Surface Tension of 1-Ethyl-3-methylimidazolium ... - ACS Publications

Article pubs.acs.org/jced

Surface Tension of 1‑Ethyl-3-methylimidazolium Ethyl Sulfate or 1‑Butyl-3-methylimidazolium Hexafluorophosphate with Argon and Carbon Dioxide Guillermo Reyes,† Marcela Cartes,† Carlos Rey-Castro,† Hugo Segura,*,† and Andrés Mejía*,† †

Departamento de Ingeniería Química, Universidad de Concepción, POB 160-C, Correo 3, Concepción, Chile Departament de Química, Universitat de Lleida, Rovira Roure 191, 25198 Lleida, Catalonia, Spain

†

S Supporting Information *

ABSTRACT: Surface tensions of two ionic liquids (IL): 1-ethyl-3methylimidazolium ethyl sulfate and 1-butyl-3-methylimidazolium hexafluorophosphate in pressurized atmospheres of argon and carbon dioxide have been measured over the temperature range (303 to 366) K and over the pressure range (0.1 to 15) MPa for the case of argon atmosphere and (0.1 to 5) MPa for the case of carbon dioxide atmosphere by using a pendant drop tensiometer. Based on the experimental measurements, the isothermal surface tension of all IL−gas systems studied decreases as the pressure increases, evidencing a gas adsorption at the IL interface. Isobaric surface tension of an IL−gas does not show a general pattern as the temperature increases. In order to verify the isothermal surface behavior, the relative Gibbs adsorption isotherms have been calculated from the surface tension data by using the theoretical Guggenheim model, corroborating the gas adsorption processes at the IL interface. Comparing the relative Gibbs adsorption isotherms, it is possible to conclude that the ILs studied have the capability to adsorb more carbon dioxide than argon. This fact provides relevant information to use the IL as a capturing agent for carbon dioxide and the use of argon to store pure ILs.

■

INTRODUCTION Ionic liquids (ILs) are part of the so-called neoteric solvents due to their tunable properties, relatively low environmental impact,1−5 and reduced volatility, which have led to a wide range of applications in biology,6,7 electrochemistry,8,9 organic chemistry, separation, extraction, and catalysis1,5−10 fields. As a rule of thumb, it is well-established that the IL cation is mostly responsible for the physical properties such as density and viscosity whereas the IL anion is predominantly responsible for the chemical properties such as reactivity and selectivity among others.5,6,10 The large versatility resulting from the possible combinations of cations and anions gives ILs their tunable behavior and transforms them into design substances. Besides, ILs exhibit negligible vapor pressure and low toxicity, which make them environmentally friendly, and so their use has been implemented within so-called green chemistry.3−5 Nowadays, mixtures of ILs with gases have gained considerable interest since ILs can be used as capturing agents for greenhouse gases such as methane (CH4), water vapor (H2O), and carbon dioxide (CO2).11−13 For the case of the ILs studied here, namely, 1-butyl-3methylimidazolium hexafluorophosphate or [bmim][PF6] and 1-ethyl-3-methylimidazolium ethyl sulfate or [emim][C2SO4], previous works have been focused on determining vapor−liquid equilibria, diffusion, and solubility of some gases. For the case of [bmim][PF6], Morgan et al.14 reported the vapor−liquid © 2013 American Chemical Society

equilibria and diffusion coefficients of this IL with CO2, oxygen (O2), nitrogen (N2), and a series of hydrocarbons from methane (CH4) to butane (C4H10). Other authors have reported the solubility of gases such as CO2 and argon (Ar).15−18 For the case of [emim][C2SO4], Blanchard et al.13 presented the high-pressure phase behavior of CO2 with this and other ionic liquids in the temperature range (313 to 333) K, while Jacquemin et al.19 reported solubilities of hydrogen and CO2 over the temperature range (283 to 343) K. For the mixture [emim][C2SO4] + Ar, no previous work has been reported previously to the best of our knowledge. Despite the importance of diffusion and solubility of gases in [bmim][PF6] and [emim][C2SO4], the surface tensions of such systems have not been thoroughly studied so far. The surface tension of pure ILs and gas pressurized ILs plays key role in science and technology as it is directly involved in solubility, miscibility and mass transfer processes.20,21 In spite of the importance of this thermophysical property, experimental surface tension data reported in the literature does usually refer to the pure IL systems only. For instance, surface tensions of pure [bmim][PF6]22−25 have been studied in the temperature range (288.14 to 393) K using techniques such as Received: December 18, 2012 Accepted: April 1, 2013 Published: April 12, 2013 1203

dx.doi.org/10.1021/je3013414 | J. Chem. Eng. Data 2013, 58, 1203−1211

Journal of Chemical & Engineering Data

Article

Table 1. Refractive Index (nD) at the Na D Line, Water Content, and Densities (ρ) of Pure Ionic Liquids as a Function of Temperature (T) at Pressure p = 0.10 MPaa nD, T/K = 298.15

ρ/g·cm−3

water content

exp.

lit.

ppm H2O

T/K

exp.

lit.c

[bmim][PF6]

1.4106

1.4094

2417.13

[emim][C2SO4]

1.4805

1.4794

1083.3

303.15 313.15 323.15 333.15 343.15 353.15 363.15 303.15 313.15 323.15 333.15 343.15 353.15 363.15

1.3613 1.353 1.3448 1.3367 1.3285 1.3205 1.3124 1.2337 1.2269 1.2202 1.2136 1.207 1.2004 1.1939

1.3626 1.3565 1.3473 1.3359 1.3285 1.3225 1.3126 1.2388 1.2319 1.2251 1.2183 1.2116 1.2050

component

b

Standard uncertainties u are u(T) = ± 0.02 K, u(nD) = ± 10−5, u(ρ) = 5·10−6 g·cm−3. bReference 39 for [bmim][PF6] and ref 40 for [emim][C2SO4]. cReference 41 for [bmim][PF6] and ref 42 for [emim][C2SO4].

a

mass densities (ρ) in the temperature range (303 to 363) K, for the studied ILs. Table 1 also includes the corresponding reference values reported in the literature39−42 for the measured pure fluid properties. The presence of relatively low amounts of water has been shown not to be systematically correlated with deviations in the measured surface tension of ionic liquids.32 Apparatus and Procedure. Density, Refractive Indexes, and Water Content Measurements. The mass density (ρ) of the pure fluids was measured over the temperature range (303 to 363) K using a DMA 5000 densimeter (Anton Paar GmbH, Austria) with an accuracy of 5·10−6 g·cm−3. In this device, the temperature was maintained constant to within ± 0.01 K. The refractive indexes (nD) of pure liquids were measured at 298.15 K using a Multiscale Automatic Refractometer RFM 81 (Bellingham and Stanley Ltd., England) with uncertainties of ± 10−5. The temperature was kept to within ± 0.01 K by means of a thermostatic bath (Haake DC3, Germany). Specific details related to the experimental procedure have been reported previously.43 The water content present in pure ionic liquids was determined using a Coulometric Karl Fischer titrator 831 (Metrohm AG, Switzerland) with an accuracy of 0.3 % mass. The water content determination is based on a Coulomb metric titration using a CombiCoulomat frit solution (Merck). Surface Tension Cell. The surface tension (γ) of an IL surrounded by a gas (CO2 or Ar) was measured by using a pendant drop tensiometer (Temco, model IFT-10). Specific details about the quoted tensiometer and its technical specifications have been reported previously.43 The main modifications implemented for the purpose of this work are the needle used (i.e., stainless steel needle of 1.4 mm i.d. and 2.45 mm o.d.) and the gas supply system. Particularly, gases (CO2 or Ar) were supplied from commercial cylinders appropriately connected to the bottom orifice of tensiometer cell by means of a high-pressure regulator. The experimental determination of γ has been carried out using the standard procedure comprehensively described by several authors,44−49 which was implemented in our experimental rig43 with the modifications described below: The cell is heated to the required experimental temperature, and then it is slightly pressurized with the gas. After degasification in an ultrasonic bath, the pure liquid is pumped

maximum bubble pressure, capillary rise, and du Noüy ring. For the case of [emim][C2SO4]26−30 data were reported in the range (278.15 to 453) K using maximum bubble pressure, pendant drop, and du Noüy ring techniques. Some recent reviews in this topic reveal that most previous works regarding the study of binary systems have focused on the liquid−liquid and/or vapor−liquid surface tensions of ILs in mixtures with, e.g., water, alcohols and other organic solvents.31,32 For the case of 1-butyl-3-methylimidazolium hexafluorophosphate, for instance, Santos and Baldelli33 studied the effect of benzene concentrations on the cation orientation and liquid−liquid tensions at atmospheric conditions, Huddleston et al.34 and Freire et al.35 reported surface tensions in mixtures with water. Łuczak et al.36 investigated the critical micelle concentration of [bmim][PF6] in aqueous solutions, whereas Modaressi et al.37 reported information on the effect of pH on the surface tension in aqueous solutions. For the case of 1-ethyl-3-methylimidazolium ethyl sulfate, Fröba et al.38 studied the influence of Coulomb interactions on the viscosity and surface tension of [emim][C2SO4] in mixtures with water and ethanol. However, for IL−gas systems there is still a lack of reliable theoretical predictions, including molecular simulations and experimental information regarding gas−liquid surface tension. Due to the relevance of these systems, the present work is devoted to the measurement of novel surface tension data for [bmim][PF6] pressurized by Ar and CO2 and [emim][C2SO4] pressurized by CO2 over the temperature range (303 to 363) K and over the pressure range (0.1 to 15) MPa for the case of argon and (0.1 to 5) MPa for the case of carbon dioxide by using a pendant drop tensiometer.

■

EXPERIMENTAL SECTION Chemicals. 1-Ethyl-3-methylimidazolium ethyl sulfate (> 98 % by RMN1H) and 1-butyl-3-methylimidazolium hexafluorophosphate (> 98 % by RMN1H) were purchased from Green Solutions Chemicals, S.L. (Spain), and were used without further purification. Ultra high purity argon and carbon dioxide were purchased from Linde S.A. (Chile) with a certified purity greater than 0.99995 in mole fraction (see Table S.1 in Supporting Information for impurities). Table 1 reports the refractive indexes (nD), water content at 298.15 K, and liquid 1204



Article

to the cell which is then pressurized up to the experimental condition. Once the drop is formed at the tip of the needle and the required experimental temperature and pressure are reached, the main geometric dimensions of the drop (i.e., the equatorial diameter (de) and the horizontal diameter (ds)) are recorded. Experimentally, it has been observed that the drop attains equilibrium in, approximately, to (2 to 4) h. After this equilibration procedure, the main geometric dimensions of the drop (de and ds) are recorded (at least during 6 h) in order to check the stability of its shape in a cyclic procedure that requires about 10 h for each measure. Considering de and ds values, the density of the liquid at the tip of the needle (ρL) and the density of the gas that fills the chamber (ρG)both evaluated at the experimental temperature and pressureγ is obtained from the following expression: γ = (ρL − ρG )gcde2f (de , ds)

Figure 1. Time evolution of experimental surface tension (γ) measured for [bmim][PF6] + Ar at T/K = 343.15, P/MPa = 1. ○, experimental data;  smoothed by eq 2 with parameters shown in Table S.1.

(1)

In eq 1, gc is the local gravitational constant (≈ 981 cm·s−2) and f(de,ds) is a function related to the silhouette of the drop, whose value is determined from numerical tables.44,50,51 It is important to establish that, in eq 1, ρL corresponds to the density of liquid saturated with the gas while ρG is the density of the gas saturated with the liquid. Consequently, besides of temperature and pressure, the indicated ρL and ρG values depend on solubility. In this work, however, ρL and ρG have been reasonably taken as the pure liquid and gas densities, respectively. Particularly, ρG for argon and carbon dioxide were taken from the NIST Chemistry WebBook,52 whereas ρL for [bmim][PF6] and [emim][C2SO4] ionic liquids were taken from Ionic Liquids Database (IL Thermo).53 Although the approximation of pure densities is well-justified in the case of the gas phase (due to the negligible vapor pressure of ILs), the density of the liquid phasewhose information is either not always available or dif fers considerably for the range of conditions investigated hererequires further examination. According to Dittmar et al.54 and Georgiadis et al.,48,49 the dependence of the phase density on concentration may be usually assumed negligible and does not have a high impact on determined γ values. In addition, in any of the experimental determinations considered here, the reported difference in density between pure [bmim][PF6] and the CO2 saturated liquid phase at 313 K and 15 MPa is less than 2 %.13

Figure 2 shows results for [bmim][PF6] + Ar system at 343.15 K over the pressure range (0.5 to 15) MPa (see Table S.2 in Supporting Information for the corresponding a and b parameters at isothermal (303 to 363) K conditions together with the correlation statistics). Figure 3 shows the surface tension for the [bmim][PF6] + CO2 system at 343.15 K and over a pressure range (1 to 5) MPa (see Table S.3 in Supporting Information for the corresponding a and b parameters). It is possible to observe that the diffusional process is longer for CO2 systems than for the case of Ar systems. This behavior can be attributed to the adsorption of CO2 and the solubility of Ar or CO2 in the [bmim][PF6]. In fact, according to the reported experimental data (see Anthony et al.55 and references therein), CO2 is considerably more soluble than Ar in [bmim][PF6]. Therefore, it is expected that diffusional process would be faster in Ar than in CO2. Finally, Figure 4 shows the surface tension for [emim][C2SO4] + CO2 system at 363.15 K and over a pressure range (0.5 to 5) MPa (see Table S.4 in Supporting Information for the corresponding a and b parameters). In this case, the temperature and pressure effects are similar than for [bmim][PF6] + Ar or CO2. However, it is possible to observe that [emim][C2SO4] + CO2 exhibits a slower diffusional dynamics than [bmim][PF6] + CO2. In fact, the CO2 diffusion process takes several hours to reach a steady state at [emim][C2SO4] interface, whereas for [bmim][PF6] it only takes minutes. This difference in dynamic behavior could be related to the difference in the solubility of CO2 and in the viscosity of the liquid phase, which, in turn, could be affected by the relative amount of impurities in each IL. The mass transfer phenomena of the gases within the studied ILs can be approximated by the one dimension Fickian mass transfer model,56 where the radial variation of the gas concentration (Cgas) with time (t) in the IL drop is governed by the following partial differential equation:

■

RESULTS AND DISCUSSION Diffusive Process and Surface Tension. When the gas is in contact with the ionic liquid drop, the gas diffusion process through the liquid is accompanied by a decay of the surface tension. Figure 1 shows the surface tension variation observed for [bmim][PF6] + Ar system at 343.15 K and 1 MPa. From this figure, it is possible to observe that the interfacial process presents a decay behavior, which can be smoothed by the following exponential function: γ(t ) = γ∞ + a ·e(−bt )

(2)

In eq 2, a and b are empirical parameters, which can be found by applying a nonlinear least-squares technique, and γ∞ represents the value of surface tension at the steady state, from which γ is determined according to:

∂Cgas ∂t

t →∞

t →∞

) = γ∞

(4)

In eq 4 Dgas‑IL is the diffusion coefficient of gas in the IL evaluated at the experimental temperature and pressure.

(−bt )

γ = lim γ(t ) = lim (γ∞ + a ·e

⎡ 1 ∂ ⎛ ∂Cgas ⎞⎤ = Dgas−IL ⎢ 2 ⎜r 2 ⎟⎥ ⎢⎣ r ∂r ⎝ ∂r ⎠⎥⎦

(3) 1205



Article

Figure 2. Time evolution of experimental surface tension (γ) measured for [bmim][PF6] + Ar at T/K = 343.15. ●, P/MPa = 0.5; ○, P/MPa = 1; ▼, P/MPa = 3; △, P/MPa = 5; □, P/MPa = 10; ■, P/MPa = 15. , smoothed by eq 2 with parameters shown in Table S.2.

Figure 4. Time evolution of experimental surface tensions (γ) measured for [emim][C2SO4] + CO2 at T/K = 363.15. ●, P/MPa = 0.5; ○, P/MPa = 1; ▼, P/MPa = 2; △, P/MPa = 3; ■, P/MPa = 5. , smoothed by eq 2 with parameters shown in Table S.4.

Figure 3. Time evolution of experimental surface tensions (γ) measured for [bmim][PF6] + CO2 at T/K = 343.15. ○, P/MPa = 1; ▼, P/MPa = 3; △, P/MPa = 5. , smoothed by eq 2 with parameters shown in Table S.3.

In eq 5, Rdrop denotes the radius of the drop, and C*gas is the saturation gas concentration at the IL drop. The analytical solution of eq 4 is given by:57

Considering that the IL drop can be modeled as a sphere, the mass transfer is governed by the following initial and boundary conditions: Cgas = 0

at t = 0

and r ≥ 0

Cgas = finite

at r = 0

* Cgas = Cgas

at r = R drop

and t > 0 and t > 0

Cgas 2 =1+ * R* Cgas

(5a)

n =∞

∑ n=1

sin(nπR *)

(5b)

1 cos(nπ ) ·exp( −n2π 2D*gas−IL t ) · nπ (6)

where R* = r/Rdrop is the reduced radial variable and D* = D/ Rdrop2 is the reduced diffusion coefficient. Considering that the

(5c) 1206



Article

diffusion coefficient of CO2 + [bmim][PF6]58 is 2.1·10−10 m2·s−1, and the drop radius is Rdrop = 1.26 mm (average radius of [bmim][PF6] pendant drops in CO2 at 343.15 K), Figure 5

Table 2. Experimental Equilibrium Surface Tension (γ) for [bmim][PF6] + Ara T/K

p/MPa

γ/mN·m−1

303.15

0.5 1 3 5 10 15 0.5 1 3 5 10 15 0.5 1 3 5 10 15 0.5 1 3 5 10 15 0.5 1 3 5 10 15

41.50 40.93 39.41 38.11 34.40 32.09 40.20 39.81 37.69 35.56 32.96 31.47 41.44 40.19 38.81 38.07 35.03 32.16 40.22 40.06 38.79 37.43 31.93 31.32 38.00 37.14 35.74 35.12 32.51 31.56

313.15

323.15

Figure 5. Concentration distribution of CO2 in [bmim][PF6] spherical drop of Rdrop = 1.26 mm at t = 10 s, t = 300 s, t = 600 s, t = 1500 s, t = 3000 s.

343.15

shows the time dependence of the concentration profiles for spatial concentration distributions of CO2 in the IL spherical drop, at different times between (10 and 3000) s. From these results, it is observed that CO2 concentration reaches equilibrium at 3000 s approximately. This result agrees well with the data presented in Figure 3, where stationary surface tension values were obtained at around 3000 s (50 min). The quoted agreement is particularly remarkable, as the simplified diffusion model does not account for the dynamics of the adsorption on the interface, the effect of the liquid phase composition on the diffusion coefficient of CO2, or the effect of impurities. Other associated effects related to the dynamic behavior observed could be related to the dehydration of the IL drop under CO2 atmosphere; however this is not consistent with two facts: (a) careful studies in literature, see for instance Feire et al.,35 indicate that [bmim][PF6] tension decreases with water traces (see Figure 4a in Freire et al. work), so the variation in the dynamic tension should follow the opposite trend if it was really controlled by the loss of water; (b) [emim][C2SO4] has a lower content in water than [bmim][PF6] and, yet, the kinetics of the tension variation is actually slower. Equilibrium Surface Tension. Tables 2 to 4 (see also Figures S1 to S3 in Supporting Information) summarize the equilibrium surface tensions for [bmim][PF6] + Ar, [bmim][PF 6 ] + CO 2 , and [emim][C 2SO 4 ] + CO 2 over the experimental temperature and pressure range considered in this work. The results of the mixtures with CO2 at higher pressures were not reported in this study due to the experimental inaccuracy derived from drop instability at low surface tensions. The values presented in Tables 2 to 4 correspond to the stationary values and have been determined from the procedure described in the previous section. From Tables 2 to 4 it is possible to conclude that, at isothermal conditions, the surface tension decreases as the pressure increases. According to these results, the surface tension of [bmim][PF6] + Ar decreases 10 mN·m−1 when the pressure increases from (0.5 to 15) MPa, whereas the tension of

363.15

Standard uncertainties are u(p) = ± 0.05 MPa, u(T) = ± 1.0 K, and u(γ) = ± 0.6 mN·m−1. a

Table 3. Experimental Equilibrium Surface Tension (γ) for [bmim][PF6] + CO2a T/K

p/MPa

γ/mN·m−1

303.15

1 3 5 1 3 5 1 3 5 1 3 5

35.19 26.59 18.79 35.19 26.15 21.47 35.31 28.42 22.33 32.09 27.54 24.99

313.15

323.15

343.15

a Standard uncertainties are u(p) = ± 0.05 MPa, u(T) = ± 1.0 K, and u(γ) = ± 0.6 mN·m−1.

[bmim][PF6] + CO2 decreases about 15 mN·m−1 when the pressure increases from (0.5 to 5) MPa. A qualitatively similar decrease is observed for [emim][C2SO4] + CO2, although in this case the surface tension is ca. (1 to 3) mN·m−1 higher than in [bmim][PF6] + CO2 at the same temperature and pressure. Figures 6 and 7 depict isobaric behavior for [bmim][PF6] + CO2 and [emim][C2SO4] + CO2, where a surface tension 1207



Article

decay is initially observed in the lowest pressure range as temperature increases. Then, as pressure increases, the observed trend experiments an inversion revealing an increase of surface tension with temperature. For mixtures containing CO2 the inversion occurs between (1 and 2) MPa, while for the mixture [bmim][PF6] + Ar no inversion was detected for pressures below 15 MPa. Gray et al.59 have mentioned some systems exhibiting a similar inversion behavior, which is explained by the authors in terms of the surface excess entropy, often associated to phase transitions (as in the case of nematic liquid crystals), or due to the presence of a strong adsorption regime at the interface. This behavior was also experimentally noted in other mixtures involving CO2. Georgadis et al.60 reported surface tension values for the mixture CO2 + n-decane, where the inversion occurs around 3 MPa. Dittmar et al.54 have studied the mixture CO2 + ethanol, where a similar behavior was observed around 2.5 MPa. For CO2 mixtures with ionic liquids, Hebach et al.61 studied CO2 + 1-ethyl-3-methylimidazolium 2-(2-ethoxyethoxy) ethylsulfate and CO2 + 1-butyl-3methylimidazolium 2-(2-methoxyethoxy) ethylsulfate, whereas Jaeger et al.62 measured CO2 + 1-butyl-3-methylimidazolium hexafluorophosphate and CO2 + 1-butylpyridinium tetrafluoroborate. In these mixtures, they observed an inversion behavior for pressures around (1 to 3) MPa. In addition to these experimental results, the effect of a high-pressure gas over a liquid surface tension has also been studied by molecular dynamics.63 Indeed, simulation results suggest that liquid− vapor interfaces exhibit a dramatic surface tension decrease with temperature by increasing the gas adsorption at the interface. Comparing the results reported in Tables 2 to 4, it is possible to observe that surface tension values in the presence of Ar are higher than for mixtures with CO2, which evidence that CO2 is strongly adsorbed at the liquid surface of ILs. In addition, it is possible to conclude that surface tensions of [emim][C2SO4] + CO2 are higher than [bmim][PF6] + CO2, which can be explained in terms of solubility of CO2 in the corresponding ILs.13−19 The high solubility of CO2 in [bmim][PF6] compared to the solubility in [emim][C2SO4] can be attributed to the strong interaction between the PF6 anion molecule and CO2 as described by Cadena et al.64 Relative Gibbs Adsorption Isotherms. As it was mentioned in the previous section, the surface tensions of the measured mixtures decrease as they are pressurized with the gas. This behavior can be attributed to the adsorption affinity of the gas at the liquid surface. A quantitative description of this trend can be obtained by considering relative Gibbs adsorption isotherms, which have been calculated from the Guggenheim’s relationship applied to gas−liquid systems:65−67

Table 4. Experimental Equilibrium Surface Tension (γ) for [emim][C2SO4] + CO2a T/K

p/MPa

γ/mN·m−1

303.15

0.5 1 2 3 5 0.5 1 2 3 5 0.5 1 2 3 5

37.10 36.50 34.15 27.23 21.94 35.15 34.38 31.70 30.77 25.70 36.80 33.24 32.50 31.94 29.32

333.15

363.15

Standard uncertainties are u(p) = ± 0.05 MPa, u(T) = ± 1.0 K, and u(γ) = ± 0.6 mN·m−1. a

Figure 6. Isobaric surface tension (γ) for [bmim][PF6] + CO2 as a function of the temperature (T) at ●, P/MPa = 1; ○, P/MPa = 3; ▼, P/MPa = 5.

Γ12 = −

1 ⎛ dγ ⎞ ⎜ ⎟ v1 ⎝ dP ⎠T

(7)

In eq 7, Γ12 is the concentration of the adsorbed gas (1) in moles per unit of surface area of the IL (2), and v1 is the molar volume of the pure gas evaluated at the experimental temperature T, and pressure P (values taken from the NIST Chemistry WebBook52). In order to determine (dγ/dP)T from experimental data at each temperature, the experimental values of γ have been correlated by a second-order polynomial with P Figure 7. Isobaric surface tension (γ) for [emim][C2SO4] + CO2 as a function of the temperature (T) at ●, P/MPa = 1; ○, P/MPa = 2; ▼, P/MPa = 3; △, P/MPa = 5.

γ = a0 + a1P + a 2P 2

(8)

where a0, a1, and a2 correspond to the polynomial coefficients (see Table S.5 in Supporting Information for ai values, which 1208



Article

have been estimated by applying a least-squares technique with a minimum coefficient of determination, r2, of 0.95). Figures 8 to 10 show the estimated values of Γ12 as a function of pressure at different isothermal conditions for the three IL +

Figure 10. Relative Gibbs adsorption isotherms (Γ12) as a function of the pressure (P) for the system [emim][C2SO4] + CO2: , T/K = 303.15; ···, T/K = 333.15; ---, T/K = 363.15.

■

CONCLUSIONS Surface tensions and relative Gibbs adsorption isotherms for binary systems have been described for the systems [bmim][PF6] + Ar over the temperature range (303.15 to 363.15) K and pressures from (0.5 to 15) MPa, [bmim][PF6] + CO2 over the temperature range (303.15 to 343.15) K and pressures from (0.5 to 5) MPa, and for [emim][C2SO4] + CO2 over the temperature range (303.15 to 363.15) K and pressures from (0.5 to 5) MPa. Experimental results indicate that the evolution of the surface tension of the mixtures in the first minutes is a process controlled by diffusion within the drop. The time required to attain diffusion equilibrium is in excellent agreement with the time required to reach stationary values of surface tension at the experimental runs. In general, it has been observed that the surface tensions decrease as pressure increases. At isothermal conditions, the decay of the surface tension with increasing pressure evidence that the gas is adsorbed at the IL surface, a condition that is confirmed by the positive values of the relative Gibbs adsorption isotherms calculated from the surface tension data and the Guggenheim model. At the same pressure and temperature, the adsorption of CO2 was greater than Ar. In addition, it is observed that the net CO2 adsorption process is greater for [bmim][PF6] than for [emim][C2SO4].

Figure 8. Relative Gibbs adsorption isotherms (Γ12) as a function of the pressure (P) for the system [bmim][PF6] + Ar: , T/K = 303.15; ···, T/K = 323.15; ---, T/K = 363.15.

Figure 9. Relative Gibbs adsorption isotherms (Γ12) as a function of the pressure (P) for the system [bmim][PF6] + CO2: , T/K = 303.15; ···, T/K = 323.15; ---, T/K = 343.15.

■

ASSOCIATED CONTENT

S Supporting Information *

gas mixtures studied here. From these figures, it is possible to conclude that Γ12 is positive in the whole pressure range thus implying that the gas is always adsorbed into the IL phase. It can also be noticed that, in good agreement with previous results and discussions, Γ12 values are much higher for CO2 adsorption than for Ar adsorption. Particularly for CO2, Γ12 decreases as temperature increases, a condition that coincides with the behavior observed for simple gases dissolved in organic fluids.67 At high pressure, the adsorption process seems to be intense enough to cause an effective increase of surface tensions with temperature increments. In fact, for some isotherms it can be observed that Γ12 reaches a maximum value and then decreases, probably due to gas saturation at the interface of the IL.

Purity of gases (Table S.1); coefficients and mean square error or coefficient of determination for the different mixtures (Tables S.2 to S.5); surface tension for the different mixtures as a function of pressure (Figures S.1 to S.3). This material is available free of charge via the Internet at http://pubs.acs.org.

■

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected] and [email protected]. Funding

This work was financed by FONDECYT, Santiago, Chile (Project 1100938). G.R. acknowledges financial support from the Comisión Nacional de Investigación Cienti fí ca y Tecnológica (CONICYT) of Chile for a Ph.D. studentship. This work is supported by Red Doctoral REDOC.CTA, 1209



Article

on cubic equations of state with the Wong-Sandler mixing rule. J. Phase Equilib. Diff. 2005, 26, 215−224. (22) Kilaru, P. K.; Condemarin, R. A.; Scovazzo, P. Correlations of Low-Pressure Carbon Dioxide and Hydrocarbon Solubilities in Imidazolium-, Phosphonium-, and Ammonium-Based Room- Temperature Ionic Liquids. Part 1. Using Surface Tension. Ind. Eng. Chem. Res. 2008, 47, 900−909. (23) Ghatee, M. H; Zolghadr, A. R. Surface Tension Measurements of Imidazolium-Based Ionic Liquids at Liquid-Vapor Equilibrium. Fluid Phase Equilib. 2008, 263, 168−175. (24) Klomfar, J.; Souckova, M.; Patek, J. Surface Tension Measurements for Four 1-Alkyl-3-methylimidazolium-Based Ionic Liquids with Hexafluorophosphate Anion. J. Chem. Eng. Data 2009, 54, 1389−1394. (25) Osada, R.; Hoshino, T.; Okada, K.; Ohmasa, Y.; Yao, M. Surface Tension of Room Temperature Ionic Liquids Measured by Dynamic Light Scattering. J. Chem. Phys. 2009, 130, 184705−184713. (26) Yang, J.-Z.; Lu, X.-M.; Gui, J.-S.; Xu, W.-G. A new Theory for Ionic Liquids-The Interstice Model Part 1. The Density and Surface Tension of Ionic Liquid EMISE. Green Chem. 2004, 4, 541−543. (27) Restolho, J.; Mata, J. L.; Saramago, B. On the Interfacial Behaviour of Ionic Liquids: Surface Tensions and Contact Angles. J. Colloid Interface Sci. 2009, 340, 82−86. (28) Santos, C. S.; Baldelli, S. Alkyl Chain Interaction at the Surface of Room Temperature Ionic Liquids: Systematic Variation of Alkyl Chain Length (R = C1−C4, C8) in both Cation and Anion of (RMIM)(R-OSO3) by Sum Frequency Generation and Surface Tension. J. Phys. Chem. B 2009, 113, 923−933. (29) Torrecilla, J. S.; Palomar, J.; García, J.; Rodríguez, F. Effect of Cationic and Anionic Chain Lengths on Volumetric, Transport, and Surface Properties of 1-Alkyl-3-methylimidazolium Alkylsulfate Ionic Liquids at (298.15 and 313.15) K. J. Chem. Eng. Data 2009, 54, 1297− 1301. (30) Wandschneider, A.; Lehmann, J. K.; Heintz, A. Surface Tension and Density of Pure Ionic Liquids and Some Binary Mixtures with 1Propanol and 1-Butanol. J. Chem. Eng. Data 2008, 53, 596−599. (31) Marsh, K. N.; Boxall, J. A.; Lichtenthaler, R. Room Temperature Ionic Liquids and Their Mixtures-a review. Fluid Phase Equilib. 2004, 219, 93−98. (32) Tariq, M.; Freire, M. G.; Saramago, B.; Coutinho, J. A. P.; Canongia Lopes, J. N.; Rebelo, L. P. N. Surface Tension of Ionic Liquids and Ionic Liquid Solutions. Chem. Soc. Rev. 2012, 41, 829− 868. (33) Santos, C. S.; Baldelli, S. The Gas Liquid Interface of Hydrophobic and Hydrophilic Room-Temperature Ionic Liquids and Benzene: Sum Frequency Generation and Surface Tension Studies. J. Phys. Chem. C 2008, 112, 11459−11467. (34) Huddleston, J. G.; Visser, A. E.; Reichert, W. M.; Willauer, H. D.; Broker, G. A.; Rogers, R. D. Characterization and Comparison of Hydrophilic and Hydrophobic Room Temperature Ionic Liquids Incorporating the Imidazolium Cation. Green Chem. 2001, 3, 156− 164. (35) Freire, M. G.; Carvalho, P. J.; Fernandes, A. M.; Marrucho, I. M.; Queimada, A. J.; Coutinho, J. A. P. Surface Tensions of Imidazolium Based Ionic Liquids: Anion, cation, temperature and water effect. J. Colloid Interface Sci. 2007, 314, 621−630. (36) Łuczak, J.; Jungnickel, C.; Łacka, I.; Stolte, S.; Hupka, J. Antimicrobial and Surface Activity of 1-alkyl-3-methylimidazolium derivatives. Green Chem. 2010, 12, 593−601. (37) Modaressi, A.; Sifaoui, H.; Mielcarz, M.; Domańska, U.; Rogalski, M. Influence of the Molecular Structure on the Aggregation of Imidazolium Ionic Liquids in Aqueous Solutions. Colloids Surf., A 2007, 302, 181−185. (38) Fröba, A. P.; Wasserscheid, P.; Gerhard, D.; Kremer, H.; Leipertz, A. Revealing the Influence of the Strength of Coulomb Interactions on the Viscosity and Interfacial Tension of Ionic Liquid Cosolvent Mixtures. J. Phys. Chem. B 2007, 111, 12817−12822.

MINEDUC project UCO1202 at U. de Concepción. C.R.C. acknowledges Comissionat per a Universitats i Recerca del Departament d’Innovació, Universitat i Empresa de la Generalitat de Catalunya (2009SGR00465). Notes

The authors declare no competing financial interest.

■

REFERENCES

(1) Plechkova, N.; Seddon, K. Applications of Ionic Liquids in the Chemical Industry. Chem. Soc. Rev. 2008, 37, 123−150. (2) Wilkes, J. A Short History of Ionic Liquids-From Molten Salts to Neoteric Solvents. Green Chem. 2002, 4, 73−80. (3) Rogers, R.; Seddon, K. Ionic Liquids − Solvents of the Future? Science 2003, 302, 792−793. (4) Seddon, K. Ionic Liquids: A taste of the future. Nat. Mater. 2003, 2, 363−365. (5) Brennecke, J. F.; Maginn, E. J. Ionic liquids: Innovative Fluids for Chemical Processing. AIChe J. 2001, 47, 2384−2389. (6) Seddon, K. Ionic Liquids for Clean Technology. J. Chem. Tech. Biotechnol. 1997, 68, 351−356. (7) Cull, S.; Holbrey, J.; Mora, V.; Seddon, K.; Lye, G. RoomTemperature Ionic Liquids as Replacements for Organic Solvents in Multiphase Bioprocess Operations. Biotechnol. Bioeng. 2000, 69, 227− 233. (8) Grätzel, M. Dye-sensitized Solar Cells. J. Photochem. Photobiol. C 2003, 4, 145. (9) Tsuda, T.; Hussey, C. Electrochemical Applications of RoomTemperature Ionic Liquids. Electrochem. Soc. Interface 2007, 16, 42− 49. (10) Chiappe, C.; Pieraccini, D. Ionic Liquids: Solvent Properties and Organic Reactivity. J. Phys. Org. Chem. 2005, 18, 275−297. (11) Ramdin, M.; de Loos, T. W.; Vlugt, T. J. H. State-of-the-Art of CO2 Capture with Ionic Liquids. Ind. Eng. Chem. Res. 2012, 51, 8149− 8177. (12) Brennecke, J. F.; Gurkan, B. E. Ionic Liquids for CO2 Capture and Emission Reduction. J. Phys. Chem. Lett. 2010, 1, 3459−3464. (13) Blanchard, L.; Gu, Z.; Brennecke, J. High-Pressure Phase Behavior of Ionic Liquid/CO2 Systems. J. Phys. Chem. B 2001, 105, 2437−2444. (14) Morgan, D.; Ferguson, L.; Scovazzo, P. Diffusivities of Gases in Room-Temperature Ionic Liquids: Data and Correlations Obtained Using a Lag-Time Technique. Ind. Eng. Chem. Res. 2005, 44, 4815− 4823. (15) Perez-Salado Kamps, A.; Tuma, D.; Xia, J.; Maurer, G. Solubility of CO2 in the Ionic Liquid (bmim)(PF6). J. Chem. Eng. Data 2003, 48, 746−749. (16) Shiflett, M. B.; Yokozeki, A. Solubility and Diffusivity of Hydrofluorocarbons in Room-Temperature Ionic Liquids. AIChE J. 2006, 52, 1205−1219. (17) Jacquemin, J.; Husson, P.; Majer, V.; Costa Gomes, M. F. Low pressure solubilities and thermodynamics of solvation of eight gases in 1-butyl-3-methylimidazolium hexafluorophosphate. Fluid Phase Equilib. 2006, 240, 87−95. (18) Anthony, J. L.; Maginn, E. J.; Brennecke, J. F. Solubilities and Thermodynamic Properties of Gases in the Ionic Liquid 1-n-Butyl-3methylimidazolium Hexafluorophosphate. J. Phys. Chem. B 2002, 106, 7315−7320. (19) Jacquemin, J.; Husson, P.; Majer, V.; Padua, A. A. H.; Gomes, M. F. C. Thermophysical properties, low pressure solubilities and thermodynamics of solvation of carbon dioxide and hydrogen in two ionic liquids based on the alkylsulfate anion. Green Chem. 2008, 10, 944−950. (20) Mejía, A.; Segura, H.; Cartes, M. Vapor-Liquid Equilibria and Interfacial Tensions of the System Ethanol + 2-Methoxy-2methylbutane. J. Chem. Eng. Data 2011, 56, 3142−3148. (21) Mejía, A.; Segura, H.; Wisniak, J.; Polishuk, I. Correlation and prediction of interface tension for fluid mixtures: An approach based 1210



Article

(59) Gray, C. G.; Gubbins, K. E.; Joslin, C. G.; Keith, K. E. Theory of Molecular Fluids: Vol. 2: Applications, 1st ed.; Oxford University Press: New York, 2011. (60) Georgiadis, A.; Llovell, F.; Bismarck, A.; Blas, F. J.; Galindo, A.; Maitland, G. C.; Martin Trusler, J. P.; Jackson, G. Interfacial Tension Measurements and Modelling of (carbon dioxide plus n-alkane) and (carbon dioxide plus water) Binary Mixtures at Elevated Pressures and Temperatures. J. Supercrit. Fluids 2010, 55, 743−754. (61) Hebach, A.; Oberhof, A.; Dahmen, N.; Griesheimer, P. Interfacial Tension of Two Ionic Liquids, 1-Ethyl-3-methylimidazolium 2-(2-Ethoxyethoxy)ethylsulfate and 1-Butyl-3-methylimidazolium 2-(2-Methoxyethoxy)ethylsulfate, with Compressed CO2. J. Chem. Eng. Data 2009, 54, 1249−1253. (62) Jaeger, P.; Eggers, R. Interfacial Tension of Ionic Liquids at Elevated Pressures. Chem. Eng. Process. 2009, 48, 1173−1176. (63) Lee, D. J.; Telo da Gama, M. M.; Gubbins, K. E. Adsorption and Surface Tension Reduction at the Vapor-Liquid Interface. J. Phys. Chem. 1985, 89, 1514−1519. (64) Cadena, C. Molecular Modeling and Determination of Properties of Ionic Liquids; VDM Verlag: Weinheim, Germany, 2009. (65) Guggenheim, E. A.; Adam, N. K. The Thermodynamics of Adsorption at the Surface of Solutions. Proc. R. Soc. London A 1933, 139, 218−236. (66) Adamson, A. W.; Gast, A. P. Physical Chemistry of Surfaces, 6th ed.; Wiley Interscience: New York, 1997. (67) Masterton, W. L.; Blanchi, J.; Slowinski, E. J., Jr. Surface Tension and Adsorption in Gas-Liquid Systems at Moderate Pressures. J. Phys. Chem. 1962, 67, 615−618.

(39) Pereiro, A. B.; Rodriguez, A. Thermodynamic Properties of Ionic Liquids in Organic Solvents from (293.15 to 303.15) K. J. Chem. Eng. Data 2007, 52, 600−608. (40) Gonzalez, E. J.; Gonzalez, B.; Calvar, N.; Dominguez, A. Physical Properties of Binary Mixtures of the Ionic Liquid 1-Ethyl-3methylimidazolium Ethyl Sulfate with Several Alcohols at T = (298.15, 313.15, and 328.15) K and Atmospheric Pressure. J. Chem. Eng. Data 2007, 52, 1641−1648. (41) Seddon, K. R.; Stark, A.; Torres, M. J. Viscosity and Density of 1-Alkyl-3-methylimidazolium Ionic Liquids. ACS Symp. Ser. 2002, 819, 34−49. (42) Krummen, M.; Wasserscheid, P.; Gmehling, J. Measurement of Activity Coefficients at Infinite Dilution in Ionic Liquids Using the Dilutor Technique. J. Chem. Eng. Data 2002, 47, 1411−1417. (43) Mejía, A.; Segura, H.; Cartes, M.; Bustos, P. Vapor−liquid equilibrium, densities, and interfacial tensions for the system ethyl 1,1dimethylethyl ether (ETBE) + propan-1-ol. Fluid Phase Equilib. 2007, 255, 121−130. (44) Andreas, J. M.; Hauser, E. A.; Tucker, W. B. Boundary tension by pendant drops. J. Phys. Chem. 1938, 42, 1001−1019. (45) Rusanov, A. I.; Prokhorov, V. A. Interfacial Tensiometry; Elsevier: Amsterdam, 1996. (46) Ambwani, D. S.; Fort, T., Jr. Pendant drop technique for measuring liquid boundary tensions. In Surface and Colloid Science, Vol. 11; Good, R. J., Stromberg, R. R., Eds.; Plenum Press: New York, 1979; pp 93−119. (47) Mobius, D.; Miller, R. Drops and Bubbles in Interfacial Research, Vol. 6; Elsevier: Amsterdam, 1997. (48) Georgiadis, A.; Maitland, G.; Trusler, J. P.; Bismarck, A. Interfacial Tension Measurements of the (H2O + CO2) System at Elevated Pressures and Temperatures. J. Chem. Eng. Data 2010, 55, 4168−4175. (49) Georgiadis, A.; Maitland, G.; Trusler, J. P.; Bismarck, A. Interfacial Tension Measurements of the (H2O + n-Decane-CO2) Ternary System at Elevated Pressures and Temperatures. J. Chem. Eng. Data 2011, 56, 4900−4908. (50) Bashforth, F.; Adams, J. C. An attempt to test the theories of capillary action by comparing the theoretical and measured forms of drops of fluid; Cambridge University Press: Cambridge, 1883. (51) Hansen, F. K. A program system for interfacial tension and contact angle measurements by image analysis, advanced ed.; Ramé-Hart Instrument Co.: Succasunna, NJ, 2005. (52) Lemmon, W.; McLinden, M. O.; Friend, D. G. Thermophysical Properties of Fluid Systems. In NIST Chemistry WebBook, NIST Standard Reference Database Number 69; Linstrom, P. J., Mallard, W. G., Eds.; National Institute of Standards and Technology: Gaithersburg MD; http://webbook.nist.gov (retrieved March 2013). (53) NIST Ionic Liquids Database, ILThermo NIST Standard Reference Database 147; National Institute of Standards and Technology, Standard Reference Data Program: Gaithersburg, MD, 2006. http://ILThermo.boulder.nist.gov/ILThermo/ (retrieved March 2013). (54) Dittmar, D.; Fredenhagen, A.; Oei, S. B.; Eggers, R. Interfacial tension of ethanol-carbon dioxide and ethanol-nitrogen. Dependence of the interfacial tension on the fluid densityprerequisites and physical reasoning. Chem. Eng. Sci. 2003, 58, 1223−1233. (55) Anthony, J. L.; Anderson, J. L.; Maginn, E. J.; Brennecke, J. F. Anion Effects on Gas Solubility in Ionic Liquids. J. Phys. Chem. B 2005, 109, 6366−6374. (56) Bird, R. B.; Stewart, W. E.; Lightfoot, E. N. Transport Phenomena, 2nd ed.; Wiley: New York, 2002. (57) Georgiadis, A. Interfacial Tension of Aqueous and Hydrocarbon Systems in the Presence of Carbon Dioxide at Elevated Pressures and Temperatures, Ph.D. Thesis, Imperial College London, 2011. (58) Shiflett, M. B.; Yokozeki, A. Solubilities and Diffusivities of Carbon Dioxide in Ionic Liquids: (bmim)(PF6) and (bmim)(BF4). Ind. Eng. Chem. Res. 2005, 44, 4453−4464. 1211


Surface Tension of 1-Ethyl-3-methylimidazolium ... - ACS Publications

Recommend Documents