Aldrich, 99%, in 50 mL of DI H2O) was prepared. The concentrations for both the thorium and 8QP were 3.9 × 10-6 M. A 50.00 ± 0.05 mL solution was prepared of 3.9 × 10-6 M 8QP stock solution at 0.1 M ionic strength containing 19.5 L of the 0.01 M stock thorium solution. This solution was placed into the flow cell apparatus as previously described and was titrated with the 0.1 M NaOH stock solution. Solutions and specifications for fluorescence studies of 8QP The 4.5 x 10-6 M (0.0014 g in 1000 mL DI H2O) stock solution of 8QP with pH ≈ 0 (1.0 M HClO4) and ionic strength of 1.0 M was used for all fluorescence studies. 5.0 mL of this solution was used to determine the intensity of fluorescence of the free ligand. The intensity was recorded for excitation wavelengths from 250-500 nm at 5 nm increments and emission wavelengths ranging from 365-800 nm at 5 nm increments. These results were plotted using the EXCEL graphing function for the data from the highest peak of the free ligand scan at an excitation of 310 nm. For each of the metal-ligand complex solutions, either a 25 or 50 mL stock was prepared at the appropriate concentration. The stock solution was separated into three separate test tubes and brought to a pH level which allowed complexation using a 1.0 M NaOH. The concentration for the copper and 8QP were at a 1:1 ratio at 4.5 × 10-6 M (7.5 L of 0.0304 M stock copper solution in 50 mL stock 8QP solution). The pH was adjusted to 1.95 by adding 1.0 M NaOH. The concentration used for the Zn2+ complex was 10:1 with 8QP (112.5 L of 0.01 M stock zinc solution in 25 mL stock 8QP solution). The pH was adjusted to 7.3 by adding 1.0 M NaOH. The concentration used for the Ni2+ complex was 10:1 with 8QP (161 L of 0.0070 M stock nickel solution in 25 mL stock 8QP solution). The pH was adjusted to 4.0 by 25 adding 1.0 M NaOH. The intensity was recorded for excitation wavelengths from 250-500 nm at 5 nm increments and emission wavelengths ranging from 365-800 nm at 5 nm increments. These results were plotted using the EXCEL graphing function for the data at an excitation of 310 nm. PDALD Solution Preparations UV/Vis spectrophotometric titrations involving PDALD The following stock solution was prepared for acid-base titrations of aqueous PDALD and metal-PDALD solutions. A 2.0 x 10-5 M (0.0047 g in 1000 mL DI H2O) solution of PDALD was prepared at a pH ≈ 2 (0.01 M HClO4), with an ionic strength of 0.1 M (0.09 M NaClO4). A 0.1 M NaOH solution was prepared in a 500 mL volumetric flask adding 5.0 mL of 10.0 M NaOH and filling to volume. Protonation constants were determined for the free ligand by preparing a 50.00 ± 0.05 mL aliquot of PDALD stock solution to be placed into the flow cell apparatus. The solution was then titrated with the 0.1 M NaOH solution. Absorbance values were noted for 209 nm, 225 nm, 256 nm, 270 nm, and 282 nm, and pH was recorded following each titrant addition. Solution for titration of PDALD with copper (II) A stock copper solution of 0.0304 M Cu(ClO4)26H2O (0.5632 g, Aldrich, 99%, in 50 mL of DI H2O) was prepared. The concentrations for both the copper and PDALD were 2.0 × 10-5 M. A 50.00 ± 0.05 mL solution was prepared of 2.0 × 10-5 M PDALD stock solution at 0.1 M ionic strength containing 33.0 L of the 0.0304 M stock copper solution. This solution was 26 placed into the flow cell apparatus as previously described and was titrated with the 0.1 M NaOH stock solution. Solution for titration of PDALD with cadmium (II) A stock cadmium solution of 0.0333 M Cd(ClO4)26H2O (0.6993 g, Aldrich, 99%, in 50 mL of DI H2O) was prepared. The concentrations for both the cadmium and PDALD were 2.0 × 10-5 M. A 50.00 ± 0.05 mL solution was prepared of 2.0 × 10-5 M PDALD stock solution at 0.1 M ionic strength containing 30.0 L of the 0.0333 M stock cadmium solution. This solution was placed into the flow cell apparatus as previously described and was titrated with the 0.1 M NaOH stock solution. Solution for titration of PDALD with calcium (II) A stock calcium solution of 0.01 M Ca(ClO4)2H2O (Aldrich, 99%, in 50 mL of DI H2O) was prepared. The concentrations for both the calcium and PDALD were 2.0 × 10-5 M. A 50.00 ± 0.05 mL solution was prepared of 2.0 × 10-5 M PDALD stock solution at 0.1 M ionic strength containing 100.0 L of the 0.01 M stock cadmium solution. This solution was placed into the flow cell apparatus as previously described and was titrated with the 0.1 M NaOH stock solution. Solution for titration of PDALD with gadolinium (III) A stock gadolinium solution of 0.0357 M Gd(ClO4)36H2O (1.0060 g, Aldrich, 99%, in 50 mL of DI H2O) was prepared. The concentrations for both the gadolinium and PDALD were 2.0 × 10-5 M. A 50.00 ± 0.05 mL solution was prepared of 2.0 × 10-5 M PDALD stock solution at 0.1 M ionic strength containing 28.0 L of the 0.0357 M stock gadolinium solution. This solution 27 was placed into the flow cell apparatus as previously described and was titrated with the 0.1 M NaOH stock solution. Solution for titration of PDALD with lead (II) A stock lead solution of 0.0106 M Pb(ClO4)26H2O (0.2454 g, Aldrich, 97%, in 50 mL of DI H2O) was prepared. The concentrations for both the lead and PDALD were 2.0 × 10-5 M. A 50.00 ± 0.05 mL solution was prepared of 2.0 × 10-5 M PDALD stock solution at 0.1 M ionic strength containing 94.0 L of the 0.0106 M stock lead solution. This solution was placed into the flow cell apparatus as previously described and was titrated with the 0.1 M NaOH stock solution. 28 RESULTS AND DISCUSSION Synthesis of PDALD In previous publications31, the synthetic technique used in this study was found to result in low yields of impure 1,10-phenanthroline-2,9-dicarboxyaldehyde (PDALD). The overall yield in our case was 22.9% and was determined to be pure based upon both 1H-NMR spectroscopy and FT-IR analysis, making further purification and product recovery unnecessary. The FT-IR spectra can be seen in Figure 13. The characteristic IR stretch was observed for the C=O at 1700 cm-1. The 1H-NMR spectra for PDALD can be seen in Figure 14a. The aldehyde protons show a peak at 10.34 ppm (H2, H9, singlet). The aromatic protons on the phenanthroline ring system can be seen at 8.79 (H4, H7, doublet), 8.29 (H3, H8, doublet), and 8.27 (H5, H6, singlet) ppm. These results correspond with the reported values in Chandler, et al.31 Figure 13 : FT-IR spectra of 1,10-phenanthroline-2,9-dicarboxyaldehyde (PDALD) 29 Figure 14a : 1 H-NMR spectrum of 1,10-phenanthroline-2,9-dicarboxyaldehyde (PDALD) in DMSO-d6. 30 Figure 14b : 1H-NMR spectrum of neocuprine in DMSO-d6. There were concerns about the small singlet peak at 3.53 ppm in the 1H-NMR spectra for PDALD. To determine that this peak was not due to the methyl hydrogen atoms of unreacted starting material, 1H-NMR analysis of neocuprine was also done in DMSO-d6. The resulting spectrum, seen in Figure 14b, shows the neocuprine methyl hydrogen peak at 3.03 ppm. A shift this large would not be typical if these were the same methyl hydrogens. These results indicate that although there may be a contaminant present in the sample, it is unlikely to be unreacted neocuprine. This scan also shows a small shift in both sets of aromatic doublets in the product spectrum. This shift is due to the electron withdrawing character of the aldehyde groups of the PDALD. 31 8QP Protonation Constants In order to determine the strength at which particular metals bind, it is necessary to determine the protonation constants (pK values) for 8QP. UV/vis spectroscopy was used as previously discussed. Titrations were performed at 25.0 ± 0.1 °C using a 4.5 x 10-6 M solution of 8QP at 1.0 M ionic strength (1.0 M NaClO4). Figure 16 shows a series of absorbance versus wavelength scans at a pH range from approximately 0.50 to 8.50. Absorbance data from 204 nm, 223 nm, 243 nm, 291 nm, and 315 nm were used to generate plots of absorbance versus pH. The points drawn in are experimental values and the solid lines are theoretical curves of absorbance versus pH calculated for the constants corresponding to the observed protonation equilibria derived using Excel.32 8QP has two separate protonation equilibria, pK1 and pK2. To determine the value of the protonation constants from the observed absorbances, it was first necessary correct each absorbance for dilution using Eq(6). AbsCorr Abs VTotal Vinitial (6) Plots of Abscorr versus pH were constructed for each wavelength selected. The total ligand concentration, [L]total, in solution can be described by Eq(7). LTotal [L] [LH] [LH2] (7) Eq(7) can be rearranged by adding the following protonation constants to get Eq(10). Each Eq(8)-(9) represent a separate protonation equilibrium. Ka1 [LH] [L][H] Ka1Ka 2 [LH2] [L][H]2 32 (8) (9) LTotal [L] Ka1[L][H] Ka1Ka 2[L][H]2 (10) By dividing out the ligand concentration, [L], Eq(10) can be simplified to Eq(11). LTotal 1 Ka1[H + ] Ka1Ka 2[H+ ]2 [L] (11) Theoretical absorbance, Abs(theor), in Eq(12) was calculated by multiplying the concentration of the species present in solution [L, LH, LH2] by the absorbance of each of these species at a 4.5 x 10-6 M concentration, as shown in Eq(11). To explain Table 2, which is an example of the spreadsheet used to calculate protonation equilibria, each term in Eq(11) was described as a function, for example L(func)1 = Ka1[H+]. 1 [Abs(L)] Ka1[H+ ][Abs(LH)] Ka1Ka 2[H + ]2[Abs(LH2)] Abs(theor) 1 Ka1[H + ] Ka1Ka 2[H + ]2 (12) Abs(L) is the absorbance where only unprotonated ligand exists in the sample solution. Abs(LH), and Abs(LH2) describe the absorbances at each protonation equilibrium. These are labeled Abs(0), Abs(1), and Abs(2) respectively in each spreadsheet constructed to determine stability constants. Plots of pH versus corrected absorbance were fit with plots of pH versus Abs(theor) using the „SOLVER‟ function of EXCEL at each wavelength selected. Figure 17 shows the plot of the 223 nm fit from Table 2 and Figure 18 shows all of the wavelengths fit simultaneously to calculate pKa. The protonation constants for 8QP were calculated using the absorbance data and pH values from this plot. The corrected protonation constants of pK1 and pK2 for 8QP were 2.40 and 6.82, respectively. A summary of the protonation equilibria for 8QP is shown in Table 4. 33 Figure 15 : Overview of the absorbance v. wavelength spectra for the 8QP free ligand at 4 different pH‟s to show the progression of the spectral curve. 34 Figure 16 : Absorbance v. wavelength spectra for the 8QP free ligand for pH ≈ 0.5 – 8.5. Figure 17 : Plot of measured and theoretical absorbance vs. pH for 8QP at 223 nm as shown in Table #. Measured points are seen in blue while the theoretical curve is plotted in pink. 35 pH mV pH (mVcorr) mL NaOH total Vadd Vtotal Abs. (223nm) Abs(corr) L(func)1 L(func)2 L(func)3 L(func)5 A(theor) 0.65 364 0.41 0.00 0.000 50.000 0.3834 0.3834 3.97E+06 4.29E+08 7.71E+04 4.33E+08 0.4610 0.77 356 0.56 25.00 25.000 75.000 0.2776 0.4164 2.85E+06 2.21E+08 2.84E+04 2.23E+08 0.4441 1.02 342 0.81 13.00 38.000 88.000 0.2406 0.4235 1.59E+06 6.89E+07 4.96E+03 7.05E+07 0.4244 1.26 328 1.06 5.00 43.000 93.000 0.2215 0.4120 8.88E+05 2.15E+07 8.65E+02 2.24E+07 0.4119 1.78 297 1.62 4.00 47.000 97.000 0.2011 0.3901 2.45E+05 1.63E+06 1.81E+01 1.88E+06 0.3902 2.21 270 2.11 1.00 48.000 98.000 0.1860 0.3646 7.97E+04 1.73E+05 6.24E-01 2.53E+05 0.3627 2.55 249 2.49 0.30 48.300 98.300 0.1723 0.3387 3.33E+04 3.02E+04 4.54E-02 6.34E+04 0.3338 3.12 218 3.05 0.20 48.500 98.500 0.1451 0.2858 9.17E+03 2.29E+03 9.51E-04 1.15E+04 0.2962 3.51 195 3.46 0.05 48.545 98.545 0.1389 0.2738 3.52E+03 3.38E+02 5.40E-05 3.86E+03 0.2809 4.65 128 4.67 0.03 48.570 98.570 0.1350 0.2661 2.17E+02 1.29E+00 1.27E-08 2.20E+02 0.2696 5.17 97 5.23 0.06 48.630 98.630 0.1364 0.2691 5.99E+01 9.78E-02 2.65E-10 6.10E+01 0.2685 5.50 78 5.58 0.02 48.650 98.650 0.1360 0.2683 2.72E+01 2.01E-02 2.48E-11 2.82E+01 0.2676 5.94 52 6.05 0.03 48.680 98.680 0.1359 0.2682 9.23E+00 2.32E-03 9.69E-13 1.02E+01 0.2648 6.32 30 6.44 0.02 48.700 98.700 0.1341 0.2647 3.70E+00 3.72E-04 6.23E-14 4.70E+00 0.2597 6.62 12 6.77 0.02 48.720 98.720 0.1318 0.2602 1.75E+00 8.33E-05 6.60E-15 2.75E+00 0.2531 6.87 -3 7.04 0.02 48.740 98.740 0.1280 0.2528 9.37E-01 2.39E-05 1.02E-15 1.94E+00 0.2465 7.08 -15 7.26 0.03 48.770 98.770 0.1214 0.2398 5.69E-01 8.83E-06 2.27E-16 1.57E+00 0.2412 7.70 -52 7.92 0.03 48.800 98.800 0.1167 0.2306 1.22E-01 4.07E-07 2.25E-18 1.12E+00 0.2301 8.14 -78 8.39 0.03 48.830 98.830 0.1129 0.2232 4.15E-02 4.68E-08 8.79E-20 1.04E+00 0.2270 8.52 -100 8.79 0.03 48.860 98.860 0.1127 0.2228 1.66E-02 7.52E-09 5.65E-21 1.02E+00 0.2260 8.90 -122 9.19 0.05 48.910 98.910 0.1174 0.2322 6.66E-03 1.21E-09 3.64E-22 1.01E+00 0.2256 10.05 -190 10.42 0.04 48.950 98.950 0.1350 0.2672 3.94E-04 4.23E-12 7.54E-26 1.00E+00 0.2253 Table 2. EXCEL spreadsheet for free ligand used to calculate protonation equilibria. 36 Solutions for each parameter solved by the „SOLVER‟ module of EXCEL in Table 3. determining pKa of 8QP. 204 nm 223 nm 243 nm 291 nm 315 nm Abs(0) Abs(1) Abs(2) Abs(0) Abs(1) 0.0003 0.0450 0.1265 -0.0002 0.0482 Abs(2) 0.1754 pK1 6.82 Abs(0) 0.0871 pK2 2.40 Abs(1) Abs(2) Abs(0) Abs(1) Abs(2) Abs(0) Abs(1) Abs(2) 0.1120 0.3135 0.2253 0.2690 0.4048 0.7649 0.6725 0.7329 Table 4. Protonation equilibria and constants for 8QP free ligand. L + H+ LH + H+ 37 LH pKa 6.82 LH2 2.40 0.5 + LH + H = LH 2 pK 2 = 2.40 absorbance 0.4 + L + H = LH pK 1 = 6.82 0.3 223 nm 291 nm 0.2 243 nm 0.1 315 nm 0.0 0.0 2.0 4.0 6.0 8.0 10.0 12.0 pH Figure 18 : Variation of absorption at the wavelengths indicated, as a function of pH for 4.0 x 10-6 M 8QP in 1.0 M NaClO4 at 25 °C. The points are experimental values of absorbance, and the solid lines are theoretical curves of absorbance vs pH calculated on the basis of two protonation constants, of 6.82 and 2.40 for 8QP. PDALD Protonation Constants The same approach was used to determine the protonation constants for the PDALD. Titrations were performed at 25.0 ± 0.1 °C at 0.1 and 1.0 M ionic strength (0.1 M and 1.0 M NaClO4). Figure 19 shows a series of absorbance versus wavelength scans at a pH range from approximately 2.25 to 12.50. Absorbance data from 209 nm, 225 nm, 256 nm, 270 nm, and 282 nm were used to generate plots of absorbance versus pH. The actual and theoretical curves of absorbance versus pH can be seen in Figure 22. The protonation constants for PDALD were calculated using the method described above. The corrected protonation constants of pK1 and pK2 for PDALD were 7.05 and 2.50, respectively. Data from v.7.0 of the 2003 NIST Standard 38 Reference Database24 for similar ligands suggests that there may be an additional pK value present in PDALD near pH = 12, however the presence of this pKa value could not be verified by the UV-vis methods used in this study. A summary of the protonation equilibria for PDALD is shown in Table 6. Figure 19 : Overview of the absorbance v. wavelength spectra for the PDALD free ligand at 4 different pH‟s to show the progression of the spectral curve. 39 Figure 20 : Absorbance v. wavelength spectra for the PDALD free ligand from pH ≈ 2.0 – 7.5. Figure 21 : Plot of measured and theoretical absorbance vs. pH for PDALD at 270 nm as shown in Table #. Measured points are seen in blue while the theoretical curve is plotted in pink. 40 Solutions for each parameter solved by the „SOLVER‟ module of EXCEL in Table 5. determining pKa of PDALD. 209 nm 225 nm 256 nm 270 nm 282 nm Abs(0) Abs(1) Abs(2) Abs(0) Abs(1) 0.5835 0.3856 0.4274 0.7957 0.4909 Abs(2) 0.3979 pK1 7.05 Abs(0) 0.2699 pK2 2.50 Abs(1) Abs(2) Abs(0) Abs(1) Abs(2) Abs(0) Abs(1) Abs(2) 0.1819 0.1652 0.3258 0.2461 0.2331 0.4277 0.3375 0.3062 Table 6. Protonation equilibria and constants for PDALD free ligand. L + H+ LH + H+ 41 LH pKa 7.05 LH2 2.50 Figure 22 : Variation of absorption at the wavelengths indicated, as a function of pH for 2.0 x 10-5 M PDALD in 0.1 M NaClO4 at 25 °C. The points are experimental values of absorbance, and the solid lines are theoretical curves of absorbance vs pH calculated on the basis of two protonation constants, of 7.05 and 2.50 for PDALD. UV-Vis spectrophotometric titrations involving metal ion complexation UV-Vis spectroscopy was used as an analytical tool to determine the stability constants (logK1) of metal-ligand complexes; the ligands in this study being 8QP and PDALD. For each titrant addition of either 0.1 M or 1.0M NaOH (specified in each case), absorbance scans were made from 200 to 350 nm. Absorbance values were recorded for all 8QP titrations at 204 nm, 42 223 nm, 243 nm, 291 nm, and 315 nm. Absorbance values for all PDALD titrations were taken at 209 nm, 225 nm, 256 nm, 270 nm, and 282 nm. These wavelengths were chosen because they corresponded to specific peaks or points of interest on the free ligand spectra. Upon complexation with a metal ion, the peaks on the absorbance spectra will shift from that of the free ligand spectra and allow determination of a logK1 value for the particular metal-ligand complex. To determine the logK1 for metal ion stability, the procedure outlined previously was performed in the presence of metal ions of interest and absorbance data were taken at the wavelengths described. The presence of the metal ions affects the protonation equilibria observed in the free ligand titrations in two ways. One way is that the protonation of the ligand now involves the displacement of the metal ion by protons: ML H+ M LH (13) Equation 13 is a simple example which shows one proton attached to the ligand and logK1 for as follows. Evidence of a protonation of the free ligand with the ML complex can be calculated no metal present is shown by an inflexion of absorbance versus pH. The pKa is at the midpoint of this inflexion. In the presence of the metal ion, if a complex is formed, an inflexion in the absorbance versus pH curve is observed, but now at lower pH. This protonation equilibrium corresponds to Eq. (13). The reaction constant in Eq. (14) can be calculated from the position of the midpoint of the inflexion corresponding to Eq. (13). Kreact [LH][M] [ML][H+ ] (14) In equation 14, [H+] is the proton concentration at pH50, which is the pH in Eq. (14) where [LH] free metal ion concentration [M] must also be included. This = [ML]. In calculating Kreact, the means that Kreact will equal the free metal ion concentration [M], which at pH50 will be 50% of 43 the total metal ion concentration ([ML] = [M]), divided by [H+] at pH50. K1 for the metal ion complex now corresponds to the constant Kreact for Eq. (14) combined with the protonation constant Ka: K1 Ka K react (15) [LH] [ML][H+ ] [ML] [L][H + ] [LH][M] [L][M] (16) 8QP is a diprotic ligand, which is calculated the same as above except that both of the protonation constants are involved. ML + 2H+ LH2 + M (17) [LH2][M] [ML][H+ ]2 (18) Ka1Ka 2 Kreact (19) [LH2] [ML][H+ ]2 [ML] + 2 [L][H ] [LH 2][M] [L][M] (20) Kreact K1 Eq. (1)-(20) were used to calculate the protonation equilibria of each metal-ligand complex with There is a difference in the pK values for the free ligand and the minimal standard deviation. complex because of the inflexion at a lower pH as previously discussed. Therefore in the presence of a metal ion, logK1 can be calculated by: logK1 (ML) = -log[M] + (pK1 – pK50) + (pK2 – pK50) (21) The other source of protonation equilibria occur as the result of a hydroxide being added or removed from the metal-ligand complex as seen in Eq. (22). MLOH + H+ 44 ML (22) The pK1 was determined for these protonation equilibria when they were evident; however they were not used in determining complex stability. MM Calculations Using HyperChem 8-quinolyl-1,10-phenanthroline Molecular mechanics calculations made using HyperChem (v.5.11) were used to determine which metal ion should have the best fit with 8QP based on the bond energy (U) of the metal-nitrogen bonds of various 8QP complexes. A plot was generated showing the plot of U (kcal/mol) versus the M-N bond length (Å) and can be seen in Figure 23. The sloping curve of this graph shows that as bond length decreases, less strain is put on the M-N bond. These calculations allow us to predict that smaller metal ions, such as Zn2+ and Cu2+, will form more stable complexes with 8QP. Larger metal ions, such as Ca2+, will have difficulty forming stable complexes with 8QP or even complexing at all. Molecular mechanics models have also provided evidence of other potential problems that larger metal ions may face when trying to complex with 8QP. In addition to having a much higher bond energy, larger metal ions are physically blocked from complexing by the arrangement of the hydrogens at the coordination site. As seen in the [Cd(8QP)(H2O)2] MM space fill diagram (Figure 24), the cadmium ion is blocked in by the two labeled hydrogens. These hydrogens may block access to the coordination site and make it difficult for larger metal ions to complex. Another thing that HyperChem models have shown is that when the metal ion in too large to fit at the coordination site properly, the ligand actually twists around the bond connecting the quinolyl group and the phenanthroline. This is particularly noticeable in the [Ca(8QP)(H2O)3] 45 MM diagram (Figure 25). This twisting is in energetically unfavorable and larger metal ions which force the ligand too far out of a planar configuration will likely not form a complex. More favorable conformations can be seen with smaller metal ions such as Zn2+ (Figure 26), which allow 8QP to remain relatively planar allow for more ideal bond length between the M-N bond. Figure 23 : Calculated strain energy (kcal/mol) versus metal-nitrogen bond length (Å) curve for 8QP. The arrows indicate points on the curve where specific metal ions fit the curve. 46 Figure 24 : [Cd(8QP)(H2O)2] MM – Space-filling model showing that a larger metal ion such as Cd2+ (0.95 Å) has difficulty coordinating with 8QP due to the difficulty getting access to the binding site because of the flanking hydrogen atoms indicated with arrows above. 47 Figure 25 : [Ca(8QP)(H2O)3] – HyperChem molecular mechanics diagram showing how the 8QP molecule twists around the bond connecting the quinolyl and phenanthroline groups when accommodating a metal ion like Ca2+ (0.99 Å) which is too large for the cleft. Figure 26 : [Zn(8QP)(H2O)2] PM3 – At 0.74 Å, the Zn2+ ion is small enough that the 8QP cleft can accommodate it easily, without having to twist very far out of the more thermodynamically stable planar conformation. 48 The stability constants (logK1) determined for metal ions with 8QP from UV-Vis spectroscopy titration experiments can be seen in Table 7. Results of the UV-Vis experiments tended to complement the expected results based upon the HyperChem MM calculations. Table 7. Protonation constants, and formation constants with a selection of metal ions with 8QP (= L in the table below), in 1.0 M NaClO4 at 25 °C. * (determined at 0.1 M NaClO4 at 25 °C) 49 1,10-phenanthroline-2,9-dicarboxyaldehyde Molecular mechanics calculations were made using HyperChem (v.5.11) to determine the likelihood of complex formation for all the metal ions used in this study. MM results showed that even large metal ions such as Pb2+ had little trouble fitting into the binding site of PDALD (Figure 27). Chelate ring size rule dictates that the five-membered ring which is formed when a metal ion complexes should be selective for larger metal ions. Figure 27 : [Pb(PDALD)(OH)2] complex designed using MM calculations in HyperChem. This diagram illustrates that even very large metal ions like Pb2+ fit well within the 5-membered cleft of PDALD. This complex is further stabilized by coordinated, aldehydic neutral oxygen donors. The stability constants (logK1) determined for metal ions with PDALD from UV-Vis spectroscopy titration experiments can be seen in Table 8. Results of the UV-Vis experiments tended to complement the expected results based upon the HyperChem MM calculations. 50 Table 8. Protonation constants, and formation constants with a selection of metal ions with PDALD (= L in the table below), in 1.0 M NaClO4 at 25 °C. 51 Titrations Involving 8QP 8QP-Copper (II) Results Cu2+ has an ionic radius of 0.57 Å, which makes it the smallest metal ion that was used in this study. Copper is categorized by Pearson as being an intermediate acid which means that it prefers to complex with intermediate donors, such as the nitrogen donors in 8QP. Because of its affinity for nitrogen donors and small size, copper was expected to form a rather stable complex with 8QP. A solution of 3.9 x 10-6 M 8QP and 3.9 x 10-5 M Cu(ClO4)2 was titrated with 0.1 M NaOH from pH = 2.15 to 7.18. UV-vis absorbance spectra are shown in Figure 28 for the titration of Cu2+ with 8QP. Absorbance values were recorded at 204, 223, 243, 291, and 315 nm because these wavelengths showed the greatest change in absorbance. Corrected absorbance and theoretical absorbance were plotted for every wavelength simultaneously against pH and can be seen in Figure 29a. Table 10 outlines the solutions of each parameter that was varied by the „SOLVER‟ function and the resulting solved pK values. The pK values were recorded once all the standard deviations were minimized and the resulting equations for the protonation equilibria observed during the 8QP-Cu titration can be seen in Table 11. No logK1 was calculated from this experiment because both pK50 values calculated were deprotonation of a water molecule attached directly to the copper. According to the UV/vis spectra, the complex had formed at the initial pH of 2.15. The spectra change as pH is increased but the spectrum for the free ligand was never seen. Results from the wavelengths described above give pK values for the formation of the metal-ligand-hydroxide (MLOH) complex as shown in Eq. (22) as well as an ML(OH)2 complex at 4.78 and 6.59 respectively. The deprotonation of a water molecule coordinated to the metal ion and occurred at 6.59 and allowed us to use Eq. (22) to determine a logK(ML(OH)2) of 52 7.19. The formation of the hydroxide complex had a pK value of 4.78 and Eq. (22) was used to calculate logK(MLOH) of 9.0. In order to find the logK values for the formation of the Cu(8QP) complex, another titration was done to much lower pH using a 1.0 M HClO4 solution of 4.0 x 10-6 M 8QP and 4.0 x 10-6 M Cu(ClO4)2 which was titrated with 1.0 M NaOH from pH = 0.55 to 3.76. UV-vis absorbance spectra are shown in Figure 30 for the titration of Cu2+ with 8QP at the lower pH. The free ligand is evident for the first two scans before the complex is formed. Absorbance values were recorded at 204, 223, 243, 291, and 315 nm. The corrected absorbance and theoretical absorbance were plotted for every wavelength simultaneously against pH and can be seen in Figure 29b. Results from the wavelengths described above show pK1 and pK2 values at 2.00 and 0.32 respectively. The pK2 at 0.32 was the first time that we were able to see the free ligand in the UV-vis spectra and was used in Eq. (21) to calculate a logK1 of 14.4. The pK1 at 2.00 corresponds to a protonation of the Cu(8QP) complex. The stability of the [Cu(8QP)] complex can be attributed to the preorganization of the 8QP ligand. Small metal ions such as Cu2+ fit well into the binding site and allow the ligand to retain a more planar configuration than metal ions of larger size. When compared with the binding constants of the ligand terpyridine (terpy) with copper, the 8QP:Cu2+ complex was significantly more stable. This was expected due to the 6-membered ring of 8QP. 53 Figure 28 : UV spectra of 3.9 x 10-6 M 8QP and 4 x 10-5 M Cu(ClO4)2 in 0.01 M HClO4 and 0.09 M NaClO4. Initial spectrum of 8QP-Cu complex at pH = 2.15. Final spectrum of 8QP- Cu hydroxide complex at pH = 7.18. 54 Figure 29 : Plot of corrected absorbance (data points) and theoretical absorbance (solid line) versus pH for titration of 8QP and Cu(ClO4)2 for varying wavelengths. a) 0.1 M from pH = 2.15 – 7.18 and b) at 1.0 M from pH = 0.55 – 3.76. 55 Figure 30 : UV spectra of 4.0 x 10-6 M 8QP and 4.0 x 10-6 M Cu(ClO4)2 in a 1.0 M (0.1 M HClO4 and 0.9 M NaClO4) solution. a) Initial spectrum of free 8QP at pH = 0.55. b) 8QP-Cu complex forming at pH = 0.73 c) 8QP-Cu complex at pH = 2.92 from a), b), and c) e) Full run spectra from pH = 0.55 – 3.76. 56 d) Combination of spectra Table 9. Solutions for the pKa and absorbance parameters solved by the „SOLVER‟ module of EXCEL in determining logK1 of 8QP-Cu complex. Left column – was a) in Figure 29. Right column – was b) in Figure 29. Overall 204 nm 223 nm 243 nm 291 nm 315 nm Parameter pK1 pK2 Abs(0) Abs(1) Abs(2) Abs(0) Abs(1) Abs(2) Abs(0) Abs(1) Abs(2) Abs(0) Abs(1) Abs(2) Abs(0) Abs(1) Abs(2) 0.1 M 6.59 4.78 0.2730 0.2720 0.2600 0.1310 0.1620 0.1780 0.0677 0.1001 0.1245 0.0247 0.0415 0.0535 0.0168 0.0422 0.0585 1.0 M 2.00 0.32 0.8803 0.8805 0.0000 0.5139 0.4989 0.1818 0.3752 0.3589 0.2713 0.1666 0.1551 0.1233 0.1774 0.1748 0.0697 Table 10. Summary of equations at each protonation equilibrium of Cu(8QP). pK [Cu(8QP)(OH)2] + H+ [Cu(8QP)OH] + H2O 6.59 [Cu(8QP)OH] + H+ [Cu(8QP)] + 2H2O 4.78 [Cu(8QP)] + H+ [Cu(8QP)H] [Cu(8QP)H] + H+ H2(8QP) 57 2.00 + Cu2+ 0.32 8QP-Nickel (II) Results Ni2+ has an ionic radius of 0.69 Å, which makes it one of the smaller metal ions explored in this study. Nickel is categorized by Pearson as being an intermediate acid which means that it prefers to complex with intermediate donors, such as the nitrogen donors in 8QP. Because of its affinity for nitrogen donors and small size, nickel was expected to form a complex with 8QP. This complex is expected to be slightly weaker than complexes formed with smaller sized metal ions (Cu2+) because the bond lengths between the M-N bond lengths in the formed fivemembered ring would be shorter than ideal and the M-N bond length in the formed sixmembered ring would be slightly longer than ideal. This prediction is according to the chelate ring size rule previously discussed. A solution of 4.5 x 10-6 M 8QP and 4.5 x 10-6 M Ni(ClO4)2 was titrated with 1.0 M NaOH from pH = 0.81 to 9.30. Absorbances values were recorded for 204, 223, 243, 291, and 315 nm because these wavelengths exhibited the greatest amount of change over the course of the titration. Corrected absorbance and theoretical absorbance were plotted for every wavelength simultaneously against pH and can be seen in Figure 31a. The solutions of each parameter were found using the „SOLVER‟ function and the resulting pK values were computed. The pK values were recorded once all the standard deviations were minimized. Results from the wavelengths described above show apparent pK1 and pK2 values at 7.16 and 2.30 respectively. These pK values are close to those found for the free ligand; however the spectra suggested that a complex may have formed so the titrations were rerun with higher concentrations of nickel. In order to determine that there was a Ni-8QP complex formed, a titration was performed at 1000 to 1 ratio Ni2+:8QP. A solution of 4.0 x 10-6 M 8QP and 4.0 x 10-3 M Ni(ClO4)2 was titrated with 1.0 M NaOH from pH = 0.64 to 7.60. UV-vis absorbance spectra are shown in 58 Figure 32 for the 1000 to 1 titration of Ni2+ with 8QP. Absorbances were again recorded at 204, 223, 243, 291, and 315 nm and corrected absorbance and theoretical absorbance were plotted for every wavelength simultaneously against pH and can be seen in Figure 31b. Table 11 outlines the solutions of each parameter that was varied by the „SOLVER‟ function and the resulting solved pK values. The pK values were recorded once all the standard deviations were minimized and the resulting equations for the protonation equilibria observed during the 8QP-Ni titration can be seen in Table 12. Results from the wavelengths described above show an additional pK value. Apparent pK1, pK2, and pK3 values were seen at 5.55, 0.78, and 0.19 respectively. The first pKa value of 5.55 corresponds to an MLOH complex as seen in Eq. (22). Using these pKa values and Eq. (21), the calculated logK1 value for the 8QP-Ni2+ complex at 1:1 and 1000:1 ratio was found to be 9.9 and 10.1 respectively. At 10.0 the binding strength of the 8QP:Ni2+ complex is very similar to that of the corresponding terpy complex (10.7). 59 Figure 31 : Plot of corrected absorbance (data points) and theoretical absorbance (solid line) versus pH for titration of 8QP and Ni(ClO4)2 for varying wavelengths. a) 1 to 1 ratio and b) 1000 to 1. 60 Figure 32 : UV spectra of 4.0 x 10-6 M 8QP and 4.0 x 10-3 M Ni(ClO4)2 in a 1.0 M (0.1 M HClO4 and 0.9 M NaClO4) solution. a) Initial spectrum of free 8QP at pH = 0.64. b) 8QP-Ni complex formed at pH = 2.00 c) at pH = 7.60 d) Combination of spectra from a), b), and c) e) Full run spectra from pH = 0.64 – 7.60. 61 Table 11. Solutions for the pKa and absorbance parameters solved by the „SOLVER‟ module of EXCEL in determining logK1 of 8QP-Ni complex at 1000:1. 204 nm 223 nm 243 nm 291 nm 315 nm Abs(0) Abs(1) Abs(2) Abs(0) Abs(1) 0.6268 0.6232 0.0000 0.2981 0.4183 Abs(2) 0.2367 pK1 0.78 Abs(0) 0.2114 pK2 0.19 Abs(1) Abs(2) Abs(0) Abs(1) Abs(2) Abs(0) Abs(1) Abs(2) 0.1907 0.4410 0.0795 0.0902 0.2973 0.0495 0.1004 0.1351 Table 12. Summary of equations at each protonation equilibrium of Ni(8QP). pK [Ni(8QP)OH] + H+ [Ni(8QP)] + H2O 5.55 [Ni(8QP)] + H+ H(8QP) + Ni2+ 0.78 H(8QP) + H+ H2(8QP) 62 0.19 8QP-Zinc (II) Results Zn2+ has an ionic radius of 0.74 Å, which slightly larger than nickel. Zinc is categorized by Pearson as being an intermediate acid which means that it prefers to complex with intermediate donors, such as the nitrogen donors in 8QP. The logK1 of the [Zn(8QP)] complex is expected to be slightly lower than that of the [Ni(8QP)] because of the increase in ionic radius. A solution of 3.9 x 10-6 M 8QP and 3.9 x 10-6 M Zn(ClO4)2 was titrated with 0.1 M NaOH from pH = 2.10 – 7.06. Absorbances values were recorded for 204, 223, 243, 291, and 315 nm because these wavelengths exhibited the greatest amount of change over the course of the titration. Corrected absorbance and theoretical absorbance were plotted for every wavelength simultaneously against pH and can be seen in Figure 33b. The solutions of each parameter were found using the „SOLVER‟ function and the resulting pK values were computed. The pK values were recorded once all the standard deviations were minimized. Results from the wavelengths described above show apparent pK1 and pK2 values at 6.00 and 2.50 respectively. Using these pKa values and Eq. (21), the calculated logK1 value for the Zn2+-8QP complex was found to be 9.73. Another titration was performed at a 10:1 Zn2+ to 8QP ratio using a solution of 3.9 x 10-6 M 8QP and 4.0 x 10-5 M Zn(ClO4)2 , titrated with 0.1 M NaOH from pH = 2.09 to 7.05. UV-vis absorbance spectra are shown in Figure 34 for the titration of Zn2+ with 8QP. Absorbance values were recorded at 204, 223, 243, 291, and 315 nm and corrected absorbance and theoretical absorbance were plotted for every wavelength simultaneously against pH and can be seen in Figure 33a. Table 13 outlines the solutions of each parameter that was varied by the „SOLVER‟ function and the resulting solved pK values. The pK values were recorded once all the standard deviations were minimized and the resulting equations for the protonation equilibria observed during the 8QP-Ni titration can be seen in Table 14. Results from the wavelengths described 63 above show apparent pK1 and pK2 values at 6.10 and 1.75 respectively. Using these pKa values and Eq. (21), the calculated logK1 value for the Zn2+-8QP complex was found to be 9.48. In order to get data from below pH ≈ 2.0 and assure accurate pK values, a 1.0 M solution of 4.0 x 10-6 M 8QP and 4.0 x 10-6 M Zn(ClO4)2 was prepared and titrated with 1.0 M NaOH from pH = 0.65 to 7.05. Absorbance values were recorded as stated above. Results from the wavelengths described above show apparent pK1 and pK2 values at 2.70 and 0.72 respectively. Using these pKa values and Eq. (21), the calculated logK1 value for the Zn2+-8QP complex was found to be 9.52. This logK1 is consistent with the findings of the 0.1 M titrations as well as expectations for how stable the [Zn(8QP)] complex should be based on metal ion size (as compared to results using similar sized Ni2+ ion), affinity for nitrogen donors, and chelate ring size rule in reference to preferred M-N bond lengths. When compared to the binding strength of the ligand terpy, the 8QP:Zn2+ complex is somewhat higher than one would expect but is consistent with other metal-8QP complexes based on the chelate ring size rule. 64 Figure 33 : Plot of corrected absorbance (data points) and theoretical absorbance (solid line) versus pH for titration of 8QP and Zn(ClO4)2 for varying wavelengths a) 0.1 M at 10:1 Zn2+:8QP b) at 1.0 M at 1:1 Zn2+:8QP. 65 Figure 34 : UV spectra of 4.0 x 10-6 M 8QP and 4.0 x 10-3 M Zn(ClO4)2 in a 1.0 M (0.1 M HClO4 and 0.9 M NaClO4) solution. a) Initial spectrum of free 8QP at pH = 0.57. b) 8QP-Zn complex formed at around pH = 2.60 c) at pH = 4.06 d) Combination of spectra from a), b), and c) e) Full run spectra from pH = 0.57 – 7.07. 66 Table 13. Solutions for the pKa and absorbance parameters solved by the „SOLVER‟ module of EXCEL in determining logK1 of 8QP-Zn complex. 204 nm 223 nm 243 nm 291 nm 315 nm Abs(0) Abs(1) Abs(2) Abs(0) Abs(1) 0.6106 0.7402 1.2635 0.3235 0.3872 Abs(2) 0.5505 pK1 2.70 Abs(0) 0.2007 pK2 0.72 Abs(1) Abs(2) Abs(0) Abs(1) Abs(2) Abs(0) Abs(1) Abs(2) 0.3126 0.5200 0.1015 0.1760 0.3395 0.0918 0.1367 0.2026 Table 14. Summary of equations at each protonation equilibrium of Zn(8QP). pK [Zn(8QP)] + H+ H(8QP) H(8QP) + H+ H2(8QP) 67 + Zn2+ 2.70 0.72 8QP-Cadmium (II) Results Cd2+ has an ionic radius of 0.95 Å and is classified as a soft acid. The logK1 of the [Cd(8QP)] complex was expected to be quite low because of the large ionic radius, if a complex formed at all. The space-fill MM model (Figure 24) of the complex shows the difficulty Cd2+ will have complexing is due to ionic radius, not only due to the longer the ideal bond lengths associated with a large ion in a six-membered ring, but also due to the two hydrogen atoms which are blocking access to the cleft. Smaller metal ions have less trouble getting past these, but once the ionic radius gets larger than that of Zn2+ and Ni2+, the rigid backbone of 8QP must twist to accommodate and allow these metals in. Steric interference is the major factor which will likely keep larger ions like cadmium from forming strong complexes with 8QP. A solution of 4.5 x 10-6 M 8QP and 4.5 x 10-5 M Cd(ClO4)2 was titrated with 1.0 M NaOH from pH = 0.89 – 9.06. UV-vis absorbance spectra were recorded for the titration of Cd2+ with 8QP. Absorbances values were recorded for 204, 223, 243, 291, and 315 nm because these wavelengths exhibited the greatest amount of change over the course of the titration. Corrected absorbance and theoretical absorbance were plotted for every wavelength simultaneously against pH. The solutions of each parameter were found using the „SOLVER‟ function and the resulting pK values were computed. The pK values were recorded once all the standard deviations were minimized. Results from the wavelengths described above show apparent pK1 and pK2 values at 6.81 and 2.42 respectively. These were very close to the pKa values for the free ligand but the UV-vis spectra suggested that there may be a very weak complex forming. Another titration at 1000 to 1 ratio was run to determine that there was in fact a [Cd(8QP)] complex forming. A solution of 4.5 x 10-6 M 8QP and 4.5 x 10-3 M Cd(ClO4)2 was titrated with 1.0 M NaOH from pH = 0.59 to 8.00. UV-vis absorbance spectra are shown in 68 Figure 36 for the 1000 to 1 titration of Cd2+ with 8QP. Absorbances were again recorded at 204, 223, 243, 291, and 315 nm and corrected absorbance and theoretical absorbance were plotted for every wavelength simultaneously against pH and can be seen in Figure 35. Table 15 outlines the solutions of each parameter that was varied by the „SOLVER‟ function and the resulting solved pK values. The pK values were recorded once all the standard deviations were minimized and the resulting equations for the protonation equilibria observed during the 8QP-Ni titration can be seen in Table 16. Results from the wavelengths described above show apparent pK1 and pK2 values at 5.95 and 2.40 respectively. The titration at higher metal concentration confirmed the presence of a weak [Cd(8QP)] complex. The lower pKa remained unchanged but the upper value dropped from 6.82 with the free ligand to 5.95 in the presence of the cadmium complex. Using these pKa values and Eq. (21), the calculated logK1 value for the 8QP-Cd2+ complex was found to be 3.2. This is significantly lower than the logK value for the terpy complex which is exactly what one would expect of a larger metal ion based on the chelate ring size rule. 69 Figure 35 : Plot of corrected absorbance (data points) and theoretical absorbance (solid line) versus pH for titration of 8QP and at 1.0 M 1000:1 ratio Cd2+ : 8QP 70 Cd(ClO4)2 for varying wavelengths Figure 36 : UV spectra of 4.0 x 10-6 M 8QP and 4.0 x 10-3 M Cd(ClO4)2 in a 1.0 M (0.1 M HClO4 and 0.9 M NaClO4) solution. a) Initial spectrum of free 8QP at pH = 0.60. b) at pH = 2.00 c) at pH = 7.60 d) Combination of spectra from a), b), and c) e) Full run spectra from pH = 0.60 – 8.00 with the [Cd(8QP)] complex evident at high pH. 71 Table 15. Solutions for the pKa and absorbance parameters solved by the „SOLVER‟ module of EXCEL in determining logK1 of 8QP-Cd complex. 204 nm 223 nm 243 nm 291 nm 315 nm Abs(0) Abs(1) Abs(2) Abs(0) Abs(1) 0.8663 0.9417 0.9758 0.3282 0.4382 Abs(2) 0.5307 pK1 5.95 Abs(0) 0.2059 pK2 2.40 Abs(1) Abs(2) Abs(0) Abs(1) Abs(2) Abs(0) Abs(1) Abs(2) 0.2504 0.4417 0.1237 0.1713 0.2736 0.1024 0.1301 0.2009 Table 16. Summary of equations at each protonation equilibrium of Cd(8QP). pK [Cd(8QP)] + H+ H(8QP) H(8QP) + H+ H2(8QP) 72 + Cd2+ 5.95 2.40 8QP-Lead (II) Results Pb2+ has an ionic radius of 1.19 Å, which is much too large to form a complex according to predictions from HyperChem molecular mechanics models. Lead is categorized by Pearson as being an intermediate metal ion which means that it prefers to complex with intermediate donors, such as the nitrogen donors in 8QP, however its size should make this irrelevant. The UV absorbance spectra were recorded for the titration of Pb2+ with 8QP in a solution of 3.9 x 10-6 M 8QP and 3.9 x 10-6 M Pb(ClO4)2 from pH = 2.09 – 7.99. The data and corresponding measured and theoretical curves against pH indicate that the UV-vis spectra seen during the equal part Pb:8QP titration was simply that of the free ligand. 8QP-Thorium (IV) Results Th4+ has an ionic radius of 1.09 Å, which is also much too large to form a complex according to predictions from HyperChem molecular mechanics models. Thorium is categorized by Pearson as being a hard acid which means that it prefers to complex with hard donors, making formation of the Th4+-PDALD complex less likely. The UV absorbance spectra were recorded for the titration of Th4+ with 8QP in a solution of 3.9 x 10-6 M 8QP and 3.9 x 10-6 M Th(NO3)4 from pH = 2.00 – 7.52. The data and corresponding measured and theoretical curves against pH indicate that the UV-vis spectra seen during the equal part Th:8QP titration was simply that of the free ligand. 73 8QP-Calcium (II) Results Ca2+ has an ionic radius of 0.99 Å which is getting too large to form a complex according to predictions from HyperChem. Calcium is categorized by Pearson as being a soft acid which means that the nitrogen donors in 8QP will not help the Ca2+ ion overcome the difficulty binding due to its size. The UV absorbance spectra were recorded for the titration of Ca2+ with 8QP in a solution of 3.9 x 10-6 M 8QP and 3.9 x 10-6 M Ca(ClO4)2 from pH = 2.04 – 7.97. Absorbance values were recorded at 204, 223, 243, 291, and 315 nm because these wavelengths showed the greatest change in absorbance. Corrected absorbance and theoretical absorbance were plotted for every wavelength simultaneously against pH. As was expected, the UV-vis spectra indicated that no complex formed between calcium and 8QP at any range of pH. The data and corresponding measured and theoretical curves against pH also indicate that the UV-vis spectra seen during the equal part Ca:8QP titration was simply that of the free ligand. 74 Fluorescence Results for 8QP The concentration for the copper and 8QP were at a 1:1 ratio at 4.5 × 10-6 M (7.5 L of 0.0304 M stock copper solution in 50 mL stock 8QP solution). The pH was adjusted to 1.95 by adding 1.0 M NaOH. The concentration used for the Zn2+ complex was 10:1 with 8QP (112.5 L of 0.01 M stock zinc solution in 25 mL stock 8QP solution). The pH was adjusted to 7.3 by adding 1.0 M NaOH. The concentration used for the Ni2+ complex was 10:1 with 8QP (161 L of 0.0070 M stock nickel solution in 25 mL stock 8QP solution). The pH was adjusted to 4.0 by adding 1.0 M NaOH. The intensity was recorded for excitation wavelengths from 250-500 nm at 5 nm increments and emission wavelengths ranging from 365-800 nm at 5 nm increments. 8QP should take advantage of the CHEF effect due to the lone pair of electrons on the nitrogen and its extended aromatic backbone. The results were plotted using EXCEL and showed that the zinc and nickel complexes both had strong fluorescence (zinc being significantly stronger), and the copper complex did not fluoresce at all (Figure 37). None of the metal-ligand complexes showed higher intensity emission than the 8QP free ligand scan. This was surprising and may be attributed to varying the pH values of the complex solutions. This was done with each solution to ensure that the complex was present during the scan, but may have led to less reliable results. Future studies may choose to include a series of scans of the metal-ligand complexes over a broad range of pH values to determine if any chelation enhanced fluorescence occurs at higher or lower pH than was included in this study. The results of the fluorescence experiments carried out in this study indicate that 8QP may not be ideal for use as a biological fluorescent indicator. The [Cu(8QP)] complex has a high logK1 value, but the complex quenches any fluorescence. The logK1 values of the Zn2+ and Ni2+ are also relatively high, at 10.0 and 9.5 respectively. However, the proximity of the logK1 values 75 of these complexes would prevent 8QP from being useful as an indicator because it would be impossible to differentiate between the fluorescence due to a Zn(8QP) complex and a Ni(8QP) complex. Figure 37 : Plot of emission v. wavelength for the 8QP free ligand (green), Zn(8QP) complex (blue), Ni(8QP) complex (pink), and Cu(8QP) complex (maroon). 76 Titrations Involving PDALD PDALD-Copper (II) Results Compared to other small metal ions, the logK1 value for the [Cu2+(PDALD)] complex is higher than one would expect considering the size of Cu2+ (r+=0.57 Å) and the complex formation of a five-membered chelate ring. Although [Cu2+(PDALD)] crystal structures could not be cultured, similarities in structure and binding trends with PDALC may indicate that the complex would include two PDALD molecules, each coordinating two nitrogens and one oxygen.33 The formation of distorted square bipyramidal coordination geometry around the Cu2+ cation using two PDALD molecules may prevent complex dissociation and account for a larger logK1 than would be expected. When comparing PDALD and PDALC to phen, one finds that the logK1 values for the Cu2+ complexes of the more highly substituted ligands (6.37, 7.56) are significantly lower than the Cu2+:phen complex (9.13). Smaller metal ion affinity decreases with the addition of neutral oxygen donors, as seen in PDALD with the addition of the carboxyaldehyde groups to the phen moiety. This accounts for the decrease in logK1 for Cu2+ from phen to PDALD by nearly three log units. The coordinating strength of alkoxide groups tend to be ruled by the affinity of the metal ion for the hydroxide ion.34 The logK1(OH-) of the Cu2+ cation is 6.3, which explains why the PDALC complex is stronger than the [Cu2+(PDALD)] complex. The negative oxygen donors in the hydroxymethyl groups of PDALC make it more polar than the carboxyaldehyde of the PDALD, which also contributes to greater PDALC affinity in Cu2+ complexes. This is a trend seen when comparing logKML values of PDALD and PDALC complexes with nearly every metal ion investigated. 77 A solution of 2.0 x 10-5 M PDALD and 2.0 x 10-5 M Cu(ClO4)2 was titrated with 0.1 M NaOH from pH = 2.15 – 10.90. UV-vis absorbance spectra are shown in Figure 39 for the titration of Cu2+ with PDALD. Corrected absorbance and theoretical absorbance were plotted for every wavelength simultaneously against pH and can be seen in Figure 38. Table 17 outlines the solutions of each parameter that was varied by the „SOLVER‟ function and the resulting solved pK values, with Table 28 showing the resulting equations for the protonation equilibria. Apparent pK1 and pK2 values were observed at 7.27 and 0.83, respectively. Using these pKa values and Eq. (21), the calculated logK1 value for the Cu2+-PDALD complex is 6.37. Figure 38 : Plot of corrected absorbance (data points) and theoretical absorbance (solid line) versus pH for titration of PDALD and Cu(ClO4)2 for varying wavelengths 78 Figure 39 : Major differences between the spectra of the free ligand and the spectra where the Cu2+ complex is present. The peak at around 280 nm is broader when there is a complex present. The two peaks between 210 and 230 nm are elevated and equal in size when the complex is present. UV spectra of 2.0 x 10-5 M PDALD and 2.0 x 10-5 M Cu(ClO4)2 in a 0.1 M (0.01 M HClO4 and 0.09 M NaClO4) solution. a) Initial spectrum of the [Cu(PDALD)] complex at pH = 2.15. b) Initial spectrum of free PDALD at pH = 2.15. c) [Cu(PDALD)] complex at pH = 2.15 – 7.37. d) Full run spectra from pH = 2.15 – 7.50 of the PDALD free ligand. 79 Table 17. Solutions for the pKa and absorbance parameters solved by the „SOLVER‟ module of EXCEL in determining logK1 of PDALD-Cu complex. 225 nm 241 nm 256 nm 270 nm 282 nm Abs(0) Abs(1) Abs(2) Abs(0) Abs(1) 0.6454 0.6546 0.6147 0.4899 0.4510 Abs(2) 0.4094 pK1 7.27 Abs(0) 0.2301 pK2 0.83 Abs(1) Abs(2) Abs(0) Abs(1) Abs(2) Abs(0) Abs(1) Abs(2) 0.2506 0.2396 0.2973 0.3595 0.3721 0.3866 0.4134 0.4246 Table 18. Summary of equations at each protonation equilibrium of Cu(PDALD). pK [Cu(PDALD)] + H+ [Cu(PDALD)H] [Cu(PDALD)H] + H+ H2(PDALD) 80 7.27 + Cu2+ 0.83 PDALD-Cadmium (II) Results Values for the [Cd2+(PDALD)] complex were determined to be slightly below the reported PDALC:Cd2+ value (7.49) and slightly higher than the phen:Cd2+ value (5.66). The PDALD complex was stronger than the phen complex because larger metal ions have an affinity toward neutral oxygen donors.34 A moderate logK1(OH-) of 3.9 and greater polarity of the negative oxygen donors explain the slightly stronger PDALC complex. A solution of 2.0 x 10-5 M PDALD and 2.0 x 10-5 M Cd(ClO4)2 was titrated with 0.1 M NaOH from pH = 2.15 – 8.82. UV-vis absorbance spectra are shown in Figure 41 for the titration of Cd2+ with PDALD. Corrected absorbance and theoretical absorbance were plotted for every wavelength simultaneously against pH and can be seen in Figure 40. Table 19 outlines the solutions of each parameter that was varied by the „SOLVER‟ function and the resulting solved pK values, with Table 20 showing the resulting equations for the protonation equilibria. Apparent pK1 and pK2 values were observed at 7.38 and 0.89, respectively. Using these pKa values and Eq. (21), the calculated logK1 value for the Cd2+-PDALD complex is 6.31. Figure 40 : Plot of corrected absorbance (data points) and theoretical absorbance (solid line) versus pH for titration of PDALD and Cd(ClO4)2 for varying wavelengths 81 Figure 41 : UV spectra of 2.0 x 10-5 M PDALD and 2.0 x 10-5 M Cd(ClO4)2 in a 0.1 M (0.01 M HClO4 and 0.09 M NaClO4) solution. a) Initial Cd2+ and PDALD spectrum at pH = 2.15. b) at pH = 5.19 c) at pH = 7.18 d) Combination of spectra from a), b), and c) e) Full run spectra from pH = 2.15 – 8.82 for [Cd(PDALD)] titration. f) Full run spectra from pH = 2.15 – 7.50 of the PDALD free ligand. 82 Table 19. Solutions for the pKa and absorbance parameters solved by the „SOLVER‟ module of EXCEL in determining logK1 of PDALD-Cd complex. 225 nm 241 nm 256 nm 270 nm 282 nm Abs(0) Abs(1) Abs(2) Abs(0) Abs(1) 0.5597 0.6254 0.5722 0.4762 0.4391 Abs(2) 0.3875 pK1 7.38 Abs(0) 0.1809 pK2 0.89 Abs(1) Abs(2) Abs(0) Abs(1) Abs(2) Abs(0) Abs(1) Abs(2) 0.2130 0.1981 0.3080 0.3818 0.3857 0.3863 0.4231 0.4152 Table 20. Summary of equations at each protonation equilibrium of Cd(PDALD). pK [Cd(PDALD)] + H+ [Cd(PDALD)H] [Cd(PDALD)H] + H+ H2(PDALD) 83 7.38 + Cd2+ 0.89 PDALD-Calcium (II) Results The [Ca2+(PDALD)] complex is of particular interest because Ca2+ was the only metal ion investigated with a higher logK1 value than the analogous PDALC (3.74) and phen (1.00) complexes. Calcium has very low electronegativity, is classified as a hard metal ion, and prefers to bind with harder donor atoms. Phen has a logK1 of 1.00 with Ca2+ because it only has intermediate nitrogen donors, which are not ideal. The addition of highly electronegative, hard, oxygen donors in PDALD and PDALC, accounts for increased logK1 values with Ca2+. Selectivity for the hydroxide ion has been shown to correlate with negative oxygen donor affinity.34 PDALC complex stability is significantly lower (ΔlogK1= -2.00) than that of PDALD because of the very low acidity of the calcium ion, logK1(OH-) = 1.3. Also important, affinity for the neutral oxygen donors of the carboxyaldehyde groups of the PDALD is increased due to the large ionic radius (r+= 0.99 Å) of the Ca2+ ion. A solution of 2.0 x 10-5 M PDALD and 2.0 x 10-5 M Ca(ClO4)2 was titrated with 0.1 M NaOH from pH = 2.19 – 8.65. UV-vis absorbance spectra are shown in Figure 43 for the titration of Ca2+ with PDALD. Corrected absorbance and theoretical absorbance were plotted for every wavelength simultaneously against pH and can be seen in Figure 42. Table 21 outlines the solutions of each parameter and the resulting solved pK values, while Table 22 shows the resulting equations for the protonation equilibria. Apparent pK1 and pK2 values were determined at 7.72 and 1.45, respectively. Using these pKa values and Eq. (21), the calculated logK1 value for the Ca2+-PDALD complex is 5.74. 84 Figure 42 : Plot of corrected absorbance (data points) and theoretical absorbance (solid line) versus pH for titration of PDALD and Ca(ClO4)2 for varying wavelengths 85 Figure 43 : UV spectra of 2.0 x 10-5 M PDALD and 2.0 x 10-5 M Ca(ClO4)2 in a 0.1 M (0.01 M HClO4 and 0.09 M NaClO4) solution. a) Initial Ca2+ and PDALD spectrum at pH = 2.19. b) at pH = 3.50 c) at pH = 7.45 d) Combination of spectra from a), b), and c) e) Full run spectra from pH = 2.19 – 7.45 for [Ca(PDALD)] titration. f) Full run spectra from pH = 2.15 – 7.50 of the PDALD free ligand. 86 Table 21. Solutions for the pKa and absorbance parameters solved by the „SOLVER‟ module of EXCEL in determining logK1 of PDALD-Ca complex. 225 nm 241 nm 256 nm 270 nm 282 nm Abs(0) Abs(1) Abs(2) Abs(0) Abs(1) 0.4785 0.5781 0.5605 0.4282 0.4150 Abs(2) 0.3846 pK1 7.72 Abs(0) 0.1534 pK2 1.45 Abs(1) Abs(2) Abs(0) Abs(1) Abs(2) Abs(0) Abs(1) Abs(2) 0.2022 0.2008 0.2660 0.3819 0.3881 0.3663 0.4229 0.4171 Table 22. Summary of equations at each protonation equilibrium of Ca(PDALD). pK [Ca(PDALD)] + H+ [Ca(PDALD)H] [Ca(PDALD)H] + H+ H2(PDALD) 87 7.72 + Ca2+ 1.45 PDALD-Gadolinium (III) Results The value for the [Gd3+(PDALD)] complex is closer to the analogous PDALC complex (6.16) than any other metal ion included in this study with a ΔlogK1= 0.21. Gadolinium, being a larger (r+= 0.97 Å), hard metal ion of intermediate acidity, is selective for neutral oxygen donors. The logK1(OH-) of 6.1 dictates a greater affinity towards the more polar oxygen donors of PDALC. The predicted Gd3+-phen complex33 logK1 of 0.9 shows how large of an effect the addition of hard oxygen donors has in both PDALD and PDALC complexes. A solution of 2.0 x 10-5 M PDALD and 2.0 x 10-5 M Gd(ClO4)3 was titrated with 0.1 M NaOH from pH = 2.18 – 8.43. UV-vis absorbance spectra are shown in Figure 45 for the titration of Gd3+ with PDALD. Corrected absorbance and theoretical absorbance were plotted for every wavelength simultaneously against pH and can be seen in Figure 44. Table 23 outlines the solutions of each parameter and the resulting solved pK values, and Table 24 gives the resulting equations for the protonation equilibria. Apparent pK1 and pK2 values were determined at 6.42 and 1.25, respectively. Using these pKa values and Eq. (21), the calculated logK1 value for the Gd3+-PDALD complex is 5.95. 88 Figure 44 : Plot of corrected absorbance (data points) and theoretical absorbance (solid line) versus pH for titration of PDALD and Gd(ClO4)3 for varying wavelengths 89 Figure 45 : UV spectra of 2.0 x 10-5 M PDALD and 2.0 x 10-5 M Gd(ClO4)3 in a 0.1 M (0.01 M HClO4 and 0.09 M NaClO4) solution. a) Initial Gd3+ and PDALD spectrum at pH = 2.18. b) at pH = 6.18 c) at pH = 7.26 d) Combination of spectra from a), b), and c) e) Full run spectra from pH = 2.18 – 7.85 for [Gd(PDALD)] titration. f) Full run spectra from pH = 2.15 – 7.50 of the PDALD free ligand. 90 Table 23. Solutions for the pKa and absorbance parameters solved by the „SOLVER‟ module of EXCEL in determining logK1 of PDALD-Gd complex. 225 nm 241 nm 256 nm 270 nm 282 nm Abs(0) Abs(1) Abs(2) Abs(0) Abs(1) 0.4953 0.5085 0.5182 0.4248 0.4125 Abs(2) 0.4012 pK1 6.42 Abs(0) 0.1640 pK2 1.25 Abs(1) Abs(2) Abs(0) Abs(1) Abs(2) Abs(0) Abs(1) Abs(2) 0.1750 0.1739 0.2354 0.3027 0.3214 0.3277 0.3973 0.4386 Table 24. Summary of equations at each protonation equilibrium of Gd(PDALD). pK [Gd(PDALD)] + H+ [Gd(PDALD)H] [Gd(PDALD)H] + H+ H2(PDALD) 91 6.42 + Gd3+ 1.25 PDALD-Lead (II) Results Pb2+ had the highest logK1 value seen in the PDALD study and can be attributed to the large ionic radius (1.19Å). Pearson classifies lead as an intermediate metal ion, so the addition of oxygen donors to phen is not necessarily the best way to increase selectivity. The addition of the neutral oxygens of the carboxyaldehyde groups in PDALD made the ligand more selective for Pb2+. An intermediate logK1(OH-) of 6.3 made PDALC slightly more selective (ΔlogK1 = +0.96). A solution of 2.0 x 10-5 M PDALD and 2.0 x 10-5 M Pb(ClO4)2 was titrated with 0.1 M NaOH from pH = 2.19 – 8.67. UV-vis absorbance spectra are shown in Figure 47 for the titration of Pb2+ with PDALD. Corrected absorbance and theoretical absorbance values were plotted simultaneously and can be seen in Figure 46. Table 25 outlines the solutions of each parameter that was varied by the „SOLVER‟ function and the resulting solved pK values. Table 26 shows the resulting equations for the protonation equilibria. Apparent pK1 and pK2 values were determined at 6.77 and 0.87, respectively. Using these pKa values and Eq. (21), the calculated logK1 value for the Pb2+-PDALD complex is 6.38. 92 Figure 46 : Plot of corrected absorbance (data points) and theoretical absorbance (solid line) versus pH for titration of PDALD and Pb(ClO4)2 for varying wavelengths 93 Figure 47 : UV spectra of 2.0 x 10-5 M PDALD and 2.0 x 10-5 M Pb(ClO4)2 in a 0.1 M (0.01 M HClO4 and 0.09 M NaClO4) solution. a) Initial Pb2+ and PDALD spectrum at pH = 2.19. b) at pH = 3.78 c) at pH = 7.37 d) Combination of spectra from a), b), and c) e) Full run spectra from pH = 2.19 – 7.37 for [Pb(PDALD)] titration. f) Full run spectra from pH = 2.15 – 7.50 of the PDALD free ligand. 94 Table 25. Solutions for the pKa and absorbance parameters solved by the „SOLVER‟ module of EXCEL in determining logK1 of PDALD-Pb complex. 225 nm 241 nm 256 nm 270 nm 282 nm Abs(0) Abs(1) Abs(2) Abs(0) Abs(1) 0.4652 0.5562 0.5618 0.3828 0.4024 Abs(2) 0.3934 pK1 6.77 Abs(0) 0.1364 pK2 0.82 Abs(1) Abs(2) Abs(0) Abs(1) Abs(2) Abs(0) Abs(1) Abs(2) 0.1706 0.1698 0.1959 0.2950 0.3039 0.3022 0.3973 0.4146 Table 26. Summary of equations at each protonation equilibrium of Pb(PDALD). pK [Pb(PDALD)] + H+ [Pb(PDALD)H] [Pb(PDALD)H] + H+ H2(PDALD) 95 6.77 + Pb2+ 0.82 PDALD-Thorium (IV) Results Thorium has an ionic radius of 1.09 Å, is classified as a hard acid, and is the most highly charged metal ion studied. This value was slightly lower than would be expected base on ionic radius alone. With a ΔlogK1 = +2.22, it had the greatest disparity between analogous PDALD (4.98) and PDALC (7.20) complexes. This is due to the acidic nature of the Th4+ ion with a logK1(OH-) of 10.8. The affect of the introduction of negative oxygen donors in PDALC was far greater than the addition of neutral oxygen donors in PDALD. Affinity for larger metal ions like the Th4+ ion, as well as the addition of hard donors caused an increase of approximately one log unit from phen to PDALD. A solution of 2.0 x 10-5 M PDALD and 2.0 x 10-5 M Th(NO4)3 was titrated with 0.1 M NaOH from pH = 2.16 – 9.35. UV-vis absorbance spectra are shown in Figure 49 for the titration of Th4+ with PDALD. Corrected absorbance and theoretical absorbance were plotted against pH and can be seen in Figure 48. Table 27 outlines the solutions of each parameter varied by the „SOLVER‟ function and the resulting solved pK values. Results from absorbance values at 225 and 241 nm were omitted from the final pKa calculations because of the broad band seen at lower wavelengths in Figure 49e caused by the nitrate. The pK values were reported in Table 28. Apparent pK1 and pK2 values were determined at 6.11 and 2.22, respectively. Using these pKa values and Eq. (21), the calculated logK1 value for the Th4+PDALD complex is 4.98. 96 Figure 48 : Plot of corrected absorbance (data points) and theoretical absorbance (solid line) versus pH for titration of PDALD and Th(NO3)4 for varying wavelengths 97 Figure 49 : UV spectra of 2.0 x 10-5 M PDALD and 2.0 x 10-5 M Th(NO3)4 in a 0.1 M (0.01 M HClO4 and 0.09 M NaClO4) solution. a) Initial Th4+ and PDALD spectrum at pH = 2.16. b) at pH = 3.94 c) at pH = 7.89 d) Combination of spectra from a), b), and c) e) Full run spectra from pH = 2.16 – 9.32 for [Th(PDALD)] titration. f) Full run spectra from pH = 2.15 – 7.50 of the PDALD free ligand. 98 Table 27. Solutions for the pKa and absorbance parameters solved by the „SOLVER‟ module of EXCEL in determining logK1 of PDALD-Th complex. 225 nm 241 nm 256 nm 270 nm 282 nm Abs(0) Abs(1) Abs(2) Abs(0) Abs(1) 0.5440 0.5282 0.6048 0.3001 0.2871 Abs(2) 0.3548 pK1 6.11 Abs(0) 0.1399 pK2 2.22 Abs(1) Abs(2) Abs(0) Abs(1) Abs(2) Abs(0) Abs(1) Abs(2) 0.1556 0.1635 0.1716 0.1910 0.2419 0.2139 0.2407 0.3289 Table 28. Summary of equations at each protonation equilibrium of Th(PDALD). pK [Th(PDALD)] + H+ [Th(PDALD)H] [Th(PDALD)H] + H+ H2(PDALD) 99 6.11 + Th4+ 2.22 PDALD-Zinc (II) Results Zn2+ has an ionic radius of 0.74 Å, which makes it one of the smaller metal ions investigated in this study, and is less than the ideal size for complexation with PDALD. Zinc is categorized by Pearson as being an intermediate acid which means that it prefers to complex with intermediate donors, so the addition of hard oxygen donors was not expected to help with Zn2+ specificity. What was seen in PDALD was that the addition of the neutral oxygens significantly destabilized the Zn2+ complex when compared with phen and PDALC. PM3 semiempirical calculation of the [Zn2+(PDALC)] complex33 indicates that upon formation of the complex, both PDALC nitrogen donors are coordinated to the metal ion but only one of the negative oxygen donors is coordinated. Three additional water molecules are coordinated, one of which is stabilized by a hydrogen bond to the uncoordinated oxygen of PDALC. Molecular mechanics calculations involving PDALD indicate that coordination of either neutral oxygen to the zinc ion is energetically unfavorable. The addition of the carboxyaldehyde groups further destabilizes the complex by preventing proper spacial arrangement of the Zn2+ ion within the PDALD binding cleft while increasing steric interference, without being able to compensate with additional coordination to the metal ion. A solution of 2.0 x 10-5 M PDALD and 2.0 x 10-5 M Zn(ClO4)2 was titrated with 0.1 M NaOH from pH = 2.17 – 11.20. UV-vis absorbance spectra are shown in Figure 51 for the titration of Zn2+ with PDALD. Corrected absorbance and theoretical absorbance were plotted against pH and can be seen in Figure 50. Table 29 outlines the solutions of each parameter and resulting solved pK values, while Table 30 shows the resulting equations for the protonation equilibria. Apparent pK1 and pK2 values were determined at 6.96 and 2.44, respectively. Using these pKa values and Eq. (21), the calculated logK1 value for the Zn2+-PDALD complex is 4.76. 100 Figure 50 : Plot of corrected absorbance (data points) and theoretical absorbance (solid line) versus pH for titration of PDALD and Zn(ClO4)2 for varying wavelengths 101 Figure 51 : UV spectra of 2.0 x 10-5 M PDALD and 2.0 x 10-5 M Zn(ClO4)2 in a 0.1 M (0.01 M HClO4 and 0.09 M NaClO4) solution. a) Initial Zn2+ and PDALD spectrum at pH = 2.17. b) at pH = 3.88 c) at pH = 7.35 d) Combination of spectra from a), b), and c) e) Full run spectra from pH = 2.17 – 7.59 for [Zn(PDALD)] titration. f) Full run spectra from pH = 2.15 – 7.50 of the PDALD free ligand. 102 Table 29. Solutions for the pKa and absorbance parameters solved by the „SOLVER‟ module of EXCEL in determining logK1 of PDALD-Zn complex. 225 nm 241 nm 256 nm 270 nm 282 nm Abs(0) Abs(1) Abs(2) Abs(0) Abs(1) 0.7327 0.5661 0.4847 0.5481 0.5363 Abs(2) 0.5062 pK1 6.96 Abs(0) 0.1780 pK2 2.44 Abs(1) Abs(2) Abs(0) Abs(1) Abs(2) Abs(0) Abs(1) Abs(2) 0.1910 0.1848 0.2437 0.2963 0.3021 0.3680 0.3959 0.3963 Table 30. Summary of equations at each protonation equilibrium of Zn(PDALD). pK [Zn(PDALD)] + H+ H(PDALD) H(PDALD) + H+ H2(PDALD) 103 + Zn2+ 6.96 2.44 PDALD-Uranyl (II) Results The uranyl ion (UO22+) is the most common species encountered in the aqueous chemistry of uranium. Uranium is categorized by Pearson as being a hard acid which means that it prefers to complex with hard bases so addition of the hard oxygen donors should be beneficial. The neutral oxygen donors of PDALD were expected to get a logK 1 close to that of the analogous PDALC complex. In actuality, the logKML for PDALD and PDALC indicate that the UO22+ ion acidity (logK1(OH-) = 8.2) affected the specificity of the negative oxygen donors of PDALC much more than the large size of the cation affected affinity for the neutral oxygen donors of PDALD. A solution of 2.0 x 10-5 M PDALD and 2.0 x 10-5 M UO2(NO3)2 was titrated with 0.1 M NaOH from pH = 2.19 – 8.95. UV-vis absorbance spectra are shown in Figure 53 for the titration of UO22+ with PDALD. Corrected absorbance and theoretical absorbance were plotted against pH and can be seen in Figure 52. Table 31 outlines the solutions of each parameter and the resulting solved pK values. Results from absorbance values at 225 and 241 nm were omitted from the final pKa calculations because of the broad band seen at lower wavelengths in Fig. 53e caused by the nitrate. The pK values were recorded once all the standard deviations were minimized and the resulting equations for the protonation equilibria observed during the UO22+PDALD titration can be seen in Table 34. Apparent pK1 and pK2 values were determined at 4.80 and 2.34, respectively. Using these pKa values and Eq. (21), the calculated logK1 value for the UO22+-PDALD complex is 4.86. 104 Figure 52 : Plot of corrected absorbance (data points) and theoretical absorbance (solid line) versus pH for titration of PDALD and UO2(NO3)2 for varying wavelengths. 105 Figure 53 : UV spectra of 2.0 x 10-5 M PDALD and 2.0 x 10-5 M UO2(NO3)2 in a 0.1 M (0.01 M HClO4 and 0.09 M NaClO4) solution. a) Initial UO22+ and PDALD spectrum at pH = 2.19. b) at pH = 3.64 c) at pH = 7.63 d) Combination of spectra from a), b), and c) e) Full run spectra from pH = 2.19 – 7.63 for [UO22+(PDALD)] titration. f) Full run spectra from pH = 2.15 – 7.50 of the PDALD free ligand. 106 Table 31. Solutions for the pKa and absorbance parameters solved by the „SOLVER‟ module of EXCEL in determining logK1 of PDALD- UO2 complex. 225 nm 241 nm 256 nm 270 nm 282 nm Abs(0) Abs(1) Abs(2) Abs(0) Abs(1) 0.9094 0.7437 0.7144 0.8388 0.6553 Abs(2) 0.6281 pK1 4.80 Abs(0) 0.2314 pK2 2.34 Abs(1) Abs(2) Abs(0) Abs(1) Abs(2) Abs(0) Abs(1) Abs(2) 0.2041 0.1834 0.2781 0.2697 0.2867 0.3315 0.3191 0.4213 Table 32. Summary of equations at each protonation equilibrium of UO2(PDALD). pK [UO2(PDALD)] + H+ [UO2(PDALD)H] [UO2(PDALD)H] + H+ H2(PDALD) 107 4.80 + UO22+ 2.34 X-Ray Crystallography Results for [Cu(8QP)2](ClO4)2 Synthesis of [Cu(8QP)2](ClO4)2 used 1.16 x 10-2g 8QP dissolved in n-butanol (25 mL), and one equivalent of Cu(ClO4)2.6H2O in DI water (495 µL 0.0304 M stock Cu(ClO4)2 solution in 10 mL DI water). The metal perchlorates were obtained from Aldrich or Alfa Aesar in 99% purity or better and used as received. All solutions were made up in deionized water (Milli-Q, Waters Corp.) of > 18 MΩ.cm-1 resistivity. The ligand solution was floated on the aqueous solution in a 50 mL beaker covered with parafilm, with a few small holes punched in the parafilm to allow for slow evaporation of the solvent. After a few days, dark green crystals formed and were filtered off. Calculated percentages for C42H26Cl2CuN6O8 were: C, 57.38%; H, 2.98%; N, 9.68%. The results of CHN analysis found: C, 57.26%; H, 2.75%; N, 9.47%. Molecular Structure Determination A Bruker SMART 1K diffractometer using the omega scan mode was employed for crystal screening, unit cell determination, and data collection at 110(2) K. The structure of [Cu(8QP)2](ClO4)2 was solved by direct methods, and refined to convergence.35 Absorption corrections were made using the SADABS program.36 All hydrogens were located in difference Fourier maps (including those at ideal positions). Some details of the structure determination are given in Table 33, and crystal coordinates and details of the structure determinations of [Cu(8QP)2](ClO4)2 have been deposited with the CSD (Cambridge Structural Database).37 A selection of bond lengths and angles for the copper complex are given in Table 34, and the structure of [Cu(8QP)2](ClO4)2 is shown in Figure 54. 108 Figure 54 : Structure of the complex cation [Cu(8QP)2]2+, showing the numbering scheme for atoms relevant to discussion of the coordination sphere around the copper. Drawing made with ORTEP38. Thermal ellipsoids drawn at the 50% level. Data obtained from the crystallographic analysis of the [Cu(8QP)2](ClO4)2 structure indicates that there are two coordinated 8QP molecules for each Cu2+ ion, with only five of six nitrogens covalently linked. This can be attributed to the steric interactions between the two ligand molecules as well as an inability for 8QP to accommodate for ideal M-L bond length due to the size of the cleft. Even when coordinating a small metal ion such as Cu2+, there is a significant amount of angle strain on the bond between the phenanthroline and the quinolyl 109 groups. This has been noted in prior studies involving the [(8QP)Pt2+] complex29 and was predicted using HyperChem molecular modeling. Figure 55 : Structure of [Pt(8QP)Cl]+ determined by Hu et al.29 illustrating the non-planarity due to angle strain characteristic of 8QP complexes. Ideal M-L bond length for a six-membered ring is approximately 1.6 Å and closer to 2.5 Å for five-membered rings. Bond lengths reported for the complex cation [Cu(8QP)2]2+ ranged from 1.936 Å to 2.306 Å. The only six-membered M-8QP(1) ring to form in the Cu2+ complex (Cu(1)-N(3b) and Cu(1)-N(2b)) has bond lengths of 2.121 Å and 1.936 Å, respectively. [Cu(1)N(3b)] is the longer of the two and is the bond between the copper ion and the nitrogen of the only coordinating quinolyl group of the complex. Upon formation of the complex, the quinolyl group is pushed out of the preferred planar conformation in order to minimize bond energy. The rigid backbone of the phenanthroline group assures that the two associated nitrogen atoms (N1b and N2b) remain in a single plane. The second 8QP molecule (2) only coordinates to the copper ion with the two nitrogens of the phenanthroline. These bond lengths (Cu(1)-N(1a) and Cu(1)-N(2a)) are 1.964 Å and 2.306 Å and form a five-membered chelate ring. Steric interactions between the quinolyl group of 8QP(2) and the phenanthroline of 8QP(1) prevent bonding of the final nitrogen. 110 Table 33. Crystal data and structure refinement for [Cu(8QP)2](ClO4)2. _______________________________________________________________________________ Empirical formula Formula weight Temperature Crystal system Space group Unit cell dimensions a b C42 H26 Cl2 Cu N6 O8 877.13 63(2) K Triclinic Pī c 13.40(2) Å Volume Z Reflections collected Independent reflections Absorption correction 112.668(19)° 108.73(2)° 98.17(2)° 1791(6) Å3 2 16360 5676 [R(int) = 0.2206] Semi-empirical from equivalents 12.36(2) Å 12.99(2) Å Final R indices [I>2sigma(I)] R1 = 0.0897, wR2 = 0.1504 R indices (all data) R1 = 0.2132, wR2 = 0.1832 ___________________________________________________________________ 111 Table 34. Bond lengths and angles of interest in the complex cation [Cu(8QP)2]2+ from Fig. 56. ________________________________________________________________________ Bond lengths (Å): Cu(1)-N(1a) 1.964(8) Cu(1)-N(1b) 2.079(7) Cu(1)-N(2a) Cu(1)-N(2b) 2.306(9) 1.936(8) Cu(1)-N(3b) 2.121(8) Bond angles (deg): N(2b)-Cu(1)-N(1a) N(1a)-Cu(1)-N(1b) 172.8(3) 90.8(3) N(2b)-Cu(1)-N(1b) N(2b)-Cu(1)-N(3b) 82.4(3) 90.1(3) N(1a)-Cu(1)-N(3b) 95.7(3) N(1b)-Cu(1)-N(3b) 161.1(3) N(2b)-Cu(1)-N(2a) N(1b)-Cu(1)-N(2a) C(12a)-N(1a)-Cu(1) C(12b)-N(1b)-Cu(1) C(10b)-N(2b)-Cu(1) 108.7(3) 111.7(3) 119.7(5) 108.9(5) 126.1(6) N(1a)-Cu(1)-N(2a) N(3b)-Cu(1)-N(2a) C(11a)-N(2a)-Cu(1) C(11b)-N(2b)-Cu(1) C(21b)-N(3b)-Cu(1) 76.0(3) 87.2(3) 108.9(5) 114.8(5) 119.9(6) ________________________________________________________________________ 112 CONCLUSIONS Studies involving cryptands and macrocyclic ligands have shown very high ligand-metal ion complex stability due to the limited number of conformations available to the free ligand. Straight chain analogues of these ligands form significantly less stable complexes, however. Introducing structures which limit unfavorable conformations increases the stability of the complex dramatically. By limiting the movement of the free ligand, the structural characteristics of the binding site have a larger impact on metal ion size-based selectivity. The structural rigidity of the aromatic backbone of both 8QP and PDALD makes the complex stability of these ligands size selective. Pearson‟s HASB theory and coordination geometry have also proven to be influential in the binding properties of these ligands. 8QP was designed as a terpy analog which has been modified to increase structural rigidity and contain both a five- and six-membered chelate ring upon complexation rather than two five-membered rings. It was hypothesized that due to the presence of the six-membered ring, 8QP would have more affinity for smaller sized metal ions than terpy. Table 35 below compares the formation constants of 8QP and terpy complexes. As expected according to the chelate ring size rule, 8QP shows an increased affinity for smaller metal ions and a decrease for larger metal ions when compared to analogous terpy complexes. Complex formation only occurred in metal ions with an ionic radius of 0.95 Å or smaller. This study has shown that the size selectivity of 8QP was driven primarily by the sixmembered ring as the calculated strain energy of the M-N bonds decreased with ionic radius. The size selectivity of 8QP was also heavily influenced by the size and sterics of the cleft, as seen in Figure 56. The HyperChem models and X-Ray crystallographic results described in this study show that even mid-sized metal ions such as Cu2+ were too large to fit properly within the 113 binding site. This caused a rotation of the bond between the phenanthroline and the quinolyl groups. This non-planar configuration is energetically unfavorable and hindered complex stability and complex formation for larger metal ions with 8QP. Table 35. Comparison of protonation constants and logK1 values with a selection of metal ions between 8QP (one six-membered and one five-membered chelate ring) and terpy (two fivemembered chelate rings). ________________________________________________________________________ Metal Ion ionic radius (Å) logK1(8QP) logK1(terpy) 24 ________________________________________________________________________ H+ - 6.8 4.7 Cu2+ 0.57 14.4 12.3 Ni2+ 0.69 10.0 10.7 Zn2+ 0.74 9.5 6.0 Cd2+ 0.96 3.2 5.1 _________________________________________________________________________ Figure 56 : The binding cleft of 8QP is significantly more crowded than that of terpy due to the creation of the six-membered ring. The presence of the bulkier quinolyl group as well as the labeled hydrogen atoms acts to make 8QP even more selective for small metal ion complexes. 114 The M-L complex stability results obtained from this study correspond with predictions made based on Pearson‟s HASB theory. 8QP had three intermediate nitrogen donors which formed complexes with all the intermediate acids studied except Pb2+ and Fe3+. Lead (1.19Å) was much too large to complex while the Fe3+ ion was likely unable to complex because it hydrolyzed in solution. A complex was formed with the Pearson soft acid Cd2+. It is not apparent whether the relative weakness of these complexes is attributed to the higher ionic radius of the metal ions or the softness of the acid. Ca2+ and the actinide Th4+ were the only hard acids tested in this study and as expected from their size, neither formed a complex. Although HASB theory fit with the results, the overall influence of these properties was likely minimal when compared with the effects of ionic radius. Fluorescence studies on 8QP revealed that upon complexation, fluorescent emission values were significantly lower than the initial free ligand emission. A wavelength shift occurred in the emission spectra between the free ligand scan (~450 nm) and subsequent M-L complex scans (~650 nm). There were visibly distinct peaks for the [Zn2+(8QP)] and [Ni2+(8QP)] complexes, while fluorescence was completely quenched in the [Cu2+(8QP)] complex. Potential uses for 8QP as a fluorescent marker in biological applications are limited. However, there may be possibilities for it in other applications where trace Zn2+ and Ni2+ are not both present. Future use of 8QP as a quantitative biological fluorescent tag is unlikely due to the relatively similar logK values of Zn2+ and Ni2+. Distinguishing between the relative fluorescence contributions of zinc against nickel emission would be impossible. PDALD is a rather non-discriminating ligand in that it complexed nearly every metal it was tested with, although none of these complexes yielded high logK1 values. Relatively open access to the binding site of this ligand makes it easy for metal ions to complex. Table 36 115 compares the logK values of PDALD and two similar ligands, 1,10-phenanthroline (phen) and PDALD. This study shows that the addition of hard oxygen donors in PDALD and PDALC to the phen backbone and relative polarity of the end groups affects the binding properties of these hemicyclic ligands. Table 36. Comparison of logK1 and ΔlogK1 values with a selection of metal ions with PDALD, phen, and PDALC. Cu2+ Zn2+ Cd2+ Gd3+ Ca2+ Th4+ UO22+ Pb2+ ionic radius (Å) 0.57 0.74 0.95 0.97 0.99 1.09 1.10 1.19 logK1(PDALD) 6.37 4.76 6.31 5.95 5.74 4.98 4.86 6.38 a 1.00 3.81 - 4.65 Metal ion: logK1(phen) 24 9.13 6.38 5.66 0.9 7.56 6.56 7.49 6.16 3.74 7.20 6.25 7.32 logK1(PDALD) - logK1(phen) -2.76 -1.62 0.65 5.15 4.74 1.17 - 1.73 logK1(PDALD)-logK1(PDALC) -1.19 -1.80 -1.18 -0.21 2.00 -2.22 -1.39 -0.94 33,39 logK1(PDALC) difference in logK1: a A value of logK1=2.6 for Nd2+with phen in 5M NaCl has been accepted,39 which indicates logK1 ∼ 0.9 as a rough figure for Ln2+ ions with phen when corrected to ionic strength (μ) 0.1 by comparison with other neutral ligands where logK1 is known both at both μ=0.1 and 5.0.40 PDALD, like phen, forms a five-membered ring when complexed with a metal ion. The addition of the carboxyaldehyde groups in PDALD adds two neutral oxygen donors and allows the formation of two additional five-membered chelate rings. As a general rule, the addition of neutral oxygen donors increases the specificity for larger metal ions over smaller metal ions. The larger metal ions included in this study (Cd2+, Gd3+, Ca2+, Th4+, and Pb2+) all showed an increase in formation constants with PDALD over phen. The smaller metal ions investigated (Cu2+ and Zn2+) showed a decrease in logK1 with PDALD compared with the reported phen complexes. Similar results have been observed in PDALC39, which indicates that the addition of these polar groups hinders complex stability with smaller ions. This can likely be attributed to 116 steric strain on the ligand associated with coordinating more donors in the proper geometry around the smaller metal ions. PDALC and PDALD binding affinities followed similar trends when compared to phen, with PDALD complexes having slightly lower formation constants in nearly every case. This is because both PDALD and PDALC have added hard oxygen donors to the phen structure. The difference in complex stability seen in PDALD and PDALC is due to the differing type of oxygen donor present in each ligand. As expected, the neutral carboxyaldehyde oxygen donors of PDALD proved to be selective for larger metal ions. The alcoholic negative oxygen donors of PDALC had a larger influence in determination of binding character of these ligands and directly correlated to the logK1(OH-) for each metal ion. The Ca2+-PDALD complex was the only PDALD complex which had a higher logK1 than both phen and PDALC. This can be explained by the affinity of the hard metal ion for the addition of hard oxygen donors, extremely low affinity for the hydroxide ion (logK1(OH-) = 1.3), and the affinity for neutral oxygen donors of PDALD due to the large ionic radius of Ca2+ (r+=0.99). 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