Measurement of Activity Coefficients of Amino Acids in Aqueous

Jun 15, 1996 - Vera, 1995). In addition, amino acids are the building blocks of other biomolecules such as peptides and proteins. Thus, it is importan...
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Ind. Eng. Chem. Res. 1996, 35, 2735-2742

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Measurement of Activity Coefficients of Amino Acids in Aqueous Electrolyte Solutions: Experimental Data for the Systems H2O + NaCl + Glycine and H2O + NaCl + DL-Alanine at 25 °C Mohammad K. Khoshkbarchi and Juan H. Vera* Department of Chemical Engineering, McGill University, Montreal, Quebec, Canada H3A 2A7

Electrochemical cells with two ion-selective electrodes, a cation and an anion ion-selective electrode, vs a double-junction reference electrode were used to measure the activity coefficients of amino acids in aqueous electrolyte solutions. Activity coefficient data were measured for the H2O-NaCl-glycine and H2O-NaCl-DL-alanine systems at 25 °C. The results of the activity coefficients of glycine were compared with values reported in the literature which were obtained by the isopiestic method and by electrochemical cells with one single sodium ion-selective electrode vs a reference electrode. It was found that the method proposed in this study provides more accurate activity coefficient data. From a theoretical point of view, it is shown that the present method is exact and it should be employed for future measurements. Introduction Separation and concentration of biochemicals are two current subjects of interest due to their high cost in comparison to the total manufacturing cost (Eyal and Bressler, 1993). The behavior of biomolecules in mixtures is affected by many factors such as chemical structure, pH, surface charge distribution, and type of electrolyte present and its concentration. The effect of the presence of electrolytes in solutions of biochemicals is of interest in a number of separation processes, such as the reverse micellar extraction of amino acids and proteins which may not occur without the presence of an electrolyte (Marcozzi et al., 1991; Khoshkbarchi and Vera, 1995). In addition, amino acids are the building blocks of other biomolecules such as peptides and proteins. Thus, it is important to study their behavior in aqueous systems containing electrolytes. The knowledge of the activity coefficient of amino acids in aqueous electrolyte solutions is of primary importance. Lack of accurate activity coefficient data makes the design of equilibrium-based separation processes such as extraction, precipitation, etc., inaccurate and empirical. The measurement of the activity coefficients of amino acids in aqueous electrolyte solutions has been the subject of some studies. Apart from some old and inaccurate methods such as the freezing point method (Cohn and Edsall, 1965), two major techniques which have been commonly used are the isopiestic method (Bower and Robinson, 1965; Schrier and Robinson, 1971, 1974) and the electrochemical method (Kelley and Lilley, 1979; Briggs et al., 1974; Rodrı´guezRaposo et al., 1994). Electrochemical cells are an essential tool for experimental studies of solutions containing electrolytes and have been widely used for the measurement of their thermodynamic properties (Haghtalab and Vera, 1991a; Manohar et al., 1992). In contrast to the isopiestic method, electrochemical methods deal directly with the thermodynamic properties of the electrolyte. A discussion of the accuracy, advantages, and drawbacks of these methods is beyond the scope of this work. In this study, we focus our attention on the electrochemical methods using ion-selective electrodes (ISEs). The application of electrochemical cells with ion-selective * To whom correspondence should be addressed.

electrodes to the measurement of the activity coefficient of electrolytes has received renewed attention as a result of recent improvements in the production of these electrodes. All studies which have reported the measurement of the activity coefficients of amino acids in aqueous electrolyte solutions using ion-selective electrodes (Briggs et al., 1974; Phang and Steel, 1974; Rodrı´guez-Raposo et al., 1994) have used a cation ion-selective electrode only, vs a reference electrode. The difference between the potentials of the ion-selective electrode and the reference electrode was calibrated with the mean ionic activity coefficient of the electrolyte in an aqueous electrolyte solution without the amino acid. Since in this method of measurement the effect of counterion of the electrolyte is neglected, it may lead to erroneous results. In this work, the activity coefficients at 25 °C of sodium chloride in the H2O-NaCl-glycine and H2ONaCl-DL-alanine systems were measured using an electrochemical cell with two ion-selective electrodes. The electrochemical potentials (emf’s) of both sodium and chloride ion-selective electrodes vs a double-junction reference electrode (DJ) were measured in the following cells: sodium ISE|NaCl (mS) + amino acid (mA)|KNO3, DJ; and chloride ISE|NaCl (mS) + amino acid (mA)|KNO3, DJ. The emf’s thus measured were then converted by means of a calibration curve to the mean ionic activity coefficients of the sodium chloride and subsequently to the activity coefficients of amino acid at different molalities. The theoretical basis of the method of measurement used here and its advantages and drawbacks over other methods are discussed below. Theoretical Basis of the Experimental Method The potential of an ion-selective electrode vs a reference electrode is related to the ionic activity coefficient of the corresponding ion by the Nernst equation. For a system containing an electrolyte at molality mS, the Nernst equations for the potentials of a cation and an anion ion-selective electrode, E+ and E-, in the electrochemical cell of type 1, ISE|electrolyte (mS)|reference electrode, can be written as

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2736 Ind. Eng. Chem. Res., Vol. 35, No. 8, 1996

E+ ) E+° +

RT ln(m+γ+) z+ F

(1)

E- ) E-° -

RT ln(m-γ-) z- F

(2)

where γ is the activity coefficient, R is the universal gas constant, T is the absolute temperature, F is the Faraday number, z is the charge number, and subscripts plus and minus sign denote the cation and anion, respectively. The terms E+° and E-° in eqs 1 and 2 are a linear combination of the junction potential, reference electrode potential, and ion-selective electrode potential: ISE

(3)

E-° ) ERef + EJ + E-ISE

(4)

Ref

E+° ) E

J

+ E + E+

where superscripts Ref and J denote the potential of the reference electrode and junction potential and the terms E+ISE and E-ISE include all asymmetry, internal solutions, and reference potentials of the cation and anion ion-selective electrodes, respectively. By subtracting eq 1 from eq 2 and considering eqs 3 and 4, we obtain a relation between the difference of the potential of a cation and an anion ion-selective electrode with the mean ionic activity coefficient of the electrolyte, γ(1)(, in cell type (1) as

E+ - E- ) (E+ISE - E-ISE) + RT z+ + zln[(ν+ν+ν-ν-)1/(ν++ν-) msγ(1)(] (5) F z+z-

(

)

For a 1:1 electrolyte, eq 5 simplifies to

E+ - E- ) (E+ISE - E-ISE) +

2RT ln(msγ(1)() (6) F

Equation 6 is only valid when the same reference electrode is used vs both the anion and cation ionselective electrodes. Equation 6 in a more general form can be written as

∆E(1) ) E° + S ln(msγ(1)()

(7)

where ∆E ) E+ - E- and E° ) E+ISE - E-ISE and the superscript (1) refers to the electrochemical cell number 1, containing electrolyte and water only. For the sake of generality, instead of the Nernstian slope of electrodes in eq 6 the general electrode slope, S, is used in eq 7. From eq 7, it is evident that the values of S and E° can be calculated from a linear fitting of the values of ∆E vs ln(mSγ(1)(), with values of γ(1)( obtained at each molality mS from the literature (Zemaitis et al., 1986). The derivation of the above equations is valid for systems containing only one electrolyte in water. The potential of an electrochemical cell changes in the presence of other solutes as a reflection of the change in the mean ionic activity coefficient of the electrolyte in the presence of other solutes. This is due to the interactions between the molecules of electrolyte, other solutes, and water. By using the same method as above, we can measure the difference between the potentials of a cation and an anion ion-selective electrode, ∆E2, in type 2 cell: ISE|electrolyte (mS) + solute (mA)|reference electrode. This potential difference is related to the mean ionic activity coefficient of the electrolyte at

molality mS in the presence of a solute at molality mA, γ(2)(, by

∆E(2) ) E° + S ln(msγ(2)()

(8)

It should be noted that the values of S and E° are the same in both eqs 7 and 8 provided that the same ionselective electrodes and reference electrode are used in both cells of type 1 and of type 2. Subtracting eq 8 from eq 7 and rearranging give

ln

( ) γ(2)(

γ(1)(

)

∆E(2) - ∆E(1) S

(9)

Equation 9 provides a relation between γ(2)( and γ(1)( useful for correlation purposes. We emphasize that the method described above uses two ion-selective electrodes, against a reference electrode, to measure the mean ionic activity coefficient of an electrolyte in the presence of other nonelectrolyte solutes. An alternative method, which is also theoretically sound, is to use the anion ion-selective electrode as the reference electrode for the cation ion-selective electrode. The governing equations of this method are the same as eqs 5 and 7 and have been discussed in detail by Haghtalab and Vera (1991a). Some investigators have applied this method to measure the mean ionic activity coefficients of electrolytes in aqueous solutions containing other electrolytes and nonelectrolytes (Haghtalab and Vera, 1991b; Han and Pan, 1993). In contrast to the last two methods discussed above, in most of the previous studies for systems containing water, electrolyte, and another solute, only one ionselective electrode (usually a cation ion-selective electrode) vs a reference electrode was used as a measurement of the mean ionic activity coefficient of the electrolyte in the presence of another solute. This method has been used to measure the mean ionic activity coefficient of the electrolyte in water + electrolyte + another electrolyte (Butler and Huston, 1970; Reddy and Ananthaswamy, 1990; Manohar et al., 1992), water + electrolyte + another solvent (Spink and Schrier, 1970), and water + electrolyte + amino acid (Briggs et al., 1974; Phang and Steel, 1974; Rodrı´guezRaposo et al., 1994). There are clear shortcomings associated with employing only one ion-selective electrode, vs a reference electrode, to monitor the mean ionic activity coefficients of an electrolyte in the presence of a nonelectrolyte. The main drawback of this method is that it identifies the mean ionic activity coefficient of the electrolyte, γ(, with the potential of a single ionselective electrode while, as shown by eqs 1-7, the value of γ( can only be obtained using two ion-selective electrodes. In fact, according to the Nernst equation, as is shown in eq 1, the response of a cation ion-selective electrode is only related to the ionic activity coefficient of the cation, whereas the mean ionic activity coefficient is a combination of the anion and cation activity coefficients defined by

γ( ) (γ+ν+γ-ν-)1/(ν++ν-)

(10)

with ν being the stoichiometric number per mole of electrolyte. From eqs 3-5 and 10, it can be inferred that attributing the response of only the cation ionselective electrode to the mean ionic activity coefficient of the electrolyte can be correct only when either the anion activity coefficient remains constant over the entire range of electrolyte concentration or the activity coefficients of both the anion and the cation are always

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Ind. Eng. Chem. Res., Vol. 35, No. 8, 1996 2737

equal, over the entire range of electrolyte concentration. We have recently reported a new method to measure the activity coefficients of ions (Khoshkbarchi and Vera, 1996) and showed that none of these assumptions hold. The anion and cation of an electrolyte in a solution have different activity coefficient values. It was also shown that the activity coefficients of both the anion and the cation vary differently with changes in electrolyte concentration. It should be mentioned that the incorrect assumption of equality of the activity coefficients of the cation and anion is widely spread in the literature (Ashrafizadeh et al., 1993) and that many models for activity coefficients of electrolytes in solutions give this prediction (Pitzer, 1980; Chen et al., 1986). Another major deficiency associated with the use of one single ion-selective electrode, vs a reference electrode, instead of using two ion-selective electrodes, each vs a reference electrode, is the uncompensated effect of the junction potential. The junction potential, which arises along the reference electrode membrane, is a sensitive function of the solution composition and varies with the change in the solution composition (Bates, 1965). In the method proposed in this work, by using two ion-selective electrodes, as shown by eqs 3-6, the effect of the junction potential is eliminated, whereas, in methods which use one single ion-selective electrode, the effect of the junction potential does not cancel and introduces additional error in the measurements. One important practical shortcoming of the method which uses only a single ion-selective electrode arises from the impossibility of keeping the reference electrode potential constant throughout the course of an experiment. The potential of the commercially available reference electrodes changes slightly with time due to factors such as outgoing of the internal fluid of the reference electrode and the clogging of membrane pores of the reference electrode. This problem, which introduces some additional error in the measurements, can be avoided by employing two ion-selective electrodes, each vs a reference electrode. In fact, by measuring the potentials of both ion-selective electrodes simultaneously vs the same reference electrode, at each electrolyte and nonelectrolyte concentration, if the potential of the reference electrode changes, when subtracting the potentials, as shown by eq 6, the effect of the change is eliminated. Materials and Methods Sodium chloride, glycine, and DL-alanine of 99.9% purity were obtained from A&C American Chemicals, Ltd. (Montreal, Quebec, Canada). Glycine and DLalanine were used as received. Sodium chloride was oven-dried for 72 h prior to use. During the drying period, the salt was taken out of the oven after 48 h and 72 h, and after cooling it in a vacuum desiccator, it was weighed. After 48 h, no change in its weight was observed. A Ross sodium ion-selective electrode glass body Model 84-11, a chloride ion-selective electrode polymer body Model 94-17, and a double-junction reference electrode Model 13-620-46 were obtained from Orion (Boston, MA). An Orion pH/ISE meter (Boston, MA) Model EA 920 with a resolution of (0.1 mV was used to monitor the emf measurements with two BNC (Bayonet Neil-Concelman) connectors for ion-selective electrodes and two pin-tip connectors for the reference electrode. All the solutions were prepared based on molality, and the water was also weighed. The compositions of the initial solutions were accurate within (0.01 wt %. In

all experiments, deionized water with a conductivity of less than 0.8 µS/cm was used. Before using it to prepare the samples, the distilled water was passed through ionexchange columns, Type Easy pure RF, Compact Ultrapure Water System, Barnstead Thermoline. The conditioning procedure of electrodes was followed exactly according to the manufacturer’s instructions. The experiments were done by measuring the emf of both the cation and the anion ion-selective electrodes against a double-junction reference electrode in a jacketed glass beaker containing 200 mL of solution. To avoid the bias potential between different reference electrodes, in each experiment, the responses of both chloride and sodium ion-selective electrodes were measured at the same time vs the same reference electrode. All the instruments were grounded prior to and during the experiments. During the experiments, in order to minimize the risk of the presence of concentration gradients in the beaker, the solutions were stirred constantly with a magnetic stirrer and the temperature was kept constant at 298.2 ( 0.1 K using a thermostatic bath. Each set of experiments was performed at fixed electrolyte concentration, and the concentration of amino acid was increased by addition of solid amino acid. The error on the molality of the electrolyte caused by the outgoing flow of the internal solution of the reference electrode was minimized by choosing KNO3 solution as the internal outerbody solution of the reference electrode. The readings of the potentiometer were made only when the drift was less than 0.1 mV. Prior to their use, the ion-selective electrodes were tested for reversibility. The tests were performed in four steps by measuring the stabilized potential of both cation and anion ion-selective electrodes vs a double-junction reference electrode. In the first step, the potential readings were performed in a 0.1 m solution of NaCl; in the second step, the electrodes were transferred to a 1.0 m solution of NaCl and their potentials were measured; in the third step, the electrodes were returned to the same 0.1 m solution of NaCl and their potentials were once more measured; and in the fourth step, they were returned to the 1.0 m solution of NaCl used in step 2 and their potentials were measured. The differences in the potential readings of the electrodes in the solutions of equal concentration were (0.05 mV, which indicates the reversibility of the electrodes. For each set of experiments, the electrodes were calibrated by measuring the emf’s of the cell shown in the previous section, without the presence of other solute, and the slope, S, of the electrodes was determined. The typical value obtained for the slope of the electrodes was 51.28 mV with a correlation coefficient of 99.99%, where its theoretical value according to the Nernst equation at 298.15 K is 51.38 mV. Most of the experiments were replicated 3 times, and the data reported are the average of the replicas. Sample variances were obtained from the replicas for each point, and a pooled standard deviation was calculated using these values. The calculated pooled standard deviations for a 95% confidence interval for the values of the ratio of the mean ionic activity coefficients of NaCl in the presence of the amino acid and in the absence of the amino acid, at the same NaCl molality, for the systems H2O + NaCl + DL-alanine and H2O + NaCl + glycine, were (0.004 and (0.003, respectively. Results and Discussion The potentials of a sodium and a chloride ion-selective electrode each, vs a double-junction reference electrode,

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2738 Ind. Eng. Chem. Res., Vol. 35, No. 8, 1996 Table 1. Experimental Data for the Ratio of the Mean Ionic Activity Coefficients of NaCl in the Presence of Glycine and in the Absence of Glycine at Different Molalities of NaCl and Glycine

Table 3. Values of the Parameters of Equation 10 C1 C2 C3 C4 C5 C6 rmsda

γ(2)(/γ(1)( for NaCl, m glycine, m

0.1

0.3

0.5

0.6

1.0

0.1 0.2 0.5 0.8 1.0 1.3 1.6 2.0 2.3 2.6

0.9913 0.9865 0.9553 0.9278 0.9134 0.8932 0.8760 0.8549 0.8417 0.8303

0.9892 0.9823 0.9564 0.9339 0.9220 0.9075 0.8923 0.8757 0.8653 0.8551

0.9923 0.9866 0.9658 0.9491 0.9364 0.9247 0.9114 0.8965 0.8861 0.8768

0.9933 0.9875 0.9695 0.9546 0.9463 0.9372 0.9237 0.9122 0.9008 0.8947

0.9952 0.9914 0.9763 0.9633 0.9559 0.9486 0.9387 0.9297 0.9182 0.9112

Table 2. Experimental Data for the Ratio of the Mean Ionic Activity Coefficients of NaCl in the Presence of DL-Alanine and in the Absence of Glycine at Different Molalities of NaCl and DL-Alanine γ(2)(/γ(1)( for NaCl, m alanine, m

0.1

0.2

0.3

0.4

0.6

0.9

0.05 0.10 0.20 0.30 0.40 0.60 0.80 1.00 1.20 1.40 1.60

0.9961 0.9922 0.9856 0.9807 0.9770 0.9655 0.9534 0.9478 0.9414 0.9377 0.9323

0.9960 0.9940 0.9920 0.9881 0.9851 0.9753 0.9666 0.9579 0.9550 0.9550 0.9490

0.9981 0.9961 0.9913 0.9836 0.9855 0.9750 0.9750 0.9694 0.9656 0.9618 0.9618

0.9952 0.9980 0.9952 0.9913 0.9894 0.9855 0.9800 0.9780 0.9779 0.9741 0.9779

0.9959 0.9960 0.9960 0.9961 0.9961 0.9911 0.9882 0.9894 0.9894 0.9880 0.9910

0.9961 0.9961 0.9961 0.9942 0.9922 0.9941 0.9961 0.9942 0.9912 0.9961 0.9961

were measured at 25 °C for the systems H2O + NaCl + glycine and H2O + NaCl + DL-alanine over the ranges of NaCl molalities from 0.05 to 1 m, DL-alanine molalities from 0.05 to 1.6 m, and glycine from 0.1 to 3 m. Throughout this work, DL-alanine is used to indicate the R form of this amino acid, namely, DL-R-alanine, unless otherwise mentioned. The potentials were measured in the following two cells: sodium ISE|NaCl (mS) + amino acid (mA)|KNO3, DJ; and chloride ISE|NaCl (mS) + amino acid (mA)|KNO3, DJ. The potential values thus obtained were converted, using eq 9 and the value of S obtained from the cells without amino acid, to the ratio of the mean ionic activity coefficients of the electrolyte, in the presence of amino acid, to that without the presence of amino acid at the same electrolyte molality. These values were collected for different electrolyte and amino acid concentrations. The experimental data thus obtained were fitted to a virial expansion series similar to that suggested by Scatchard and Pentiss (1934). To obtain a better curve fit for the experimental data, an extra term was added to account for the higher order terms of the virial expansion:

ν ln

γ( γ(1)(

) C1mA + C2mSmA + C3mA2 + C4mAmS2 + C5mA3 + C6mSmA2 (11)

The experimental data measured are presented in Tables 1 and 2. The activity coefficient of amino acid is related to the mean ionic activity coefficient of the electrolyte through the cross-differential relation:

(

ν

)

∂ ln γ( ∂mA

mS,T,P

)

(

)

∂ ln γA ∂mS

(12)

mA,T,P

Combining eqs 11 and 12, the activity coefficient of

glycine

DL-alanine

-0.232 720 0.178 438 0.037 054 -0.050 052 -0.002 981 -0.016 483 0.001 0

-0.165 701 0.298 762 0.038 005 -0.159 047 -0.002 202 -0.024 528 0.000 7

amino acid, γ(2)A, can be calculated as

ln

γ(2)A (1)

γ

A

1 1 ) C1mS + C2mS2 + 2C3mAmS + C4mS3 + 2 3 3C5mA2mS + C6mS2mA (13)

The values of the parameters of eq 11 were evaluated by a least-squares error analysis of the experimental data. The results of the evaluation of parameters using eq 11, together with the root-mean-square deviation for both H2O + NaCl + glycine and H2O + NaCl + DL-alanine systems, are presented in Table 3. As can be seen from Table 3, the C1 coefficient in eqs 11 and 13, which represents the pairwise interactions between amino acid molecules and electrolyte molecules, is negative for both glycine and DL-alanine. This indicates an attractive interaction between NaCl and both DLalanine and glycine. The negative values of the C1 coefficient also indicate that the presence of NaCl decreases the activity coefficient and consequently represent a salt-in effect for both amino acids. The values of the ratio of the mean ionic activity coefficients of NaCl for the system H2O + NaCl + glycine to that, at the same electrolyte molality, for the system H2O + NaCl obtained using three different approaches were compared. The experimental data compared are those measured in this study using an electrochemical cell with a sodium and a chloride ionselective electrode each vs a reference electrode, those measured in this study using only sodium ion-selective electrode vs a reference electrode, and those reported by Schrier and Robinson (1971) measured by the isopiestic method. The absolute mean relative difference between the data obtained by the isopiestic method and those obtained in this work using a cation and an anion ion-selective electrode was found to be 0.4%. The absolute mean relative difference between the data obtained by the isopiestic method and those obtained in this work using only a cation ion-selective electrode was found to be 1.7%. The absolute mean relative deviation (amrd) is defined as

amrd )

(

∑

1 n

)

(γ(2)(/γ(1)()ISO - (γ(2)(/γ(1)()ISE (γ(2)(/γ(1)()ISO

The ratios of the mean ionic activity coefficients of NaCl in the presence of glycine and in the absence of glycine, at 0.5 m NaCl, as a function of glycine molality obtained from the abovementioned methods are compared in Figure 1. As shown in Figure 1, due to the elimination of the reference electrode potential in the method using two ion-selective electrodes, the experimental data obtained were more smooth and less scattered. This comparison shows that the values measured using the isopiestic method are in better agreement with those obtained using two ion-selective electrodes and are different from those obtained using only one ionselective electrode. This can be considered to be proof for the reliability of the method proposed in this work

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Ind. Eng. Chem. Res., Vol. 35, No. 8, 1996 2739

Figure 1. Comparison of the ratio of the mean ionic activity coefficients of NaCl in the presence of glycine and in the absence of glycine, at 0.5 m NaCl. (*) Isopiestic method (Schrier and Robinson, 1971). (b) An electrochemical cell with a sodium and a chloride ion-selective electrode each vs a reference electrode and (9) an electrochemical cell with only a sodium ion-selective electrode vs a reference electrode.

for measuring the activity coefficients of electrolytes in the presence of a nonelectrolyte. Phang and Steel (1974) also measured the mean ionic activity coefficients of NaCl in the system H2O + NaCl + glycine using only a sodium ion-selective electrode against a reference electrode and compared their measured values with those obtained by the isopiestic method. Their comparison also showed an inconsistency between values obtained by their method and those obtained by the isopiestic method. It should be mentioned that, although the method which uses two ion-selective electrodes has an exact theoretical basis and leads to more accurate measurements, the experimental data obtained using one ion-selective electrode are close to those obtained by the other methods. This suggests that, in the absence of experimental data measured with either the isopiestic method or electrochemical cells with a cation and an anion ion-selective electrode, the experimental data obtained using a single ion-selective electrode can be used with caution. At this point, it is important to mention that the reliability of the measurements performed by ion-selective electrodes depends to a large extent on their thermodynamic reversibility. The irreversibility of an ion-selective electrode can be originated from different sources such as the type of cell and electrode reactions, the unsuitability of the electrode for certain measurements, and even deficiencies in the fabrication of the electrode. Unfortunately, apart from some approximate methods, it is not usually an easy task to experimentally monitor the thermodynamic reversibility of an ion-selective electrode. This suggests that the measurements with ion-selective electrodes for complex systems should be performed cautiously. However, investigations have shown that anion ion-selective electrodes are generally more reversible than cation ionselective electrodes (Covington, 1969). Therefore, once more, it can be argued that the use of an anion and a cation ion-selective electrode instead of one single cation ion-selective electrode leads to more accurate measurements. Figure 2 shows, using the molality of NaCl as a parameter, the ratio of the mean ionic activity coef-

Figure 2. Effect of glycine and NaCl concentrations on the ratio of the mean ionic activity coefficients of NaCl in the presence of glycine and in the absence of glycine, at the same NaCl molality.

Figure 3. Effect of glycine and NaCl concentrations on the ratio of the activity coefficients of glycine in the presence of NaCl and in the absence of NaCl, at the same glycine molality.

ficients of NaCl in the presence of glycine and in the absence of glycine, at the same NaCl molality, as a function of glycine molality. The solid lines are those obtained by curve fitting the experimental data using eq 11. From Figure 2, it can be seen that the presence of glycine affects the mean ionic activity coefficient of the NaCl. As can be seen in Figure 2, for a fixed NaCl molality, the mean ionic activity coefficient of NaCl decreases as the molality of glycine in the solution increases. The decrease in the mean ionic activity coefficient of NaCl with an increase in glycine molality is smaller at higher NaCl molalities. This suggests that higher electrolyte concentrations screen more electrostatic ion-dipole interactions and also increase the importance of the short-range interactions, resulting in a smaller effect of the amino acid on the mean ionic activity coefficient of NaCl. Higher electrolyte concentrations also favor the formation of ion-dipole pairs which reduce the dipole-dipole and ion-dipole interactions. Figure 3 shows, using the molality of glycine as a parameter, the ratio of the activity coefficients of glycine

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2740 Ind. Eng. Chem. Res., Vol. 35, No. 8, 1996

Figure 4. Effect of DL-alanine and NaCl concentrations on the ratio of the mean ionic activity coefficients of NaCl in the presence of DL-alanine and in the absence of DL-alanine, at the same NaCl molality.

in the presence of NaCl and in the absence of NaCl, at the same glycine molality, as a function of NaCl molality. As can be seen in Figure 3, the presence of NaCl, for a fixed glycine molality, decreases the activity coefficient of glycine. This in turn suggests a salt-in effect for the glycine in the presence of NaCl. This decrease in the activity coefficient of the glycine with an increase in NaCl concentration is smaller at higher molalities of glycine and NaCl. A reason for this behavior, as mentioned before, can be that higher electrolyte and amino acid concentrations lead to the formation of ion-dipole pairs. This, in turn, weakens the forces between the charged ions and charged amine and carboxyl groups of amino acid and makes the effect of forces between the organic part of the amino acid and ions more pronounced. Figure 4 shows, using the molality of NaCl as a parameter, the ratio of the mean ionic activity coefficients of NaCl in the presence of DL-alanine and in the absence of DL-alanine, at the same NaCl molality, as a function of DL-alanine molality. From Figure 4, it is evident that the presence of DL-alanine, compared to that of glycine, has a smaller effect on the mean ionic activity coefficient of NaCl. This is reflected in the small absolute value of C1, the coefficient of the leading term in eq 11, which represents pairwise interactions between NaCl (considered undissociated and as one species) and the amino acid. The absolute value of C1 in the case of DL-alanine is smaller than that of glycine, and this indicates bigger pairwise interactions between NaCl and glycine than between NaCl and DL-alanine. A reason for this can be that DL-alanine has a larger hydrocarbon backbone than glycine. This results in a compensation of the attractive ion-dipole forces with repulsive ion-nonpolar hydrocarbon chain forces. As is shown in Figure 4, similar to the effect of glycine on the mean ionic activity coefficient of the NaCl, for a fixed NaCl molality, the mean ionic activity coefficient of NaCl decreases as the molality of DL-alanine in the solution increases. The decrease in the mean ionic activity coefficient of NaCl with an increase in the DLalanine concentration is smaller at higher NaCl molalities. In general, the system H2O + NaCl + DL-alanine exhibits a behavior close to that of an ideal solution.

Figure 5. Effect of DL-alanine and NaCl concentrations on the ratio of the activity coefficients of DL-alanine in the presence of NaCl and in the absence of NaCl, at the same DL-alanine molality.

Figure 6. Comparison of the effect of DL-R-alanine and DL-βalanine concentrations at two different NaCl concentrations on the ratio of the mean ionic activity coefficients of NaCl in the presence of an amino acid and in the absence of an amino acid, at the same NaCl molality.

This kind of ideal behavior can also be seen in the H2O + DL-alanine system, in which the activity coefficient of DL-alanine, as reported by Fasman (1976), varies from 1.004 to 1.044 in the molality range of DL-alanine between 0.2 and 1.86 m. Figure 5 shows, using the molality of DL-alanine as a parameter, the ratio of the activity coefficients of DLalanine in the presence of NaCl and in the absence of NaCl, at the same DL-alanine molality, as a function of NaCl molality as obtained from eq 13. As is shown in Figure 5, the presence of NaCl, for a fixed DL-alanine molality, decreases the activity coefficient of DL-alanine. This behavior is similar to that of the system H2O + NaCl + glycine but with smaller intensity. This suggests a salt-in effect for DL-alanine in the presence of NaCl which can again be a result of the bigger hydrocarbon chain in DL-alanine than in glycine. Figure 6 compares, at two different NaCl concentrations, the ratio of the mean ionic activity coefficients of NaCl in the presence of two isomeric forms of DL-alanine to that, at the same NaCl molality, without the presence

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Ind. Eng. Chem. Res., Vol. 35, No. 8, 1996 2741

trace activity coefficient, all amino acid molality terms in eq 13 were set equal to zero. The comparison of the trace activity coefficients of the glycine and DL-alanine in the solutions with different NaCl concentrations shows that the trace activity coefficients of DL-alanine have generally closer to unity values than those of glycine over the whole NaCl concentration range. This is also in agreement with the comparison reported by Schrier and Robinson (1974) for the systems H2ONaCl-glycine and H2O-NaCl-L-R-alanine. Conclusions

Figure 7. Effect of NaCl concentration on the trace activity coefficients of DL-alanine and glycine.

of the amino acid, as a function of the amino acid molality. The amino acids compared are DL-R-alanine and DL-β-alanine. These amino acids are isomers and differ in the position of the amino groups on their hydrocarbon backbones. Experimental data for the system H2O + NaCl + DL-β-alanine are reported by Schrier and Robinson (1971). As is shown in Figure 6, β-alanine shows a greater interaction with NaCl than R-alanine. This difference remains over the whole range of amino acid and NaCl concentrations. A reason for this can be that the charges on β-alanine, arising from its charged amine and carboxyl groups, are farther apart from each other than in the case of R-alanine. This, in turn, results in a higher dipole moment and thus in larger interactions with both water molecules and charged ions. It has been shown that aqueous solutions of D-, L-, and DL-alanine at the same concentration of amino acid have identical vapor pressures, within the accuracy of the isopiestic method (Robinson et al., 1942). This indicates that the different optical isomers of the alanine have identical activity coefficients. We have also compared the results of the measurements of the mean ionic activity coefficients of the NaCl for the system H2O + NaCl + DL-alanine obtained in this study with those reported by Schrier and Robinson (1974) for the system H2O + NaCl + L-alanine. This comparison showed an absolute average relative difference of 0.8% and an absolute maximum difference of 2.1% between the mean ionic activity coefficient of NaCl in both systems. The small difference in the mean ionic activity coefficients in the systems with L-alanine and DL-alanine confirms that the activity coefficients of the optical isomers of alanine are identical within the precision of the measurements. However, although the optical isomers have identical physicochemical properties such as boiling point, melting point, etc., their solubilities have been reported to be different (Jin and Chao, 1992). In the absence of a solid phase, for all engineering purposes, the equality of the activity coefficients of optical isomers seems to be a good assumption. Figure 7 shows the trace activity coefficients of glycine and DL-alanine, γtr, as a function of NaCl molality. The trace activity coefficient of an amino acid in the aqueous solution of NaCl is the value of its activity coefficient when its molality approaches zero. To calculate the

The theoretical basis of the method for the measurement of the mean ionic activity coefficients of electrolytes in the presence of nonelectrolytes using electrochemical cells with two ion-selective electrodes has been discussed. From this analysis, it was concluded that, in spite of previous measurements made in cells using only a cation ion-selective electrode, precise measurements require cells consisting of a cation and an anion ion-selective electrode, each measured vs a reference electrode. Activity coefficient data for the systems H2O + NaCl + DL-alanine and H2O + NaCl + glycine were obtained using an electrochemical cell. The cell consisted of a sodium and a chloride ion-selective electrode, each measured vs a double-junction reference electrode. The results of the measurements of the mean ionic activity coefficients of NaCl in the H2O + NaCl + glycine system were compared with those obtained by the isopiestic method and with cells using only a sodium ion-selective electrode against a reference electrode. The results obtained in cells with two ion-selective electrodes are in better agreement with isopiestic data than those obtained with one single ion-selective electrode. The results of the measurement of the mean ionic activity coefficients of NaCl in the H2O + NaCl + DL-R-alanine system were compared with those obtained from the literature for the H2O + NaCl + DL-β-alanine and H2O + NaCl + L-R-alanine systems. The results showed that, although DL-R-alanine and DL-β-alanine are isomers, the mean ionic activity coefficients of NaCl have a larger change in the system containing DL-R-alanine than in the system containing DL-β-alanine. This difference was attributed to the difference in the position of the charged amine and carboxyl groups on the amino acid. The mean ionic activity coefficients of NaCl in the systems of the DL and L optical isomers of alanine were found to be equal within experimental error. Acknowledgment We are grateful to the Natural Sciences and Engineering Research Council of Canada for financial support. Nomenclature amrd ) absolute mean relative deviation Ci ) adjustable parameters E+ ) potential of the cation ion-selective electrode E+ISE ) internal potential of the cation ion-selective electrode E+° ) standard-state potential of the cation ion-selective electrode E- ) potential of the anion ion-selective electrode E-° ) standard-state potential of the anion ion-selective electrode

+

+

2742 Ind. Eng. Chem. Res., Vol. 35, No. 8, 1996 E-ISE ) internal potential of the anion ion-selective electrode EJ ) junction potential ERef ) reference electrode potential ∆E(1) ) potential difference in the electrochemical cell with electrolyte but without the presence of solutes ∆E(2) ) potential difference in the electrochemical cell with both electrolyte and other solutes F ) Faraday number ISE ) ion-selective electrode mA ) molality of amino acid mS ) molality of electrolyte m+ ) molality of cation m- ) molality of anion n ) number of data points rmsd ) root-mean-square deviation S ) slope of electrode potential T ) absolute temperature P ) pressure R ) universal gas constant z+ ) charge number of the cation z- ) charge number of the anion Greek Letters γA ) activity coefficient of the amino acid γ+ ) activity coefficient of the cation γ- ) activity coefficient of the anion γ( ) mean ionic activity coefficient of the electrolyte γ(1)( ) mean ionic activity coefficient of the electrolyte in the electrochemical cell without the presence of other solutes γ(2)( ) mean ionic activity coefficient of electrolyte in the electrochemical cell with other solutes ν ) stoichiometric number of the electrolyte, ν+ + νν+ ) stoichiometric number of cations ν- ) stoichiometric number of anions Subscripts ISO ) isopiestic ISE ) ion-selective electrode

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Received for review September 21, 1995 Accepted April 16, 1996X IE950581E

X Abstract published in Advance ACS Abstracts, June 15, 1996.