Implementing f-block metal ions in medicine: tuning the size selectivity of expanded macrocycles

Article abstract of DOI:10.1021/acs.inorgchem.9b01277

The f-block elements, which comprise both the lanthanide and actinide series, possess interesting spectroscopic, magnetic, and nuclear properties that make them uniquely suited for a range of biomedical applications. In this Forum Article, we provide a co

Full text of DOI:10.1021/acs.inorgchem.9b01277

Forum Article
pubs.acs.org/IC
Cite This: Inorg. Chem. XXXX, XXX, XXX−XXX
Implementing f‑Block Metal Ions in Medicine: Tuning the Size
Selectivity of Expanded Macrocycles
Nikki A. Thiele,†,§ Joshua J. Woods,†,‡,§ and Justin J. Wilson*,†
†Department of Chemistry and Chemical Biology and ‡Robert F. Smith School for Chemical and Biomolecular Engineering, Cornell
University, Ithaca, New York 14853, United States
S
* Supporting Information
ABSTRACT: The f-block elements, which comprise both the lanthanide and actinide series, possess interesting spectroscopic,
magnetic, and nuclear properties that make them uniquely suited for a range of biomedical applications. In this Forum Article,
we provide a concise overview on the different ways that these elements are employed in medicine, highlighting their dual
implementation in both diagnostic and therapeutic applications. A key requirement for the use of these labile metal ions in
medicine is a suitable chelating agent that controls their in vivo biodistribution. Toward this goal, we also report our research
describing the synthesis and characterization of a rigid 18-membered macrocycle called CHX-macropa, an analogue of the
previously reported nonrigid ligand macropa (J. Am. Chem. Soc. 2009, 131, 3331). The lanthanide coordination chemistry of
CHX-macropa is explored in detail by pH potentiometry and density functional theory (DFT) calculations. These studies reveal
that CHX-macropa exhibits an enhanced thermodynamic selectivity for large over small lanthanides in comparison to its
nonrigid analogue macropa. DFT calculations suggest that a key factor in the enhanced selectivity of this ligand for the large f-
block ions is its rigid macrocyclic core, which cannot adequately distort to interact effectively with small ions. On the basis of its
high affinity for large f-block ions, the design strategies implemented in CHX-macropa may be valuable for applying these
elements in the diagnosis or treatment of disease.
including cancer, heart disease, infection, arthritis, and
gastrointestinal disease.6,7
INTRODUCTION
■
The U.S. Food and Drug Administration (FDA) approval of
the simple coordination compound cis-diamminedichlorido-
platinum(II), or cisplatin, in 1978 for the treatment of cancer
was a significant milestone in the field of medicinal
chemistry.1,2 The demonstration of the efficacy of this
compound in a wide range of cancer types was key in
supporting the idea that metal-containing compounds may
have significant value in the development of new medicines.3,4
Since then, the field of metals in medicine has dramatically
expanded. Researchers have capitalized on the unique chemical
and physical properties afforded by metal centers that are
otherwise inaccessible by conventional organic molecules to
develop an extensive library of novel medicinal agents
comprising elements spanning the Periodic Table.5 Under-
scoring the diverse utility of metals in medicine, dozens of
metallodrugs are currently employed in the diagnosis, imaging,
and treatment of a wide range of human diseases and ailments,
Lanthanides in Medicine. Among the different classes of
metal ions in the Periodic Table, lanthanides (Ln), comprising
the 15 4f elements, have played a central role in both
diagnostic and therapeutic medicine.8,9 Although lanthanides
have no endogenous biological function in humans,10 their
trivalent ions can act as Ca2+ ion mimics by virtue of their
similar ionic radii.11 As such, the Ln ions may be useful for the
treatment of bone resorption disorders,12 where they behave as
surrogates for Ca2+ in the hydroxyapatite-like matrix of
bone.13,14 Notably, the largest lanthanide, lanthanum, was
efficiently incorporated into bone in preclinical studies,
Special Issue: Celebrating the Year of the Periodic Table: Emerging
Investigators in Inorganic Chemistry
Received: May 1, 2019
© XXXX American Chemical Society
A
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Table 1. Clinically Relevant Radioisotopes of f-Block Elements and Their Properties
isotope
decay mode(s)
half-life (days)
application
clinical status
ref
134La
135La
134Ce
149Pm
153Sm
149Tb
152Tb
155Tb
161Tb
166Ho
170Tm
177Lu
225Ac
227Th
230U
β+
4.48 × 10−3
therapy/imaging
therapy
a
EC
0.79
52
EC
β−
β−
α/β+
β+
3.16
therapy
41
2.21
therapy
preclinical
approved (1997)
preclinical
preclinical
preclinical
preclinical
phase I/II
preclinical
approved (2018)
phase I
44, 46
46, 48
1.95
therapy
0.17
therapy/imaging
imaging
imaging
therapy
42, 43, 46, 62, 63
42, 46
0.73
EC
β−
β−
β−
β−
α
5.32
42, 46
6.88
42, 44, 46, 51
44−46
46, 47
1.12
therapy
128.63
6.65
therapy
therapy
44, 46, 49
64−68, 77, 78, 80−82
68−76, 79
84, 85
9.92
therapy
α
18.69
20.83
0.97 × 103
0.84
therapy
phase I
α
therapy
b
252Cf
255Fm
SF
therapy
phase II
89−91
86−88
α
therapy
a
b
EC, electron capture. SF, spontaneous fission.
demonstrating the potential use of this ion in diseases such as
osteoporosis.15−17 Furthermore, La3+ is a key component of
the drug Fosrenol [La2(CO3)3] for the treatment of hyper-
phosphatemia, a serious condition that is caused by elevated
levels of phosphate in the blood. Administered as a simple
carbonate salt, La3+ treats hyperphosphatemia by forming an
insoluble phosphate salt that is cleared through the gastro-
intestinal tract, preventing its absorption into the body.18 The
FDA approval of this drug in 2004 highlights how a
fundamental property of the lanthanides, namely, their poor
aqueous solubility upon complexing phosphate, can be
exploited in medicine.
Similarly, the unique magnetic and spectroscopic properties
of lanthanides have made them indispensable in the field of
imaging.19 The most well-known Ln ion in medicine, Gd3+, is
used extensively in magnetic resonance imaging contrast
agents; notably, all eight of the T1-based contrast agents
currently approved for use in the United States contain Gd3+ as
the metal center.20−22 Owing to the S = 7/2 ground state of this
ion afforded by its seven unpaired 4f electrons, Gd3+ enhances
T1 or the longitudinal relaxation rate of the 1H nuclei of water
surrounding tissues of interest, enabling high-resolution
imaging of internal body structures such as cancerous lesions.
Additionally, the lanthanides possess narrow emission energies,
long luminescence lifetimes, and large Stokes shifts that arise
from Laporte forbidden f−f electronic transitions.23−28
Researchers have taken advantage of these unique photo-
physical properties to develop luminescent lanthanide-based
imaging agents for subcellular structure,29−31 bone integ-
rity,32−34 and tumor morphology.35,36
imaging agents for PET or SPECT include 149Tb, 152Tb, 155Tb,
and 134La.41−43
The therapeutic applications of f-block ions can be achieved
by using radioisotopes that emit β particles, α particles, or
Auger electrons. β particles have the same mass and charge as
electrons and are ejected with high energy (0.5−2.3 MeV)
from a decaying nucleus. A consequence of their high energy
and low mass is a relatively high biological penetration range of
these particles of up to 12 mm, which makes them useful for
treating large solid tumors.37 β-particle-emitting isotopes 177Lu,
149Pm, 153Sm, 161Tb, and 166Ho have shown promise for the
treatment of solid tumors in a wide range of cancers, with
several drug candidates reaching clinical trials.42,44−46
Preclinical studies involving the β emitter 170Tm have
demonstrated that 170Tm-ethylenediaminetetramethylene
phosphonic acid (170Tm-EDTMP) is a cost-effective alter-
native to 89SrCl2 for bone pain palliation associated with bone
metastases.47 The clinical utility of the β-emitting radio-
lanthanides has been further validated by the FDA approval of
153Sm-lexidronam (Quadramet) in 1997 and 177Lu-dotatate
(Lutathera) in 2018 for the treatment of pain associated with
bone metastases48 and gastroenteropancreatic neuroendocrine
tumors,49 respectively.
The emission of Auger electrons occurs after radioactive
decay by either electron capture or internal conversion. These
decay modes generate vacancies in the 1s, 2s, and 2p core
atomic orbitals. The relaxation of these high-energy core hole
states occurs when an electron in a higher-energy orbital
transitions downward to an orbital of lower energy, ejecting
another electron in the process. These electron cascades
release cytotoxic amounts of highly localized radiation that can
cause significant damage to DNA in the cell nucleus.50 By
virtue of their low energy, Auger electrons have a very small
biological penetration range that can be harnessed for precise
intracellular targeting. Within the Ln series, 134La, 135La, 134Ce,
and 161Tb have been identified as potential Auger emitters for
therapeutic use.41,51,52 At this point, however, there are no
preclinical studies describing the full cytotoxic efficacy of these
radionuclides.
Lanthanides also possess a number of radioisotopes that are
suitable for diagnostic and therapeutic applications in nuclear
medicine (Table 1).37−40 For diagnosis, radioisotopes that
emit β+ particles, or positrons, can be leveraged for use in
positron emission tomography (PET) imaging. The emitted
positrons rapidly annihilate with electrons in the surrounding
tissue to release two γ photons that can be detected by imaging
cameras. Single-photon emission tomography (SPECT)
imaging can be carried out using radioisotopes that emit a
single γ photon. Radiolanthanides that have been explored as
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Chart 1. Structures of the Ligands Discussed in This Work
α-Particle-emitting radionuclides have garnered significant
interest for their use in targeted α therapy (TAT)53−55 based
on several advantages that they possess over conventional β
emitters.56−58 Compared to the high penetration range of β
particles, the significantly more massive α particles can only
travel up to 100 μm in tissue. Despite this short range, α
particles possess a high linear energy transfer and are quite
effective at inducing cell death through DNA double-strand
breaks.59−61 These properties make α-emitting isotopes
favorable for the treatment of small tumors for which there
is high risk for nonspecific irradiation of nearby healthy tissue.
Within the Ln series, 149Tb is currently the only radionuclide
that has been evaluated as an α emitter for clinical use.42,43,62,63
Actinides in Medicine. Moving further down the Periodic
Table, the radioactive actinides (An) are most commonly
considered in the context of nuclear energy or weapons.
However, their radioactivity can also be harnessed for nuclear
medicine (Table 1). More specifically, the α-emitting radio-
nuclides 225Ac64−67 and 227Th68−76 are currently under intense
investigation for use in TAT. In conjunction with different
biological targeting vectors, these An radionuclides have
progressed to clinical trials for the treatment of a range of
cancers, including metastatic castration-resistant prostate
cancer,77−79 acute myeloid leukemia,80,81 neuroendocrine
tumors,82 and tumors expressing mesothelin.83 Additionally,
a radioisotope of uranium, 230U, has been identified as a
promising candidate for TAT based, in part, on its relatively
short half-life of 20.83 days.84,85 Because the most stable form
2+
of uranium in oxygenated aqueous solution is the linear UO2
ion, the implementation of 230U in medicine will require the
development of specifically tailored chelating agents.
Lastly, there may also be applications in medicine for the
late synthetic actinides. For example, fermium, a man-made
element that was first discovered in the debris of the
detonation of the first hydrogen bomb,86,87 has a radioisotope,
255Fm, with suitable properties for use in TAT.88 Although no
serious efforts have been undertaken yet to evaluate this
radionuclide for its therapeutic properties, it does highlight the
intriguing possibility of using these exotic man-made elements
to benefit society in the biomedical realm. Another trans-
plutonium element that may have therapeutic value is
californium. Specifically, 252Cf, a radioisotope that was also
discovered in the debris from the first thermonuclear test,87 has
been investigated for the treatment of a range of cancer types.
252Cf undergoes decay by spontaneous fission (3.1%
probability), rendering this radionuclide a strong source of
neutrons that can be leveraged for brachytherapy, or internal
radiation therapy.89 It is anticipated that the promising clinical
efficacy of 252Cf demonstrated to date90,91 will spark future
research efforts directed toward the more widespread
implementation of this synthetic actinide in medicine.
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Scheme 1. Synthesis of CHX-macropa
Ligand Design for f-Block Metals. To leverage the
distinct properties of the 4f and 5f metal ions for medical
applications, they need to be administered in a chelated form
in vivo. Under most circumstances, chelators for these ions will
be conjugated to a biological targeting vector, such as an
antibody or peptide, that can deliver them to the desired
locations in vivo. For these types of constructs to be effective,
the chelating agents need to be appropriately matched to the
ion of interest.92 Generally, these ligands must possess two key
features to be suitable for in vivo applications. First, the
chelator should rapidly complex the metal center under mild
conditions of pH and temperature. Fast metal-binding kinetics
is particularly essential for applications that employ short-lived
radioactive f-block ions, which can undergo considerable decay
during the preparation of the radiopharmaceutical agent. The
ability to chelate under mild conditions is also beneficial
because the biological targeting vectors that are often used will
decompose under conditions of extreme pH and temperature.
A second requirement for these chelators is that they form
complexes of high stability. Stable chelation in vivo is
imperative because it will prevent the redistribution of these
nonendogenous metal ions throughout the patient and the
resulting toxicity from this off-target localization.
A key challenge in designing chelating agents for the
biological use of the f-block elements resides in the general
lability of these ions that results from the electrostatic nature of
their bonding interactions. In particular, effective chelation
strategies for large f-block metal ions such as La3+ and Ac3+ are
scarce owing to their low charge density, which gives rise to
longer and weaker metal−donor bonds that serve to destabilize
the resulting complexes.66 As such, the “normal” trend
observed for the thermodynamic stability of f-block complexes
is one that is inversely related to the ionic radius, where smaller
ions form more stable ionic bonds with chelators than their
larger counterparts. Thus, few options are available for the
chelation of large lanthanides and actinides, a limitation that
has hindered their implementation in medicine. The most
widely used chelator for binding large 4f and 5f ions is the
tetraazamacrocycle 1,4,7,10-tetraazacyclododecane-1,4,7,10-
tetraacetic acid (DOTA; Chart 1). Although DOTA can
adequately chelate the majority of trivalent lanthanides and
actinides of medical interest, the stability of its complexes
decreases significantly for larger ions,93−96 like most conven-
tional ligands. Given the lack of well-defined systems for the
chelation of large f-block metal ions, there exists a critical need
for new ligands with high affinity for large metals in order to
capitalize on the valuable properties of these ions in medicine.
Tuning Size-Selectivity Patterns. In ongoing efforts to
develop ligands that form stable complexes with large metal
ions, we have investigated the coordination properties of
macrocycles based on the 1,7,10,16-tetraoxa-4,13-diazacyclooc-
tadecane (D18C6) core.97,98 This 18-membered macrocycle
(Chart 1) gives rise to ligands that exhibit an unusual
thermodynamic preference for large over small metal
ions.99−108 As discussed above, this trend is in contrast to
the majority of conventional poly(aminocarboxylate) ligands
that form more stable complexes with smaller ions. Notably,
our studies have revealed that macropa (Chart 1), a D18C6-
based macrocycle possessing two picolinate pendent arms, can
rapidly form a stable complex with 225Ac3+, the largest 3+ ion
in the Periodic Table.97 Leveraging the therapeutic properties
of this α-particle-emitting radionuclide (Table 1), we recently
reported that a single dose of 225Ac-RPS-074, an 225Ac−
macropa bifunctional construct that targets prostate cancer,
induces complete tumor ablation in mice.109 In addition to
macropa’s remarkable Ac-chelating properties, we and others
have demonstrated its ability to selectively bind large Ln,110
alkaline earth,98,111 and p-block111 metal ions over their smaller
congeners. The unusual reverse-size selectivity of macropa fills
an important niche in medicine, where bifunctional chelators
for large metal ions are greatly needed.
On the basis of the success of macropa for chelating large
metal ions of biomedical relevance, we sought to explore how
further structural modifications of this ligand could alter or
improve its reverse-size selectivity. In this Forum Article, we
describe these efforts with a report on the synthesis and Ln-
chelating properties of the novel ligand CHX-macropa, an
analogue of macropa in which the ethylene bridge connecting
two oxygen donors in the D18C6 backbone is replaced with a
rigid cyclohexylene bridge. Our detailed thermodynamic
evaluation of CHX-macropa with the Ln series reveals that
complexes of this ligand are progressively destabilized relative
to complexes of macropa as the ionic radius of the Ln ion
decreases. Consequently, CHX-macropa is substantially more
selective than macropa for large lanthanides. The under-
pinnings of this size-dependent destabilization were probed by
density functional theory (DFT). Collectively, our studies
show that the conformational rigidity in the D18C6 backbone
can significantly influence f-block ion-selectivity patterns and
be applied for the development of novel expanded macrocycles
for use in medicine.
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carboxylic acid group of the corresponding picolinate arm
attached to the nitrogen atom. Crystallographic parameters
and interatomic distances and angles for [H2CHX-macropa]-
Cl2 are provided in Tables S1−S3.
RESULTS AND DISCUSSION
■
Synthesis and Characterization of CHX-macropa. The
synthesis of the decadentate macrocycle CHX-macropa was
accomplished in six steps (Scheme 1). trans-1,2-Cyclo-
hexanediyldioxydiacetic acid (1) was prepared from the
alkylation of trans-1,2-cyclohexanediol with bromoacetic acid
and then converted to the corresponding acyl chloride (2)
according to established literature procedures.112,113 Cycliza-
tion of 2 with commercially available 2,2′-(ethylenedioxy)bis-
(ethylamine) in benzene gave the diamide macrocycle 3 in
good yield. Subsequent reduction of the amide groups using
lithium aluminum hydride (LAH) in tetrahydrofuran (THF)
afforded the novel trans-cyclohexyl-D18C6 macrocyclic scaf-
fold (4). Picolinate groups were appended to the secondary
amine nitrogen atoms of the macrocycle via alkylation with 6-
(bromomethyl)pyridine-2-carboxylic acid methyl ester in
acetonitrile (ACN) at reflux. Following hydrolysis of the
methyl ester groups in 6 M HCl at 90 °C and purification of
the product by semipreparative high-performance liquid
chromatography (HPLC), CHX-macropa was isolated as a
Thermodynamic Stability Studies. Increasing the
conformational rigidity of a ligand can have a marked effect
on its affinity for metal ions, giving rise to complexes of greater
thermodynamic stability compared to complexes of the parent
ligand.114,115 A classic example of the successful implementa-
tion of this principle is the acyclic chelator trans-(1,2-
diaminocyclohexane)tetraacetic acid (CDTA), an analogue of
ethylenediaminetetraacetic acid (EDTA), where the ethyl-
enediamine component is replaced with a rigid trans-
diaminocyclohexane bridge (Chart 1). The stability constants
of complexes of CDTA with most metal ions, including the
lanthanides and actinides, are 1−5 orders of magnitude higher
than those of EDTA.116,117 This enhanced complex stability is
attributed to the preorganization conferred by the cyclohexyl
group, which locks the nitrogen atoms of the ligand in the
gauche conformation that is required for metal chelation.
Accordingly, the thermodynamics of complex formation are
more favorable for CDTA than for EDTA, which has to
undergo a configurational rearrangement to attain the correct
geometry for metal binding.118
crystalline hydrochloride salt and characterized by H and 13C
NMR spectroscopy, mass spectrometry, analytical HPLC, and
elemental analysis (Figures S1−S11, in the Supporting
Information, SI).
1
Despite many examples in the literature of constrained
macrocycles,119,120 few efforts to date have been directed at
decreasing the conformational mobility of ligands containing
the D18C6 scaffold to modulate metal-binding affinity.
Notably, the ligands (2B, 2),121 in which one benzene ring is
fused into the D18C6 backbone, and CY2-K22,122 which has
two rigid 2-hydroxycyclohexyl arms appended to D18C6,
exhibit increased affinity for small metal ions and decreased
affinity for large metal ions compared to their respective
nonrigid parent ligands D18C6 and BHE-K22 (Chart 1).
Because of this divergence in affinity, both conformationally
constrained ligands have augmented size selectivity for small
metal ions. Other rigidified D18C6 ligands have been
synthesized (Chart 1), but their coordination chemistry has
not been evaluated in detail.112,123−125 These limited studies
underscore the need for further exploration of the effects of
conformational flexibility on the size selectivity of this class of
macrocycles.
Single crystals of CHX-macropa suitable for X-ray diffraction
were grown by vapor diffusion of diethyl ether into a methanol
solution of the ligand. CHX-macropa crystallizes in the
centrosymmetric P21/c space group, which requires that both
enantiomers of this ligand be present. This result is consistent
with the use of racemic trans-1,2-cyclohexanediol as a starting
reagent for the synthesis of the ligand. The asymmetric unit,
shown in Figure 1, comprises the ligand dication and two
To systematically probe the thermodynamic consequences
of reinforcing the bridge of one chelate ring of macropa’s
D18C6 scaffold with a fused cyclohexyl group, we determined
the protonation constants of CHX-macropa and the stability
constants of its metal complexes with the entire Ln series
(except Pm) using pH potentiometry. The results are compiled
in Table 2. The protonation constants and Ln stability
constants previously reported for macropa are provided for
comparison.110 From the titration of CHX-macropa in the
absence of metal ions, four protonation constants (log Kai; eq
1) were determined over a pH range of 2.3−11. These values
correspond to the sequential protonation of the two amine
nitrogen atoms of the macrocycle (log Ka1 and log Ka2) and the
two pyridinecarboxylate pendent groups (log Ka3 and log Ka4).
The fifth and sixth ligand protonation constants are out of the
range of measurement for our potentiometric system (log Ka <
2) and therefore could not be determined. A comparison of the
ligand protonation constants reveals that the macrocyclic
nitrogen atoms of CHX-macropa are slightly more basic
relative to macropa; the first protonation constant is increased
by 0.36 log K units, and the second protonation constant is
Figure 1. X-ray crystal structure of [H2CHX-macropa]Cl2. Hydrogen
atoms attached to carbon centers are omitted for clarity. Thermal
ellipsoids are drawn at the 50% probability level. One chloride
counteranion (Cl2) is disordered over two positions and is shown at
the highest occupied position (occupancy = 0.82). Dotted lines
represent hydrogen-bonding interactions.
chloride counteranions. The picolinic acid pendent arms of
CHX-macropa are situated on opposite sides with respect to
the plane of the cyclohexyl-D18C6 ring, which adopts a slightly
puckered conformation. The cyclohexyl group of the macro-
cycle backbone achieves the chair conformation, the most
stable configuration for cyclohexane. Each chloride anion
engages in a hydrogen-bonding interaction with a protonated
nitrogen atom of the macrocyclic ring, as well as with the
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preventing the accurate determination of log KMHL. Removing
the MHL species from the model did not significantly change
the values of log KML. Therefore, in contrast to macropa, log
KMHL values have not been reported for CHX-macropa
because of our uncertainties regarding the significance of this
equilibrium.
Table 2. Protonation Constants of CHX-macropa and
a
Macropa and Thermodynamic Stability Constants of Their
Ln3+ Complexes Determined by pH Potentiometry (25 °C
and I = 0.1 M KCl)
b
CHX-macropa2−
macropa2−
c
log Ka1
log Ka2
log Ka3
log Ka4
7.77(1)
7.41
6.85
3.32
In comparison to macropa, the stability constants of the
lanthanide complexes of CHX-macropa are consistently lower
across this series. Despite this general reduction of the affinity
for the lanthanides, CHX-macropa retains selectivity for large
over small lanthanides, with stability constants for the large
lanthanides approaching those of macropa (Table 2 and Figure
2a). For example, the log KML value of 14.54 measured for La−
CHX-macropa is similar to that of 14.99 for La−macropa.
Although the stability constants of macropa and CHX-macropa
are comparable for the early lanthanides, these values become
more disparate for the mid and late lanthanides, which possess
increasingly smaller ionic radii (Figure 2a). Notably, the
stability constant of the complex of CHX-macropa with Lu3+,
the smallest lanthanide ion, is 1.55 log K units lower than that
of Lu−macropa. A plot of Δlog KML, the difference in the
stability constants between macropa and CHX-macropa, versus
the ionic radii of the lanthanides highlights the differences
between the size selectivity of these ligands (Figure 2b).
Specifically, this plot reveals that complexes of CHX-macropa
are increasingly destabilized relative to their macropa
analogues as the ionic radius of the Ln ion decreases, a trend
that is reflected by an increase in the absolute value of Δlog
KML upon moving across the Ln series. The poorer affinity of
CHX-macropa for the small lanthanides renders this ligand 10
times more selective than macropa for La3+ over Lu3+ (log KLaL
− log KLuL = 7.84 for CHX-macropa and 6.74 for macropa).
Thus, the integration of a rigid cyclohexyl group into the
D18C6 scaffold of macropa leads to a significant enhancement
in the reverse-size selectivity for large lanthanides.
6.95(1)
3.17(4)
2.51(7)
2.36
log Ka5
1.69
log KLaL
log KLaHL
log KCeL
log KCeHL
log KPrL
log KPrHL
log KNdL
log KNdHL
log KSmL
log KSmHL
log KEuL
log KEuHL
log KGdL
log KGdHL
log KTbL
log KTbHL
log KDyL
log KDyHL
log KHoL
log KErL
log KTmL
log KYbL
log KLuL
14.54(1)
14.38(3)
13.89(1)
13.57(5)
12.79(1)
12.25(3)
11.53(2)
10.98(4)
10.34(3)
14.99
2.28
15.11
2.07
14.70
2.96
14.36
2.08
13.80
2.70
13.01
1.97
13.02
2.48
11.79
2.91
11.72
2.42
10.59
10.10
9.59
9.67(5)
8.96(6)
8.19(3)
7.44(4)
8.89
8.25
6.70(1)
a
b
c
Data provided for comparison. Reference 110. The standard
deviation is given in parentheses and corresponds to the last digit of
the stability constant.
X-ray Structure of La−CHX-macropa. To assess the
impact of the cyclohexyl bridge in CHX-macropa on the
structural architecture of its metal complexes, the solid-state
structure of the La3+ complex of CHX-macropa was obtained
using single-crystal X-ray diffraction. The complex was formed
in situ by mixing equimolar amounts of the ligand and LaCl3·
6.8H2O in water (∼37 mM) and adjusting the pH to 7 with
dilute potassium hydroxide (KOH). Potassium hexafluoro-
phosphate (KPF6, ∼7.5 equiv) was added to the solution, and
single crystals suitable for X-ray diffraction were isolated from
slow evaporation of the solution. The complex crystallized in
increased by 0.1 log K units. The overall basicities126 of the
two ligands, however, are very similar, with the sum of their log
Ka values equaling 20.4 for CHX-macropa and 21.6 for
macropa.
[HiL]
Kai =
[Hi−1L][H+]
(1)
[ML]
−
KML
K
=
the P21/c space group with one PF6 counteranion and two
[M][L]
(2)
water molecules in the outer sphere, giving a formula of
[La(CHX-macropa)(OH2)]PF6·2H2O. Although the initial
refinement of the crystal structure proceeded as expected, a
large residual electron density remained near the La center.
Further analysis of these data revealed that the origin of this
residual density could be explained by whole-molecule
disorder. This whole-molecule disorder was refined, as
described in detail in the Supporting Information (see section
5 and Figures S13 and S14), to give two components of the
[La(CHX-macropa)(OH2)]+ cation with occupancies of 0.85
and 0.15. Both components of the whole-molecule disorder
were identical with respect to coordination geometries and
ligand conformation. As such, our discussion here is based on
the major occupancy component. Crystallographic parameters
and interatomic distances and angles for both components are
given in Tables S1, S4, and S5.
[MHiL]
[MHi−1L][H+]
=
MHiL
(3)
Stability constants, or log KML values (eq 2), for complex
formation were measured from potentiometric titrations of
CHX-macropa carried out in the presence of equimolar Ln3+
ions (Table 2 and Figure S12). For the smaller lanthanides
Tb−Lu, monoprotonated complexes (MHL) could not be
reasonably modeled based on the titration data over the pH
range investigated. For the larger lanthanides Ce−Gd, log
KMHL values (eq 3) ranging from 1.23 to 2.23 could be refined
in the model, but these values carried a high standard deviation
in most cases. We attribute this to the fact that only a minor
amount of monoprotonated species (<10%) would be present
in solution across the pH range investigated, thereby
F
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Figure 2. Effects of substituting an ethylene bridge in macropa’s D18C6 core with a cyclohexylene bridge in CHX-macropa on the thermodynamic
complex stability and size selectivity, plotted as a function of the 6-coordinate ionic radius of each Ln3+ ion. (a) Comparison between the Ln
stability constants (log KML values) of CHX-macropa (blue circles) and macropa (red diamonds). (b) Plot of the change in the complex stability,
Δlog KML, in passing from macropa to CHX-macropa. Δlog K = log KLn−CHX‑macropa − log KLn−macropa
.
The crystal structure reveals that all 10 atoms of the N4O6
donor set are coordinated to the La3+ center, with the
picolinate pendent arms positioned on the same side of the
macrocycle in a syn conformation (Figure 3). A water
molecule, penetrating from underneath the macrocycle,
completes the undecadentate coordination sphere. The
positions of the hydrogen atoms of the water molecules
(inner and outer sphere) could not be definitively located on
the difference Fourier map, most likely because of the whole-
molecule disorder. Their presence, however, is inferred based
on close contacts within the crystal lattice that reflect
hydrogen-bonding interactions. The absolute configuration of
the ligand can be assigned using Δ or Λ symbols to denote the
helical twist of the picolinate pendent arms and δ or λ to
designate the chirality of each of the six 5-membered chelate
rings.127 The conformations of the La−CHX-macropa complex
in the crystal structure are Δ(δλδ)(δλδ) and its enantiomer
Λ(λδλ)(λδλ). As expected, the trans-cyclohexyl group is
positioned in the chair configuration, which is required for
the oxygen atoms of the macrocycle to attain the correct
orientation for metal binding. Notably, the structural arrange-
ment of [La(CHX-macropa)(OH2)]+ is consistent with that
reported previously97 for [La(Hmacropa)(OH2)]2+ and other
complexes of macropa with large metal ions.98,111 This
similarity suggests that the incorporation of a cyclohexyl
group in the ligand backbone does not significantly perturb the
overall conformation of the La3+ complex.
Although the whole-molecule disorder of [La(CHX-
macropa)(OH2)]+ prevents a detailed discussion of the
interatomic distances within this complex, some general
observations can be made. The La1A−N1A and La1A−N2
distances in the structure of [La(CHX-macropa)(OH2)]+ are
2.943 and 2.878 Å, respectively. The similarity of these values
reflects the fact that the metal center is situated centrally within
the macrocycle. However, the interatomic distances between
La1A and the ether donor atoms of the macrocycle, O1A,
O2A, O3A, and O4A, vary more significantly, ranging from
2.779 to 2.988 Å. Notably, both the shortest and longest La−O
interatomic distances reside on the side of the D18C6 ring
bearing the cyclohexyl group. For comparison, the crystal
structure of [La(macropa)(OH2)]+ shows that all four of the
La−Oether interatomic distances are within 0.1 Å of each other.
The disparate La−O interatomic distances in [La(CHX-
macropa)(OH2)]+ are corroborated by DFT calculations
below and may be a consequence of the rigidity of the
G
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Figure 3. X-ray crystal structure of [La(CHX-macropa)(OH2)]PF6·2H2O viewed from the (a) side and (b) top. Hydrogen atoms attached to
−
carbon centers, the PF6 counteranion, and outer-sphere water molecules are omitted for clarity. Thermal ellipsoids are drawn at the 50%
probability level. Owing to a whole-molecule disorder (see the main text), the positions of the hydrogen atoms of the coordinated water molecule
could not be assigned. Only the major component (occupancy = 0.85) is shown.
cyclohexylene group that prevents all of the donor atoms of the
macrocycle from optimally interacting with the metal center.
DFT Calculations. The thermodynamic studies described
above indicate that the addition of a cyclohexyl group into a
ligand scaffold can confer enhanced reverse-size selectivity to
CHX-macropa by disfavoring the formation of stable
complexes with small metal ions. Similar results were reported
for related macrocyclic ligands in which adjacent nitrogen
donors were bridged to form piperazine rings.128,129 For these
previously reported ligands, the decreased stability of
complexes of smaller metal ions was attributed to the
conformational rigidity of the ligands, which inhibited their
abilities to attain geometries suitable for small ion binding. To
probe the role of the conformational rigidity on the Ln-binding
properties of CHX-macropa, DFT calculations were carried
out at the TPSSh/LC RECP/6-31G(d,p) level of theory. We
selected the TPSSh functional and large-core relativistic
effective core potential (LC RECP) approach based on
previous studies that have demonstrated this functional and
core potential to yield more accurate results for Ln complexes
in comparison to other common functionals.130,131 Coordi-
nates for the optimized structures are included in the SI. To
verify the accuracy of our computational approach, we first
calculated the aqueous Gibbs free energy (ΔGaq) for the
exchange reaction shown in eq 4:
Scheme 2. Thermodynamic Cycle Showing the Relationship
between ΔGg and ΔGaq
in subsequent computations (α = 1.0). Using this approach,
ΔGaq can be calculated from eq 5:
ΔGaq = ΔGg + ΔGsol([Ln(L)]+) + ΔGsol(La3+)
− ΔGsol(Ln3+) − ΔGsol([La(L)]+)
(5)
The calculated ΔGaq values for both ligands (Figure S15 and
Table S6) are in very good agreement with the experimental
data and previous DFT-calculated values for macropa,131
verifying the accuracy of this computational approach.
To investigate the origin of the enhanced size selectivity of
CHX-macropa compared to macropa, we first considered the
conformational preference of the ligands. A previous analysis of
the stability of Ln−macropa complexes demonstrated a rather
abrupt change in the conformational preference of the ligand
as the lanthanide series was traversed.110 For the larger
lanthanides, the Δ(δλδ)(δλδ) conformation was observed to
be the most energetically stable, whereas the smaller
lanthanides were most stable with the Δ(λδλ)(λδλ) con-
formation. To determine if the Ln complexes of CHX-macropa
had a similar trend in conformational preferences, we
computationally determined the relative energies of these
conformations for the complexes of La3+, Gd3+, and Lu3+. The
relative energies of eight different ligand conformations,
calculated with the solvation model based on density (SMD)
implicit solvation model for water, are reported in Figure 4.
ΔGaq
[La(L)]+(aq) + Ln3+(aq) ⎯⎯⎯⎯→ [Ln(L)]3+(aq) + La3+(aq)
(4)
This calculation was achieved through the use of the
thermodynamic cycle131−134 presented in Scheme 2, where
ΔGg is the Gibbs free energy of the gas-phase reaction and
ΔGsol is the solvation free energy for the complex in water.
For the Ln3+ ions, previously optimized polarizable
continuum model (PCM) radii were used in all solvation
calculations.131,135 No scaling factor was applied to these radii
H
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Figure 5. Calculated strain energies for macropa (red diamonds) and
CHX-macropa (blue circles) as a function of the ionic radius. Values
are reported relative to the energy of the optimized free ligand, which
was normalized to zero. Dashed lines represent the linear fit of the
data. The ionic radii of the Ln ions are reported based on a
coordination number of 6.
Figure 4. Relative SMD free energies (ΔGsol) of the eight
conformations of [Ln(CHX-macropa)]+ complexes (Ln = La, Gd,
Lu) in aqueous solution. Energies are reported relative to the
Δ(δλδ)(δλδ) conformation.
These calculations reveal that, for CHX-macropa, the Δ(δλδ)-
(δλδ) conformation is the lowest-energy conformation for all
of the lanthanides evaluated. For La3+, the Δ(δλδ)(δλδ)
conformation is calculated to be 3.35 kcal mol−1 lower in
energy than the next-lowest-energy conformation Δ(δλλ)-
(δλλ). This preference is consistent with the crystal structure
of [La(CHX-macropa)(OH2)]+, which attains this Δ(δλδ)-
(δλδ) conformation. Similarly, the Lu3+ complex of CHX-
macropa is calculated to be most stable in the Δ(δλδ)(δλδ)
conformation. The Δ(λδλ)(λδλ) conformation is only margin-
ally higher in energy (1.27 kcal mol−1). By contrast, the Lu3+
complex of macropa energetically favors the Δ(λδλ)(λδλ)
conformation over the Δ(δλδ)(δλδ) conformation by 6.1 kcal
mol−1.110 Thus, macropa has strong conformational prefer-
ences in a manner that is dependent on the size of the Ln ion.
CHX-macropa, on the other hand, maintains a thermodynamic
preference for a single conformation across the entire Ln series
regardless of the ionic radii. We hypothesize that this effect
arises from the enhanced rigidity of the diazacrown backbone
provided by the cyclohexyl functionality, which restricts its
conformational freedom and prevents the ligand from attaining
an optimal configuration for interacting with the small Ln ions.
Thus, this effect may contribute to the enhanced reverse size
selectivity of CHX-macropa.
The energy required for a ligand to distort from its relaxed
equilibrium conformation to bind a metal ion is its strain
energy. The magnitude of this strain energy can have a
pronounced effect on the metal-ion size selectivity of
polydentate ligands.93,115,136 To determine the role of the
strain energy in the Ln ion selectivity of CHX-macropa, this
energy was computed via DFT for several representative Ln-
ion complexes of macropa and CHX-macropa.137 The strain
energies for both chelators increase linearly (R2 ⩾ 0.98) with
decreasing ionic radius of the Ln ion (Figure 5). The strain
energies of CHX-macropa are consistently larger than those for
macropa, which is most likely a consequence of the greater
rigidity of the cyclohexyl backbone. The slopes of these linear
regressions, however, are different. The values of these slopes
are −1.36 and −1.08 kcal mol−1 pm−1 for CHX-macropa and
macropa, respectively. The larger magnitude of the slope in the
linear regression for CHX-macropa indicates that the strain
energy of this ligand is more sensitive to ionic radius changes
than macropa. This result highlights the influence of the
increased conformational rigidity of CHX-macropa on the
steric strain and size selectivity within this class of ligands.
The coordination geometry of metal complexes of macro-
cycles can also reveal structural features that may reflect their
relative stabilities. For example, complexes of macropa with
large metals such as Sr2+, Pb2+, La3+, and Ba2+ hold these ions
centrally in the macrocycle; interatomic distances between
chemically equivalent donor atoms and metal centers are all
nearly equal.97,98,111 By contrast, thermodynamically less stable
complexes of smaller metal ions such as Lu3+ exhibit a
somewhat asymmetric coordination environment; interatomic
distances between the donor atoms and metal centers are
disparate.97 Thus, there appears to be a correlation between
the thermodynamic stability and structural features of these
complexes. Analysis of the DFT-optimized structures may
therefore provide additional insight on the stabilities of these
complexes. The DFT-optimized structure of [Lu(macropa)]+
captures the incongruous Lu−OC (OC = crown ether oxygen)
distances observed experimentally, which span values ranging
from 2.596 to 2.8 Å. For [La(macropa)]+, the La−OC
distances are all nearly equivalent, suggesting that macropa
can optimally bind the large La3+ ion but not the small Lu3+
ion.97 Analysis of the [Ln(CHX-macropa)]+ DFT-optimized
structures reveals that the metal coordination environment
becomes increasingly disparate as the f block is traversed from
large to small lanthanides (Figure S16). In comparison to
[Lu(macropa)]+ (max−min: Lu−NAM = 0 Å; Lu−OC = 0.208
Å), the calculated range of metal−donor distances is greater for
[Lu(CHX-macropa]+ (max−min: Lu−NAM = 0.143 Å; Lu−OC
= 0.218 Å). For both [La(macropa)]+ and [La(CHX-
macropa)]+, however, these distances are nearly equivalent.
This analysis of the geometric parameters shows that the donor
atom interactions of CHX-macropa with small metal ions are
less effective than those of macropa.
Kinetic Inertness of Complexes. The kinetic inertness of
metal complexes is of utmost importance when using f-block
I
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Figure 6. Stability of La−CHX-macropa in the presence of a competing ligand. (a) Time course of the UV spectra of La−CHX-macropa in pH 7.4
MOPS buffer at room temperature with 1000 equiv of DTPA. The final spectrum matches that of a solution of CHX-macropa with 1000 equiv of
DTPA (black dotted line), confirming that no intact La−CHX-macropa complex remains. (b) Plot of the change in absorbance at 282 nm versus
time. Inset: Linear fit of the data to the first-order integrated rate law.
ions in medicine. Free metal ions released from metallodrugs
in vivo accumulate in various tissues of the body, such as the
bone, giving rise to toxic side effects that can prevent their
clinical use in therapeutic and diagnostic applications.
stability afforded by other cyclohexyl-rigidified chela-
tors.139−142 Thus, these studies highlight that more subtle
kinetic effects may be operational during the decomplexation
of metal ions.
Although the affinity of macropa for La3+ (log KLaL
=
14.99)110 is 4.5 orders of magnitude lower than that of
CONCLUSIONS
■
diethylenetriaminepentaacetic acid (DTPA; log KLaL
=
Although the f-block elements reside at the bottom of the
Periodic Table in a region that is largely neglected by many
pharmaceutical scientists, these ions have important applica-
tions in the realm of medicine, as highlighted in this Forum
Article. A key requirement for their biological use is the
development of appropriate chelating agents that can stably
retain and deliver them in vivo. In this work, we have described
our efforts to optimize the large f-block ion chelator macropa
by installing a rigid cyclohexyl group on the macrocycle
backbone. Our results have shown that rigidification of the
macrocycle was generally ineffective for increasing both the
thermodynamic and kinetic stabilities of Ln complexes. These
results contrast with previous studies that have shown that
ligand rigidification using the cyclohexyl group is a highly
effective strategy for developing more stable complexes,
indicating that additional factors may be operational. Despite
the poorer stabilities of Ln complexes of CHX-macropa, this
ligand exhibited an enhanced thermodynamic selectivity for
large lanthanides over small lanthanides, a property that was
supported by DFT calculations to be a consequence of its
increased structural rigidity. Therefore, we have shown that
rigidifying the diaza-18-crown-6 scaffold of macropa with the
cyclohexylene group is a promising strategy to tune the ion size
selectivity of this class of ligands. Although CHX-macropa may
not be directly useful for f-element chelation in the biomedical
realm, it may be a valuable ligand for the separation of large
lanthanides from small lanthanides. Ongoing efforts are
directed toward exploring the potential of CHX-macropa for
this application.
19.48),138 we previously observed that La−macropa is stable
for up to 21 days when challenged with a 1000-fold excess of
DTPA.97 This remarkable inertness of La−macropa, despite its
lower thermodynamic stability, underscores the fact that
kinetic stability cannot be predicted on the basis of
thermodynamic affinity.
On the basis of this rationale, we investigated the stability of
La3+, Gd3+, and Lu3+ complexes of CHX-macropa against
transchelation by DTPA using UV−vis spectrophotometry.
Challenges were performed at room temperature in pH 7.4 3-
(N-morpholino)propanesulfonic acid (MOPS) buffer (100
mM) by adding excess DTPA to the complexes formed in situ.
For comparison, these challenges were also carried out with
macropa. The transchelation of the complexes was monitored
using changes with time of the π−π* absorption at 282 and
285 nm. Under the pseudo-first-order conditions employed,
La−CHX-macropa exhibited moderate kinetic stability when
challenged with 1000 equiv of DTPA, undergoing trans-
chelation with a half-life of 39 1 h (Figure 6). By contrast,
La−macropa was approximately 87% intact after 23 days
(Figure S17). The stability of La−macropa is reduced slightly
from our previous study, in which we used 10 mM instead of
100 mM MOPS.97 Therefore, there may be a small effect of the
buffer composition on the rate of transchelation. Gd−CHX-
macropa displayed 10-fold reduced kinetic stability compared
to Gd−macropa, corresponding to half-lives of 52 1 and 555
11 s, respectively, in the presence of 1000 equiv of DTPA
(Figures S18 and S19). For complexes of both CHX-macropa
and macropa with Lu3+, the amount of DTPA employed in the
study was reduced to only 10 equiv. Nonetheless, the kinetics
of transchelation was so rapid that half-lives could not be
determined (<20 s; Figures S20 and S21). Collectively, these
results indicate that the increased rigidification of CHX-
macropa leads to an increase in the kinetic lability of its Ln
complexes in comparison to macropa. This result is somewhat
unexpected based on the successful enhancement of kinetic
EXPERIMENTAL SECTION
■
Synthesis. All solvents and reagents, unless otherwise noted, were
of ACS grade or higher and were purchased from commercial sources.
Solvents noted as dry were obtained following storage over 3 Å
molecular sieves. [N(CH3)4][OH] was purchased as a 25 wt %
solution in water (H2O; trace metals basis, Beantown Chemical,
Hudson, NH). Concentrated hydrochloric acid (HCl; BDH Aristar
J
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Plus, VWR, Radnor, PA) and metal salts (various commercial
suppliers) were of trace metal grade. Deionized H2O (⩾18 MΩ cm)
was obtained from an Elga Purelab Flex 2 water purification system.
High-performance liquid chromatography (HPLC) consisted of a
CBM-20A communications bus module, an LC-20AP (preparative)
or LC-20AT (analytical) pump, and an SPD-20AV UV−vis detector
monitoring at 270 nm (Shimadzu, Japan). Analytical chromatography
was carried out using an Ultra Aqueous C18 column, 100 Å, 5 μm,
250 mm × 4.6 mm (Restek, Bellefonte, PA) at a flow rate of 1.0 mL
min−1. Semipreparative purification was performed using an Epic
Polar preparative column, 120 Å, 10 μm, 25 cm × 20 mm (ES
Industries, West Berlin, NJ) at a flow rate of 14 mL min−1. NMR
spectra were recorded at 25 °C on a Varian Inova 600 MHz
spectrometer or on a Bruker AV III HD 500 MHz spectrometer
equipped with a broadband Prodigy cryoprobe. Chemical shifts are
reported in parts per million. Spectra acquired in CDCl3 were
referenced to a tetramethylsilane internal standard (0 ppm). Samples
prepared in D2O were spiked with ACN, and the spectra were
referenced to the corresponding signal at either 2.06 ppm (1H) or
1.47 ppm (13C). The splitting of proton resonances in the reported
1H NMR spectra is defined as s = singlet, d = doublet, t = triplet, m =
multiplet, and br = broad. High-resolution mass spectra were obtained
on an Exactive Orbitrap mass spectrometer in positive-ion electro-
spray ionization (ESI) mode (ThermoFisher Scientific, Waltham,
MA). UV−vis spectra were recorded on a Cary 8454 UV−vis
spectrometer (Agilent Technologies, Santa Clara, CA) using 1 cm
quartz cuvettes. Elemental analysis was performed by Atlantic
Microlab, Inc. (Norcross, GA).
Macropa (2HCl·3.5H2O) was synthesized according to a literature
procedure,110,143 except 6-(bromomethyl)pyridine-2-carboxylic acid
methyl ester144 was used instead of 6-(chloromethyl)pyridine-2-
carboxylic acid methyl ester. trans-1,2-Cyclohexanediyldioxydiacetic
acid (1)112 and trans-1,2-cyclohexanediyldioxydicetyl dichloride
(2)113 were prepared according to literature procedures, except
bromoacetic acid was used in place of chloroacetic acid for the
synthesis of 1. Purification of 1 by recrystallization from ethyl acetate,
as previously reported,112 was unsuccessful. Instead, the crude residue
of this compound was purified by silica gel column chromatography
using dichloromethane (DCM) as the eluent. Fractions were pooled
and concentrated to give an amber oil, from which a white solid
precipitated upon standing overnight. The solid was filtered, washed
with DCM, and dried in vacuo, furnishing 1 in 7% yield.
rac-(16aS,20aS)/(16aR,20aR)-Dodecahydro-2H,11H-benzo[b]-
[1,4,10,13]tetraoxa[7,16]diazacyclooctadecine-3,14(4H,15H)-dione
(3). The synthesis of 3 was adapted from a general procedure for the
preparation of similar macrocycles.145 Compound 2 (1.54 g, 5.7
mmol) in benzene (39 mL) and 2,2′-(ethylenedioxy)bis(ethylamine)
(1.58 mL, 10.8 mmol) in benzene (39 mL) were simultaneously
added dropwise over approximately 10 min to a flask containing
vigorously stirring benzene (95 mL) at room temperature. The
resulting pale-yellow suspension was stirred for an additional 30 min
and then filtered. The flask and filtered solid were rinsed with benzene
(3 × 20 mL), and the faintly yellow filtrate was concentrated under
reduced pressure to an off-white solid. Column chromatography
[silica gel, 0−2.5% methanol (MeOH) in CHCl3] afforded the title
compound as a white solid (1.34 g, 68% yield). 1H NMR (500 MHz,
CDCl3): δ 7.68 (br t, 2H), 4.18−3.99 (m, 4H), 3.70−3.63 (m, 2H),
3.62−3.47 (m, 8H), 3.46−3.35 (m, 2H), 3.31−3.20 (m, 2H), 2.21−
2.06 (m, 2H), 1.83−1.66 (m, 2H), 1.25−1.07 (m, 4H). 13C{1H}
NMR (126 MHz, CDCl3): δ 170.50, 81.83, 70.01, 69.38, 67.59, 39.25,
29.59, 23.87. ESI-MS. Found: m/z 711.37820. Calcd for
[C32H56N4O12Na]+: m/z 711.37869. Found: m/z 367.18373. Calcd
for [C16H28N2O6Na]+: m/z 367.18396. Found: m/z 345.20185. Calcd
for [C16H29N2O6]+: m/z 345.20201.
mL). A hot suspension of 3 (1.31 g, 3.8 mmol) in dry THF (11 mL)
was slowly added to the chilled suspension of LAH, and the gray
mixture was heated at reflux for 24 h. The reaction mixture was
subsequently cooled in an ice bath, and the excess LAH was destroyed
by the careful addition of 50:50 H2O/THF (∼8 mL). The mixture
was vacuum filtered on sintered glass equipped with a nylon
membrane (0.22 μm), and the filtered solid was washed with 3
portions of hot benzene (100 mL total). The filtrate was concentrated
under reduced pressure to an oil, which was redissolved in DCM
(∼25 mL), dried over sodium sulfate, and concentrated to give
1
macrocycle 4 as a pale-yellow oil (1.00 g, 83% yield). H NMR (600
MHz, CDCl3): δ 3.79−3.70 (m, 2H), 3.66−3.56 (m, 10H), 3.21−
3.14 (m, 2H), 2.90−2.73 (m, 8H), 2.11−2.03 (m, 2H), 1.71−1.62
(m, 2H), 1.21−1.10 (m, 4H). 13C{1H} NMR (126 MHz, CDCl3): δ
81.71, 70.00, 69.66, 67.75, 49.57, 49.20, 30.37, 24.14. ESI-MS. Found:
m/z 317.24279. Calcd for [C16H33N2O4]+: m/z 317.24348. Found:
m/z 159.12507. Calcd for [C16H34N2O4]2+: m/z 159.12538.
rac-6,6′-[[(16aS,20aS)/(16aR,20aR)-Hexadecahydro-4H,13H-
benzo[b][1,4,10,13]tetraoxa[7,16]diazacyclooctadecine-4,13-diyl]-
bis(methylene)]dipicolinic acid (CHX-macropa). 6-(Bromomethyl)-
pyridine-2-carboxylic acid methyl ester144 (0.740 g, 3.2 mmol) and
Na2CO3 (1.06 g, 10.0 mmol) were added to a slight suspension of 4
in dry ACN (33 mL). The resulting mixture was heated at reflux and
stirred overnight. It was then concentrated under reduced pressure,
and the pale-yellow residue was dissolved in DCM (30 mL) and
washed with H2O (15 mL). The organic layer was separated, and the
aqueous layer was extracted two more times with DCM (30 mL). The
combined organics were backextracted with H2O (5 mL), dried over
sodium sulfate, and concentrated under reduced pressure to give a
cloudy orange-yellow oil. ESI-MS. Found: m/z 637.32460. Calcd for
[C32H46N4O8Na]+: m/z 637.32079. Found: m/z 615.34266. Calcd for
[C32H47N4O8]+: m/z 615.33884. Found: m/z 308.17495. Calcd for
[C32H48N4O8]2+: m/z 308.17306.
The oil was redissolved in 6 M HCl (4 mL) and heated at 90 °C
for 15 h. The orange-brown reaction solution was then concentrated
on a rotary evaporator, and the crude residue was purified by reverse-
phase semipreparative HPLC using a binary mobile phase containing
0.1% CF3COOH (program: 10% MeOH/H2O for 5 min, followed by
a linear gradient to 100% MeOH over 35 min). Pure fractions were
combined and concentrated under reduced pressure. To prepare the
hydrochloride salt, the residue was dissolved repeatedly in 6 M HCl
and concentrated on a rotary evaporator, and then dried in vacuo for
48 h to give the title compound as a off-white, crystalline solid (0.494
g, 42% yield). 1H NMR (500 MHz, D2O, pD ≈ 1−2): δ 8.17 (d, J =
7.8 Hz, 2H), 8.10 (t, J = 7.8 Hz, 2H), 7.74 (d, J = 7.7 Hz, 2H), 4.77−
4.67 (m, 4H), 4.04−3.84 (m, 8H), 3.73−3.50 (m, 12H), 3.33−3.22
(m, 2H), 1.99−1.88 (m, 2H), 1.60−1.48 (m, 2H), 1.13−1.00 (m,
2H), 0.98−0.81 (m, 2H). 13C{1H} NMR (126 MHz, D2O, pD ≈ 1−
2): δ 168.05, 150.43, 147.56, 140.71, 128.74, 126.25, 82.37, 70.28,
64.73, 62.36, 58.48, 55.25, 55.18, 29.58, 23.57. ESI-MS. Found: m/z
587.31030. Calcd for [C30H43N4O8]+: m/z 587.30754. Found: m/z
294.15875. Calcd for [C30H44N4O8]2+: m/z 294.15741. Elem anal.
Found: C, 49.16; H, 6.56; N, 7.47. Calcd for C30H42N4O8·3.2HCl·
1.8H2O: C, 48.97; H, 6.69; N, 7.61. HPLC: tR = 18.70 min.
X-ray Diffraction Measurements. Low-temperature X-ray
diffraction data for crystals of [H2CHX-macropa]Cl2 and [La(CHX-
macropa)(OH2)]PF6·2H2O were collected on a Rigaku XtaLAB
Synergy diffractometer coupled to a Rigaku Hypix detector with Cu
Kα radiation (λ = 1.54184 Å) from a PhotonJet microfocus X-ray
source at 100 K. The diffraction images were processed and scaled
using the CrysAlisPro software (2015, Rigaku OD, The Woodlands,
TX). The structures were solved through intrinsic phasing using
SHELXT146 and refined against F2 on all data by full-matrix least
squares with SHELXL147 following established refinement strat-
egies.148 All non-hydrogen atoms were refined anisotropically. All
hydrogen atoms bound to carbon atoms were included in the model
at geometrically calculated positions and refined using a riding model.
Hydrogen atoms bound to oxygen and nitrogen atoms were located in
the difference Fourier map, if possible, and subsequently refined
semifreely with the help of distance restraints. The isotropic
rac-(16aS,20aS)/(16aR,20aR)-Hexadecahydro-2H,11H-benzo[b]-
[1,4,10,13]tetraoxa[7,16]diazacyclooctadecine (4). The synthesis of
4 was adapted from a general procedure for the reduction of similar
macrocyclic diamides.145 A three-necked round-bottom flask under a
positive atmosphere of N2 was charged with lithium aluminum
hydride (LAH; 1.03 g, 27.1 mmol) and dry tetrahydrofuran (THF, 26
K
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displacement parameters of all hydrogen atoms were fixed to 1.2 times
the Ueq value of the atoms that they are linked to (1.5 times for
methyl groups). In the structure of [H2CHX-macropa]Cl2, one
chloride counterion was disordered over two positions, which was
modeled using the PART command, with the occupancies summing
to 1 (Cl2A 0.18; Cl2B 0.82). The structure of [H2CHX-macropa]Cl2
contains one heavily disordered methanol molecule, which is located
on a crystallographic center of inversion and has been modeled using
the solvent mask routine in Olex2 (volume, 246.8 Å3; e− count, 32).149
The rationale and procedure for the structure refinement of
[La(CHX-macropa)(OH2)]PF6·2H2O, which exhibited whole-mole-
cule disorder, are detailed in the SI. Our approach to addressing this
crystallographic problem in a systematic manner followed procedures
outlined for a previously reported osmium(VI) nitrido complex that
also crystallized with whole-molecule disorder.150 Crystallographic
data collection and refinement parameters, interatomic distances, and
interatomic angles are given in Tables S1−S5.
Solution Thermodynamics. Potentiometric titrations were
carried out using a Metrohm Titrando 888 titrator equipped with a
Ross Orion combination electrode (8103BN, ThermoFisher
Scientific), a Metrohm 806 exchange unit with an automatic buret
(10 mL), and Tiamo 2.5 software. The titration vessel was fitted with
a removable glass cell (∼70 mL) and thermostated at 25 °C using a
Thermomix 1442D circulating H2O bath. CO2 was excluded from the
vessel using a small positive pressure of argon scrubbed with 30 wt %
KOH. Carbonate-free KOH (∼0.1 M) was prepared by the
dissolution of pellets (semiconductor grade, 99.99% trace metals
basis, Sigma-Aldrich) or by the dilution of Baker Dilut-It analytical
concentrate (Avantor) using freshly boiled H2O (⩾18 MΩ cm) and
standardized by potentiometric titration against potassium hydrogen
phthalate. HCl (0.1 M, J.T. Baker) was standardized against Tris base
(J.T. Baker). All titration solutions were maintained at a constant
ionic strength of 0.1 M using KCl (BioUltra, ⩾99.5%, Sigma-Aldrich)
and equilibrated for 25 min prior to the addition of titrant.
Before every ligand or ligand−metal titration, the electrode was
calibrated in terms of the hydrogen-ion concentration by titrating a
solution of standardized HCl (5 mM) containing a supporting
electrolyte (H/KCl = 0.1 M) with standardized KOH. Data within the
pH ranges of 2.3−3.2 and 10.8−11.2 were analyzed using the program
Glee (version 3.0.21)151 to obtain the standard electrode potential
(E0) and slope factor. The H2O ion product (pKw = 13.78) was taken
from the literature.152 Prior to titrations containing La3+ and Ce3+
metal ions, which form complexes of high stability with CHX-
macropa, the electrode was calibrated to pH 2.0 by using a higher
concentration of HCl (10 mM).
LnCl3 salts were dissolved in 0.1 M HCl to give stock solutions of
the metal ions. The concentrations of these solutions were
determined by complexometric titration in a NaOAc buffer (100
mM, pH 5.5) with a standardized solution of ethylenediaminetetra-
acetic acid disodium salt (Alfa Aesar, 57.5 mM). Xylenol orange (1
mg mL−1, 4 drops) was used as the indicator. A stock solution of
CHX-macropa was prepared in H2O, and its exact concentration was
calculated in triplicate from the volume difference between the two
sharp end points of its potentiometric titration curve with KOH.
The ligand protonation constants and stability constants of Ln3+
complexes were measured by adding standardized KOH to an
aqueous solution (20 mL) containing ligand (∼0.022 mmol), 0.1 M
HCl (0.1 mmol), and 1 M KCl (1.9 mmol) in the absence and
presence of metal ion (∼0.022 mmol), respectively. The titration
method employed a 0.1 mV min−1 drift limit and a maximum wait
time of 180 s (protonation constants) or 300 s (stability constants)
between the addition of aliquots of base. For metal−ligand titrations,
a minimum wait time of 120 s was also enforced. If the drift criterion
was not satisfied within the maximum wait time, the point was
excluded from the titration data set. Additionally, data beyond the end
point (pH ∼7) in metal−ligand titrations were omitted owing to
precipitation of Ln(OH)3 at high pH. Back titrations with 0.1 M HCl
were performed once for each Ln−ligand combination to verify that
equilibrium was attained at each data point under the given
conditions. Only the back titration of CHX-macropa with La3+
displayed poor reversibility. A titration in which the minimum wait
time between the addition of aliquots of base was increased to 10 min
gave rise to the same calculated stability constant as when the
minimum wait time was 120 s, indicating that the hysteresis between
forward and back titrations arises from slow complex dissociation
rather than formation kinetics.
The protonation and stability constants were refined using the
program Hyperquad2013.153 Only the proton concentration was
admitted as a refineable parameter. The protonation constants,
defined in eq 1, were calculated from the average of seven titrations
using data points from pH 2.3 to 11. With the protonation constants
in hand, the stepwise stability constants (eq 2) and protonation
constants of the metal complexes (eq 3) were calculated from the
average of three forward titrations. Hydrolysis constants for the
formation of Ln(OH)2+ in aqueous solution were included in the
model.154 The errors provided correspond to 1 standard deviation.
Computational Methods. DFT calculations were performed
using the Gaussian09 software,155 following procedures similar to
those previously described.131,134,156 Macropa97 and CHX-macropa
were optimized in the deprotonated state starting from their
respective crystal structures. The starting geometries for the metal
complexes were taken from previous computational studies on
macropa.110 Inner-sphere H2O molecules were not included in the
optimization of the metal complexes. Geometries of the complexes
were optimized in the gas phase at 298 K using the TPSSh
functional.157 Light atoms (hydrogen, carbon, nitrogen, and oxygen)
were treated with the 6-31G(d,p) basis set, and metal atoms were
assigned the LC RECP and related (7s6p5d)/[5s4p3d] basis
set.158,159 The LC RECP calculations include 46 + 4f n electrons in
the core of the metal center, leaving the outermost 5s, 5p, 5d, and 6s
electrons to be handled explicitly. As such, all calculations were
performed in the pseudosinglet state. This method has been
previously justified by the corelike properties of the f elec-
trons.131,160−162 Frequency calculations were performed to confirm
that the calculated structures were local minima on the potential
energy surface and to determine the thermodynamic properties of the
complexes. Solvation energies were calculated using the SMD
method, and thermodynamic parameters in H2O were estimated by
taking the thermodynamic corrections from gas-phase calculations as
previously described.163 Hydration free energies of the Ln3+ ions were
calculated as previously described using the SMD solvation model and
parametrized PCM radii.131,135 Basis set superposition errors, which
arise from the use of finite basis sets and lead to an overestimation of
the binding energy, were calculated in the gas phase using the
Counterpoise method,164,165 as implemented in Gaussian09, and these
energies were included in the metal-exchange reaction calculations.
Optimized coordinates for the complexes are given in the SI. The
strain energy is reported as the difference in energy between the
optimized structure of the free ligand and the single-point energy of
the ligand in the geometry of the metal complex, with the metal ion
removed at the TPSSh/6-31G(d,p) level of theory. All free ligand
structures were handled in the deprotonated state.
Transchelation Challenges. The kinetic stabilities of complexes
of La3+, Gd3+, and Lu3+ with CHX-macropa and macropa were
evaluated in the presence of excess diethylenetriaminepentaacetic acid
(DTPA) in a manner similar to that previously described.97 All
challenges were carried out in 100 mM pH 7.4 MOPS (I = 1 M with
[N(CH3)4][Cl], with the pH adjusted using aqueous [N(CH3)4]-
[OH]). An aqueous solution of DTPA (125 mM) was prepared at pH
7.4 using [N(CH3)4][OH] and diluted to 1.25 mM with MOPS
buffer. Ln−L complexes were formed in situ by combining equimolar
amounts (500 μM) of standardized LnCl3 solution (0.1 M HCl) and
ligand (H2O) in MOPS buffer. Challenges were initiated by adding an
aliquot of the Ln−L complex (500 μL) to a cuvette containing 125 or
1.25 mM DTPA (2000 μL). La−L and Gd−L complexes were
challenged with 1000-fold excess DTPA (100 mM), and Lu−L
complexes were challenged with 10-fold excess DTPA (1 mM). UV−
vis spectra were acquired at various intervals (at least 10 data points)
until there were no further spectral changes, except for samples
containing La−macropa, which were monitored for 23 days. Pseudo-
L
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first-order rate constants (kobs) for transchelation of the complexes
were calculated from the slopes of the plots of ln(At − A∞) versus
time using linear regression. A∞ is the final absorbance value, and At is
the absorbance value at time t. Absorbance values were monitored at
282 nm for La3+ and 285 nm for Gd3+. Half-lives (t1/2) were calculated
using the equation t1/2 = ln(2)/kobs and are reported as the mean of
three replicates 1 standard deviation.
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ASSOCIATED CONTENT
* Supporting Information
The Supporting Information is available free of charge on the
ACS Publications website at DOI: 10.1021/acs.inorg-
chem.9b01277.
■
S
Spectral characterization of synthesized compounds,
details of computational and crystallographic studies,
supporting figures for DFT and transchelation studies,
and optimized coordinates from DFT calculations
(PDF)
Accession Codes
CCDC 1913153 and 1913154 contain the supplementary
crystallographic data for this paper. These data can be obtained
free of charge via www.ccdc.cam.ac.uk/data_request/cif, or by
emailing data_request@ccdc.cam.ac.uk, or by contacting The
Cambridge Crystallographic Data Centre, 12 Union Road,
Cambridge CB2 1EZ, UK; fax: +44 1223 336033.
(17) Cawthray, J. F.; Weekes, D. M.; Sivak, O.; Creagh, A. L.;
Ibrahim, F.; Iafrate, M.; Haynes, C. A.; Wasan, K. M.; Orvig, C. In
Vivo Study and Thermodynamic Investigation of Two Lanthanum
Complexes, La(dpp)3 and La(XT), for the Treatment of Bone
Resorption Disorders. Chem. Sci. 2015, 6, 6439−6447.
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nol®): A Novel Agent for the Treatment of Hyperphosphataemia in
Renal Failure and Dialysis Patients. Int. J. Clin. Pract. 2005, 59, 1091−
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Probes for Bioresponsive Imaging. Chem. Rev. 2014, 114, 4496−4539.
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Resonance Imaging. Chem. Soc. Rev. 2006, 35, 557−571.
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Multimeric MR Contrast Agents. Acc. Chem. Res. 2009, 42, 893−903.
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Chemistry of MRI Contrast Agents: Current Challenges and New
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AUTHOR INFORMATION
Corresponding Author
*E-mail: jjw275@cornell.edu.
ORCID
Nikki A. Thiele: 0000-0003-3301-0849
Joshua J. Woods: 0000-0002-6213-4093
Justin J. Wilson: 0000-0002-4086-7982
■
Author Contributions
§These authors contributed equally.
Notes
The authors declare no competing financial interest.
ACKNOWLEDGMENTS
■
J.J.W. is supported by the NSF-GRFP (Grant DGE-1650441).
This work was supported by Cornell University. This research
made use of the NMR Facility at Cornell University, which is
supported, in part, by the NSF under Award CHE-1531632. A.
S. Ivanov and V. S. Bryantsev (Chemical Sciences Division,
Oak Ridge National Laboratory, Oak Ridge, TN) are thanked
for their helpful correspondence involving DFT calculations.
(23) Evans, C. H. Biochemistry of the Lanthanides; Springer: Boston,
MA, 1990.
(24) Blasse, G.; Grabmaier, B. C. How Does a Luminescent Material
Absorb Its Excitation Energy? In Luminescent Materials; Springer:
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Actinides. In Physical Inorganic Chemistry: A Coordination Chemistry
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to Assay: Lessons Learned in Lanthanide Luminescence. Acc. Chem.
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