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J. Chem. Eng. Data 2005, 50, 817-821

817

Solubility of Carbon Dioxide in Aqueous HCl and NaHCO3 Solutions from 278 to 298 K Chi Shing Wong,*,† Pavel Ya. Tishchenko,‡ and Wm. Keith Johnson† Institute of Ocean Sciences, P.O. Box 6000, 9860 West Saanich Road, Sidney, British Columbia V8L 4B2, Canada, and V.Il’ichev Pacific Oceanological Institute, 43 Baltiyskaya Street, Vladivostok 690041, Russia

The solubility of carbon dioxide in aqueous HCl and NaHCO3 solutions has been determined by the coulometric technique at ambient pressure and at 278, 283, 287, and 298 K. The maximum molality of HCl and NaHCO3 solutions was 1 and 0.5 mol‚kg-1, respectively. The error in CO2 solubility in HCl solutions estimated from parallel experiments was about 0.3%. For bicarbonate solutions, the overall error was larger because the solubility is the difference between the determined dissolved inorganic carbon and sodium bicarbonate molality. Under investigated conditions, errors in solubility data varied from (0.3% for dilute solutions to (1.3% for 0.5 m sodium bicarbonate. Obtained data were treated by the Pitzer method. Pitzer parameters λCO2,NaHCO3 and λCO2,HCl were derived as a function of temperature.

Introduction The salting-out effect of electrolytes on the solubility of CO2 and the reverse effect of dissolved CO2 on the behavior of electrolyte solutions have been studied as a valuable step in developing reliable chemical models for technology and for many geological fluids, including seawater, hydrothermal water, and oil-field brines.1-5 In these studies, the Pitzer method6,7 was successfully used for the treatment of experimental data as well as for developing predictive models. It is obvious that dissolved CO2 gas may have an effect on the activity coefficients of carbon dioxide system species (CO2(aq), H+, HCO3-, and CO32-) and as a result produce a shift in stoichiometric acid-base equilibrium. This possible effect was eliminated in classical experiments to determine the first ionization constants of carbonic acid by empirical extrapolation of CO2 to zero content.8 For investigations of the carbon dioxide system in seawater, this effect is negligible because of the very low molalities of dissolved CO2, H+, HCO3-, and CO32- ions in seawater and is not considered. However, for experiments involving the ocean disposal of fossil fuel CO29 or for the study of carbon dioxide-rich hydrothermal fluids,10 the effect of high CO2 molality on activity coefficients of components of the carbon dioxide system may be appreciable and should be taken into account. The influence of dissolved CO2 on carbon dioxide equilibrium in relation to liquid CO2 disposal in the ocean can be taken into account by the Pitzer method if the needed Pitzer parameters are available at low temperatures. The most important parameters are λCO2,CO2 and λCO2,H + λCO2,HCO3. The second virial coefficient λCO2,CO2 can be derived from high-pressure CO2 solubility data in pure water.4,11,12 The electroneutral combination of virial coefficients of interest (λCO2,H + λCO2,HCO3) may be obtained from CO2 solubility data in NaCl, HCl, and NaHCO3 solutions at ambient pressure. There are accurate CO2 solubility data in aqueous NaCl solutions at low * To whom correspondence should be addressed. E-mail: [email protected] pac.dfo-mpo.gc.ca. Phone: +1-250-363-6407. Fax: +1-250-363-6476. † Institute of Ocean Sciences. ‡ Pacific Oceanological Institute.

temperatures8 that were used to determine λCO2,NaCl12 as a function of temperature, but data are scarce on the solubility of CO2 in HCl5,13,14 and NaHCO315 aqueous solutions at low temperatures. Our aim in this work was to determine CO2 solubility in aqueous HCl and NaHCO3 solutions experimentally for molality ranges of 0.2 to 1 and 0.1 to 0.5 mol‚kg-1, respectively, and at temperatures of 278, 283, 287, and 298 K. The Pitzer method was used to treat data for the evaluation of λCO2,HCl and λCO2,NaHCO3 parameters as a function of temperature. Experimental Procedures Measurements of CO2 solubility were carried out by the coulometric technique, which is commonly used to measure dissolved inorganic carbon.16,17 A stock solution of HCl of approximately 1 mol‚kg-1 was prepared from ACS-grade concentrated acid. The exact concentration was determined by the coulometric titration method. HCl solutions in other concentrations were prepared by diluting the stock solution. All experiments were carried out using a Mettler AT400 analytical balance, taking into account buoyancy corrections. Sodium bicarbonate (ACS grade) was used without any additional purification and was assayed by the Fisher Scientific Co. to be 100.1% NaHCO3. Results of our titration were 100.12%, which agreed well with the value given by Fisher Scientific. Solutions of sodium bicarbonate were prepared by weight. Barometric pressures were measured by means of a digital pressure gage (Setra model 320 with a resolution of 0.01 mbar). High-purity CO2 gas (99.99%) was saturated with water vapor by bubbling first through a chamber containing a test solution before being equilibrated with the same solution in a second chamber. Both chambers were immersed in a water bath at a temperature that was constant to within (0.01 K. After bubbling CO2 gas through the equilibration chamber for over 2 h, we drew about 2 g of the sample from the second chamber using a glass syringe. The sample was injected into a scrubber via a rubber septum and acidified (NaHCO3 solutions) with CO2-free orthophosphoric acid. The CO2 in the solution was then purged, using ultra-high-purity

10.1021/je049716q CCC: $30.25 © 2005 American Chemical Society Published on Web 03/10/2005

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Journal of Chemical and Engineering Data, Vol. 50, No. 3, 2005

Table 1. Measured Solubility Coefficients of CO2 (r/mol‚kg-1‚bar-1) in NaHCO3 and HCl Solutions 278.15 K

283.15 K R

m

m

288.15 K R

293.15 K R

m

m

298.15 K R

R

m

mol‚kg-1 mol‚kg-1‚bar-1 mol‚kg-1 mol‚kg-1‚bar-1 mol‚kg-1 mol‚kg-1‚bar-1 mol‚kg-1 mol‚kg-1‚bar-1 mol‚kg-1 mol‚kg-1‚bar-1 0.09973 0.20017 0.29874 0.40523 0.51435

0.0606 0.0573 0.0539 0.0509 0.0462

0.10333 0.20179 0.30049 0.39971 0.50043

0.0497 0.0474 0.0436 0.0414 0.0381

NaHCO3 Solutions 0.09957 0.0430 0.20046 0.0410 0.30001 0.0379 0.39899 0.0352 0.50885 0.0321

0.10314 0.20110 0.30064 0.39754 0.49825

0.0366 0.0341 0.0312 0.0288 0.0271

0.10843 0.19909 0.29908 0.40206 0.49876

0.0324 0.0305 0.0281 0.0265 0.0238

0.19031 0.40314 0.61220 0.81950 1.09303

0.06231 0.06171 0.06092 0.06069 0.06032

0.19976 0.40209 0.61174 0.84022 1.09303

0.05220 0.05181 0.05153 0.05139 0.05104

HCl Solutions 0.20021 0.04439 0.40368 0.04411 0.61040 0.04392 0.81047 0.04384 1.09303 0.04361

0.20487 0.40240 0.61184 0.80407 1.09303

0.03832 0.03816 0.03796 0.03791 0.03756

0.19988 0.39690 0.59983 0.81097 1.09303

0.03339 0.03329 0.03314 0.03310 0.03301

nitrogen, and transferred into a coulometric cell where it was titrated. The exact amount of sample was determined by weighing the syringe before and after injection. Each of the samples was run a minimum of four times. The results of the first run were not used because they gave systematically lower values, as a rule, because of trapped air in the needle of the syringe, which was flushed out after the first injection. Analytical runs were continued until three runs agreed to within 0.15% of the total dissolved inorganic carbon (DIC). In the case of sodium bicarbonate solutions, molalities of dissolved CO2 gas were calculated as the difference between total dissolved inorganic carbon (measured) and the prepared (calculated) molality of NaHCO3. For all investigated solutions, determined molalities of CO2 were normalized to standard CO2 pressure, P° ) 0.1 MPa. During our experiments, atmospheric pressure (Ptot) varied; therefore, solubility coefficients were calculated by means of the following equation

R)

mCO2

(1)

Ptot - Psw

where mCO2 is the measured molality of dissolved CO2, Ptot is atmospheric pressure, and Psw is the vapor pressure of water. The solubility coefficients are expressed in mol‚ kg-1‚bar-1 and are related to a standard pressure of 0.1 MPa. Each value is an average of three runs, and the results are listed in Table 1. DIC was determined to have a standard deviation of about (0.07%. Therefore, the error in CO2 solubility in HCl solutions is about 0.3% at the 0.99 confidence level. For bicarbonate solutions, the overall error is larger because the solubility is the difference between DIC and the sodium bicarbonate molality. This is the most important distinction of this CO2 solubility study from other solution solubility studies. Under our conditions, errors in the solubility data varied from (0.3% for dilute solutions to (1.3% for 0.5 m sodium bicarbonate.

The experimental conditions for determining CO2 solubility in HCl and NaHCO3 aqueous solutions imply that there are two coexisting phases in equilibrium for the system under consideration. At vapor-liquid equilibrium, the chemical potential of CO2 is the same in both phases, allowing us to write the following equation

(

ln

kH,mm

o

)

+

∫

Psw

Po

V∞CO2(T, P) RT

dP ) ln

(

)

mCO2γCO2 mo

Psw yCO2 ) 1 P

(3)

Psw ) Pow exp(-MAφm

∑m) i

(4)

i

Here, Pow is the vapor pressure of pure water,19 φm is the osmotic coefficient, and MA is the molar mass of water. The osmotic coefficient has been calculated using the Pitzer method, where the Pitzer parameters for NaHCO3 and HCl solutions have been taken from Millero20 and Peiper.21 The Pitzer method7 for neutral species gives the equation for the activity coefficient of dissolved CO2 as follows:

ln(γCO2) ) 2mCO2λCO2,CO2 + 2

∑m λ c

c CO2,c

6

+2

∑m λ

a CO2,a

a

∑∑m m ζ c

c

a

a CO2,c,a

+

(5)

For single 1-1 electrolyte (our case) solutions, eq 5 yields

ln(γCO2) ) 2mCO2λCO2,CO2 + 2m(λCO2,c + λCO2,a) + 6m2ζCO2,c,a (6)

Thermodynamic Treatment and Discussion

yCO2PtotφCO2

where R is the gas constant, T is the temperature in Kelvin, mo is the standard molality (1 mol‚kg-1), φCO2 is the fugacity coefficient of CO2 for which values were calculated using Weiss’s relationship,18 kH,m is the Henry’s law constant (molality scale), γCO2 is the activity coefficient of CO2 referenced to Henry’s law (molality scale), V∞CO2 is the partial molar volume of CO2 at infinite dilution, and yCO2 is the molar fraction of CO2 in the gas phase and is calculated from the following relationship:

(2)

where m is the electrolyte molality (NaHCO3 or HCl). Using Weiss’s relationship for the solubility coefficient in seawater18 and solubility data of liquid CO2,22 we12 obtained empirical relations for λCO2,CO2 and for the solubility coefficient of CO2 in seawater, K/o,

λCO2,CO2/kg‚mol-1 ) -0.2546 + 75.571/T - 5.428 × 10-10P/Pa (7)

Journal of Chemical and Engineering Data, Vol. 50, No. 3, 2005 819

ln(K/o/mol‚kg-1‚bar-1) ) -5.97508 × 101 + 9.34517 × 103 K/T + 2.33585 × 101 ln(T/100 K) - 3.2874 × 10-3T/K + 5.380 × 10-6T2/K2 + S[2.2843 × 10-2 - 2.2534 × 10-4T/K + 4.518 × 10-7T2/K2] (8) where S is the salinity of seawater, expressed as the mass fraction per mil. The solubility coefficient K/o corresponds to trace molalities of CO2 in solution and relates to the Henry’s law constant as follows:

ln{K/okH,m}

)

∫

Psw

V∞CO2(T, P)

Po

RT

dP - ksetS

(9)

Here, kset is the Setschenow coefficient for seawater. From eqs 1, 2, 6, and 9, we have

( }

ln

K/oφCO2 R

) 2mCO2,CO2λCO2,CO2 + 2(λCO2,c + λCO2,a)m + 6ζCO2,c,am2 (10)

When the electrolyte molality is less than 1, the contribution of the third term of eq 10 is not important, as a rule, and for ambient pressure conditions, the slope of the dependence ln(RPo/mo) versus m is roughly 2(λCO2,Na + λCO2,HCO3). This dependence, including all experimental points, is demonstrated in Figure 1. There is a curvature in the dependencies shown in Figure 1a despite there being comparatively low concentrations of sodium bicarbonate in the studied solutions. This gives evidence that the third term of eq 10 is important for sodium bicarbonate solutions. According to eq 10, we defined function f as

[ ] φCO2mo

f ≡ ln

RPo

- 2mCO2λCO2,CO2

Figure 1. Logarithm of the solubility coefficient of carbon dioxide in (a) NaHCO3 and (b) HCl solutions at the following temperatures: b, 278.15 K; 9, 283.15 K; 2, 288.15 K; [, 293.15 K; ×, 298.15 K. The solubility coefficients for pure water (O) were calculated using eq 8.

(11)

Using this equation, we were able to determine extrapolated values of the solubility coefficients (Ro) at zero molality using the least-squares method and two- and three-term statistical hypotheses. Relative deviations between values of Ro obtained from our experimental data for NaHCO3 and K/o, calculated by eq 8, demonstrate better agreement when the three-term hypothesis was used (Figure 2). However, this improvement is not significant (