Equilibrium partial pressures of strong acids over concentrated saline solutions—I. HNO3

Afmospheric Printed

Environmenr

Vol.

22, No.

I, pp.

91-100.

00Md981/88

1988.

S3.00+0.00

Pergamon Journals Ltd.

in Great Britain

EQUILIBRIUM PARTIAL OVER CONCENTRATED

PRESSURES OF STRONG ACIDS SALINE SOLUTIONS-I. HN03

S. L. CLEGG*~ and P. BRIMBLECOMBE School of Environmental Sciences, University of East Anglia, Norwich NR4 7TJ, U.K. (First received 19 February 1987 and received for publication 8 July 1987)

Abstract-Equilibrium partial pressures of HN03 have been measured at 25°C over concentrated aqueous solutions containing the ions H+, NO;, SO:-, K+, NH: and Mg*+. Measurements agree closely with predictions made using a Henry’s Law constant of 2.45 x lo6 mol’ kg-’ atm-’ for the reaction HNO!@, = H+ + NOr,,,, and the Pitzer activity coefficient model. The model is recommended for the calculation of so ute and so vent activities in aqueous aerosols at high ionic strength, corresponding to equilibrium relative humidities greater than about 70 %. The following new parameters have been estimated for the mode1 using the partial pressure measurements and other thermodynamic data: $H,+N~~, -0.003; ~NO,,SO,, -0.00% *No,,So,,Na, O.W& JIH,NH~,NO,, -0.01; *H,K,NO,, -0.01. o.0926; tiNO,,SO.,,K. Key word index: Aerosols, nitric acid, _ partial _ pressure, Henry’s Law, activity coefficient, Pitzer model, saline sol;tions, solubility.

1982). It has been shown that the dissolution of HN03 into the seasalt aerosol can completely displace chloride as HCl (Martens et al., 1973). Recent measurements of HNO, in the remote free troposphere imply that anthropogenic HN03 may be transported to remote parts of the atmosphere with relatively high efficiency (Galasyn et al., 1988) and constitutes a significant source of strong acidity. Reactions such as that above have been incorporated into aerosol models (Bassett and Seinfeld, 1983; Clegg and Brimblecombe, 1985a), usually as an equilibrium process since ambient data and extensive thermodynamic predictions suggest that equilibrium generally exists between aerosol and gas phases (Stelson et al., 1979; Stelson and Seinfeld, 1982; Tanner, 1982). The chemistry of the aerosol is complicated by the non-ideality inherent in very concentrated multicomponent solutions, whose chemical properties are poorly understood. Modelling the behaviour of soluble acids is often complex and the approach taken varies from author to author (e.g. Clarke and Williams, 1983; Saxena and Peterson, 198 1; Stelson et al., 1984). In this work and a companion paper (Clegg and Brimblecombe, 1988) we present new determinations of equilibrium HN03 and HCl partial pressures above concentrated acidified salt solutions containing the major ions presetit in tropospheric and marine aerosols. Test solutions were chosen to produce highly non-ideal solution behaviour, and yet were simple enough to show the effect of individual ions on acid gas solubility. The partial pressure data obtained from the experiments were used to validate a theoretical treatment of acid gas solubility. This should enable both equilibrium acid gas partial pressures and water activities (equivalent to equilibrium relative humidity) to

1. INTRODUCTION

Very concentrated solutions occur in the atmosphere as pm-sized aerosol particles and droplets. Since the aerosol contains very small amounts of water relative to the surrounding gas phase the ambient relative humidity controls solute concentrations in the droplets. These range from very dilute near 100 “/, relative humidity to supersaturated with respect to soluble constituents (such as NaCl) below about 75 “/, relative humidity. Under such conditions ionic strengths greater than 10 mol kg-’ may be reached in the aerosol. Interactions between aqueous aerosols and other atmospheric constituents are important for the behaviour of anthropogenic smog and aerosols, and on a large scale in the cycling of elements such as S, N and the halogens. For example the accumulation of excess sulphate and depletion of chloride in the marine aerosol is well known (Buat-Menard et al., 1974; Meinert and Winchester, 1977; Hitchcock et al., 1980). It is thought to be the result of the dissolution of H$04 and its precursor SO2 into the aerosol, and the expulsion of volatile HCl to the gas phase (Cauer, 1949). Eriksson (1960) used a simple mass balance equation to account for this: HzS04 + 2NaCl = NazSOa + 2HCl,,,.

(1)

Both laboratory and field data suggest that similar reactions involving HN03 occur, particularly in polluted environments (Robbins et al., 1959; Tanner,

* Present address: The Marine Biological Association of the United Kingdom, The Laboratory, Citadel Hill, Plymouth PLl 2PB, U.K. t To whom correspondence should be addressed. 91

S. L. CLEGGand P.

92

be predicted for multicomponent solutions over a wide range of composition and concentration.

2. THEORY The equilibrium of a strong acid HX between aqueous and gas phases is represented by: HX(,, = Hi,., + X,&J. The thermodynamic Henry’s Law constant (mo12 kg-* atm-‘) is given by: K, = mH+.mX-.yZ,,/pHX

(2) K,

BRIMBLECOMBE

are similar in form to series used to describe the properties of non-ideal gases. The theory is developed in detail by Pitzer and co-workers in 12 papers of which the following are most relevant to this work: Pitzer, (1973, 1975); Pitzer and Mayorga, (1973, 1974); Pitzer and Kim, (1974); Pitzer and Silvester, (1976); Pitzer et al., (1977). An equation is postulated for the excess Gibbs energy from which other functions can be obtained: Ge"/n,RT= F(I)+ E[iij(l)]mimj + CXZpijkmimjmk.

(3)

where prefix ‘m'represents molality, yux is the mean activity coefficient of H’ and X- ions in solution and pHX the equilibrium partial pressure of HX. The theoretical justification for this approach is discussed by Clegg and Brimblecombe (1986), and also by Denbigh (1971). A Henry’s Law constant of 2.45 x lo6 mol* kg-* atm-’ (standard deviation 0.18 x 106) was determined for HNOs at 25°C from available activity coefficient and partial pressure data (Brimblecombe and Clegg, 1988). Measurements by the present authors agree with this value. It should be noted that Schwartz and White (1981) derive an equivalent value of 3.26 x lo6 mol’ dm - 6 atm - 1 from essentially the same data, though by a less direct method. The reason for the discrepancy is not clear. However, the measurements presented below suggest that the former value is more nearly correct. In order to calculate partial pressures of HN03 above solutions of known composition, values of the mean activity coefficient of H+ and NO; ions in the solution are required [Equation (3)]. In model calculations for realistic aerosol systems the water activity of the droplets is also needed, as this is equivalent to the ambient relative humidity for droplets at equilibrium with their environment. Neither solvent nor solute activities can be calculated exactly, but there are a variety of models available to estimate them. In solution mixtures of strong electrolytes where association of ions is assumed not to occur, an ioninteraction rather than an ion-pairing model is the appropriate choice. The most powerful of these is that of Pitzer and co-workers (summarized in Pitzer, 1979; Harvie and Weare, 1980). It has been extensively used by geochemists to model activity coefficients in seawater (Whitfield, 1975; Millero, 1982), phase equilibria in brines (Harvie and Weare, 1980; Harvie et al., 1984) and vapour-liquid equilibria in electrolyte systems (Chen et al., 1979). The Pitzer model can be used to calculate both water and solute activities, thus meeting two of the most important requirements of atmospheric modellers. It has therefore been adopted in the present work and will now be briefly described. In Pitzer’s model the Gibbs excess free energy of a mixed electrolyte solution and the derived properties osmotic and mean activity coefficients are represented by a virial expansion of terms in concentration. These

(4)

Gex/n,,, is the excess Gibbs energy kg-’ of solvent and m,,mj,mk the molalities of the various ions or neutral solutes present. The Debye-Huckel term F(I) represents long-range electrostatic forces which dominate in very dilute solutions. Short-range interactions are accounted for by virial coefficients iij for binary and p,jk for ternary interactions. The virial coefficients ibij are dependent upon ionic strength (I), something not recognized in earlier treatments (e.g. Guggenheim, 1935; Scatchard, 1968). They are also used for interactions between ions of like sign. Ternary interactions pijk are used but are not considered to depend on ionic strength. They are omitted for all triplets ofions of like they do appear for most applications (Harvie and Weare, 1980). In expressions for and mean activity coefficients from Equation (4) the virial coefficients are into observable combinations so that interaction parameters can from available This representation only a terms has for modelling For example, most of required virial coefficients for a complex mixture such as seawater are from measurements pure electrolyte solutions. Others are derived from data solution mixtures containing three or four can be modelled data obtained from relatively few very simple mixtures. The following parameters are the model to calculate solute activity and osmotic coefficient, the water activity the solution. For ions of opposite sign four parameters PC*) C4 used. are obtained the equations For solutes except 2:2 metal sulphates the /I’*) parameter is to zero. For ions of like sign and ions (two like sign and one of opposite two parameters are used, These are from activity Bij and made on three ion and Kim, 1974). It should out that the six parameters are not merely ‘empirical’ but physical significance, and the same parameter set used in calculation both osmotic

Equilibrium

partial

pressures

of strong

Table 1. Values of the pure electrolyte

acids over concentrated parameters

saline solutions

/?‘O’,/3”), /?‘2’ and C*

93

used

in the Pitzer model

H+

Na’

K+

NH:

HSO; SO:c10; NO; CH;SO; Cl_ F_ HSO; so:c10; NO; CH,SO; Cl_ FHSO; so:c10; NO; CH,SO; ClFHSO, so:-

c10; NO;

CH,SO, ClF_ Mg2+ HSO; so:c10; NO;

Ca2 +

CH;SO; ClFHSO; so:c10; NO; :pso; F-

0

0 0.0438 0.00819 - 0.00539 - 0.00409 0.0008 -

0.2065 0.0298 0.1747 0.1168 0.1544 0.1775 -

0.5556 0.0000 0.293 1 0.3546 0.4775 0.2945 -

0.0454 0.0196 0.0554 0.0068 0.0787 0.0765 0.0215 - 0.0003 0.0500

0.3980 1.1130 0.2755 0.1783 0.2740 0.2664 0.2 107 0.1735 0.7793 -

-

0 0.00497 -0.00118 - 0.00072 - 0.0024 0.00127 0 0 0 -

-0.0816 0.0581 0.0484 0.0809 -

0.0494 0.1650 0.2122 0.202 1 -

0 0 0 0 -

0.0066 - 0.0046 - 0.00084 0.00093 -

0

-

-0.001161 0 - 0.00003 -0.0041 -0.00301 0 0.025 0.00958 - 0.02062 -

0 0

0 0 0

0.0409 - 0.0103 -0.0154 0.066 1 0.0522 -

0.6590 -0.0194 0.1120 0.1910 0.1918

0.4746 0.22 10 0.4961 0.3671 -

1.7290 3.3430 2.0090 1.5847 -

0 - 37.23 0 0 -

0.3524 -

1.6815 -

0 -

0.2145 0.2000 0.4511 0.2108

2.5300 3.1973 1.7560 1.4090 -

0 - 54.24 0 0 -

0.3159 -2.3000

1.6140 0.0000

0 0

0 0

0 0

0.00519 0 0 - 0.005 -0.02014 - 0.000339 1.9

Data principally from Pitzer (1979) and Harvie et al. (1984). Parameters for KF, HN03 and CH,SOSH have been refitted to ionic strength 6 mol kg-r.

and activity coefficients. While most parameters have been determined from EMF measurements other data can also be used, for example solubilities in multicomponent salt solutions (e.g. Harvie and Weare, 1980) and also the partial pressure measurements in the present work. For all conventional strong electrolytes complete dissociation is assumed by the model. However, explicit recognition of association in solution is required for weak acids such as H3P04, and also for H2S04. In this work both an equilibrium constant for the reaction HSO; = H+ + SO:- and appropriate ion interaction parameters are used, following Harvie et al. (1984). For solutions where both H+ and SOi- are present the equations must first be solved iteratively to determine H+, HSO; and SO:- concentrations, then

the activity coefficients of the ions are calculated. Parameters used in this work are listed in Tables 1 and 2. They are drawn from the compilations of Pitzer (1979) and Harvie et al. (1984). The parameters for pure aqueous HN03 have been recalculated to fit osmotic coefficient data to a concentration of 6.0 mol kg- I. The f&n, parameter has also been reevaluated (Clegg and Brimblecombe, 1988). The approach of Pitzer is essentially pragmatic in that no formal conclusion as to ion speciation is implied. For most electrolytes an explicit definition of ion complex species is not required even when they are known to exist. In the case of nitric acid, for example, the stoichiometric mean activity coefficients are known (Hamer and Wu, 1972) and the degree of association in aqueous solution has been measured over a wide range

S. L. CLEGGand P. BRIMBLECOMBE

94

Table 2. Values of the mixed electrolyte parameters Bijand lLijrused in the Pitzer equations

e H+

NH:

Na+ Kf NH: Mg2 + CaZ + Kf NH: Mg’ + Ca’+ NH: Mg2+ Ca’ + $++

Mg’ +

Ca2+

Na+

K+

0.036 0.005 -0.01 0.1 0.092 -0.012 0.07 0.07 0 0.032 0.007

e HSO;

so:c10; NO; :IH_Iso;

HSO;

SO:-

c10;

NO;

CH$O;

-0.0129 - 0.0265 -0.0178 0 0 0 0 0 0

0 0.197 0 0 -0.01 -0.015 -0.055 -0.048 0

-0.016 -

-

-

so:-

0

0.024

c10;

NO;

CH,SO; Cl-

zY”; FNO; CHaSO; ClFCH$O; ClFclFF-

0.02

Mg2+

Ca2 +

- 0.0425

0

-

-

-

H+

Na+

K+

NH:

0

- 0.0094 - 0.006

- 0.0677 0

-

-

-

-

-

-

0.013 Fc10; NO;

-

0 -

0.0014 -

0 -

F-

-0.004 -0.011 - 0.009 -0.011 -0.015 -0.0018 0 -0.012 - 0.007 0 -0.022 -0.025 0 0 -0.012

-

-0.001 -

Cl-

0

-

-

-

-

0

-

-0.018

-

0.016

- 0.006

- 0.006

-

0

-0.017

-

-

-

Data principally from Pitzer (1979) and Harvie er al. 1984). See text for estimation of new parameter values.

of concentrations (Redlich and Hood, 1957). Since equilibrium partial pressures of HN03 over aqueous solutions are also known (Davis and DeBruin, 1964), the activity coefficients of all species in solution can be calculated, including that of the undissociated molecule. However, HN03 is treated by the Pitzer model as a completely dissociated electrolyte.

3. EXPERIMENTAL Using Equation (3) in conjunction with the Pitzer activity coefficient model it should be possible to calculate equilibrium partial pressures of HN03 over acidified salt mixtures. However, for many common ions the B,, and I(lillparameters are unknown, and no suitable partial pressure data exist to test predictions. Therefore HN03 vapour pressures have been measured over high ionic strength solution mixtures at 25°C. 1 atm pressure to provide a severe test of the model and, if possible, estimate unknown model parameters. In addition to HNOJ the test solutions contained the following ions: NH:, SO:-, Na+, K+ and Mg*+. Both

NH: and SOi- are major tinstituents of the tropospheric background aerosol and are also important in areas affected by pollution. The other three ions are major constituents of seasalt. Chloride was not included because the experimental ;;ipa&us could not distinguish between different gas phase All test solutions were prepared at a constant total ionic strength of 6.0 mol kg- I. Experiments conducted at constant water activity might be considered more relevant, since this represents equilibrium with a constant r.h. However, water activities cannot be calculated with complete accuracy. More importantly, this approach would be inconsistent with that universally adopted by solution chemists, i.e. experiments conducted at constant ionic strength. The reason for this is that ion interactions and/or ion pairing in solution are functions of ionic strength, and so results at constant water activity, but variable ionic strength, would be very difficult to interpret. No attempt was made to use atmospherically realistic test solutions. As an aid to model parameterization, and for experimental simplicity, the solutions contained only a few ions, with very high H+ ion concentrations. Each solution was prepared by weight from stock solutions made up from dried analytical grade chemicals and deionized water. The apparatus used in these experiments has been fully

Equilibrium partial pressures of strong acids over concentrated saline solutions described in Clegg and Brimblecombe (1985b). It is based on the dynamic method of Scarano et al. (1971). Briefly, a flow of inert Nz gas was used to strip small amounts of acid gas from above 50 ml of test solution. Care was taken to ensure that equilibrium betweenaqueous and gas phases was maintained. The stripped acid was removed from the equilibrated gas stream by dissolution into a known volume of unbuffered solution (KCl) and the pH change measured. The relationship between the hydrogen ion concentration in the KC1 solution and pH was known from a calibration experiment (Clegg, 1986). The equilibrium partial pressure of the acid gas was calculated from the pH change and volume of gas stripped, with appropriate corrections for temperature and pressure. The apparatus may be used to measure partial pressures between about lo-’ and lo-“ atmospheres. It was developed and tested using solutions of pure aqueous HCI for which partial pressure data already exist (Clegg, 1986; Clegg and Brimblecombe, 1986). Measurements of the partial pressure of HN03 over aqueous nitric acid solutions agree well with the data of Davis and DeBruin (1964). It was found previously that measurement errors have both random and systematic components (Clegg and Brimblecombe, 1985b; Clegg, 1986). The most important systematic error was caused by a transfer of small quantities of test solution as an aerosol into the solution containing the pH electrode. The magnitude of this error was estimated from experiments using non-volatile acids. The results

were treated as for normal test solutions and an apparent partial pressure, due to aerosol transfer, was calculated. Its value in these experiments was estimated at 2.8 x lo-’ atm (mol kg-’ .+)-I. Random errors were estimated from test data for 10 pure aqueous HCl solutions of concentration 0.5a.85 mol kg- I, and vary from approximately 30 to SoA, respectively over the range of solutions studied. A complete discussion of experimental error is given by Clegg (1986). At low measured partial pressures the systematic error becomes large. Accordingly all measured partial pressures were corrected for this. An upper error bar was then assigned, equal to the greater of (i) the uncorrected measured partial pressure or (ii) the corrected partial pressure plus random error. The lower error bar was assigned equal to the corrected partial pressure less the random variation. The estimated errors in these experiments range from about 1.7 x 10M6atm (8 “/,) for a measured partial pressure of 21 x 10e6 atm, to 0.47 x 10e6 atm (104 %) for a measured partial pressure of 0.45 x 10m6atm. It must be emphasized that there is no experimental point for which the uncorrected partial pressure lies outside the error bars of the corrected value.

4. RESULTS The results of the experiments are presented graphically, and in table form in Appendix 1 where the composition of each test solution is given, together with its calculated water activity. The graphs show partial pressures, with error bars, and also partial pressures calculated using the Pitzer model and two values of the Henry’s Law constant: an upper value of 2.37 x lo6 mol’kg-* atm-’ derived from the data of Davis and DeBruin (1964), and a lower value of 2.19 x lo6 mo12 kgW2 atm-‘, calculated from a partial pressure measurement over pure aqueous HNOj made with the apparatus used in this study. It is likely that the higher value of K,, based on a larger number of measurements, is more nearly correct.

0,O2,O 4,O mH+--

95

6,0

Fig. 1. Partial pressure measurements for System 1. The measured (corrected) partial pressures are marked as open circles, with error bars. The boxes show the partial pressures calculated for each solution using the Pitzer equations and two values of K : 2.37 x lo6 and 2.19 x 106mo12kgm2atm-‘. These partial pressures are calculated with all unknown interaction parameters set to zero. Best fit partial pressure curve calculated using: $H,.,~,NO equal to - 0.003, K, equal to 2.35 x 10d mo12kg-’ atm-‘.

4.1. System 1: H+Na+NO; The results for this system (Fig. 1) agree well with predicted partial pressures, although there appears to be a progressively greater error toward higher Na+ concentrations. All necessary parameters except $H.NhNO,are known for this system. The deviation of the predicted from the measured partial pressure indicates that this parameter is probably significant. Its value was estimated as - 0.003 simply by substituting different values into the Pitzer equations to obtain the best fit. 4.2. System 2: H+NO;

SOi-

The results show a progressive underestimate of predicted partial pressure with increasing H2S04 concentration, Fig. 2. However, the agreement between measured and predicted partial pressures is good considering the following parameters for the system are unknown: eN03,S04t eHS04,N039 *HSO,.NO,,H~ $H,so,,No,. The corresponding system for HCl was found to be insensitive to variations in the parameter $so,,c,,H (Harvie et nl., 1984), so it was assumed that * SO,,NO,,His also equal to zero. There is insufficient data to uniquely determine the remaining parameters. However, eNO,,SO, and also $NO,,SO,,Kmay be determined from published solubility data for the system K+/NO;/SOi(Linke, 1965). The parameters take the values &o,,so.equal to 0.0926and +No,,so&equal to - 0.0046. The validity ofthe parameter eNo,,so.may be checked using similar solubility data for the system Na+/NO; /SO:- saturated with respect to Na2S04 (Linke, 1965). The value of ONO,,SO,derived above,

S.

96

L.

CLEW

and P.

BRIMBLECOMBE

show that the contribution of $Na,~u,,~o, to Yu~o,can probably be disregarded. Because the parameter 0 N&NH,makes no contribution to yuNoI, I//n,Nu,,No, must have a significant value. It was estimated to be equal to -0.01 by substitution. 4.4. System

4: H+ K+ Na+ NO;

to System 3, with K+ present instead of the NH: ion. These two ions have very similar effects, as was the case for the corresponding mixtures with the Cl- anion (Clegg and Brimblecombe, 1988; Clegg, 1986). Fitting of the unknown parameter $nxNo, suggests a value of about -0.01, the same as for $H,NH,,NO,. This value has been adopted in subsequent calculations. This system, Fig. 4, is analogous

mNO$ A Fig. 2. Results for System 2. Fitted curve calculated using B~o,,so, equal

to 0.0926, Ku equal to 2.35 x lo6 mol’ kg-’ atm ‘. Seecaption of Fig. 1 for explanation of symbols.

together with I(/NO,,sO,,Naequal to 0.0066, yields a constant activity product for the saturated salt Na2S04, suggesting that the parameters derived are not merely specific to the K+/NO; /SO:- system. A complete description of the procedure for determination of these parameters is given in Clegg (1986). Using ONO,,SO,in the Pitzer model to predict the partial pressures in System 2 slightly improves the fit to the data but the other unknown parameters are also likely to have significant values. 4.3. System

4.5. System

5: H+ Na+ NH: NO; SOi-

The data for this complex mixture, Fig. 5, are fitted well without any of the previously unknown interaction parameters. This result is particularly noteworthy since the system, essentially a mixture of HN03 and (NH,)$04, is of great interest in certain polluted environments (Bassett and Seinfeld, 1983). The use of the interaction parameters estimated above does not significantly improve the fit to the data. 4.6. System

6: H+Mg’+NO;

SOi-

Measurements were made for this system, Fig. 6, to determine the effects of the major seawater cation

3: H+ NH: Na+ NO;

This system shows the effect of two positive ions Na+ and NH: on the partial pressure of HN03, Fig. 3. The presence of NH: depresses the partial pressure relative to Na+ in a similar way to that found for the corresponding system containing the chloride anion (Clegg and Brimblecombe, 1988). There appears to be a small overprediction of partial pressure for the systems containing the highest concentrations of NHf The unknown parameters for this system are: ONa,NH, and $NaNu,,NoJ. Sensitivity tests h.NH,,NO,~

mNHiFig. 3. Results for System 3. Rest fit curve calculated with: $ H,NH,,No, equal to -0.01, K, equal to 2.35 x lo6 mot’ kg-* atm- r. See caption of Fig. 1 for explanation of symbols.

Fig. 4. Results for System 4. Rest fit curve calculated using: +u,k,NO, equal -0.01, K, equal to 2.35 to x lo6 mol’ kg-’ atm- I. See caption of Fig. 1 for explanation of symbols.

=

0.01 o,o

I,0 mNa’-

2,o

Fig. 5. Results for System 5. See caption of Fig. 1 for explanation of symbols.

Equilibrium partial pressures of strong acids over concentrated saline solutions

91

dated for solution mixtures up to ionic strengths of the order of 10 mol kg-’ (Harvie et al., 1984), and reproduces the properties of pure electrolytes to concentrations of about 6 mol kg-‘. These concentrations correspond to equilibrium relative humidities of about 70-75 “i,. However, aerosols often contain solutes such as NH,NO, and (NH,)$O, which remain in aqueous of greater than solution to concentrations 20 mol kg-‘. For such systems it is necessary to use polynomial equations to calculate pure electrolyte properties, coupled with a simple model for mixtures, for example that of Bromley (1972). However, it has 2,0 4,0 6,0 already been pointed out that such models are marmH’. kedly less accurate than that of Pitzer, even at low ionic strengths. In addition they have no formal method of Fig. 6. Results for System 6. See captreating ion association equilibria which are important tion of Fig. 1 for explanation of symbols. in acidified sulphate systems. Further data are needed for HNO,/H,SO, mixturesand thosecontaining the NH: and K+ ions. Data Mg2 + on an HN03/H2S04 mixture. In addition to are available for the osmotic coefficient of NH,HSO, those listed in section 4.2 the parameters $NOJ,sO,,Mg, (Tang and Munkelwitz, 1977) but have not been used I(/H,Mg,.o,and +No,,Hso,,Naare unknown and therefore to estimate interaction parameters in this work. Partial set to zero. The experimental measurements and pressure measurements for systems containing both calculated partial pressures agree closely, apparently Cl- and NO; would also be valuable from the point of suggesting that these parameters may be negligible. view of realistic atmospheric solution mixtures. However, this must remain uncertain as the results for However, it is clear from the many partial pressure System 2, similar to System 6 except for the presence of measurements made in this study that calculations can Mg 2+, showed that some of the interaction parameters be made with some confidence for these mixtures even common to the two systems have significant values. when some of the required interaction parameters are unavailable. Measurements of the partial pressure of HCl over acidified salt solutions, together with an 5. DISCUSSION identicai theoretical treatment, are described in a companion paper, Clegg and Brimblecombe (1988). The systems studied here are particularly relevant to important atmospheric systems, for example the NH,/HNO,/H,SO, mixtures modelled by Bassett 6. CONCLUSIONS and Seinfeld (1983, 1984). The results show good agreement between measured and predicted partial The partial pressure measurements of HNO, that have been made in this work show that the Pitzer pressures, despite the lack of many interaction parameters. model may be used, with a thermodynamic Henry’s The Pitzer activity coefficient model is more accurate Law constant, to calculate partial pressures of HNO, than, for example, that of Meisner and Kusik (1973) over acidified salt mixtures to high concentrations. used by Bassett and Seinfeld (1983), and that of The model is accurate and robust, and has been verified Bromley (1972) used by Saxena et al. (1986). The here for systems of ionic strength up to 6.0 mol kg- ‘. behaviour of many electrolytes which are known to be Its adoption for atmospheric chemical calculations is partially associated in solution, such as the bivalent recommended. The simple theoretical description of metal sulphates and HNO, itself, are well represented strong acid solubility presented here should enable by the model, thus simplifying calculations. All the strong acid gas equilibria in atmospheric aerosol work presented here has all been carried out at 25°C systems to be modelled more accurately than by other although for modelling realistic atmospheric systems a methods. temperature range of about &3o”C is required. Values for the following interaction parameters have Temperature derivatives of the pure electrolyte par- been estimated from the partial pressure results and ameters are available for most solutes (Pitzer, 1979). published solubility data: tju,N=,No,, -0.003; &o,,so, , Variations in the Bijand tiije parameters are thought to 0.0926; J&NO,,SO,,K, -0.0046; +&No,,so,,N~~ O.OCW be small close to 25°C (Pitzer, 1979). Henry’s Law * H.NH,,NOI, -0.01; *H,K,N~,, -0.01. Values are reconstants for HNO, and the hydrohalic acids over the quired for the parameters tinso,,No,and $ uso,,~o,,n in are presented by order to improve the fit for H+/NO;/SO: temperature range 0-40X - systems. Brimblecombe and Clegg (1987). The results presented above do not test the model at Acknowledgements-This work was supported by NERC extreme concentrations. It has previously been vali- grant GT4/82/APPS/20.

S. L. CLEGGand P. BRIMBLECOMBE

98 REFERENCES

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99

Table 3A. Results for System 3 (H’ NH: Nat f = 6.0 mol kg- ’ mH+

mNa’

mNH2

2.80 2.80 2.80 2.80 2.80 2.80

3.20 2.50 1.80 1.20 0.60 0.00

0.00 0.70 1.40 2.00 2.60 3.20

mNO; 6.00 6.00 6.00 6.00 6.00 6.00

YHNO,

1.005 0.955 0.907 0.868 0.830 0.794

a,

0.794 0.800 0.805 0.810 0.814 0.819

NO; ),

PHNOJ,

6.83 5.78 4.83 4.57 4.04 3.42

Refer to Table 1A caption for explanation of column headings.

APPENDIX 1 Table IA. Partial pressure measurements for System 1 (H’ Na+ NO; ), ionic strength (I) is equal to 6.0 mol kg-’ r?lH+

mNa+

mN0;

YHNO,

a,

PHNO~

6.00 5.20 4.40 3.60 2.80 2.00

0.00 0.80 1.60 2.40 3.20 4.00

6.00 6.00 6.00 6.00 6.00 6.00

1.169 1.126 1.084 1.044 1.005 0.968

0.159 0.766 0.714 0.783 0.794 0.807

20.90 16.80 12.30 9.01 6.83 4.16

Units: all concentrations (m)mol kg- ‘, all partial pressures 1O-.6atm. The values of YuNoJ and water activity (a,) given for each solution were calculated using the Pitzer model and the parameter set given in Tables 1 and 2. The calculated concentration of free H * ions (H$) and true ionic strength are listed for solutions containing SO:-. Superscript ‘T’ denotes a stoichiometric or total concentration.

Table 4A. Results

mH+ 2.80 2.80 2.80 2.80 2.80 2.80

for System 4 (H + K” Na’ NO;), I = 6.0 mol kg-’

mNa+ mK* 3.20 2.85 2.50 1.80 1.50 1.20

mN0;

YHN~,

a,

6.00 6.00 6.00 6.00 6.00 6.00

t.005 0.978 0.952 0.902 0.882 0.862

0.794 0.798 0.803 0.8 10 0.813 0.816

0.00 0.35 0.70 1.40 1.70 2.00

Refer to Table 1A caption for explanation headings.

Table 2A. Results for System 2 (H + NO; SO: - ), stoichiometric I = 6.0 mol kg- 1 mHT

mH;

mS0:

6.00 5.63 5.27 4.90 4.53 4.17

6.00 5.41 4.80 4.16 3.51 2.82

O.ooO 0.367 0.733 1.100 1.467 1.833

mN0; 6.00 4.90 3.80 2.70 1.60 0.50

YHNO,

a,

I

1.169 1.103 1.021 0.921 0.824 0.713

0.759 0.783 0.810 0.839 0.869 0.899

6.00 5.56 5.07 4.53 3.94 3.31

PHNO, 20.90 15.10 9.99 5.71 2.26 0.82

Refer to Table 1A caption for explanation of column headings.

Table 5A. Results for System 5 (H+Na+ NHf NO; SO:- ), stoicbiometric I = 6.0 mol kg’ mHT

mH$

mSO,7

mNa+

mNH=

mN0;

4.40 4.40 4.40 4.40 4.40 4.40

4.40 4.3 1 4.24 4.17 4.10 4.03

0.000 0.133 0.233 0.333 0.433 0.533

1.60 1.20 0.90 0.60 0.30 0.00

0.000 0.267 0.467 0.667 0.867 1.070

6.00 5.60 5.30 5.00 4.70 4.40

YHNO,

a,

f

1.084 1.041 1.010 0.979 0.948 0.919

0.774 0.784 0.792 0.799 0.807 0.815

6.00 5.81 5.68 5.54 5.40 5.26

Refer to Table 1A caption for explanation of column headings.

PHNO, 12.30 11.40 9.64 9.11 7.53 6.97

PHNO, 6.83 4.99 4.68 3.76 5.63 4.20 of cohrmn

S. L.

100 Table 6A. Results

CLEGG

and P.

BRIMBLECOMBE

for System 6 (H+ Mg2+ NO; SO:-), = 6.0 mol kg- I

mHT

mHf

mSOf

mMg’+

6.00 5.30 4.60 3.90 3.20 2.50 1.70

6.00 5.20 4.38 3.56 2.14 1.94 1.07

0.000 0.175 0.350 0.525 0.700 0.875 1.080

0.000 0.175 0.350 0.525 0.700 0.875 1.075

mN0; 6.00 5.30 4.60 3.90 3.20 2.50 1.70

stoichiometric

YHNO,

a,

1

1.169 1.145 1.107 1.056 0.988 0.897 0.148

0.159 0.171 0.191 0.819 0.843 0.869 0.901

6.00 5.79 5.56 5.32 5.08 4.87 4.74

Refer to Table 1A caption for explanation of column headings.

W-3

20.90 15.00 11.40 7.09 4.24 2.16 0.45

I