Ionic liquids containing an alkyl sulfate group as potential electrolytes

Ionic liquids containing an alkyl sulfate group as potential electrolytes

Electrochimica Acta 55 (2010) 4475–4482 Contents lists available at ScienceDirect Electrochimica Acta journal homepage: www.elsevier.com/locate/elec...

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Electrochimica Acta 55 (2010) 4475–4482

Contents lists available at ScienceDirect

Electrochimica Acta journal homepage: www.elsevier.com/locate/electacta

Ionic liquids containing an alkyl sulfate group as potential electrolytes Tzi-Yi Wu a,b , Shyh-Gang Su a , Shr-Tusen Gung a , Ming-Wei Lin a , Yuan-Chung Lin c , Chao-Anx Lai a , I-Wen Sun a,d,∗ a

Department of Chemistry, National Cheng Kung University, Tainan, 701, Taiwan Department of Polymer Materials, Kun Shan University, Tainan 71003, Taiwan c Institute of Environmental Engineering, National Sun Yat-Sen University, Kaohsiung 804, Taiwan d Sustainable Environment Research Center, National Cheng Kung University, Tainan 701, Taiwan b

a r t i c l e

i n f o

Article history: Received 28 October 2009 Received in revised form 22 February 2010 Accepted 27 February 2010 Available online 6 March 2010 Keywords: Ionic liquid Molten salt Potential window Conductivity Sulfate Polarity

a b s t r a c t Various types of ionic liquid (IL) containing an alkyl sulfate group are synthesized and their physical and electrochemical properties are investigated. The temperature dependency of dynamic viscosity and ionic conductivity are measured for these ILs. The low temperature phase behaviour of the ethylsulfate salts is investigated using differential scanning calorimetry. Three ethylsulfate-containing ionic liquids exhibit wide electrochemical windows of about 5.0 V, and one pyrrolidinium-containing ionic liquid shows a conductivity of 3.8 mS cm−1 . The various cations of alkylsulfate-containing ionic liquids are shown to greatly influence viscosity, density, and conductivity. Absorbance solvatochromic probe, Nile Red is used to investigate the relative polarity of alkylsulfate base ionic liquids compared with several organic solvents. The electrochemical and thermal stabilities of these ILs make them promising electrolytes for use in electrochemical devices. © 2010 Elsevier Ltd. All rights reserved.

1. Introduction During the past decade, numerous studies have discussed the properties and applications of room-temperature ionic liquids (RTILs) [1]. Ionic liquids can serve as solvents in organic chemistry [2] and electrochemistry [3], and as catalysts [4]. They can also be used to extract organic or ionic solutes [5,6]. The first IL extensively studied was 1-ethyl-3-methylimidazolium chloroaluminate (EMImAlCl4 ) [7]. However, chloroaluminate ionic liquids are easily hydrolyzed by traces of water, which limits their use. More recently, ILs with nitrogen, sulfur, or phosphorus as the central atom of cations have been extensively investigated. These include imidazolium, pyrrolidinium, tetraalkylammonium, pyridinium, piperidinium, sulfonium and phosphonium-based ILs [8,9]. Moreover, ILs consisting of large organic, asymmetrical ions, such as 1-ethyl-3-methylimidazolium, 1,3-dialkylimidazolium, 1-alkyl2,3-dimethylimidazolium, 1-alkylpyridinium, 1-alkylpyrazolium, tetraalkylammonium, or tetraalkylphosphonium cations and BF4 − , PF6 − , CF3 SO3 − , or N(SO2 CF3 )2 − anions have been developed [10,11]. The easy modification of the cations and anions in RTILs allows the development of task-specific RTILs for catalysis,

∗ Corresponding author at: Department of Chemistry, National Cheng Kung University, Tainan 70101, Taiwan. Tel.: +886 6 2757575x65355. E-mail address: [email protected] (I.-W. Sun). 0013-4686/$ – see front matter © 2010 Elsevier Ltd. All rights reserved. doi:10.1016/j.electacta.2010.02.089

organic synthesis, nanoparticles, extraction and dissolution [12,13]. However, structural modifications affect RTIL physical–chemical properties. As a consequence, it is of great importance to understand the relationship between structural changes and the properties of ILs. This paper reports the synthesis of seven ILs, which contain one of three cations (pyrrolidinium, piperidinium, and morpholinium) and one of two alkyl length anions (methylsulfate and ethylsulfate). The thermal and physicochemical properties of cationic and anionic structures were studied in detail. Polarity is an important property of ILs because it affects the solubility of a solute, the reaction efficiency of a solvent, and miscibility with other solvents [14]. To determine the polarity of ILs, several suitable probes, mainly solvatochromic dyes, such as Reichardt’s dye and Nile Red, have been reported [15,16]. Nile Red displays positive solvatochromism, which leads to a large red shift in the absorption and emission maxima when going from nonpolar solvents to polar solvents [17]. Nile Red is a particularly effective solvatochromic dye, consisting of a rigid aromatic group and an exocyclic diethylamine group [17]. Its absorption and fluorescence depends on the physical properties of the surrounding solvent environment, such as polarity. To understand more about the nature of ambient-temperature ionic liquids, and how they compare with conventional solvents or other RTILs, we studied their polarity using the solvatochromic dye Nile Red.

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Compare the alkyl sulfate with other anion for representative ILs reported, such as halide (X), metallic halide complex ([MmXn]), [PF6 ]− , [BF4 ]− , [NTf2 ]− , [OTf]− , and dialkylphosphate [18]. Halides are highly corrosive to steel [19], [MmXn] is sensitive to moisture [20], [PF6 ] is subject to hydrolysis [21], and [NTf2 ] and [OTf] type ILs are expensive. In contrast, ILs with alkyl sulfate anions are suitable for practical applications because they can be produced in a one-pot reactor under mild conditions with very high yield. More importantly, they are biodegradable and less toxic than other ILs [19]. To determine the suitability of ionic liquids for electrochemistry applications, the thermal properties, electrochemical windows, viscosity, conductivity, and polarity of the ILs synthesized in this work were also studied.

2. Experimental 2.1. Chemicals All starting materials were purchased from Aldrich, Lancaster, TCI, and Acros and used as received. Solvents were freshly distilled prior to use according to the literature procedure.

2.2. Measurement The conductivity () of the ionic liquid was systematically measured with a conductivity meter (LF 340) and a standard conductivity cell (TetraCon 325, Wissenschaftlich-Technische Werkstätten 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−1 . The viscosities () of the ILs were measured using a calibrated modified Ostwald viscometer (Cannon-Fenske glass capillary viscometers, CFRU, 9721-A50) with inner diameters 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-lmer Pyris 1) in the temperature range of −140 ◦ C to a predetermined temperature. The sample was sealed in an aluminum pan, and then heated at a scan rate of 10 ◦ C min−1 under a flow of nitrogen. The thermal data were collected during heating in the second heating–cooling scan. The thermal stabilities were measured with TGA (Perkin-lmer, 7 series thermal analysis system). The sample was heated at 20 ◦ C min−1 from room temperature to 800 ◦ C under nitrogen. The water content of the dried ILs was measured with a moisture titrator (Metrohm 73KF coulometer) using the Karl–Fischer method; the content was less than 200 ppm. The NMR spectra of the synthetic ionic liquids were recorded on a BRUKER AV500 spectrometer in D2 O and calibrated with tetramethylsilane (TMS) as the internal reference. Cyclic voltammetry was performed at 25 ◦ C using an electrochemical workstation (CH instruments Inc., CHI, model 750A). The working electrode was a glassy carbon electrode, the counter electrode was a Pt wire, and a Pt quasi-reference electrode. All electrochemical experiments were performed under a dry argon atmosphere to remove oxygen and air humidity. The absorption spectra of Nile Red dissolved in the ionic liquids were obtained with a Jasco V-550 spectrometer.

2.3. Synthetic procedure of alkyl sulfate-containing ILs 2.3.1. 1,1-Dimethylpyrrolidinium methyl sulfate ([MeMePyr][MeSO4 ]) Dimethyl sulfate (45.41 g, 360 mmol) was added dropwise to a solution of equal molar amounts of 1-methylpyrrolidine (30.65 g, 360 mmol) in 200 mL toluene, and then cooled in an ice-bath under nitrogen at a rate that maintained the reaction temperature below 313.15 K (highly exothermic reaction). The reaction mixture was stirred at room temperature for 1–4 h. After the reaction stopped, the upper organic phase of the resulting mixture was decanted, and the lower ionic liquid phase was washed with ethyl acetate (4 × 70 mL). After the last washing, the remaining ethyl acetate was removed by rotavapor under reduced pressure. The IL obtained was dried by heating at 343.15–353.15 K and stirring under a high vacuum 2 × 10−1 Pa for 48 h. The IL was kept in bottles under an inert gas. In order to reduce the water content to negligible values (lower than 0.03 mass%), a vacuum (2 × 10−1 Pa) and moderate temperature (343.15 K) were applied to the IL for several days. Yield: 93%. 1 H NMR (300 MHz, D2 O, ppm): 3.65 (s, 3H, CH3 SO4 ), 3.43 (s, 4H, N-CH2 CH2 ), 3.05 (s, 6H, N-CH3 ), 2.14 (s, 4H, N-CH2 CH2 ). 1 H NMR (300 MHz, CDCl , ppm): 3.69 (s, 3H, CH SO ), 3.56 (s, 3 3 4 4H, N-CH2 CH2 ), 3.19 (s, 6H, N-CH3 ), 2.19 (s, 4H, N-CH2 CH2 ). Elem. Anal. calcd. for C7 H17 NO4 S: C, 39.79%; H, 8.11%; N, 6.63%. Found: C, 39.58%; H, 8.07%; N, 6.51%. 2.3.2. 1-Ethyl-1-methylpyrrolidinium ethyl sulfate ([MeEtPyr][EtSO4 ]) Yield: 82.4%. 1 H NMR (300 MHz, D2 O, ppm): 4.04–3.99 (m, 2H, CH3 CH2 SO4 ), 3.41–3.33 (m, 6H, N-CH2 CH2 and N-CH2 CH3 ), 2.96–2.91 (m, 3H, N-CH3 ), 2.14 (m, 4H, N-CH2 CH2 ), 1.31–1.21 (m, 6H, N-CH2 CH3 and CH3 CH2 SO4 ). 1 H NMR (300 MHz, CDCl3 , ppm): 3.89–3.85 (m, 2H, CH3 CH2 SO4 ), 3.51–3.45 (m, 6H, N-CH2 CH2 and N-CH2 CH3 ), 3.00–2.93 (m, 3H, N-CH3 ), 2.12 (m, 4H, N-CH2 CH2 ), 1.28–0.99 (m, 6H, N-CH2 CH3 and CH3 CH2 SO4 ). Elem. anal. calcd. for C9 H21 NO4 S: C, 45.17%; H, 8.84%; N, 5.85%. Found: C, 44.98%; H, 8.87%; N, 5.71%. 2.3.3. 1,1-Dimethylpiperidinium methyl sulfate ([MeMePip][MeSO4 ]) Yield: 92.6%. 1 H NMR (300 MHz, D2 O, ppm): 3.64 (s, 3H, CH3 SO4 ), 3.25–3.22 (t, 4H, N-CH2 CH2 CH2 ), 2.99 (s, 6H, N-CH3 ), 1.77 (m, 4H, NCH2 CH2 CH2 ), 1.59–1.53 (m, 2H, N-CH2 CH2 CH2 ). Elem. anal. calcd. for C8 H19 NO4 S: C, 42.65%; H, 8.50%; N, 6.22%. Found: C, 42.45%; H, 8.47%; N, 6.11%. 2.3.4. 1-Ethyl-1-methylpiperidinium ethyl sulfate ([MeEtPip][EtSO4 ]) Yield: 87.7%. 1 H NMR (300 MHz, D2 O, ppm): 4.04–3.97 (m, 2H, CH3 CH2 SO4 ), 3.35–3.21 (m, 6H, N-CH2 CH2 and N-CH2 CH3 ), 2.93–2.83 (m, 3H, N-CH3 ), 1.79–1.57 (m, 6H, N-CH2 CH2 CH2 and N-CH2 CH2 CH2 ), 1.26–1.17 (m, 6H, N-CH2 CH3 and CH3 CH2 SO4 ). 1 H NMR (300 MHz, CDCl , ppm): 3.89–3.80 (m, 2H, CH CH SO ), 3 3 2 4 3.43–3.31 (m, 6H, N-CH2 CH2 and N-CH2 CH3 ), 2.96–2.94 (m, 3H, N-CH3 ), 1.70–1.56 (m, 6H, N-CH2 CH2 CH2 and N-CH2 CH2 CH2 ), 1.11–1.02 (m, 6H, N-CH2 CH3 and CH3 CH2 SO4 ). Elem. anal. calcd. for C10 H23 NO4 S: C, 47.41%; H, 9.15%; N, 5.53%. Found: C, 47.29%; H, 9.17%; N, 5.41%. 2.3.5. 4,4-Dimethylmorpholin-4-ium methyl sulfate ([MeMeMor][MeSO4 ]) Yield: 90.5%. 1 H NMR (300 MHz, D2 O, ppm): 3.95 (m, 4H, O-CH2 ), 3.65–3.61 (s, 3H, CH3 SO4 ), 3.41–3.39 (m, 4H, N-CH2 ), 3.16–3.12 (m, 6H, N-CH3 ). Elem. anal. calcd. for C7 H17 NO5 S: C, 36.99%; H, 7.54%; N, 6.16%. Found: C, 36.85%; H, 7.47%; N, 6.01%.

T.-Y. Wu et al. / Electrochimica Acta 55 (2010) 4475–4482

2.3.6. 4-Ethyl-4-methylmorpholin-4-ium methyl sulfate ([MeEtMor][MeSO4 ]) Yield: 91.7%. 1 H NMR (400 MHz, D2 O, ppm): 3.98 (m, 4H, OCH2 CH2 ), 3.68 (s, 3H, CH3 SO4 ), 3.49–3.40 (m, 6H, N-CH2 CH2 and N-CH2 CH3 ), 3.09 (s, 3H, N-CH3 ), 1.32 (t, 3H, N-CH2 CH3 ). Elem. anal. calcd. for C8 H19 NO5 S: C, 39.82%; H, 7.94%; N, 5.80%. Found: C, 39.77%; H, 8.00%; N, 5.66%.

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Table 1 Ionic liquid cations and anions used in this study. Ionic liquids

Cation

Anion

[MeMePyr][MeSO4 ]

2.3.7. 4-Ethyl-4-methylmorpholin-4-ium ethyl sulfate ([MeEtMor][EtSO4 ]) Yield: 85.7%. 1 H NMR (300 MHz, D2 O, ppm): 4.12–3.99 (m, 6H, O-CH2 and CH3 CH2 SO4 ), 3.60–3.35 (m, 6H, N-CH2 CH2 and NCH2 CH3 ), 3.10 (s, 3H, N-CH3 ), 1.33 (t, 3H, CH3 CH2 SO4 ), 1.19 (t, 3H, N-CH2 CH3 ). Elem. anal. calcd. for C9 H21 NO5 S: C, 42.34%; H, 8.29%; N, 5.49%. Found: C, 42.17%; H, 8.20%; N, 5.36%.

[MeEtPyr][EtSO4 ]

[MeMePip][MeSO4 ]

3. Results and discussion

[MeEtPip][EtSO4 ]

3.1. Synthesis and thermal properties of alkyl sulfate-based ILs The structures of the ionic liquids prepared in the present work are shown in Table 1. A previously reported method [22] was used to prepare these ionic liquids. Pyrrolidinium, piperidinium, and morpholinium-based ionic liquids were prepared by reacting dimethyl sulfate or diethyl sulfate with N-methylpyrrolidine, N-methylpiperidine, and N-methylmorpholine, respectively. The RTILs were obtained in medium-to-high yields (80–95%). The physical and thermal properties of seven alkyl sulfate-based products, including melting point, viscosity, density, conductivity, and thermal decomposition temperature, are shown in Table 2. The ionic liquid used for physical property measurements was dried under vacuum at 393 K for 24 h before use; the viscometer had to be airtight to avoid influence from air. Thermogravimetric analysis (TGA) provides information concerning thermal stability. Fig. 1 shows the thermogravimetric traces of alkyl sulfate-containing ILs. The thermal decomposition temperatures (Td ) of 10% weight loss are listed in Table 2, with no decomposition observed in the range of temperatures tested (up to 150 ◦ C). The first stage thermal decomposition temperatures (Td ) of these ILs are related to the structure of anions; the thermal stability of anions at Td is methylsulfate > ethylsulfate. Moreover, the second stage thermal decomposition temperatures of these ILs are related to the structure of cations in the range of 370–700 ◦ C; the thermal stability of these cations in the range of 370–700 ◦ C is pyrrolidinium > piperidinium > morpholinium. The thermal properties of the salts were investigated using differential scanning calorimetry. The melting transitions were detected during heating in the second heating–cooling scan. Methylsulfate-containing salts are solids at room temperature; however, ethylsulfate-containing salts are liquids at room temperature. The melting transition of [MeEtMor][MeSO4 ] is lower than that of [MeMeMor][MeSO4 ], implying that the ethyl group substitutes for the methyl group in cation decreases the melting transitions significantly.

[MeMeMor][MeSO4 ]

[MeEtMor][MeSO4 ]

[MeEtMor][EtSO4 ]

Fig. 1. Thermogravimetric trace for the alkyl sulfate-containing ILs.

Table 2 Physical and thermal properties of alkyl sulfate-based ILs. Ionic liquids

Mw /g mol−1

[MeMePyr][MeSO4 ] [MeEtPyr][EtSO4 ] [MeMePip][MeSO4 ] [MeEtPip][EtSO4 ] [MeMeMor][MeSO4 ] [MeEtMor][MeSO4 ] [MeEtMor][EtSO4 ]

211.28 239.33 225.31 253.36 227.28 241.31 255.33

a b

Tm /◦ C 95 9 99 −80 113 45 −68

da /g cm−3

C/mol dm−3

/mPa s

/mS cm−1

/S cm2 mol−1

Td b /◦ C

– 1.174 – 1.177 – – 1.242

– 4.905 – 4.646 – – 4.864

– 116 – 868 – – 3284

– 3.8 – 0.54 – – 0.13

– 0.775 – 0.116 – – 0.0266

322 299 329 308 325 324 301

Density (d), concentration (C), viscosity (), conductivity () and molar conductivity () are measured at 30 ◦ C. Decomposition temperature of 10% weight loss.

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T.-Y. Wu et al. / Electrochimica Acta 55 (2010) 4475–4482 Table 3 Fulcher equation parameters of viscosity for three ILs ( = o exp[−B/(T − To )]). Ionic liquids

o /mPa s

To /K

B/Ka

R2 b

[MeEtPyr][EtSO4 ] [MeEtPip][EtSO4 ] [MeEtMor][EtSO4 ]

2.543 2.059 1.974

218.2 230.2 227.4

331.7 443.6 561.7

0.999 0.999 0.999

a b

Fig. 2. (a and b) Dynamic viscosity () as a function of temperature for the ionic liquids [MeEtPyr][EtSO4 ], [MeEtPip][EtSO4 ], and [MeEtMor][EtSO4 ].

3.2. Viscosities, densities, and conductivites of ILs The viscosity of an ionic liquid is a very important parameter in electrochemical studies due to its strong effect on the rate of mass transport within a solution. For ethyl sulfate anions, the viscosity of RTILs is determined by the nature of the cation. For cations, differences in viscosity are due to hydrogen bonding and van der Waals interactions [23,24]. Fig. 2 shows the temperature dependence of viscosity data for [MeEtPyr][EtSO4 ], [MeEtPip][EtSO4 ], and [MeEtMor][EtSO4 ], whose viscosities at 30 ◦ C are 116, 868, and 3284 mPa s, respectively. The viscosity order is [MeEtMor][EtSO4 ] > [MeEtPip][EtSO4 ] > [MeEtPyr][EtSO4 ], which can be attributed to an increase in the length of alkyl substituents in cations (pyrrolidine and piperidine) or to the substitution of carbon by oxygen atoms in cations (piperidine and morpholine). The ln −1 versus 1/T plots conform to the Vogel–Tamman– Fulcher (VTF) relationship [25]:



−B A −1 = √ exp (T − To ) T

Activation energy (kJ mol−1 ). Correlation coefficient.

where T is the absolute temperature, and o , B, and To are adjustable parameters. The best-fit o (cP), B (K), and To (K) parameters are given in Table 3. The density of ionic liquids falls typically in the range 1.2–1.6 g cm−3 [30]. Fig. 3 shows the temperature dependence of density data for three selected ILs. The morpholine-based ionic liquid has a higher density than those of pyrrolidine- and piperidine-based ionic liquids. This shows that higher density liquids have a denser structure. In morphiline-based RTIL, the electron rich oxygen part likely interacts with other organic cations. The piperidine-based ionic liquid has a higher density than that of the pyrrolidine-based ionic liquid, this can be attributed to the higher density of N-methylpyrrolidine (0.819 g mL−1 ) compared to that of N-methylpiperidine (0.816 g mL−1 ). Ionic conductivity is one of the most important properties of ILs when they are considered as electrolytes. Ionic liquids have reasonably good ionic conductivities compared to those of organic solvents/electrolyte systems (up to 10 mS cm−1 ) [18]. The conductivities of [MeEtPyr][EtSO4 ], [MeEtPip][EtSO4 ], and [MeEtMor][EtSO4 ] are 3.80, 0.54, and 0.13 mS cm−1 at 30 ◦ C, respectively. As expected, [MeEtPip][EtSO4 ] has a lower conductivity than that of [MeEtPyr][EtSO4 ]. Long chains decelerate the movement of cations so ILs with longer chains have lower electrical conductivity [31–33]. [MeEtMor][EtSO4 ] has a lower conductivity than that of [MeEtPip][EtSO4 ], implying that the substitution of carbon by oxygen atoms in cations significantly decreases electrical conductivity. The ionic conductivities are shown as a function of temperature in Fig. 4. The conductivities of ionic liquids increased with increasing temperature, possibly due to the decrease in viscosity at higher temperatures. According to the hole theory [34] for transport properties in molten salts, an Arrhenius type equation for the temperature dependence of the electrical conductivity is expected [35]:  = ∞ exp

 −E  a

kB T



(1)

where A and B are constants, T is the absolute temperature, and To is the thermodynamic Kauzmann temperature. The values of viscosity were fitted using the Fulcher equation (or more commonly called VTF equation), which is commonly used for ILs [26–29]:  = o exp



B (T − To )



(2)

Fig. 3. Temperature dependence of density data for three selected ILs.

(3)

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Table 4 Fulcher equation parameters of conductivity for three ILs ( =  o exp[−B /(T − To )]). Ionic liquids

 o /mS cm−1

To /K

B /K

R2 a

[MeEtPyr][EtSO4 ] [MeEtPip][EtSO4 ] [MeEtMor][EtSO4 ]

64.4 33.6 12.7

247.3 250.2 255.5

165.5 221.9 219.3

0.999 0.999 0.999

a

Correlation coefficient.

The interdependence of ionic conductivity and viscosity is usefully considered by reference to the Walden plot of the data, as described by Angell et al. [37]. The Walden rule relates ion mobility to the fluidity of the medium: if a liquid mostly consists of independent ions, then the Walden plot will be close to an ideal line, which is typically represented by potassium chloride [38]. The Walden rule is represented by the equivalent conductivity  to the fluidity −1 of the medium through which the ions move [39,40]. The equivalent conductivities  are calculated using  = Ve , where Ve is the molar volume. The Walden plots obtained for [MeEtPyr][EtSO4 ], [MeEtPip][EtSO4 ], and [MeEtMor][EtSO4 ] over a temperature range of 30–80 ◦ C are shown in Fig. 5. The position of the ideal line was established using aqueous KCl solutions at high dilution. The curves for [MeEtPyr][EtSO4 ], [MeEtPip][EtSO4 ] and [MeEtMor][EtSO4 ] lie closely to the ideal KCl line. A similar result was reported for hydrogen sulfate-based protic ionic liquids [41]. Most ionic liquids lie below the ideal line [39,40]. The slopes of the lines at various temperatures for [MeEtPyr][EtSO4 ], [MeEtPip][EtSO4 ] and [MeEtMor][EtSO4 ] are 1.074, 0.931 and 0.939, respectively, which are close to 1, implying that the conductivity–fluidity relationship remains constant. 3.3. Electrochemical properties

Fig. 4. (a and b) Arrhenius plot of specific conductivity () for the ionic liquids [MeEtPyr][EtSO4 ], [MeEtPip][EtSO4 ], and [MeEtMor][EtSO4 ].

where Ea is the activation energy for electrical conduction (which indicates the energy needed for an ion to jump to a free hole),  ∞ is the maximum electrical conductivity (that it would have at infinite temperature), and kB is the Boltzman constant. However, some observed temperature dependence of conductivity are not linear but polynomial (n = 2 or 3), and are often best fitted by the empirical Vogel–Tammann–Fulcher (VTF) or Fulcher equation [27,29,36]. The dominant conduction mechanism of ionic liquid electrolytes can be obtained from the ln  versus 1/T plots. When ionic transport is coupled to the relaxation of the host polymer, the ln  versus 1/T plots conform to the Vogel–Tamman–Fulcher (VTF) relationship [36]:



−B A  = √ exp (T − To ) T

The electrochemical window, the potential interval where no charge transfer processes take place, is the most important parameter with regard to possible electrochemical applications of ionic liquids. The electrochemical stabilities of the ionic liquids were analyzed using cyclic voltammetry. Fig. 6 shows the cyclic voltammograms for three types of ionic liquid containing a sulfate group recorded at a glassy carbon electrode. Similar results were obtained for a platinum electrode. The cathodic stability of these materials is mainly determined by the potential at which the reduction of the cations takes place. The reduction process is expected to mainly involve the positively charged nitrogen in the pyrrolidinium, piperidiniuim, and morpholinium rings. The anodic stability of



(4)

where  is the specific conductivity, A is a pre-exponential factor proportional to T−1/2 and somewhat related to the number of charge carriers, and B (Ea ) has the dimension of energy and is related to the critical free volume for ion transport. To is a reference temperature, which closely related to the glass transition temperature Tg . The Fulcher equation was also fitted to the RTILs, which behaved as (r2 > 0.97) [27–29]. The values of conductivity were fitted using the equation:  = o exp



−B (T − To )



(5)

where  o (mS cm−1 ), B (K), and To (K) are constants. The best Fulcher equation fitting parameters are given in Table 4.

Fig. 5. Walden plots for sulfate-containing ionic liquids. The straight line plot is generated from data obtained in aqueous 0.01 M KCl solution.

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Fig. 6. Cyclic voltammogram of [MeEtPyr][EtSO4 ], [MeEtPip][EtSO4 ], [MeEtMor][EtSO4 ], [MeEtMor][MeSO4 ] relative to Pt reference electrode, with glassy carbon working electrode and Pt counter electrode at 30 ◦ C. Potential was reported against the redox potential of ferrocene/ferricenium (Fc/Fc+ ) redox couple measured in each ionic liquid.

these materials is mainly determined by the potential at which the oxidation of the anions takes place. With regard to the anodic limit, the sulfate anion is oxidized at the anodic potential (1.53 V vs. Fc/Fc+ ). The estimated cathodic and anodic limiting potentials are summarized in Table 5. The redox potential of ferrocene dissolved in the ionic liquids was used as a standard. All the ionic liquids exhibited wide electrochemical windows of at least 4.5 V, indicating the contribution of the sulfate group to the electrochemical stability of these salts. Moreover, the electrochemical windows of these ionic liquids are wider than those of imidazole containing ionic liquids [19], implying that pyrrolidinium, piperidinium, and morpholinium units significantly increase the electrochemical window.

Fig. 7. (a) Cyclic voltammograms obtained for 3.00 mM ferrocene in [MeEtPyr][EtSO4 ] at a glassy carbon electrode (1 mm diameter) at scan rates of 0.01, 0.02, 0.05, 0.10, 0.20, 0.35, and 0.5 V s−1 . (b) The dependence of peak current on square root of the scan rate for oxidation of ferrocene.

3.4. Diffusion coefficient of ferrocene

process:

Representative cyclic voltammograms of ferrocene in [MeEtPyr][EtSO4 ] are shown in Fig. 7(a). The voltammograms illustrate that as the potential was scanned to more positive value, an anodic peak current (ipa ) is observed at the anodic peak potential (Epa ), indicating that ferrocene Fe(C5 H5 )2 is converted to its oxidized form Fe(C2 H5 )2 + . During the reverse scan, reduction of Fe(C2 H5 )2 + occurs and a cathodic peak current (ipc ) is observed at the cathodic peak potential (Epc ). The anodic and cathodic peak separation (Ep ) in the cyclic voltammograms was found to be 0.06 V (the ideal value that is indicative of one-electron oxidation). The linear dependence of cathodic peak current (ip ) on square root of scan rate (1/2 ), shown in Fig. 7(b) confirms the oxidation process of ferrocene in [MeEtPyr][EtSO4 ] is diffusion controlled. In general, the peak current of diffusion controlled reversible electrochemical reaction follows Randles-Sevcik equation [42], which assumes that mass transport occurs only by a diffusion

ip = 0.4463nF

 nF 1/2 RT

CAD1/2 1/2

(6)

where ip is the peak current, n is the number of electron equivalents exchanged during the oxidation/redox reversible process (electron stoichiometry), A is the active surface area of the working electrode (cm2 ), D is the diffusion coefficient (cm2 s−1 ), c is the bulk concentration of the diffusing species (mol cm−3 ),  is the voltage scan rate (V s−1 ), F is Faraday’s constant, and R is the universal gas constant. The diffusion coefficient of ferrocene in [MeEtPyr][EtSO4 ] was determined to be 1.33 × 10−7 cm2 s−1 at 323 K. The electrochemical behavior of ferrocene in [MeEtPyr][EtSO4 ] was further studied using chronoamperometry, the chronoamperograms for the oxidation of ferrocene in [MeEtPyr][EtSO4 ] at 323 K at various constant potentials are shown in Fig. 8(a). Chronoamperometry was made through the potential step from −0.253 V to various values of potential ranging from −0.203 to 0.247 V in

Table 5 Electrochemical windows of ionic liquids. Ionic liquid

Cathodic limiting potential vs. (Fc/Fc+ )/V

Anodic limiting potential vs. (Fc/Fc+ )/V

Electrochemical window/V

[MeEtPyr][EtSO4 ] [MeEtPip][EtSO4 ] [MeEtMor][EtSO4 ] [MeEtMor][MeSO4 ]

−3.38 −3.44 −3.43 −3.45

1.53 1.54 1.53 1.53

4.91 4.98 4.96 4.98

T.-Y. Wu et al. / Electrochimica Acta 55 (2010) 4475–4482

4481

Table 7 Wavelengths of maximum absorption (max ) and molar transition energies (ENR ) for Nile Red dissolved in ionic liquids. Ionic liquid

Fig. 8. (a) Chronoamperograms and (b) Cottrell plot for ferrocene (3 mM) in [MeEtPyr][EtSO4 ] at a glassy carbon electrode at various applied potentials.

the time domain from 0.002 to 10 s. As is shown in Fig. 8(b), the Cottrell plot showed an approximately linear relation in the time domain from 1 to 3 s. The current at the time shorter than 1 s deviated upward from the linear line probably because of a delay of the potentiostat due to extremely large Faradaic current at a short time. From the Cottrell equation [42]: i=

nFAD1/2 C

(7)

1/2 t 1/2

where i is current (unit: A), n is number of electrons (to reduce/oxidize one molecule of analyte), F is Faraday constant, A is the area of electrode (unit: cm2 ), C is initial concentration of the reducible analyte (unit: mol cm−3 ), D is diffusion coefficient (unit: cm2 s−1 ), t is time (unit: s), the diffusion coefficient of ferrocene in [MeEtPyr][EtSO4 ] at 0.6 V was determined to be 2.57 × 10−7 cm2 s−1 at 323 K.

max /nm

ENR /kcal mol−1 a

Reference

565.4

50.6

48

564.4

50.7

48

562.5

50.8

48

556.0

51.4

47

555.7

51.4

47

550.8

51.9

47

547.5

52.2

47

548.7

52.1

47

[MeEtPyr][EtSO4 ]

570

50.2

This work

[MeEtPip][EtSO4 ]

564

50.7

This work

[MeEtMor][EtSO4 ]

580

49.3

This work

Water

593.2

48.2

47

Methanol

549.8

52.0

47

Ethanol

548.8

52.1

47

DMF

541.5

52.8

47

CHCl3

537.4

53.2

47

CH3 CN

531.4

53.8

47

Hexane

484.6

59

47

ENR = (hcNA /max ) × 10 , where h is Planck’s constant, c is the speed of light, NA is Avogadro’s number, and max is the wavelength of maximum absorption (nm). a

6

The Stokes–Einstein relationship [43] (Eq. (8)) predicts a linear dependence of D on the reciprocal of viscosity (): D=

kB T 6˛

(8)

where kB is the Boltzmann constant, T is the temperature, and ˛ is the radii of the diffusing entities. Table 6 summarizes the calculated Stokes–Einstein product, D/T of ferrocene in [MeEtPyr][EtSO4 ], these values are comparable to other Stokes–Einstein product of ferrocene in functionalized

Table 6 The calculated Stokes–Einstein product, DT−1 of ferrocene in ionic liquids. Ionic liquid

D/×107 cm2 s−1

/cp

Temp/K

DT−1 /×1010 g cm s−2 K−1

Reference

[emim][NTf2 ] [mimSBu][NTf2 ] [MeEtPyr][EtSO4 ] [MeEtPyr][EtSO4 ] [TEtA][Ac] [TEtA][Of]

5.09a 1.35a 1.33a 2.57b 8.1a 8.7a

34.7 98.5 62.43 62.43 11 10

294 294 323 323 298 298

6.01 4.52 2.57 4.96 2.99 2.92

44 44 This work This work 45 45

a b

Cyclic voltammetry. Chronoamperometry.

4482

T.-Y. Wu et al. / Electrochimica Acta 55 (2010) 4475–4482

ionic liquids [emim][NTf2 ] (1-ethyl-3-methylimidazolium bis(trifluoromethanesulfonyl)amide) [44], [mimSBu][NTf2 ] (2butylthiolonium bis(trifluoromethanesulfonyl)amide) [44], [TEtA][Ac] (triethylammonium acetate) [45] and [TEtA][Of] (triethylammonium formate) [45].

Acknowledgement

3.5. Polarity study of some sulfate-containing ILs with solvatochromic dye Nile Red

References

The polarity of ambient-temperature ionic liquids based on [MeEtPyr][EtSO4 ], [MeEtPip][EtSO4 ], and [MeEtMor][EtSO4 ] was probed using solvatochromic dye Nile Red. Nile Red is extensively used in many applications where the dipolarity of the medium needs to be explored due to its photochemical stability and strong fluorescence nature [46]. Nile Red is a positive solvatochromic dye and shows one of the largest shifts in excitation and emission maxima in going from nonpolar solvents (in hexane, max ∼ 484.6 nm) to polar solvents (in water, max ∼ 593.2 nm) (Table 7). For polarity measurements, a Nile Red concentration was chosen such that absorbances would fall in the range 0.5–1.0. The wavelengths of maximum absorption (max ) and the molar transition energies (ENR ) for Nile Red dissolved in the ionic liquids are summarized in Table 7, along with the values for several widely used solvents and imidazole cation containing ionic liquid [47,48] for comparison. The dipolarity (ENR value) of alkyl sulfatebased ionic liquids is lower than those of general organic solvents, such as methanol, ethanol, DMF, CHCl3 , CH3 CN, and hexane, but it is similar to those of some imidazole-based ionic liquids. The morpholine-based ionic liquid has a lower ENR value (larger polarity) than those of pyrrolidine- and piperidine-based ionic liquids. Large polarity on morpholine-based cations increases the interactions between ions; therefore, morpholine-based ionic liquids have a higher density than those of pyrrolidine- and piperidinebased ionic liquids. The strong solvatochromic behavior observed for Nile Red arises from the fact that it undergoes large dipole moment changes during the transition between ground and excited states [49]. As these ILs become alternatives for conventional solvents, polarity data will become important to the synthetic chemist.

4. Conclusions Seven ILs containing pyrrolidinium, piperidinium, morphinium units as the cations, and the alkyl sulfate group as the anions, were designed and synthesized via a halide-free approach. The thermal properties, density, viscosity, conductivity, electrochemical window, and polarity were measured and investigated in detail. The obtained physical properties show that the effect of cation identity is preponderant. Polarities were measured spectroscopically using Nile Red dye. The dipolarity (ENR value) of alkyl sulfatebased ionic liquids is lower than those of general organic solvents, such as methanol, ethanol, DMF, CHCl3 , CH3 CN, and hexane, but it is similar to those of some imidazole-based ionic liquids. Cyclic voltammetry measured in neat alkyl sulfate-based ionic liquids at a glassy carbon electrode indicated at least 4.5 V wide potential windows. The presented IL characterization study is of importance to obtain a better understanding of their application scope. These results could facilitate the development of new task-specific ILs.

The authors would like to thank the National Science Council of the Republic of China for financially supporting this project under grant: 96-2113-M-006-014-MY3.

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