Formation of ionic liquids of divalent metal complexes comprising N?alkylethylenediamines and the solvation of the nickel(II) complexes

Article abstract of DOI:10.1016/j.molliq.2018.07.124

A series of divalent N?alkylethylenediamine (alkyl?en) metal(II) (alkyl = hexyl, 2?ethylhexyl, octyl, dodecyl; metal = Ni, Cu, and Zn) complexes was prepared and their phase behavior was studied using differential scanning calorimetry. This kind of metal complexes is very useful for systematically investigating the relationship between phase behavior and molecular structures. It was found that several of the zinc(II) and nickel(II) complexes form room-temperature ionic liquids (RTILs), despite the divalent cation. Although the solid-to-liquid transition temperatures of metal(II) complex-based ILs are typically higher than those of the more common monovalent ionic liquids (ILs), they are dependent on the nature and combination of the metal(II) ions, alkyl chains, and counter anions. The zinc(II) complexes coupled with weakly coordinating bis(fluorosulfonyl)amide (FSA) or bis(trifluoromethanesulfonyl)amide (Tf₂N) anions have significantly lower melting points, which is attributable to the longer distance between the zinc(II) ions and the counter anions upon the formation of tetrahedral bis(alkyl?en)zinc(II) complexes as compared to the corresponding distances in NO₃ and trifluoroacetate (TFA) complexes. The correlation of melting points with the molecular structures of the zinc(II) complexes is similar to that for the silver(I) alkyl?en complexes. The tris(alkyl?en)nickel(II) complexes coupled with Tf₂N counter ions do not readily solidify and have glass transition temperatures below 0 °C, whereas the corresponding bis complexes have much higher melting points despite the counter-ions residing in the outer-spheres of the nickel(II) ions. The interactions of NO₃, TFA, FSA, and Tf₂N anions with the nickel(II) ions of the bis(alkyl?en) complexes and their solvation behaviors were also studied in organic solvents using visible absorption spectroscopy based on the structures of the neat states. The characteristic solvation behaviors of the nickel(II) complexes were rationalized in terms of the counter anions and solvents.

Full text of DOI:10.1016/j.molliq.2018.07.124

Accepted Manuscript
Formation of ionic liquids of divalent metal complexes
comprising N'alkylethylenediamines and the solvation of the
nickel(II) complexes
Haruka Yasuda, Chisayo Nakayama, Masayasu Iida
PII:
DOI:
Reference:
S0167-7322(18)33433-0
doi:10.1016/j.molliq.2018.07.124
MOLLIQ 9445
To appear in:
Journal of Molecular Liquids
Received date:
Revised date:
Accepted date:
4 July 2018
29 July 2018
31 July 2018
Please cite this article as: Haruka Yasuda, Chisayo Nakayama, Masayasu Iida , Formation
of ionic liquids of divalent metal complexes comprising N'alkylethylenediamines and the
solvation of the nickel(II) complexes. Molliq (2018), doi:10.1016/j.molliq.2018.07.124
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ACCEPTED MANUSCRIPT
Formation of Ionic Liquids of Divalent Metal Complexes Comprising
N-alkylethylenediamines and the Solvation of the Nickel(II) Complexes
Haruka YASUDA, Chisayo NAKAYAMA, and Masayasu IIDA*
Department of Chemistry, Faculty of Science, Nara Women’s University, Kitauoya-nishi-machi, Nara 630-8506,
Japan
e-mail: iida@cc.nara-wu.ac.jp
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Abstract
A series of divalent N-alkylethylenediamine (alkyl-en) metal(II) (alkyl = hexyl, 2-ethylhexyl, octyl,
dodecyl; metal = Ni, Cu, and Zn) complexes was prepared and their phase behavior was studied using
differential scanning calorimetry. This kind of metal complexes is very useful for systematically investigating
the relationship between phase behavior and molecular structures. It was found that several of the zinc(II) and
nickel(II) complexes form room-temperature ionic liquids (RTILs), despite the divalent cation.
Although the solid-to-liquid transition temperatures of metal(II) complex-based ILs are typically higher
than those of the more common monovalent ionic liquids (ILs), they are dependent on the nature and
combination of the metal(II) ions, alkyl chains, and counter anions. The zinc(II) complexes coupled with
weakly coordinating bis(fluorosulfonyl)amide (FSA) or bis(trifluoromethanesulfonyl)amide (Tf2N) anions
have significantly lower melting points, which is attributable to the longer distance between the zinc(II) ions
and the counter anions upon the formation of tetrahedral bis(alkyl-en)zinc(II) complexes as compared to the
corresponding distances in NO3 and trifluoroacetate (TFA) complexes. The correlation of melting points with
the molecular structures of the zinc(II) complexes is similar to that for the silver(I) alkyl-en complexes. The
tris(alkyl-en)nickel(II) complexes coupled with Tf2N counter ions do not readily solidify and have glass
transition temperatures below 0 ºC, whereas the corresponding bis complexes have much higher melting
points despite the counter-ions residing in the outer-spheres of the nickel(II) ions. The interactions of NO3,
TFA, FSA, and Tf2N anions with the nickel(II) ions of the bis(alkyl-en) complexes and their solvation
behaviors were also studied in organic solvents using visible absorption spectroscopy based on the structures
of the neat states. The characteristic solvation behaviors of the nickel(II) complexes were rationalized in terms
of the counter anions and solvents.
Key words: ionic liquids; divalent metal complexes; phase behavior; chelate amines; solvation of
alkylethylenediamine nickel(II) complexes
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1. Introduction
Ionic liquids (ILs) and ionic liquid crystals containing transition-metal ions show a variety of specific properties
that are not exhibited by conventional organic compounds. [1-10] It is of particular note that the ILs of metal
complexes (met-ILs) are regarded as soft materials in which the metal ions are highly condensed in nano-sized
regions. Therefore, met-ILs are expected to become increasingly utilized as functional soft materials in fields such
as nanochemistry, electrochemistry, and magnetochemistry.[11-23] It is important for the rational molecular design
of met-ILs to be able to relate their phase behavior to the molecular structures of the ligands and the nature of the
central metal ions, and counter ions, as well as the relationships between them. However, systematic studies on
these relationships have not been performed. From this viewpoint, a suitable series of alkylethylenediamine
(alkyl-en) metal complexes would be a model upon which to perform intensive studies on the above relationship
based on the different molecular moieties and structures within the series. In previous studies, our research group
has prepared several alkyl-en silver(I) complexes related to the formation of silver(0) nanoparticles (AgNPs)
revealing the close correlation between their molecular structures and phase behavior. [24-26] Silver(I) complexes
as well as alkyl-en monoprotic ILs (PILs) form ILs more easily than multivalent compounds owing to their smaller
inter-ionic interactions. For divalent metal complexes, although the melting points are generally higher, several
structural factors, e.g., molecular sizes and geometries are expected to act more synergistically on the formation of
ILs compared to those for monopositive silver(I) complexes and monoprotonated PILs.
In the present study, several divalent metal (Ni(II), Cu(II), and Zn(II)) alky-en complexes coupled with the
—
counter anions of NO3 , TFA—, FSA—, and Tf2N— were prepared and their phase behaviors were studied using
differential scanning calorimetry (DSC). The preparation of the [Zn(NO3)2(octen)2] complex and its X-ray
crystallographic analysis has been previously reported in a short communication by us [27], revealing that the NO3
anions directly coordinate to the zinc(II) ion to form an octahedral complex. However, the phase behavior of most
divalent metal complexes comprising alkyl-en moieties in the neat states has not been studied, although it is
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expected to be dependent on various molecular structures. ILs are generally defined as normal salts with melting
points below 100 ºC, with room temperature ionic liquids (RTILs) being attractive as functional soft materials.
The melting points of the metal alkyl-en complexes studied herein cover a wide range, from below 0 ºC to 200 º

C. Therefore, a series of these complexes is a suitable system for studying the relationship between molecular
structures and melting points.
The structures of nickel(II) complexes can be mainly categorized into two types, i.e., those with octahedral and
those with square-planar geometry. These two structure types can be clearly discriminated by their visible
absorption spectra and can be expected to exhibit distinct phase behaviors. Based on their neat states, the structures
of the bis(alkyl-en) nickel(II) complexes are also studied in terms of their solvation behaviors in different organic
solvents, with particular attention being given to the differences between the complexes with strongly coordinating
anions and those with weakly coordinating anions.
2. Experimental
2.1 Materials. The N-hexyl-, N-2-ethylhexyl-, N-octyl-, and N-dodecyl-ethylenediamines (C6, C2C6,
C8, and C12, respectively) were prepared by reacting the corresponding alkyl bromides (purchased from
Tokyo Chemical Industry) with ethylenediamine (guaranteed reagent, purchased from Wako Pure Chemical
Industry) 1:5 molar ratio according to the literature, [28] followed by distillation at 1 kPa. The MX2 salts (M =
Ni(II) Cu(II), or Zn(II), and X = Tf2N or TFA) were isolated as follows. 5×10-3 mol NiCO3 (594 mg) CuO
(398mg), or ZnO (407mg) was dispersed in 10 mL water, and 102 mol HTf2N or HTFA (acidic form)
dissolved in 10 mL water was gradually added, and then the mixture was stirred until the solid was almost
dissolved. After the solution was passed through #5C filter paper, the filtrate was condensed by an evaporator
at 40 ºC until a small amount of solid was precipitated. The solution with precipitates was then cooled in a
refrigerator; the solids were collected and dried in vacuo. The dehydration was performed with Hitachi VR-16
high-vacuum pump over 48 h. Ni(TFA)2 was obtained as a trihydrate and the other MX2 salts were as an
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anhydride. The purities of the isolated salts were confirmed by CHN elemental analyses.
The NO3 complexes of [MX2(alkyl-en)2] (or [M(alkyl-en)2]X2) for Ni(II), Zn(II), and Cu(II) were prepared
following previously reported procedures and obtained as solids.[27] The FSA salts were obtained from the
nitrate (for Ni(II), Zn(II), and Cu(II)) of each complex by the counter ion exchange using potassium
bis(fluorosulfonyl)amide (KFSA) (purchased from Mitsubishi Materials Electronic Chemicals Co.) following
the way for the Ag(I) complexes [26], whereas the solids of FSA salts were obtained as described below as for
the other solids. The alkyl-en complexes were prepared by the reaction of the MX2 (X = Tf2N or TFA) salt
with alky-en in 1:2.05 molar ratio for the [M(alkyl-en)2]X2 salt or 1:3.05 molar ratio for the [Ni(alkyl-en)3]X2
salt following the previously reported procedure [24-27, 29-31]. The solvents used for the syntheses were
methanol or methanol/dichloromethane mixtures. To solidify the complexes, dichloromethane or diethyl ether
was gradually added to the methanol solutions. For the [Ni(alkyl-en)2]X2 complexes where X = Tf2N or FSA,
the solids obtained from the solution are purple blue and by dryness under high vacuum they changed to
hygroscopic orange solids. We herein used the orange solids, which gradually change to the color of purple
blue by an absorption of moisture. In the case of the RTILs of the [Ni(alkyl-en)3]X2 and of some zinc(II)
complexes, the residual liquids of the ligands as well as the organic solvents in the products were completely
removed using Hitachi VR-16 high-vacuum pump over a few days.
The organic solvents used for the solvation study of the nickel(II) complexes were dichloromethane (DCM),
diethyl ether (Et2O), acetone, acetonitrile (MeCN), methanol (MeOH), ethanol (EtOH),
N,N-dimethylformamide (DMF), and dimethylsulfoxide (DMSO). All the solvents were purchased from Wako
Chemicals, dehydrated by passing over molecular sieves (3A 1/16, Wako Chemicals), and then once distilled
before use.
2.2 Measurements.
The purities of the obtained metal complexes were determined using CHN elemental analyses with a
Perkin Elmer 2400II Elemental Analyzer (equipped in the division of Instrumental Analysis of Nagasaki
University) andC NMR spectra with a JEOL EX-270 spectrometer for the Zn(II) complexes. The CHN
elemental analyses for the alkyl-en metal(II) complexes used in the study of the phase behavior are listed in
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Table S1. The concentrations of the nickel(II) complexes in organic solvents are expressed as molality unit
(mol kg-1). The solid-liquid transitions for the newly synthesized metal(II) complexes were examined using
differential scanning calorimetry with a Shimadzu DSC-50 equipped with an LTC-50 apparatus for
low-temperature measurements; the temperature was increased at a constant rate (10 K min-1) from 80 to
200 or 250 (only for the Ni(II)C2C6(NO3) complex) °C. The transition behavior was observed only for the
temperature-increasing process. The sample (2–3 mg) was placed in an aluminum pan and the measurement
was performed under nitrogen at a flow rate of 20 cm3 min. Indium metal (purchased from Wako Chemicals)
was used as a standard for the transition temperature and enthalpy. The melting point was defined as the
minimum point in the endothermic curve and the glass transition point follows the common definition as shown
in Fig. 2. [24,25] Several systems showed double melting points on the DSC profiles as described below. In
order to confirm the transition temperatures from solid to liquid, we followed the melting processes using
polarizing microscope (Olympus BH2).
The TG/DTA measurements were performed to detect the thermal stabilities of the complexes using a
Shimadzu DTG-60/60H with the temperature increasing at a constant rate (10 K min-1) from room
temperature to 300 °C.
2.3 Electronic spectral measurements. The VIS absorption spectra of the nickel(II) complexes in solution
were acquired in the 400–1200 nm wavelength range at 0.2 nm intervals using a JASCO V-670iRM
(single-beam) instrument at ambient temperature. The precision of the wavelength was ±0.3 nm. The cell
length was 0.05, 0.2, or 1.0 cm. The concentrations were 0.05 mol kg unless otherwise stated and the molar
extinction coefficient () is expressed in molkg cm unit. For the solid samples, the spectra were measured
by the diffuse reflectance method with a JASCO ISN-723 (integrating sphere) attachment for the typical
samples. The rough measurements to decide whether the octahedral or square-planar geometry for the
nickel(II) complexes in the solid states were performed by nujol mull method [32] in the 400800 nm
wavelength region.
3. Results and discussion
Characterization of the alkyl-en metal(II) complexes.
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As shown in Table S1 the calculated (anhydride) and experimental CHN elemental analysis values for
most of the complexes agree within 0.3% with several agreeing within 0.5%. The exceptions are
[Ni(doden)2](FSA)2 (the differences are 3.34 % and 0.55 % for C and H, respectively) and [Ni(doden)2](NO3)2
(the differences are 1.16 and 0.53 % for C and H, respectively).
The purities of the diamagnetic Zn(II) complexes were also assessed using 13C NMR spectra on the basis
of previous studies. [24, 25, 33]
Coordination of the counter anions in the neat-state alkyl-en nickel(II) complexes.
The nickel(II) complexes in the neat states typically exhibit one of two colors depending on the counter
anions: with the NO3 and TFA salts being purple-blue and the FSA and Tf2N complexes being orange. The
absorption spectra of the solids were measured using the diffuse reflectance method and those of
[Ni(hexen)2(TFA)2] and [Ni(hexen)2](Tf2N)2 are given as typical examples in Fig. 1. These results clearly
indicate that the geometry of the main component of the TFA salt is octahedral and that of the Tf2N salt is
square planar. [34,35] The absorption spectra of the neat nickel(II) complexes were also measured using a
nujol mull method in order to classify the complexes into two types depending on whether the counter anion is
directly coordinating to the nickel(II) ion or not. These results indicate that the NO3 and TFA anions directly
coordinate to the nickel(II) ion to form octahedral complexes while the FSA and Tf2N anions do not directly
coordinate to the nickel(II) ion, forming instead square- planar complex. These results are also related to the
phase behavior, as discussed below.
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Fig. 1. Diffuse reflectance AB spectra of solid bis(hexen) nickel(II) complexes of Tf2N ( — ) and TFA (---)
salts.
Phase behavior
The features of the DSC curves can be classified into the following four types of solid-liquid
transitions. 1) a crystal-to-liquid transition above 100 °C. 2) a crystal-to-liquid transition with double
melting peaks. 3) a glass transition below 0 °C. 4) a crystal-to-liquid transition below 100 °C. Typical DSC
curves for each type of the Zn(II) complexes are shown in Fig. 2. The solid-to-liquid phase transition
temperatures (TMP) are listed in Table 1 together with the transition enthalpies (in parentheses). Here, glass
transitions are indicated by (g) and the other temperatures are melting points or double melting points. The
TMP values for the analogous monopositive silver(I) complexes and protic alkyl-en ILs reported previously
are also listed for comparison. Double endothermic peaks are observed in many of the NO3 and TFA
metal(II) complexes (Fig. 2b, Table 1). For these complexes, it was confirmed using polarized microscope
(PM) that melting occurs at the higher of the two TMP values’ while the PM image does not change across
the lower transition temperatures. This kind of double minimum has been reported as an onset of liquid
crystalline formation [36,37]. In the present system, formation of liquid crystals was not clearly detected.
However, formation of liquid crystals was well characterized for several of the silver(I) complexes. [38]
Herein, for simplicity we take the higher temperature in the double transitions as the melting point.
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Before addressing the phase transitions, we obtained the TG profiles of each complex. The TG profiles of
the hexyl(C6) and ethylhexyl(C2C6) complexes are presented in Fig. 3 as typical examples. The temperatures
at which significant weight losses are observed are higher than the transition temperature (Ttr) except for
Ni(II)-C2C6(NO3) (Fig. 3b) and are generally dependent on the counter anions in the order TFA < NO3 < FSA
 Tf2N. The lower thermal stability of the TFA salt has also been reported for the monopositive PILs of the
alkyl-en derivatives.[33] In terms of the alkyl chains, the ethylhexyl (C2C6) complex tends to be degraded at a
lower temperature compared to the other alkyl derivatives. However, the Ni(II)-C2C6(NO3) complex is
unique, having a particularly-higher melting point (above 200 °C) and a lower decomposition temperature.
(Fig. 3b)
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Fig. 2. Typical DSC curves for the zinc(II) complexes. (a) bis(hexen)Zn(II) nitrate, (b) bis(ethexen)Zn(II) nitrate,
(c) bis(ethexen)Zn(II) Tf2N salt, and (d) bis(doden)Zn(II) Tf2N salt.
Table 1. Phase transition temperatures (TMP/ °C) and enthalpies (HMP/ kJ mol-1, in parentheses)
Metal(II)
/ alkyl
counter anion
(a) Zn(II)
C6
NO3
TFA
FSA
Tf2N
144.1 (30.4)
61.7 (10.5), 122.8 (13.8)
61.8 (1.8), 87.3 (17.2)
68.6 (8.7), 102.3 (13.2)
71.7 (40.6), 118.5 (46.7)
TFA
48 (g)
44 (g)
37 (g)
44 (g)
52.8 (65.6)
Tf2N
C2C6
C8
141.5 (4.5), 168.6 (23.2)
150.4 (35.5)
45 (g)
64.8 (27.9)
58.8 (67.0)
FSA
C12
83.4 (55.9), 139.0 (36.1)
NO3
(b) Ni(II)
C6
137.1 (39.1)
94.7 (8.1), 154.2 (27.9) a
98.6 (23.5)
85.9 (22.7)
tris(C6)
C2C6
C8
7 (g)b
237.9 (39.4)
93.2 (1.2), 124.6 (26.2)
78.1 (17.7)
76.0 (28.4)
90.0 (14.8)
37 (g)
123.8 (7.2), 143.8 (35.4) 90.0 (8.1), 125.5 (20.8)
118.1 (43.3)
tris(C8)
C12
93.0 (73.5)
124.3 (53.6)
TFA
114.3 (55.0)
FSA
101.9 (43.5)
a : glass transition
(c) Cu(II)
NO3
Tf2N
C6
122.7 (5.4),179.0 (44.3)
111.4 (16.8),168.4 (20.5) 118.0 (43.3) 109.3 (48.8)
C2C6 130.9 (18.0)
48.3 (6.9), 88.9 (4.5)
86.2 (22.8)
108.2 (53.6)
84.5 (38.1)
C8
171.8 (2.8), 188.7 (35.0)
98.8 (31.5), 170.7 (36.4)
85.8 (20.8), 151.5 (28.2)
85.5 (43.6), 143.8 (39.0)
122.8 (43.6)
C12
132.5 (63.6) 44.8(18.0), 100.9(31.4)
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Ag(I) (ref. 26)
C6
NO3
44
TFA
11
FSA
Tf2N
75 (g)
72 (g)
87 (g)
78 (g)
C2C6
54 (g)
7 (g)
C8
48
59
30
45
26
37
10 (g)
38
C12
H+ (ref. 25)
Tf2N
C6
16
C2C6
C8
70 (g)
18
C12
30
a Underlined temperatures are identified as melting points from solid-to-liquid using PM. (see text)
b Glass transition temperatures.
Fig. 3. Thermo-gravimetric curves for the typical systems of the [Ni(alkyl-en)2X2] complexes depending on the
counter anions (X = NO3, TFA, FSA, and Tf2N) for (a) alkyl = hexyl, and (b) alkyl = 2-ethylhexyl.
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Fig. 4. Melting points or glass transition temperatures (TMP) vs. counter anions for each bis(alkyl-en)
metal(II) complex. alkyl : C6 = hexyl, C2C6 = 2-ethylhexyl, C8 = octyl, C12 = dodecyl.
For the dipositive alkyl-en complexes, the effect of molecular structures on the phase transition is more
irregular compared to that for the monopositive alkyl-en silver(I) complexes and PILs, since several
intermolecular interactions act on the phase behavior synergistically due to the stronger ionic interactions. We
can, however, identify several significant trends based on the principle of intermolecular interactions. The
results are interpreted in terms of the two molecular components, i.e., the anion and the alkyl-chain of the
cation, for each metal(II) complex. Figures 4 and 5 illustrate the respective effects on the phase transition
temperatures (TMP) for each metal complex.
(1) Anion effects (Fig. 4): The abscissa in Fig. 4 is taken as an approximate indicator of the anion size.
(a) Zn(II) complexes. (Fig. 4(a)) Among the three metal(II) complexes, the zinc(II) complexes show the clearest
trends concerning the anion effect, i.e., the melting points decreases with an increase in anion size and decrease in
TMP is significant for the weakly coordinating FSA and Tf2N anions. A similar trend is seen in the silver(I) alkyl-en
complexes as previously reported (Table 1). It is remarkable that, despite the dipositive charge, the Zn(II)-C6(FSA),
C6(Tf2N), C2C6(FSA), C2C6(Tf2N), and C8(Tf2N) complexes form RTILs as also observed in the corresponding
Ag(I) complexes. X-ray crystallographic analysis of the Zn(II)-C8(NO3) complex revealed that the NO3 anions
directly coordinate to the zinc(II) center.[27] Conversely, although X-ray structural study has not been performed
for the FSA and Tf2N complexes, the FSA and Tf2N anions most likely do not coordinate to the zinc(II) center.
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Here, comparing the Zn(II) system to the Ni(II) system is useful, in that, as described above, FSA and Tf2N anions
do not directly coordinate to the nickel(II) ion while the NO3 and TFA anions coordinate to the nickel(II) ion to
form octahedral complexes in the neat states. Therefore, the large difference in the melting points between the NO3
and TFA group and the FSA and Tf2N group for the zinc(II) complexes can be attributed to the counter anions
directly coordinating to the zinc(II) ion or not. Such significantly lower melting points for shorter alkyl-chained
complexes have also been observed in the Ag(I) complexes coupled with FSA and Tf2N anions. Thus, the structural
bulkiness of the anion seems more effective on reducing melting points when coupled with shorter alkyl-chained
cations (Table 1)[26]. Low melting points for tetrahedral and bulky zinc(II) complexes are, furthermore, similarly
observed for imidazolium zinc(II) complexes, as reported by Lin et al.[3, 4]. In the present study, the significantly
low melting points for the zinc(II) complexes are achieved by the combination of bulky tetrahedral alkyl-en ligands
and weakly coordinating bulky anions. A similar trend has been reported for imidazolium nickel(II) and cobalt(II)
complexes. [20, 23]
A remarkable difference between the Ag(I) and Zn(II) ions is observed for the C2C6(NO3) complexes, where the
zinc(II) complex has a higher melting point than that of the other alkyl(NO3) complexes while the Ag(I) complex
has a remarkably low melting point. For the nickel(II) C2C6(NO3) complex, a similarly high melting point to that
of the zinc(II) complex is observed, as discussed below.
(b) Ni(II) complex. Figure 4(b) shows that, except for the two NO3 complexes (C6 and C12), the melting point
decreases with an increase in anion size. As described above, the geometrical structure of the nickel(II) complex in
the neat state is clearly dependent on the counter anion, with the NO3 and TFA salts being octahedral and the FSA
and Tf2N salts being square planar. However, Fig. 4(b) shows that, even for the FSA and Tf2N complexes, the
decrease in melting point is not so large as that observed for the zinc(II) FSA and Tf2N complexes (Fig. 4(a)). This
difference between the nickel(II) and zinc(II) complexes indicates that, although the FSA and Tf2N counter anions
are in the outer-sphere of the nickel(II) complexes as well, these anions should be located on the z-axis of the
square-planar complex, and that the distance between the nickel(II) ion and the anions should not be as large as
those for the non-coordinating anions in the tetrahedral bis(alkyl-en) zinc(II) complexes. From this viewpoint, we
can also understand the appreciably lower glass transition points (– 7 and – 37 °C, respectively. See Table 1) of the
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tris(C6 and C8) nickel(II) Tf2N anion complexes. The remarkable difference between the tris and bis complexes is
attributable to the significantly larger distance between the nickel(II) ion and the counter anion in the tris complex.
The inverse relationship concerning the anion size between the NO3 and TFA salts in the C6 and C12
complexes is remarkable and can be explained in terms of the stronger interaction for the nickel(II) ion and
the TFA anion rather than that for the NO3 anion, which is also observed in solution (as discussed in the next
section). However, the modes of interaction for the NO3 and TFA anions with the metal(II) ion seem to be
somewhat complicated, as evidenced by the double endothermic peaks that are observed for some complexes
(Fig. 2 and Table 1). Furthermore, it is not clear why the inverse relationship occurs in the two unique
nickel(II) complexes. Although CHN elemental analysis indicates poorer purities for the C12(NO3) and
C12(FSA) complexes, only the C12(NO3) complex shows irregularly lower melting points. (Fig. 4b)
Furthermore, it is possible that the higher melting points of the C12 complexes containing the FSA and
Tf2N counter anions than that for the NO3 salt may be due to the stronger van der Waals interactions of the
longer alkyl chain with the larger anion.
The higher melting point of the C2C6(NO3) complex is also seen for the zinc(II) complex. In the divalent
complexes, the bulkiness of the ethyl-branch is not always sufficient to lower the melting point but rather in
some cases enhances the intermolecular (van der Waals) interactions. This trend is in contrast to that observed
in the mono positive Ag(I) and proton analogs (Table 1) and can be interpreted in terms of the interaction
mode, where the van der Waals interactions due to the ethyl branch may become synergistically effective on
the interactions in combination with the stronger electrostatic interaction. The Ni(II)C2C6(NO3) complex is
unique: although it has a significantly higher melting point (Fig. 4(b)), it is gradually decomposed at a lower
temperature than the other complexes. (Fig. 3(b))
(c) Cu(II) complex. An increase in the anion size tends to decrease in the melting point (except for the
C2C6(Tf2N) complex) as demonstrated in Fig. 4(c). This trend indicates that the electrostatic interaction
dominates the anion size effect. The increase in the melting point of the C2C6(Tf2N) complex can be
attributed to the van der Waals interactions between the bulky cation and the bulky anion. It is notable that the
C2C6(NO3) complex shows a different trend to those of the other metal(II) alkyl-chain complexes, i.e., the
14

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ethyl branch lowers the melting point. This trend is similarly observed in the monopositive Ag(I) complex and
PIL (Table 1), where it is suggested that the ethyl branch may cause a steric hindrance for the intermolecular
interactions. However, it is not clear why there is a difference between the copper(II) complexes and the other
divalent metal complexes. The ambivalent factors (steric hindrance and the van der Waals attraction force) of
the ethyl branch may affect the melting point in a complex manner in combination with the counter anions and
molecular structures.
Fig. 5. Melting point or glass transition temperature (TMP) vs. alkyl-chain length of the ligand in each
bis(alkyl-en) metal(II) complex with four different counter anions.
Alkyl chain length of the cations. Figure 5 shows the dependence of phase transition temperatures (TMP)
on alkyl chain length for each metal complex. The effect of the alkyl chain on the molecular interaction and
on the melting point becomes more irregular in comparison with those for the alkyl-en monopositive
silver(I)-ILs, PILs, and common AILs (Table 1 [25, 26, 39-41]) where an increase in TMP with an increase in
the alkyl-chain length (from C6 to C12) is observed, which is basically attributable to an increase in the van
der Waals interaction between the cationic units. Conversely, if the electrostatic force between the cation and
the anion rather the van der Waals interactions dominates the overall interaction, the phase transition
temperatures vs. alkyl chain length plot present a negative slope due to a decrease in the Coulombic force
[40,41]. This trend can also be interpreted in terms of a decrease in the entropic effect, i.e., as the ordering of
molecular assemblies increases with increasing alkyl-chain length, the stability of the molecular assemblies
15

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decreases, leading to a lower melting point [39]. Figure 5 demonstrates that, although the two opposite driving
forces affect TMP in a complex manner, there is a rough trend for each metal(II) system, i.e., the lower melting
complexes comprising FSA and Tf2N counter anions tend to have positive slopes while the higher melting
complexes comprising the smaller NO3 and TFA counter anions tend to have a negative slope. From the
standpoint of the intermolecular interactions, this trend is reasonable since the former system tends to be van
der Waals interaction-driven and the latter system tends to be electrostatic interaction-driven.
The other outstanding feature concerning the alkyl chain effect is observed for the C2C6 derivatives as
described above.
Interactions of the NO3, TFA, FSA, and Tf2N anions with the nickel(II) ion and the resultant solvation
behavior of the nickel(II) complexes in organic solvents
As described above, there are two structures for the neat bis(alkyl-en) nickel(II) complexes depending on
the counter anions. They can be unambiguously distinguished by their visible absorption spectra, so we can
easily compare the coordination modes between the neat state and the solution state. Solvatochromism of
[Ni(acac)(tmen)]+ (acac = acetylacetonato, tmen = tetramethylethylenediamine) with the weakly coordinating
—
—
counter anions ClO4 or BF4 has been studied in various solvents, revealing that this nickel(II) complex is
useful as a good donor number (DN) probe for common molecular solvents.[34,35,42,43] In those studies, in
order to focus on the solvation of the nickel(II) complex, the counter anions chosen coordinate weakly to the
nickel(II) ion. In the present study, bis(alkyl(C6, C2C6, C8, and C12)-en) nickel(II) complexes with the anions
(NO3, TFA, FSA, and Tf2N) were dissolved in various organic solvents as listed in the experimental section (or
below). Their visible absorption spectra were measured and the results correlate with solvation behavior as
well as in terms of the counter-anion binding to the cationic complex except for several poorly soluble
complexes. Several of the absorption spectra for the octahedral [Ni(alkyl-en)2X2] complexes (where X is the
monodentate anion or solvent) have been previously studied [30] on the basis of the established absorption
spectra for nickel(II) complexes. [34,35,44-48] These results are useful for this study, whereas the
discrimination between the cis and trans geometries using the first absorption peak (960 nm) is not discussed
16

ACCEPTED MANUSCRIPT
here due to ambiguities. However, there is a significant difference in the absorption spectra of

[Ni(alkyl-en)2X2] for X = NO3 , TFA, and molecular solvent for the first absorption peak (max is around 960
nm) as shown in Fig. 6, which shows the spectrum of the hexyl complex as a typical example. These
absorption spectra depend significantly on the counter anions and solvents while being almost independent of
the alkyl moieties of the ligands. In the two anion-coordinating systems (NO3 and TFA) the intensity of the
first absorption peak is significantly smaller than that of the solvated systems and also present splitting, as
exemplified in Fig. 1 and Figs. 6(a) and 6(b) (in DCM).
Fig. 6. Visible absorption spectra of the [Ni(hexen)2]X2 complexes in solution for each counter anion (X =
Anion) depending on the solvents.
Conversely, for the second absorption peak (max is around 570 nm, max = 6–7 kg cm1mol1) the peak
intensities for [Ni(solv)2(alkyl-en)2]2+ and [NiX2(alkyl-en)2] are comparable (X = NO3 or TFA) (Figs. 6(a) and
(b), and ref. 30). Therefore, in the following discussion, we focus on the relative intensity and shape of the first
absorption peak compared to those for the neat state or in DCM (for the TFA salt) and in DCM (for the NO3
salt), and use them as a rough indicator for the extent of solvation of the [NiX2(alkyl-en)2] complex (where X
= NO3 or TFA), i.e., the [Ni(solv)2(alkyl-en)2]2+/[NiX2(alkyl-en)2] ratio. For the FSA and Tf2N salts, on the
other hand, as the neat state shows no coordination of counter anions to form the square-planar complex (the
spectra of the Tf2N salt for the neat states and in DCM are shown in Fig. 1 and in Fig. 6(c), respectively), the
extent of solvation can be easily monitored by using the peak height of the single large peak for max = 70 kg
cm1mol1 at max = 460 nm, where the absorption peak becomes much weaker upon solvation to form an
17

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octahedral complex. (as exemplified in Figs. 6(a)–(c), in comparison with Fig. 6(c) (DCM))
As described in the experimental section, all the measurements except for some systems in DCM, Et2O, and
acetone were performed at 0.05 mol kg. The absorption spectra are largely dependent on the relative
dielectric constant (r) as well as on the DN of the solvents. The two parameters (r and DN) for each solvent
are as follows (in the order of increasing dielectric constant (r)). Et2O: r = 4.3, DN = 19.2; DCM: r = 8.9,
DN = 1; acetone: r = 20.7, DN = 17; EtOH: r = 24.5, DN = 20; MeOH: r = 32.7, DN = 19; MeCN: r = 36.0,
DN = 14.1; DMF: r = 36.7, DN = 26.6; DMSO: r = 46.7, DN = 29.8. [49, 50] The results allowed the
solvents to be categorized into the following five groups (a – e) based on the spectral solvation modes for the
nickel(II) complexes.
(a) DCM: This has the smallest DN and the second lowest r among the solvents used here resulting in the
lowest solvation. Among the nickel(II) complexes, C6(NO3) (0.01 mol kg), C6(FSA) (2 mmol kg), and
C8(FSA) (2 mmol kg) were measured at lower concentrations due to solubility limitations. The other
complexes were measured at 0.05 mol kg. For the alkyl-en complexes, the salts show very similar
absorption spectra to that in the neat state as shown in Figs. 1 and 6. Therefore, direct solvation of the
nickel(II) ion hardly occurs. Thus, the spectra of the nickel(II) complexes in DCM can be taken as
representative of zero solvation.
(b) Et2O: Although the DN of this solvent is larger than those of MeCN and acetone and is comparable to that
of MeOH, its r value is the smallest among the solvents used here. Thus, the counter anions (either inner
sphere or outer sphere) hardly dissociate from the nickel(II) ion in this solvent. As a result, the NO3 and
FSA salts except for C2C6(FSA) do not dissolve in this solvent, whereas the TFA and Tf2N salts are
slightly soluble ( 1–5 mmol kg-1). This solubility difference may be due to whether a hydrophobic CF or

3
ethyl branch is present or not. For the TFA salts, the absorption spectra (measured at 1–5 mmol kg-1)
indicate coordination of the TFA anion to form the [Ni(TFA)2(alkyl-en)2] complex, despite the low
concentrations, whereas for all the Tf2N complexes and the C2C6(FSA) complex the major component has
square-planar geometry (max = 35–40 kg cm1mol1 at max = 460 nm). However, the significantly smaller
18

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
max values than those in DCM (max = 70–80 kg cm1mol1 at max = 460 nm) suggest that Et2O may
slightly coordinate to the nickel(II) ion and thus the ratio of the square planar complex is less than those for
the neat state and in DCM.
(c) acetone : As both the r and DN values for acetone are intermediate among the solvents used here, the NO3,
FSA, and Tf2N salts show characteristic behaviors in this solvent. The TFA salt, on the other hand, present
a similar absorption spectrum as that in DCM or Et2O, indicating that most of the TFA anions still
coordinate to the nickel(II) ion in solution. The spectrum of the NO3 complex is different from that in
DCM and is like that in MeCN. (Fig. 6a) The first absorption peak is significantly less intense than the
second (at around 570 nm). This feature is different from those of the solvated [Ni(solv)2(alkyl-en)2]2+ in
polar solvents (group(e), (Figs. 6(a–c) for DMF)) and of the FSA and Tf2N salts in this solvent at infinite
time, as shown in Fig. 7, where the intensity of the first band is larger than that of the second. This
difference in spectral shape indicates that the NO3 counter anion is partially dissociated from the nickel(II)
ion and replaced by acetone.
It is a particularly interesting phenomenon that when the FSA and TF2N salts are dissolved in this solvent,
the color gradually changed from orange to blue. An example of the change in the absorption spectra is
shown in Fig. 7.
Fig. 7. Typical spectral change of the square planar [Ni(octen)2]2+ complex in acetone, where counter anion =
FSA. The concentration is ca. 0.03 mol kg and at 20 °C.
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The spectrum at 15h (approximating infinite time) is similar to that of [Ni(solv)2(alkyl-en)2]2+ (where solv
is MeOH, EtOH [30], or DMF(Fig. 6)). (Only the C6(Tf2N) complex exhibits the presence of slight
amount of the residual square planar complex whose absorption appears at 460 nm.) This change clearly
indicates that the square planar bis(alkyl-en) complex is gradually solvated to form the octahedral
[Ni(solv)2(alkyl-en)2]2+. The absorbance at 460 nm was monitored at 20 °C and the —ln(At  A)/(A0  A)
vs. t plot, where At, A, and A0 are the absorbances at time t, an infinite time, and time zero, respectively,
presents a straight line for each system as illustrated in Fig. S-1. The linearity indicates that the reaction is
nearly first-order. Accordingly, the first-order rate constant k [h-1] was determined for each system and the
results are listed in Table 2. Note that the (dod)(FSA)2 system was not included here due to its poor
elemental analysis result.
Table 2. First-order rate constants (k) for the acetone solvation of the square planar bis
(alkyl-en)Ni(II) complexes.
k/ h
FSA
0.37
0.33
0.29
n/a
Tf2N
1.31
1.29
1.22
1.38
alkyl
C6
/
anion =
C2C6
C8
C12
Table 2 reveals a trend concerning the anion effect, i.e., the higher reaction rate for the Tf2N anion system
than that for the FSA system. As both the FSA and Tf2N anions do not coordinate to the nickel(II) ion in
the inner sphere, these anions may have an obstructing effect on the solvation reaction for the
[Ni(alkyl-en)2]2+A outer-sphere complex. This is consistent with the melting point result, because the
larger anion experiences a smaller electrostatic interaction with the nickel(II) ion and may thus obstruct
20

ACCEPTED MANUSCRIPT
solvation of the nickel(II) ion less. On the other hand, the alkyl chain moiety appears to have no effect on
the kinetic stability of the ion-pair.
(d)
MeCN : As this solvent has a larger r value than acetone, the counter anion is more easily
dissociated from the nickel(II) ion, whereas the solvation will be weaker than that by acetone due to the
smaller DN. Typical spectra for the hexyl complex in this solvent are shown in Fig. 6(a–c), demonstrating
that the spectra are largely independent of the alkyl moiety.
For the nitrate (Fig. 6a), the first absorption peak is significantly larger than that in DCM and does not
double. This feature is not observed for the TFA system (Fig. 6b), where the spectra in MeCN and in DCM
are similar. The difference between the NO3 and TFA systems indicates that the TFA anion may more
strongly coordinate to the nickel(II) ion than the NO3 anion does. This is consistent with the results
observed for acetone.
For the FSA and Tf2N salts, no time dependency of the spectra is observed and their shapes (Fig. 6c, Fig.
S2) are close to that at infinite time (15h) in acetone or in solvents having larger DN values, such as MeOH,
EtOH, DMF (Fig. 6(c)) and DMSO. The nickel(II) complex is kinetically more easily solvated by MeCN
to form [Ni(solv)2(alkyl-en)2]2+ than in the less polar acetone (smaller r). However, we can see that the
second absorption peak (at around 520 nm) is broader and shows a blue shift in this solvent compared to
those for the five polar solvents above having larger DN values (exemplified for the DMF system in Fig.
6(c)). This difference indicates that solvation by MeCN of the nickel(II) ion may be weaker than those in
the solvents having larger DN values at equilibrium and that the spectrum contains a contribution by the
square-planar complex. As shown in Fig. S2, the shape of the spectrum somewhat depends on the
alkyl-chain and the hexyl complex (also shown in Fig. 6(c)) exhibits the largest blue shift of the second
absorption peak. Here, it should be noted that although acetone coordinates to the nickel(II) ion through the
nitrogen atom, whether the coordination atom is nitrogen or oxygen does not significantly affect the
spectral shape. For example, [Ni(alkyl-en)3]2+ presents a very similar spectrum to that of the
[Ni(solv)2(alkyl-en)2]2+ complex where ‘solv’ is the (e) group solvent, as previous reported [30].
21

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(e) EtOH, MeOH, DMF, DMSO: These solvents with largerr and DN values all present similar nickel(II)
ion solvation modes. Spectra typical of this group are shown in Fig. 6(a–c) for the hexyl complex in DMF.
A significant difference between the (e) group and the (a)–(d) groups is observed for the first absorption
peak (980 nm) of the TFA salt. (Fig. 6(b)) The intensity of this peak for the (e) group solvents is
comparable to (EtOH) or higher (other solvents) than that of the second peak (570 nm). This result
indicates that, in these solutions, the TFA anion is largely dissociated from the nickel(II) ion and replaced
by the solvents. The first peak height is slightly smaller than that of the second peak in EtOH, having the
smallest r in this group, which indicates that some TFA anions are still coordinated to the nickel(II) ion in
EtOH. Comparing the results for the EtOH (smaller r but slightly larger DN) and MeOH systems suggests
that r governs the coordination of the TFA anion rather than DN. A similarly dominant effect of r over
DN on the solvation of the FSA and Tf2N salts is revealed upon comparing the MeCN and acetone solvents
from the kinetic viewpoint, as described above.
Although most of the complexes with different counter anions are solvated to form Ni(solv)2(alkyl-en)2
in this group of solvents, a small peak at 470 nm, which is attributable to the square-planar configuration,
is presented by some of the FSA and Tf2N complexes (up-arrow in Fig. 6(c) and Fig. S3). This small peak
is observed in the solvents of this group as well as in group (c) even at nearly infinite time (15h, Fig.7).
The magnitude of this peak depends on the alkyl moiety of the metal complexes in the order, C6 > C8 >>
C12 > C2C6, as shown in Fig. S3. This may be the reverse of the order of solvation. The magnitude of this
small peak is larger in MeOH and EtOH than in DMF and DMSO. This trend indicates poorer solvation in
the former two solvents, which is consistent with the solvent parameters (i.e., lower r and DN values).
As the spectra of Ni(DMF)62+ and Ni(DMSO)62+ are significantly different from the present spectra in
DMF and DMSO, respectively, [51,52] most of the alkyl-en ligands remain coordinated to the nickel(II)
ion.
Conclusions
The solid-to-liquid transition temperatures for dipositive bis(alkyl-en)metal(II) (metal = nickel, copper,
22

ACCEPTED MANUSCRIPT
and zinc) and tris(alkyl-en)nickel(II) complexes fall in the wide range of –44 to 200 °C depending on the
molecular moieties. Some bis(alkyl-en)zinc(II) and tris(alkyl-en)nickel(II) complexes are liquids at room
temperature despite the dipositive charge. The distance between the metal ion and the counter anion essentially
governs the transition temperatures in these dipositive complexes. The extent of solvation of the
bis(alkyl-en)nickel(II) complexes is dominated not only by the strength of the counter-anion binding to the
nickel(II) ion but also by the polarity and the DN of the solvents. As a result, nickel(II) complexes containing
the FSA or Tf2N anion show a time-dependent solvation in acetone.
Acknowledgments
This work was financially supported by a Grant-in-Aid for Scientific Research No. 25410191. We are
grateful to Mr. Kazufumi Chifuku at Division of Instrumental Analysis of Nagasaki University for the
measurements of CHN elemental analyses and to Mr. Takuya Funao at Analytical Center of JASCO
Engineering Co. Ltd. for the measurements of diffuse reflection spectra of the nickel(II) complex solids. We
are also grateful to Prof. Sumio Kaizaki of Osaka University for the advice in the discussion of VIS absorption
spectra of the nickel(II) complexes.
Notes
The authors declare no competing financial interest.
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Synopsis: The solid-to-liquid transition temperatures (TMP) for dipositive analogous bis(alkyl-en)metal(II)
(metal = nickel, copper, and zinc) and tris(alkyl-en)nickel(II) complexes fall in a wide range of –50 to 250 °C
depending on the molecular structures.
Graphical abstract
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Highlights:




Some divalent zinc(II) and nickel(II) complexes form room temperature ionic liquids
Analogous divalent complexes take a wide range of melting temperatures
Some rules were found concerning the formation of ionic liquids of divalent complexes
Alkylethylenediamine nickel(II) complexes show time-dependent acetone solvation
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