Enthalpy and Solubility Data of H2S in Aqueous Salt Solutions at

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Enthalpy and Solubility Data of H2S in Aqueous Salt Solutions at Conditions of Interest for Geological Sequestration Diana Koschel,† Jean-Yves Coxam, and Vladimir Majer* Clermont University, Blaise Pascal University, Institute of Chemistry of Clermont-Ferrand, BP 10448, F-63000 Clermont-Ferrand, France CNRS, UMR 6296, ICCF, BP 80026, F-63171 Aubiere, France S Supporting Information *

ABSTRACT: Describing the dissolution of H2S in aqueous electrolyte solutions under high pressures and at moderately elevated temperatures is of great interest in connection with geological storage of acid gases in deep saline aquifers. The dissolution in NaCl(aq) and CaCl2(aq) was studied by direct calorimetric measurements of the heat of mixing ΔHmix at conditions where hydrogen sulfide is gaseous, liquid, or supercritical. The measurements were carried out using a customized mixing unit developed for an isothermal differential heat flux calorimeter Setaram BT 2.15D. The experimental technique was specifically adapted for operating with H2S taking into consideration its high toxicity. The heats of mixing were measured with the 1, 3, and 5 mol·kg−1 NaCl(aq) at target temperatures of 323, 353, and 393 K and at pressures up to 13, 31, and 20 MPa, respectively. Analogous measurements were performed with 0.33 and 1 mol·kg−1 CaCl2(aq) at the medium temperature and pressures up to 31 MPa. The experiments were always carried out in the regions where the solution is unsaturated and saturated by H2S. The obtained concentration dependence of ΔHmix allowed one to derive simultaneously both the limiting enthalpy of solution ΔHsol of H2S and the concentration of its solubility limit. Comparison with the earlier published data for H2S in pure water at identical conditions (Koschel et al. Ind. Eng. Chem. Res. 2007, 46, 1421−1430) gives the information on the salting-out effect and its dependence on temperature and pressure, as well as on the concentration and nature of the salt. The newly obtained data are also compared with those resulting from the measurements with supercritical CO2 at analogous conditions (Koschel et al. Fluid Phase Equilib. 2006, 247, 107−120).

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INTRODUCTION

lower reliability than for carbon dioxide where much more information is available. From over 700 experimental solubility data points available for the H2S−H2O−NaCl system in the literature, about 40% are at atmospheric pressure only and about 75% are at pressures lower than 5 MPa. The data reaching above 15 MPa are available from one literature source only (see Table 1). The most extensive studies are those of Drummond,1 reporting about one-half of all available experimental values reaching both high temperatures and pressures, and of Barrett et al.2 who presented numerous data points, however, in a limited temperature range and at atmospheric pressure only. Drummond has encountered serious problems due to corrosion of his experimental equipment, and considerable internal discrepancies in the values for dissolution of H2S taken in different regimes at identical temperature and pressure are notable. These results should be therefore considered with extreme caution as well as the limited data of Kozintseva3 who reported also similar experimental difficulties. The most representative sources are the publications of Suleimenov and Krupp4 and of Xia et al.5 complementary regarding the temperature ranges covered. In addition, the latter source is the only publication where H2S solubilities are presented for

Hydrogen sulfide is an important natural component of geological fluids but highly toxic and hazardous for living organisms and the environment in general when released in an uncontrolled manner during technological operations. Being one of the constituents of natural gas, it must be captured and stored similarly like carbon dioxide, preferentially close to the production site for minimizing transport costs and ecological concerns. A geological sequestration in deep saline aquifers is one of the promising processes that is in active development by petroleum industries. In this connection, the data on solubility of H2S in electrolyte solutions as well as the enthalpic effects during dissolution at conditions encountered in geological storage are of great interest. It is expected that the temperatures typically encountered in deep aquifers are up to 423 K and occasionally higher while the fluid to be stored is introduced at pressures that may reach up to 80 MPa. The broader thermodynamic understanding of the dissolution process of H2S in aqueous systems as a function of temperature, pressure, and liquid phase composition is therefore desirable for modeling the sequestration process. Having consistent solubility and calorimetric data at high pressures has, besides its direct practical value, a more fundamental interest. Information on enthalpic effects during dissolution is useful in establishing and testing correlations of fluid solubility as a function of temperature. This is particularly important for hydrogen sulfide, where the data in literature are scarce and of © XXXX American Chemical Society

Received: June 20, 2013 Revised: August 31, 2013 Accepted: September 9, 2013

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dx.doi.org/10.1021/ie401947h | Ind. Eng. Chem. Res. XXXX, XXX, XXX−XXX

Industrial & Engineering Chemistry Research

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Table 1. Literature Review of the Solubility Data of Hydrogen Sulphide in Aqueous NaCl Solutions

a

mel (mol kg−1)

T (K)

p (MPa)

No. points

reference

0.5−1.2 1.0−6.0 1.0−5.0 0.2−2.5 4, 6

475−535 284−673 299−368 429−594 313−393

0.2−1.3a 0.6−30 0.1 2.7−14 0.3−10

5 351 239 23 57

Kozintseva3 Drummond1 Barrett et al.2 Suleimenov and Krupp4 Xia et al.5

Partial pressure of H2S.

compare our data for H2S with those for CO2 in NaCl(aq) determined at analogous conditions11 and to distinguish the differences in behavior of the two fluids during dissolution in the salt solution. Our determinations were performed over a relatively wide range of pressure (up to 31 MPa) to allow particularly for one to examine the extent of pressure influence on the gas solubility decrease due to the presence of an electrolyte. This is important for storage in saline aquifers where the dissolution takes place in fairly concentrated solutions at high pressures. The measurements were performed at temperatures of 323, 353, and 393 K and at NaCl(aq) concentration of 1 and 3 m (and also 5 m at the lowest temperature). The pressure ranges were from 2 to 13 MPa, 14 to 31 MPa, and 14 to 20 MPa for the lowest, medium, and highest temperatures, respectively. In addition, we have also obtained, at temperature of 353 K and pressures between 13 and 31 MPa, the data for dissolution in 0.33 and 1 m solutions of CaCl2(aq) which allowed one to examine the effect of ionic charge on the salting-out effect. The determinations with the latter system are the first of its kind reported in the literature. The flow calorimetry was the experimental technique used for studying the dissolution of H2S in aqueous solutions as a function of temperature and pressure. As has been already shown in our previous publications,10,11 the advantage of this technique is the possibility to determine simultaneously both the thermal effect, having its own practical value, and the solubility in a liquid phase. The analysis of the calorimetric signal versus the flow rate of dissolving hydrogen sulfide allows one to obtain both the enthalpy of solution and the limit where the solution is saturated. The applicability and reliability of our method for determination of solubility in electrolyte solutions was tested previously11 by comparing our carbon dioxide solubility data with those resulting from direct phase equilibria measurements reported in the literature.

concentrated NaCl solutions. The same group at the University of Kaiserslautern in Germany published also the data of H2S solubility in aqueous solutions of seven other salts at conditions analogous with the NaCl(aq) study, composed namely of Na+ and NH4+ cations and Cl−, NO3−, CH3COO−, and SO42− anions (Xia at al.5−7). The above literature data indicate that the presence of an electrolyte leads in most cases to the decrease of H2S solubility in water (salting-out effect). For H2S, the Setchenov constant

ks = log(xo/x)/mel

(1)

reflecting the salting-out effectiveness (xo and x stand for molar fractions of H2S in pure water and in a salt solution of molality mel) does not change much with temperature at least up to 400 K but decreases with increasing NaCl concentration (Xia et al.5). This indicates that the salting-out of H2S gets less pronounced with increasing ionic presence in NaCl(aq) similarly as in the case of CO2 (Malinin and Savelyeva,8 Malinin and Kurovskaya9). The salting-out effect as a function of pressure is more difficult to discuss for H2S due to the limited literature data. The values from Xia et al.5 at 393 K indicate a slight increase of salting-out at pressures up to about 3 MPa, while at pressures higher than 3 MPa the salting-out effect seems to be independent of pressure at least up to 10 MPa. The salting-out effect decreases with the size of an ion and increases with its valence as apparent from the results published by Xia et al.5 The ion with higher charge density is more effective in salting-out; therefore, H2S is more soluble in NH4Cl(aq) than in NaCl(aq) at the same salt concentrations due to the size difference between the sodium and ammonium ions. Higher charged ions enhance salting-out, which is demonstrated by the lower solubility of hydrogen sulfide in (NH4)2SO4(aq) compared to NaCl(aq). These differences in solubility due to the size and charge of ions get more pronounced with increasing pressure. A more complex behavior was observed for the salts containing the acetate ion.6 The increase in solubility of H2S (salting-in effect) was observed for low pressures (up to 0.5 MPa) for both sodium and ammonium acetates while, at higher pressures, hydrogen sulfide is salted-out in the case of NaCH3COO. This inversion can be explained by the fact that hydrogen sulfide and the acetate ion can interact chemically. This effect favoring salting-in weakens with increasing concentration of H2S in the solution and is also dependent on the nature of cation present. The aim of this study is to contribute to better understanding the influence of pressure, temperature, salt nature, and its concentration on the solubility and enthalpic data in connection with the H2S sequestration processes. This is facilitated by the availability of the data on H2S dissolution in pure water obtained at identical pressure and temperature conditions and published earlier.10 In addition, it is possible to

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EXPERIMENTAL SECTION Experimental Arrangement. The hydrogen sulfide dissolution in the salt solutions was studied using a customized mixing flow unit adapted to a Setaram BT 2.15D heat conduction differential calorimeter described in detail earlier by Koschel et al.11 The operation mode was analogous with that for experiments carried out with H2S in pure water;10 therefore, only a very succinct description is given here. The hydrogen sulphide and aqueous solution enter the mixing unit supplied from two high-pressure syringe ISCO pumps 100 DM and 260 D, with flow rates ranging from 0.005 to 0.8 mL/min and 0.2 to 0.7 mL/min, respectively. The H2S pump is connected to a nitrogen tank that enables the system to be flushed. All the flow lines are stainless steel tubes of 1.6 mm o.d. and 1.0 mm i.d. A specially adapted set of valves and check valves is added to ensure cleaning or replacement of a part of the mixing flow unit in safe conditions. B



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Figure 1. Enthalpies of mixing of H2S and aqueous NaCl solutions at a temperature of 323.1 K and pressures of 2 MPa □, 8 MPa ◇, and 13 MPa ○; at a temperature of 353.1 K and pressures of 14 MPa □, 20 MPa △, 26 MPa ◇, and 31 MPa ○; and at a temperature of 393.1 K and pressures of 14 MPa □ and 20 MPa △.

The experiments are carried out in an isobaric mode; the maximum working pressure for the unit conditioned by the tubes diameter is about 40 MPa. The system pressure is maintained constant to 0.02 MPa using a mechanic back pressure regulator placed at the end of the flow line. The pressure is measured by three electronic pressure gauges with accuracy to 0.25% of the full-scale. Their location is at the outlets of the acid gas and aqueous phase pumps and between the mixing cell and the back pressure regulator. Hydrogen sulfide dissolves in the aqueous phase in a customized mixing cell that is the core of the mixing unit. The two phases get in contact at the bottom of the cell where the input tubes bringing H2S and aqueous solution are connected to the mixing tube of the cell. The quantitative mixing occurs in this 2.8 m long tube, made also of a stainless steel tube of 1.6 mm o.d., coiled in good thermal contact with the inner wall of a confinement cylinder (about 20 mm of i.d., 80 mm height). The mixing cell is located inside the measuring calorimeter’s well, surrounded by the thermopiles that measure the heat power during mixing. The experiments are carried out at constant temperature; the entering fluids must reach the working temperature before entering the mixing cell with the help of three preheaters, one

external to the calorimeter and two inside it, the last one being located just above the calorimetric block. The temperature of the calorimeter is controlled with a stability of 0.01 K. When the solution is saturated, the excess H2S is scrubbed from the line, at the outlet of the back pressure regulator, using an alkaline solution. Specific measures were taken when operating with hydrogen sulfide. This is a toxic and corrosive gas. Being a cellular poison, it affects the central nervous and pulmonary systems. The permissive exposure limit is 10 ppm; the short-term exposure limit (S.T.E.L.) is 15 ppm, and immediately dangerous to life and health (I.D.L.H.) is 100 ppm. For that reason, the concentration of H2S in room atmosphere was monitored permanently by a single gas detector (Gas Alert Extreme Single Gas Monitor by BW Technologies). Because of the risks of gas leaks resulting from possible corrosion all along the flow lines of the mixing unit, a specially adapted set of valves and check valves was added to ensure the replacement of a part of the mixing flow unit in safe conditions. An extractor fan was permanently operated in order to allow rapid removal of polluted atmosphere in case of H2S leaks. Since the most substantial portion of H2S is located inside the high pressure C



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pump, this one was placed inside a permanently operating fume hood. Operation Procedure. The heat of mixing ΔHmix is defined here as a heat per one mole of the final mixture. This quantity is directly obtained from the thermopile signal S (μV) of the calorimeter and the molar flow rate ṅ (mol·s−1) of H2S and the aqueous solution. Before measuring the thermopile signal SM (μV) during the mixing process, a baseline signal SBL (μV) is recorded with only the aqueous phase flowing through the calorimeter. The baseline signal is kept close to zero by adjusting the thermoregulation of the entering fluids. The heat of mixing ΔHmix (J·mol−1) is then calculated following eq 2 as the ratio of the difference in the thermopile signals and the overall molar flow rate: ΔHmix =

(SM − SBL) Kn ̇

(2)

The molar flow rate ṅ is derived from the volumic pump flow rates, the densities, and the salt concentration. The thermopile sensitivity K (μV·mW−1), converting the thermopile signal to heat power, can be determined from electric calibration by the Joule effect using a calibration heater or by a chemical calibration with two fluids where ΔHmix is well-known. The latter method was adopted in this study with the reference {ethanol + water} system, using the enthalpy data provided by Ott et al.12,13 All the measured values of the heat of mixing for the aqueous solution of NaCl and CaCl2 are listed in tables in the Supporting Information. The expected experimental uncertainty in δΔH is determined as a statistical estimate from the expected error in K, uncertainty in the molar flow rate, and the heat power based on eq 2. The uncertainty of the molar flow rate, estimated to be smaller than 1.5%, depends on the error of the volumic flow rates of pumps and the accuracy of the fluid densities. The error connected with the heat power is related to the accuracy of the thermopile sensitivity of the calorimeter, estimated to 2%, and to the fluctuations of the calorimetric signal. The uncertainty in the difference between SM and SBL is between 1 and 3% and can reach 5% for the smallest measured heat effects. The experiments are carried out at constant temperature and pressure and heats of mixing are determined at different concentrations resulting from the flow rate ratios change of H2S versus aqueous solution. The enthalpy of solution ΔHsol (defined here as heat per one mole of solute) and the solubility are determined simultaneously from the plot of ΔHmix versus the molar fraction xH2S in the aqueous solution, as shown in Figures 1 and 2 for solutions of NaCl and CaCl2, respectively. These graphs make it possible to distinguish the sections before and after saturation of aqueous solution by H2S. First, the absorption in the solution is complete, and therefore, the absolute value of the heat of mixing increases strongly and approximately linearly ΔHmix = a1 xH2S. Second, ΔHmix changes much less when the aqueous solution does not absorb H2S any more ΔHmix = a2 xH2S + b. At the lowest temperature, the heat of mixing remains approximately constant after saturation but continually increases at the two higher temperatures due to an additional heat effect, related to the dissolution of water in the undissolved H2S. It is also apparent that dissolution of gaseous hydrogen sulfide in aqueous phase is exothermic at conditions where H2S is gaseous, i.e., at a temperature of 323 K and pressure 2 MPa, the saturation pressure psat(H2S) being 3.5 MPa. The dissolution process is endothermic at the medium

Figure 2. Enthalpies of mixing of H2S and aqueous CaCl2 solutions at a temperature of 353.1 K and pressures of 13 MPa □, 26 MPa Δ, and 31 MPa ○.

and highest temperature where hydrogen sulfide is liquid and supercritical, respectively, the critical parameters of H2S being Tc = 373.4 and pc = 8.94 MPa. This can be understood when imagining dissolution of a gas in the liquid phase as a combination of a strongly exothermic condensation process with a subsequent endothermic mixing, the first effect being thermally dominant. The solubility limit is determined from the intersection of the curves fitting the data before and after saturation as xH2S = b/(a1 − a2). The enthalpy of solution ΔHsol is defined exactly as the difference between the enthalpies of an infinitely dilute and pure solute at the same temperature and pressure. It is reported to one mole of dissolved gas and corresponds thus to the slope of the ΔHmix curve in the limit of infinite dilution. In most cases, the concentration dependence of ΔHmix is close to linear before reaching the domain of saturation and the determination of the slope is thus straightforward ΔHsol = a1. The uncertainties of solubility and enthalpy of solution were determined from the statistical errors in a1 and b as δxH2S (%) = (Δa1/a1 + Δb/b)·100, since a1 ≫ a2, and δΔHsol (%) = (Δa1/a1)·100. Finally, the average uncertainty in the solubility and the enthalpy of solution for H2S are typically between 4% and 8% and between 3% and 6%, respectively.

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RESULTS AND DISCUSSION This section shows and interprets the influence of salt presence (NaCl or CaCl2) in water on the solubility and enthalpy of solution of hydrogen sulfide at different combinations of temperature and pressure where H2S is in the gaseous, liquid, or supercritical states. A comparison with analogous data reported in the literature for carbon dioxide is made where possible. Solubility. The numerical values of the hydrogen sulfide solubility in aqueous solutions of NaCl at the three experimental temperatures and several pressures are reported D



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in Table 2; values also were added for comparison of solubilities in pure water, determined during the same campaign of Table 2. Solubility of H2S in Water and Aqueous NaCl Solutions Obtained from the Calorimetric Data mNaCl (mol kg−1) 0.000 0.000 0.000 0.000 1.000 1.000 1.000 3.000 3.000 3.000 5.000 5.000 5.000 T = 353.1 K 0.000 0.000 0.000 0.000 1.000 1.000 1.000 1.000 3.000 3.000 3.000 3.000 T = 393.1 K 0.000 0.000 0.000 0.000 1.000 1.000 3.000 a

p (MPa)

xH2S

δxH2S (%)

T = 323.1 K 1.72 0.0175a 1.84 0.0187a 7.83 0.0388a 13.6 0.0390a 1.76 0.0148 7.89 0.0333 13.18 0.0344 1.77 0.0119 7.89 0.0231 13.32 0.0250 1.77 0.0097 7.88 0.0217 12.89 0.0231

59 6 7 7 5 6 5 5 7 7 6 7 5

14.02 19.16 25.45 30.86 14.02 19.92 26.00 31.17 13.77 19.87 26.05 30.93

0.0438a 0.0443a 0.046a 0.0471a 0.0375 0.0381 0.0385 0.0394 0.0260 0.0265 0.0282 0.0300

5 5 5 4 6 5 6 5 6 6 5 6

13.58 19.86 25.48 30.27 13.51 19.18 19.59

0.0554a 0.0570a 0.0585a 0.0603a 0.0462 0.0500 0.0326

6 7 6 6 8 6 6

Solubility values of H2S in pure water reported by Koschel et al.10.

measurements but published earlier.10 The last column lists the expected uncertainty in percent of the mole fraction estimated from the scatter of the experimental points used for the indirect determination of solubility. Analogous data for H2S in the aqueous solution of CaCl2 at a temperature of 353 K are listed in Table 3. Figures 3 and 4 are the graphic representation of the data for NaCl(aq) and CaCl2(aq), respectively, that allow one to visualize the evolution of solubility with pressure at each

Figure 3. Solubility of H2S in water and aqueous NaCl solutions as a function of pressure at temperatures of 323.1 K (I), 353.1 K (II), and 393.1 K (III).

experimental temperature as well as the solubility decrease with increasing concentration of an electrolyte. This salting-out effect (s.o.) can be quantified as the percentage decrease in solubility of H2S in aqueous solution of electrolyte x, related to solubility in pure water xo, s.o. = 100(xo − x)/xo, or by the value of the Setchenov constant; see eq 1. It is apparent first of all that the solubility of gaseous hydrogen sulfide both in pure water and in a salt solution is substantially lower (T = 323 K, p = 2 MPa) than that of liquid or supercritical H2S. When we analyze the data for NaCl(aq), our results suggest that, similar to CO2 at experimental conditions of this

Table 3. Solubility of H2S in Aqueous CaCl2 Solutions Obtained from the Calorimetric Data at 353.1 K mCaCl2 (mol kg−1)

p (MPa)

xH2S

δxH2S (%)

0.333 0.333 0.333 1.000 1.000 1.000

13.48 26.23 30.50 13.79 26.03 30.90

0.0393 0.0411 0.0422 0.0310 0.0342 0.0343

5 6 5 5 5 5 E



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concentrations (0.141 kg·mol−1). Thus, when comparing salting-out in 1 m NaCl and CaCl2 solutions, it is two times higher for the latter salt. The ratio of the Setchenov constants ks(CaCl2)/ks(NaCl) resulting from this investigation is consistent with the value of 1.8 calculated for a temperature of 298 K from the correlation of Weisenberger and Schumpe.14 This is logical since the CaCl2 solution contains a higher number of ions where in addition Ca2+ has a higher charge, thus binding more molecules of water. However, when studying the influence of the salt nature on the salting-out effect, it is of interest to also make comparisons at the identical ionic strength I = 1/2∑mizi2, where mi and zi are ionic molalities (mols of ion per kilogram of water) and their valence, respectively. This is the case for 1 and 0.333 m solutions (I = 1 mol·kg−1) and for 3 and 1 m solutions (I = 3 mol·kg−1) of NaCl and CaCl2, respectively. Comparison of our results for the two salts indicates that 1−1 electrolyte is more efficient in salting-out than 2−1 electrolyte of the same ionic strength. The influence of pH on the solubility of hydrogen sulfide in a 1 m NaCl solution was also examined at two different pressures (14 and 31 MPa). The pH of a 1 m solution, originally about 5.6, was changed by adding HCl until a pH of 3.2 was reached. The comparison of the solubilities of hydrogen sulfide at two different pH values shows that there is apparently no influence on the solubility. This finding is consistent with the results of Douabul and Riley15 who noted that the solubility of H2S in 5 m sulfuric acid was essentially the same as in pure water. The assumption that the acid would drive out the hydrogen sulfide from the solution was not confirmed. Carroll16 analyzed the influence of pH on the solubility of hydrogen sulfide in water and investigated the speciation at higher pH values. He found that the predominant form of the sulfide species at low pH (acidic solution) is the molecular H2S until a pH of about 6 when the bisulfide (S2−) becomes present. With a pH slightly less than 7, there are equal amounts of the molecular form and the disulfide anions. Table 4 lists salting-out effects and Setchenov constants for dissolution of CO2 and H2S at similar experimental conditions allowing a quantitative comparison between the two solutes at temperatures of 323 and 393 K; our earlier data11 and those

Figure 4. Solubility of H2S in water and aqueous CaCl2 solutions as a function of pressure at a temperature of 353.1 K.

investigation, the pressure does not have much of an affect on the salting-out. No systematic trend can be observed, taking into account uncertainty of our solubility values. The s.o. values for dissolving gaseous or liquid H2S at a temperature of 323 K are on average 15%, 37%, and 44% for 1, 3, and 5 m NaCl solutions, respectively. The value for the highest salt concentration is in reasonable agreement with 47% in solubility decrease calculated for 1.7 MPa from the data of by Xia et al.5 The averaged Setchenov constants ks at 323 K are 0.069, 0.067, and 0.050 kg·mol−1 for 1, 3, and 5 m solutions, respectively. It can therefore be considered that the Setchenov law is approximately valid at least until a concentration of 3 mol· kg−1. The drop in ks at the highest concentration documents the decrease in the salting-out efficiency. The pressure range of our measurements was the largest at a temperature of 353 K reaching up to 31 MPa, yet again, no systematic pressure dependence was observed with the average s.o. values of 15% and 39% for 1 and 3 m NaCl solutions, respectively, and the averaged Setchenov constants practically constant (0.072 kg· mol−1). Less data are available at the highest temperature of 393 K giving the average s.o. and ks values of 14% and 0.068 kg· mol−1, respectively, for the lower concentration, that are close to those for two lower temperatures. For the 3 m solution, the s.o. and ks values are 43% and 0.081 kg·mol−1, respectively, yet these two values should be regarded with caution since they are based on determination at one pressure only and more measurements would be needed for verification. Suleimenov and Krupp4 noted that the influence of temperature on the salting-out effect is negligible up to around 490 K .Our study confirms this finding in the temperature range of our measurements where the average s.o. values are between 14% and 15% and 37−39% for 1 and 3 m NaCl solutions, respectively, and the Setchenov constants are between 0.067 and 0.072 kg·mol−1 for salt concentrations at least up to 3 m (the data at temperature of 393 K for 3 m solution were not taken into account). The solubility of hydrogen sulfide is listed in Table 3 for 0.333 and 1 m CaCl2 solutions at a temperature of 353 K for demonstrating the influence of the salt valence. A strong decrease of the hydrogen sulfide solubility in the aqueous CaCl2 solution compared to that in pure water can be observed. The salting-out effect in 0.333 and 1 m CaCl2 solutions is on average about 10% and 27%, respectively. Again, no significant trend of the salting-out as a function of pressure can be noted. The Setchenov constants are practically identical for both

Table 4. Comparison of Salting-Out Effects for Carbon Dioxide and Hydrogen Sulphide in Aqueous NaCl Solutions at Temperatures of 323 and 393 K CO2

a b

F

T (K)

mNaCl

p (MPa)

323 323 323 323 323 323 323 323 323 323 323 393 393

1 1 1 1 1 1 3 3 3 3 3 1 1

2 5 8 10 13 14 2 5 8 10 14 13 19

H2S b

s.o. (%)

ks

22

0.108

18

0.086

17

0.081

42

0.079

43

0.081

21a 21a

0.102 0.102

s.o. (%)

ksb

15

0.071

14

0.066

12

0.055

32

0.056

40

0.075

36 17 12

0.065 0.079 0.057

Calculated values from the correlation of Duan and Sun.17 Setchenov constant calculated with eq 1. dx.doi.org/10.1021/ie401947h | Ind. Eng. Chem. Res. XXXX, XXX, XXX−XXX


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Table 5. Enthalpy of Solution ΔHsol,, Residual Enthalpies ΔH*res and ΔH∞ res, and Enthalpy of Hydration ΔHhyd of H2S in Water and Aqueous NaCl Solutions

a

mNaCl (mol kg−1)

P (MPa)

ΔHsol [T, p]

0.000 0.000 0.000 0.000 1.000 1.000 1.000 3.001 3.001 3.001 5.000 5.000 5.000

1.72a 1.84a 7.83a 13.36a 1.76 7.89 13.18 1.77 7.89 13.32 1.77 7.88 12.89

−13.5 −13.1 1.4 1.4 −13.6 1.2 1.1 −13.0 1.5 1.3 −12.1 1.3 1.3

0.000 0.000 0.000 0.000 1.000 1.000 1.000 1.000 3.000 3.000 3.000 3.000

14.02a 19.16a 25.45a 30.86a 14.02 19.92 26.00 31.17 13.77 19.87 26.05 30.93

2.9 3.2 3.3 3.3 2.9 3.0 3.1 3.3 2.8 3.0 3.1 3.1

0.000 0.000 0.000 0.000 1.000 1.000 3.000

13.58a 19.86a 25.48a 30.27a 13.51 19.18 19.59

3.9 5.2 5.6 5.9 3.4 4.7 4.3

δΔHsol [T, p] T = 323 K 0.5 0.4 0.1 0.1 0.4 0.0 0.0 0.6 0.1 0.1 0.6 0.1 0.1 T = 353 K 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1 T = 393 K 0.1 0.2 0.2 0.3 0.2 0.2 0.2

ΔH*res (kJ mol−1 H2S)

ΔH∞ res

ΔHhyd [T, po]

−0.8 −0.8 −14.6 −14.7 −0.8 −14.6 −14.7 −0.8 −14.6 −14.7 −0.8 −14.6 −14.7

0 0 0.2 0.3 0 0.2 0.3 0 0.2 0.3 0 0.2 0.3

−14.2 −13.9 −13.4 −13.6 −14.4 −13.6 −13.9 −13.8 −13.3 −13.7 −12.9 −13.5 −13.7

−13.2 −13.4 −13.5 −13.5 −13.2 −13.4 −13.5 −13.5 −13.2 −13.4 −13.5 −13.5

0.2 0.3 0.4 0.5 0.2 0.3 0.4 0.5 0.2 0.3 0.4 0.5

−10.5 −10.5 −10.6 −10.5 −10.6 −10.7 −10.8 −10.7 −10.6 −10.7 −10.8 −10.9

−9.2 −10.8 −11.4 −11.6 −9.1 −10.7 −10.8

0.1 0.2 0.2 0.2 0.1 0.2 0.2

−5.4 −5.8 −6.0 −6.0 −5.8 −6.2 −6.6

Enthalpic values for H2S in pure water reported by Koschel et al.10

Table 6. Enthalpy of Solution ΔHsol, Residual Enthalpies ΔH*res and ΔH∞ res, and Enthalpy of Hydration ΔHhyd of Hydrogen Sulphide in Aqueous CaCl2 Solutions

a

mCaCl2 (mol kg−1)

p (MPa)

ΔHsol [T, p]

δΔHsol [T, p]

ΔH*res (kJ mol−1 H2S)

ΔH∞ res

ΔHhyd [T, po]

0.000 0.000 0.000 0.333 0.333 0.333 1.000 1.000 1.000

14.02a 19.16a 25.45a 13.48 26.23 30.50 13.79 26.03 30.9

2.9 3.2 3.3 2.8 3.1 3.2 2.4 2.7 2.9

0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1

−13.2 −13.4 −13.5 −13.1 −13.5 −13.5 −13.2 −13.5 −13.5

0.2 0.3 0.4 0.2 0.4 0.5 0.2 0.4 0.5

−10.5 −10.5 −10.6 −10.6 −10.8 −10.8 −11.0 −11.2 −11.2

Enthalpic values for H2S in pure water reported by Koschel et al.10

from the correlation of Duan and Sun17 were used for CO2. Generally, the salting-out affects more significantly nonelectrolytes with a higher partial molar volume and a lower polarity. Therefore, carbon dioxide should be salted-out more than hydrogen sulfide. When the experimental uncertainty of our values are taken into account, which can be estimated

between 5 and 10% for s.o., the table suggests comparable salting-out of both solutes by NaCl, yet the salting-out seems to be somewhat lower for hydrogen sulfide as expected. Carbon dioxide is certainly much closer in salting-out behavior to hydrogen sulfide than to methane. G



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⎛ ∂ln kH ⎞ ⎟ ΔHhyd(T , p0 ) ≅ ΔHhyd(T , psat,H O ) = −RT 2⎜ 2 ⎝ ∂T ⎠ p

Enthalpy of Solution. The enthalpy of solution ΔHsol (in Joule per mole of dissolved H2S) at constant temperature and pressure in NaCl(aq) and CaCl2(aq) was determined as described above from the concentration dependence of ΔHmix represented by the curves in Figures 1 and 2. The obtained enthalpic data for ternary systems are therefore analogous to those characterizing the dissolution process of hydrogen sulfide in pure water at temperatures of 323, 353, and 393 K and in a pressure range between 2 and 31 MPa. The ΔHsol data both for binary and ternary systems are listed in Tables 5 and 6. The discontinuity in the pressure dependence of enthalpy of solution of hydrogen sulfide in water at 323 K and the more elevated increase of ΔHsol with pressure at 393 K compared to 353 and 323 K is analogous to the binary system. No experimental data exist for direct comparison of our enthalpies of solution with literature at conditions of this investigation. The enthalpies of hydration, based on the obtained enthalpies of solution, were calculated similarly as for dissolution of H2S in pure water.10 This allows one to test the internal consistency of our results (binary and ternary data) and the agreement with literature, since the enthalpies of hydration can be obtained also from the temperature derivatives of the reported Henry’s law constants. The enthalpy of hydration ΔHhyd at a standard pressure p0 = 0.1 MPa corresponds to the difference between the enthalpies of H2S in the states of infinite dilution and of ideal gas as follows:

⎛ dln kH ⎞ ⎟ ≅ −RT 2⎜ ⎝ dT ⎠sat

The enthalpies of hydration of hydrogen sulfide in water and NaCl solutions calculated from the correlation of Drummond show a different trend than that obtained by our measurements. Nevertheless, our data (unlike those of Drummond) are consistent with the enthalpy of hydration of H2S in pure water calculated from the recommended Henry’s law constants of Fernandez-Prini et al.20 via eq 4 and from the SOCW hydration model19 using the parameters in ref 21. Comparison of ΔHsol and ΔHhyd values for dissolution of H2S in NaCl(aq) and CaCl2(aq) at a temperature of 353 K indicates that there is apparently no significant influence of the salt nature on the enthalpy of solution or hydration since the obtained values agree within several percent.

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CONCLUSIONS In this study, we wanted to contribute to understanding the relationship between H2S solubility and enthalpic effects on one side and conditions of dissolution on the other side, with an accent put to investigation of the salting-out effect. We can conclude that the evolutions of solubility of hydrogen sulfide in water and salt solution are analogous in the temperature and pressure range of this investigation. It means that the saltingout effect does not depend much on pressure and temperature but mainly on the salt concentration. The solubility of hydrogen sulfide is affected strongly by pressure when H2S is gaseous but is not very pressure dependent when two liquid phases coexist. However, when hydrogen sulfide is supercritical, the solubility increases somewhat faster with pressure again. Since the solubility of H2S increases with temperature, the conditions for the sequestration process are favorable mainly at supercritical conditions of hydrogen sulfide (T > 373 K and p > 10 MPa). Generalizing with caution, we can say that the saltingout efficiency decreases at concentrations above 3 m for NaCl(aq), where a drop in the Setchenov constant is observed. The salting-out effect is stronger for electrolytes with ions of higher valence yet 1−1 electrolyte produces higher salting-out effect than a 2−1 electrolyte at the constant ionic strength. The enthalpies of solution are negative at conditions where hydrogen sulfide is gaseous and positive for liquid or supercritical H2S increasing slightly with temperature. They are not significantly affected by salt presence and the valence of the electrolytes investigated.

* (p → p) + ΔHsol(p) ΔHhyd(p0 ) = ΔHres 0 ∞ − ΔHres (p0 → p)

(4)

(3)

where ΔHres * and ΔH∞ res are the residual enthalpies of pure solute and solute in the state of infinite dilution that were calculated from software ALLPROPS18 and the SOCW hydration model by Sedlbauer et al.,19 respectively, and are listed in Tables 5 and 6. The residual enthalpy of solute at infinite dilution was calculated neglecting the presence of salt in water. A constancy of the calculated enthalpies of hydration based on the enthalpy of solution at different pressures can be observed which indicates good internal consistency of the measurements. Our ΔHhyd of hydrogen sulfide in water and NaCl solutions are ploted in Figure 5 where also are those by Drummond1 resulting from the correlation of Henry’s law constant kH using relationship

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ASSOCIATED CONTENT

S Supporting Information *

Heats of mixing ΔHmix of H2S and aqueous NaCl solutions (1, 3, and 5 m) at 323.1 K, from 2 to 13 MPa (Table SI 1). Heats of mixing ΔHmix of H2S and aqueous NaCl solutions (1 and 3 m) at 353.1 K, from 14 to 31 MPa (Table SI 2). Heats of mixing ΔHmix of H2S and aqueous NaCl solutions (1 and 3 m) at 393.1 K, from 14 to 20 MPa (Table SI 3). Heats of mixing ΔHmix of H2S and CaCl2 solutions (0.33 and 1 m) at 353.1 K, from 13 to 31 MPa (Table SI 4). This information is available free of charge via the Internet at http://pubs.acs.org.

Figure 5. Enthalpy of hydration of H2S in water and aqueous NaCl solutions at temperatures of 323.1, 353.1, and 393.1 K; our work ■, Drummond1 dashed lines, Fernandez-Prini et al.20 ○, and Sedlbauer et al.19 ▲. H



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Article

(14) Weisenberger, S.; Schumpe, A. Estimation of gas solubilities in salt solutions at temperatures from 273−363K. AIChE J. 1996, 42, 298−300. (15) Douabul, A. A.; Riley, J. P. The solubility of gases in distilled water and seawater-hydrogen sulphide. Deep Sea Res. 1979, 26A, 259− 268. (16) Carroll, J. J. What is Henry’s law? Chem. Eng. Prog. 1991, 87 (9), 48−52. (17) Duan, Z.; Sun, R. An improved model for calculating CO2 solubility in pure water and aqueous NaCl solutions from 273 to 533 K and from 0−2000 bar. Chem. Geol. 2003, 193, 257−271. (18) ALLPROPS software Thermodynamic properties of fluids, Version 4.1. University of Idaho: Moscow, ID, 1995. (19) Sedlbauer, J.; O’Connell, J. P.; Wood, R. H. A new equation of state for correlation and prediction of standard molal properties of aqueous electrolytes and nonelectrolytes at high temperatures and pressures. Chem. Geol. 2000, 163, 43−63. (20) Fernandez-Prini, R.; Alvarez, J.; Harvey, A. Henry’s constants and vapor-liquid distribution constants for gaseous solutes in H2O and D2O at high temperatures. J. Phys. Chem. Ref. Data 2003, 32, 903−916. (21) Majer, V.; Sedlbauer, J.; Bergin, G. Henry’s law constant and related coefficients for aqueous hydrocarbons, CO2 and H2S over a wide range of temperature and pressure. Fluid Phase Equilib. 2008, 272, 65−74.

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Phone + 33 4 73407188. Present Address

† D.K.: Vallourec & Mannesmann Tubes, Vallourec Research Aulnoye, 60 route de Leval, 59620 Aulnoye-Aymeries, France.

Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This research was partly funded by TOTAL in the frame of the thesis of D. Koschel. The authors thank J. Sedlbauer from the University of Liberec for calculating the enthalpies of hydration from the SOCW model.

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REFERENCES

(1) Drummond, D. E. Boiling and mixing of hydrothermal fluids: Chemical effects on mineral precipitation. Ph.D. Thesis, Pennsylvania State University, State College, PA, 1981. (2) Barrett, T. J.; Anderson, G. M.; Lugowski, J. The solubility of hydrogen sulphide in 0−5 m NaCl solutions at 25−95°C and one atmosphere. Geochim. Cosmochim. Acta 1988, 52, 807−811. (3) Kozintseva, T. N. Rostvorimost serovodoroda v vode i solevykh rastvorakh pri povyshenykh temperaturakh. In Geokhimichskie issledovaniya v oblasti povyshennykh davlenii i temperatur; Khitarova, N. I., Ed; Izdatelstvo Nauka: Moscow, 1963; pp. 121−134. (4) Suleimenov, O. M.; Krupp, R. E. Solubility of hydrogen sulphide in pure water and in NaCl solutions, from 20 to 320°C and at saturation pressures. Geochim. Cosmochim. Acta 1994, 58, 2433−2444. (5) Xia, J.; Perez-Salado Kamps, A.; Rumpf, B.; Maurer, G. Solubility of hydrogen sulfide in aqueous solutions of the single salts sodium sulfate, ammonium sulfate, sodium chloride, and ammonium chloride at temperatures from 313−393 K and total pressures up to 10 MPa. Ind. Eng. Chem. Res. 2000, 39, 1064−1073. (6) Xia, J.; Perez-Salado Kamps, A.; Rumpf, B.; Maurer, G. Solubility of H2S in (H2O+CH3COONa) and (H2O+CH3COONH4) from 313−393 K and at pressures up to 10 MPa. J. Chem. Eng. Data 2000, 45, 194−201. (7) Xia, J.; Perez-Salado Kamps, A.; Rumpf, B.; Maurer, G. Solubility of hydrogen sulfide in aqueous solutions of single strong electrolytes sodium nitrate, ammonium nitrate, and sodium hydroxide at temperatures from 313−393 K and total pressures up to 10 MPa. Fluid Phase Equilib. 2000, 167, 263−284. (8) Malinin, S. D.; Savelyeva, N. I. Solubility of CO2 in NaCl and CaCl2 solutions at 25, 50, 75°C under elevated CO2 pressures. Geochem. Int. 1972, 9, 410−418. (9) Malinin, S. D.; Kurovskaya, N. A. Solubility of CO2 in chloride solutions at elevated temperatures and CO2 pressures. Geochem. Int. 1975, 12, 199−201. (10) Koschel, D.; Coxam, J.-Y.; Majer, V. Enthalpy and solubility data of H2S in water at conditions of interest for geological sequestration. Ind. Eng. Chem. Res. 2007, 46, 1421−1430. (11) Koschel, D.; Coxam, J.-Y.; Rodier, L.; Majer, V. Enthalpy and solubility data of CO2 in water and NaCl(aq) at conditions of interest for geological sequestration. Fluid Phase Equilib. 2006, 247, 107−120. (12) Ott, J. B.; Cornett, G. V.; Stouffer, C. E.; Woodfield, B. F.; Guanquan, C.; Christensen, J. J. Excess enthalpies of (ethanol + water) at 323.15, 333.15, 348.15 and 373.15 K from 0.4 to 15 MPa. J. Chem. Thermodyn. 1986, 18, 867−875. (13) Ott, J. B; Stouffer, C. E.; Cornett, G. V.; Woodfield, B. F.; Guanquan, C.; Christensen, J. J. Excess enthalpies for (ethanol+water) at 398.15, 423.15, 448.15 and 473.15 K at pressures of 5 and 15 MPa. Recommendations for choosing (ethanol+water) as an HE reference. J. Chem. Thermodyn. 1987, 19, 349−351. I


Enthalpy and Solubility Data of H2S in Aqueous Salt Solutions at

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