Ind. Eng. Chem. Res. 1998, 37, 3519-3523
3519
Solubility of Methane in Aqueous Solutions of 2-(2-Aminoethoxy)ethanol Fang-Yuan Jou,† John J. Carroll,‡ Frederick D. Otto,† and Alan E. Mather*,† Department of Chemical and Materials Engineering, University of Alberta, Edmonton, Alberta, Canada T6G 2G6, and Gas Liquids Engineering, #300, 2749 39th Avenue NE, Calgary, Alberta, Canada T1Y 4T8
In this study experimental data are presented for the solubility of methane in 3.0 and 6.0 kmol/ m3 (30.5 and 59.5 wt %) solutions of 2-(2-aminoethoxy)ethanol (which is also called diglycolamine). Temperatures in this study ranged from 25 to 125 °C and pressures to 13 MPa. The data were incorporated into a rigorous thermodynamic model that has been applied to other similar systems. The new data were then compared with the data available in the literature. In pure water and low amine concentrations it is demonstrated that the isobaric solubility exhibits a minimum. However, for high amine concentration there is a somewhat surprising resultsthe isobaric solubility is an increasing function of temperature over the range of temperature studied here. Introduction Aqueous solutions of alkanolamines are commonly used in the hydrocarbon processing industry to strip acid gases (carbon dioxide and hydrogen sulfide) from raw hydrocarbon streams. One of the amines used in this process is 2-(2-aminoethoxy)ethanol (commonly called diglycolamine or DGA, which is a registered trademark of Huntsman Corp.). The process for the removal of acid gases takes place in two steps. In the first the raw gas, which is called sour if it contains hydrogen sulfide, is contacted with a lean solvent. This produces an enriched solvent and a sweet gas. In the second stage the enriched solvent is regenerated and a sour gas stream, made up largely of hydrogen sulfide and carbon dioxide, is produced. The regenerated solvent is returned to the first step, and the acid gas product is sent for further treating. If there is a large volume of acid gas, it is typically sent to a Claus plant where the hydrogen sulfide is converted to elemental sulfur. Small volumes of acid gas are often flared where the hydrogen sulfide is converted to sulfur oxides and emitted to the atmosphere. Because of environmental concerns alternative schemes are being sought to dispose of the low volume of acid gas produced at these plants. One such scheme is the injection into a suitable underground formation. This paper is a part of an ongoing project to provide experimental data useful for the design of plants for treating these hydrocarbon gases and liquids. Although data have been published for the solubility of alkanes in aqueous solutions of amines, the only data available for the solubility of methane in a solution of DGA are the few points by Dingman (1986). He measured the solubility of methane in a 50 wt % (5.0 kmol/m3) DGA solution at 22.8, 24.4, and 76.7 °C (73, 76, and 170 °F) at pressures to 20 MPa (3000 psia). He measured only a single point at the first two temperatures and six points at the highest temperature. In this work we present additional data for this system. Two amine concentrations, 3 and 6 kmol/m3, * To whom correspondence should be addressed. † University of Alberta. ‡ Gas Liquids Engineering.
were studied here. The model of Carroll and Mather (1997) will be used to correlate the new data. The model will also be used to make a comparison with the data collected by Dingman (1986). Experimental Work The experimental apparatus and procedure have been used in this laboratory for many years (see, for example, Jou et al., 1985). Thus, they will only be described here briefly. The equipment consists of two major components: an equilibrium cell and a recirculation pump. The cell is a Jerguson liquid level gauge with a cylindrical reservoir attached to the top. The purpose of the reservoir is to ensure that there is sufficient mass in the system. The recirculation pump is a magnetically driven piston pump similar to the one devised by Ruska et al. (1970). The temperature of the contents of the cell was measured using an iron-constantan thermocouple. The thermocouple was calibrated against a platinum resistance thermometer and was found to be accurate to within (0.1 °C over the range of temperature investigated. The pressure was measured using two digital Heise pressure gauges. The gauges were calibrated against a dead-weight gauge and were found to be accurate to within 0.1% of full scale. The water used in the experiments was laboratory distilled. The DGA was obtained from Aldrich, had a stated purity of 99%, and was used without further purification. The amine solutions were made up on the laboratory bench and thus were 3.0 and 6.0 kmol/m3 at room conditions. The methane was UHP grade from Linde with a minimum purity of 99.97%. A small sample of the methane was injected into a gas chromatograph, and no significant impurities were detected. To perform a run, about 50 mL of the amine solution was charged to the cell. Methane was added in an amount determined by observation of the pressure. The pump was started, and the fluids were left in the cell for a time sufficiently long for them to equilibrate. The vapor phase was analyzed using gas chromatography. The column used was 3 m long by 6.35 mm o.d. and was packed with Porapak Q. The liquid phase was sampled into a 50 cm3 sample bomb. During the liquid-
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3520 Ind. Eng. Chem. Res., Vol. 37, No. 8, 1998 Table 1. Solubility of Methane in a 3 kmol/m3 Aqueous Solution of 2-(2-Aminoethoxy)ethanol
Table 2. Solubility of Methane in a 6 kmol/m3 Aqueous Solution of 2-(2-Aminoethoxy)ethanol
solubility
solubility
temp (°C)
pressure (MPa)
mole fraction × 1000
mol of CH4/100 kg of solvent
temp (°C)
pressure (MPa)
mole fraction × 1000
mol of CH4/100 kg of solvent
25.0
13.59 9.65 6.43 3.14 0.819 0.284 0.096 13.42 9.93 6.79 2.92 0.942 0.296 13.36 9.60 6.47 3.39 0.920 0.309 13.08 9.36 6.63 3.31 1.017 0.344 12.67 9.62 6.58 3.64 1.073 0.455
2.67 2.16 1.57 0.902 0.254 0.0830 0.0290 2.57 2.02 1.47 0.697 0.267 0.0799 2.56 2.07 1.46 0.852 0.250 0.0736 2.73 2.06 1.54 0.835 0.261 0.0720 2.91 2.36 1.65 0.949 0.248 0.0804
11.1 8.98 6.51 3.74 1.05 0.343 0.122 10.7 8.40 6.11 2.89 1.11 0.331 10.6 8.61 6.04 3.53 1.03 0.305 11.3 8.56 6.37 3.46 1.08 0.298 12.1 9.78 6.84 3.94 1.03 0.333
25.0
12.93 9.39 6.36 2.95 0.699 0.286 0.096 13.49 9.15 6.10 2.82 0.882 0.294 13.19 9.75 6.33 2.80 0.917 0.305 12.86 9.78 6.56 3.31 1.091 0.365 13.15 9.77 6.88 3.52 1.129 0.485
3.75 3.06 2.25 1.21 0.323 0.132 0.0418 4.55 3.23 2.53 1.30 0.408 0.156 4.92 3.98 2.78 1.37 0.484 0.0157 5.97 4.57 3.35 2.00 0.597 0.181 6.77 5.51 4.02 2.22 0.685 0.224
10.5 8.57 6.29 3.37 0.902 0.369 0.117 12.8 9.05 7.08 3.64 1.14 0.435 13.8 11.2 7.77 3.84 1.35 0.440 16.8 12.8 9.40 5.58 1.67 0.506 19.0 15.5 11.3 6.21 1.91 0.626
50.0
75.0
100.0
125.0
phase sampling, the pressure in the cell was held constant by the careful addition of more gas. The bomb was weighed before and after sampling to determine the mass of sample. It was then attached to a mercuryfilled buret, and the pressure of the sample was brought to atmospheric. The gas evolved from the liquid was collected in a calibrated 50-mL buret. The moles collected were calculated from volumetric data after subtracting the vapor pressure of the solution. In addition, the solubility of the gas remaining in the solution at atmospheric pressure was taken into account. Results The solubilities of methane in 3 kmol/m3 DGA are presented in Table 1 and those for 6 kmol/m3 in Table 2. For convenience the solubilities are listed both in mole fraction and in moles of solute per 100 kg of solvent. The data were correlated using the model of Carroll and Mather (1997). Basically the model uses Henry’s law for the aqueous phase and the Peng-Robinson (1976) equation of state for the vapor phase. The model used in this work is virtually identical with that presented by Carroll and Mather (1997). Furthermore, most of the required parameters were taken from that paper. Those details will not be repeated here. The only new parameters calculated from the data in this work are the salting-in ratios and the Setchenow coefficients. The salting-in ratios are defined as the mole fraction solubility in the amine solution divided by the mole fraction solubility in pure water. Because there are no experimental data for the solubilities of methane at the
50.0
75.0
100.0
125.0
exact conditions of the data presented in this work, the model of Carroll and Mather (1997) was used to calculate the solubilities. In an earlier work (Carroll and Mather, 1997), it was demonstrated that the effect of concentration was accurately modeled using the Setchenow relation for the solubility of ethane in aqueous solutions of triethanolamine. This equation will be applied to the new data. The Setchenow equation is
ln S ) kCa
(1)
where S is the salting-in ratio, k is the Setchenow coefficient, and Ca is the amine concentration. The amine concentration must be expressed as molarity, kmol/m3. Table 3 summarizes the salting-in and Setchenow coefficients derived from the data presented in this paper. The ranges given for the salting-in ratios are the standard deviations. For comparison, those derived from the data of Dingman (1986) are also included in this table. It is clear from this table that the Setchenow coefficients are an increasing function of the temperature. They have been correlated with the following equation:
kC1-DGA ) -2.466 + 1.398 × 10-2T - 1.816 × 10-5T2 (2) where T is the absolute temperature. Figure 1 is a plot of the Setchenow coefficients as a function of temperature. From this plot it can be seen that eq 2 is a good fit of the data. However, the correlation predicts a maximum in the Setchenow
Ind. Eng. Chem. Res., Vol. 37, No. 8, 1998 3521 Table 3. Salting-in and Setchenow Coefficients for Methane in Aqueous Solutions of 2-(2-Aminoethoxy)ethanol Setchenow coefficient temp (°C)
DGA conc
22.8 24.4 25.0 25.0 50.0 50.0 75.0 75.0 76.7 100.0 100.0 125.0 125.0 a
(kmol/m3)
no. of points
salting-in ratio
from data
from eq 2
ref
1 1 6 7 6 6 6 6 6 6 6 6 6
1.45 1.48 1.27 ( 0.03 1.87 ( 0.05 1.56 ( 0.07 2.82 ( 0.19 1.77 ( 0.06 3.49 ( 0.22 2.64 ( 0.06 1.85 ( 0.08 4.11 ( 0.30 1.89 ( 0.19 4.33 ( 0.29
0.074 0.078 0.080 0.104 0.148 0.173 0.190 0.208 0.194 0.205 0.236 0.212 0.244
0.081 0.086 0.088 0.088 0.155 0.155 0.200 0.200 0.202 0.222 0.222 0.211 0.211
a a b b b b b b a b b b b
5.0 5.0 3.0 6.0 3.0 6.0 3.0 6.0 5.0 3.0 6.0 3.0 6.0
Dingman (1986) (see Carroll and Mather, 1997). b This work.
Figure 1. Setchenow coefficients for the solubility of methane in aqueous solutions of 2-(2-aminoethoxy)ethanol.
Figure 2. Solubility of methane in water and aqueous solutions of 2-(2-aminoethoxy)ethanol at 50 °C.
Table 4. Summary of the Errors in the Calculated Solubility of Methane in 2-(2-Aminoethoxy)ethanol salting-in ratio from Table 3
3 kmol/m3 6 kmol/m3 overall
salting-in ratio from eq 2
average error (%)
average absolute error (%)
average error (%)
average absolute error (%)
-1.57 -1.30 -1.44
3.45 3.85 3.65
-3.12 7.16 2.25
4.67 9.58 7.13
a Error ) (measured - calculated)/measured × 100%. Absolute error ) |measured - calculated|/measured × 100%.
coefficient that is probably not justified by the data. Thus, it would be unwise to extrapolate eq 2 beyond the range of temperature examined here. In addition, it can be seen from Table 3 that the Setchenow coefficients from this work appear to be a function of the composition. The Setchenow coefficients for the 6 kmol/m3 solution are larger than those for the 3 kmol/m3 solutions. On the other hand, the values from the solubility measurements of Dingman (1986), which are at 5 kmol/m3, are in good agreement with those for 3 kmol/m3 from this work. Table 4 lists the errors in the calculated solubilities using the salting-in ratios given in Table 3 and those calculated using eq 2. Using the optimum salting-in ratios the average absolute error is less than 4%. On the other hand, the salting-in ratios from eq 2 tend to overpredict the solubility for the 3 kmol/m3 solution and underestimate it for the 6 kmol/m3 solution. This is the
Figure 3. Solubility of methane in aqueous solutions of 2-(2aminoethoxy)ethanol at 75 and 76.7 °C.
composition effect mentioned earlier. However, in both cases, the errors are not unreasonable. The solubilities of methane in water and the aqueous solutions of DGA at 50 °C are shown in Figure 2. This graph is typical of the other isotherms, which were not shown. First, this plot demonstrates that the solubility in the amine solution is greater than that in water. Second, it shows graphically the error that one might expect when using the approximate Setchenow coefficients from eq 2. Figure 3 shows the calculated solubility of methane in solutions of DGA. The 3 and 6 kmol/m3 concentrations are at 75 °C, and the 5 kmol/m3 concentration is
3522 Ind. Eng. Chem. Res., Vol. 37, No. 8, 1998
Figure 4. Solubility of methane in water as a function of temperature at 13, 9.5, and 6.5 MPa.
Figure 5. Solubility of methane in a 3 kmol/m3 aqueous solution of 2-(2-aminoethoxy)ethanol as a function of temperature at 13, 9.5, and 6.5 MPa.
at 76.7 °C. The curves on the figure were generated using the salting-in ratios from eq 2. This graph provides a comparison between the data of Dingman (1986) and those obtained in this study. It is difficult to make direct comparisons between the data obtained in this work and those of Dingman (1986) because the two sets are for a different amine concentration and are for different temperatures. However, this plot shows a good degree of agreement between the two works. The Effect of Temperature on the Solubility Conventional wisdom is that solubility tends to decrease with increasing temperature. Although this is probably a reasonable rule of thumb, it is not always the case. For example, it is well-known that the isobaric solubility of methane in water at high pressure exhibits a minimum (Culberson and McKetta, 1951). Figure 4 shows the solubility of methane in water for three isobars as predicted by the model of Carroll and Mather (1997) showing these minima. The minima are a function of the pressure: at 13.5 MPa it is approximately 80 °C, at 9.5 MPa it is 83 °C, and finally at 6.5 MPa it is 87 °C. The exact location of these minima is a little difficult to determine because the curves are shallow through the minimal region. The presence of the amine changes the location of these minima. Figure 5 shows the same three isobars except these are for a 3 kmol/m3 DGA solution. The curves on this plot were generated using Setchenow coefficients from eq 2. Also plotted on this figure are the experimental data from this work. Deviations between the curves and the data can be explained by two effects. First, as explained earlier, the model, especially based on eq 2, is not an exact fit of the data. Second, the data are not exactly at the pressures of the calculated values. However, both the data and the calculations show the minima. For this amine concentration, the minima are at significantly lower temperatures. The minima, as predicted by the model, are as follows: 13 MPa, 46 °C; 9.5 MPa, 50 °C; 6.5 MPa, 55 °C. In addition, the minima in this case are not as shallow as those for methane in pure water. For an amine concentration of 6 kmol/m3 the effect is even more surprising. Figure 6 shows the same isobars for this amine concentration. The minima have disap-
Figure 6. Solubility of methane in a 6 kmol/m3 aqueous solution of 2-(2-aminoethoxy)ethanol as a function of temperature at 13, 9.5, and 6.5 MPa.
peared, and the solubility is an increasing function of the temperature over the entire range studied here. Both the data and the predictions exhibit this behavior. It should be noted that the minima may be at lower temperatures than those examined here. As was noted for the previous figure, the seemingly large deviations between the correlation and the data are easily explained.
Conclusions Additional data for the solubility of alkanes in gastreating solvents have been presented. In this work the solubility of methane in DGA is reported. These data were modeled using a rigorous thermodynamic model. The new data compared well with the data of Dingman (1986), the only other data available for this system. Both the experimental data and the model presented in this work show that, for high amine concentration, the solubility is an increasing function of temperature. This is a somewhat surprising result and is counter to the conventional wisdom that solubility tends to decrease with increasing temperature.
Ind. Eng. Chem. Res., Vol. 37, No. 8, 1998 3523
Literature Cited Carroll, J. J.; Mather, A. E. A Model for the Solubility of Light Hydrocarbons in Water and Aqueous Solutions of Alkanolamines. Chem. Eng. Sci. 1997, 52, 545. Culberson, O. L.; McKetta, J. J. Phase Equilibria in HydrocarbonWater Systems. IIIsThe Solubility of Methane in Water at Pressures to 10, 000 PSIA. Petrol. Trans. A. I. M. E. 1951, 192, 223. Dingman, J. C. Don’t Blame Hydrocarbon Solubility for Entrainment Problems in Amine Treating Systems. Annual AIChE Meeting, Miami Beach, FL, 1986; Paper 140e. Jou, F.-Y.; Otto, F. D.; Mather, A. E. Equilbria of H2S and CO2 in Triethanolamine Solutions. Can. J. Chem. Eng. 1985, 63, 122.
Peng, D.-Y.; Robinson, D. B. A New Two-Constant Equation of State. Ind. Eng. Chem. Fundam. 1976, 15, 59. Ruska, W. E. A.; Hurt, L. J.; Kobayashi, R. Circulating Pump for High Pressure and -200 to +400 C Application. Rev. Sci. Instrum. 1970, 41, 1444.
Received for review December 15, 1997 Revised manuscript received May 19, 1998 Accepted May 26, 1998 IE970912I
