Generation of Superoxide Ion in Pyridinium, Morpholinium

Article pubs.acs.org/IECR

Generation of Superoxide Ion in Pyridinium, Morpholinium, Ammonium, and Sulfonium-Based Ionic Liquids and the Application in the Destruction of Toxic Chlorinated Phenols Maan Hayyan,○,# Farouq S. Mjalli,§,* Mohd Ali Hashim,○,# and Inas M. AlNashef∥ ○

University of Malaya Centre for Ionic Liquids (UMCiL), University of Malaya, 50603 Kuala Lumpur, Malaysia Department of Chemical Engineering, University of Malaya, 50603 Kuala Lumpur, Malaysia § Petroleum and Chemical Engineering Department, Sultan Qaboos University, Muscat 123, Oman ∥ Chemical Engineering Department, King Saud University, Riyadh 11421, Saudi Arabia #

S Supporting Information *

ABSTRACT: Generation of superoxide ion (O2•−) was carried out in four ionic liquids (ILs) having the same anion, bis(trifluoromethylsulfonyl)imide [N(Tf)2]−, and different cations, N-hexylpyridinium [HPy]+, N-methoxyethyl-N-methylmorpholinium [MO1,1O2]+, N-ethyl-N,N-dimethyl-2-methoxyethylammonium [N112,1O2]+, and triethylsulfonium [S222]+. Cyclic voltammetry (CV) and chronoamperometry (CA) electrochemical techniques were used in this investigation. It was found that O2•− is not stable in the [HPy]+-based IL. On the other hand, CV showed that the electrochemically generated O2•− is stable in [MO1,1O2]+-, [N112,1O2]+-, and [S222]+-based ILs for the time duration of the experiment. The long-term stability of the generated O2•− was then investigated by dissolving potassium superoxide (KO2) in dimethyl sulfoxide (DMSO) in the presence of the corresponding IL. It was found that ILs containing [MO1,1O2]+ and [N112,1O2]+ offer a promising long-term stability of O2•− for various reactions to be used for several applications. However, it was found that after 2 h, about 92.5% of the generated O2•− in [S222]+ based IL was consumed. The diffusion coefficient and solubility of O2 in the studied ILs were then determined using CV and CA techniques simultaneously. It was found that diffusion coefficients and CA steady-state currents increase with temperature increases, while the solubility of O2 decreased. To our best knowledge, this is the first time that morpholinium and sulfoniumbased ILs were utilized as media for chemical and electrochemical generation of O2•−. Additionally, the chemically generated O2•−, by dissolving KO2, was then used for the destruction of 2,4-dichlorophenol (DCP) in [MO1,1O2][N(Tf)2] under ambient conditions. The destruction percentage was higher than 98%. This work represents a novel application of the chemically generated O2•− for the destruction of toxic chlorinated phenols in ILs media.

1. INTRODUCTION O2 reduction is considered one of the important electrochemical reactions because of its role in electrochemical energy conversion, several industrial processes, and corrosion reactions. Consequently, for many years, it has been under the consideration of electrochemical research studies.1 The reduction of O2 by electron transfer yields O2•−, which has a negative charge and an electronic spin density delocalized between the two oxygen atoms.2 Due to the vital importance in biological systems and electrochemical applications (e.g. fuel cells), the protonation reaction of the O2•− species has been widely investigated in various solutions containing a proton source, eq 1.3,4 O2•− + (proton source) → HO2−

low boiling points, and the negative ecological effects. Hence, ILs are more preferable for O2•− generation. ILs have become popular alternative media for the reduction of O2 to O2•−.6−8 Many ILs have been investigated on the basis of their structures (cation and anion) such as hexafluorophosphate, chloroaluminate, trifluoromethylsulfonyl, and tetrafluoroborate, combined with alkylated imidazolium or ammonium cations.8,9 Many research groups studied the structure effects of the anion and the cation on the stability of the generated O2•−.4−6,8−11 It was found that the cations of some ILs tend to react with O2•− reduced from the dissolved O2, such as quaternary phosphonium cation,8,12 pyridinium,5 and imidazolium cation.4,13,14 In spite of the fact that all researchers agree that O2•− can be generated chemically and electrochemically in ILs,13,15 the stability of this radical anion in the media under consideration is still being investigated.16

(1)

O2•−

The is a safe and mild oxidizing agent that has several potential applications. Different aprotic solvents were utilized as conventional media for O2•− generation, such as dimethyl formamide (DMF), dimethyl sulfoxide (DMSO), acetonitrile (AcN), propylene carbonate, pyridine, methylene chloride, and acetone.5 However, although O2•− is a long-lived species in aprotic media, the used solvents are not benign. Therefore, so far, there are no industrial applications for O2•− due to the restrictions of using aprotic solvents such as their high volatility, © 2012 American Chemical Society

Received: Revised: Accepted: Published: 10546

March 14, 2012 June 18, 2012 July 26, 2012 July 26, 2012 dx.doi.org/10.1021/ie3006879 | Ind. Eng. Chem. Res. 2012, 51, 10546−10556

Industrial & Engineering Chemistry Research

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Scheme 1. Structures of the Cations and Anions of Used ILs

commonly wide EW and hydrophobic characteristic of the anion.22,23 Furthermore, the long-term stability of O2•− was tested, and the kinetic rate constant of O2•− was determined. In addition, the chemically generated O2•− by dissolving KO2 was used for the destruction of 2,4-dichlorophenol (DCP) in [MO1,1O2][N(Tf)2] under ambient conditions. It should be emphasized that O2•− is very sensitive to the presence of small traces of water, since O2•− disproportionates in the presence of water. Therefore, the generation of O2•− and its reaction with chlorinated phenol must be conducted under controlled atmosphere (e.g., glovebox). However, it should be noted that this is not a serious limitation for the use of O2•− in the industry because there are many industrial applications that need controlled atmosphere, such as processes using catalysts sensitive to water or oxygen, production of sodium, and other alkali metals, etc.

The cyclic voltammetry (CV) technique was utilized to study the short-term stability of O2•− in ILs. Though, it is difficult to give a real indication for O2•− stability as the time needed to run CV is less than a few minutes in most cases. To use O2•− in practical applications, the long-term stability is necessary since some important applications of O2•− need longer time than that used in CV. Islam et al. (2009) and AlNashef et al. (2010) reported the long-term stability of O2•− in imidazolium-based ILs.4,13 Recently, we have shown the long-term stability of O2•− in piperidinium, pyrrolidinium, and phosphonium cation-based ILs.8,10 As an application for the use of O2•−, we selected the destruction of chlorinated hydrocarbons (CHCs), in particular, chlorinated phenols. The disposal of chlorinated and brominated industrial waste is a matter of increasing concern.17 Typically, there are several recognized methods for the degradation of CHCs, including incineration, biological, photochemical, chemical, and electrochemical oxidation.10 However, there are many disadvantages in the use of these methods such as high energy consumption, production of toxic byproduct [acid gases (HCI, HF, or HBr) and/or free halogen gases (Cl2, F2, or Br2)], economical viability, and high degree of sophistication. The use of O2•− generated chemically from its salts (e.g. KO2) and electrochemically by reduction of O2 for the destruction of CHCs in aprotic solvents was investigated.18,19 It was concluded that O2•− can degrade CHCs via nucleophilic substitution reaction to carbonates and halide ions.20 However, the volatility of used aprotic solvents as well as the limited solubility of KO2 in these organic solvents make the utilizing of O2•− not applicable industrially. ILs based on [N(Tf)2]− are attractive media for many applications because of their good properties, such as wide electrochemical window (EW), high thermal stability, high ionic and electrical conductivity, and lower viscosity compared to other types of anions. Moreover, they are hydrophobic and stable against moisture,21 and hence, their water content will be less than that of hydrophilic ILs. It is well-known that the presence of water decreases the EW.22 Therefore, ILs based on [N(Tf)2]− are expected to be more suitable for the use as media for the reduction of O2 to O2•−. In this article, four ILs were used as media for the electrochemical and chemical generation of O2•−. These ILs are pyridinium [HPy]+-, morpholinium [MO1,1O2]+-, ammonium [N112,1O2]+-, and sulfonium [S222]+-based ILs, with the common anion of bis(trifluoromethylsulfonyl)imide [N(Tf)2]−. ILs based on [N(Tf)2]− were selected for their

2. EXPERIMENTAL AND MATERIAL SECTION All studied ILs were supplied by Merck and were of synthesis grade. Scheme 1 shows the structures of the cations and anion that make up the ILs. Formulas, molecular weights and melting points of the ILs are listed in Table 1S in the Supporting Information. 2.1. Electrochemical Generation of Superoxide Ion. CV tests were performed as the electrochemical analysis technique, since this method is extremely powerful and is among the most widely practiced of all electrochemical methods.24 The IL to be used was dried overnight in a vacuum oven at 50 °C. It should be noted that some used ILs were acidic without pretreatment, pH = 4−6. The pH of ILs was measured using pH strips (Merck) and a very small quantity of KO2 was added to the acidic IL until its pH became 7. AlNashef et al. (2001) reported that O2•− was unstable in some ILs because of the acidity of these ILs.6 Therefore, small additions of KO2 can neutralize the acidic ILs. The electrochemistry was performed using EG&G 263A potentiostat/galvanostat (PAR) connected to computer with data acquisition software. CVs were conducted in a one compartment cell because the time of the experiment is relatively small to affect the ILs. The cell was a jacketed vessel (10 mL volume) with a Teflon cap including 4 holes for the three electrochemical electrodes and a gas sparging tube. A glassy carbon (GC) macroelectrode (BASi, 3 mm diam.) was used as working electrode for CV, while carbon fiber (CF) ultra-microelectrode (BASi, 11 μm diam.) was used for chronoamperometry (CA). A platinum electrode was used as a counter electrode, and an aqueous Ag/AgCl electrode (BASi) 10547

dx.doi.org/10.1021/ie3006879 | Ind. Eng. Chem. Res. 2012, 51, 10546−10556


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was used as a reference electrode. The macroelectrodes were polished using alumina solution (BASi) and sonicated in distilled water for 10 min before each experiment. This was done to ensure that there are no impurities on the surface of the working electrode. The ultra-microelectrodes were polished using diamond solution (BASi) and then sonicated in distilled water for 10 min. All experiments were performed in a dry glovebox, with tight humidity control of less than 1 ppm water, under either an argon or helium atmosphere. Prior to O2•− generation, a background voltammogram was obtained after removal of O2. The O2 removal was achieved by purging the IL with dry N2. This particular method is quite effective and also simple to be employed. Purging a solution with an inert gas can reduce the partial pressure of O2 above the solution, and consequently, the solubility of dissolved O2 in the solution is decreased.25,26 O2 was then bubbled into tested IL for at least 30 min to ensure that equilibrium was achieved.5,27 To confirm that the tested IL is saturated with O2, CVs at different time intervals were conducted and the final measurement was taken when the cathodic peak current of the CV is constant. Between consecutive CV runs, O2 was bubbled briefly to refresh the system and to remove any concentration gradients. N2 or O2 sparging was discontinued during the CV runs. CA measurements were conducted inside a Faraday cage to avoid any interference. The value of the steady state background current was deducted from the value obtained after sparging with O2 to provide the net steady-state current. The net value of the current was then used in the calculations. 2.2. Long-Term Stability of O2•−. DMSO (Fisher, 99.98%) of spectroscopic grade was dried overnight in a vacuum oven. KO2 (Sigma Aldrich, 99.9%) was kept in a sealed vial filled with molecular sieves. The chemical generation of O2•− was performed by dissolving KO2 in DMSO while stirring with a magnetic stirrer. Subsequently, a certain amount of IL was added to the generated O2•− in DMSO to investigate the stability of O2•− with time. A computer-controlled UV−visible spectrophotometer (PerkinElmer-Lambda 35) was used to measure the absorption spectra of O2•− every 10 min until reaching 2 h of reaction time. The reference solution of spectral measurements was DMSO or DMSO solution containing an appropriate amount of IL. 2.3. Destruction of Chlorinated Hydrocarbon Using Potassium Superoxide. An appropriate amount of selected chlorinated hydrocarbon, 2,4-dichlorophenol (ALDRECH 99%) was added to a labeled vial containing the IL. KO2 (Sigma Aldrich, 99.9%) was then added gradually to the vial containing the IL and DCP under vigorous stirring. Samples, before and after the addition of KO2, were taken by dissolving 0.1 g of IL-DCP in 1 g of AcN (UNICHROM, high pressure liquid chromatography (HPLC) grade 99.9%). The samples were then analyzed using an HPLC instrument (Shimadzu, Japan); the conditions are illustrated in Table 2S in the Supporting Information. This procedure was repeated and more KO2 was added until the peak of chlorinated hydrocarbon was not detected or no significant change between the last two peaks after additions could be detected.

Table 1. EWs of ILs at Cut-off Current Density of 0.5 mA/ cm2, Using GC Macroelectrode at Sweep Rate 100 mV/s IL

EC,L

EA,L

EW

[HPy][N(Tf)2] [MO1,1O2][N(Tf)2] [N112,1O2][N(Tf)2] [S222][N(Tf)2]

−1.22 −1.93 −3.47 −1.30

2.74 2.75 2.52 2.21

3.96 4.68 5.99 3.51

consistent with other reported EWs of ILs.28,29 The typical EWs for ILs are 4.5−5 V although some ILs such as 1-n-butyl3-methylimidazolium cation with tetrafluoroborate and hexafluorophosphate have wider EWs (up to 7.00 V).30 The [N112,1O2][N(Tf)2] EW is close to the EWs reported in the literature, where the reductive limit of ammonium based ILs was found to be approximately −3 V, irrespective of the substituents attached to the ammonium cation.31 It can be clearly noticed that the cation affects the EW, since all the tested ILs have the same anion, [N(Tf)2]−. This is in agreement with results reported by other groups. It was shown in the literature that the oxidation of the anion and the reduction of the cation are responsible for the anodic and cathodic limits observed in ILs.32,33 Nevertheless, it was shown that, in a few cases, there are some ILs with the same anion or cation gave different anodic or cathodic stability.8,33 From Table 1, it can be noticed that the [N112,1O2]+-based IL, followed by the [MO1,1O2]+-based IL, has the widest reductive windows, with all studied ILs having EW greater than 3.00 V. This is in agreement with results reported in the literature.33 The reductive windows indicate the feasibility of using all studied ILs as medium for O2•− generation, since the potential for O2 reduction is ≅ −1 V.34 It can be noticed that [HPy][N(Tf)2] and [S222][N(Tf)2] have smaller cathodic potential limits compared to the other two ILs. As previously mentioned, the reductive limit is presumed to be determined by the reduction of the cation so it would be expected that varying the cation of an IL with a common anion would provoke a variation in the reductive potential window depending on the stability of the cation to reduction.33 3.2. Electrochemical Generation of O2•−. CV for the reduction of O2 to O2•− in ILs was studied at sweep rates of 9, 36, 64, 81, 100, and 144 mV/s. The resulting CVs are shown in Figures 1 and 2. The absence of any peaks after sparging of

3. RESULTS AND DISCUSSION 3.1. Electrochemical Window. Table 1 illustrates the EW of the studied ILs (Figure 1S in the Supporting Information). In general, they have wide potential ranges, and this is

Figure 1. Cyclic voltammogram of [HPy][N(Tf)2] at 25 °C and sweep rate of 100 mV/s. 10548



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Figure 2. Cyclic voltammograms of (a) [MO1,1O2][N(Tf)2], (b) [N112,1O2][N(Tf)2], and (c) [S222][N(Tf)2] at 25 °C and sweep rates of 9, 36, 64, 81, 100, and 144 mV/s.

the solvent employed and, in particular, its acidity. The mechanism of O2 reduction critically depends on the pH of the medium as well as the nature of the electrode material. For instance, in acid media, O2 electro-reduction proceeds by a twoelectron reduction of O2 to H2O2 on polycrystalline Au and single crystal Au electrode surfaces. Conversely, the reduction mechanism changes to a four-electron process of O2 to H2O at pH> 6 on Au electrode. In nonaqueous media, polycrystalline Au electrode supports 1-electron reduction of O2 to O2•−.36 In contrast, Figure 2 shows that both the forward reduction peak and the backward oxidation peak are present. The presence of the backward peak confirms that the generated O2•− is stable within the time limits of the experiment. The reduction and oxidation peak currents, at the same sweep rate, are different for the three studied ILs. This shows the effect of the structure of the cation of the IL on the solubility and diffusivity of O2 in ILs. It can be clearly seen that [S222]+ based IL has higher reduction peak currents than [N112,1O2]+ and [MO1,1O2]+ based ILs. From the analysis of the CVs at various sweep rates, it can be seen that the separation between the cathodic peak, Epc and the half peak potential, Ep/2c, |Epc − Ep/2c|, is not 56.5 mV, Table 2.24 Moreover, the difference between the reduction and

nitrogen in the studied ILs confirms that there are no electrochemically active impurities in the ILs and that the IL is electrochemically stable in this range of potential. From Figure 1, it can be seen that there is a reduction peak indicating the generation of O2•− at approximately −1.06 V vs Ag/AgCl in [HPy][N(Tf)2]. However, the absence of any oxidation peak in the backward sweep indicates that generated O2•− is not stable in this IL. The instability may be due to the reaction of O2•− with the cation of the IL or with products of the decomposition of the cation since the voltage at which the reduction of O2 occurs is very close to the negative limit of the EW of the IL. This is in agreement with previous work, where it has been shown that O2•− was unstable in N-(3-hydroxypropyl) pyridinium bis(trifluoromethylsulfonyl)imide, [HPPy][N(Tf)2].5 This is also in agreement with Martiz et al. (2004), where they reported that after grafting of pyridinium-based IL with aliphatic perfluorinated chain, [C5H5N+C8F18][N(Tf)2], a 105% increase of O2 solubility was noticed and a stable oxidation signal of O2•− was observed.35 The existence of a proton source leads to the spontaneously rapid disproportionation of the electrogenerated superoxide species.14 Various studies have proven that reduction of O2 is a complex process, the outcome of which is highly dependent on 10549



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Table 2. Net Values of |Epc − Ep/2c| for Sweep Rates 9−144 mV/s at 25, 35, and 45 °C Using GC Macroelectrode |Epc − Ep/2c| (mV) [MO1,1O2][N(Tf)2]

[N112,1O2][N(Tf)2]

[S222][N(Tf)2]

sweep rate (mV/s)

25 °C

35 °C

45 °C

25 °C

35 °C

45 °C

25 °C

35 °C

45 °C

144 100 81 64 63 9

150.60 150.60 137.60 141.45 125.05 106.10

148.53 136.15 131.64 127.71 115.72 93.70

130.35 125.83 120.25 111.98 109.40 88.58

119.38 113.16 109.40 111.00 99.15 88.55

127.05 132.03 118.26 116.45 105.22 90.47

136.63 120.25 113.60 111.97 105.20 93.38

120.32 115.53 114.92 107.50 100.33 100.91

122.60 114.92 109.75 111.20 106.11 112.38

141.43 140.82 139.85 138.63 137.05 117.58

separations in all studied ILs decreased with increasing temperature, suggesting faster electrode kinetics.9 The stability of [N112,1O2]+ is in accordance with Katayama et al. (2004) where they deduced that O2•− is substantially stable in N-hexyltrimethylammonium bis(trifluoromethanesulfone)imide. The high stability of O2•− in the ILs consisting of ammonium cations is also anticipated because O 2 •− is known to form an ionic salt of tetramethylammonium superoxide.40 The linear relationship between the reduction peak currents and the square root of sweep rate, Figure 3, indicates that the electrochemical reaction is diffusion controlled; this is in agreement with O’Toole et al. (2007).38 Figure 4 shows that the reduction peak currents of CVs increase as temperature increases. This shows that the increase in diffusion coefficient of O2 due to the decrease of viscosity of IL is much higher than the decrease of solubility of O2. This is in agreement with results reported by Huang et al. (2009).9 As the concentration of O2 in the studied ILs is unknown, CV technique alone is not enough to determine the diffusion coefficient of O2 in the studied ILs. Therefore, in addition to CV, CA experiments using a CF ultra-microelectrode were conducted to find the solubility and diffusion coefficient of O2 simultaneously. Figure 5 shows that the steady-state current of CA increases with the increase of temperature. Again, this is due to the reduction of the viscosity of the ILs with increasing temperature and consequently the increase of O2 diffusivity. 3.3. Calculation of Diffusion Coefficient and Solubility of O2. Both CV at macroelectrode and CA at ultramicroelectrode techniques were used to determine the diffusion coefficients and solubility of O2 in the studied IL. The following three equations were used:24

oxidation peaks changes with sweep rate. This indicates that the electrochemical generation of O2•− in the studied ILs is not reversible. The peak currents and peak potentials are proportional to the square root and the log of the sweep rate, respectively, Figure 3. This shows that the process is mass

Figure 3. Peak current vs square root of sweep rates [MO1,1O2][N(Tf)2], [N112,1O2][N(Tf)2], and [S222][N(Tf)2] at 25 °C.

diffusion controlled and that it is irreversible.24 This is in agreement with results reported for other ILs.6,37 It should be noted that strong (bond forming) ion-paring or protonation leads to irreversibility.38 Sawyer discussed the variation in the peak-separation values ΔEp for the CV of the generation of O2•− in different classical aprotic solvents. He attributed that to the heterogeneous electron-transfer kinetics; however, the most reasonable explanation is the effect of the uncompensated resistance and surface reactions especially for the metal electrodes.19 Similarly, O’Toole et al. (2007) ascribed the slightly larger ΔEp observed in their CVs to the effect of some residual uncompensated solution resistance (iR).38 The separation between the oxidative and reductive peaks also depends largely on the electrode substrate, and a broadening of voltammetry on Pt compared to Au and GC in the order Pt > Au > GC has been reported, suggesting possibly coupled chemical reaction or product adsorption on Pt.14 Electrolysis of a dissolved species at an electrode creates a concentration gradient, triggering diffusion of the species from bulk solution and establishing a diffusion layer extending out from the electrode surface.39 Huang et al. (2009) studied the generation of O2•− in different types of ILs. They confirmed that the peak-to-peak

i p = (2.99 × 105)α 0.5ACoDo0.5v 0.5

Ep = E 0 ′ −

⎡ ⎛ D1/2 ⎞ ⎛ αFv ⎞1/2 ⎤⎥ RT ⎢ ⎟ 0.780 + ln⎜⎜ o0 ⎟⎟ + ln⎜ ⎝ RT ⎠ ⎥⎦ αF ⎢⎣ ⎝ k ⎠

iss = 4nFDoCoro

(2)

(3) (4)

where, in eq 2, ip is the cathodic peak current of CV in A; α is the charge transfer coefficient; A is the surface area of the macro-working electrode in cm2; Co is the bulk concentration of O2 in mol/mL; Do is the diffusion coefficient of O2 in cm2/s; and v is the potential sweep rate in V/s. In eq 3, Ep is the peak potential for the cathodic current in V; E0′ is the formal potential of the reaction; R is the universal gas constant in J/ (mol·K); T is the absolute temperature in K; F is Faraday’s constant 96485 C/mol; and k0 is the standard heterogeneous rate constant in cm/s. In eq 4, iss is the steady-state current of 10550



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Figure 4. Cyclic voltammograms of (a) [MO1,1O2][N(Tf)2] and (b) [N112,1O2][N(Tf)2] at sweep rate of 100 mV/s and temperatures of 25, 35, and 45 °C.

The charge transfer coefficient α was found, in general, to be independent of temperature. This is in agreement with that reported previously.36 The relatively low α values, which are not in the typical region of 0.5, suggest sluggish kinetics, which may be due to the formation of a passive oxide layer on the surface of the electrode.41 Although the mechanism for O2 reduction is found to be the same in ILs and conventional aprotic solvents, the different behavior in ILs suggests that the kinetics of diffusion of O2 and superoxide is sensitive to the choice of the IL solvent; this is in agreement with Silvester et al. (2008).42 Faster diffusion of the neutral O2 leads to more steady-state-like behavior, whereas the slower diffusion of the charged radical anion, O2•− results in a more transient reverse peak.9,12 The fast diffusion of O2 can be explained by the weak interaction between O2 and the IL because there is no Coulombic interaction between the organic ions and neutral O2. Additionally, the diffusion of O2 seems not to be affected by the apparent viscosity of IL, probably because O2 is small enough to move through the interstices between the bulky organic ions.43 However, significant differences have been

CA in A, n is the number of electrons, and ro is the radius of ultra-microelectrode. Using CV results, the reduction peak potential versus the log of the sweep rate was first plotted to a give a straight line. The α was calculated by plotting Ep vs. ln(ν), eq 3, which gives a straight line with a slope equals to RT/(2αF). This was then used in eq 2 to get a value for CoDo1/2. CA experiments were then used to determine a value for CoDo, eq 4. This provides two equations in terms of Co and Do. These two equations were solved simultaneously to determine both Co and Do. The numerical values of the diffusion coefficient and solubility of O2 at different temperatures are listed in Table 3. From Table 3, it can be seen that the diffusion coefficients and CA steady state currents are increasing with increasing temperature while the solubility of O2 is decreasing. Furthermore, Table 3 illustrates that [MO1,1O2]+-based IL has the lowest solubility of O2 at 35 and 45 °C, although it is the highest at 25 °C, whereas [N112,1O2]+-based IL has the highest diffusion coefficients followed by [S222]+- and [MO1,1O2]+-based ILs, respectively. In contrast, [S222]+based IL has the highest steady current in comparison to [N112,1O2]+- and [MO1,1O2]+-based ILs. 10551



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Figure 5. CA of (a) [MO1,1O2][N(Tf)2], (b) [N112,1O2][N(Tf)2], and (c) [S222][N(Tf)2] at temperatures of 25, 35, and 45 °C.

Table 3. Diffusion Coefficient (Do) × 1010 (m2/s), Solubility of O2 (C) (mM), CA Steady-State Current (nA), Charge Transfer Coefficient α at 25, 35, and 45 °C

The presence of a solute such as O2 in an IL can modify the mass transport properties. The reason that N2 does not cause a measurable change in the D value is probably due to its low solubility in the IL.45 Nevertheless, in principle, N2 also could alter the mass transport properties.26 As mentioned, the increase in the diffusion coefficient is observed with increasing temperature, owing, at least in part, to the reduction in viscosity of the solvent media as described by Stokes−Einstein relationship, eq 5. This is in agreement with Huang et al. (2009).9

temp. 25 °C Do C α CA current Do C α CA current Do C α CA current

35 °C

[MO1,1O2][N(Tf)2] 0.1 0.9 31.4 9.7 0.41 0.40 0.4 1.9 [N112,1O2][N(Tf)2] 0.9 2.4 14.9 11.2 0.64 0.49 2.9 5.6 [S222][N(Tf)2] 0.6 1.9 24.6 16.2 0.44 0.44 3.1 6.4

45 °C 1.9 7.4 0.51 2.9

D= 3.1 10.6 0.53 7.0

KBT 6πηr

(5) −23

where KB is the Boltzmann constant (1.38 × 10 m kg s−2 K−1), T is the temperature in K, and r is the hydrodynamic radius of the diffusing species, assuming the molecule is spherical.9,46 It has been reported by several studies that the diffusion coefficients of oxidized and reduced forms of a redox couple can be very different in ILs while in molecular solvents the differences are small.12,26 The diffusion coefficient values of O2 in [MO1,1O2]+ and [N112,1O2][N(Tf)2] are in agreement with that reported by Evans and co-workers, where they stated that, in the two

2.3 15.2 0.38 7.5

reported for the two species involved, O2 and O2•−, in terms of reversibility and diffusion coefficients.44 10552

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Figure 6. Change of superoxide absorbance peak with time in (a) [MO1,1O2][N(Tf)2], (b) [N112,1O2][N(Tf)2], and (c) [S222][N(Tf)2].

that was observed by Islam et al. (2009), with two absorption bands, can be attributed to the deprotonated imidazolium cation or other products.4 It is noteworthy that the long-term stability result of O2•− in [S222][N(Tf)2] provided an opponent indication to the CV of electrochemical generation of O2•− (short-term stability) as it indicated a stable generated O2•−. Clearly, this supports our hypothesis in terms of the need and usefulness of long-term stability and its high importance for industrial applications. The pseudo-first-order rate constants of O 2 •− in [MO1,1O2][N(Tf)2] and [N112,1O2][N(Tf)2] were estimated to be 0.672 × 10−5 s−1 and 0.419 × 10−5 s−1, respectively. These results support the CV results of O2 electrochemical reduction. These two ILs show the best and highest stability of O2•− in comparison to other reported ILs based on piperidinium, pyrrolidinium, imidazolium, and phosphonium cations.4,5,8,10,13 However, these values are much lower, by 2 orders of magnitude, than that reported in DMSO solution with the presence of 1-n-butyl-2,3-dimethylimidazolium tetrafluoroborate, which was estimated to be 2.5 × 10−3 s−1.4 Nevertheless, the determined rate constants are in accordance with those reported by AlNashef et al. (2010) which were 2.1 × 10−5 s−1 and 1.7 × 10−5 s−1 in [BMIm][PF6] and in 1-ethyl-3-methylimidazolium ethylsulfate, respectively,13 as well as those in [MOPMPip][N(Tf)2], [HMPyrr][N(Tf)2], and [P14,666][N(Tf)2], which were estimated to be 4.8 × 10−5 s−1, 1.9 × 10−5 s−1, and 16.2 × 10−5 s−1, respectively.10 Furthermore, it can clearly be observed that the rate constants of [MO1,1O2][N(Tf)2] and [N112,1O2][N(Tf)2] are lower

studied quaternary ammonium ILs, O2 diffusion was more rapid than that in the pyrrolidinium salt, and this is again due to its lower viscosity in comparison to that of [N6222][N(Tf)2] 95 mPa·s vs 220 mPa·s at 293 K.12 The obtained diffusion coefficient values were also in good agreement with those published previously in various ILs, since they have the same order of magnitude.9,12 3.4. Chemical Generation and Long-Term Stability of O2•−. With the aim of avoiding the possibility of mass controlling process while investigating the long-term stability, KO2 was dissolved in DMSO first, and then the tested IL was added to DMSO containing generated O2•−. Afterward, any consumption in O2•− generated can be ascribed to the reaction of O2•− with IL or impurities that could not be removed, in the IL.10 Figure 6 shows the effect of time on the peak height of O2•− in the DMSO/IL solutions. A slight decrease in the absorbance peak of O2•− with time can be observed in [MO1,1O2][N(Tf)2] and in [N112,1O2][N(Tf)2], Figure 6a and b, which indicates that O2•− is very stable in these ILs and that they can be used as good media for the generation of a stable O2•−. In contrast, the consumption rate of O2•− in [S222][N(Tf)2] is high, Figure 6c. However, there is no significant development of any new band in the 150−450 nm range of wavelength in the UV−visible spectra in this IL similar to that reported in the literature for the reaction of O2 •− and 1-n-butyl-2,3dimethylimidazolium tetrafluoroborate.4 Marcinek et al. (2001) detected two absorption bands [λmax 320 (main) and 250 nm (weak)] for the neutral radical of imidazolium ring generated by pulse radiolysis. Hence, the UV−visible spectrum 10553



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O2•− with chlorine atoms in DCP. It was observed that the peak of the CHC became smaller with the addition of KO2. This confirmed the destruction of CHC by the O2•− that was generated by dissolving KO2. This is in accordance with that reported by Kalu and White (1991) for the destruction of hexachlorobenzene (HCB) in DMSO, DMF, and DMSO/ DMF mixture by the electrochemically generated O2•−.48 This destruction process is considered as an in situ reaction; when the O2•− is generated, it reacts with CHCs as O2•− is a free radical anion.19 The result obtained is in agreement with Furman et al. (2010) who reported that 98% of HCB was degraded over 30 s in the crown ether-KO2-DMSO system and HCB degradation percentage decreased with increasing water concentrations. The presence of 5% and 10% of water in the DMSO systems lowered the reactivity of superoxide, resulting in 62% and 31% HCB degradation, respectively, after 30 s. In systems containing 15%, 20% and 25% of water, HCB degradation after 5 min was 90%, 40% and 15%, respectively.49 This is also in accordance with earlier reported work, where the chemically generated O2•− was utilized by dissolving KO2 for the destruction of chlorobenzenes in 1-(3-methoxypropyl)-1-methylpiperidinium and 1-hexyl-1-methyl-pyrrolidinium cations, structured with bis(trifluoromethylsulfonyl)imide anion and the destruction of chloroethanes in 1-butyl-1-methylpyrrolidinium trifluoroacetate50 under ambient conditions.10 However, work is in progress to extend this method for different types of toxic chlorinated hydrocarbons in various ILs media.

than those of the above-mentioned ILs, which indicates that O2•− is more stable in them. Results show that only 2.6% and 2.8% of the initial O2•− in [MO1,1O2][N(Tf)2] and [N112,1O2][N(Tf)2] were consumed after two hours of reaction time. The consumption rate of O2•− in [MO1,1O2][N(Tf)2] was comparable to that in [N112,1O2][N(Tf)2] with 0.651 × 10−3 mM/min and 0.698 × 10−3 mM/min, respectively. The differences in consumption rate can be attributed to the difference in type and quantity of impurities that could not be removed from the ILs. It was found that 92.5% of the generated O2•− in [S222][N(Tf)2] was consumed in 2 h. On the other hand, it was reported that only 27.32% and 13.28% of the initial O2•− in [MOPMPip][N(Tf)2] and [HMPyrr][N(Tf)2] were consumed after 2 h of reaction time and 70.68% of generated O2•− was consumed in [P14,666][N(Tf)2].10 The results of [N112,1O2][N(Tf)2] are in agreement with that reported by Laoire et al. (2010), where they studied the influence of solvents on O2 reduction reaction in nonaqueous electrolytes for elucidating the mechanism of the O2 electrode processes in the rechargeable Li-air battery. Using either tetrabutylammonium hexafluorophosphate [TBAmm][PF6] or lithium hexafluorophosphate [Li][PF6] electrolyte solutions in four different solvents, namely, DMSO, AcN, dimethoxyethane, and tetraethylene glycol dimethyl ether. They found that O2 reduction in TBAmm salt solutions having superior stability of O2•− in the presence of [TBAmm]+ in various solvents involving the formation of the [TBAmm]+... O2•− complex.47 3.5. Destruction of 2,4-Dichlorophenol. To test the possibility of using the generated O2•− in relevant applications, it was investigated the application of O2•− in the destruction of chlorinated phenols as this is one of the most important issues related to the environment. In addition, this subject was investigated by many other groups but in volatile aprotic solvents. [MO1,1O2][N(Tf)2] was selected as a medium for destruction of 2,4-dichlorophenol using O2•− generated by dissolving KO2 in the IL. Figure 7 shows the HPLC chromatograms of DCP in the selected IL before and after the addition of KO2. Remarkably, it was found that O2•− is capable to destroy DCP in [MO1,1O2][N(Tf)2] under ambient conditions with a percentage higher than 98%. The destruction can be ascribed to the nucleophilic substitution of

4. CONCLUSION Four ILs with [N(Tf)2]− anion were investigated as possible media for the generation of O2•−. The studied ILs containing the corresponding cations [HPy] + , [MO1,1O2] + , [N112,1O2]+, and [S222]+. All four ILs were found to possess wide EWs. The [MO1,1O2]+ based IL has the widest EW followed by the [N112,1O2]+-, [HPy]+-, and [S222]+-based ILs, respectively. CV and CA techniques were utilized to study the electrochemical generation of O2•−. It was found that the generated O2•− in the [HPy]+-based IL was not stable. On the other hand, it was found that [MO1,1O2]+-, [N112,1O2]+-, and [S222]+-based ILs can be used as successful media to generate a stable O2•−. CV showed that the superoxide generation is an irreversible process. Both CV using GC macroelectrode, and CA, using CF ultra-microelectrode, were used simultaneously to determine the solubility and diffusion coefficient of O2 in the studied ILs under different temperatures. It was found the diffusion coefficients followed the trend [N112,1O2]+- > [S222]+- > [MO1,1O2]+-based ILs and the solubility followed the sequence [MO1,1O2] + - < [N112,1O2]+- < [S222]+-based ILs at 35 and 45 °C. The results showed that the structure of the cation of the IL has an important role in determining both the solubility and diffusion coefficient of O2. Long-term stability of the generated O2•− was then investigated in [MO1,1O2]+-, [N112,1O2]+-, and [S222]+-based ILs by chemical generation of O2•−, in DMSO, in the presence of the corresponding IL. A UV−visible spectrophotometer was used at an absorbance range of 190− 400 nm to determine the stability of O2•−. It was found that the IL containing [MO1,1O2]+ gives the highest stability of O2•−, followed by [N112,1O2]+, while it was found that O2•− is unstable in [S222]+ based IL. The chemically generated O2•− by dissolving KO2 was used for the destruction of DCP in [MO1,1O2][N(Tf)2] under ambient conditions. The destruc-

Figure 7. HPLC chromatograms of DCP in [MO1,1O2][N(Tf)2] (a) before KO2 addition (b) after KO2 addition. 10554



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(12) Evans, R. G.; Klymenko, O. V.; Saddoughi, S. A.; Hardacre, C.; Compton, R. G. Electroreduction of oxygen in a series of room temperature ionic liquids composed of group 15-centered cations and anions. J. Phys. Chem. B 2004, 108, 7878. (13) AlNashef, I. M.; Hashim, M. A.; Mjalli, F. S.; Ali, M. Q.; Hayyan, M. A novel method for the synthesis of 2-imidazolones. Tetrahedron Lett. 2010, 51, 1976. (14) Rogers, E. I.; Huang, X. J.; Dickinson, E. J. F.; Hardacre, C.; Compton, R. G. Investigating the mechanism and electrode kinetics of the oxygen| superoxide (O2|O2•−) couple in various room-temperature ionic liquids at gold and platinum electrodes in the temperature range 298−318 K. J. Phys. Chem. C 2009, 113, 17811. (15) Barnes, A. S.; Rogers, E. I.; Streeter, I.; Aldous, L.; Hardacre, C.; Wildgoose, G. G.; Compton, R. G. Unusual voltammetry of the reduction of O2 in [C4dmim][N(Tf)2] reveals a strong interaction of O2•− with the [C4dmim]+ cation. J. Phys. Chem. C 2008, 112, 13709. (16) Navarro-Suárez, A. M.; Hidalgo-Acosta, J. C.; Fadini, L.; Feliu, J. M.; Suárez-Herrera, M. F. Electrochemical oxidation of hydrogen on basal plane platinum electrodes in imidazolium ionic liquids. J. Phys. Chem. C 2011, 115, 11147. (17) Noradoun, C.; Engelmann, M. D.; McLaughlin, M.; Hutcheson, R.; Breen, K.; Paszczynski, A.; Cheng, I. F. Destruction of chlorinated phenols by dioxygen activation under aqueous room temperature and pressure conditions. Ind. Eng. Chem. Res. 2003, 42, 5024. (18) Afanas’ev, I. B. Superoxide Ion: Chemistry and Biological Implications; CRC Press: Boca Raton, 1989; Vol. 1. (19) Sawyer, D. T. Oxygen Chemistry; Oxford University Press, New York/Oxford, 1991. (20) Sugimoto, H.; Matsumoto, S.; Sawyer, D. T. Degradation and dehalogenation of polychlorobiphenyls and halogenated aromatic molecules by superoxide ion and by electrolytic reduction. Environ. Sci. Technol. 1988, 22, 1182. (21) Hapiot, P.; Lagrost, C. Electrochemical reactivity in roomtemperature ionic liquids. Chem. Rev. 2008, 108, 2238. (22) O’mahony, A. M.; Silvester, D. S.; Aldous, L.; Hardacre, C.; Compton, R. G. Effect of water on the electrochemical window and potential limits of room-temperature ionic liquids. J. Chem. Eng. Data 2008, 53, 2884. (23) Kato, H.; Nishikawa, K.; Koga, Y. Relative hydrophobicity and hydrophilicity of some “ionic liquid” anions determined by the 1propanol probing methodology: A differential thermodynamic approach. J. Phys. Chem. B 2008, 112, 2655. (24) Bard, A. J.; Faulkner, L. R. Electrochemical Methods: Fundamentals and Applications, 2 ed.; Wiley: New York, 2001. (25) Rollie, M. E.; Patonay, G.; Warner, I. M. Deoxygenation of solutions and its analytical applications. Ind. Eng. Chem. Res. 1987, 26, 1. (26) Zhao, C.; Bond, A. M.; Compton, R. G.; O’Mahony, A. M.; Rogers, E. I. Modification and implications of changes in electrochemical responses encountered when undertaking deoxygenation in ionic liquids. Anal. Chem. 2010, 82, 3856. (27) Buzzeo, M. C.; Hardacre, C.; Compton, R. G. Use of room temperature ionic liquids in gas sensor design. Anal. Chem. 2004, 76, 4583. (28) MacFarlane, D. R.; Meakin, P.; Sun, J.; Amini, N.; Forsyth, M. Pyrrolidinium imides: A new family of molten salts and conductive plastic crystal phases. J. Phys. Chem. B 1999, 103, 4164. (29) Suarez, P. A. Z.; Consorti, C. S.; Souza, R. F.; Dupont, J.; Gonçalves, R. S. Electrochemical behavior of vitreous glass carbon and platinum electrodes in the ionic liquid 1-n-butyl-3-methylimidazolium trifluoroacetate. J. Braz. Chem. Soc. 2002, 13, 106. (30) Ignat’ev, N. V.; Welz-Biermann, U.; Kucheryna, A.; Bissky, G.; Willner, H. J. Fluorine Chem. 2005, 126, 1150. (31) Matsumoto, H.; Sakaebe, H.; Tatsumi, K.; Kikuta, M.; Ishiko, E.; Kono, M. Fast cycling of Li/LiCoO2 cell with low-viscosity ionic liquids based on bis(fluorosulfonyl)imide [FSI]. J. Power Sources 2006, 160, 1308.

tion percentage was higher than 98%. This shows that the generated O2•− is stable enough to be used in various relevant applications.

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ASSOCIATED CONTENT

S Supporting Information *

The formulas, molecular weights and melting points of ILs, HPLC specifications and analysis conditions, and the Figures of EWs. This information is available free of charge via the Internet at http://pubs.acs.org/.

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AUTHOR INFORMATION

Corresponding Author

*Phone: + 968 2414 2558. Fax: + 968 2414 1354. E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The authors would like to express their thanks to University of Malaya HIR-MOHE (D000003-16001), FRGS, Centre for Ionic Liquids (UMCiL), and the National Plan for Science, Technology, and Innovation at King Saud University (10ENV1010-02), and Sultan Qaboos University for their support to this research.

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REFERENCES

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Generation of Superoxide Ion in Pyridinium, Morpholinium

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