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J. Phys. Chem. 1989, 93, 7237-1248
Solubility of Ammonia in Pure Aqueous and Multicomponent Solutions S. L. Clegg* and P. Brimblecombet Plymouth Marine Laboratory, Citadel Hill, Plymouth PLl 2PB, UK (Received: January 23, 1989)
The thermodynamic Henry's law constant (KH/mol kg-' atm-I) of the weak electrolyte NH, is described by the equation In (KH)= -8.09694 + 3917.5071T- O.O0314T, from 273 to 313 K. Measured NH3 solubilities in both pure aqueous and multicomponent solutions agree well with calculations using the Pitzer thermodynamic model to predict osmotic and activity coefficients. Partial pressure and heat of dilution data were used to determine the activity coefficient of NH, (YNH,). In pure aqueous solutions where dissociation is not significant this is given by yNH, = eXP(2"H,XNH,,NH3), where XNH,,NH~ 0.033161 - 21.12816/T+ 4665.14611p from 273 to 313 K. Ion-NH, interaction parameters were obtained by using partial pressure, salt solubility, and partitioning data for the following ions at 298.15 K: Li', Na+, K+,NH4+,Mg2+,Ca2+,Sr2+, Ba2+,F,Cl-, Br-, I-, OH-, CNS-, NO3-, NO2-, C103-, C104-, S2-, S032-,CH3COO-, HCOO- and (COO)22-. Parameter values are found to be simply related to ion charge and size.
I. Introduction In order to calculate gas solubility in typical industrial and natural geochemical systems, a thermodynamic model of solution behavior must be used in conjunction with basic thermodynamic properties such the Henry's law constant and the dissociation constant. An analogous approach has been successfully used to predict mineral solubilities in concentrated brines,'-3 and similar systems are also important in the atmosphere where very concentrated solutions occur as micrometer-sized aerosol particles and droplets. In this work thermodynamic Henry's law constants for N H 3 are presented, and we describe a treatment of N H 3 solubility in pure aqueous and multicomponent solutions in terms of the Pitzer thermodynamic model, utilizing a wide range of partial pressure, salt solubility, and NH3 partitioning data. 11. Theoretical Approach
The solubility of a weak electrolyte such as N H 3 is described in two stages, as dissolution followed by base dissociation N H 3 ( d = NH3(aq) (1) NH,(aq)
are a ~ a i l a b l e . ~ ~ ~The * ' ~ equations .'~ for the osmotic coefficient (+), and activity coefficients (y) of cation (M), anion (X), and neutral species (N) are given below:
(4- 1) = (2/Cmi)[-MP/*/(1 I
+ 1 . 2 W ) + CCmcma(lPca + ZCca) + c a
C Cmcmcl(@+cc'+ Xma+ccTa) + C Cmamai(@+aai+ a < a'
a
c 8 mol dm-3 N H 3 are included, both with and without the triplet parameter pN,N,K in eq 17. It is clear that at such high concentrations of N H 3 this term is significant. Recommendations for ion-neutral parameter values for general use are based principally upon salt solubility data, since they are to very high aqueous phase concentrations and thus the effect of experimental error on derived parameter values is minimized. Values for Na+, K+,NH4+, S042-,and C0:- have been adopted and are listed in Table VI. Of the parameters obtained only from partial pressure and partitioning data (Tables 111-V), those for Li+, I-, and to a lesser extent Br- and OH- agree well. For the first three of these ions, values listed in Table VI are simple averages. Independent fitting of the partial pressure data suggested XN,OH equal to 0.103-higher than that given in Table 111. Parameter values for F,Clod-, CH,COO-, HCOO-, and (C00)22are adopted unchanged. The use of parameter values based on salt solubility data causes some difficulties for series of anions such as S042-, S032-,S2-. The partial pressure data (Table 111) show that the parameter values increase monotonically, similarly to those of the alkali metals and halides. It was therefore assumed that the best fit slopes correctly represent the relative magnitudes of the parameters, and estimates given in Table VI are based on equal to 0.1 4, and cation-ammonia interaction parameters taken from the same data set. Estimates of XN,NO, obtained from both partial pressure and NH3 partitioning data are significantly greater than that in Table 11, so the mean of -0.008 has been adopted for this parameter. To remain consistent with this, a slightly reduced value of XN,,-1oj is listed in Table VI.
The chief remaining difficulty concerns XN,Na, as the estimate using salt solubility data (Table 11) is significantly lower than those obtained from other data. The original value of XN,Na (0.0175) has been retained in Table VI as it was found to be robust. However, it is recognized that for some purposes parameter values obtained from partial pressure or partitioning data may be preferred. In this case XN,Na equal to 0.031 should be adopted. Note that this may imply a somewhat lower since this parameter was obtained by using salt solubility data with XN,Na equal to 0.017 51. What effect does the higher value of XN,Na have upon estimates of Na+ salt solubility in aqueous NH3 media? In test calculations it was found that Na2S04solubilities at 25 OC ranged from 10% lower than measured values in 4 mol kg-I NH,, to 35% lower in 20 mol kg-' NH,. Differences such as these are substantial, and the value of XN,Na therefore needs to be known more precisely, particularly as it is important in the calculation of the dissociation constant of N H 3 in brines and natural waters. This also applies to the alkaline earth cations, since the interaction parameters listed in Table VI for these ions are based upon single determinations of N H 3 partial pressure.
VI. Calculation of the Dissociation Constants of Ammonia In many systems the dissolved concentration of N H 3 is such that dissociation according to eq 2 must be taken into account when the solubility is calculated. In such systems, estimation of the degree of dissociation, rather than the value of YNH,, may be the principal source of error. As a test of this application of the Pitzer equations, emf measurements for Na+/Cl-/NH3 and K+/Cl-/NH3 aqueous solutions71containing 0.503 mol kg-I total N H 3 were used to derive Kb values, which were compared with calculation. The data of Harned and Robinson,71 obtained by using a cell incorporating a liquid junction, were related to the free concentration of OH- by an expression (eq 19),71which includes the activity coefficient ratios of C1-, OH-, and H 2 0 in the two half-cells. Harned and Robinson assumed that these activity ratios were unity, since the half-cell compositions were matched as closely as possible. In the present work the activity coefficients were calculated by using the Pitzer model, since small differences do arise, and an iterative solution of eq 1g7' was obtained. The results are compared (as pKb of NH,) in Figure 18 with stoichiometric dissociation constants calculated by using the Pitzer model with available ion-ion interaction parameters together with ion-neutral parameters determined in this study. Agreement between measurement and calculation is to better than 0.05 in pKb. The systematic deviations that occur may be due to liquid junction potentials at low salt concentrations, or the fact that these systems are not fully parametrized for ion-ion interactions involving NH4+,Na+, and CI-. Based on previous measurements it was assumed that, in the NaCl system, the &a,NH4 and +Na,NH4,CI parameters take the same values as corresponding (71) Harned, H. S.; Robinson, R. A. J . Am. Chem. SOC.1928,50, 3 1 5 7 .
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J. Phys. Chem. 1989, 93. 7248-7252
coefficients involving K', Na+, and C1-. Maeda et al.62969have determined potentiometrically the equilibrium constant (K,) for the dissociation of NH4+ in single salt solutions containing the ions Li+, Na+, K+, C1-, NO), and C104-. These authors compared their results with calculation, using yNHo and Y~ estimated by the Bronsted, Scatchard, and Guggenheim specific interaction theory,72 and experimentally determined yNH,. Measured and calculated pK, values agreed to within 0.05 for Na+ and K+ salt solutions. Estimates of pK, made in this work using the Pitzer model were found to have a similar degree of accuracy. Differences between measured and calculated values were attributed to the fact that not all of the higher ion-ion interaction parameters (0, and qijk)are available. More recent work of Maeda et aL6*on the dissociation constant of NH4+ in LiCl solutions has emphasized the importance of these terms and shown that predictions using the Pitzer model fully parametrized for the system yield values of pK, in excellent agreement with experiment.
VII. Conclusions This work has shown that the activity of N H 3 in pure aqueous and multicomponent solutions, and the effect of N H 3 on ionic solutes, are satisfactorily described by the complete Pitzer thermodynamic model for a wide range of solution composition. As has been noted, in its simplest form using only doublet interaction parameters the model expression for the activity coefficient of a neutral species is analogous to the Setschenow equation. However, the advantage of using the Pitzer model to describe the behavior of neutral species is ready integration with equations describing the activities of both water and ionic solutes. The results presented here, in addition to being of practical use in calculating the properties of NH,, suggest that the Pitzer model may usefully be applied to the wide range of partially dissociating and nondissociating solutes for which solubility and other thermodynamic data exist. Registry No. NH,, 7664-41-7; Li', 17341-24-1; Na*, 17341-25-2; K+, 24203-36-9; NH4+, 14798-03-9; Mg2+,22537-22-0; Ca2+,14127-61-8; Sr2+,22537-39-9; Ba2+, 22541-12-4; F, 16984-48-8; C1-, 16887-00-6;
(72) Guggenheim, E. A. Application of Statistical Mechanics; Clarendon Press: Oxford, 1966. (73) Wagman, D. D.; Evans, W. H.; Parker, V. B.; Schumm, I. H.; Bailey, S. M.; Churney, K. L.; Nuttall, R. L. J . Phys. Chem. ReJ Data 1982, 1 1 , 1.
Br-, 24959-67-9; I-, 2046 1-54-5; OH-, 14280-30-9; CNS-, 302-04-5; NO
