Thermodynamic Stability and Speciation of Ga(III) and Zr(IV) Complexes with High-Denticity Hydroxamate Chelators

Article abstract of DOI:10.1021/acs.inorgchem.1c01622

Increasing attention has been recently devoted to 89Zr(IV) and 68Ga(III) radionuclides, due to their favorable decay characteristics for positron emission tomography (PET). In the present paper, a deep investigation is presented on Ga(III) and Zr(IV) comp

Full text of DOI:10.1021/acs.inorgchem.1c01622

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Article
Thermodynamic Stability and Speciation of Ga(III) and Zr(IV)
Complexes with High-Denticity Hydroxamate Chelators
Yuliya Toporivska, Andrzej Mular, Karolina Piasta, Małgorzata Ostrowska, Davide Illuminati,
Andrea Baldi, Valentina Albanese, Salvatore Pacifico, Igor O. Fritsky, Maurizio Remelli, Remo Guerrini,
and Elzbieta Gumienna-Kontecka*
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ABSTRACT: Increasing attention has been recently devoted to 89Zr(IV) and
68Ga(III) radionuclides, due to their favorable decay characteristics for positron
emission tomography (PET). In the present paper, a deep investigation is presented
on Ga(III) and Zr(IV) complexes with a series of tri-(H3L1, H3L3, H3L4 and
desferrioxamine E, DFOE) and tetrahydroxamate (H4L2) ligands. Herein, we
describe the rational design and synthesis of two cyclic complexing agents (H3L1
and H4L2) bearing three and four hydroxamate chelating groups, respectively. The
ligand structures allow us to take advantage of the macrocyclic effect; the H4L2
chelator contains an additional side amino group available for a possible further
conjugation with a biomolecule. The thermodynamic stability of Ga(III) and Zr(IV)
complexes in solution has been measured using a combination of potentiometric and
pH-dependent UV−vis titrations, on the basis of metal−metal competition. The
Zr(IV)-H4L2 complex is characterized by one of the highest formation constants
reported to date for a tetrahydroxamate zirconium chelate (log β = 45.9, pZr = 37.0), although the complex-stability increase derived
from the introduction of the fourth hydroxamate binding unit is lower than that predicted by theoretical calculations. Solution
studies on Ga(III) complexes revealed that H3L1 and H4L2 are stronger chelators in comparison to DFOB. The complex stability
obtained with the new ligands is also compared with that previously reported for other hydroxamate ligands. In addition to
increasing the library of the thermodynamic stability data of Ga(III) and Zr(IV) complexes, the present work allows new insights
into Ga(III) and Zr(IV) coordination chemistry and thermodynamics and broadens the selection of available chelators for 68Ga(III)
and 89Zr(IV).
To the best of our knowledge, the most widely used chelator
for 68Ga is 1,4,7,10-tetraazacyclododecane-1,4,7,10-tetraacetic
acid (DOTA).11 In 2016 the FDA approved a 68Ga-
DOTATATE (NETSPOT) kit for the preparation of a 68Ga-
DOTATATE injection;12 clinical trials revealed the superiority
of 68Ga-DOTATATE with respect to 111In-pentetreotide in
imaging neuroendocrine tumors.13 Currently, the most
successfully used 89Zr(IV) chelator is DFOB, but some
decomposition has been observed over time in vivo, and 89Zr
slowly accumulates in bones.14−16
INTRODUCTION
■
Recent research confirms an increasing interest in the use of
gallium and zirconium radioiosotopes for medical diagnostic
techniques such as PET or single-photon emission computed
tomography (SPECT).1−8
The interest in the use of 68Ga (t1/2 = 1.13 h, Eβ+avg = 830
keV, 89%) for clinical PET comes from the accessibility of its
production via an easily portable and long-lived 68Ge/68Ga
generator system.9 The favorable decay characteristics of 89Zr
(t1/2 = 78.4 h, Eβ+avg = 395.5 keV, 23%) allow high PET image
resolution to be obtained, since the sufficiently long half-life is
an optimal match for the pharmacokinetics of most
monoclonal antibodies.10
Many chelators have been already suggested for 68Ga and
89Zr on the basis of in vivo assays.14,17−23 While scientists are
currently devoting considerable efforts toward the design of
To be applied to molecular imaging, the metal isotope must
be converted into a radiolabeled probe that can specifically
reach the target of interest in vivo and remain there long
enough to be detected. Therefore, the metal ion must be
bound by an efficient chelator to overcome metal hydrolysis
and transchelation, and linked to a biologically active targeting
molecule, to be properly directed to the desirable molecular
target in vivo.
Received: June 2, 2021
© XXXX The Authors. Published by
American Chemical Society
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Scheme 1. Structures of the Ligands Investigated and Discussed in This Paper
more efficient 89Zr and 68Ga chelators by increasing the in vivo
stability of the corresponding complexes, the solution
equilibrium chemistry, especially of Zr(IV) systems, has still
rarely been investigated.2,24−27 Solution studies on the
coordination chemistry and formation constants of Zr(IV)
complexes are not trivial for several reasons: an extremely high
thermodynamic stability (requiring application of indirect
competition methods with the use of other strong ligands with
known stability constants), strong hydrolysis (occurring over
almost the entire pH range), and the lack of spectral activity of
the complexes. On the other hand, knowledge of the speciation
of such complexes, especially at physiological pH, could
provide information concerning the actual chemical form of
the complex in biological media, and this can contribute to
both a better understanding of the in vivo speciation and an
explanation of the differences in the biological activity.
Our laboratory has recently reported the thermodynamic
properties of Zr(IV)-DFOB complexes, suggesting the
formation of three mononuclear complexes, i.e. [ZrHL]2+
[ZrL]+, and [ZrLH−1], over the pH range 1−11.24 The
stability constants and pZr value determined for the Zr(IV)-
DFOB system place DFOB among the best Zr(IV) chelators,
although the formation of six-coordinate unsaturated com-
plexes (i.e., with the coordination sphere of Zr(IV) completed
by two water molecules28 or hydroxide ligands29) and the
susceptibility of coordinated water molecules to deprotonation
were suggested to be the reason for the in vivo lability of
89Zr(IV)-DFOB complexes. The thermodynamic stability of
Zr(IV)-DFOB complexes is in line with in vivo re-
search15,16,30,31 and also with Holland’s recent DFT calcu-
lations,32 indicating that our experimental approach was
appropriate.
By capitalizing upon our earlier works on siderophore
mimics,33−36 in this work we have designed, synthesized, and
fully characterized tri- and tetrahydroxamic H3L1 and H4L2
chelators (Scheme 1), analogues of DFOE,37 and their Ga(III)
and Zr(IV) complexes. In addition, the physicochemical
properties of Zr(IV) complexes with three further trihydroxa-
mic ligands, i.e. H3L3 (FOXE 2-5),38 H3L4 (T4),35 and DFOE
(Scheme 1), were also investigated for the sake of comparison.
The ligands employed here were selected to investigate the
influence of some structural elements on the physicochemical
properties of Zr(IV) complexes: i.e., (i) the number of binding
groups required to complete the coordination sphere of the
metal cation; (ii) the possible advantage of the macrocyclic
effect; (iii) the size of the ligand cavity, which should be large
enough to minimize ring strain; (iv) the symmetry of the
ligand, which could limit the probability of a complex
challenge. Ga(III) (and Fe(III)) binding to all of these ligands
was also performed and will be discussed here, as the relation
between the ligand structure and the stability of the complexes
has been much more studied and is better understood for
trivalent metal ions. Moreover, Fe(III) complexes are used as a
tool to determine the thermodynamic stability of Ga(III) and
Zr(IV) analogues.
According to DFT studies, to minimize the ring strain, the
cyclic tetrahydroxamic ligand should consist of a minimum of
36 atoms, which gives at least 7 chemical groups or 8
bonds.39,40 The tetrahydroxamic ligand H4L2, designed to
completely saturate the oxophilic coordination sphere of
Zr(IV), meets this criterion, possessing at least 9 bonds
B
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a
Scheme 2. Synthetic Pathway of the Ligand H3L1
a
Reagents and conditions: (i) (Boc)2O, K2CO3, H2O/dioxane; (ii) ethyl 4-bromobutyrate, NaH, DMF; (iii) TFA; (iv) Boc-γ-aminobutyric acid,
HATU, DIPEA, DMF; (v) LiOH, H2O/dioxane; (vi) 5, HATU, DIPEA, DMF; (vii) 6, HATU, DIPEA, DMF; (viii) HATU, DIPEA, DMF; (ix)
H2, Pd/C, MeOH.
between binding groups. Of note, H4L2 has been designed
with the aim of introducing in the chelator a primary amine
group, useful to improve the solubility of the ligand but,
primarily, to allow the easy conjugation of the chelator to a
targeting molecule for future in vivo studies. Trihydroxamic
H3L1 is a symmetrical, macrocyclic ligand, comprising 9 bonds
between binding groups; it allows the determination and direct
comparison of the influence of the fourth binding group of
H4L2 on Zr(IV) stability. This ligand also reveals the effect of
the shortening of the chain in comparison to DFOE, H3L3,
and DFOB, all with 10 bonds (Scheme 1). Tripodal H3L4,
used earlier as a good mimic of a ferrichrome siderophore, was
investigated here in order to compare its Zr(IV) binding
capacity to those of other tri- and tetrahydroxamic ligands.35 In
H3L1, H4L2, and H3L338 we have used retrohydroxamic units,
with a reversed order with respect to the native hydroxamic
moiety.
Until now, the stability constants for tetrahydroxamate
Zr(IV) complexes have only been estimated from computa-
tional calculations32 for linear DFO* (log β[Zr(DFO*)] = 51.56)
and cyclic CTH36 (log β[Zr(CTH36)] = 52.84). DFO*41 and
CTH3639 possess 10 and 8 bonds between two hydroxamic
groups, respectively.
Therefore, an experimental verification of the order of the
stability increase between tri- and tetrahydroxamic chelators
should be very useful. The same is true for examining the
effects of a number of other variable structural elements, such
as the symmetry of the structure and the length of the chain
between the hydroxamate binding units. Although the
thermodynamic formation constants do not predict in vivo
stability and kinetic inertness, they are very useful and
interesting parameters that may allow for a better design of
efficient chelators for Zr(IV) ions.
RESULTS AND DISCUSSION
■
Design and Synthesis of the Ligands. The synthetic
approaches for the preparation of H3L1 and H4L2 are depicted
in Schemes 2 and 3, respectively. The hydroxamate-based
monomers protected either at the carboxylic group as an ethyl
ester (5) or at the amino function with Boc (6) were employed
as common synthetic precursors of H3L1 and H4L2. The
building blocks 5 and 6 were synthesized by starting from O-
benzylhydroxylamine hydrochloride that was first reacted with
di-tert-butyl dicarbonate to give compound 1 and then
alkylated with ethyl 4-bromobutyrate in the presence of NaH
to provide compound 2 (Scheme 2). Boc removal with TFA
furnished the intermediate 3, which was coupled with Boc-
protected γ-aminobutyric acid using HATU as a coupling
reagent. This allowed us to obtain the orthogonally protected 4
as a suitable precursor of both 5 and 6, which were
alternatively isolated after acidic and basic treatments,
respectively.
C
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a
Scheme 3. Synthetic Pathway of the Ligand H4L2
a
Reagents and conditions: (i) TFA; (ii) Z-Lys(Boc)-OH, HATU, DIPEA, DMF; (iii) LiOH, H2O/dioxane; (iv) HATU, DIPEA, DMF; (v) H2,
Pd/C, MeOH.
For the synthesis of H3L1, 5 and 6 were linked together via a
standard amide coupling followed by TFA treatment to give 8.
The reaction with another unit of 6 gave 9 as the linear
protected precursor of H3L1. A head−tail HATU-mediated
cyclization of 9 was realized under dilute conditions (0.5 mg/
mL) after removal of the protection at the C and N terminal
positions. A final Pd-catalyzed hydrogenolysis afforded the
desired macrocycle H3L1 in good yields.
For the synthesis of the tetrahydroxamic derivative H4L2
(Scheme 3), an appropriate hydroxamate-bearing lysine
derivative was first prepared as a building block (14). This
was obtained by a coupling reaction between 5 and a residue of
Z-Lys(Boc)-OH followed by saponification of the ester
function. The monomeric unit 14 was coupled to 12, which
resulted from TFA-mediated Boc deprotection of 9. The
resulting intermediate 15 underwent head−tail HATU-
mediated cyclization after removal of the protections at the
C and N terminal positions as described above, leading to 16.
In addition in this case, the final macrocycle was obtained after
Pd-catalyzed removal of the benzyl functions from the
hydroxamic groups. These conditions led also to CBz cleavage,
leaving the free amino group suitable for future bioconjugation
strategies.
Thermodynamic Solution Studies. Ligand Protonation
Constants. The metal affinity of a ligand depends on its acid−
base properties; therefore, the protonation equilibria of H3L1
and H4L2 were first investigated. The proton-dissociation
processes of the ligands were followed by potentiometric
titrations in the pH range 2−11. The hydrolytic stability of the
ligands was monitored by a second titration of the same sample
with NaOH, following back-acidification of the initially titrated
sample. The recorded titration curves were almost exactly
superimposed; consequently, the protonation constants
calculated from the two consecutive titrations were found to
be equal within 0.05 log unit, which indicated that no
decomposition occurred.
Data analysis allowed the determination of three protonation
constants for H3L1 and four protonation constants for H4L2;
all of them fall in the pH range 8−10 and can be attributed to
the hydroxamate groups (Table 1). For each ligand, the
protonation constants were assigned by comparing them with
the known protonation constants of hydroxamate li-
gands.24,35,37,42 When the changes in temperature, ionic
strength, and ligand structures are allowed for, the protonation
constants of H3L1 and H4L2 are in excellent agreement with
the literature values of the cyclic hydroxamate siderophore
DFOE (log K1 = 9.89, log K2 = 9.42, and log K3 = 8.65)
reported by Anderegg et al.37 The pH-dependent UV−vis
titrations of H3L1 and H4L2 (Figure S1) revealed the
development of a strong band with λmax = 230 nm, when the
pH was increased from 7 to 11, which is usually observed for a
hydroxamic group deprotonation process.24,43 The amino
group protonation of H4L2 was not detectable in the
H3L1, H4L2, and their precursors were fully characterized by
1H and 13C NMR (see the Supporting Information). The
degree of purity of the final product was evaluated by analytical
HPLC assays (see the Supporting Information), showing a
purity of higher than 95%.
D
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a
Table 1. Protonation Constants of Ligands and log β Values of Complexes Formed with Fe(III), Ga(III), and Zr(IV) Ions
b
c
b
,
H3L3 (FOXE 2−5)38
H3L4 (T4)35
DFOE36,38
,
,
H3L1
H4L2
assignt
log β
log K
log β
log K
log β
log K
log β
log K
log β
log K
d
d
d
d
d
d
d
LH
9.89(1)
19.13(1)
27.44(1)
9.89
9.24
8.31
10.06(1)
19.65(1)
28.59(1)
36.77(1)
10.06
9.59
8.94
8.18
9.89
19.31
27.96
9.89
19.31
27.96
9.50
18.47
26.73
9.50
8.97
8.26
9.89
19.31
27.96
9.89
9.42
8.65
d
d
d
d
d
d
LH2
LH3
LH4
d
b
c
,
b
,
H3L3 (FOXE 2−5)38
,
H3L1
H4L2
H3L4 (T4)35
log β
DFOE36,38
log β
pKa
log β
pKa
log β
pKa
pKa
log β pKa
e
d
[FeH2L]
[FeHL]
[FeL]
39.96(3)
3.00
3.12
d
e
e
d
e
31.80(3)
28.59(2)
3.21
3.1
36.96(7)
36.82(6)
29.82
2.48
2.36
d
d
32.43
32.43
27.34
32.43
d
28.7(1)
f
d
[GaH2L]
[GaHL]
[GaL]
38.11(3)
35.91(7)
27.60(6)
2.20
e
2.15(4)
f
d
d
d
d
f
d
29.44(7)
26.79(2)
2.65
8.13
27.92(3)
25.56(1)
e
e
2.53(2)
2.24(1)
d
d
d
9.79
29.79
29.79
29.79
f
d
f
d
[ZrHL]
[ZrL]
45.9(3)
2.6
36.4(5)
2.15
e
e
2.2(1)
2.2(2)
f
d
d
f
d
d
d
f
34.8(2)
5.48
43.3(1)
35.46(5)
28.31(7)
7.03
34.25(5)
5.43
35.54(9)
e
e
e
5.34(5)
7.2(1)
5.5(2)
d
d
d
[ZrLH−1
]
29.32(8)
28.8(1)
a
b
Charges omitted for clarity. The protonation constants of the ligands together and the stability constants of Fe(III) and Ga(III) complexes were
c
taken from the literature. The protonation constants of the ligands and the stability constants of Fe(III) complexes were taken from the literature.
d
e
f
Determined by potentiometric titrations. Determined by pH-dependent UV−vis titrations. Determined by metal−metal competition titrations;
all measurements performed at 25 °C and I = 0.1 M NaClO4.
experimental pH range. We are aware that the electron-
withdrawing character of both the −NH3 and −NHOH
details see Figure S3 and Table S1 in the Supporting
Information).
+
groups affects the acidity of the other group in comparison
with that in the related nonsubstituted compound H3L1. As
was previously reported in the example of α- and β-alanine
hydroxamic acids, the amino group may be more acidic than
the hydroxamic group, or vice versa. According to the
literature, the amino group may be more acidic than
hydroxamic one, or vice versa.44,45 On the basis of the
protonation constant of the amino group of DFOB (log Kamine
= 10.9724), we assume for H4L2 that it is >11. However, we
keep in mind that the deprotonation processes of amino and
hydroxamate groups overlap and cannot be distinguished by
potentiometry. To elucidate the protonation microequilibria of
Determination of Complex Stability. To evaluate the
thermodynamic stability of Ga(III) and Zr(IV) complexes of
the investigated ligands, the binding properties and speciation
of Fe(III) complexes have first been determined (all the details
are given in the Supporting Information). There are several
reasons for this protocol. First, (i) the electron configuration of
Ga(III) (d10) and Zr(IV) (d0) hinders the attainment of
spectral information for most of the complexes. Furthermore,
(ii) both metal ions are highly acidic and they are readily
hydrolyzed over almost the entire pH range. In addition, (iii)
the high charge to size ratio of the Zr(IV) ion implies the
formation of complexes with an exceptional thermodynamic
stability (already at very low pH); as a consequence, the
stability constants cannot be directly determined using
standard potentiometric titrations. Thus, the thermodynamic
stability constants of the Fe(III)-H3L1 and Fe(III)-H4L2
systems were first determined (using a combination of
potentiometric and pH-dependent UV−vis titrations), fol-
lowed by Fe(III)−Ga(III) and Fe(III)−Zr(IV) metal−metal
competition experiments. Of importance, in order to get
accurate results, an experiment where two metal ions compete
for a ligand must fulfill two basic requirements: (i) one of the
metal chelates should have a strong absorption band in teither
he visible or ultraviolet region of the spectrum, with an
extinction coefficient much different from that of the free metal
ion, while the second metal complex should not absorb in the
same region of the spectrum; (ii) the equilibrium constant for
the competition reaction must not be too small or too large.
1
H4L2, H NMR titrations should be carried out. However,
such a precise analysis is not needed for the determination of
the stability of H4L2-metal complexes and therefore was not
performed. The species distribution diagrams of H3L1 and
H4L2 are presented in Figure S2. The protonation constants of
H3L3 and H3L4 were reported elsewhere (Table 1).35,38
ESI-MS: Stoichiometry Evaluation. The stoichiometry of
the complexes was evaluated by ESI-MS, frequently used as the
first step in the determination of metal complex stoichiometry
and already previously employed.35,46,47 When the fact that
ESI-MS is not able to distinguish the ionizable protons in the
species is taken into account, this method can be successfully
applied to evaluate the metal to ligand stoichiometry directly
from the m/z values. For all of the investigated systems, an
analysis of the ESI-MS data (collected for various metal to
ligand molar ratios) revealed only mononuclear complexes (for
E
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Figure 1. UV spectra of Ga(III)-H3L1 (a, c) and Ga(III)-H4L2 (b, d) systems at a metal to ligand molar ratio of 1:1 in the pH range 1.0−11.0.
Conditions: cL1 = 0.05 mM, cL2 = 0.05 mM, 0.1 M NaClO4, T = 25 °C.
These requirements are fulfilled in the competition experi-
ments between Fe(III) and Ga(III) for both ligands. In the
case of Fe(III)−Zr(IV) competition, the difference between
the stability constants of Fe(III) and Zr(IV) complexes was
too high; thus, an additional competitor ligandnitrilotriacetic
acid (NTA)was used in the titrations.
the hydroxamate group protonation state, were monitored.51
The appearance of a band with a maximum at 225−230 nm
with an increase in pH was observed and associated with
complex formation.
The pH-dependent UV−vis titration experiments for the
Ga(III)-H3L1 system revealed an increase in a 230 nm
transition band starting from pH 1 up to pH 4.4, with pKa =
2.53(2) (Figure 1 and Table 1). For the Ga(III)-H4L2 system,
the 225 nm band development was observed starting from pH
1 up to pH 3.5, with pKa = 2.15(4) (Figure 1 and Table 1),
while for Ga(III)-H3L4, it continued to increase up to pH 4.6
with pKa = 2.36(1) (Figure S7 and Table 1). Similar behavior
was observed in the acidic range for Ga(III)-DFOB (and was
NTA is one of the most widely investigated and often used
chelating agents.48,49 It was selected for the current studies, as
both Zr(IV)-NTA and Fe(III)-NTA complexes remain stable
until pH 4, even at a metal to ligand molar ratio of 1:1.27,48
Moreover, as an additional competing agent, NTA prevents the
hydrolysis of the metal ions present in solution and weakens
the transchelation observed in the case of H3L1 and H4L2
Fe(III) complexes titrated directly by Zr(IV) ions. The
accuracy of the metal−metal competition titration with NTA
was checked on the Zr(IV)-DFOB system, for which
experimental data gave log β[ZrHDFOB] = 46.1(2) (see the
Supporting Information), in very good agreement with our
recently reported data (log β[ZrHDFOB] = 47.7, allowing for
changes in the ionic strength).24 Similar competition
procedures are widely used for the evaluation of the stability
constants of spectroscopically blind metal complexes.25,50
Ga(III) Complex Formation Equilibria. The evaluation of
thermodynamic stability constants of Ga(III) and Zr(IV)
complexes with H3L1, H4L2, and H3L4 started from pH-
dependent UV−vis spectrophotometric titrations in the pH
range 1−11 (Figure 1 and Figure S7, respectively). The
spectral changes in the 200−300 nm range, corresponding to
assigned to two protonation constants of the complex, pKa1
=
0.78 and pKa2 = 1.1043) and Th(IV)-DFOB (with pKa = 1.942)
complexes. For the three investigated systems, UV spectra did
not reveal any additional changes up to pH 9, indicating that
the fully coordinated complex [GaL], in the case of the
Ga(III)-H3L1 and Ga(III)-H3L4 systems, and the monop-
rotonated [GaHL2]+ complex, in the Ga(III)-H4L2 system,
are the dominant species in solution. When the pH was
increased to above 9, an increase in absorbance below 240 nm
was observed (Figure 1 and Figure S7). Considering that the
spectrophotometric titrations of the free ligands showed the
same absorption curves at higher pH, we can suppose that the
sharp band at 230−240 nm arises from unbound deprotonated
hydroxamate chromophores. Most probably, at higher pH the
complex is dissociated, yielding the hydrolyzed gallium species
F
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Figure 2. UV−vis metal−metal competition experiment for (a) Fe(III)-H3L1+Ga(III) (cFe(III) = 0.075 mM, cL = 0.075 mM) and (b) Ga(III)-
H3L1+Fe(III) (cFe(III) = 0.08 mM, cL1 = 0.08 mM) systems at pH 1.5, I = 0.1 M (NaClO4), and T = 25 °C.
Figure 3. UV spectra of the Zr(IV)-H3L1 (a, c) and Zr(IV)-H4L2 (b, d) systems at a metal to ligand molar ratio of 1:1 in the pH range 0.1−11
Conditions: cL1 = 0.05 mM, cL2 = 0.037 mM, 0.1 M NaClO4.
[Ga(OH)4]−. A similar observation has already been noted for
hydroxamate ligands.43,52
[FeHL1]+ (λmax = 470 nm, Figure 2a) slowly disappeared as a
result of the [GaHL1]+ complex formation. In the next
competition experiment, the appearance of an LMCT
transition band (λmax = 470 nm) upon addition of Fe(III) to
the Ga(III)-H3L1 solution was observed (Figure 2b). The data
refinement using H3L1 protonation constants (Table 1),
Fe(III)-H3L1 stability constants (Table 1), and stability
constants of hydroxocomplexes of both metals (see the
Assuming the domination of the [GaHL1]+ complex below
pH 2, its stability was determined by metal−metal competition
titrations, (i) Fe(III)-H3L1 + Ga(III) and (ii) Ga(III)-H3L1 +
Fe(III), both performed at pH 1.5. This pH was chosen in
order to prevent the hydrolysis of the free metal ions and
decomposition of the ligand, which is common for hydroxamic
acids at a very acidic pH.37,51 Upon addition of Ga(III) (up to
600 equiv) to the Fe(III)-H3L1 solution, the UV−vis band of
+
Experimental Section) yielded a log β[GaHL1] value of
29.44(7) for Fe(III)-H3L1 + Ga(III) (Table 1) and of
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Figure 4. UV−vis spectra of competition titrations of the Fe(III)-H3L1 + Zr(IV)-NTA (cFe(III) = 0.098 mM, cL1 = 0.098 mM) (a) and Fe(III)-
H4L2 + Zr(IV)-NTA (cFe(III) = 0.055 mM, cL2 = 0.055 mM) (b) systems at pH 1.5 with 0.1 M NaClO4.
28.3(6) in the case of Ga(III)-H3L1 + Fe(III) competition
titrations.
pH 7.0, the spectra do not reveal any significant changes until
pH 9.0, where the hydrolysis probably started. Since
information about hydrolysis of the hydroxamate ligands at
acidic pH has beenwidely described in the literature,51,52 pKa1
= 0.4 indicates that the three hydroxamate groups are
dissociated and therefore most probably bound to the Zr(IV)
ion, already at pH <2.0. Assuming the formation of only
monomeric complexes, the stability of [ZrL1]+ was determined
via UV−vis competition batch experiments, using a Zr(IV)-
NTA solution as a competing system for the Fe(III)-H3L1
complex (Figure 4a). The large LMCT band centered at 470
nm characteristic of dihydroxamate [FeHL1]+ species
decreased gradually, and the refinement of the titration data,
using the Fe(III)-H3L1 stability constants (Table 1) together
with the Fe(III) and Zr(IV) hydrolysis constants, yielded a log
For the Ga(III)-H4L2 system, [GaH2L2]+ was assumed to
be the most abundant species at pH 1.5, and its stability was
again determined via (i) Fe(III)-H4L2 + Ga(III) and (ii)
Ga(III)-H4L2 + Fe(III) titrations (Figure S8). The refinement
+
of the titration data yielded log β[GaH L2] = 38.11(3) and
2
39.06(4) for Fe(III)-H4L2 + Ga(III) and Ga(III)-H4L2 +
Fe(III) competition titrations, respectively.
For the Ga(III)-H3L4 system, [GaHL4]+ was assumed to be
the major complex at pH 1.5, and its stability was determined
via an Fe(III)-H3L4 + Ga(III) competition titration (Figure
+
S9). The refinement of the titration data yielded log β[GaHL4]
=
27.92(3).
For the Ga(III)-H3L1 and Ga(III)-H4L2 systems, the
constants obtained with the two types of titrations are not
far from each other but are significantly different. Most likely,
the constants measured by means of Ga(III)-H3L1/H4L2 +
Fe(III) competition experiments are endowed with a greater
error, attributable to the overlapping absorption of free iron,
present in excess. Therefore, we retained the constants
obtained from the Fe(III)-H3L1/H4L2 + Ga(III) competition
+
β[ZrL1] value of 34.8(2).
The UV−vis titration data of the Zr(IV)-H4L2 system
(Figure 3b,d) showed the development of a 230 nm shoulder
upon an increse in pH from 0.1 to 3.0 and allowed us to
calculate a pKa value of 2.2(1). Further, the spectra did not
reveal any significant changes over the pH range 3.1−9.0,
wherethe hydrolysis probably started. [ZrHL2]+ was assumed
to be the major complex at pH 1.5, and its stability was
determined via Fe(III)-H4L2 + Zr(IV)-NTA (Figure 4b)
competition titrations. The refinement of the titration data
+
+
experiments (log β[GaHL1] = 29.44(7) and log β[GaH L2]
=
2
38.11(3)) as fixed values in the subsequent potentiometric
data calculations. The best-fitted speciation model for Ga(III)-
H3L1 and Ga(III)-H3L4 systems revealed the presence of one
additional complex, [GaL], with log β values of 26.79(2) and
25.56(1), respectively (pKa1 = 2.65 for Ga(III)-H3L1 and 2.36
for Ga(III)-H3L4; Table 1). For the Ga(III)-H4L2 system, in
addition to [GaH2L2]+, [GaHL2] and [GaL2]− complexes
+
yielded log β[ZrHL2] = 45.9(3).
The UV−vis spectra of the Zr(IV)-H3L3 and Zr(IV)-H3L4
systems were very similar to the spectra described above
(Figure S11); at a pH of around 7 a slight blue shift of the 230
nm band appeared together with the isosbestic point, allowing
us to calculate pKa = 7.2(1) for Zr(IV)-H3L3 and 5.5(2) for
Zr(IV)-H3L4 (Table 1). For Zr(IV)-H3L4, an additional pKa =
2.2(2) was calculated, probably corresponding to the
formation of the three-hydroxamate complex (Table 1). An
evaluation of the competition titration data for Fe(III)-H3L3 +
−
were found with log β[GaHL2] = 35.91(7) and log β[GaL2]
=
27.60(6) (pKa1 = 2.20 and pKa2 = 8.13; Table 1). The pKa2
value of 8.13 is in good agreement with the pKa values of the
free ligand (Table 1) and could be assigned to the
deprotonation of an unbound hydroxamate group.
+
Zr(IV)-NTA revealed log β[ZrL3] = 35.46(5), while for Fe(III)-
Zr(IV) Complex Formation Equilibria. The UV−vis
titrations of Zr(IV)-H3L1 equimolar solution over the pH
range 0.1−11 (Figure 3a,c) showed a well-defined absorbance
band in the 200−300 nm range. The significant changes and
the presence of an isosbestic point observed when the pH was
increased from 0.1 to 0.9, allowed us to calculate a pKa1 value
of 0.4(2); afterward, the observed 230 nm shoulder remained
stable up to pH 4.6. When the pH was increased to 7.0, the
development of a 220 nm shoulder with an isosbestic point at
230 nm was observed, characterized by pKa2 = 5.34(5). From
2+
H3L4 + Zr(IV)-NTA log β[ZrHL4] = 36.4(5) (Figure S12 and
Table 1). Additionally, we have performed the same kind of
titration for Fe(III)-DFOE + Zr(IV)-NTA, which gave log
+
β[ZrDFOE] = 35.54(9) (Figure S12 and Table 1). It is worth
noting that, during the competition titrations for the H3L3
ligand and DFOE, a decrease in the 430 nm (typical for three-
hydroxamate iron complexes) band was observed, confirming
that all three hydroxamate groups are bound already at pH
<2.38
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Using the constants obtained from competition titrations
−log[Zr(IV)free] were calculated at pH 7.4 with cL = 10 μM
and cGa(III)/Zr(IV) = 1 μM (Table 2).
+
+
(log β[ZrL1] = 34.8(2) and log β[ZrHL2] = 45.9(3)) as fixed
values, the potentiometric data were processed. The best-fitted
speciation models revealed the presence of one additional
complex for the Zr(IV)-H3L1 system ([ZrL1H−1], log
β[ZrL1H = 29.32(8)) and one additional complex for the
Table 2. pGa and pZr Values for Various Synthetic and
a
Natural Chelators
]
−1
ligand
H3L1
pGa
22.5
pZr
32.4
chelating groups and ligand geometry
Zr(IV)-H4L2 system ([ZrL2], log β[ZrL2] = 43.3(1)) (Table 1).
The [ZrL1]+ complex dominates the solution from pH 3 up
to pH ∼5, where its deprotonation to [ZrL1H−1] occurs, with
pKa = 5.48 (Figure S13a). These results are in line with the
spectroscopic data (pKa = 5.34(5)) and could be ascribed to
the dissociation of a water molecule from the coordination
sphere of the Zr(IV) ion. For the Zr(IV)-EDTA system, the
pKa attributed to the deprotonation of a water molecule is
6.2,27 while for Zr(VI)-DFOB it is 6.36.24
The stability constants calculated for the Zr(IV)-H4L2
system reveal pKa = 2.6, which is in good agreement with the
spectroscopic data (pKa = 2.2(1)) and could be assigned to the
deprotonation of the hydroxamic group and the formation of a
fully coordinated tetrahydroxamate complex. The [ZrL2]
complex dominates the solution from pH 3 up to pH ∼9,
when the dissociation of the complex probably occurs (Figure
S13b).
3 hydroxamate groups in a cyclic
arrangement
H4L2
22.3
37.0
31.5
32.6
32.224
31.0
4 hydroxamate groups in a linear
arrangement
3 hydroxamate groups in a cyclic
arrangement
3 hydroxamate groups in a linear
arrangement
3 hydroxamate groups in a linear
arrangement
3 hydroxamate groups in a cyclic
arrangement
4 macrocyclic amine groups and
4 carboxylate pendant arms
3 macrocyclic amine groups and 3
carboxylate pendant arms
2 8-hydroxyquinoline groups and 2 amino
groups in a linear arrangement
H3L3
25.4
H3L4
21.9
DFOB
DFOE
DOTA
NOTA
H2hox
PrP9
21.643
25.238
20.511
27.454
28.455
23.156
28.057
3 macrocyclic amine groups and 3
carboxylate pendant arms
2 hydroxyaromatic donor groups and
2 carboxylate pendant arms
An evaluation of the potentiometric data of Zr(IV)-H3L3
(with log β[ZrL3] = 35.46 fixed) revealed the presence of an
additional complex, [ZrL3H−1] (log β[ZrL3H = 28.35(5))
HBED
THPN
3,4,3-LI-
HOPO
42.72
4 3-hydroxy-4-pyridinone pendant arms
]
−1
44.025,58 4 1-hydroxy-2-pyridinonates in a linear
arrangement
(Figure S13c), while for Zr(IV)-H3L4 (with log β[ZrHL4] = 36.4
fixed), there are two additional complexes, [ZrL4] (log β[ZrL4]
DTPA
32.325,59 3 amino groups in linear arrangement and
4 carboxylate pendant arms
= 34.25(5)) and [ZrL4H−1] (log β[ZrL4H ] = 28.8(1)) (Figure
−1
a
S13d). For both systems, the pKa values calculated from
potentiometric experiments are in excellent agreement with
those from the pH-dependent UV−vis titrations (Table 1) and
match the data obtained for the whole series of hydroxamate-
based ligands. Of importance, the pKa value attributed to the
deprotonation of a water molecule in the Zr(IV)-H3L3 system
(pKa = 7.2) is higher than those for Zr(VI)-DFOB (pKa =
6.36)24 and Zr(IV)-H3L1 (pKa = 5.48), suggesting that a
longer linker between the binding groups is advantageous. In
Zr(IV)-H3L4, this process is not observed, as all the
coordinating positions of the Zr(IV) ion are occupied by
four hydroxamate ligands.
Ligand Sequestering Ability. Despite a large variety of PET
chelators synthesized and tested in order to provide strong
coordination of Ga(III) and Zr(IV) in vivo, until now it is has
been hard to avoid the release of these metal ions in the body.
Here we report two new Ga(III) and Zr(IV) hydroxamate
chelators, designed to achieve an efficient sequestering of these
metal ions but also to understand how the cyclization and the
introduction of an additional hydroxamate group influences
the stability of Zr(IV) complexes. Therefore, it is important to
evaluate and compare the Ga(III)- and Zr(IV)-sequestering
abilities of H3L1 and H4L2 with those of other chelators.
However, the direct comparison of the stability constants of
metal complexes is not straightforward, and other tools taking
into account all the physicochemical properties of the ligands,
i.e. their denticity, coordination modes, and acid−base
properties, should be used.52 In order to reliably compare
the chelating abilities of H3L1 and H4L2 toward Ga(III) and
Zr(IV) ions, the pM values were calculated. pFe was originally
introduced by Raymond for the comparison of iron-side-
rophore systems;53 pGa = −log[Ga(III)free] and pZr =
Values (re)calculated at pH 7.4 and cL = 10 μM and cGa(III)/Zr(IV) = 1
μM, on the basis of the protonation and stability constants given in
original publications. The hydrolysis constants of Ga(III) and Zr(IV)
ions were taken from the literature60 and are given in the
Experimental Section.
The pGa values for Ga(III)-H3L1 and Ga(III)-H4L2
systems are on the same order of magnitude as those of the
well-known gallium chelators DFOB and PRP9 but are higher
than that of the clinically used DOTA (Table 2). On the other
hand, H2hox,55 NOTA,54 and HBED57 present much higher
Ga(III) chelating efficacy. The observed effect reflects the
differences in the number and type of chelating groups present
in the ligands (and therefore the number and type of donor
atoms), as well as the ligand dimensions. Of importance, there
is only about a 1 order of magnitude increase between pGa
values for the linear trihydroxamate ligand DFOB and the
cyclic tri- and tetrahydroxamates H3L1 and H4L2, respectively.
This indicates that the cyclization of the structure only slightly
influences the complex stability. It is worth underlining that
H3L1 has shorter spacers between the hydroxamate groups (9
bonds) in comparison to those in DFOB and DFOE (10
bonds). In H3L3, the spacers have the same length of 10 bonds
and the Ga(III) complexes reach the stability of DFOE.38
Additionally, there is almost no difference in complex stability
between Ga(III)-H3L1 and Ga(III)-H4L2, even though the
cavity of H4L2 is much larger than that of H3L1, allowing
higher flexibility and entropy of the complex structure.
For the Zr(IV) complexes of trihydroxamate H3L1, H3L3,
and H3L4 systems, the pZr value is on the same order as those
for DFOB24 and DTPA25 chelators. This suggests that ligand
cyclization does not provide any increase in complex stability
with respect to its linear analogue; for H3L3 and DFOE one
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may even claim a slight decrease in relation to DFOB. A similar
conclusion was drawn from the comparison of Zr(IV)
complexes of DFOB with fusarinine C (FCS, Scheme 1),
where only minor differences in complex stability were
observed in in vivo studies.61 An elongation of the chain
between hydroxamate binding units from 9 bonds in H3L1 to
10 in H3L3 is not reflected in the corresponding pZr values.
The flexibility of the tripodal H3L4 ligand does not produce a
higher stability of Zr(IV) complexes. Of importance, the pZr
value for the tetrahydroxamate analogue H4L2 is >4 units
higher than the values calculated for Zr(IV)-DFOB and
Zr(IV)-H3L1 systems, reflecting the higher affinity of
tetrahydroxamate H4L2 for Zr(IV), as expected. This feature
was already observed in biological studies of other tetrahy-
droxamate chelators.4,39,41 The high thermodynamic stability is
certainly the result of the involvement of the fourth
hydroxamate coordinating group of the ligand moiety.
The H3L1 and H4L2 ligands are very good examples to
directly compare the stability of Zr(IV) complexes formed with
tri- and tertrahydroxamate compounds. For these two ligands,
we observe an increase in log β of 8.5 orders of magnitude for
the Zr(IV)-H4L2 complex with respect to Zr(IV)-H3L1 (Table
1). However, this increase is lower than that predicted from
the computational calculations performed by Holland32 for tri-
dependent on the ligand geometry and emphasize the demand
for a thermodynamic solution study in order to understand this
dependence. Other octadentate hydroxy-pyridinone chelators,
such as THPN2 and 3,4,3-LI-HOPO,25 form the strongest
complexes (Table 2). The reason could be not only the type of
chelating groups present in the ligands but also the ligand
architecture and dimensions.
CONCLUSIONS
■
In the present work we have developed new chelating agents
for the complete saturation of the coordination spheres of
Ga(III) and Zr(IV) metals. The trihydroxamic (H3L1) and
tetrahydroxamic (H4L2) ligands were successfully synthesized,
and the thermodynamic properties of their Ga(III) and Zr(IV)
complexes were evaluated. In addition, a series of other
synthetic (H3L3, H3L4) and natural (DFOE) compounds was
investigated. H3L1 proved to be an efficient Ga(III) chelator,
but the stability of its Zr(IV) complexes is about 1 order of
magnitude lower than that reported for the Zr(IV)-DFOB
system. H4L2 is the first tetrahydroxamate ligand for which the
formation constants and speciation with Zr(IV) were
experimentally determined. Of importance, it revealed an
enhanced stability of 8.5 orders of magnitude (log β[ZrL2]
43.3, pZr = 37.0) with respect to Zr(IV)-H3L1 (log β[ZrL1]
=
=
and tetrahydroxamate chelators: i.e., DFOB (log β[Zr(DFOB)]
=
41.20) versus linear DFO*41 (log β
= 51.56) and
34.8, pZr = 32.4), as a consequence of the introduction of a
fourth hydroxamate binding unit. However, the stability
increase is lower than that predicted by computational
calculations for the tetrahydroxamate chelators DFO* (log
β[Zr(DFO*)] = 51.56) and CTH36 (log β[Zr(CTH36)] = 52.84), and
this effect can be ascribed to the structural alterations of the
H4L2 ligand.
Overall, the determination of the thermodynamic stability of
metal complexes coupled with a suitable chelator design will
help further developments of optimal chelators for PET
imaging applications. However we are aware that there are still
a great number of tests to do in order for these ligands to be
used as a PET chelators, such as radiolabeling and kinetics
studies, biodistribution assays, etc. Current efforts are focused
on the design and studies of tetrapodal hydroxamate ligands, to
interrogate how their shape and size tune the thermodynamic
stability of Zr(IV) complexes.
[Zr(DFO*)]
cyclic CTH3639 (log β[Zr(CTH36)] = 52.84). Also, log β[ZrL2]
=
43.3(1) does not reach the values predicted from the above
calculations for eight-coordinate Zr(IV)-DFO* and Zr(IV)-
CTH36. Of note, the log β[ZrH(DFOB)OH] value previously
determined by us for DFOB (40.04)24 matches very well the
computationally predicted log β[Zr(DFOB)] value (41.20).32 Still,
in the [ZrH(DFOB)OH]+ complex, dominating at pH 6.5−
10.5, we have suggested the presence of an unbound
protonated amino group and a hydrolyzed water molecule
bound to Zr(IV). For the cyclic trihydroxamate ligand H3L1,
characterized by 9-bond linkers between its hydroxamate units
(Scheme 1), log β[ZrL1] = 34.8(2) is not far from the value
estimated for the trihydroxamate cyclic siderophore FSC,7,62
log β[Zr(FSC)] = 38.92;32 this difference can be ascribed to the
alterations in geometry and dimensions of the ligands.
Zirconium is known to form complexes with a complicated
geometry of a dodecahedron or square antiprism,63,64 and
numerous DFT studies have revealed that minor variations in
ligand geometry (such as a pendant arm elongation or a
modification of the ligand cavity size) could result in significant
changes in the stability of Zr(IV) complexes.39,65,66 H4L2
presented in this work possesses four hydroxamate units and
amide units in the linker, but a larger cavity size and a
significant asymmetry (coming from one much longer linker
and with two amides and amino group) with respect to
CTH36 (Scheme 1).32,39 These structural alterations are most
probably the reason for the lower thermodynamic stability of
H4L2 complexes, as the coordination sphere might not be
uniformly closed around the central ion. Unfortunately, the
thermodynamic characterization of Zr(IV)-CTH36 complexes
has not yey been reported; thus, the pZr value cannot be
quantified. Another cyclic hydroxamic ligand, PPDDFOT1,
that possesses four hydroxamic groups in a symmetrical
arrangement and a cavity size even larger than that of H4L2
(11 bonds between hydroxamic groups) showed superior
stability vs DFOB in EDTA challenging assays.67 These results
confirm that the Zr(IV) complex stability is strongly
According to the literature, the 89Zr radiolabeling strategies
for hydroxamate ligands is usually simple, robust, and relatively
rapid. They are performed under mild pH conditions, at room
temperature, and take around 1 h for DFOB derivatives on
their own without being attached to any targeting vectors and
1−3 h for DFO derivatives attached to targeting vectors, such
as trastuzumab.14,68,69 Cyclic hydroxamate ligands such as
C740 and CTH3639 (both with 8 bonds between hydroxamate
groups) have demonstrated excellent complexation abilities at
ambient temperature (>99% complexation after 120 min for
C7 and >90% of the activity within 5 min reaction time for
CTH36, respectively). The slightly smaller ligands C6 and C5,
reported by Guerard et al.,40 appear to be less suitable for
radiolabeling, with higher temperatures being required to
obtain high complexation yields. The cyclic ligands reported in
this paper possess a larger cavity size, with at least 9 bonds
between hydroxamate groups, which allows us to assume that
the radiolabeling process will be highly efficient and performed
under mild conditions. Preliminary radiolabeling of artificial
FOXE siderophores with 68Ga, represented here by FOXE 2−
5, was achieved after 10 min at room temperature with
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moderate yields and high specific activities of 68Ga.38 Further
characterization is ongoing.
The crude product was purified by column chromatography using
ethyl acetate/petroleum ether (from 1/4 to 1/1 by volume) as an
eluent mixture. Compound 4 was obtained as a slightly yellowish oil
(4.35 g, 76.6% yield). ESI-MS: calcd for C22H35N2O6, 423.53 [M +
H]+; found, 423.25 [M + H]+, 445.23 [M + Na]+, 867.47 [2M +
Na]+. TR = 21.11 min. 1H NMR (400 MHz, CDCl3): δ 7.44−7.32 (m,
5H), 4.80 (s, 2H), 4.11 (qd, J = 7.1, 2.9 Hz, 2H), 3.70 (t, J = 6.8 Hz,
2H), 3.12 (t, J = 6.6 Hz, 2H), 2.42 (t, J = 7.2 Hz, 2H), 2.32 (t, J = 7.3
Hz, 2H), 1.99−1.91 (m, 2H), 1.80−1.73 (m, 2H), 1.42 (s, 9H),
1.26−1.20 (m, 3H).13C NMR (CDCl3): δ 172.9, 156.0, 134.3, 129.2,
129.0, 128.7, 79.1, 60.4, 44.6, 40.3, 31.4, 29.6, 28.4, 24.7, 22.3, 14.2.
Synthesis of Ethyl 10,20-Bis(benzyloxy)-2,2-dimethyl-4,9,14,19-
tetraoxo-3-oxa-5,10,15,20-tetraazatetracosan-24-oate (7). The
Boc-deprotected derivative 5 (3.40 g, 7.8 mmol) was obtained as
previously described for 3. Compound 6 was synthesized by
dissolving the ethyl ester 4 (3.0 g, 7.1 mmol) in a 1,4-dioxane/H2O
mixture in the presence of LiOH (1 M aqueous solution, 12.5 mmol).
The mixture was stirred at rt for 20−30 min. Once the reaction was
complete, dioxane was evaporated and the crude product was acidified
using 1 M HCl to reach pH 6. Then, the aqueous phase was extracted
using ethyl acetate. Compound 6 (0.91 g, 2.31 mmol) was used in the
next step without further purification. The coupling reaction was
conducted as previously described for 4, and derivative 7 was
obtained as a yellowish oil (1.24 g, 77% yield) after column
chromatography. ESI-MS: calcd for C37H55N4O9, 699.87 [M + H]+;
EXPERIMENTAL SECTION
■
Synthesis. General Considerations. Unless stated otherwise, all
commercially available reagents and solvents were of analytical grade.
For the synthesis of H3L1 and H4L2, solvents and reagents were
purchased from Bachem, BLDpharm, and Fluka. Crude products were
purified via flash column chromatography on silica gel (Merck, 230−
400 Mesh) or, for compounds 10, 11, 16, and 17, by semipreparative
RP-HPLC using a Waters Prep 600 system equipped with a C18
Jupiter column (250 × 30 mm, 300 Å, 15 μm spherical particle size).
Gradients were established each time by considering the analytical
HPLC profile of the crude product. The column was perfused at a
flow rate of 20 mL/min over 30 min with a binary system of solvent A
(H2O + 0.1% v/v TFA) and solvent B (60% CH3CN in water +0.1%
v/v TFA). Analytical RP-HPLC analyses were performed on a
XBridge C18 column (4.6 × 150 mm, 5 μm particle size) using a flow
rate of 0.7 mL/min and a linear gradient of acetonitrile (and 0.1%
TFA) in water (and 0.1% TFA) from 0% to 100% over 25 min. The
mass spectra were recorded on an ESI-Micromass ZMD 2000
instrument. TLC was performed on precoated plates of silica gel F254
(Merck, Darmstadt, Germany). 1H NMR analyses were obtained
using a Varian spectrometer (400 MHz) and were referenced to
residual 1H signals of the deuterated solvents (δ(1H) 7.26 for CDCl3;
δ(1H) 2.50 for DMSO). The following abbreviations are used to
describe the shape of the peaks: s, singlet; d, doublet; dd, doublet of
doublets; t, triplet; m, multiplet.
1
found, 699.96 [M + H]+. TR = 21.18 min. H NMR (400 MHz,
CDCl3): δ 7.50−7.31 (m, 10H), 7.04 (bs, 1H), 4.81 (d, J = 5.3 Hz,
4H), 4.12 (q, J = 7.1 Hz, 2H), 3.71−3.69 (m, 4H), 3.26 (dd, J = 11.9,
6.2 Hz, 2H), 3.17−3.09 (m, 2H), 2.53−2.39 (m, 4H), 2.33 (t, J = 7.3
Hz, 2H), 2.20 (t, J = 6.8 Hz, 2H), 1.98−1.92 (m, 4H), 1.84−1.73 (m,
4H), 1.42 (s, 9H), 1.24 (t, J = 7.1 Hz, 3H). 13C NMR (CDCl3): δ
174.7, 173.3, 173.0, 134.2, 129.3, 129.1, 129.1, 128.8, 60.5, 44.7, 44.3,
40.0, 39.6, 33.1, 31.4, 30.0, 29.4, 28.5, 24.8, 23.9, 23.2, 22.3, 14.3.
Synthesis of Ethyl 10,20,30-Tris(benzyloxy)-2,2-dimethyl-
4,9,14,19,24,29-hexaoxo-3-oxa-5,10,15,20,25,30-hexaazatetratria-
contan-34-oate (9). Compound 9 was synthesized under the same
coupling conditions used for 4 by starting from the acid derivative 6
(0.91 g, 2.31 mmol) and the amino derivative 8 (1.81 g, 2.54 mmol).
The desired product was obtained as a colorless oil (1.89 g, 84%
yield) after column chromatography. ESI-MS: calcd for C52H75N6O12,
976.20 [M + H]+; found, 975.94 [M + H]+. TR = 17.83.1H NMR (400
MHz, CDCl3): δ 7.39−7.35 (m, 15H), 5.05 (bs, 3H), 4.83−4.77 (m,
6H), 4.11 (q, J = 7.1 Hz, 2H), 3.72−3.67 (m, 6H), 3.34−3.19 (m,
4H), 3.16−3.11 (m, 2H), 2.49−2.45 (m, 6H), 2.32 (t, J = 7.3 Hz,
2H), 2.25−2.12 (m, 4H), 2.02−1.89 (m, 6H), 1.87−1.72 (m, 6H),
1.42 (s, 9H), 1.24 (t, J = 7.1 Hz, 3H). 13C NMR (CDCl3): δ 175.3,
174.1, 172.5, 157.3, 135.4, 128.9, 128.3, 128.2, 80.7, 73.8, 61.2, 48.6,
41.2, 40.4, 35.2, 31.8, 31.5, 28.4, 24.1, 21.7, 18.5, 14.7.
Synthesis of tert-Butyl(benzyloxy)carbamate (1). To an ice-
cooled solution of O-benzylhydroxylamine·HCl (4.00 g, 25 mmol) in
a 1,4-dioxane/H2O mixture (60 mL, 1/1 v/v) was added K2CO3
(10.37 g, 75 mmol). Boc2O (8.18 g, 37.5 mmol), previously dissolved
in dioxane, was then added dropwise, and the reaction mixture was
stirred at room temperature overnight. The solvent was removed
under vacuum, and the crude product extracted using ethyl acetate
(30 mL) and water (3 × 15 mL). The organic phase was dried over
Na2SO4, filtered, and evaporated. Compound 1 (4.86 g, 87% yield)
was obtained as a colorless oil, which was used without any further
purification. NMR data match those reported in the literature (PMID:
11906271). ESI-MS: calcd for C12H18NO3, 224.28 [M + H]+; found,
224.13 [M + H]+. TR = 19.70 min.
Synthesis of Ethyl 4-((Benzyloxy)(tert-butoxycarbonyl)amino)-
butanoate (2). To a solution of 1 (4.86 g, 21.79 mmol) in DMF (15
mL) was added NaH (60% dispersion in mineral oil, 1.20 g, 23.94
mmol). The mixture was initially stirred at rt for 30 min, and then the
reaction mixture was warmed to 60 °C and ethyl 4-bromobutyrate
was added dropwise. At the completion of the reaction, the solvent
was removed, and the residue was extracted with ethyl acetate and
water (3 × 30 mL), dried over Na2SO4, and concentrated under
vacuum. The product (2) was obtained as a yellowish oil (5.36 g, 73%
yield). The NMR data correspond to those in the literature (PMID:
28715615). MS (ESI): calcd for C18H28NO5, 338.20 [M + H]+;
found, 360.18 [M + Na]+, 697.37 [2M + Na]+. TR = 23.51 min.
Synthesis of Ethyl 4-(N-(Benzyloxy)-4-((tert-butoxycarbonyl)-
amino)butanamido)butanoate (4). Compound 2 (5.36 g, 15.90
mmol) was dissolved in trifluoroacetic acid (TFA, 6 mL), and the
mixture was stirred at room temperature for 2 h. The reaction mixture
was monitored by MS (ESI) before being concentrated under
vacuum. The deprotected amino ester 3 was used without further
purification in the next step. To an ice-cold solution of Boc-γ-
aminobutiric acid (2.7 g, 13.45 mmol) in DMF (20 mL) were added
1-[bis(dimethylamino)methylene]-1H-1,2,3-triazolo[4,5-b]-
pyridinium 3-oxide hexafluorophosphate (HATU, 5.6 g, 14.75 mmol)
and DIPEA (2.6 mL, 14.75 mmol). A portion of 3 (3.5 g, 14.75
mmol) was dissolved in DMF (10 mL), and this solution was added
dropwise to the first one. Then the reaction mixture was warmed to
room temperature and stirred for 1 h. After removal of the solvent, the
residue was dissolved in ethyl acetate and washed with a 5% aqueous
solution of citric acid, a 10% aqueous solution of NaHCO3, and brine.
Synthesis of 1,11,21-Tris(benzyloxy)-1,6,11,16,21,26-hexaazacy-
clotriacontane-2,7,12,17,22,27-hexaone (10). Compound 9 was
Boc-deprotected as described for 3. Then, the ethyl group was
hydrolyzed by LiOH as for 6. To a dilute solution of the fully
deprotected trimer (1.92 g, 2.0 mmol) in DMF (100 mL) were added
HATU (0.84 g, 2.2 mmol) and DIPEA (0.38 mL, 2.2 mmol)
dropwise at 0 °C. The reaction mixture was stirred for 3 h. Then, the
solvent was removed, and the residue was extracted with ethyl acetate
and an aqueous solution of citric acid (10%), a solution of NaHCO3
(5%), and brine. The crude product was purified via semipreparative
HPLC, giving the desired product as a colorless oil (0.70 g, 42%
yield). ESI-MS: calcd for C45H61N6O9, 830.02 [M + H]+; found,
1
829.90 [M + H]+. TR = 21.47 min. H NMR (400 MHz, CDCl3): δ
7.39−7.31 (m, 18H), 4.77 (s, 6H), 3.69−3.67 (m, 6H), 3.25−3.23
(m, 6H), 2.46 (t, J = 6.7 Hz, 6H), 2.19 (t, J = 7.0 Hz, 6H), 1.97−1.91
(m, 6H), 1.85−1.70 (m, 6H). 13C NMR (CDCl3): δ 174.9, 173.8,
133.8, 129.3, 128.8, 44.2, 39.4, 33.0, 29.5, 23.7, 23.2.
Synthesis of 1,11,21-Trihydroxy-1,6,11,16,21,26-hexaazacyclo-
triacontane-2,7,12,17,22,27-hexaone (11, H3L1). To a solution of
the benzyl-protected derivative 10 (0.70 g, 0.84 mmol) in MeOH (30
mL) was added glacial acetic acid (1 mL). The mixture was treated
with a catalytic amount (0.084 mmol) of palladium on activated
K
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charcoal (10% Pd basis) under a hydrogen atmosphere. After 24 h,
the reaction mixture was filtered through Celite, concentrated under
reduced pressure, diluted with water, and alkalinized with saturated
sodium bicarbonate. The aqueous phase was extracted with ethyl
acetate (4 × 10 mL), and the combined organic layers were dried over
Na2SO4, filtered, and concentrated under vacuum. Compound 11 was
obtained as a light yellow oil after preparative HPLC purification
(0.43 g, 91% yield). ESI-MS: calcd for C24H43N6O9, 559.64 [M +
H]+; found, 559.79 [M + H]+. TR = 15.79 min. 1H NMR (400 MHz,
CDCl3): δ 9.59 (bs, 3H), 7.83−7.81 (m, 3H), 3.53−3.38 (m, 6H),
3.15−2.93 (m, 6H), 2.40−2.24 (m, 5H), 2.20−2.16 (m, 2H), 2.02 (t,
J = 7.1 Hz, 5H), 1.76−1.70 (m, 6H), 1.64−1.50 (m, 6H). 13C NMR
(CDCl3): δ 174.5, 173.0, 172.2, 169.5, 158.9, 158.5, 129.9, 50.5, 47.3,
47.1, 38.7, 38.6, 33.0, 31.5, 31.2, 30.3, 29.7, 24.9, 24.8, 23.0, 22.3,
19.7. HR-ESI-MS m/z 559.30894; calcd for C24H43N6O9 ([M + H]+)
559.30860. Anal. Calcd for C24H42N6O9: C, 51.6; H, 7.6; N, 15.0.
Found: C, 51.4; H, 7.5; N, 14.9.
Synthesis of Ethyl 17-(Benzyloxy)-10-(((benzyloxy)carbonyl)-
amino)-2,2-dimethyl-4,11,16-trioxo-3-oxa-5,12,17-triazahenico-
san-21-oate (13). Compound 13 was synthesized under the same
coupling conditions used for compounds 4 and 9 bystarting from Z-
Lys(Boc)-OH (1.05 g, 2.75 mmol) and the amino derivative 5 (1.09
g, 2.50 mmol). The desired product was obtained as a yellowish oil
(1.40 g, 80% yield) after column chromatography. ESI-MS: calcd for
C36H53N4O9, 685.84 [M + H]+; found, 685.73 [M + H]+. TR = 26.16
min. 1H NMR (400 MHz, CDCl3): δ 7.37−7.35 (m, 5H), 7.31−7.26
(m, 5H), 6.99 (s, 1H), 5.85 (d, J = 7.7 Hz, 1H), 5.13−4.96 (m, 2H),
4.76 (s, 2H), 4.07 (q, J = 7.1 Hz, 2H), 3.73−3.59 (m, 2H), 3.28−3.13
(m, 2H), 3.03−2.97 (m, 2H), 2.47−2.40 (m, 2H), 2.28 (t, J = 7.3 Hz,
2H), 1.94−1.88 (m, 2H), 1.84−1.69 (m, 3H), 1.62−1.58 (m, 1H),
1.39 (s, 11H), 1.31 (dd, J = 19.1, 12.2 Hz, 2H), 1.19 (t, J = 7.1 Hz,
3H). 13C NMR (CDCl3): δ 13C NMR (101 MHz, cdcl3) δ 174.3,
172.9, 172.0, 156.2, 136.2, 134.1, 129.2, 128.9, 128.6, 128.4, 128.0,
127.9, 78.9, 76.2, 66.8, 60.4, 54.8, 44.4, 39.9, 39.2, 38.5, 32.2, 31.2,
29.7, 29.4, 28.3, 23.7, 22.4, 22.1, 14.1.
Synthesis of Ethyl 17,27,37,47-Tetrakis(benzyloxy)-10-
( ( ( b e n z y l o x y ) c a r b o n y l ) a m i n o ) - 2 , 2 - d i m e t h y l -
4 , 1 1 , 1 6 , 2 1 , 2 6 , 3 1 , 3 6 , 4 1 , 4 6 - n o n a o x o - 3 - o x a -
5,12,17,22,27,32,37,42,47-nonaazahenpentacontan-51-oate (15).
The tetramer 15 was synthesized under the same coupling conditions
used for compounds 4, 9, and 13 by starting from the acid derivative
14 (0.61 g, 0.93 mmol) and the amino derivative 12 (0.90 g, 1.02
mmol). The desired product was obtained as a colorless oil (0.97 g,
69% yield) after column chromatography. ESI-MS: calcd for
C81H113N10O18, 1514.85 [M + H]+; found, 1514.16 [M + H]+,
775.95 [M + 2H]2+. TR = 25.88 min. 1H NMR (400 MHz, CDCl3): δ
7.47−7.28 (m, 25H), 5.13−4.96 (m, 2H), 4.86−4.65 (m, 8H), 4.09
(dd, J = 13.8, 6.8 Hz, 3H), 3.80−3.53 (m, 8H), 3.32−3.10 (m, 7H),
3.04−3.00 (m, 2H), 2.54−2.34 (m, 7H), 2.29 (t, J = 7.0 Hz, 2H),
2.23−2.05 (m, 6H), 1.96−1.90 (m, 8H), 1.82−1.68 (m, 8H), 1.66−
1.53 (m, 2H), 1.40 (s, 12H), 1.25−1.19 (m, 4H), 0.94−0.85 (m, 2H).
13C NMR (CDCl3): δ 174.5, 172.9, 136.2, 134.0, 129.2, 129.1, 128.8,
3H), 3.09−2.88 (m, 10H), 2.84−2.61 (m, 1H), 2.43−2.25 (m, 8H),
2.09−1.93 (m, 9H), 1.75−1.71 (m, 9H), 1.59−1.53 (m, 8H), 1.39−
1.08 (m, 7H), 1.07−0.76 (m, 2H). 13C NMR (DMSO-d6): δ 171.8,
156.4, 135.2, 129.8, 129.1, 128.9, 128.7, 128.2, 128.1, 75.7, 65.8, 55.1,
44.5, 38.5, 33.0, 29.5, 24.7, 23.2.
Synthesis of (S)-43-Amino-6,16,26,36-tetrahydroxy-
1,6,11,16,21,26,31,36,41-nonaazacycloheptatetracontane-
2,7,12,17,22,27,32,37,42-nonaone (17, H4L2). Compound 17 was
synthesized as previously described for 11 by starting from derivative
16 (0.13 g, 71% yield). ESI-MS: calcd for C38H69N10O13, 874.03 [M +
H]+; found, 873.75 [M + H]+. TR = 21.28 min. 1H NMR (400 MHz,
DMSO-d6): δ 9.64−9.59 (m, 3H), 8.07−8.04 (m, 2H), 7.86−7.71
(m, 4H), 3.47−3.44 (m, 9H), 3.04−2.99 (m, 10H), 2.42−2.22 (m,
8H), 2.04−2.00 (m, 8H), 1.76−1.49 (m, 20H), 1.47−1.16 (m, 5H).
13C NMR (DMSO-d6): δ 172.9, 172.2, 168.7, 158., 56.5, 52.7, 47.2,
38.7, 33.0, 31.2, 29.7, 29.1, 24.8, 24.5, 23.0, 22.1. HR-ESI-MS m/z
873.50488; calcd for C38H69N10O13 ([M + H]+) 873.50401. Anal.
Calcd for C38H68N10O13: C, 52.3; H, 7.9; N, 16.0. Found: C, 52.3; H,
7.8; N, 16.1.
Thermodynamic Solution Studies. General Considerations.
Unless otherwise stated, all commercially available reagents and
solvents were of analytical grade, were purchased from commercial
suppliers (Sigma-Aldrich, Titripur, Merck, Fisher Scientific, Fluka),
and were used as received without further purification. All solutions
were prepared in doubly distilled water. A stock solution of Fe(III)
was prepared immediately before use from Fe(ClO4)3·xH2O in 0.01
M HClO4 and standardized by an inductively coupled plasma−optical
emission spectrometer (ICP-OES; iCAP 7400 Duo ICP-OES) along
with spectrophotometric determination, on the basis of the molar
extinction coefficient ε = 4160 M−1 cm−1 at 240 nm.70,71 Stock
solutions of Ga(III) and Zr(IV) were prepared immediately before
use from Ga(ClO4)3·xH2O and anhydrous ZrCl4, respectively, in 0.1
M HClO4 to prevent hydrolysis and standardized by ICP-OES (iCAP
7400 Duo ICP-OES) along with direct titration with ethyl-
enediaminetetraacetic acid (EDTA).72,73 The HClO4 solutions were
titrated with standardized NaOH (0.1 N). The carbonate-free NaOH
solution was standardized by titration with potassium hydrogen
phthalate (KHP). All stock solutions were prepared using a R200D
Sartorius analytical balance (with 0.01 mg precision).
All measurements were performed at 0.1 M NaClO4 ionic strength,
which was chosen instead of 1.0 M NaClO4 ionic strength in order to
increase the solubility of the investigated ligands and their complexes.
We are aware that some measurements were performed at a very
acidic pH (<1), where the ionic strength 0.1 M is not enough to keep
the ionic activity stable, but due to the decomposition of hydroxamate
ligands in strong acids,43,51 all measurements performed below pH 1
were assumed to be endowed with a large error and (i) were not taken
into account during data evaluation or (ii) precluded from the
discussion.
Electrospray Ionization Mass Spectrometry (ESI-MS). ESI-MS
data were recorded on a Bruker Q-FTMS spectrometer. The
instrumental parameters were as follows: scan range, m/z 200−
1600; dry gas, nitrogen; temperature, 170 °C; capillary voltage, 4500
V; ion energy, 5 eV. The capillary voltage was optimized to the
highest signal to noise ratio. The spectra were recorded in the positive
mode. Compounds were dissolved in a MeOH/H2O solution (80/20
by weight); the same solvent mixture was used to dilute the matrix
solutions to the concentration range of 0.01 mM. The Fe(III),
Ga(III), and Zr(IV) and stock solutions were prepared as described
previously and added to the ligand solutions in 1/1, 2/1 and 1/3
mixtures for Fe(III) and Ga(III) and 1/1 and 1/3 mixtures for
Zr(IV), all at pH 3 (the pH was adjusted by using acetic acid). The
free hydrogen ion concentration was measured with a Mettler-Toledo
InLab Semi-Micro combined glass electrode filled with NaCl in
MeOH/H2O (80/20 by weight). Potential differences were measured
with a Beckman ϕ72 pH meter, standardized according to the classical
methods with buffers prepared according to reported procedures in
MeOH/H2O solvent (80/20 by weight).74,75
128.5, 128.2, 128.0, 67.0, 60.5, 55.0, 44.1, 39.4, 39.0, 32.9, 32.0, 31.3,
29.8, 29.7, 29.5, 28.4, 23.9, 23.0, 22.5, 22.2, 14.2.
Synthesis of Benzyl (6,16,26,36-Tetrakis(benzyloxy)-
2,7,12,17,22,27,32,37,42-nonaoxo-1,6,11,16,21,26,31,36,41-nonaa-
zacycloheptatetracontan-43-yl)carbamate (16). Compound 15 was
Boc-deprotected as described for 3. Then, the ethyl group was
hydrolyzed by LiOH as for 6. To a dilute solution of the fully
deprotected tetramer (0.55 g, 0.367 mmol) in DMF (40 mL) were
added HATU (0.154 g, 0.40 mmol) and DIPEA (0.07 mL, 0.40
mmol) dropwise at 0 °C. The reaction mixture was stirred for 3 h.
Then, the solvent was removed, and the residue was extracted with
ethyl acetate and an aqueous solution of citric acid (10%), a solution
of NaHCO3 (5%), and brine. The crude product was purified via
semipreparative HPLC, giving the desired product as a colorless oil
(0.29 g, 57% yield). ESI-MS: calcd for C74H99N10O15, 1368.66 [M +
H]+; found, 1368.36 [M + H]+, 684.74 [M+2H]2+. TR = 25.73 min.
1H NMR (400 MHz, DMSO-d6): δ 7.96−7.63 (m, 6H), 7.56−7.13
Potentiometric Titrations. The potentiometric titrations of ligands
and their complexes were carried out using a Titrando 905
(m, 25H), 5.06−4.89 (m, 2H), 4.85−4.71 (m, 8H), 4.00−3.76 (m,
L
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(Metrohm) automatic titrator system, equipped with a combined
glass electrode (Mettler Toledo, InLab Semi-Micro, with XEROLYT
EXTRA Polymer filling) and a 800 Dosino dosing system, equipped
with a a 2 mL micro buret. The ionic strength was fixed at I = 0.1 M
with NaClO4. The electrode was calibrated daily in terms of hydrogen
ion concentration using HClO4 (0.1 M) with CO2-free NaOH
solutions (0.1 M).76 A stream of high-purity argon, presaturated with
water vapor, was passed over the surface of the solution cell, the cell
was filled with 50 mL of the studied solution, and the system was
Spectrophotometric titrations were performed on samples with a
concentration of the ligand of ∼0.05−0.1 mM and I = 0.1 M
(completed by adding NaClO4), at 25.0 0.1 °C; pH 1.5 or 2.0 was
adjusted by adding the proper volume of HClO4. In general, the stock
solution of the starting complex (Fe(III)-L, Ga(III)-L, or Zr(IV)-L,
respectively) was divided into several aliquots to which an excess of
titrant (Ga(III), Fe(III), or Zr(IV)-NTA, respectively) was added.
After preparation, each solution was allowed to equilibrate for about 1
h, and then its UV−vis spectrum was recorded. The vials were kept in
the dark, and the absorbance was measured again after 24, 48, and 120
h. Changes were observed only between the spectra collected after 1
and 24 h, indicating that an equilibrium was attained.
thermostated at 25.0
0.2 °C. At least three titrations were
performed for each system, with a starting concentration of the ligand
of 1 mM and a 1:1 metal to ligand molar ratio with a 10% excess of
the ligand in the pH range 2−11. The purity and exact concentration
of the ligand solutions were determined using the Gran method.77
Special care was taken to ensure that complete equilibration was
attained. The titration curves were carefully checked and did not
display any pH fluctuations that often accompany the precipitation of
metal hydroxides. The potentiometric data were refined with the
SUPERQUAD78 and HYPERQUAD79 programs, which use nonlinear
least-squares methods. The successive protonation constants of the
ligand were calculated from the cumulative constants determined with
the program and defined by eqs 1 and 2 (charges are omitted for
clarity).
In order to determine the log β values of [GaHL]+ for H3L1 and
H3L4 and of [GaH2L2]+, competition experiments at pH 1.5 of (i)
Fe(III)-H3L + Ga(III) and (ii) Ga(III)-L + Fe(III) were performed.
For (i) 15 samples with a constant concentration of Fe(III) ions and
L (1:1) were titrated by up to 600 equiv of Ga(III) ions; for Ga(III)-L
+ Fe(III), 18 samples with a constant concentration of Ga(III) ions
and H3L (1:1) were titrated by up to 4 equiv of Fe(III) ions.
In order to determine the log β values of [ZrL]+ for H3L1, H4L2,
H3L3, H3L4, and DFOE and for [ZrHL2]+, competition experiments
at pH 1.5 or 2, Fe(III)-L + Zr(IV)-NTA, were performed. In each
experiment 18 samples with a constant concentration of Fe(III) ions
and ligands were titrated by up to 50 equiv of a Zr(IV)-NTA solution
with a metal to ligand molar ratio of 1:3, starting from 0 equiv.
Data Treatment. In the calculations of complex stability constants,
the protonation constants of free ligands (Table 1) and the constants
were related to hydrolytic species being taken into account: Ga(III),60
Hn−1L + H ⇆ HnL
(1)
[HnL]
[Hn−1L][H]
KnH
=
(2)
+
Ga(OH)2+ log βGaH = −3.11, Ga(OH)2 log βGaH = −7.66,
−1
−2
The stability constants calculated for metal complexes are defined
by eqs 3 and 4:
−
Ga(OH)3 log βGaH = −11.94, Ga(OH)4 logβGaH = −15.66, to
−3
−4
Fe(III),80 Fe(OH)2+ log βFeH = −2.56, Fe(OH)2 log βFeH = −6.2,
+
−1
−2
pM + qH + rL ⇆ MpHqLr
(3)
−
Fe(OH)3 log βFeH = −11.44, Fe(OH)4 log βFeH = −21.88,
−3
−4
60
Fe(OH)5 log βFeH = −2.74, Fe(OH)63−, and Zr(IV), Zr(OH)3+
2−
[MpHqLr]
−5
βM H L =
q r
[M]p[H]q[L]r
(4)
log βZrH = −0.56, Zr(OH)22+ log βZrH = −1.44, Zr(OH)4 log βZrH
p
−1
−2
−4
2−
8+
= −8.85, Zr(OH)6 log βZrH = −30.6, Zr3(OH)4 log βZr H
=
−6
3
−4
The uncertainties in the log K values correspond to the added
standard deviations in the cumulative constants.
−6.96, Zr4(OH)88+ log βZr H = 6.52, Zr3(OH)93+ log βZr H = 12.19.
4
−8
3
−9
The Zr(IV) hydrolysis constants for the species Zr(OH)3+, Zr-
pH-Dependent UV−Vis Titrations. The pH-dependent UV−vis
spectrophotometric experiments for the Fe(III)-H3L1, Fe(III)-H4L2,
Ga(III)-H3L1, Ga(III)-H4L2, Ga(III)-H3L4, Zr(IV)-H3L1, Zr(IV)-
H4L2, Zr(IV)-H3L3, and Zr(IV)-H3L4 systems were carried out as a
function of concentration with a Varian Cary 300 Bio spectropho-
tometer in the 300−700 nm range for iron complexes and 200−300
nm range for gallium and zirconium solutions using Hellma quartz
optical cells with a 1 cm path length. To calculate the stability
constants for investigated systems, two sets of pH-dependent UV−vis
titrations were carried out: in the pH ranges (i) 0.1−2 and (ii) 2−11.
In the (i) series, the experiments were performed by making 20
samples, differing by 0.1 pH unit, with a constant total volume of 0.7
mL and concentration of metal ion of ∼0.05−0.1 mM and metal to
ligand molar ratio of 1:1; for all samples, the ionic strength was
adjusted to 0.1 M by the addition of NaClO4 and the pH (range 0.1−
2.0) was controlled by the concentration of HClO4. After preparation,
each solution was allowed to equilibrate for about 1 h, and then its
UV−vis spectrum was recorded. This was necessary to minimize the
effects of hydroxamate ligand hydrolysis, which occurs in strong
acid.43,51 In the (ii) set of experiments, 3 mL of a solution containing
a 1:1 Fe(III):ligand molar ratio, where the ferric concentration was
around 0.20 mM, was introduced into a cell and the pH was adjusted
by adding the proper microvolume of HClO4; the solutions were
allowed to equilibrate (up to 30 min) and checked with a Mettler
Toledo Super Easy pH meter with an accuracy of 0.01, and then the
spectra were recorded.
(OH)22+, Zr(OH)4, and Zr3(OH)4 were recalculated for 0.1 M
8+
NaClO4 ionic strength according to literature parameters.60 The pKw
value used in the calculation at the 0.1 M NaClO4 ionic strength was
−13.77.81
The UV−vis data were refined to obtain the overall binding
constant using SPECFIT/32 software82−84 that adjusts the
absorptivity and the stability constants of the species formed at
equilibrium. Specfit uses factor analysis to reduce the absorbance
matrix and to extract the eigenvalues prior to the multiwavelength fit
of the reduced data set according to the Marquardt algorithm.82−84
Uncertainties in log β were calculated from the standard deviation.
The competition data were refined to obtain the overall binding
constant using SPECFIT/32 software.82−84 The protonation con-
stants of ligands and formation constants for iron complexes (Table 1
and the literature24) were used as fixed parameters during data
analysis. The concentration of iron complexes was calculated from the
absorbance spectra (collected in the 300−700 nm range). Hydrolytic
forms of the ferric ion in the studied pH range are characterized by an
absorption band with a λmax value of below 300 nm, and therefore
they are beyond the experimental wavelength window. However, the
spectrum of Fe(III) in solution at the pH of the experiment was fixed
in the calculations.85,86 The stability constants of the Fe(III)-NTA48
and Zr(IV)-NTA27 complexes were taken from the literature and
were used as fixed constants during the evaluation of the stability
constants of the zirconium complexes.
Metal Competition Batch UV−Vis Titrations. In order to calculate
the stability constants of the investigated complexes, several
competitive titrations were performed. All of them were carried out
as a function of concentration with a Varian Cary 300 Bio
spectrophotometer in the 300−650 nm range (with 1 nm precision)
using Hellma quartz optical cells with a 1 cm path length.
The competition equilibrium is described by eqs 5 and 6:
(5)
FeL + ZrNTA V ZrL + FeNTA
[FeL][ZrNTA]
K =
[FeNTA][ZrL]
(6)
M
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Article
The molecular charges are omitted for clarity. The data were
processed using Origin 7.0. The species distribution diagrams were
computed with the HYSS program.79
2015/19/B/ST5/00413) for financial support. A.M. was
supported by the NCN (UMO-2017/26/A/ST5/00363).
REFERENCES
■
ASSOCIATED CONTENT
* Supporting Information
The Supporting Information is available free of charge at
https://pubs.acs.org/doi/10.1021/acs.inorgchem.1c01622.
■
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AUTHOR INFORMATION
Corresponding Author
Elzbieta Gumienna-Kontecka − University of Wroclaw,
Faculty of Chemistry, 50-383 Wrocław, Poland;
orcid.org/0000-0002-9556-6378;
■
Email: elzbieta.gumienna-kontecka@chem.uni.wroc.pl
Authors
Yuliya Toporivska − University of Wroclaw, Faculty of
Chemistry, 50-383 Wrocław, Poland
Andrzej Mular − University of Wroclaw, Faculty of Chemistry,
50-383 Wrocław, Poland
Karolina Piasta − University of Wroclaw, Faculty of
Chemistry, 50-383 Wrocław, Poland; orcid.org/0000-
0003-0160-6920
Małgorzata Ostrowska − University of Wroclaw, Faculty of
Chemistry, 50-383 Wrocław, Poland
Davide Illuminati − University of Ferrara, Dipartimento di
Scienze Chimiche, Farmaceutiche ed Agrarie, 44121 Ferrara,
Italy
Andrea Baldi − University of Ferrara, Dipartimento di Scienze
Chimiche, Farmaceutiche ed Agrarie, 44121 Ferrara, Italy
Valentina Albanese − University of Ferrara, Dipartimento di
Scienze Chimiche, Farmaceutiche ed Agrarie, 44121 Ferrara,
Italy; orcid.org/0000-0002-1947-2644
Salvatore Pacifico − University of Ferrara, Dipartimento di
Scienze Chimiche, Farmaceutiche ed Agrarie, 44121 Ferrara,
Italy
Igor O. Fritsky − Taras Shevchenko National University of
Kyiv, Department of Chemistry, 01601 Kyiv, Ukraine;
orcid.org/0000-0002-1092-8035
Maurizio Remelli − University of Ferrara, Dipartimento di
Scienze Chimiche, Farmaceutiche ed Agrarie, 44121 Ferrara,
Italy
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Complete contact information is available at:
https://pubs.acs.org/10.1021/acs.inorgchem.1c01622
Notes
The authors declare no competing financial interest.
ACKNOWLEDGMENTS
■
This contribution is based upon work from COST Action
CA18202, NECTAR−Network for Equilibria and Chemical
Thermodynamics Advanced Research, supported by COST
(European Cooperation in Science and Technology). We
acknowledge the Polish National Science Centre (NCN, UMO
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Larson, S. M.; Lewis, J. S. Zr-89-DFO-J591 for ImmunoPET of
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