Responsive nature of 1,2,4,5-Tetrakis(2-pyridylthiomethyl)benzene toward group 12 metal nitrates: Activity of coordinated nitrate in metal complexes

Article abstract of DOI:10.1021/ic500746e

This study provides a detailed analysis on the responsive behavior of 1,2,4,5-tetrakis(2-pyridylthiomethyl)benzene (L) toward group 12 metal nitrates in both aqueous and nonaqueous media. The ligand L proved to be an environmentally responsive species, and structural investigations of its complexes with respective M(NO₃)₂ (M = Zn, Cd, and Hg) allowed one to remark on the inherent activity of the nitrate ion, resulting in a distinctively higher coordination number and dimensionality to the metal cations.

Full text of DOI:10.1021/ic500746e

Communication
pubs.acs.org/IC
Responsive Nature of 1,2,4,5-Tetrakis(2-pyridylthiomethyl)benzene
Toward Group 12 Metal Nitrates: Activity of Coordinated Nitrate in
Metal Complexes
Neeru,† Dolly Yadav,† Geeta Tiwari,† Jai Deo Singh,*,† and Raymond J. Butcher‡
†Department of Chemistry, Indian Institute of Technology Delhi (IITD), Hauz Khas, New Delhi 110 016, India
‡Department of Chemistry, Howard University, Washington, D.C. 20059, United States
S
* Supporting Information
addition of M2+ ion (in H2O) to a solution of the ligand (in
ABSTRACT: This study provides a detailed analysis on
the responsive behavior of 1,2,4,5-tetrakis(2-
pyridylthiomethyl)benzene (L) toward group 12 metal
nitrates in both aqueous and nonaqueous media. The
ligand L proved to be an environmentally responsive
species, and structural investigations of its complexes with
respective M(NO3)2 (M = Zn, Cd, and Hg) allowed one
to remark on the inherent activity of the nitrate ion,
resulting in a distinctively higher coordination number and
dimensionality to the metal cations.
CH3CN), were notable. However, the “turn-on” responses in the
fluorescence measurements continued to a 8:1 M/L stoichiom-
etry. The availability of donor sites in ligand L reasonably offers
the possibility of coordination to a maximum of two metal ions.
The unusual spectroscopic behavior of the ligand toward M2+ (M
= Zn, Cd, and Hg) ions in a mixed-solvent system was intriguing,
and conclusions were drawn only when a crystal obtained from a
solution titrated to a 8:1 M/L stoichiometry in mixed-solvent
system was analyzed by X-ray crystallography. The single-crystal
X-ray structure of this species revealed the formation of a
protonated species (H4L) in the “aabb” conformation with
protonation of all nitrogen atoms of the four thiopyridyl ring of
the ligand. The influence of protonation on pyridyl nitrogen
atoms is well noticeable as CN−C bond angles of the pyridyl
ring of the free ligand L expand from 117.0(1)−117.2(1)° to
123.7(2)−123.0(2)° in H4L and well comparable to several
pyridine-based receptors familiar for their selective nitrate-ion
recognition.7 In the present study, the possibility of protonation
of the ligand likely occurs in the mixed-solvent system H2O/
he chemistry of group 12 metal ions with various organic
Tdonors and structural aspects of their complexes has always
been a fascinating area of research interest.1 This has led to
considerable effort in studying the interactive behavior of group
12 metal ions with various organic donor bases, in particular,
having pyridyl nitrogen atoms as major coordination sites.
Additionally, they also exhibit selective binding affinities toward
ZnII, CdII, and HgII ions when devised as chemosensors and
biosensors.2 In the past few decades, the design of organic donor
bases bearing a combination of both nitrogen and sulfur donor
atoms also received considerable attention because they show
improved selectivity and have an exceptional background to
explain the reactivity−selectivity of these metal ions in chemical
and biological processes.3 Despite the fact that widespread
progress has been made in understanding the chemistry of group
12 metal ions, challenges remain in controlling the activity of the
anion. In fact, the anion plays a crucial role in regulating the
functional behavior of metal ions in chemical and biological
processes.4 Specified by our interest5 in the chemistry of
conformationally flexible organochalcogen ligands, herein we
report the synthesis and structural aspect of 1,2,4,5-tetrakis(2-
pyridylthiomethyl)benzene (L) bearing nitrogen and sulfur as
donor sites and its complexation behavior with group 12 metal
nitrates. The species L was synthesized by a known procedure
and revealed as the “aabb” (a = above; b = below) pattern rather
than the reported “abab” conformation for substituted
thiopyridyl donor arms.6 Primarily, the species L exhibited its
preferred selectivity for M2+ (M = Zn, Cd, and Hg) ions over
other competitive first-row transition-metal ions. The changes in
the UV−vis spectra for L and further “turn-on” fluorescence
responses (λex = 280 nm; λem = 340 nm) with striking red shifts in
emissions, with a shift of up to 5−10 nm upon the successive
CH3CN because of the partial hydrolysis of metal salts Mn+
+
xH2O ↔ [M(H2O)x]n+ ↔ M(OH)x(n−x)+ + xH+ (n = 2; x = 4 or
6) during titration. In the present conditions, the H+ ions
produced in solution were enough to protonate the pyridyl
nitrogen atoms of the ligand. The behavior of L with other metal
ions such as Ni2+ and Pd2+ was also examined in the same
reaction conditions but led to recovery of the ligand in a new
conformation (L1). These results reveal a complicated solution
behavior, where possibilities of complexation followed by
decomplexation, partial or complete hydrolysis of metal salts,
and protonation of the ligand in an aqueous organic solvent are
likely to occur. Therefore, metal−ligand interactions in an
aqueous organic solvent medium may not be suggestive and are
inconclusive, especially when quantification of the metal ion is
being suggested on the basis of spectroscopic changes. The
reaction event of L with different cations in a mixed-solvent
system is outlined in Scheme 1. To explore and understand the
coordinative behavior of the nitrogen center of L with group 12
metal nitrates, the metal/ligand titration experiments were also
performed in an aprotic nonaqueous organic medium, i.e., in
acetonitrile, and monitored by spectroscopic changes.
Received: March 31, 2014
Published: June 2, 2014
© 2014 American Chemical Society
5887
dx.doi.org/10.1021/ic500746e | Inorg. Chem. 2014, 53, 5887−5889

Inorganic Chemistry
Communication
The crystals of 1−3 were grown under strictly anhydrous and
similar conditions upon the slow evaporation of a clear solution
of the respective complexes 1−3 dissolved in acetonitrile. The
complex 1 (M = Zn) revealed a discrete neutral unit, where the
ligand forms a bridge between two metal centers by adopting its
energetically least preferred “aaaa” conformation for all four of its
thiopyridyl donor arms. Each ZnII center is coordinated by the
nitrogen atoms of two adjoining thiopyridyl arms of the bridging
ligand. The remaining coordination site on zinc is completed by
the involvement of two nitrate ions bonded in a unidentate
manner. This overall coordination scenario around the ZnII
center provides a discrete molecular species in a slightly distorted
tetrahedral configuration. The most interesting feature of this
structure is the observation of the unexpected and rarest anion−
solvent (ONO2···CH3CN) interactions as one of the nitrate ions
on each ZnII center in complex 1 interacts with the solvent
CH3CN molecule (Figure 1).
Scheme 1. Coordination Behavior of L with Metal Cations
The ligand L itself revealed temperature-dependent emissive
changes in the fluorescence measurements and may be suggestive
of rapid conformational changes above 35 °C (Figure S5 in the
Ssupporting Information, SI). Indeed, the ligand L in acetonitrile
proved to be a temperature-sensitive species. Considering the
responsive nature of the ligand, the interactive behavior of L with
M2+ (M = Zn, Cd, and Hg) ions was also monitored
spectroscopically in CH3CN at room temperature. The
appearance of isosbestic points at ∼235 nm (nearly at the
same positions for respective M2+ ions) in UV−vis spectra and
titration curves plotting the absorbance as a function of the added
metal ions indicated a 2:1 M/L binding stoichiometry. Similarly,
the acquisition of fluorescence emission spectra with the
successive addition of M2+ ions (2 × 10−4 M in CH3CN) to a
solution of L (1 × 10−4 M in CH3CN) revealed significant “turn-
on” fluorescence responses (λex = 280 nm; λem = 340 nm) to a
maximum 2:1 M/L stoichiometry. Although spectroscopic
changes can be taken into consideration for metal-ion speciation,
real insight into the interactive behavior of L with M2+ ions and
the structural parameters was required to establish the overall real
reaction scenario. Complexes of ZnII, CdII, and HgII (1−3) were
also isolated as solid crystalline materials by reaction of the ligand
with respective metal nitrate salts in acetonitrile at room
temperature (Scheme 2).
Figure 1. Perspective view of discrete ZnII complex species 1 with
ONO2···CH3CN interactions.
The influence of the solvent on metal-bonded nitrate is quite
noteworthy because N−O and O−N−O bond lengths and bond
angles respectively were observed in the ranges of 1.216(3)−
1.282(2) Å and 117.6(2)−123(2)°. We believe that the
observation of ONO2···CH3CN interactions is a positive
indication of the perturbed symmetry of nitrate ions and
suggests that complex 1 should remain active in the solution.
Nevertheless, these interactions are valuable for new fundamen-
tals because they should provide an experimental basis in
understanding the upcoming activity of the anion, i.e., the
structural outcome of a M/L combination attributable to
structural and dynamical activities of nitrate in the present
reaction conditions. Sensibly, structural investigations of CdII
and HgII complexes obtained under conditions similar to those of
ZnII were necessary to study the structural trends in the same
family group. A structural comparison of the crystal structures of
CdII and HgII complexes revealed the close conformational and
coordinative resemblance of the ligands in complexes 1−3. The
“aaaa” conformational pattern of the ligand upon metal-ion
coordination provides a constrained coordination environment
around metal centers. The M−N bonds of 2.001(18)−2.042
(17) Å (M = Zn, 1), 2.281(7)−2.304(5) Å (M = Cd, 2), and
2.18(1) Å (M = Hg, 3) predict an increase of the M−N bond
upon moving from ZnII to CdII but a decrease upon continuing
from CdII to HgII. On the other hand, the M−O bonds (for
nitrate ion bonded to metal in a unidentate mode) range from
1.989(16)−2.017(17) Å (1) to 2.331(7) Å (2) to 2.57(1) Å (3).
Similarly, the N−M−N angles in the complexes increases from
126.93(7)° to 137.0(2)° to 147.6(4)° upon moving from 1 to 2
Scheme 2. Synthesis and Coordination Environment in 1−3
Physicochemical and spectroscopic data for complexes 1−3
were consistent and in favor of 2:1 M/L complex formation.
These complexes show moderate-to-poor solubility in common
organic solvents. The IR spectra of complexes 1−3 evidence
coordination of the ligand through its nitrogen-to-metal centers,
yet it was not possible to distinguish unambiguously the bonding
of nitrate because all three complex species exhibited comparable
peak patterns in the range of ν 1490−1250 cm−1 for nitrate. The
structural confirmation for 1−3 came from single-crystal X-ray
diffraction studies.
5888
dx.doi.org/10.1021/ic500746e | Inorg. Chem. 2014, 53, 5887−5889

Inorganic Chemistry
Communication
to 3, respectively. The solvated nitrate ions, as seen in 1 with
perturbed symmetry, corresponding to the nitrates in complexes
2 and 3 advance them to an extended structure in such a way that
they bind in chelating bidentate mode to the same metal center
and at the same time to the adjacent metal centers in bridging
mode. The overall activity of the nitrate ion thus results in anion-
assisted 2D structures for 2 and 3 (Figures S13−S16 in the SI).
The shortest inter- and intramolecular (M−O) bonds due to
nitrate ions bonded to neighboring metal centers were observed
in the ranges of 2.378(5)−2.410(6) (M = Cd, 2) and 2.44(1)−
2.67(1) (M = Hg, 3). The structural complexities due to nitrate
ions are found to be similar in 2 and 3 and exhibit a close
crystallographic resemblance. We considered all possible
experimental factors and reasoned that the structural complex-
ities in 2 and 3 were a result of the inherent activity of the nitrate
ion, which is deeply related to solvation around it. The behavior
of the nitrate ions in 2 and 3 could be rationalized in view of the
fundamentals associated with nitrate chemistry. For instance, the
from acetonitrile and rationalized in terms of the coordination
flexibility of the anion.12
Conclusively, the present findings provide a systematic
experimental basis in understanding the “symmetry breaking”,
dynamics, and interactions of the anion with solvent molecules
and serve as a tool in recognizing the “inherent activity” of the
nitrate ion to some extent. These findings also provide some basic
insights into a rather arguable issue in inorganic chemistry, over the
presence of nitrate ion as a mere counterion or play a vital role in
determining the dimensionality and structure of the coordination
polymer.
ASSOCIATED CONTENT
* Supporting Information
X-ray crystallographic data in CIF format, experimental details,
and spectroscopic data. This material is available free of charge
via the Internet at http://pubs.acs.org.
■
S
AUTHOR INFORMATION
Corresponding Author
*E-mail: jaideo@chemistry.iitd.ernet.in.
−
■
nitrate ion (NO3 ) in the gas phase is frequently encountered as
a planar ion in D3h symmetry, with the negative charge uniformly
distributed on the three oxygen atoms.8 In polar solvents,
however, the D3h symmetry of the nitrate ion is expected to be
broken because of the nonsymmetric solvation environment and
partially localizes the negative charge. There is also evidence in
which nitrate bonded to a metal ion in its complexes exhibits a
reduction in the local symmetry9 from D3h to C2v symmetry. On
the contrary, while working in aqueous or protic media, the
nitrates usually act as charge-balancing counterions in their
respective metal complexes. Considering the nature of the nitrate
ion in different environments, it is more likely that a
nonsymmetrical solvation environment of the anion in complex
1 adopts a nonplanar feature and partially localizes the negative
charge on their nonbonded end. The partially localized negative
charge on nitrate inevitably allows it to act as a weak nucleophile
(the anion becomes a ligand), which, in turn, interacts with
solvent molecules, as seen in the crystal structure of 1.
Apparently, the advances activity of the anion can be visualized
provided that the geometrical preferences and charge-accepting
ability of a metal ion allow it to proceed. On the basis of
comparable experimental conditions and structural information
obtained for complexes 1−3, the extended coordination of
nitrate in 2 and 3 may be anticipated because of the relativistic
effect.10 As relativistic effects become significant for the heaviest
ions, the relativistic stabilization of the d and s orbitals of the
heaviest ions enhances their charge-accepting ability relative to
those of lighter ions in the same group of the periodic table.
Because of the counterions bearing negative charges and the
increased charge-accepting properties of metal ions in 2 and 3,
the structural characteristics are fundamentally different
compared to 1. The overall nitrate ion activity in these complexes
results in an extended 2D coordination structure. Definitely, the
role of external influences in the activity of the nitrate ion cannot
be ignored because the nitrate ion in atmospheric and
geochemical sciences is considered to be a photochemically
active species and is believed to remain active in the dark.11 It is
difficult to make any further comment on the exact nature of the
symmetry breaking of the nitrate ion in the solution state except
that complexes 1−3 were prepared and crystallized from
acetonitrile in comparable reaction conditions. It is important
to note that the extended bonding of the nitrate ion has also been
seen in a few of its metal complexes derived from silver(I),
cadmium(II), and lead(II) nitrates either prepared or crystallized
Notes
The authors declare no competing financial interest.
ACKNOWLEDGMENTS
■
We thank CSIR and UGC, India, for financial support and FIST
and IITD for X-ray diffraction instrumentation. The authors are
grateful to Dr. Yoshiaki Tanabe for determining the crystal
structure of complex 2.
REFERENCES
■
(1) (a) Brand, U.; Vahrenkamp, H. Inorg. Chem. 1995, 34, 3285−3293.
(b) Holm, R. H.; Kennepohl, P.; Solomon, E. I. Chem. Rev. 1996, 96,
2239−2314. (c) Lipscomb, W. N.; Strater, N. Chem. Rev. 1996, 96,
̈
2375−2434. (d) Parkin, G. Chem. Rev. 2004, 104, 699−768. (e) Zastrow,
M. L.; Pecoraro, V. L. J. Am. Chem. Soc. 2013, 135, 5895−5903.
(2) (a) Kim, H. N.; Ren, W. X.; Kim, J. S.; Yoon, J. Chem. Soc. Rev.
2012, 41, 3210−3244. (b) Lee, M. H.; Yang, Z.; Lim, C. W.; Lee, Y. H.;
Dongbang, S.; Kang, C.; Kim, J. S. Chem. Rev. 2013, 113, 5071−5109.
(3) (a) Bebout, D. C.; Stokes, S. W.; Butcher, R. J. Inorg. Chem. 1999,
38, 1126−1133. (b) Cordes, D. B.; Hanton, L. R. Inorg. Chem. Commun.
2005, 8, 967−970. (c) Dudev, T.; Lim, C. Chem. Rev. 2014, 114, 538−
556.
(4) (a) Dowling, C.; Murphy, V. J.; Parkin, G. Inorg. Chem. 1996, 35,
2415−2420. (b) Pettinari, C.; Marinelli, A.; Marchetti, F.; Ngoune, J.;
Galindo, A.; Alvarez, E.; Gom
́
ez, M. Inorg. Chem. 2010, 49, 10543−
10556.
(5) Kumar, A.; Singh, J. D. Inorg. Chem. 2012, 51, 772−774.
(6) McMorran, D. A.; Steel, P. J. Tetrahedron 2003, 59, 3701−3707.
(7) Horiuchi, S.; Kumai, R.; Tokura, Y. J. Am. Chem. Soc. 2013, 135,
4492−4500.
(8) Waterland, M. R.; Kelley, A. M. J. Chem. Phys. 2000, 113, 6760−
6773.
(9) (a) Trautner, F.; Kirsch, R.; Minkwitz, R.; Oberhammer, H. Inorg.
Chem. 2002, 41, 7049−7052. (b) Addison, C. C.; Sutton, D. J. Chem. Soc.
A 1966, 1524−1528.
(10) (a) Pyykko, P. Chem. Rev. 1988, 88, 563−594. (b) Bayler, A.;
̈
Schier, A.; Bowmaker, G. A.; Schmidbaur, H. J. Am. Chem. Soc. 1996,
118, 7006−7007.
(11) (a) Noxon, J. F.; Norton, R. B.; Marovich, E. Geophys. Res. Lett.
1980, 7, 125−128. (b) Platt, U.; Winer, A. M.; Biermann, H. W.;
Atkinson, R.; Pitts, J. N. Environ. Sci. Technol. 1984, 18, 365−369.
(12) Barnett, S. A.; Champness, N. R. Coord. Chem. Rev. 2003, 246,
145−168.
5889
dx.doi.org/10.1021/ic500746e | Inorg. Chem. 2014, 53, 5887−5889