Download metal ion complexing properties of the highly preorganized ligands

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)26H2O (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)26H2O (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)26H2O (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)26H2O (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)26H2O (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)26H2O (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)26H2O (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)4H20 (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)26H2O (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)26H2O (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)2H2O (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)36H2O (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)26H2O (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).
Bond length and angle in PDALC may also be slightly more ideal for the five-membered
ring structure of the complex because of the greater attraction between the negative oxygen of
the alcohol group and the bound cation. However, comparison to the bond lengths and angles of
PDALD complexes was not possible as numerous attempts to culture crystals were unsuccessful.
117
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