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Simultaneous solubility of ammonia and hydrogen sulfide in water at temperatures from 313 K to 393 K
Fluid Phase Equilibria 158–160 Ž1999. 923–932
Simultaneous solubility of ammonia and hydrogen sulfide in water at temperatures from 313 K to 393 K ´ Perez-Salado B. Rumpf, A. Kamps, R. Sing, G. Maurer ´
)
Lehrstuhl fur ¨ Technische Thermodynamik, UniÕersitat ¨ Kaiserslautern, D-67653 Kaiserslautern, Germany Received 19 March 1998; accepted 11 September 1998
Abstract The simultaneous solubility of ammonia and hydrogen sulfide in water was measured at temperatures from 313 to 393 K and total pressures up to 0.7 MPa. An adapted Pitzer model is used to correlate the new data. Experimental results are reported and compared to the limited literature data and the correlation. q 1999 Elsevier Science B.V. All rights reserved. Keywords: Vapour–liquid equilibrium; Weak electrolytes; Ammonia; Hydrogen sulfide; Water; Experimental data; Correlation
1. Introduction The simultaneous solubility of weak electrolyte gases like ammonia, carbon dioxide, hydrogen sulfide or hydrogen cyanide in aqueous solutions is known in many applications. Typical examples are the cleaning of raw gases in power stations, the production of fertilizers or applications in the field of environmental protection. The design of separation equipment requires reliable models to calculate vapour–liquid equilibria, caloric effects and often also kinetic effects. Due to chemical reactions in the liquid phase and a strong deviation from ideality, correlating and predicting the thermodynamic properties of aqueous systems containing ammonia and sour gases is an extremely difficult task. Besides this, experimental data on the simultaneous solubility of ammonia and sour gases in aqueous phases are in most cases scarce, e.g., for systems containing hydrogen sulfide. Therefore, continuing earlier work on the solubility of ammonia and sour gases in aqueous phases Ž cf. Refs. w1,2x. , the simultaneous solubility of ammonia and hydrogen sulfide in water was measured in the temperature range from 313 to 393 K and for total pressures up to 0.7 MPa. Experimental results for the partial )
Corresponding author. Tel.: q49-631-205-2410; fax: q49-631-205-3835; e-mail: [email protected]
0378-3812r99r$ - see front matter q 1999 Elsevier Science B.V. All rights reserved. PII: S 0 3 7 8 - 3 8 1 2 Ž 9 9 . 0 0 1 1 0 - 7
B. Rumpf et al.r Fluid Phase Equilibria 158–160 (1999) 923–932
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Table 1 Experimental results for the solubility of ammonia and hydrogen sulfide in water T ŽK.
m NH 3 Žmolrkg.
m H 2 S Žmolrkg.
10 p NH 3 ŽMPa.
10 p H 2 S ŽMPa.
10 p ŽMPa.
313.19 313.17 313.17 313.17 313.16 313.17 313.15 313.15 313.15 313.16 313.15 313.19 313.17 313.18 313.17 313.17 313.18 313.14 313.15 313.15 313.14 313.14 313.12 353.17 353.14 353.14 353.13 353.13 353.12 353.13 353.14 353.13 353.13 353.14 353.13 353.13 353.14 353.14 353.13 353.13 353.14 353.14 353.14 353.13 353.14 353.14 353.14
3.002 3.002 3.003 3.005 3.007 3.008 3.009 3.011 3.011 3.011 3.011 4.938 4.938 4.940 4.942 4.945 4.947 4.949 4.952 4.952 4.952 4.952 4.952 3.039 3.040 3.042 3.045 3.048 3.052 3.057 3.062 3.066 3.069 3.071 3.072 5.980 5.981 5.987 5.993 5.999 6.004 6.010 6.015 6.020 6.024 6.029 6.033
0 0.223 0.485 0.861 1.270 1.406 1.721 2.215 2.693 3.101 3.477 0 0.257 0.800 1.525 2.094 2.680 3.292 3.936 4.628 5.024 5.228 5.403 0 0.116 0.387 0.623 0.824 1.147 1.545 1.936 2.248 2.535 2.755 2.924 0 0.246 0.735 1.212 1.654 2.084 2.539 2.958 3.327 3.732 4.209 4.647
0.109 0.102 0.095 0.077 0.055 0.044 0.028 0.007 - 0.005 - 0.005 - 0.005 0.185 0.177 0.160 0.131 0.095 0.064 0.037 0.008 - 0.005 - 0.005 - 0.005 - 0.005 0.453 0.450 0.416 0.381 0.345 0.286 0.216 0.151 0.104 0.068 0.040 0.028 0.961 0.942 0.868 0.778 0.698 0.619 0.537 0.465 0.400 0.338 0.266 0.204
0 - 0.003 - 0.003 0.019 0.051 0.068 0.115 0.266 0.755 2.744 6.717 0 - 0.003 - 0.003 0.023 0.072 0.142 0.269 0.568 1.675 3.567 5.143 6.765 0 - 0.003 - 0.003 0.046 0.112 0.262 0.575 1.117 1.891 3.129 4.657 6.320 0 - 0.003 0.016 0.102 0.214 0.372 0.607 0.913 1.293 1.848 2.836 4.231
0.180 0.172 0.167 0.164 0.174 0.181 0.211 0.337 0.819 2.812 6.784 0.255 0.245 0.231 0.224 0.234 0.273 0.372 0.645 1.743 3.636 5.209 6.832 0.916 0.901 0.885 0.892 0.916 1.002 1.236 1.711 2.436 3.635 5.131 6.775 1.386 1.353 1.310 1.302 1.334 1.414 1.568 1.802 2.112 2.607 3.520 4.851
B. Rumpf et al.r Fluid Phase Equilibria 158–160 (1999) 923–932
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Table 1 Žcontinued. T ŽK.
m NH 3 Žmolrkg.
m H 2 S Žmolrkg.
10 p NH 3 ŽMPa.
10 p H 2 S ŽMPa.
10 p ŽMPa.
353.14 393.17 393.16 393.15 393.14 393.13 393.13 393.14 393.15 393.14 393.14 393.15 393.14 393.14 393.14 393.14 393.14 393.13 393.14 393.15 393.17 393.14 393.18
6.037 3.176 3.179 3.184 3.189 3.196 3.204 3.210 3.216 3.223 3.227 3.232 5.745 5.749 5.751 5.758 5.767 5.776 5.785 5.795 5.804 5.813 5.822
5.062 0 0.129 0.280 0.454 0.636 0.856 1.056 1.259 1.483 1.648 1.832 0 0.134 0.332 0.533 0.778 1.043 1.335 1.615 1.897 2.167 2.474
0.154 1.508 1.462 1.383 1.294 1.201 1.079 0.984 0.888 0.789 0.725 0.656 2.674 2.623 2.580 2.472 2.341 2.212 2.066 1.930 1.787 1.660 1.529
6.269 0 - 0.003 0.062 0.229 0.446 0.780 1.186 1.706 2.428 3.101 4.051 0 0.003 0.026 0.150 0.318 0.559 0.890 1.293 1.793 2.362 3.183
6.824 3.348 3.316 3.324 3.386 3.516 3.770 4.113 4.583 5.277 5.933 6.849 4.511 4.467 4.439 4.447 4.502 4.623 4.839 5.137 5.548 6.046 6.790
pressures and the total pressure above the solutions are reported. A model based on the Pitzer equation for the excess Gibbs energy of an aqueous electrolyte solution is used to correlate the new data. The correlation is also used to compare the new results to the limited literature data.
2. Experimental The apparatus and procedure are basically the same as used in previous investigations w1,3x, therefore only a short outline is given here. A thermostated, evacuated cell is filled with a known amount of water. After adding a known amount of ammonia, hydrogen sulfide is added stepwise. After each addition of a gas, the mixture is equilibrated and thereafter temperature, pressure and the gas phase volume are measured. A small amount of the gas phase is withdrawn from the cell and analyzed by gas chromatography. From the results for the composition of the vapour phase, the volume of the vapour phase and the total amounts of each substance charged into the cell, the overall amounts of the gases dissolved in the liquid phase are calculated. The temperature is measured by a calibrated resistance thermometer with a maximum uncertainty of "0.1 K. The pressure is determined by a pressure transducer mounted at the bottom of the cell with an uncertainty of "0.5 kPa. The composition of the vapour phase was determined by using a gas chromatograph ŽGC. equipped with a thermal conductivity cell. The GC was calibrated before and
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after each series of measurements with the pure components ammonia and hydrogen sulfide as well as with gaseous binary mixtures ammonia–water and hydrogen sulfide–water. The calibration was checked by experiments using ternary mixtures of known composition.
3. Substances Ammonia ŽG 99.999 mol%. and hydrogen sulfide ŽG 98 mol%. were purchased from MesserGriesheim, Ludwigshafen, Germany. Ammonia was used without further purification whereas hydrogen sulfide was degassed under vacuum at temperatures well below 210 K. Water was deionized and further purified by vacuum distillation.
4. Results The experimental results for the solubility of hydrogen sulfide in aqueous ammoniacal solutions are given in Table 1. Three isotherms Ž313, 353 and 393 K. were investigated. At each temperature, two overall molalities of ammonia Ž about 3 and 6 molrkg. in the liquid phase were investigated. The overall molality of hydrogen sulfide and the total pressure ranged up to about 5.4 molrkg and 0.7 MPa, respectively. In Figs. 1–3, for the two series of measurements at 393 K, the experimental results for the total pressure and the partial pressures of ammonia and hydrogen sulfide are plotted vs. the overall amount of hydrogen sulfide in the liquid phase. Figs. 1–3 show the typical behaviour observed when a sour
Fig. 1. Simultaneous solubility of NH 3 and H 2 S in water: total pressure above the solution. `,I Experimental results, this work, T s 393 K; Correlation, this work.
B. Rumpf et al.r Fluid Phase Equilibria 158–160 (1999) 923–932
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Fig. 2. Simultaneous solubility of NH 3 and H 2 S in water: partial pressure of H 2 S. `,I Experimental results, this work, T s 393 K; Correlation, this work.
gas is dissolved in an aqueous ammoniacal solution: adding hydrogen sulfide at first results in a decrease of the total pressure as ammonia is converted into ammonium ions and the sour gas is mainly dissolved in ionic form. When nearly the whole amount of ammonia is converted into ionic form, the total pressure increases rapidly. A similar behaviour is observed for the other temperatures investigated.
Fig. 3. Simultaneous solubility of NH 3 and H 2 S in water: partial pressure of NH 3. `,I Experimental results, this work, T s 393 K; Correlation, this work.
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5. Modeling Fig. 4 shows a scheme of the model applied to correlate the data. The model is a modification of the thermodynamic framework developed by Edwards et al. w4x. Due to chemical reactions in the liquid phase, NH 3 and H 2 S are not only present in neutral, but also in ionic form. To calculate the true molalities in the liquid phase, the condition for chemical equilibrium for each reaction R is applied: K R Ž T . s Ł ani i ,R
R s I . . . IV.
Ž1.
i
Together with the balance equation for the true number of moles of species i n i s n i q Ý n i ,R j R
Ž2.
R
where n i is the overall amount of species i dissolved in the liquid phase, n i,R is the stoichiometric factor of species i in reaction R and j R is the extent of reaction R, this set of equations is solved in an iterative procedure, thus yielding the true number of moles n i of each species in the liquid phase. The total pressure and the composition of the vapour phase were calculated from extended Raoults law for water pyw w
Y s w s pw
w
s w exp
ž
Õw Ž p y pws . RT
/
aw
Ž3.
and extended Henrys law for the dissolved gases pyi w iY s HiŽ,wm. Ž T . exp
ž
Õi`,w Ž p y pws . RT
/
m i g iU
i s NH 3 , H 2 S.
Ž4.
The calculation requires the knowledge of the equilibrium constants K I to K IV , the activity coefficients g iU of all species present in the liquid phase as well as the activity a w of water, Henry’s
Fig. 4. VLE and chemical reactions in the system NH 3 –H 2 S–H 2 O.
B. Rumpf et al.r Fluid Phase Equilibria 158–160 (1999) 923–932
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Ž m. constants Hi,w for NH 3 and H 2 S in pure water, the vapour pressure pws and molar volume Õw of ` water as well as information on the vapour phase nonideality and the partial molar volumes Õi,w of the dissolved gases. The equilibrium constant K I was taken from Bieling et al. w5x, K II to K IV from Edwards et al. w4x. Henry’s constants for ammonia and hydrogen sulfide were obtained from Bieling et al. w5x and Edwards et al. w4x, respectively. The vapour pressure and molar volume of water were taken from Saul and Wagner w6x. A virial equation of state truncated after the second virial coefficient was used to calculate the fugacity coefficients. Pure component second virial coefficients Bw,w and BNH 3,NH 3 were calculated from a correlation based on data recommended by Dymond and Smith w7x. Pure component virial coefficient BH 2 S,H 2 S and mixed second virial coefficients Bi, j were taken from Hayden and O’Connell w8x. The partial molar volumes of the dissolved gases were calculated as recommended by Brelvi and O’Connell w9x. To calculate the activity coefficients of both neutral and ionic species, the Pitzer w10x equation for the excess Gibbs energy of an aqueous electrolyte mixture was used:
GE RTn w M w
s f 1Ž I . q
Ý
m i m j bŽi ,0j. q biŽ1. , j f2 Ž I . q
ž
Ž i , j ./w
/
Ý
m i m j m kt i , j, k
Ž5.
Ž i , j, k ./w
where f 1 and f 2 are functions of ionic strength I and bi,Ž0.j , bi,Ž1.j and t i, j, k are binary and ternary interaction parameters. The resulting expressions for the activity of a dissolved species and of water are given elsewhere Ž cf. for example Ref. w11x.. The calculation requires the dielectric constant of pure water which was taken from Bradley and Pitzer w12x. Parameters describing interactions in NH 3 –H 2 S–H 2 O were determined as follows Ž1. For the binary systems NH 3 –H 2 O and H 2 S–H 2 O, interaction parameters were taken from Rumpf et al. w13x and Kuranov et al. w14x. Ž2. Due to the very small amounts of hydrogen and hydroxide ions, all interaction parameters involving these species were set to zero. Ž3. As the equilibrium constant for the formation of S 2y ions Žreaction III in Fig. 4. is rather small, all interaction parameters involving S 2y ions were set to zero. Ž4. As molecular dissolved NH 3 and H 2 S are simultaneously present only in very small amounts, Ž0. interaction parameter b NH and all ternary interaction parameters involving two or more molecular 3,H 2 S dissolved species were set to zero. Ž5. According to Kurz et al. w1x and Rumpf and Maurer w11x, in a solution containing dissolved gases and ionic species only certain observable interaction parameters may be determined. For the present case, the following combinations of binary and ternary interaction parameters appear Ž G stands for NH 3 or H 2 S.: BGŽ0.,NH 4 HS s bGŽ0.,NHq4 q bGŽ0.,HSy
Ž6.
GG ,NH 4 HS,NH 4 HS s t G ,NHq4 ,NHq4 q 2t G ,NHq4 ,HSyq t G ,HSy ,HSy
Ž7.
GG ,G ,NH 4 HS s t G ,G ,NHq4 q t G ,G ,HSy.
Ž8.
Ž6. The remaining parameters describe interactions between charged species. Following Pitzer w10x, it is common practice to neglect all interaction parameters involving only charged species of the same sign.
B. Rumpf et al.r Fluid Phase Equilibria 158–160 (1999) 923–932
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Table 2 Interaction parameters to describe VLE in the system NH 3 –H 2 S–H 2 O Ž313FT rK F 393., eŽT . s q1 q q2 rŽT rK.q q3 lnŽT rK. Parameter
q1
Ž0. b NH 3 ,NH 3
y0.01979 5.539Ey03 y0.26156
t NH 3 ,NH 3 ,NH 3 b HŽ0.2 S,H 2 S B HŽ0.2 S,NH 4 HS Ž0. q y b NH 4 ,HS Ž1. q y b NH 4 ,HS
t NHq4 ,NHq4 ,HSy
q2
0.05147 y0.47256 6.4495 0.03111
q3
Fitted to experimental results for
– y8.61Ey04 –
NH 3 –H 2 O
y21.2928
–
NH 3 –H 2 S–H 2 O
174.8838 y2073.6627 y10.83011
– – –
9.864 y0.1789 69.751
H 2 S–H 2 O
Ž0. Ž1. q y, b y A sensitivity study showed that amongst the remaining parameters BHŽ0.2 S,NH 4 HS , b NH NHq 4 ,HS 4 ,HS and t NHq4 ,NHq4 ,HSy are sufficient to describe the new experimental data within the experimental uncertainty. The temperature dependence of those parameters was approximated by
e Ž T . s q1 q q2r Ž TrK .
Ž e s b ,t , B . .
Ž9.
Parameters q1 and q2 were then fitted to the new results for the total pressures above NH 3 –H 2 S–H 2 O. The resulting set of parameters Žcf. Table 2. correlates the new data with an average relative deviation in the total pressure of 1.3% Žfor a detailed comparison cf. Ref. w15x. .
6. Comparison with literature data Literature data on the simultaneous solubility of ammonia and hydrogen sulfide in water are scarce. Van Krevelen et al. w16x measured the partial pressure of hydrogen sulfide above NH 3 –H 2 S–H 2 O in the temperature range from 293 to 333 K. Measured partial pressures range from about 0.04 kPa to about 42 kPa. The average absolute deviation between their data for p H 2 S and results calculated from the correlation of the new data is about 1 kPa. Terres et al. w17x reported 13 data points for the partial pressures above NH 3 –H 2 S–H 2 O at 293, 313 and 333 K. The average absolute deviation between their data for p NH 3 and p H 2 S is about 17 and 6 kPa, respectively. Miles and Wilson w18x reported 16 data points for the total and partial pressures above NH 3 –H 2 S–H 2 O at 353 and 393 K. Although the concentration range investigated by those authors Ž m NH 3,max s 22.6 molrkg. is by far larger than that of our work, the average relative deviation between their data for the total pressure and the results from the correlation is only 6%. Good agreement is also observed with the data of Ginzburg et al. w19,20x, the average relative deviation in the total pressure is only 2.4% Žfor a detailed comparison with literature data cf. Ref. w15x. .
7. Conclusions The simultaneous solubility of ammonia and hydrogen sulfide in water was measured at temperatures from 313 to 393 K and total pressures up to about 0.7 MPa. An adapted Pitzer model for the
B. Rumpf et al.r Fluid Phase Equilibria 158–160 (1999) 923–932
931
excess Gibbs energy of an aqueous electrolyte solution is used to correlate the new data. The model is capable to describe the experimental data mostly within the experimental uncertainty. A comparison with the limited literature data reveals a good agreement.
8. List of symbols ai qi Ž0. BG,MX Bi, j e f 1, f 2 GE Hi,wŽ m. I KR Mw mi mi ni ni p pi R T Õ y Greek letters b Ž0., b Ž1. w gU n i,R
activity of component i coefficients for the temperature dependence of interaction parameters ‘observable’ combination of binary interaction parameters second virial coefficient for interactions between species i and j function for the temperature dependence of interaction parameters functions in Pitzer’s equation excess Gibbs energy Henry’s constant for the solubility of gas i in water Žon molality scale. ionic strength on molality scale equilibrium constant for chemical reaction R Žon molality scale. molar mass of water Ž kgrmol. overall molality of component i true molality of component i overall number of moles of component i true number of moles of component i total pressure partial pressure of component i universal gas constant temperature partial molar volume mole fraction in vapour
t jR
binary interaction parameters in Pitzer’s equation fugacity coefficient activity coefficient normalized to infinite dilution Žon molality scale. stoichiometric coefficient of component i in reaction R Ž n i,R ) 0 for products, n i,R - 0 for educts. ternary interaction parameter in Pitzer’s equation extent of reaction R
Subscripts G i, j, k MX R w
gas G components i, j, k salt MX reaction R water
B. Rumpf et al.r Fluid Phase Equilibria 158–160 (1999) 923–932
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Superscripts s U
` X Y
saturation normalized to infinite dilution infinite dilution liquid phase gas phase
Acknowledgements Financial support of this investigation by the government of the Federal Republic of Germany ŽBMFT Grant No. 0326558 C., BASF, Ludwigshafen, Bayer, Leverkusen, Degussa, Hanau, Hoechst, Frankfurt, Linde, Dresden and Lurgi, Frankfurt is gratefully acknowledged.
References w1x w2x w3x w4x w5x w6x w7x w8x w9x w10x w11x w12x w13x w14x w15x w16x w17x w18x
F. Kurz, B. Rumpf, G. Maurer, Fluid Phase Equilibria 104 Ž1995. 261–275. B. Rumpf, F. Weyrich, G. Maurer, Fluid Phase Equilibria 83 Ž1993. 253–260. U. Goppert, G. Maurer, Fluid Phase Equilibria 41 Ž1988. 153–158. ¨ T.J. Edwards, G. Maurer, J. Newman, J.M. Prausnitz, AIChE J. 24 Ž1978. 966–976. V. Bieling, B. Rumpf, F. Strepp, G. Maurer, Fluid Phase Equilibria 53 Ž1989. 251–259. A. Saul, W. Wagner, J. Phys. Chem. Ref. Data 16 Ž1987. 893–901. J.H. Dymond, E.B. Smith, The Virial Coefficients of Pure Gases and Mixtures, Oxford Univ. Press, Oxford, 1980. J.G. Hayden, J.P. O’Connell, Ind. Eng. Chem. Proc. Des. Dev. 14 Ž1975. 209–216. S.W. Brelvi, J.P. O’Connell, AIChE J. 18 Ž1972. 1239–1243. K.S. Pitzer, J. Phys. Chem. 77 Ž1973. 268–277. B. Rumpf, G. Maurer, Ber. Bunsenges. Phys. Chem. 97 Ž1993. 85–97. D.J. Bradley, K.S. Pitzer, J. Phys. Chem. 83 Ž1979. 1599–1603. B. Rumpf, F. Weyrich, G. Maurer, Thermochim. Acta 303 Ž1997. 77–91. J. Kuranov, B. Rumpf, N.A. Smirnova, G. Maurer, Ind. Eng. Chem. Res. 35 Ž1996. 1959–1966. R. Sing, PhD. dissertation, Universitat ¨ Kaiserslautern, 1998. D.W. van Krevelen, P.J. Hoftijzer, F.J. Huntjes, Rec. Trav. Chim. 68 Ž1949. 191–216. E. Terres, W. Attig, F. Tscherter, Gas und Wasserfach 98 Ž1957. 512–516. D.H. Miles, G. Wilson, Vapor–liquid equilibrium data for design of sour water strippers, Ann. Rep. to API Ž1975. 1–25. w19x D.M. Ginzburg, N.S. Pikulina, V.P. Litvin, Zh. Prikl. Khim. 38 Ž1965. 2117–2119. w20x D.M. Ginzburg, N.S. Pikulina, V.P. Litvin, Zh. Prikl. Khim. 39 Ž1966. 2371–2373.
Simultaneous solubility of ammonia and hydrogen sulfide in water at temperatures from 313 K to 393 K ´ Perez-Salado B. Rumpf, A. Kamps, R. Sing, G. Maurer ´
)
Lehrstuhl fur ¨ Technische Thermodynamik, UniÕersitat ¨ Kaiserslautern, D-67653 Kaiserslautern, Germany Received 19 March 1998; accepted 11 September 1998
Abstract The simultaneous solubility of ammonia and hydrogen sulfide in water was measured at temperatures from 313 to 393 K and total pressures up to 0.7 MPa. An adapted Pitzer model is used to correlate the new data. Experimental results are reported and compared to the limited literature data and the correlation. q 1999 Elsevier Science B.V. All rights reserved. Keywords: Vapour–liquid equilibrium; Weak electrolytes; Ammonia; Hydrogen sulfide; Water; Experimental data; Correlation
1. Introduction The simultaneous solubility of weak electrolyte gases like ammonia, carbon dioxide, hydrogen sulfide or hydrogen cyanide in aqueous solutions is known in many applications. Typical examples are the cleaning of raw gases in power stations, the production of fertilizers or applications in the field of environmental protection. The design of separation equipment requires reliable models to calculate vapour–liquid equilibria, caloric effects and often also kinetic effects. Due to chemical reactions in the liquid phase and a strong deviation from ideality, correlating and predicting the thermodynamic properties of aqueous systems containing ammonia and sour gases is an extremely difficult task. Besides this, experimental data on the simultaneous solubility of ammonia and sour gases in aqueous phases are in most cases scarce, e.g., for systems containing hydrogen sulfide. Therefore, continuing earlier work on the solubility of ammonia and sour gases in aqueous phases Ž cf. Refs. w1,2x. , the simultaneous solubility of ammonia and hydrogen sulfide in water was measured in the temperature range from 313 to 393 K and for total pressures up to 0.7 MPa. Experimental results for the partial )
Corresponding author. Tel.: q49-631-205-2410; fax: q49-631-205-3835; e-mail: [email protected]
0378-3812r99r$ - see front matter q 1999 Elsevier Science B.V. All rights reserved. PII: S 0 3 7 8 - 3 8 1 2 Ž 9 9 . 0 0 1 1 0 - 7
B. Rumpf et al.r Fluid Phase Equilibria 158–160 (1999) 923–932
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Table 1 Experimental results for the solubility of ammonia and hydrogen sulfide in water T ŽK.
m NH 3 Žmolrkg.
m H 2 S Žmolrkg.
10 p NH 3 ŽMPa.
10 p H 2 S ŽMPa.
10 p ŽMPa.
313.19 313.17 313.17 313.17 313.16 313.17 313.15 313.15 313.15 313.16 313.15 313.19 313.17 313.18 313.17 313.17 313.18 313.14 313.15 313.15 313.14 313.14 313.12 353.17 353.14 353.14 353.13 353.13 353.12 353.13 353.14 353.13 353.13 353.14 353.13 353.13 353.14 353.14 353.13 353.13 353.14 353.14 353.14 353.13 353.14 353.14 353.14
3.002 3.002 3.003 3.005 3.007 3.008 3.009 3.011 3.011 3.011 3.011 4.938 4.938 4.940 4.942 4.945 4.947 4.949 4.952 4.952 4.952 4.952 4.952 3.039 3.040 3.042 3.045 3.048 3.052 3.057 3.062 3.066 3.069 3.071 3.072 5.980 5.981 5.987 5.993 5.999 6.004 6.010 6.015 6.020 6.024 6.029 6.033
0 0.223 0.485 0.861 1.270 1.406 1.721 2.215 2.693 3.101 3.477 0 0.257 0.800 1.525 2.094 2.680 3.292 3.936 4.628 5.024 5.228 5.403 0 0.116 0.387 0.623 0.824 1.147 1.545 1.936 2.248 2.535 2.755 2.924 0 0.246 0.735 1.212 1.654 2.084 2.539 2.958 3.327 3.732 4.209 4.647
0.109 0.102 0.095 0.077 0.055 0.044 0.028 0.007 - 0.005 - 0.005 - 0.005 0.185 0.177 0.160 0.131 0.095 0.064 0.037 0.008 - 0.005 - 0.005 - 0.005 - 0.005 0.453 0.450 0.416 0.381 0.345 0.286 0.216 0.151 0.104 0.068 0.040 0.028 0.961 0.942 0.868 0.778 0.698 0.619 0.537 0.465 0.400 0.338 0.266 0.204
0 - 0.003 - 0.003 0.019 0.051 0.068 0.115 0.266 0.755 2.744 6.717 0 - 0.003 - 0.003 0.023 0.072 0.142 0.269 0.568 1.675 3.567 5.143 6.765 0 - 0.003 - 0.003 0.046 0.112 0.262 0.575 1.117 1.891 3.129 4.657 6.320 0 - 0.003 0.016 0.102 0.214 0.372 0.607 0.913 1.293 1.848 2.836 4.231
0.180 0.172 0.167 0.164 0.174 0.181 0.211 0.337 0.819 2.812 6.784 0.255 0.245 0.231 0.224 0.234 0.273 0.372 0.645 1.743 3.636 5.209 6.832 0.916 0.901 0.885 0.892 0.916 1.002 1.236 1.711 2.436 3.635 5.131 6.775 1.386 1.353 1.310 1.302 1.334 1.414 1.568 1.802 2.112 2.607 3.520 4.851
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Table 1 Žcontinued. T ŽK.
m NH 3 Žmolrkg.
m H 2 S Žmolrkg.
10 p NH 3 ŽMPa.
10 p H 2 S ŽMPa.
10 p ŽMPa.
353.14 393.17 393.16 393.15 393.14 393.13 393.13 393.14 393.15 393.14 393.14 393.15 393.14 393.14 393.14 393.14 393.14 393.13 393.14 393.15 393.17 393.14 393.18
6.037 3.176 3.179 3.184 3.189 3.196 3.204 3.210 3.216 3.223 3.227 3.232 5.745 5.749 5.751 5.758 5.767 5.776 5.785 5.795 5.804 5.813 5.822
5.062 0 0.129 0.280 0.454 0.636 0.856 1.056 1.259 1.483 1.648 1.832 0 0.134 0.332 0.533 0.778 1.043 1.335 1.615 1.897 2.167 2.474
0.154 1.508 1.462 1.383 1.294 1.201 1.079 0.984 0.888 0.789 0.725 0.656 2.674 2.623 2.580 2.472 2.341 2.212 2.066 1.930 1.787 1.660 1.529
6.269 0 - 0.003 0.062 0.229 0.446 0.780 1.186 1.706 2.428 3.101 4.051 0 0.003 0.026 0.150 0.318 0.559 0.890 1.293 1.793 2.362 3.183
6.824 3.348 3.316 3.324 3.386 3.516 3.770 4.113 4.583 5.277 5.933 6.849 4.511 4.467 4.439 4.447 4.502 4.623 4.839 5.137 5.548 6.046 6.790
pressures and the total pressure above the solutions are reported. A model based on the Pitzer equation for the excess Gibbs energy of an aqueous electrolyte solution is used to correlate the new data. The correlation is also used to compare the new results to the limited literature data.
2. Experimental The apparatus and procedure are basically the same as used in previous investigations w1,3x, therefore only a short outline is given here. A thermostated, evacuated cell is filled with a known amount of water. After adding a known amount of ammonia, hydrogen sulfide is added stepwise. After each addition of a gas, the mixture is equilibrated and thereafter temperature, pressure and the gas phase volume are measured. A small amount of the gas phase is withdrawn from the cell and analyzed by gas chromatography. From the results for the composition of the vapour phase, the volume of the vapour phase and the total amounts of each substance charged into the cell, the overall amounts of the gases dissolved in the liquid phase are calculated. The temperature is measured by a calibrated resistance thermometer with a maximum uncertainty of "0.1 K. The pressure is determined by a pressure transducer mounted at the bottom of the cell with an uncertainty of "0.5 kPa. The composition of the vapour phase was determined by using a gas chromatograph ŽGC. equipped with a thermal conductivity cell. The GC was calibrated before and
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after each series of measurements with the pure components ammonia and hydrogen sulfide as well as with gaseous binary mixtures ammonia–water and hydrogen sulfide–water. The calibration was checked by experiments using ternary mixtures of known composition.
3. Substances Ammonia ŽG 99.999 mol%. and hydrogen sulfide ŽG 98 mol%. were purchased from MesserGriesheim, Ludwigshafen, Germany. Ammonia was used without further purification whereas hydrogen sulfide was degassed under vacuum at temperatures well below 210 K. Water was deionized and further purified by vacuum distillation.
4. Results The experimental results for the solubility of hydrogen sulfide in aqueous ammoniacal solutions are given in Table 1. Three isotherms Ž313, 353 and 393 K. were investigated. At each temperature, two overall molalities of ammonia Ž about 3 and 6 molrkg. in the liquid phase were investigated. The overall molality of hydrogen sulfide and the total pressure ranged up to about 5.4 molrkg and 0.7 MPa, respectively. In Figs. 1–3, for the two series of measurements at 393 K, the experimental results for the total pressure and the partial pressures of ammonia and hydrogen sulfide are plotted vs. the overall amount of hydrogen sulfide in the liquid phase. Figs. 1–3 show the typical behaviour observed when a sour
Fig. 1. Simultaneous solubility of NH 3 and H 2 S in water: total pressure above the solution. `,I Experimental results, this work, T s 393 K; Correlation, this work.
B. Rumpf et al.r Fluid Phase Equilibria 158–160 (1999) 923–932
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Fig. 2. Simultaneous solubility of NH 3 and H 2 S in water: partial pressure of H 2 S. `,I Experimental results, this work, T s 393 K; Correlation, this work.
gas is dissolved in an aqueous ammoniacal solution: adding hydrogen sulfide at first results in a decrease of the total pressure as ammonia is converted into ammonium ions and the sour gas is mainly dissolved in ionic form. When nearly the whole amount of ammonia is converted into ionic form, the total pressure increases rapidly. A similar behaviour is observed for the other temperatures investigated.
Fig. 3. Simultaneous solubility of NH 3 and H 2 S in water: partial pressure of NH 3. `,I Experimental results, this work, T s 393 K; Correlation, this work.
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5. Modeling Fig. 4 shows a scheme of the model applied to correlate the data. The model is a modification of the thermodynamic framework developed by Edwards et al. w4x. Due to chemical reactions in the liquid phase, NH 3 and H 2 S are not only present in neutral, but also in ionic form. To calculate the true molalities in the liquid phase, the condition for chemical equilibrium for each reaction R is applied: K R Ž T . s Ł ani i ,R
R s I . . . IV.
Ž1.
i
Together with the balance equation for the true number of moles of species i n i s n i q Ý n i ,R j R
Ž2.
R
where n i is the overall amount of species i dissolved in the liquid phase, n i,R is the stoichiometric factor of species i in reaction R and j R is the extent of reaction R, this set of equations is solved in an iterative procedure, thus yielding the true number of moles n i of each species in the liquid phase. The total pressure and the composition of the vapour phase were calculated from extended Raoults law for water pyw w
Y s w s pw
w
s w exp
ž
Õw Ž p y pws . RT
/
aw
Ž3.
and extended Henrys law for the dissolved gases pyi w iY s HiŽ,wm. Ž T . exp
ž
Õi`,w Ž p y pws . RT
/
m i g iU
i s NH 3 , H 2 S.
Ž4.
The calculation requires the knowledge of the equilibrium constants K I to K IV , the activity coefficients g iU of all species present in the liquid phase as well as the activity a w of water, Henry’s
Fig. 4. VLE and chemical reactions in the system NH 3 –H 2 S–H 2 O.
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Ž m. constants Hi,w for NH 3 and H 2 S in pure water, the vapour pressure pws and molar volume Õw of ` water as well as information on the vapour phase nonideality and the partial molar volumes Õi,w of the dissolved gases. The equilibrium constant K I was taken from Bieling et al. w5x, K II to K IV from Edwards et al. w4x. Henry’s constants for ammonia and hydrogen sulfide were obtained from Bieling et al. w5x and Edwards et al. w4x, respectively. The vapour pressure and molar volume of water were taken from Saul and Wagner w6x. A virial equation of state truncated after the second virial coefficient was used to calculate the fugacity coefficients. Pure component second virial coefficients Bw,w and BNH 3,NH 3 were calculated from a correlation based on data recommended by Dymond and Smith w7x. Pure component virial coefficient BH 2 S,H 2 S and mixed second virial coefficients Bi, j were taken from Hayden and O’Connell w8x. The partial molar volumes of the dissolved gases were calculated as recommended by Brelvi and O’Connell w9x. To calculate the activity coefficients of both neutral and ionic species, the Pitzer w10x equation for the excess Gibbs energy of an aqueous electrolyte mixture was used:
GE RTn w M w
s f 1Ž I . q
Ý
m i m j bŽi ,0j. q biŽ1. , j f2 Ž I . q
ž
Ž i , j ./w
/
Ý
m i m j m kt i , j, k
Ž5.
Ž i , j, k ./w
where f 1 and f 2 are functions of ionic strength I and bi,Ž0.j , bi,Ž1.j and t i, j, k are binary and ternary interaction parameters. The resulting expressions for the activity of a dissolved species and of water are given elsewhere Ž cf. for example Ref. w11x.. The calculation requires the dielectric constant of pure water which was taken from Bradley and Pitzer w12x. Parameters describing interactions in NH 3 –H 2 S–H 2 O were determined as follows Ž1. For the binary systems NH 3 –H 2 O and H 2 S–H 2 O, interaction parameters were taken from Rumpf et al. w13x and Kuranov et al. w14x. Ž2. Due to the very small amounts of hydrogen and hydroxide ions, all interaction parameters involving these species were set to zero. Ž3. As the equilibrium constant for the formation of S 2y ions Žreaction III in Fig. 4. is rather small, all interaction parameters involving S 2y ions were set to zero. Ž4. As molecular dissolved NH 3 and H 2 S are simultaneously present only in very small amounts, Ž0. interaction parameter b NH and all ternary interaction parameters involving two or more molecular 3,H 2 S dissolved species were set to zero. Ž5. According to Kurz et al. w1x and Rumpf and Maurer w11x, in a solution containing dissolved gases and ionic species only certain observable interaction parameters may be determined. For the present case, the following combinations of binary and ternary interaction parameters appear Ž G stands for NH 3 or H 2 S.: BGŽ0.,NH 4 HS s bGŽ0.,NHq4 q bGŽ0.,HSy
Ž6.
GG ,NH 4 HS,NH 4 HS s t G ,NHq4 ,NHq4 q 2t G ,NHq4 ,HSyq t G ,HSy ,HSy
Ž7.
GG ,G ,NH 4 HS s t G ,G ,NHq4 q t G ,G ,HSy.
Ž8.
Ž6. The remaining parameters describe interactions between charged species. Following Pitzer w10x, it is common practice to neglect all interaction parameters involving only charged species of the same sign.
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Table 2 Interaction parameters to describe VLE in the system NH 3 –H 2 S–H 2 O Ž313FT rK F 393., eŽT . s q1 q q2 rŽT rK.q q3 lnŽT rK. Parameter
q1
Ž0. b NH 3 ,NH 3
y0.01979 5.539Ey03 y0.26156
t NH 3 ,NH 3 ,NH 3 b HŽ0.2 S,H 2 S B HŽ0.2 S,NH 4 HS Ž0. q y b NH 4 ,HS Ž1. q y b NH 4 ,HS
t NHq4 ,NHq4 ,HSy
q2
0.05147 y0.47256 6.4495 0.03111
q3
Fitted to experimental results for
– y8.61Ey04 –
NH 3 –H 2 O
y21.2928
–
NH 3 –H 2 S–H 2 O
174.8838 y2073.6627 y10.83011
– – –
9.864 y0.1789 69.751
H 2 S–H 2 O
Ž0. Ž1. q y, b y A sensitivity study showed that amongst the remaining parameters BHŽ0.2 S,NH 4 HS , b NH NHq 4 ,HS 4 ,HS and t NHq4 ,NHq4 ,HSy are sufficient to describe the new experimental data within the experimental uncertainty. The temperature dependence of those parameters was approximated by
e Ž T . s q1 q q2r Ž TrK .
Ž e s b ,t , B . .
Ž9.
Parameters q1 and q2 were then fitted to the new results for the total pressures above NH 3 –H 2 S–H 2 O. The resulting set of parameters Žcf. Table 2. correlates the new data with an average relative deviation in the total pressure of 1.3% Žfor a detailed comparison cf. Ref. w15x. .
6. Comparison with literature data Literature data on the simultaneous solubility of ammonia and hydrogen sulfide in water are scarce. Van Krevelen et al. w16x measured the partial pressure of hydrogen sulfide above NH 3 –H 2 S–H 2 O in the temperature range from 293 to 333 K. Measured partial pressures range from about 0.04 kPa to about 42 kPa. The average absolute deviation between their data for p H 2 S and results calculated from the correlation of the new data is about 1 kPa. Terres et al. w17x reported 13 data points for the partial pressures above NH 3 –H 2 S–H 2 O at 293, 313 and 333 K. The average absolute deviation between their data for p NH 3 and p H 2 S is about 17 and 6 kPa, respectively. Miles and Wilson w18x reported 16 data points for the total and partial pressures above NH 3 –H 2 S–H 2 O at 353 and 393 K. Although the concentration range investigated by those authors Ž m NH 3,max s 22.6 molrkg. is by far larger than that of our work, the average relative deviation between their data for the total pressure and the results from the correlation is only 6%. Good agreement is also observed with the data of Ginzburg et al. w19,20x, the average relative deviation in the total pressure is only 2.4% Žfor a detailed comparison with literature data cf. Ref. w15x. .
7. Conclusions The simultaneous solubility of ammonia and hydrogen sulfide in water was measured at temperatures from 313 to 393 K and total pressures up to about 0.7 MPa. An adapted Pitzer model for the
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excess Gibbs energy of an aqueous electrolyte solution is used to correlate the new data. The model is capable to describe the experimental data mostly within the experimental uncertainty. A comparison with the limited literature data reveals a good agreement.
8. List of symbols ai qi Ž0. BG,MX Bi, j e f 1, f 2 GE Hi,wŽ m. I KR Mw mi mi ni ni p pi R T Õ y Greek letters b Ž0., b Ž1. w gU n i,R
activity of component i coefficients for the temperature dependence of interaction parameters ‘observable’ combination of binary interaction parameters second virial coefficient for interactions between species i and j function for the temperature dependence of interaction parameters functions in Pitzer’s equation excess Gibbs energy Henry’s constant for the solubility of gas i in water Žon molality scale. ionic strength on molality scale equilibrium constant for chemical reaction R Žon molality scale. molar mass of water Ž kgrmol. overall molality of component i true molality of component i overall number of moles of component i true number of moles of component i total pressure partial pressure of component i universal gas constant temperature partial molar volume mole fraction in vapour
t jR
binary interaction parameters in Pitzer’s equation fugacity coefficient activity coefficient normalized to infinite dilution Žon molality scale. stoichiometric coefficient of component i in reaction R Ž n i,R ) 0 for products, n i,R - 0 for educts. ternary interaction parameter in Pitzer’s equation extent of reaction R
Subscripts G i, j, k MX R w
gas G components i, j, k salt MX reaction R water
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Superscripts s U
` X Y
saturation normalized to infinite dilution infinite dilution liquid phase gas phase
Acknowledgements Financial support of this investigation by the government of the Federal Republic of Germany ŽBMFT Grant No. 0326558 C., BASF, Ludwigshafen, Bayer, Leverkusen, Degussa, Hanau, Hoechst, Frankfurt, Linde, Dresden and Lurgi, Frankfurt is gratefully acknowledged.
References w1x w2x w3x w4x w5x w6x w7x w8x w9x w10x w11x w12x w13x w14x w15x w16x w17x w18x
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