New VaporâLiquidâLiquid Equilibrium Solubility Data for iso-Butane, n

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New Vapor−Liquid−Liquid Equilibrium Solubility Data for iso-Butane, n‑Butane, n‑Pentane, and n‑Hexane in Alkanolamine Aqueous Solutions Salim Mokraoui,*,† Mohamed Kamel Hadj-Kali,‡ Alain Valtz,§ and Dominique Richon∥,⊥ †

Sustainable Energy Technology Center, and ‡Chemical Engineering Department, King Saud University, P.O Box 800, 11421 Riyadh, Saudi Arabia § Mines ParisParis Tech, CTP, 35, Rue Saint Honoré, 77305 Fontainebleau, France ∥ Thermodynamics Research Unit, School of Chemical Engineering, University of KwaZulu-Natal, Howard College Campus, King George V Avenue, Durban 4041, South Africa ⊥ Department of Biotechnology and Chemical Technology, Aalto University, School of Science and Technology, P.O. Box 16100, 00076 Aalto, Finland ABSTRACT: Recently, we have published the solubility data of ethane and propane in aqueous diethanolamine (DEA) and methyldiethanolamine (MDEA) solutions at different concentrations (Mokraoui et al., J. Chem. Eng. Data 2013, 58, 2100−2109). In this work, new measurements are provided for longer aliphatic light hydrocarbons, namely, isobutane, n-butane, n-pentane, and n-hexane. The same concentrations of alkanolamines (wDEA = 0.35, wMDEA = 0.25, and wMDEA = 0.50) have been considered in three phases coexisting conditions. MDEA has shown more affinity than DEA for a given weight-based concentration. Within the temperature window explored in this paper (293−353) K, the solubility increases with the amine concentration but it decreases with an increase in the hydrocarbon molecular size. Furthermore, the higher is the carbon atom number of the hydrocarbon, the smaller is the solubility dependence with temperature.

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INTRODUCTION Natural gases, liquefied petroleum gases (LPGs), and refinery process streams contain variable amounts of acid gases that must be removed. The traditional treatment for carbon dioxide and sulfur compounds removal is the absorption/stripping process using aqueous alkanolamine solutions. The main advantage of using alkanolamines is their highly selective absorption of acid gases from the gas or liquid hydrocarbon streams. Two of the alkanolamines, methyldiethanolamine (MDEA) and diethanolamine (DEA), have received much attention because of their low heat of absorption and low corrosiveness. In addition, as stated by Sada et al.,1 the solubility of any hydrocarbon in aqueous amine solutions remains higher than its solubility in pure water. In most LPG production facilities, alkanolamine solutions are particularly convenient for use in sweetening both gas and liquid hydrocarbon streams since the two absorbers in the LPG sweetening process share a common amine regenerator,2 as highlighted in Figure 1. In the petroleum industry, the mutual solubility of hydrocarbons and amines has been investigated in several gas mixtures for gas treating conditions in order to estimate the equilibrium loss and to diagnose plant-operating problems that could result in excess losses. In particular, for a better design of LPG sweetening plants, the solubility of the hydrocarbons part of the LPG streams, such as propane, butane, pentane, and hexane, in © 2014 American Chemical Society

amine solutions should be considered in order to minimize the hydrocarbon losses and optimize the overall process. Data needed for purification process of LPG often involve three phase equilibrium systems consisting of an aqueous alkanolamine liquid phase and two distinct hydrocarbon rich phases (liquid and vapor). In this framework, the GPA research project 021 has been developed for the investigation of the mutual solubility of the most common liquid treating solvents, DEA (w = 0.35) and MDEA (w = 0.25 and w = 0.5) in liquid ethane, propane, isobutane, n-butane, n-pentane, and n-hexane at different temperatures.3 Although the process of LPGs sweetening using amines is well established, the solubility data of hydrocarbons in these amines are still scarce, especially, in vapor−liquid−liquid (VLL) conditions. The methane, ethane, and propane solubilities in amine have been widely studied. However, our literature survey indicates a limited amount of data for n-butane and n-pentane solubility in amines.4,5 The i-butane and n-hexane solubility with amines has only been available to a limited audience.3,6 Jou et al.5 have investigated the aqueous MDEA (M = 3) + n-butane system. They reported experimental results involving the vapor−liquid (VL), liquid−liquid (LL), and VLL equilibrium Received: February 7, 2014 Accepted: April 10, 2014 Published: April 21, 2014 1673

dx.doi.org/10.1021/je500130j | J. Chem. Eng. Data 2014, 59, 1673−1684

Journal of Chemical & Engineering Data

Article

Figure 1. Typical LPG sweetening process.

(E). As with propane in LLE conditions,7 the solubility of liquid n-butane increases with temperature but an opposite behavior is observed in VLE conditions. Similarly, Critchfield et al.4 have studied the solubility of n-butane and n-pentane in different amines solutions and in the LLE conditions. Specifically, they measured the solubility of n-butane and n-pentane in DEA and MDEA at 313.15 K and 1.7 MPa. They have shown that the solubility of C3−C5 paraffins increases with the molarity of the amines and is an increasing function of amines strength according to the following order: MEA (monoethanol amine), DEA, DGA (diglycol amine), MDEA, and DIPA (diisopropanol amine). In this paper, new solubility data of isobutane, n-butane, n-pentane, and n-hexane in aqueous methyl-diethanolamine (MDEA) and diethanolamine (DEA) solutions are presented at different temperatures and pressures. The data were performed at vapor−liquid−liquid equilibrium (VLLE) conditions. Two weight-based concentrations of aqueous MDEA (w = 0.25 and w = 0.50) and one aqueous solution of DEA (w = 0.35) were considered. These concentration types are commonly used in the gas processing industry.8 On the basis of the experimental results and the thermodynamics analysis developed subsequently, the three-phase pressure, the solubility data, and the Henry’s constants are correlated, for each system, as a function of temperature.

balance (0.1 mg precision). The apparatus and the experimental procedure adopted in this work are those related to the “staticanalytic” method with vapor and/or liquid-phase samplings. Figure 2 shows the flow diagram of the apparatus used. The experimental protocol is identical to the one developed by Laugier and Richon9 and is thoroughly described in our previous works.10,11 Samplings are completed through two online ROLSI samplers, one for each phase (liquid and vapor). ROLSI is a trademark of Armines (France), protected by several patents (PCT patent 2004/090508, PCT patent 2000/011462, EPO patent EP 1105722). The phase equilibrium is achieved inside a cylindrical cell of about 28 cm3, made of a sapphire tube. The cell is immersed inside a constant−temperature liquid bath (Ultra-Kryomat Lauda) which controls and maintains the desired temperature within ± 0.01 K. Two 100 Ω platinum resistance thermometer devices (Pt 100), fixed at two different locations (vapor and liquid phases), are used after careful calibrations against a 25 Ω reference platinum resistance thermometer (Tinsley Precision Instruments). Pressures are measured using a Druck pressure transducer (0 to 6) MPa, kept at constant temperature (higher than the maximum temperature of the study) and against a dead-weight pressure balance (Desgranges & Huot 5202S, CP 0.3-40 MPa, Aubervilliers, France). The resulting uncertainty is estimated to be within ± 0.02 K for temperature measurements and within ± 0.3 kPa for pressure measurements. Analyses were carried out using a GC (Varian model CP-3800) equipped with two detectors in series: (i) thermal conductivity detector (TCD) used to quantify water and amine contents and (ii) flame ionization detector (FID) for isobutane, n-butane, n-pentane, and n-hexane detection. Both detectors were calibrated by injecting known masses of each compound. Experimental uncertainties on mole numbers of each compound, resulting from calibration, are estimated to be within ± 2 % for water, ± 2.5 % for DEA and MDEA, ± 2 % for isobutane, ± 1.3 % for n-butane, ± 1.5 % for n-pentane and ± 1 % for n-hexane. The mole fractions uncertainties were deduced by combining the mole numbers uncertainty (systematic error part) and the uncertainty due to the repeatability of the sequential analysis of the aqueous phase (random error part). The combined

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EXPERIMENTAL SECTION Materials. All the chemicals were used as received without any further purification, except for careful degassing of liquid components. Isobutane and n-butane were purchased from Air Liquide (France) with traces of impurities (H2O, CO2, O2, N2, and CnHm) less than 500 ppm mol. n-Pentane is from Fluka (Germany) with a certified GC purity greater than 99.5 %. n-Hexane is from Merck (Germany) with a certified gas chromatograph (GC) purity greater than 99 %, while MDEA and DEA are from Aldrich (Germany) with purity greater than 99 vol %. The carrier gas (helium), used for GC analyses, was supplied by Air Liquide; it is pure grade with only traces of water (3 ppm) and of hydrocarbons (0.5 ppm). Deionized water was used after careful degassing to prepare amine aqueous solutions. Apparatus and Experimental Procedure. The water− amine solutions were prepared gravimetrically using an analytical 1674



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Figure 2. Flow diagram of the apparatus. C, carrier gas; d. a. s., data acquisition system; DA, degassed amine solution; DDD, displacement digital display; DT, displacement transducer; EC, equilibrium cell; FV, feeding valve; HC, hydrocarbon cylinder; LB, liquid bath; LS, ROLSI liquid sampler; MS, magnetic stirring; PN, pressurized nitrogen; PP, platinum probe; PT, pressure transducer; SM, sampler monitoring; ST, sapphire tube; TC, thermal container; Th, thermocouple; TR, temperature regulator; VS, ROLSI vapor sampler; VSS, variable speed stirring; VVCA, variable volume cell for amine solution; Vi, shut-off valve; VP, vacuum pump.

(L) and vapor (V) phases are considered as hydrocarbon-rich phases. Consequently, to model the phase equilibrium of these systems, we have made the following hypothesis: (i) The amount of the solvent in the hydrocarbon liquid phase is ignored. Thus, the water and alkanolamine are assumed to escape into the vapor phase according to the Raoult’s law. (ii) The effect of the total pressure on the molar volume of each hydrocarbon is negligible. Therefore, the Poynting correction factor approaches unity. The first assumption is supported by the experimental work of Carroll et al.14 studying the system propane−aqueous MDEA. The second assumption is preferred due to the difficulty of getting the molar volume variation with pressure. Following the above assumptions, the fugacities of the hydrocarbon component in each of the three phases are given by the following equations:

uncertainties’ results are presented in the results/discussion section. The separation of hydrocarbons, water, and amine compounds is accomplished thanks to a gas chromatograph circuit analogous to that presented in our previous work.12 Hydrocarbon and water are separated through column A2 (Hayesep T 100/120 mesh column; silcosteel tube, length, 1.5 m; diameter, 1/8 in.) then sent to TCD 1 (for water) and FID (for hydrocarbons), whereas amine is retained by column A1 (Porapak Q 100/120 mesh column; silcosteel tube, length, 0.06 m; diameter, 1/8 in.) and directly sent to TCD 2 after the commuting valve was rotated.

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THERMODYNAMIC TREATMENT Equilibrium Equations. At constant temperature and pressure, the phase equilibrium of a multicomponent system is expressed by the equality between the chemical potential (alternatively the fugacity) of each component within each phase.13 In practice, the fugacity calculation is commonly performed either by the homogeneous or the heterogeneous approach: (1) In the homogeneous approach, (also called the φ−φ approach) both the liquid and the vapor phases are described by the same cubic equation of state such as Peng− Robinson (PR) or Soave−Redlich−Kwong (SRK). (2) In the heterogeneous approach, (known as the γ−φ approach), the vapor phase is described by an equation of state while the nonideality of the liquid phase is expressed by an activity coefficient model such as NRTL, UNIQUAC, or UNIFAC. The systems studied in this work are in liquid − liquid − vapor equilibrium. The first liquid phase (denoted A) is mainly composed of the aqueous amine solution with the measured amount of hydrocarbon diluted on it. While, the second liquid

A A A sat A f hcA = xhc γhc(x ̅ , T )φhcsatPhc = xhc /(T , P)

(1)

L L L sat sat f hc = xhc γhc(x ̅ L , T )φhcsatPhc ≅ φhcsatPhc

(2)

f hcV = yhc PφhcV (y ̅ , T , P)

(3)

where, x ̅ A , x ̅ L , y,̅ are, respectively, the compositions of the aqueous, liquid hydrocarbon, and vapor phases. In eq 1, the Henry’s law constant “/ ” has been used to represent the solubility data. Because of the hydrocarbon rich phase is mostly composed of hydrocarbon, the composition and activity coefficient, in eq 2, are both set to unity. The equality condition of the fugacities, imposed by the equilibrium between the three phases, leads to the following two equations: 1675



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temperature, the VLLE occurs between the dew and the bubble pressures of the hydrocarbon phase. Experimentally, these pressures are almost indistinguishable and considered as one point pressure named the three-phase vapor pressure which is substantially close to the vapor pressure of the pure hydrocarbon. As a consequence, we considered only the variation of this pressure with the temperature, ignoring the overall composition effect. To demonstrate this statement, measurements of the three phase vapor pressure have been performed for each system and then compared to the pure hydrocarbon vapor pressures (Figure 3).

Equilibrium between the aqueous and the hydrocarbon-rich liquid phases: A A A L L xhc γhc(x ̅ , T ) = xhc γhc(x ̅ L , T ) ≅ 1

(4)

Equilibrium between the aqueous and the vapor phases: A xhc /(T , P) = yhc PφhcV (y ̅ , T , P)

(5)

Hence, eq 5 has been used to calculate the Henry’s law constant of the hydrocarbon in aqueous alkanolamine solution. The vapor phase’s fugacity coefficient, φVhc, is estimated by means of the Peng−Robinson15 equation of state with the binary interaction parameters neglected. The mole fraction of the hydrocarbons in the vapor phase is approached as

yhc = (P − Psolv)/P

(6)

where Psolv is the partial pressure of the solvent, composed of water and alkanolamine, and approximated by Raoult’s law through the following equation: Psolv = x wAPwsat + xaAPasat

(7)

where (xAw) and (xAa ) represent the mole fractions of water and alkanolamine in the aqueous phase, respectively, approached by the initial molar concentrations of the aqueous solvent. Thermodynamic Properties. The thermodynamic properties such as the critical temperature (Tc), critical pressure (Pc), and acentric factor (ω) of each component are provided13 in Table 1. The vapor pressure is calculated according to the formula suggested by Reid et al.16

Figure 3. Three phase vapor pressure as a function of temperature: ⧫, isobutane−aqueous MDEA (w = 0.5); ▲, n-butane−aqueous MDEA (w = 0.5); ●, n-pentane−aqueous MDEA (w = 0.5); and ■, n-hexane− aqueous MDEA (w = 0.5). The dashed lines represent the pure compounds vapor pressures.

Table 1. Critical Temperatures, Critical Pressures, And Acentric Factors of the Pure Compounds17 isobutane n-butane n-pentane n-hexane water DEA MDEA

CAS No.

Tc/K

Pc/MPa

ω

75-28-5 106-97-8 109-66-0 110-54-3 7732-18-5 111-42-2 105-59-9

407.8 425.1 469.7 507.6 647.1 736.6 675.0

3.640 3.796 3.370 3.025 22.05 4.270 3.829

0.184 0.200 0.252 0.301 0.345 0.953 1.165

ln(P sat /Pa) = A + B /T + C ln(T ) + DT E

T /K

(8)

The values of the parameters A, B, C, D, and E are extracted from Daubert et al.17 and presented in Table 2.

Figure 4. VLLE experimental solubility of ◆, isobutane; ▲, n-butane; ●, n-pentane; and ■, n-hexane in MDEA aqueous solutions (w = 0.25).

Table 2. Parameters17 of Equation 8 isobutane n-butane n-pentane n-hexane water DEA MDEA

A

B

C

D

E

108.43 66.343 78.741 104.65 73.649 106.38 253.07

−5039.9 −4363.2 −5420.3 −6995.5 −7258.2 −13714 −18378

−15.012 −7.046 −8.8253 −12.702 −7.3037 −11.06 −33.972

2.27·10−02 9.45·10−06 9.62·10−06 1.24·10−05 4.17·10−06 3.26·10−18 2.33·10−05

1 2 2 2 2 6 2

Tables 3 to 6 report the measured and calculated three-phase vapor pressures for isobutane, n-butane, n-pentane, and n-hexane. For n-hexane, the measurements were restricted to the system involving aqueous MDEA (w = 0.5). The calculated data were correlated using eq 9: ln(P3/MPa) = A iP + BiP /T

APi ,

T /K

(9)

BPi

Parameters (given in Table 7) were obtained by minimizing the sum of the quadratic difference between the calculated and experimental pressures. For each hydrocarbon, a unique correlation is obtained by gathering the experimental data of all aqueous amine solutions. The maximum relative deviations are 0.7 % for isobutane, 2.6 % for n-butane, 1.9 % for n-pentane,

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RESULTS AND DISCUSSION The Three-Phase Vapor Pressure. Theoretically, for the light hydrocarbon-aqueous amine ternary systems and for each 1676



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Table 3. Experimental and Calculated Three-Phase Vapor Pressures for Isobutane−Aqueous Amine Systemsa aqueous MDEA (w = 0.25) T

a

Pexp

Pcal

aqueous MDEA (w = 0.5) RD

T

Pexp

Pcal

aqueous DEA (w = 0.35) RD

T

Pexp

Pcal

RD

K

MPa

MPa

%

K

MPa

MPa

%

K

MPa

MPa

%

293.36 295.40 297.35 299.32 301.25 303.36 305.28 307.31 309.36 311.31 313.28 315.27 317.28 319.26 321.23 323.20 325.29 327.18 329.20 331.19 333.21 335.17 337.16 339.22 341.13 343.18 345.15 347.18 349.13 351.07 353.13

0.307 0.326 0.346 0.366 0.387 0.411 0.434 0.459 0.486 0.512 0.540 0.569 0.599 0.631 0.663 0.697 0.734 0.769 0.808 0.847 0.889 0.931 0.975 1.022 1.067 1.117 1.167 1.221 1.274 1.328 1.388

0.307 0.326 0.346 0.366 0.387 0.411 0.434 0.459 0.486 0.512 0.540 0.569 0.600 0.631 0.664 0.698 0.735 0.770 0.808 0.848 0.889 0.931 0.975 1.021 1.066 1.116 1.166 1.219 1.271 1.324 1.383

0.00 0.02 −0.02 −0.03 −0.05 −0.08 −0.06 −0.06 −0.10 −0.09 −0.11 −0.11 −0.13 −0.12 −0.11 −0.11 −0.10 −0.08 −0.07 −0.06 −0.02 0.01 0.02 0.05 0.07 0.07 0.11 0.19 0.24 0.27 0.35

293.53 295.44 297.34 299.28 301.22 303.25 305.37 307.39 309.36 311.29 313.29 315.27 317.28 319.25 321.22 323.21 325.27 327.18 329.18 331.20 333.19 335.17 337.16 339.24 341.15 343.20 345.17 347.12 349.12 351.05 353.17

0.311 0.329 0.348 0.368 0.389 0.412 0.437 0.462 0.487 0.513 0.541 0.570 0.601 0.632 0.664 0.698 0.735 0.770 0.808 0.847 0.889 0.931 0.974 1.022 1.066 1.116 1.166 1.218 1.272 1.325 1.386

0.308 0.327 0.346 0.366 0.387 0.410 0.435 0.460 0.486 0.512 0.540 0.569 0.600 0.631 0.664 0.698 0.734 0.770 0.808 0.848 0.889 0.931 0.975 1.022 1.067 1.117 1.166 1.217 1.270 1.324 1.384

0.75 0.70 0.62 0.53 0.46 0.40 0.37 0.32 0.26 0.25 0.20 0.17 0.14 0.11 0.06 0.02 0.02 −0.01 −0.02 −0.04 −0.01 −0.02 −0.03 −0.03 −0.05 −0.11 −0.06 0.10 0.11 0.12 0.15

293.33 295.36 297.22 299.27 301.21 303.31 305.26 307.26 309.26 311.25 313.28 315.26 317.29 319.37 321.24 323.22 325.28 327.22 329.19 331.20 333.21 335.16 337.15 339.22 341.14 343.19 345.16 347.19 349.11 351.06 353.16

0.306 0.325 0.344 0.365 0.386 0.410 0.433 0.458 0.484 0.510 0.539 0.568 0.599 0.632 0.662 0.696 0.733 0.768 0.806 0.846 0.887 0.929 0.973 1.020 1.065 1.116 1.166 1.220 1.271 1.326 1.387

0.307 0.326 0.344 0.366 0.387 0.411 0.434 0.459 0.485 0.512 0.540 0.569 0.600 0.633 0.664 0.698 0.735 0.770 0.808 0.848 0.889 0.931 0.974 1.021 1.067 1.117 1.166 1.219 1.270 1.324 1.384

−0.22 −0.22 −0.22 −0.25 −0.23 −0.25 −0.28 −0.27 −0.25 −0.29 −0.30 −0.29 −0.26 −0.24 −0.26 −0.27 −0.26 −0.27 −0.23 −0.22 −0.18 −0.16 −0.13 −0.12 −0.12 −0.04 −0.01 0.08 0.09 0.15 0.19

Standard uncertainties u are u(T) = 0.02 K, u(P) = 0.3 kPa.

Figure 5. VLLE experimental solubility of ◆, isobutane; ▲, n-butane; ●, n-pentane; and ■, n-hexane in MDEA aqueous solutions (w = 0.50).

Figure 6. VLLE experimental solubility of ◆, isobutane; ▲, n-butane; ●, n-pentane; and ■, n-hexane in DEA aqueous solutions (w = 0.35).

and 11.3 % for n-hexane. The highest deviations for n-hexane are observed at the lowest pressures, below 0.04 MPa. Solubility. The VLLE solubility measurements are presented in Table 8 (for iso- and n-butane) and Table 9 (for n-pentane and n-hexane). Each measurement was repeated between five and ten times. The experimental uncertainties due to the repeatability of the sequential analyses are included. It is worth pointing out that

for a convenient sampling with the ROLSI samplers during the experiments, the pressure in the equilibrium cell was always kept greater than 0.3 MPa. Therefore, for low pressure systems, especially n-pentane and n-hexane, helium was added in order to increase the pressure to about 0.5 MPa. The presence of helium has a negligible effect on the solubility within the experimental uncertainty, as it was verified with the n-butane−amine systems. 1677



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Table 4. Experimental and Calculated Three-Phase Vapor Pressures for n-Butane−Aqueous Amine Systemsa aqueous MDEA (w = 0.25) T

a

Pexp

Pcal


T

Pexp

Pcal

aqueous DEA (w = 0.35) T

RD

Pexp

Pcal

RD

K

MPa

MPa

%

K

MPa

MPa

%

K

MPa

MPa

%

293.49 295.53 297.44 299.42 301.38 303.42 305.35 307.37 309.34 311.33 313.30 315.28 317.28 319.26 321.21 323.20 325.27 327.17 329.21 331.20 333.22 335.17 337.16 339.22 341.13 343.14 345.11 347.17 349.11 351.05 353.14

0.213 0.227 0.241 0.257 0.272 0.290 0.307 0.326 0.345 0.366 0.387 0.409 0.432 0.456 0.481 0.508 0.536 0.564 0.594 0.625 0.658 0.690 0.725 0.763 0.799 0.838 0.878 0.921 0.963 1.006 1.056

0.215 0.229 0.243 0.259 0.275 0.292 0.310 0.329 0.348 0.368 0.390 0.412 0.435 0.459 0.484 0.510 0.539 0.566 0.596 0.627 0.660 0.692 0.727 0.764 0.800 0.838 0.878 0.920 0.962 1.005 1.053

−1.03 −0.99 −0.96 −0.93 −0.88 −0.85 −0.78 −0.76 −0.75 −0.71 −0.69 −0.66 −0.63 −0.60 −0.57 −0.53 −0.50 −0.45 −0.44 −0.41 −0.34 −0.28 −0.22 −0.19 −0.13 −0.07 0.01 0.06 0.12 0.13 0.28

293.30 295.37 297.32 299.26 301.21 303.26 303.27 305.19 307.23 309.20 309.20 311.18 313.26 315.25 317.24 319.24 321.36 323.33 325.38 327.27 329.26 331.26 333.27 335.23 337.20 343.17 345.13 347.18 349.12 351.06 353.13

0.212 0.227 0.241 0.256 0.271 0.289 0.289 0.306 0.325 0.344 0.344 0.364 0.386 0.409 0.431 0.456 0.483 0.509 0.537 0.564 0.594 0.625 0.657 0.690 0.724 0.835 0.874 0.917 0.959 1.003 1.051

0.214 0.228 0.243 0.258 0.273 0.291 0.291 0.308 0.327 0.347 0.347 0.367 0.389 0.411 0.435 0.459 0.486 0.512 0.540 0.567 0.597 0.628 0.661 0.693 0.728 0.839 0.878 0.921 0.962 1.005 1.053

−0.65 −0.69 −0.68 −0.71 −0.73 −0.72 −0.73 −0.72 −0.71 −0.73 −0.73 −0.74 −0.73 −0.68 −0.72 −0.71 −0.70 −0.68 −0.65 −0.61 −0.57 −0.54 −0.52 −0.50 −0.47 −0.47 −0.47 −0.41 −0.33 −0.23 −0.20

293.35 295.37 297.34 299.31 301.24 303.28 305.25 307.29 309.24 311.24 313.21 315.22 317.21 319.32 321.28 323.23 325.30 327.22 329.22 331.21 333.23 335.19 337.16 339.22 341.14 343.19 345.12 347.18 349.14 351.06 353.14

0.220 0.234 0.248 0.263 0.279 0.296 0.314 0.333 0.352 0.372 0.393 0.416 0.439 0.464 0.489 0.514 0.543 0.570 0.600 0.631 0.664 0.696 0.731 0.768 0.804 0.844 0.883 0.926 0.968 1.012 1.060

0.214 0.228 0.243 0.258 0.274 0.291 0.309 0.328 0.347 0.367 0.389 0.411 0.434 0.460 0.485 0.511 0.539 0.567 0.597 0.627 0.660 0.693 0.727 0.764 0.800 0.839 0.878 0.921 0.963 1.005 1.053

2.57 2.34 2.18 2.01 1.88 1.75 1.64 1.51 1.41 1.33 1.24 1.15 1.11 0.87 0.79 0.67 0.61 0.55 0.54 0.52 0.57 0.54 0.55 0.54 0.54 0.53 0.54 0.58 0.59 0.64 0.69


Figure 8. VLLE experimental solubility of n-butane in water and aqueous amine solutions: ◇, pure water, Mokraoui et al.;11 △, w = 0.25 of MDEA; □, w = 0.5 of MDEA; ○, w = 0.35 of DEA; ●, w = 0.35 of MDEA, Jou et al.5 The continuous lines through experimental data are drawn for visual clarity.

Figure 7. VLLE experimental solubility of isobutane in water and aqueous amine solutions: ◇, pure water, Mokraoui et al.;11 △, w = 0.25 of MDEA; □, w = 0.5 of MDEA; ○, w = 0.35 of DEA. The continuous lines through experimental data are drawn for visual clarity.

Figures 4 to 6 (−ln (xAhc) versus 1/T) show a comparison of the hydrocarbons solubilities in each aqueous amine solution. As it can be seen, the logarithm of solubility varies linearly with the inverse of temperature over the explored temperature range of each system. Hence, the solubility data were correlated using the following equation:

A ln(xhc ) = A ix + Bix /T

T /K

(10)

The fitted parameters Axi , Bxi including the temperature validity ranges are reported in Table 10. These parameters are obtained by minimizing the sum of quadratic difference between calculated 1678



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Table 5. Experimental and Calculated Three-Phase Vapor Pressures for n-Pentane−Aqueous Amine Systemsa aqueous MDEA (w = 0.25) T

a

Pexp

Pcal


T

Pexp

Pcal

aqueous DEA (w = 0.35) RD

T

Pexp

Pcal

RD

K

MPa

MPa

%

K

MPa

MPa

%

K

MPa

MPa

%

293.31 295.32 297.30 299.26 301.20 303.29 305.22 307.26 309.20 311.20 313.19 315.18 317.18 319.38 321.34 323.31 325.37 327.27 329.26 331.25 333.26 335.21 337.19 339.23 341.13 343.15 345.10 347.15 349.10 351.03 353.08

0.062 0.067 0.072 0.077 0.082 0.089 0.095 0.102 0.109 0.116 0.124 0.133 0.142 0.153 0.162 0.173 0.184 0.195 0.208 0.221 0.235 0.249 0.264 0.280 0.296 0.313 0.331 0.350 0.370 0.390 0.412

0.061 0.066 0.071 0.076 0.082 0.088 0.095 0.102 0.109 0.116 0.124 0.133 0.142 0.152 0.162 0.173 0.184 0.195 0.208 0.220 0.234 0.248 0.263 0.279 0.294 0.311 0.328 0.348 0.366 0.386 0.407

1.95 1.72 1.34 0.90 0.72 0.33 0.31 0.22 0.03 0.01 0.05 −0.02 −0.01 0.05 0.05 0.04 0.05 0.04 0.11 0.18 0.24 0.30 0.36 0.44 0.57 0.66 0.77 0.83 0.92 1.04 1.10

293.32 295.35 297.28 299.23 301.18 303.23 305.16 307.22 309.19 311.17 313.28 315.19 317.18 319.18 321.16 323.21 325.32 327.23 329.21 331.17 333.21 335.15 337.12 339.12 341.10 343.09 345.04 347.09 349.03 350.96 352.97

0.061 0.066 0.071 0.076 0.082 0.088 0.094 0.101 0.108 0.115 0.124 0.132 0.141 0.150 0.160 0.171 0.183 0.193 0.205 0.218 0.232 0.246 0.261 0.276 0.292 0.310 0.327 0.346 0.366 0.385 0.407

0.061 0.066 0.071 0.076 0.082 0.088 0.094 0.101 0.109 0.116 0.125 0.133 0.142 0.151 0.161 0.172 0.184 0.195 0.207 0.220 0.234 0.248 0.262 0.278 0.294 0.311 0.328 0.347 0.366 0.385 0.406

0.72 0.34 0.17 −0.11 −0.17 −0.40 −0.36 −0.44 −0.46 −0.63 −0.87 −0.80 −0.72 −0.74 −0.71 −0.75 −0.80 −0.86 −0.99 −0.78 −0.71 −0.57 −0.55 −0.68 −0.46 −0.32 −0.21 −0.17 0.01 0.08 0.20

293.36 295.38 297.34 299.29 301.23 303.29 305.23 307.28 309.25 311.22 313.20 315.19 317.22 319.19 321.16 323.17 325.33 327.25 329.23 331.22 333.24 335.23 337.21 339.26 341.19 343.18 345.14 347.30 349.18 351.08 353.12

0.061 0.066 0.071 0.076 0.082 0.088 0.094 0.102 0.109 0.116 0.124 0.132 0.141 0.151 0.161 0.171 0.183 0.195 0.207 0.220 0.233 0.248 0.262 0.279 0.295 0.312 0.329 0.350 0.368 0.389 0.410

0.061 0.066 0.071 0.076 0.082 0.088 0.095 0.102 0.109 0.116 0.124 0.133 0.142 0.151 0.161 0.172 0.184 0.195 0.207 0.220 0.234 0.248 0.263 0.279 0.294 0.311 0.329 0.349 0.367 0.386 0.408

0.39 0.06 0.15 0.09 0.03 −0.06 −0.09 −0.08 −0.24 −0.31 −0.35 −0.40 −0.51 −0.51 −0.50 −0.57 −0.47 −0.40 −0.31 −0.34 −0.26 −0.14 −0.15 −0.07 0.02 0.13 0.19 0.31 0.32 0.63 0.61


and experimental solubilities. The maximum relative deviations are less than 6.6 % for isobutane, 5.4 % for n-butane, 8 % for n-pentane and 7.8 % for n-hexane, by comparing all amine solutions. Furthermore, Figures 7 to 10 highlight the temperature influence as well as the amine type and concentration effects on the solubility for each hydrocarbon. For comparison, we have included the solubility data in pure water taken from our previous work11 for all hydrocarbons and those from Jou et al.5 for the n-butane−MDEA aqueous solution in particular (Figure 8). As compared with the values in the literature,1,18 Figures 7 to 10 confirm the solubility increase with the amine concentration. Regarding the amine type effect, Figure 8 shows clearly that MDEA7 has greater affinity for n-butane than DEA for the same concentration (w = 0.35). Similar observation has been mentioned by Critchfield et al.4 Since the pressure increases systematically with temperature to maintain the VLLE conditions, the temperature effect on the solubility is inevitably combined with that of pressure. Therefore, for all hydrocarbons and within the temperature range of this study, the solubility increases with temperature. As mentioned by Carroll at al.,18 Tsonopoulos,19 and Mokraoui et al.,12 for hydrocarbon−water binary systems and hydrocarbon−water− amine ternary systems, increasing the temperature (at constant pressure) tends first to decrease the solubility because of the predominance condensation effects over the hydrophobic

Figure 9. VLLE experimental solubility of n-pentane in water and aqueous amine solutions: ◇, pure water, Mokraoui et al.;11 △, w = 0.25 of MDEA; □, w = 0.5 of MDEA; ○, w = 0.35 of DEA. The continuous lines through experimental data are drawn for visual clarity.

interactions. Then, the solubility crosses a minimum when the two effects are equal. As the temperature continues to increase, the hydrophobic interaction effect predominates and the solubility increases. In our case, the solubility is always increasing, favored by the pressure increase. The minimum may exist at a lower temperature range; however, additional experiments are necessary to strengthen this assumption. 1679



Article

Table 6. Experimental and Calculated Three-Phase Vapor Pressures for n-Hexane−Aqueous MDEA Solution (w = 0.5)a aqueous MDEA (w = 0.5)

a

T

Pexp

Pcal

RD

K

MPa

MPa

%

293.24 295.53 297.45 299.39 301.35 303.39 305.33 307.37 309.33 311.31 313.27 315.26 317.26 319.26 321.19 323.18 325.29 327.19 329.19 331.18 333.22 335.19 337.15 339.22 341.15 343.15 345.14 347.16 349.13 351.04 353.17

0.024 0.027 0.030 0.032 0.034 0.036 0.040 0.042 0.045 0.046 0.050 0.053 0.056 0.060 0.065 0.070 0.075 0.079 0.085 0.091 0.098 0.104 0.112 0.120 0.127 0.135 0.143 0.153 0.162 0.172 0.184

0.023 0.025 0.027 0.029 0.031 0.034 0.037 0.040 0.043 0.046 0.049 0.053 0.057 0.061 0.065 0.070 0.075 0.080 0.086 0.092 0.098 0.104 0.111 0.118 0.126 0.134 0.142 0.151 0.160 0.169 0.180

6.44 7.04 11.22 8.74 7.02 7.03 8.45 6.39 4.82 1.36 0.56 0.06 −0.94 −1.43 −0.79 −0.68 −0.99 −1.25 −1.39 −1.12 −0.12 0.05 0.51 0.98 1.17 1.15 0.81 1.14 1.49 1.78 2.24

Figure 10. VLLE experimental solubility of n-hexane in water and aqueous amine solutions: ◇, pure water, Mokraoui et al.;11 △, w = 0.25 of MDEA; □, w = 0.5 of MDEA; ○, w = 0.35 of DEA. The continuous lines through experimental data are drawn for visual clarity.

Figure 11. VLLE solubility of n-alkanes in aqueous MDEA (w = 0.5) as a function of carbon atom number and temperature: ●, 298 K; □, 313 K; ▲, 323 K; ◇, 333 K. For Nc = 3, the data are taken from our previous work12


Table 7. Values of Parameters for Calculating Three-Phase Vapor Pressures (eq 9) i

APi

BPi

ARD/%

MRD/%

T range/K

isobutane n-butane n-pentane n-hexane

7.72 7.87 8.42 22.27

−2609.93 −2761.54 −3290.23 −3590.05

0.17 0.73 0.43 2.88

0.74 2.55 1.91 11.31

293 to 353 293 to 353 293 to 353 293 to 353

Another general trend for n-alkanes is that the solubility decreases with an increase in the hydrocarbon’s carbon atoms number due to the molecular size difference. This is observed in Figure 11 which represents the solubility of propane, n-butane, npentane, and n-hexane in aqueous MDEA solution (w = 0.5) at different temperatures. Moreover, from the same Figure, it should be pointed out that the higher is the carbon atom number the smaller is the temperature influence. On the other hand, it is worth noting that the solubility of isobutane and n-butane are quite different (Figures 7 and 8). Indeed, because of the isomerism structure of isobutane, its density is higher than that of n-butane. So, the smaller size of isobutane leads it to be more soluble than n-butane in aqueous amine solutions. Henry’s Constant. The Henry’s law is a limiting law expressing the ratio between the fugacity of a solute at infinite

Figure 12. Henry’s constant of ◆, isobutane; ▲, n-butane; ●, npentane; and ■, n-hexane in aqueous MDEA (w = 0.5) as a function of temperature. The error bars correspond to the relative uncertainties.

dilution in a solvent to its mole fraction. It is used for estimating the solubility of a gas in a liquid solvent through some assumptions. As developed earlier, eq 5 is one form of this law retained for our calculations. From the experimental solubility data, Henry’s constants / expressed in MPa were calculated for each of the hydrocarbons−aqueous amine system. Table 11 presents the calculated Henry’s constants for all of the hydrocarbons in the different aqueous amine solutions including the corresponding estimated uncertainties. The general trend 1680



Article

Table 8. VLLE Experimental and Calculated Solubility Data of Isobutane and n-Butane in MDEA and DEA Aqueous Solutionsa n-butane

isobutane T/K

Pexp/MPa

xexp·104

xcal·104

RD/%

298.29 303.31 313.24 323.20 333.20 343.16

0.355 0.410 0.539 0.696 0.888 1.116

1.80 1.87 2.06 2.22 2.45 2.83

1.77 1.87 2.07 2.28 2.50 2.73

1.84 0.16 −0.55 −2.88 −2.23 3.49

298.23 303.30 313.29 323.22 333.24 343.15

0.354 0.409 0.539 0.695 0.887 1.116

5.25 5.81 6.96 8.62 10.83 14.56

5.03 5.72 7.26 9.07 11.20 13.63

4.14 1.62 −4.26 −5.18 −3.69 6.62

298.28 303.28 313.26 323.20 333.28 343.18

0.354 0.409 0.539 0.696 0.889 1.116

2.15 2.28 2.59 2.96 3.30 3.74

2.13 2.29 2.61 2.95 3.33 3.71

0.83 −0.27 −0.79 0.17 −0.77 0.82

δx·106 c

T/K

Aqueous MDEA (w = 0.25) 4 298.27 3 303.27 2 313.22 2 323.32 4 333.25 5 343.17 Aqueous MDEA (w = 0.50) 5 298.28 3 303.29 3 313.19 5 323.16 19 333.22 21 343.16 Aqueous DEA (w = 0.35) 2 298.26 3 303.31 3 313.30 4 323.27 4 333.21 5 343.16

Pexp/MPa

xexp·104

xcal·104

RD/%

δx·106 c

0.494b 0.502b 0.523b 0.509 0.658 0.839

0.95 1.02 1.10 1.19 1.35 1.53

0.94 1.00 1.11 1.24 1.36 1.49

0.84 1.92 −1.33 −3.82 −0.65 2.87

1 1 1 1 1 2

0.509b 0.521b 0.530b 0.506 0.655 0.835

2.99 3.10 3.82 4.70 5.89 7.31

2.83 3.17 3.95 4.85 5.90 7.08

5.42 −2.37 −3.32 −3.21 −0.14 3.21

1 2 5 6 6 13

0.506b 0.512b 0.508b 0.507 0.664 0.843

1.21 1.22 1.29 1.39 1.56 1.81

1.16 1.22 1.34 1.46 1.59 1.71

4.34 0.26 −3.65 −5.06 −1.68 5.26

2 1 1 1 1 2

Standard (combined) uncertainties are u(T) = 0.02 K, u(P) = 0.3 kPa, uc(x) = 1·10−5 for isobutane and uc(x) = 6·10−6 for n-butane. bTotal pressures Pexp are adjusted with helium in order to allow convenient sampling. cδx: deviation observed through repeatability tests of the sequential analyses of the aqueous phase. a

Table 9. VLLE Experimental and Calculated Solubility Data of n-Pentane and n-Hexane in MDEA and DEA Aqueous Solutionsa n-pentane T/K

Pexp/MPa

xexp·10

298.26 303.32 313.20 323.28 333.23 343.17

0.512 0.493 0.508 0.511 0.503 0.508

298.41 303.38 313.26 323.19 333.21 343.17 298.28 303.29 313.13 323.18 333.22 343.32

n-hexane

xcal·10

RD/%

0.27 0.30 0.34 0.40 0.46 0.56

0.27 0.29 0.35 0.41 0.47 0.55

0.51 1.66 −1.48 −1.49 −1.96 2.66

0.514 0.512 0.508 0.503 0.510 0.514

1.35 1.40 1.57 1.84 2.50 3.16

1.24 1.38 1.70 2.06 2.47 2.93

8.02 1.21 −8.02 −11.73 1.29 7.35

0.493 0.503 0.505 0.502 0.527 0.509

0.32 0.35 0.41 0.46 0.52 0.62

0.32 0.34 0.40 0.47 0.53 0.61

−0.77 0.97 1.27 −1.00 −2.42 1.87

4

4

δx·10

7

T/K

Aqueous MDEA (w = 0.25) 5 5 298.17 3 313.13 5 333.14 5 353.18 5 Aqueous MDEA (w = 0.50) 11 20 298.21 18 313.17 22 333.24 41 353.17 75 Aqueous DEA (w = 0.35) 4 2 298.16 8 313.16 4 333.15 7 353.15 7

Pexp/MPa

xexp·104

xcal·104

RD/%

δx·107

0.501 0.514 0.503 0.513

0.09 0.12 0.18 0.29

0.09 0.13 0.19 0.28

6.88 −7.78 −5.50 5.42

2 2 3 3

0.503 0.505 0.503 0.518

0.44 0.66 1.15 1.98

0.42 0.67 1.18 1.93

2.97 −2.95 −2.62 2.43

2 3 9 20

0.503 0.500 0.501 0.502

0.09 0.14 0.20 0.34

0.09 0.13 0.22 0.33

0.48 2.04 −5.76 2.98

3 4 5 6

a Standard (combined) uncertainties are u(T) = 0.02 K, u(P) = 0.3 kPa, uc(x) = 3·10−6 for n-pentane and uc(x) = 1·10−6 for n-hexane. The total pressures Pexp for these systems are adjusted with helium in order to allow convenient sampling. δx: deviation observed through repeatability tests of the sequential analyses of the aqueous phase.

due to its high solubility excepting the aqueous MDEA (w = 0.5) system; this may be because the increase in solubility for n-pentane and n-hexane is more pronounced for higher amine concentrations. Henry’s constant is also an increasing function

that can be drawn from this table is that the Henry’s constant decreases with an increase in the amine concentration and the number of carbon atom of the normal alkanes, especially at high temperatures. Isobutane has the lowest Henry’s constant values 1681



Article

Table 10. Values of Parameters for Calculating Solubility Data in VLLE (eq 10) Axi

solvent

Bxi

isobutane −993.54 1.86 n-pentane −1629.92 1.63 isobutane −2270.39 4.25 n-pentane −1962.25 6.27 isobutane −1262.47 0.61 n-pentane −1482.86 1.38

−5.310

aqueous MDEA (w = 0.25)

−5.067 aqueous MDEA (w = 0.50)

0.018 −2.418 −4.221

aqueous DEA (w = 0.35)

−5.386

Axi

Trange/K

ARD/%

Bxi

298 to 343

−5.807

298 to 343

−4.331

298 to 343

−1.154

298 to 343

−0.311

298 to 343

−6.060

298 to 343

−3.225

ARD/%

n-butane −1032.04 1.91 n-hexane −2178.94 6.40 n-butane −2093.05 2.94 n-hexane −2910.37 2.74 n-butane −896.08 3.37 n-hexane −2503.33 2.81

Trange/K 298 to 343 298 to 353 298 to 343 298 to 353 298 to 343 298 to 353

Table 11. Henry’s Constants of Isobutane, n-Butane, n-Pentane and n-Hexane in Various Amine Aqueous Solutionsa n-butane

isobutane

/ /MPa

T/K

a

298.29 303.31 313.24 323.20 333.20 343.16 u(/ ) = 87 MPa

1785 1967 2293 2675 3009 3165

298.23 303.30 313.29 323.22 333.24 343.15 u(/ ) = 19 MPa

613 633 680 690 680 616

298.28 303.28 313.26 323.20 333.28 343.18 u(/ ) = 60 MPa

1500 1616 1828 2009 2238 2393

n-pentane

/ /MPa

T/K 298.27 303.27 313.22 323.32 333.25 343.17 u(/ ) = 68 MPa 298.28 303.29 313.19 323.16 333.22 343.16 u(/ ) = 21 MPa 298.26 303.31 313.30 323.27 333.21 343.16 u(/ ) = 62 MPa

T/K

n-hexane

/ /MPa

T/K

/ /MPa

2383 2561 3066 3506 3962 4190

298.17 313.13 333.14 353.18

1802 2584 3102 3175

Aqueous MDEA (w = 0.25) 2374 298.26 2553 303.32 3092 313.20 3729 323.28 4136 333.23 4520 343.17 u(/ ) = 86 MPa Aqueous MDEA (w = 0.50) 758 298.41 845 303.38 894 313.26 945 323.19 950 333.21 946 343.17 u(/ ) = 19 MPa Aqueous DEA (w = 0.35) 1869 298.28 2136 303.29 2657 313.13 3206 323.18 3625 333.22 3854 343.32 u(/ ) = 75 MPa

u(/ ) = 62 MPa 476 546 671 769 740 752

298.21 313.17 333.24 353.17

398 466 499 485

u(/ ) = 8 MPa 2031 2189 2570 3050 3531 3811

298.16 313.16 333.15 353.15

1903 2203 2762 2749

u(/ ) = 85 MPa

Temperature standard uncertainty is u(T) = 0.02 K.

Table 12. Values of Parameters for Calculating Henry’s Law Constants, VLLE (eq 11) solvent

Axi


186.93

−9600

172.33

−8891

334.61

−15 544

558.47

−26 929

77.06

−4189

104.88

−5846


aqueous DEA (w = 0.35)

Bxi

Cxi isobutane −25.84 n-pentane −23.65 isobutane −48.45 n-pentane −81.08 isobutane −9.78 n-pentane −13.63

ARD/%

Trange/K

Axi

0.83

298 to 343

173.68

−9112

0.91

298 to 343

385.40

−18 942

0.91

298 to 343

250.49

−11 893

1.55

298 to 343

214.20

−10 259

0.44

298 to 343

313.40

−15 851

1.26

298 to 343

197.17

−9696

1682

Bxi

Cxi n-butane −23.76 n-hexane −55.17 n-butane −35.79 n-hexane −30.50 n-butane −44.35 n-hexane −27.58

ARD/%

Trange/K

1.22

298 to 343

1.16

298 to 353

1.11

298 to 343

0.30

298 to 353

0.61

298 to 343

2.59

298 to 353



Article

Axi , Bxi

with temperature, except for the systems with aqueous MDEA at w = 0.5, where we can observe a tendency to culminate at a maximum value around 330 K, independently of the hydrocarbon type, as shown in Figure 12. Considering that the Henry’s constant variation is proportional to the pressure and inversely proportional to the solubility (eq 5), the behavior observed for MDEA solution (w = 0.5) can be explained by the fact that the hydrocarbon solubility is much higher compared to the other solutions, in which case its increase with temperature overrides that of pressure. This suggests that the Henry’s constant maximum will be more pronounced for higher amine concentrations. The temperature dependence of the Henry’s constant is represented by the following correlation: ln(/i/MPa) = A iH + BiH /T + CiH ln(T )

AHi ,

BHi ,

T /K

AHi , BHi , CHi

fi / MRD Nc P RD T x y w u uc φi γi Superscripts A related to aqueous phase L related to the hydrocarbon rich liquid phase sat at saturation V related to hydrocarbon rich vapor phase W related to pure water phase Subscripts a amine c critical properties cal calculated exp experimental hc hydrocarbon species solv solvent w water

(11)

CHi

Coefficients for each system were obtained by minimizing the sum of the square difference between the calculated (eq 11) and estimated (eq 5) Henry’s constant. They are summarized in Table 12 which includes the average relative deviations. The ARD does not exceed 3 % for all the systems: less than 0.91 % for isobutane, 1.22 % for n-butane, 1.55 % for n-pentane, and 2.59 % for n-hexane, which is satisfactory in the temperature range of this study. The fitted parameters in this paper (from eq 9 to eq 11) were obtained using the regression tool available within Simulis thermodynamics environment; a software provided by ProSim company20 for phase equilibria and thermo-physical properties calculations. Simulis Thermodynamics was also used to estimate the fugacity coefficients in eq 5.

■

CONCLUSION LPGs are commonly used for heating and as a transportation fuel. Scrubbing LPG with amines, to reduce acid gas concentration, may cause some hydrocarbon losses in the amine aqueous solutions. These losses need to be quantified. In this paper, new solubility data of the usual components of LPGs (n-butane, isobutane, n-pentane, and n-hexane) in aqueous alkanolamine solutions (DEA and MDEA) were provided. The static analytic method used to achieve our experiments has allowed an accurate determination of the solubilities with a combined uncertainty not higher than 4 % for all systems. It was found that the solubility increases with the amine concentration as well as with the temperature, whereas it decreases with the hydrocarbon’s carbon atom number. Moreover, practical correlations for the threephase pressure, the solubility data, and the Henry’s constants have been generated from the experimental data. The relative deviations obtained with those correlations are satisfactory.

■

parameters of each hydrocarbon i in the solubility correlation parameters of each hydrocarbon i in the Henry’s constant correlation fugacity of species (MPa) Henry’s law constant (MPa) maximum relative deviation carbon atom number of the hydrocarbon species pressure (MPa) relative deviation temperature (K) liquid mole fraction vapor mole fraction mass fraction standard uncertainty standard combined uncertainty fugacity coefficient of species i activity coefficient of species i

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REFERENCES

(1) Sada, E.; Kumazawa, H.; Butt, M. A. Solubilities of gases in aqueous solutions of amine. J. Chem. Eng. Data 1977, 22 (2), 272−273. (2) Fleming, K. B.; Spears, M. L.; Bullin, J. A. Design Alternatives for Sweetening LPG’s and Liquid Hydrocarbons with Amine; Bryan Research and Engineering, Inc.: Bryan, TX, 1988. (3) Mokraoui, S.; Coquelet, C.; Valtz, A.; Richon, D. Mutual solubility of hydrocarbons and amines, GPA Research Report 195 (RR-195). Gas Processors Association: Tulsa, OK, 2008. (4) Critchfield, J.; Holub, P.; Ng, H. J.; Mather, A. E.; Bacon, T. Solubility of hydrocarbons in aqueous solutions of gas treating amines. Proceedings of the Laurence Reid Gas Conditioning Conference; University of Oklahoma: Norman, OK, 2001; pp 199−227. (5) Jou, F. Y.; Carroll, J. J.; Mather, A. E.; Otto, F. D. Phase equilibria in the system n-butane−water−methyldiethanolamine. Fluid Phase Equilib. 1996, 116, 407−413. (6) Darryl M.; Kevin F.. Hydrocarbon Solubility in Glycol and Amine Solutions; GPA Research Report (RR-206); Trimeric Corporation: Buda, TX, 2012. (7) Jou, F.-Y.; Ng, H.-J.; Mather, A. E. Solubility of propane in aqueous alkanolamine solutions. Fluid Phase Equilib. 2002, 194, 825−830. (8) Engineering Data Book, 13th ed.; Gas Processors Suppliers Association: Tulsa, OK, 2012. (9) Laugier, S.; Richon, D. New apparatus to perform fast determinations of mixture vapor−liquid equilibria up to 10 MPa and 423 K. Rev. Sci. Instrum. 1986, 57, 469−472. (10) Chapoy, A.; Mokraoui, S.; Valtz, A.; Richon, D.; Mohammadi, A. H.; Tohidi, B. Solubility measurement and modeling for the system propane−water from 277.62 to 368.16 K. Fluid Phase Equilib. 2004, 226, 213−220. (11) Mokraoui, S.; Coquelet, C.; Valtz, A.; Hegel, P. E.; Richon, D. New solubility data of hydrocarbons in water and modelling concerning

AUTHOR INFORMATION

Corresponding Author

*Tel.: +966-1-4676832. Fax: +966-1-4697122. E-mail: [email protected]. Funding

Authors gratefully acknowledge the Gas Processors Association (GPA, USA) for the financial support of this research. Notes

The authors declare no competing financial interest.

■

NOMENCLATURE ARD average relative deviation APi , BPi parameters of each hydrocarbon i in the three-phase pressure correlation 1683



Article

vapor−liquid−liquid binary systems. Ind. Eng. Chem. Res. 2007, 46, 9257−9262. (12) Mokraoui, S.; Hadj-Kali, M. K.; Valtz, A.; Richon, D. New vapor− liquid−liquid equilibrium data for ethane and propane in alkanolamine aqueous solutions. J. Chem. Eng. Data 2013, 58, 2100−2109. (13) Sandler, S. Chemical, Biochemical, and Engineering Thermodynamics, 4th ed.; Wiley: New York, 2006. (14) Carroll, J. J.; Jou, F.-Y.; Mather, A. E.; Otto, F. D. Phase equilibria in the system water−methyldiethanolamine−propane. AIChE J. 1992, 38, 511−520. (15) Peng, D. Y.; Robinson, D. B. A new two-constant equation of state. Ind. Eng. Chem.: Fundam. 1976, 15, 59. (16) Reid, R. C.; Prausnitz, J. M.; Poling, B. E. The Properties of Gases and Liquids, 4th ed.; McGraw-Hill Book Company: New York, USA. 1987. (17) Daubert, T. E.; Danner, R. P.; Sibel, H. M.; Stebbins, C. C. Physical and Thermodynamic Properties of Pure Chemicals; Data Compilation; Taylor & Francis: Washington, D.C, 1997. (18) Carroll, J. J.; Maddocks, J.; Mather, A. E. The solubility of hydrocarbons in amine solutions. Laurence Reid Gas Conditioning Conference; University of Oklahoma: Norman, OK, 1998. (19) Tsonopoulos, C. Thermodynamic analysis of the mutual solubilities of normal alkanes and water. Fluid Phase Equilib. 1999, 156, 21−33. (20) ProSim software website, http://www.prosim.net (accessed Feb. 03, 2014).

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New VaporâLiquidâLiquid Equilibrium Solubility Data for iso-Butane, n

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