High conductivity and low viscosity Brønsted acidic ionic liquids with oligomeric anions

Journal of the Taiwan Institute of Chemical Engineers 42 (2011) 874–881

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High conductivity and low viscosity Brønsted acidic ionic liquids with oligomeric anions Tzi-Yi Wu a,b, I.-Wen Sun a, Shr-Tusen Gung a, Bor-Kuan Chen b, H. Paul Wang c, Shyh-Gang Su a,* a

Department of Chemistry, National Cheng Kung University, Tainan 70101, Taiwan Department of Polymer Materials, Kun Shan University, Tainan 71003, Taiwan c Department of Environmental Engineering, National Cheng Kung University, Tainan, Taiwan b

A R T I C L E I N F O

A B S T R A C T

Article history: Received 7 September 2010 Received in revised form 23 December 2010 Accepted 11 January 2011 Available online 22 March 2011

Two series of Brønsted acidic room temperature ionic liquids (RTILs) comprising of oligomeric anions and cyclic amine based cations are prepared. All RTILs studied here are liquids and have a high ionic conductivity (up to 20.96 mS/cm) and a low viscosity (down to 10.4 mPa s) at room temperature. The incorporation of oligomeric anions in ILs significantly increases the density and conductivity, and decreases the viscosity. The temperature-dependent electrolyte viscosity with molar conductivity is correlated through applying an empirical Walden’s rule. A relatively larger deviation of the plots from the ‘‘ideal’’ Walden line is observed for the ILs synthesized by equal molar acid and base, although the deviation decreases significantly when the ILs contain oligomeric anions. Spectroscopic measurements of solvatochromic probe molecules in RTILs provide insights into solvent intermolecular interactions through interpretation of the different and generally uncorrelated polarity scales. ß 2011 Taiwan Institute of Chemical Engineers. Published by Elsevier B.V. All rights reserved.

Keywords: Ionic liquids Viscosity Conductivity Walden rule Polarity

1. Introduction One of the biggest problems in the chemical industry is the fact that all chemical plants rely heavily on toxic, hazardous, and flammable organic solvents. Room temperature ionic liquids (RTILs) have certain properties, such as negligible vapor pressure, wide liquid range, and stability at high temperatures, make them good replacements for conventional volatile, flammable, and toxic organic solvents in the chemical industry (Earle and Seddon, 2000). Moreover, the option of fine-tuning the physicochemical properties by an appropriate choice of cations and anions has stimulated much of the current excitement with respect to these compounds and has led to them being termed ‘‘designer solvents’’ (Lazzu´s, 2009; Soriano et al., 2009, 2010a,b; Tien and Teng, 2009; Tsao et al., 2011; Wang et al., 2010; Wong et al., 2009; Wu et al., 2010a,b,c,d,e, 2011a,b; Yu et al., 2009). Ionic liquids can be thought of as falling into one of two distinct groups: aprotic and protic. Of these, protic ionic liquids (PILs) are easily produced through the combination of a Brønsted acid and Brønsted base. PILs have recently been studied in fuel cells, as electrolytes in cell compartments for H2 oxidation and O2 reduction (Nakamoto and Watanabe, 2007), and for impregnating proton diffusion membranes (Yamada and Honma, 2005). In both cases, these studies were motivated by the need to replace water as

* Corresponding author. Tel.: +886 6 2757575x65330. E-mail address: [email protected] (S.-G. Su).

the proton conductor to increase the operating temperature. Protic ILs are also emerging as useful materials due to various applications in double layer capacitors (Rochefort and Pont, 2006), catalysts (Qiao et al., 2009), and surfactants (Ma et al., 2010). These promise large new fields of application for protic ILs. In recent years, many studies have focused on the physicochemical properties of ionic liquids. The viscosities of most ionic liquids are one to three orders of magnitude greater than those of traditional organic liquids, which affects the rate of mass transfer, and the lower electrical conductivity of some pure ionic liquids due to their higher viscosity is another obstacle to progress in numerous fields, so it is necessary to develop ILs with high conductivity and low viscosity. Polarity can be viewed as the capacity of a solvent with regard to all the intermolecular interactions between the solvent and solute (Reichardt, 2005), such as hydrogen bonding, induction, and a probe molecule is capable of providing a quantitative scale of solvent polarity. Nile Red is a neutral, hydrophobic, and solvatochromic dye with interesting photophysical and lasing properties (Trivedi et al., 2009). Nile Red dissolved in protic ILs has its characteristic maximum absorption in the appropriate visible region, and thus this work will evaluate it for use as a solvatochromic dye and estimate the polarity of the protic ILs. Formic acid (methanoic acid, HCOO) is the simplest carbonic acid, formates are considered to have comparatively low toxicity, as LD50 (mouse) has been determined to be 11.2 g/kg (MSDS of sodium formate). In this paper, we report two series of ionic liquids based on piperidinium and morpholinium as cations, and

1876-1070/$ – see front matter ß 2011 Taiwan Institute of Chemical Engineers. Published by Elsevier B.V. All rights reserved. doi:10.1016/j.jtice.2011.01.008

T.-Y. Wu et al. / Journal of the Taiwan Institute of Chemical Engineers 42 (2011) 874–881

oligomeric formate as anions. Their thermal and physicochemical properties, such as viscosity, density, conductivity, and polarity, were evaluated and studied in detail. Furthermore, their conductivity and viscosity data are plotted under Walden’s rule, where ion association plays a major role. It is expected that the incorporation of oligomeric formate containing ILs into the ionic liquid family may facilitate the utility of ionic liquids and open up a new field of ionic liquid chemistry. 2. Experimental 2.1. Materials and measurements All starting materials were purchased from Aldrich, Alfa Aesar, TCI, and Acros and used as received. Solvents were freshly distilled prior to use. The conductivity (s) of the ionic liquid was systematically measured with a conductivity meter (LF 340) and standard conductivity cell (TetraCon 325, WissenschaftlichTechnische Werksta¨tten GmbH, Germany). The cell constant was determined by calibration after each sample measurement using an aqueous 0.01 M KCl solution. The densities of ionic liquids were measured gravimetrically with a 1 mL volumetric flask. Values for the densities are given as 0.01 g/mL. The viscosities (h) of the ILs were measured using a calibrated modified Ostwald viscometer (Cannon-Fenske glass capillary viscometers, CFRU, 9721-A50) with an inner diameter of 1.2  2% mm. The viscometer was placed in a thermostatic water bath (TV-4000, TAMSON) whose temperature was regulated to within 0.01 K. The flow time was measured using a stop watch with a time resolution of 0.01 s. For each IL, the experimental viscosity was obtained by averaging three to five flow time measurements. The melting point of each IL was analyzed using a differential scanning calorimeter (DSC, Perkin–Elmer Pyris 1) in the temperature range of 140 8C to a predetermined temperature. The sample was sealed in an aluminum pan, and then heated at a scan rate of 10 8C/min under a nitrogen flow. The thermal data were collected during heating in the second heating–cooling scan. The thermal stabilities were measured with a TGA (Perkin–Elmer, 7 series thermal analysis system). The sample was heated at 20 8C/min from room temperature to 800 8C under nitrogen. The water content of the dried ILs was measured with a moisture titrator (Metrohm 73KF coulometer) using the Karl–Fischer method, and the content was found to be less than 300 ppm. The NMR spectra of the synthetic ionic liquids were recorded on a BRUKER AV500 spectrometer in D2O and calibrated with tetramethylsilane (TMS) as the internal reference. The absorption spectra of Nile Red dissolved in the ionic liquids were obtained with a Jasco V-550 spectrometer. IL solutions were examined in a quartz cuvette with 1-cm light path length. For polarity measurements, a Nile Red concentration was chosen such that absorbances would fall in the range 0.5–1.0. The measurement of the acidic scale of these Brønsted acidic ILs was conducted on a Jasco V-550 spectrometer with a Methyl Red indicator.

875

removal procedure (Anouti et al., 2008; Bicak, 2005; Burrell et al., 2010). The synthetic reactions were carried out without any solvent. The structure of each protic IL was identified using nuclear magnetic resonance (NMR) spectroscopy. The ionic liquid samples were kept sealed in vials using thick layers of paraffin and stored in an argon atmosphere glovebox (VAC, O2 < 1 ppm, H2O < 1 ppm) before use. 2.2.1. [MePip][HCOO] Feed ratio of N-methylpiperidine:formic acid = 1:1. Yield: 77%. 1 H NMR (400 MHz, DMSO-d6, ppm): 11.9 (br, 1H, N–H), 8.30 (s, 1H, HCOO–), 2.98–2.95 (t, 4H, N–CH2–), 2.59 (s, 3H, N–CH3), 1.69–1.61 (m, 4H, N–CH2–CH2–CH2), 1.48–1.42 (m, 2H, N–CH2–CH2–CH2). 1H NMR (400 MHz, D2O, ppm): 8.18 (s, 1H, HCOO–), 3.27–3.23 (t, 2H, N–CH2–), 2.75–2.68 (t, 2H, N–CH2–), 2.61 (s, 3H, N–CH3), 1.70–1.62 (m, 2H, N–CH2–CH2–CH2), 1.31–1.17 (m, 4H, N–CH2–CH2–CH2 and N–CH2–CH2–CH2). 2.2.2. [MePip][2HCOO] Feed ratio of N-methylpiperidine:formic acid = 1:2. Yield: 73%. 1 H NMR (400 MHz, D2O, ppm): 8.22 (s, 1H, HCOO–), 3.34 (t, 2H, N– CH2–), 2.80 (t, 2H, N–CH2–), 2.70 (s, 3H, N–CH3), 1.84–1.79 (m, 2H, N–CH2–CH2–CH2), 1.65–1.34 (m, 4H, N–CH2–CH2–CH2 and N– CH2–CH2–CH2). 2.2.3. [MePip][4HCOO] Feed ratio of N-methylpiperidine:formic acid = 1:4. Yield: 70%. 1 H NMR (400 MHz, D2O, ppm): 8.19 (s, 1H, HCOO–), 3.32 (t, 2H, N– CH2–), 2.79 (t, 2H, N–CH2–), 2.68 (s, 3H, N–CH3), 1.82–1.77 (m, 2H, N–CH2–CH2–CH2), 1.63–1.32 (m, 4H, N–CH2–CH2–CH2 and N– CH2–CH2–CH2). 2.2.4. [MeMor][HCOO] Feed ratio of N-methylmorpholine:formic acid = 1:1. Yield: 71%. 1 H NMR (400 MHz, DMSO-d6, ppm): 8.98 (br, 1H, N–H), 8.19 (s, 1H, HCOO–), 3.65 (t, 4H, O–CH2–), 2.70 (t, 4H, N–CH2–), 2.42 (s, 3H, N– CH3). 1H NMR (400 MHz, D2O, ppm): 8.15 (s, 1H, HCOO–), 3.94 (t, 2H, O–CH2–), 3.63 (t, 2H, O–CH2–), 3.31 (t, 2H, N–CH2–), 3.02 (t, 2H, N–CH2–), 2.74 (s, 3H, N–CH3). 2.2.5. [MeMor][2HCOO] Feed ratio of N-methylmorpholine:formic acid = 1:2. Yield: 68%. 1 H NMR (400 MHz, D2O, ppm): 8.19 (s, 1H, HCOO–), 3.96 (t, 2H, O– CH2–), 3.65 (t, 2H, O–CH2–), 3.34 (t, 2H, N–CH2–), 3.04 (t, 2H, N– CH2–), 2.77 (s, 3H, N–CH3). 2.2.6. [MeMor][4HCOO] Feed ratio of N-methylmorpholine:formic acid = 1:4. Yield: 61%. 1 H NMR (400 MHz, D2O, ppm): 8.18 (s, 1H, HCOO–), 3.97 (t, 2H, O– CH2–), 3.65 (t, 2H, O–CH2–), 3.34 (t, 2H, N–CH2–), 3.05 (t, 2H, N– CH2–), 2.77 (s, 3H, N–CH3).

2.2. Synthesis of cyclic amine-based Brønsted acidic ionic liquids

3. Results and discussion

Amine compounds were placed in a three-necked glass flask equipped with a reflux condenser and a dropping funnel. The flask was mounted in an ice bath. The formic acid was added dropwise to the flask under stirring with a magnetic bar. The addition was done at room temperature (25  3 8C) while stirring. Stirring was continued for 24 h at ambient temperature in order to obtain a final viscous liquid, the produced ILs were washed repeatedly with diethyl ether to remove unreacted materials. The same general process was used for the synthesis of all ILs. The product was then dried at 80 8C for 12 h in a vacuum oven containing phosphorus pentoxide (P2O5) to remove any excess water. Relatively low yield for synthesized ionic liquids can be attributed to the product evaporation during water

3.1. Thermal properties of ILs Good thermal stability is important for these compounds to be used as phase-transfer catalysts and electrolytes. The thermal properties were measured for two compositions in the formic acid + N-methylpiperidine (or N-methylmorpholine) system, and the feed ratio for the mixtures was 1:1, 2:1, and 4:1. Thermogravimetric analysis experiments were conducted to determine the thermal stabilities of the proposed ILs. Fig. 1 shows the thermogravimetric traces of these ILs. The decomposition temperatures of 5% weight loss are listed in Table 1, the thermal decomposition temperature (Td) of the ILs with oligomeric anions

T.-Y. Wu et al. / Journal of the Taiwan Institute of Chemical Engineers 42 (2011) 874–881

876

100 [MePip][HCOO] [MePip][2HCOO] [MePip][4HCOO] [MeMor][HCOO] [MeMor][2HCOO] [MeMor][4HCOO]

80

Weight / %

range 98 to 102 8C, which is higher than that of piperidinium-based ILs (107 to 110 8C). Moreover, inspection of the transition temperatures as a function of the anion shows that the glass transition temperature of ILs with oligomeric anions is slightly lower than that of ILs synthesized by an equal molar acid and base.

60

3.2. Densities of ILs 40

20

0

0

100

200

300

400

500

600

Temperature / K

The basic physicochemical data of ILs, such as density, are important for a better understanding of the interactions in this kind of compound. Fig. 2 shows the temperature dependency of density for the RTILs, and the density has been found to decrease linearly with increasing temperature. The temperature-dependent densities (r) can be well fitted by the following equation:

r ¼ A þ BT

(1)

Fig. 1. Thermogravimetric trace for these ILs.

([MePip][4HCOO], Td = 86 8C) is lower than that of ILs synthesized by an equal molar acid and base ([MePip][HCOO], Td = 98 8C). The thermal properties of the salts were also investigated using differential scanning calorimetry. The phase transitions were observed on a second heating. All ILs were liquid at room temperature and most of them only exhibited glass transition point (Tg), the melting point (Tm) and Tg are summarized in Table 1. An inspection of the transition temperatures as a function of the cation shows that morpholinium-based ILs varied in the

where fitting parameters B and A are related to the coefficient of volume expansion (g/cm3/K) and extrapolated density at 0 K (g/cm3), respectively, and T is the temperature (K). The adjustable parameters of Eq. (1) for the density of these ILs are summarized in Table 2. The densities increase in the order [MePip][HCOO] < [MePip][2HCOO]<[MePip][4HCOO]<[MeMor][HCOO]<[MeMor] [2HCOO] < [MeMor][4HCOO] at all measured temperatures. The obtained densities depend not only on the cationic structure, but also on the oligomeric anions. In morphiline-based RTIL, the electron rich oxygen part likely interacts with other organic cations. For dimeric or oligomeric anions in morpholinium-based ILs, [MeMor][4HCOO]

Table 1 Physicochemical properties of the studied ionic liquids. Tda(8C)

rb(g/cm3)

h(mPa s)

s(mS/cm)

L(S cm2/mol)

98

1.0711

13.5

16.28

2.389

31

86

1.0772

12.8

16.64

3.038

110

29

86

1.0889

10.4

20.96

5.492

[MeMor][HCOO]

98

15

91

1.1517

20.6

12.98

1.738

: HCOOH (1:2)

[MeMor][2HCOO]

99

87

1.1650

19.06

13.34

2.222

: HCOOH (1:4)

[MeMor][4HCOO]

102

82

1.1700

16.9

16.34

3.814

Base:acid (base:acid ratio)

Name

Tg(8C)

: HCOOH (1:1)

[MePip][HCOO]

107

: HCOOH (1:2)

[MePip][2HCOO]

110

: HCOOH (1:4)

[MePip][4HCOO]

: HCOOH (1:1)

Tm(8C)

N

N

N

N O N O N O a b

Decomposition temperature of 5% weight loss. Density (r), viscosity (h), conductivity (s) and molar conductivity (L) are measured at 30 8C.

T.-Y. Wu et al. / Journal of the Taiwan Institute of Chemical Engineers 42 (2011) 874–881

-3

1.20

24 [MePip][HCOO] [MePip][2HCOO] [MePip][4HCOO] [MeMor][HCOO] [MeMor][2HCOO] [MeMor][4HCOO]

22 20 18

1.15

η / mPa.s

Density / g cm

(a)

[MePip][HCOO] [MePip][2HCOO] [MePip][4HCOO] [MeMor][HCOO] [MeMor][2HCOO] [MeMor][4HCOO]

877

1.10

16 14 12 10 8

1.05

6

300

310

320

330

340

350

4

360

2 300

Temperature / K Fig. 2. Temperature dependence of density data for ILs.

(b)

310

3.2

The viscosity is an important property of ionic liquids, because it strongly influences the diffusion of species, which are dissolved or dispersed in the ionic liquid. The viscosity of these ILs are studied from 303.0 to 353.8 K, and reported herein in Fig. 3. The viscosity of these neat ILs decreases with the temperature rising from 303 to 353 K, and the viscosity values, h, were fitted using the Arrehenius-like law and the Vogel–Tamman–Fulcher (VTF) equations. The most commonly used equation to correlate the variation of viscosity with temperature is the Arrhenius-like law:   Ea h ¼ h1 exp RT

T



 B ðT  T o Þ

(3)

where T is the absolute temperature and ho, B, and To are adjustable parameters. The best-fit ho (cP), B (K), and To (K) parameters are given in Table 3. The six ILs were very well fit by the VTF model over the temperature range studied. The ho, B, and To values did not show a distinct relationship with the viscosity value.

Ionic liquids

A

104B

R2a

[MePip][HCOO] [MePip][2HCOO] [MePip][4HCOO] [MeMor][HCOO] [MeMor][2HCOO] [MeMor][4HCOO]

1.269 1.262 1.286 1.384 1.380 1.386

6.529 6.088 6.496 7.671 7.107 7.132

0.999 0.999 0.999 0.999 0.999 0.999

Correlation coefficient.

360

2.6 2.4 2.2 2.0 1.8 1.6 1.4 1.2 1.0 2.8

2.9

3.0

3.1

3.2

3.3

1000 T-1 / K-1 Fig. 3. Dynamic viscosity (h) as a function of temperature for ILs. (a) h and T plot, (b) ln h and 1000/T plot.

When the feed ratio of N-methylpiperidine:formic acid increases from 1:1 to 1:4, the viscosity of ILs decrease significantly (from 13.5 to 10.4 mPa s at 30 8C), and this can be attributed to lower viscosity of formic acid (1.36 mPa s at 30 8C) (Cases et al., 2001) than [MePip][HCOO] (13.5 mPa s at 30 8C). Under similar conditions, [MeMor][4HCOO] exhibits lower viscosity than [MeMor][HCOO]. Inspection of the viscosity as a function of the cation shows that [MeMor][HCOO] has larger viscosity than [MePip][HCOO]. The transport property plays a key role in the viscosity data, and it seems to be determined by several factors, such as the size and shape of the ionic structures and intermolecular interactions. The interactions include the van der Waals and Coulombic interactions (Tokuda et al., 2005). In morphiline-based RTILs, the electron rich oxygen part likely interacts with other organic cations, which causes [MeMor][HCOO] to have greater viscosity than [MePip][HCOO]. These Brønsted acidic RTILs show Table 3 h i hoffiffi B VTF equation parameters of viscosity for ILs (h1 ¼ p exp ðTT ). oÞ T

Table 2 The adjustable parameters of density for ILs.

a

350

(2)

Viscosity at infinite temperature, h1, and the activation energy, Ea, are characteristic parameters generally adjusted from experimental data. According to Seddon et al. (2004), the Arrhenius law can generally be applied when the cation presents only a limited symmetry. If this is not the case, and especially in the presence of symmetrical cations, then the VTF equation is recommended (Wilkes, 2004).

h h1 ¼ poffiffiffi exp

340

[MePip][HCOO] [MePip][2HCOO] [MePip][4HCOO] [MeMor][HCOO] [MeMor][2HCOO] [MeMor][4HCOO]

2.8

ln (η/mPa.s)

3.3. Viscosities of ILs

330

T/K

3.0

has a higher density (r = 1.1700 g/cm3 at 30 8C) than [MeMor] [HCOO] (r = 1.1517 g/cm3 at 30 8C), implying more oligomeric anions in ILs increases the intermolecular packing and therefore increases the density.

320

Ionic liquids

ho (mPa s)

To (K)

B (K)a

R2b

[MePip][HCOO] [MePip][2HCOO] [MePip][4HCOO] [MeMor][HCOO] [MeMor][2HCOO] [MeMor][4HCOO]

0.1056 0.1607 0.1146 0.1053 0.2101 0.2001

158.8 165.6 147.6 162.2 190.5 183.7

704.4 605.8 704.7 747.8 517.0 535.2

0.999 0.999 0.999 0.999 0.999 0.999

a b

Activation energy (kJ/mol). Correlation coefficient.

T.-Y. Wu et al. / Journal of the Taiwan Institute of Chemical Engineers 42 (2011) 874–881

878

(a)

Table 4 The Ea, DS and DH evaluated by Eyring equation and the relationships of viscosity vs. temperature. Ea(kJ/mol)

DS(J/mol/K)

DH(kJ/mol)

[MePip][HCOO] [MePip][2HCOO] [MePip][4HCOO] [MeMor][HCOO] [MeMor][2HCOO] [MeMor][4HCOO]

22.22 20.808 19.53 24.655 24.928 23.267

305.69 301.49 298.97 310.23 311.5 307.31

24.943 23.531 22.252 27.376 27.649 25.993

50

σ / mS cm-1

Ionic liquids

60 [MePip][HCOO] [MePip][2HCOO] [MePip][4HCOO] [MeMor][HCOO] [MeMor][2HCOO] [MeMor][4HCOO]

40

30

20

10 300

310

320

330

340

350

360

T/K

(b)

4.2 [MePip][HCOO] [MePip][2HCOO] [MePip][4HCOO] [MeMor][HCOO] [MeMor][2HCOO] [MeMor][4HCOO]

4.0 3.8

ln (σ /mS cm-1)

lower viscosity (10.4–20.6 mPa s at 30 8C) than other reported ILs ([BMPy][NTf2], h = 73.9 mPa s; [BMIm][BF4], h = 77.1 mPa s at 25 8C (Stepniak and Andrzejewska, 2009)), and such low viscosity ILs can be good candidates for fuel cell devices (Nakamoto and Watanabe, 2007) and thermal transfer fluids. For the relationship of h vs. T, the Ea, DS, and DH values evaluated using the slope (Ea/R) of the Arrhenius equation and Eyring equation in these ILs are summarized in Table 4, and the absolute values of Ea, DS, and DH for these ILs show the order: morpholinium-based ILs ([MeMor][4HCOO]: Ea = 23.267 kJ/mol, jDSj = 307.31 J/mol/K, and DH = 25.993 kJ/mol) > piperidiniumbased ILs ([MePip][4HCOO]: Ea = 19.53 kJ/mol, jDSj = 298.97 J/ mol/K, and DH = 22.252 kJ/mol), there is no notable relationship of the Ea, DS, and DH values between ILs with oligomeric anions and ILs synthesized by equal molar acid and base.

3.6 3.4 3.2 3.0 2.8

3.4. Conductivites of ILs

2.6

According to the equation suggested by Every et al. (2000), solution conductivity can be expressed by Eq. (4): X s¼ ni qi ui (4) where ni is the number of charge carriers of type i, qi is the charge and ui is the mobility of each species, which is related to the viscosities of the mixtures. It can be seen that the increase in the conductivity of a given system must be due to the increase in ion mobility and/or the number of charge carriers. A comparison of the temperature dependence of the ionic conductivity of these ILs is made in Fig. 4. Some temperature dependences of the ionic conductivity have concave-curved profiles. To explore this behavior, the natural logarithm of s vs. the inverse of absolute temperature, i.e. ln s vs. 1/T, is plotted in Fig. 4(b). The observed temperature dependence of conductivity is often best described by the empirical Vogel–Tammann–Fulcher (VTF) equation (Wilkes, 2004): 

s s ¼ poffiffiffi exp T

B0 ðT  T o Þ



(5)

where so, B0 , and T0 are the fitting parameters, a factor related to the activation energy, and the ideal glass transition temperature, respectively. The VTF fitting parameters of the ionic conductivity for the ILs are summarized in Table 5. The temperature dependence of conductivity for the six ILs was very well fitted by the VTF model over the temperature range studied. Table 1 shows that the conductivity values of piperidiniumbased ILs are higher than those of the corresponding morpholinium salts at 30 8C. For instance, [MePip][HCOO] has a higher conductivity (16.28 mS/cm) than that of [MeMor][HCOO] (12.98 mS/cm). The effect of the anion species on the ionic conductivity was also revealed. A comparison of the conductivity shows the order: piperidinium-based ILs ([MePip][4HCOO], s = 20.96 mS/cm > [MePip][2HCOO], s = 16.64 mS/cm >

2.4 2.8

2.9

3.0

3.1 -1

3.2

3.3

-1

1000 T / K

Fig. 4. Specific conductivity (s) as a function of temperature for the ILs. (a) s and T plot, (b) ln s and 1000/T plot.

[MePip][HCOO], s = 16.28 mS/cm); morpholinium-based ILs ([MeMor][4HCOO], s = 16.34 mS/cm > [MeMor][2HCOO], s = 13.34 mS/cm > [MeMor][HCOO], s = 12.98 mS/cm). All of these six ILs have a high conductivity, and among the salts, [MePip][4HCOO] shows the highest ionic conductivity of 20.96 mS/cm at 30 8C, as shown in Fig. 4, which is larger than that of other reported ILs ([BMIM][BF4], s = 5.77 mS/cm (Liu et al., 2006), [C4mim][PF6], s = 2.40 mS/cm (Geng et al., 2008); PYR1n4TFSI, s = 1.8 mS/cm (Appetecchi et al., 2009); [amim][CI], s = 4.4 mS/cm (Xu et al., 2005)). ILs with high ionic conductivity are potential candidates for use as electrolytes in batteries or supercapacitors (Shah et al., 2002; Sun et al., 1998). For the relationship of s vs. T, the Ea, DS, and DH values evaluated using the slope (Ea/R) of the Arrhenius and Eyring equations in these ILs are summarized in Table 6, and the Ea, DS, and DH values of the ILs show the order: morpholinium-based ILs ([MeMor][4HCOO]: Ea = 19.002 kJ/mol, DS = 168.29 J/mol/K, and Table 5 h i s offiffi exp B0 ). VTF equation parameters of conductivity for ILs. (s ¼ p ðTT o Þ T Ionic liquids

so (mS/cm)

To (K)

B0 (K)

R2a

[MePip][HCOO] [MePip][2HCOO] [MePip][4HCOO] [MeMor][HCOO] [MeMor][2HCOO] [MeMor][4HCOO]

303.03 303.03 434.78 714.29 573.39 806.45

198.2 197.8 185.4 171.6 185.0 167.3

306.4 305.3 356.6 526.6 446.8 535.9

0.999 0.999 0.999 0.999 0.999 0.999

a

Correlation coefficient.

T.-Y. Wu et al. / Journal of the Taiwan Institute of Chemical Engineers 42 (2011) 874–881 Table 6 The Ea, DS and DH evaluated by Eyring equation and the relationships of conductivity vs. temperature. Ionic liquids

Ea/kJ/mol

DS/J/mol/K

DH/kJ/mol

[MePip][HCOO] [MePip][2HCOO] [MePip][4HCOO] [MeMor][HCOO] [MeMor][2HCOO] [MeMor][4HCOO]

16.425 16.487 16.246 19.652 19.928 19.002

176.13 175.91 174.91 167.34 166.6 168.29

13.707 13.768 13.549 16.934 17.211 16.285

DH = 16.285 kJ/mol) > piperidinium-based ILs ([MePip][4HCOO]: Ea = 16.246 kJ/mol, DS = 174.91 J/mol/K, and DH = 13.549 kJ/ mol), there is no remarkable relationship of the Ea, DS, and DH values between ILs with oligomeric anions and ILs synthesized by an equal molar acid and base. The Walden plot was used to evaluate the ionicity in ionic liquids. Recent work by Xu et al. (2003) has shown that the empirical Walden’s rule applies rather well to pure ionic liquids. The Walden plot (ln(molar conductivity) vs. ln(1/h)) shows the relationship between conductivity and viscosity. The molar conductivity L of the electrolyte is defined as L = Ves, where Ve is the molar volume of the salt. The molecular volume of the ILs was calculated from the experimental density using the following equation: Ve ¼

M

(6)

r

where M is the molar mass and r is the density, which in this case is a function of temperature. When the conductivity is strongly correlated with the viscosity, the relationship can be expressed as:

Lha ¼ C

(7)

where C is a constant and a is the slope of the line in the Walden plot, which reflects the decoupling of the ions. The Walden plots in Fig. 5 show the variations of ln(L) vs. ln(h1) for the ILs. The fitted a values of the ionic liquids are [MePip][HCOO] (a = 0.763), [MePip][2HCOO] (a = 0.816), [MePip][4HCOO] (a = 0.846), [MeMor][HCOO] (a = 0.817), [MeMor][2HCOO] (a = 0.819), and [MeMor][4HCOO] (a = 0.838). All of these values are smaller than unity (a < 1), and this indicates that the ionic conductivities of these liquid salts is somewhat diminished as a result of ion-pair

log(Λ/S cm2 mol-1)

formation. The Walden plot is calibrated using the data for a 0.01 M KCl aqueous solution, where the ions are known to be fully dissociated and to have equal mobility. As shown in Fig. 5, deviation from the KCl line depends on the ionic structure of the ILs. For the anionic structures with a fixed [MePip] cation, the deviations follow the order: [MePip][HCOO] > [MePip][2HCOO] > [MePip][4HCOO]. A relatively larger deviation of the plots from the ‘‘ideal’’ Walden line was observed for the ILs synthesized by an equal molar acid and base ([MePip][HCOO] and [MeMor][HCOO]). This suggests that ion-ion correlations in these two ILs are stronger, and more ion-pairs or aggregates occurred in them. 3.5. Determination of Ho values of Brønsted acidic ILs The Hammett acidity function can effectively express the acidity strength of an acid in organic solvents, according to previous work (Thomazeau et al., 2003). It can be calculated by the equation below: Ho ¼ pkðIÞaq þ logð½Is =½IHþ s Þ

(8)

Here, ‘‘I’’ represents the indicator base, [IH+]s and [I]s are the molar concentrations of the protonated and unprotonated forms of the indicator, respectively. According to Lambert–Beer’s Law, the value of [I]s/[IH+]s can be determined and calculated through UV–visible spectrum. In our experiment, Methyl Red was chosen as the basic indicators and CH2Cl2 as the solvent, the acidity order of the six ILs with the following Ho values (Table 7): (a) [MePip][4HCOO] (5.2411) > [MePip][2HCOO] (5.3041) > [MePip][HCOO] (5.3554); (b) [MeMor][4HCOO] (5.3420) > [MeMor][2HCOO] (5.4991) > [MeMor][HCOO] (5.7808); (c) [MePip][4HCOO] (5.2411) > [MeMor][4HCOO] (5.3420). The acidity of the ILs depended both on the characteristics of the cations and anions. The ILs of piperidinium-based cations have stronger acidity than those of morpholinium-based cations. When the cations of the ILs were the same, the dependence of the acidity of the IL on anion was significant, the Brønsted acidity of ILs with oligomeric anions were relatively stronger than those of ILs synthesized by an equal molar acid and base. 3.6. Polarity study of ILs with the solvatochromic dye Nile Red In this study, Nile Red was used instead of Reichardt’s dye 30 (Deye et al., 1990), because it is more stable under acidic conditions and has been used previously for studying the polarity of ILs (Carmichael and Seddon, 2000; Ogihara et al., 2004). Nile Red is a positive solvatochromic dye and shows one of the largest shifts in excitation maxima in going from nonpolar solvents (in hexane, lmax  484.6 nm) to polar solvents (in water, lmax  593.2 nm) (see Table 8). These ENR values are simply defined as the molar transition energies (in kcal/mol; 1 kcal = 4.184 kJ) of the Nile Red dye,

4 [MePip][HCOO] [MePip][2HCOO] [MePip][4HCOO] [MeMor][HCOO] [MeMor][2HCOO] [MeMor][4HCOO]

2

879

0

Table 7 Calculation and comparison of Ho values of different ILs in CH2Cl2 (30 8C).

-2

-4 -4

-2

0 -1

2

4

-1

log(η /P ) Fig. 5. Walden plots for ILs, where L is the equivalent conductivity and h1 is the fluidity. The solid line plot is generated from data obtained in aqueous 0.01 M KCl solution.

[MePip][HCOO] [MePip][2HCOO] [MePip][4HCOO] [MeMor][HCOO] [MeMor][2HCOO] [MeMor][4HCOO]

Abs1

Abs2

[I]

[IH+]

Ho

1.8221 1.866 0.9057 1.7943 3.5106 1.3072

1.2643 1.2469 0.5754 1.5393 2.6658 0.8984

0.6939 0.6682 0.6353 0.8579 0.7594 0.6873

0.3061 0.3318 0.3647 0.1421 0.2406 0.3127

5.3554 5.3041 5.2411 5.7808 5.4991 5.3420

Indicator: Methyl Red (pk(I)aq = 5). Abs1: concentration is 0 M; Abs2: concentration is 1 M.

T.-Y. Wu et al. / Journal of the Taiwan Institute of Chemical Engineers 42 (2011) 874–881

880

Table 8 Wavelengths of maximum absorption (lmax) and molar transition energies (ENR) for Nile Red dissolved in ionic liquids.

lmax

Ionic liquid

N Me

BF 4

N

ENR (kcal/mol)a

Reference

(nm) 562.3

50.9

Ogihara et al., 2004

ENR = 52.1 kcal/mol), which can be attributed to the higher polarity of formic acid (ENR = 45.09 kcal/mol) (Deye et al., 1990). The polarities of these ILs are lower than those of some imidazole and tetrafluoroborate based protic ILs (ENR = 50.8–50.9 kcal/mol) (Ogihara et al., 2004). As these ILs become alternatives for conventional solvents, this polarity data will become important for the chemical industry.

H

4. Conclusions N Et

H

N Bu

BF 4

N

BF4

N H

[MePip][HCOO] [MePip][2HCOO] [MePip][4HCOO] [MeMor][HCOO] [MeMor][2HCOO] [MeMor][4HCOO] Water Methanol Ethanol DMF CHCl3 CH3CN Hexane

562.9

50.8

Ogihara et al., 2004

562.8

50.8

Ogihara et al., 2004

549 550 552 552 553 554 593.2 549.8 548.8 541.5 537.4 531.4 484.6

52.1 52.0 51.8 51.8 51.7 51.6 48.2 52.0 52.1 52.8 53.2 53.8 59.0

This work This work This work This work This work This work Carmichael Carmichael Carmichael Carmichael Carmichael Carmichael Carmichael

and and and and and and and

Seddon, Seddon, Seddon, Seddon, Seddon, Seddon, Seddon,

2000 2000 2000 2000 2000 2000 2000

a ENR = (hcNA/lmax)  106, where h is Planck’s constant c is the speed of light, NA is Avogadro’s number and lmax is the wavelength of maximum absorption (nm).

measured in solvents of different polarity at room temperature (30 8C) and normal pressure (1 bar), according to Eq. (8): ENR ðkcal=molÞ ¼

hcNA

lmax

 106 ¼

28; 591

lmax

(9)

where h is the Planck’s constant, c is the speed of light, NA is the Avogadro’s number and lmax is the wavelength of maximum absorption (nm). The wavelengths of maximum absorption (lmax) and the molar transition energies (ENR) for Nile Red dissolved in these ILs are summarized in Table 8, along with the values of the absorption behavior of Nile Red dissolved in a variety of methylimidazolium-based RTILs (Carmichael and Seddon, 2000; Ogihara et al., 2004) for comparison. The ENR values calculated for these ILs are in the range of 51.6–52.1 kcal/mol, which shows higher polarity than those of the apolar solvent hexane (ENR = 59.0 kcal/mol), and polar solvents CHCl3 (ENR = 53.2 kcal/mol) and CH3CN (ENR = 53.8 kcal/mol). The polarity of these ILs is comparable with that of general organic solvents, such as methanol (ENR = 52.0 kcal/mol), ethanol (ENR = 52.1 kcal/mol), and DMF (ENR = 52.8 kcal/mol), showing that they are suitable for technical development as polar solvents. An inspection of the effect of cation species on polarity at 30 8C shows that morpholinium-based ILs ([MeMor][4HCOO], ENR = 51.6 kcal/mol) have a lower ENR value (higher polarity) than those of piperidinium-based ILs ([MePip][4HCOO], ENR = 51.8 kcal/mol). This can be attributed to the fact that morpholinium-based ILs have a higher density than those of piperidinium-based ILs, which increases the interactions between ions. An inspection of the effect of anion species on polarity at 30 8C shows that ILs containing oligomeric anions ([MePip][4HCOO]; ENR = 51.8 kcal/mol) have a larger polarity than those of ILs synthesized by an equal molar acid and base ([MePip][HCOO];

We have synthesized two series of Brønsted acidic ionic liquids containing oligomeric anions. Physico-chemical properties, such as glass transition temperature, decomposition temperature, viscosity, and density, conductivity, and polarity have been measured and discussed in terms of the component ion structure. RTILs contains oligomeric anions that have lower viscosity, lower density, and higher conductivity compared to those ILs prepared by an equal molar acid and base. These RTILs have a relatively low cost, low toxicity, lower viscosity, lower density, and higher conductivity as compared to other aprotic RTILs. Absorbance solvatochromic probe Nile Red is used to investigate the relative polarity of the ionic liquids, and the results are compared with those of imidazole-based RTILs and traditional organic solvents, demonstrating that these ILs are suitable for technical development as polar solvents. Acknowledgements The authors would like to thank the National Science Council of the Republic of China for financially supporting this project. The authors also gratefully acknowledge the contributions of Chao-anx Lai, Department of chemistry, National Cheng Kung University, for helping with the laboratory work. References Anouti, M., M. Caillon-Caravanier, C. Le Floch, and D. Lemordant, ‘‘AlkylammoniumBased Protic Ionic Liquids. II. Ionic Transport and Heat-Transfer Properties: Fragility and Ionicity Rule,’’ J. Phys. Chem. B, 112, 9412 (2008). Appetecchi, G. B., M. Montanino, D. Zane, M. Carewska, F. Alessandrini, and S. Passerini, ‘‘Effect of the Alkyl Group on the Synthesis and the Electrochemical Properties of NAlkyl-N-Methyl-Pyrrolidinium Bis(trifluoromethanesulfonyl)imide Ionic Liquids,’’ Electrochim. Acta, 54, 1325 (2009). Bicak, N., ‘‘A New Ionic Liquid: 2-Hydroxy Ethylammonium Formate,’’ J. Mol. Liq., 116, 15 (2005). Burrell, G. L., I. M. Burgar, F. Separovic, and N. F. Dunlop, ‘‘Preparation of Protic Ionic Liquids with Minimal Water Content and 15N NMR Study of Proton Transfer,’’ Phys. Chem. Chem. Phys., 12, 1571 (2010). Carmichael, A. J. and K. R. Seddon, ‘‘Polarity Study of Some 1-Alkyl-3-Methylimidazolium Ambient-Temperature Ionic Liquids with the Solvatochromic Dye, Nile Red,’’ J. Phys. Org. Chem., 13, 591 (2000). Cases, A. M., A. C. G. Marigliano, C. M. Bonatti, and H. N. So´limo, ‘‘Density, Viscosity, and Refractive Index of Formamide, Three Carboxylic Acids, and Formamide + Carboxylic Acid Binary Mixtures,’’ J. Chem. Eng. Data, 46, 712 (2001). Deye, J. F., T. A. Berger, and A. G. Anderson, ‘‘Nile Red as a Solvatochromic Dye for Measuring Solvent Strength in Normal Liquids and Mixtures of Normal Liquids with Supercritical and Near Critical Fluids,’’ Anal. Chem., 62, 615 (1990). Earle, M. J. and K. R. Seddon, ‘‘Ionic Liquids Green Solvents for the Future,’’ Pure Appl. Chem., 72, 1391 (2000). Every, H., A. G. Bishop, M. Forsyth, and D. R. MacFarlane, ‘‘Ion Diffusion in Molten Salt Mixtures,’’ Electrochim. Acta, 45, 1279 (2000). Geng, Y. F., S. L. Chen, T. F. Wang, D. H. Yu, C. J. Peng, H. L. Liu, and Y. Hu, ‘‘Density, Viscosity and Electrical Conductivity of 1-Butyl-3-Methylimidazolium Hexafluorophosphate + Monoethanolamine and +N, N-Dimethylethanolamine,’’ J. Mol. Liq., 143, 100 (2008). Lazzu´s, J. A., ‘‘r–T–P Prediction for Ionic Liquids Using Neural Networks,’’ J. Taiwan Inst. Chem. Engrs., 40, 213 (2009). Liu, W. W., T. Y. Zhao, Y. M. Zhang, H. P. Wang, and M. F. Yu, ‘‘The Physical Properties of Aqueous Solutions of the Ionic Liquid [BMIM][BF4],’’ J. Solution Chem., 35, 1337 (2006). Ma, F. M., X. Chen, X. D. Wang, Y. R. Zhao, Q. H. Li, X. Yue, and C. Lv, ‘‘Synthesis and Characterization on a Novel Series of Protic Pyrrolidinium Surfactants,’’ Chin. Chem. Lett., 21, 385 (2010). MSDS Sodium Formate, CAS #141-53-7, available at http://www.sciencelab.com/ xMSDS-Sodium_formate-9927596.

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