THE SOLUBILITIES OF CERTAIN AMINO ACIDS IN WATER, THE DENSITIES OF THEIR SOLUTIONS AT TWENTY- FIVE DEGREES, AND THE CALCULATED HEATS OF SOLUTION AND PARTIAL MOLAL VOLUMES* BY JOHN B. DALTON AND CARL L. A. SCHMIDT (From the Division of Biochemistry, University of California Medical School, Berkeley) (Received for publication, June 13, 1933) It is a striking fact that, although the properties of the amino acids have been studied from many standpoints, viz. mkitional, optical, and physicochemical, data on two of the most fundamental properties of their aqueous solutions are fragmentary. The more important of these properties is that of solubility in water and the effect of temperature thereon, and the other is that of the densities of their solutions. The data which have been obtained have in- cidentally permitted the calculation of such quantities as heat of solution and molecular volume. These relationships have also furnished some clue as to the state of the racemic material. It is unfortunate that two terminologies have been used in expressing concentration values; namely, one in gm. or in moles per liter of solution, and the other in gm. or moles per 1000 gm. of solvent. One objection to the first method of expression is that the concentration is not the same at all temperatures. The volume of the solution may change with the temperature, but the amount of solute remains constant. Another objection is that on dilution the concentration change is not necessarily proportional to the volume change due to changes in the partial molal volumes of the constit,uents. In order to obviate these difhculties, the measurements reported here have been expressed in moles or gm. per 1000 gm. of water. Density measurements at 25’ which were carried out on solutions of the amino acids will facilitate the con- version of one method of expressing concentration into the other. * Aided by grants from The Chemical Foundation, Inc., and the Re- search Board of the University of California. 549 by guest on May 28, 2018 http://www.jbc.org/ Downloaded from
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THE SOLUBILITIES OF CERTAIN AMINO ACIDS IN WATER, THE DENSITIES OF THEIR SOLUTIONS AT TWENTY- FIVE DEGREES, AND THE CALCULATED HEATS OF SOLUTION AND PARTIAL MOLAL VOLUMES*
BY JOHN B. DALTON AND CARL L. A. SCHMIDT
(From the Division of Biochemistry, University of California Medical School, Berkeley)
(Received for publication, June 13, 1933)
It is a striking fact that, although the properties of the amino acids have been studied from many standpoints, viz. mkitional, optical, and physicochemical, data on two of the most fundamental properties of their aqueous solutions are fragmentary. The more important of these properties is that of solubility in water and the effect of temperature thereon, and the other is that of the densities of their solutions. The data which have been obtained have in- cidentally permitted the calculation of such quantities as heat of solution and molecular volume. These relationships have also furnished some clue as to the state of the racemic material.
It is unfortunate that two terminologies have been used in expressing concentration values; namely, one in gm. or in moles per liter of solution, and the other in gm. or moles per 1000 gm. of solvent. One objection to the first method of expression is that the concentration is not the same at all temperatures. The volume of the solution may change with the temperature, but the amount of solute remains constant. Another objection is that on dilution the concentration change is not necessarily proportional to the volume change due to changes in the partial molal volumes of the constit,uents. In order to obviate these difhculties, the measurements reported here have been expressed in moles or gm. per 1000 gm. of water. Density measurements at 25’ which were carried out on solutions of the amino acids will facilitate the con- version of one method of expressing concentration into the other.
* Aided by grants from The Chemical Foundation, Inc., and the Re- search Board of the University of California.
In general, the solubilities of the dl forms of the amino acids reported in the literature are reasonably concordant and checked roughly with the data obtained in this investigation. This is probably due to the fact that the racemic amino acids can be prc- pared synthetically quite free from contaminants.
The solubility data of the optically active materials are, how- ever, very discordant. The difficulty probably arose from several sources. First, the compounds may have been slightly contami- nated with neutral salts. The work of Pfeiffer and Wiirgler (l), Pfeiffer and Angern (2), von Euler and Rudberg (3), and ot,hers (4) has shown that the solubilities of amino acids are affected by the presence of neutral salts.
Another difficulty is probably due to contamination with other amino acids during the process of isolation. Thus Bayliss (5) has shown that the solubility of I-aspartic acid in a 1.25 per cent solution of I-leucine is 30 times that in pure water. The effect of one amino acid on the solubility of the other is a tiommon phenome- non observed during isolation procedures when mixtures of quite insoluble amino acids are concentrat.ed to a thick syrup before any evidence of crystallization is noted. This factor is also evident in the experimental work dealing with the separation of the isomeric leucines. In fact it has been so difllcult to isolate d-norleucine that its presence in protein hydrolysated has only recently been established (6-10). On the other hand, von Euler and Rudberg (3) found that in a solution saturated with respect to bot.h dl- leucine and Z-tyrosine the solubilities of both amino acids were less than when each was dissolved alone in pure water.
The third difsculty is that of partial raeemization which proba- bly occurs during the hydrolysis of the protein from which the amino acids are obtained. As Pellini and Coppola (11) have shown, the solubility of a mixture of d-alanine and dl-alanine is greater than either one of the two separately. Bayliss (5) ad- mitted his d-glutamic acid was partially racemized. His solubility value is greater than that interpolated from the present data. In some instances the difhculty may be attributed to the inadequacy of the analytical methods employed for the estimation of the amino acids.
The determinations herein reported were undertaken with a view of obtaining accurate solubility and densit,y data on carefully purified amino acids under well controlled experimental conditions.
The experimemal technique used t,o carry out solubility and density determinations is so well known that a detailed descript,ion here is unnecessary. It suffices to state that the weights, volumet- ric glassware, and thermometers were carefully calibrated against known standards. The solubilities of the amino acids were determined by permitting a supersaturated solution and an under- saturated solution in the presence of an excess of the particular amino acid to come to equilibrium by shaking the mixture in an inverted Pyrex glass T-tube having closed ends and of about 100 cc. capacity, at the desired temperature. Filtration was accom- plished by the use of filters so devised that air under a pressure of 150 mm. of mercury was applied to the solution so that the liquid was forced through a filter of cotton and asbestos and into a weighing bottle. In this way evaporation losses were kept at a minimum, and as the filtration was rapid, there was little or no crystallization before the solution reached the weighing bottle. In the case of l-aspartic acid, d-glutamic acid, and l-tyrosine, the solubilities were determined by estimating the nitrogen content of the solution by the micro-Kieldahl procedure of Parnas and Wagner (12). For higher solubilities the estimation was based on the weight of the residue when dried at 95-97”. In the case of dLglutamic acid the residue was dried at 60” to prevent the forma- tion of pyrrolidone-a-carboxylic acid. The solubility of diiodo- I-tyrosine was determined by estimating the amount of iodine present in the solution according to the procedure of Kelly and Husband (13).
The purity of the materials used is all-important insolubility work. The amino acids were decolorized by treating the solutions with purified decolorizing carbon. They were twice recrystallized from hot distilled water and once from conductivity water. The products were then dried in vucuo at 37“ and finally at 95”.
In general, several criteria of purity were required of the amino acid under investigation. First, it must give a negative reaction for the ammonium ion-a common contaminant-when tested with Nessler’s reagent. Second, it,s nitrogen content must agree with the theoretical value within the limits of the errors of the determination. Third, in the case of the optically active amino acids, its degree of rotation must agree reasonably well with the values quoted in the literature. This is very important and for
that reason, under each optically active ammo acid, the rotation found, together with that given by other workers, is given. Fourth, particularly in the case of I-leucine, which can always be suspected of being comaminated with amino acids of the same molecular weight, it must show the same solubility when a small excess of the solid phase is added as when a large excess is present. Also, after further recrystallization, the solubility of the amino acid must be the same as before.
The thermodynamic nomenclature used in this paper, unless otherwise stated, is the same as given by Lewis and Randall1 (14). For the sake of space economy, the complete solubility data relat- ing to the fifteen amino acids studied have not been presented, but rat.her equations in which the logarithms of the solubilities of the amino acids are expressed as functions of the temperature. These equations were developed by using the method of least squares as outlined by Lipka ((15) p. 127). When the relationship was not linear, the procedure indicated by t,he same aut,hor ((15) p. 151) was followed with the exception that the method of least squares instead of the method of averages was employed to evaluate the point.s. It is probable that the calculated solubility values are more accurate than any single determined value since the former represent a mean of all of the values over the entire range of temperatures studied.
In order to make the solubility data readily available, a table of solubilities has been included (see Table I) in which the solubilities of the amino acids have been calculated at 5” intervals. With the aid of Table I intermediate values may be readily obt,ained by interpolation. All of the solubility values are given in terms of gm. per 1000 gm. of solvent. In order that comparisons could be made, it was necessary in certain instances to recalculate the data given by others. This was done by the use of the density data obtained at 25”. For other temperatures, t,he corresponding density was calculated on the assumption that the density of the
* The following is the nomenclature used throughout this paper. no, molality in moles per 1000 gm. of water; nl, number of moles of solvent; 7~~: number of moles of solute; Ne, mole fraction of solute; 8, solubility in gm. per 1000 gin. of water; PI, molsl volume of solvent; 51, partial molal volume of solvent; 52, partial molal volume of solute; V, total volume; (p, apparent molal volume; y, activity coefficient; R, gas constant.
amino acid solution changed with temperature in the same way as water. In order to economize on space, these data have not been included. They are on file at the University of California library.
For the calculation of heats of solution2 of the amino acids, use was made of the well known van’t Hoff equation as motied by Schriider (21). According to the present nomenclature his equa- tion is
b In N&T = AH/RT2 (1)
This equation, however, applies only to perfect solutions. Since not all of the solutions of the amino acids are perfect solutions as shown by the freezing point measurements of Fr%nkel (22), Lewis (23)) Hoskins, Randall, and Schmidt (24), and Cann (25), it was necessary to apply the relationships as formulated by Brijnsted (26). His equations may be written
AH = RT2 b In m dlnr dm -+-.-
bT dm dT
for un-ionized solutes, and
b In m b In y dm bT+dm-
dT
(2)
(3)
for ionized solutes. In order to calculate the heat of solution of any particular amino acid, it is necessary to use the activity coeffi- cient. Consideration of this as well as the calculation of the heat of solution is given under each amino acid.
The density data of the amino acid solutions are given in Tables IV to VI. These data have enabled us to compute the partial molal volumes of the amino acids. For this purpose we have followed essentially the procedure given by Lewis and Randall ((14) p, 37) in which a term, (o, is defined as the apparent volume occu- pied by 1 mole of the solute. This may be expressed its Q = (V - Wl)/n2. If the weight of the water is 1000 gm., itlvl becomes 1002.94 cc. at 25” and nz becomes the molaJity, m. Values for (p are obtained by dividing the sum of the weight of the
* For the development of the ideas which relate the heat of solution to the solubility see also Le Chatelier (16), Goldschmidt (17), Hildebrand (18), Lewis and Randall ((14) p. 229), Butler and Hiscocks (19), and Linder- str#m-Lang (20).
solvent, 1096 gm., and the weight of the solute that it contains by t,he density. If q is plotted against log 17a, the slope of the curve may be designated as s. This value s is divided by 2.303 and to the quotient is added the corresponding value of cp at t,hat point. This sum can be demonstrated to be equal to the partial molal volume of the solute, &. Now it is known ((14) p. 42) that, T’ = n& + w&. Then since ~2, m, and n1 are also known, and since V can be calculated from t#he weight of a solution and its density, then uI, the partial molal volume of the solvent, can be calculated.
The values for t,he partial molal volumes of the amino acids so obtained can be compared with t,he apparent. molal volumes calcu- lated from the empirical atomic volumes given by Traube (27). The “molekulares Losungsvolumen” computed from his figures refer to the apparent molal volume, p, rather than to the t,rue partial molal volume, VZ. Since in the present, case the two are only slightly different, a rough comparison can be made between the calculated and observed values. Traube’s data were com- put,ed for 15”. Since his values have an insignificant temperature coefficient, the assumption has been made t,hat the values calcu- lated for 15” hold also for 25”. The following volumes occupied by the amino acids studied were calculated: alanine, 58.8 cc.; aspartic acid, 74.7 cc. ; glycine, 42.7 cc. ; glutamic acid, 90.8 cc. ; leucine and its isomers, 107.1 cc. ; valine, 91.1 cc. ; phenylalanine, 68.9 and 130.6 cc. The value of 139.6 cc., which was computed for phenyl- alanine, was obt,ained by assuming that t.he carbon and hydrogen at,oms maintained t,heir usual volumes. Traube (27) noted the fact t,hat the calculated molar volumes of certain organic com- pounds in aqueous solutions are in general at least 12 cc. per mole less than the observed values. Cohn and his coworkers (28) have found that the observed molar volumes of glycine, alanine, valine, and leucine agree well with those which they calculated from Traube’s empirical volumes. Our data harmonize with those pub- lished by Cohn and his coworkers. The t,ighter packing of the amino acids is believed by them to depend upon electrost,rict,ion of solvent molecules due to the zwitter ions.
Discussion of Volubility and Density Bata
d-Alanine-The only data dealing with the solubility of this amino acid were obtained by Pellini and Coppola (11). The
specific rotatory power of the hydrochloride dissolved in water was .reported by them as [a]:’ = +10.3”. This agrees well with [or]:’ = + 10.2” found for the hydrochloride of the present material, and with (cy]:’ = + 10.1” reported by Huffman and Borsook (29). The present series of solubility determinations were carried out at nine different temperatures between O-65’. A total of eight,een determinations was made.
The following equations, which may be used for both practical as well as theoretical purposes, express the solubility relationships for d-alanine as a function of the temperat,ure.3 Coefficients for equations of a similar type expressing the corresponding relation- ships, together with the maximum deviation from the equat.ions, for this and for certain other amino acids are summarized in Table II. The equations for alanine are:
log S = 2.1048 + 0.004669 t log m = 0.1551 + 0.004669 t In m = -2.5792 + 0.01075 T In N2 = -6.5150 + 0.01037 T
The present. data for d-alanine fit these expressions very closely, the maximum deviation of t.he observed figures from those calcu- lated from the equations being of the order of ho.5 per cent. The values found by Pellini and Coppola (11) for d-alanine are between 0.6 and 2.0 per cent higher than those reported in this paper.
According to Friinkel (22), the freezing point lowering, and hence the activity coefficients, for d-alanine were found to be normal up to concentrat,ions of 1.512 M. It has been assumed in the present calculations that t,his relationship holds at higher con- centrations. The differential heat of solution of d-alanine in a saturated solution was then computed directly from Equation 1. From the appr0priat.e solubility equation given above is obtained the equation b In N2/bT = 0.01037, and at 25’, AHZss = 1830 calories per mole.
3 The data may be presented in graphical form by plotting In N, against either l’or l/T. We prefer the former for the sake of simplicity since seven amino acids (dl-alanine, d-alanine, I-aspartic acid, dl-glutamic and d- glutamic acid, I-tyrosine, and diiodo-dl-tyrosine) when plotted in this way give straight lines rather than curves. The method of plotting -R In N2 against l/T is theoretically preferable since the slope of the line is directiy equal to the differential heat of solution. For the sake of space economy, the curves have not been included in this paper.
According to the theory underlying the fundamental relation- ship; the melting point of d-alanine should be t,hat temperature at which Ns = 1. This temperature computed from the appropriate equation above proved to be approximat.ely 355’. As the decom- position point recorded by Fischer (30) for this compound was
TABLE 111
Density of Aqueous So1ution.r o/ CertaifL Amino Acids at 25”
* In Tables III, V, and VI the concentration is given as gm. per 1000 gm. of water. rn, is given as moles per 1000 gm. of water. d is the absolute density.
297”, it is considered probable that decomposition occurred before the melting point was reached.
The densities of a series of aqueous solutions of d-alanine and the calculated partial molal volumes are given in Tables III and IV. The values agree roughly with 58.8 cc. calculated according to Traube (27).
dl-Alan&e.--Eighteen solubility measurements on dl-alanine were carried out at nine different temperatures between O-60”. Considerable difhculty was encountered in measuring the solubility at lower temperatures due to the tendency of the amino acid to supersaturate. At 0” samples taken up to 85 hours of shaking showed a successive drop in solubility. The value finally obt.ained was neglected in the calculation of the solubility equations. With
TABLE IV
Partial Molal Volumes of Solvent, ii,, and of Solute, 02, for Solutions of d-Alanine, dl-Alanine, and Glycine at 2.5”
the exception of the measurement at O”, the maximum deviation of the observed values from those calculated from the equation was found to be f1.2 per cent. Previous measurements by Pellini and Coppola (11) and by Holleman and Antusch (31) agree to within 1.5 per cent of the figures interpolated from the present data. Cohn (4) and Pfeiffer and Angern (2) have report.ed their solubility values in terms of moles per liter. By assuming an appropriate density for these solutions comparative figures were
obt,ained. They are from 3.2 to 4.6 per cent lower than those found in the present investigation.
dl-Alanine offers an excellent experimental verification of the relationship between the solubility and the differential heat of solution. Frankel (22) found that the freezing point lowering of this compound was normal up t,o a concentration of 1.680 M.
Therefore Equation 1 was applied directly. From the equation, In NZ = -7.1317 + 0.01245 T, and with
the aid of Equation 1 AH = 0.01245 RT2 is obtained. Baur (32) reported a value of 2020 calories for the heat of solu-
tion of dl-alanine. The init,ial and final concentrations of the solute were not given and the temperature was merely described as room temperature. On the assumption that his temperature was 15”, on the basis of the present, solubility measuremen@ AHzss = 2050 calories. Unpublished calorimetric measurements by Mr. Zittle of this laboratory at 25” indicate that for the heat of solution of 1 mole of dl-alanine in its saturated solution AHXM = 2140 f 30 calories. On the basis of the present solubility meas- urements, AHZSS = 2200 calories. The two values are in quit,e good agreement.
A point of interest was noted relative to the melting point of dl- alanine. According to Schroder’s (21) theory, the temperature at which Nz = 1 is the melting point of the compound. From the appropriate equation this temperature was calculated to be 299” which is in close agreement with 297” reported by Dunn and Brophy (33) and 295” found by Fischer and Leuchs (34) for the decompo- sition point.
The only density value for an aqueous solution of &alar&e at 25” reported in the literature is that of Holleman and Antusch (31). For a solution containing 1.849 moles of solute per 1000 gm. of water they report.ed a. value of 1.0421. By use of linear interpo- lation between the appropriate concentrations in the present data (see Table III), a value of 1.0422 was obtained for a similar solution.
The data presented in Table IV give the partial molal volumes of the solvent and of the solute. It will be noted that there is an approximate agreement between these figures and the value of 58.8 cc. computed from Traube’s atomic volumes.
A comparison of the data for dl- and d-alanine shows that the
solubility of the d form is greater than that of the dl form at lower temperatures. The slope of the solubility curve and hence the heat of solution is much less for d- than for dl-alanine. The partial molal volumes of the two solutes are also different, although the difference is only slight.
According to the rules which have been set down by Meyerhoffer (35), which are based on a comparison of th;c: properties of various optically active compounds and their inact,ive isomers, on the basis of the present data dl-alanine is to be regarded as a racemic compound. This conclusion was also reached by Pellini and Coppola (11).
Glyczke-Despite the fact that glycine is perhaps the most easily obtained amino acid, the solubility data contained in the literature are of the most fragmentary type. The value reported by Lewis (23) agrees very well with the present data. Dehn (36) reported an irrational value of twice the amount found in the present in- vestigation. Values reported by Pfeiffer and Wirgler (l), Pfeiffer and Angern (2), Cohn (4), Sano (37), and Ando (38) are from 2.4 to 9.0 per cent lower than those reported in the present paper.4
The equations which express the solubility relationships of glycine (see Table II) are based on a series of thirty solubility measurements at ten different temperatures over the range of O- 60”.
There are three sources of cryoscopic data for the calculation of the activity coefficients of glycine. Those given by Lewis (23) are more self-consistent. than those published by l?riinkel (22) and by Cann (25). As these data indicate diverging values for the a&ivity coefficients at higher concentrations, each set of data was examined separately. The activity coefficients were calculated by the method indicated by Lewis and Ftandall ((14) p. 286). Cann (25) calculated the activity coefficients for glycine assuming that two ions were formed at Unite dilution. According to the zwitter ion theory this is correct. Nevertheless, one particle only is present which carries the two charges; and the osmotic effect therefrom is not that which would be expected from two ions but
4 New solubility values found by E. J. Cohn and T. L. McMeekin check quite well with the values reported in this paper. Their values at 25” are: glycine, 249.9; dl-valine, 70.7; dl-leucine, 9.86; dl-norleucine, 11.5; and l-leucine, 24.7 gm. per 1000 gm. of water. Private communication.
only one. Cann’s (25) figures were therefore recalculated to conform with the convention that at infinite dilution glycine depresses t.he freezing point at the rate of 1.858 degrees per mole of solute. As there are no adequate thermal data regarding the relative heat contents, z1 and &, for glycine solutions, no correc- tion could be made for the heat of dilution ((14) p. 288). Data published by Naude (39) concerning the heats of dilution for various glycine solutions indicate the probability of a significant correction. It is doubtful, however, from the character of the measurements cited that the data warrant this refinement.
Empirical equations were devised to express the logarithm of the activity coefficient as a function of the concentration.6 It was found that
In 7 = -0.1814 In m - 0.2538 In y = -0.1616 In m. - 0.2165 In y = -0.0607 In m - 0.0396
expressed Frankel’s (22), Lewis’ (23), and Cann’s (25) data respectively. ,
By modifying Equation 2 t,o
the above equations, together with the solubility equation, can be differentiated and substitut,ed directly. Thus, by use of Frankel’s data, the equation for In na (Table II), and Equation 4 we obtain AH = RT2 (7.676 X 10e2 - 1.895 X 1O-4 T) (1 - 0.1814), and at. 25”, AHzes = 2930 calories per mole. Similarly from Lewis’ (23) data, AH29s = 3010 calories per mole and from Cann’s (25) data ~~~~~ = 3370 calories per mole. As the above relations rigidly hold only in the neighborhood of the freezing point, these values may be somewhat in error.6 Since at higher temperatures the solu-
6 In this and other equations which express the activity coefficient as a function of concentration, the equations were calculated and hold only for concentration ranges near saturation.
6 The assumptions which have been made in calculating values for AH in the case of glycine, c&alanine, and dl-valine appear to be justifiable inas- much as the values calculated for AH from solubility measurements check quite well with those which were determined by calorimetric measurements.
tion tends to become more nearly perfect, at T = 298”, the true differential heat of solution should lie between the values indicated above and AH298 = 3590 calories, which is obtained by assuming that the perfect solution laws are obeyed. Calorimetric measure- ments which have been carried out in this laboratory by Mr. Zittle indicate t.hat AHw = 3530 calories when. 1 mole of glycine is dis- solved in an infinite amount of a solution of glycine saturated at 25”.
Louguinine (40) measured the integral heat of solution of glycine at 16”, the final concentration being 0.266 M. The value, AHsss = 3580 f 40 calories, may be regarded as identical with the differ- ential heat of solution as is customary at those concentrations. This value also lies between the limits calculated on the basis of
Frlinkel’s (22) data (2990), Lewis’ (23) data (3060), and the data of Cann (25) (3430) and the theoretical value of AHzss = 3660 calories.
The densities of a number of aqueous solutions of glycine were determined at 25’ (see Table V). From these the partial molal volumes of the solvent and of the solute were calculated in the usual manner. These are given in Table IV. It will be noted that the values for the partial molal volumes are in approximate agreement with the value 42.7 cc. calculated on the basis of Traube’s atomic volumes.
dEVuZ&--Solubility values for dZ-valine have been reported in the literature by Cohn (4), Slimmer (41), and Hoppe-Seyler (42).
They are, however, very discordant, ranging from 28 per cent lower to 30 per cent higher than the values interpolated from the present data.
52 solubility measurements were carried out at fifteen different temperatures between O-80”. The data were summarized by means of equations, the c.oefficients of which are found in Table II.
The cryoscopic data for dZ-valine are limited solely to the meas- urements given by Fr&nkel (22). On the basis of t,hese data, the activity coefficients were calculated in the usual manner ((14) p. 286). The equation, In 7 = 0.0549 In m + 0.1669, was devised to express these values. From this, from the solubility equation for In m (see Table II), and from Equation 4 we obtain AR = RT2 (-2.729 x 102 + 1.201 x 10-4 T) 1.055. At 25“, AHzes = 1590.
The true value for the differential heat of solution probably lies between this value and that obtained by assuming that the perfect solution laws were obeyed. This would indicate that the value of AH298 lies between 1590 and 1500 calories. Calorimetric measure- ments by Mr. Zittle in this laboratory indicate that AH298 = 1590 f 30 calories for the differential heat of solution of dZ-valine in a saturated solution.
From densities of a number of aqueous solutions of ukaline at 25” (see Table III) the partial molal volumes, g1 and G, were calculated. In the preliminary plot of cp against log m it was found that p was a constant. The term s in the equation, c2 = ((p + s/2.303), becomes equal to zero and c2 = 8. Since (a = ~~ = (V - Wl)/nz, (V - n1vl)/n2 may be substituted for 3 in the fundamental relationship, V = nlii, + nzfi2, which gives nlvl = n&k The partial molal volume of the solvent & then becomes equal to ~1, the molal volume of the pure solvent. Thus for the concentrations from 0 to 0.6814 M, 51 = 18.069 cc. and ij2 = 91.06 cc. The value for & agrees well with 91.1 cc. calculated from Traube’s atomic volumes.
dLIsoleucine-Apparently no accurate solubility data for this amino acid have appeared in the literature.
Thirty-two solubility measurements were carried out on pure &isoleucine at nine different temperatures between O-65”. By the usual methods equations were derived to express the relation- ship between the solubility and the temperature. The coefficients of these equations are summarized in Table II.
Since there are no data from which the activity coefficients of dl- isoleucine can be calculated, the estimation of the differential heat of solution becomes only approximate.
Using the equation for In NZ (see Table II), we obtain AH = RT2 (-4.134 X*1O-2 + 1.726 X lo4 T). At 25’, AH29, = 1790 calories per mole.
Since the shape of the curve of dl-isoleucine in which In NZ was plotted against T resembles that of dl-valine, dl-leucine, and L leucine which were known to possess activity coefficients which increased with the concentration, it was assumed iir Equation 3 that b In y/&n is positive. As dm/dT is also positive, 1790 calo-
TABLE VI
lhnsity of Aqueous Solutions of Several Amino Acids at .W
ries is regarded as a minimum value for the differential heat of solution.
The density determinations were carried out on two solutions of this material, one at approximately saturation, and one at half saturation (see Table TV.). From these values the apparent molal volumes at 25’ were computed as follows: when m = 0.1647, Q = 106.2 f 0.2 cc.; when m = 0.0798, Q = 166.4 f 0.2 cc. Follow- ing the reasoning described from dZ-valine, (p = & = 106.3 cc. and til = v1 = 18.069 cc. The value of 106.3 cc. compared roughly wit,h 107.1 cc. computed from Traube’s empirical atomic volumes.
I-Leueine-The inconsistencies in the data reported in the
literature concerning the solubility of Z-leucine in water are, in all probability, due to contamination with various amounts of the other leucine isomers. The former workers, with one exception, report only isolated determinations. These include Pfeiffer and Angern (2), Cohn (4), Bayliss (5), Schulze (43), Schulze and Bosshard (44), Schulze and Likiernik (45), Sano (46), and Nencki (47). The most extensive work on Lleucine has been reported by Takahashi and Yaginuma (48) whose values agree well with the present data.
The material used for the present investigation had a rotatory power of [oL]: = +15.7” for a 7 per cent solution of I-leucine in 20 per cent hydrochloric acid. This value agrees well with [a]:’ = +15.7” reported by Ehrlich (49) and with [&’ = i-15.6” found by Fischer and Warburg (50). A series of twenty solubility measurements at nine different temperatures between O-65” were carried out upon this compound. The usual equations relating to the solubility and temperature were devised and their coefficients tabulated (see Table 11).
Since the freezing point data for the leucines are of rather doubt- ful character, it was realized that only a rough estimate of t.he activity coefficients could be made. As it was not,ed in Frankel’s (22) work that the freezing point depression of I- and of dl-leucine was about the same at the same concentraticn, it was assumed that the activity coefficients are the same for both compounds. The empirical equation, ln y = 0.382 m, was devised to express the relation between the activity coefficient of the leucines and the concentration.
From the equation for ln na (see Table II) is derived the equation drn/dT = (-4.683 X lOwa + 1.716 X 10-J T) m. Fror the appropriate equation above b In r/&n = 0.382 is obtained. The equation for the heat of solution then becomes AH = RT2 (1 + 0.382 m) (-4.683 X 10” + 1.716 X lo+ T). At 25”, na = 0.1851, and hence AHags = 830 calories per mole.
The true value probably lies somewhere between this value and that obtained by assuming that the laws of the perfect solution are obeyed. This would indicate that AH298 = 830 - 770 calories per mole at saturation.
From density determinations (see Table VI) which were carried out on two aqueous solutions of Meucine, one at approximately
saturation concentration and the other at half of t,his value, the following apparent molal volumes were calculated: when m = 0.1829, ‘cp = 106‘9 f 0.2 cc., and when m = 0.0885, (a = 107.2 f 0.2 cc. Since the values for (p are identical, from the reasoning described for dl-valine, they were considered to be identical with the true part.ial molal volume, ~2. The mean of these two values, 107.1 cc., is identical with t,he value calculated on the basis of Traube’s atomic volumes.
dl-Leucine-The solubility values reported in the literature concerning dl-leucine, akhough fragmentary, agree reasonably well with the values found in the present investigation. Measure- ments have been reported by Pfeiffer and Wtirgler (l), von Euler and Rudberg (3), Schulze and Bosshard (44), Hiifner (51), and Schulze (52, 53).
Forty solubility measurements at ten different temperatures between O-70” were carried out on dl-leucine. Equations were derived in the usual manner to express the relationship between the solubility and the temperature. The coefficients are summarized in Table II.
In the previous section the assumpt,ion WM made on the baais of freezing point data that the activity coefficients of I-leucine and dl-leucine are t.he same. On this basis, In y = 0.382 m also holds for dl-leucine. Hence b In y/bm = 0.382. In the manner previously described we obt,ain for Equation 2, AH = RT2 (-5.167 X 1O-2 + 2.114 X 10” T) (1.0 + 0.382 m). At 25”, and m = 0.0756, the true value of the differential heat of solution at saturation probably lies between the value indicated by the above equation and that obtained by assuming that the perfect solution laws hold for t,his amino acid. Hence AHzgR = 2070 - 2000 calories per mole.
By comparing the solubility data for Z-leucine and dZ-leucine it was noted t,hat the solubility of t.he dZ form is less and the tempera- ture coefficient is greater t,han the corresponding values for the Z form. The heat of solution of the dZ form is also greater than that of the 1 form. According to the rules set down by Meyerhoffer (35), dhleucine should be regarded as a racemic compound.
The apparent molal volume of dl-leucine was calculated from the density data which were obtained at 25” (see Table VI). At these concentrations Q and 62 may be considered to be identical. The
observed value of 110.0 of 0.2 cc. agrees roughly with 107.1 cc. cal- culated on t,he basis of Traube’s atomic volumes.
It will be noted t.hat the value of Bz = 110 cc. for dbleucine is significantly different from tiz = 107 cc. reported for l-leucine. This fact lends furt.her support to t,he idea that dl-leucine may be regarded as a racemic compound.
l-Aspartic Acid-l-Aspartic acid has been studied to a much greater extent from the st.andpoint of solubility relationships than any of the other amino acids. The values reported, however, are quite variable, agreeing neither with the present dat,a nor are they consistent among themselves. It was noted that the different optical rotations reported were accompanied by different solu- bility values, the solubility becoming less as the rotation value increases. This fact suggests the possibility that some of the observat,ions were made on partially racemized material.
Observations conducted over a number of temperatures were made by Guareschi (54), Engel (55), Cook (56), and Bressler (57), while isolated values were reported by Pfeiffer and Wtirgler (l), Pasteur (58), and Marshall (59).
The optical activity of the Z-aspartic acid used in the present investigation was [al E5 = +25.1” when 1 mole of aspartic acid was dissolved in 3 moles of hydrochloric acid. This agrees well with [(r]:’ = + 25.5O reported by Fischer (60) and [LY] z = 25.0” reported by Clough (61). Thirty-eight observations were carried out at nine evenly spaced temperatures between O-60’. The usual solubility-temperature equations were devised and t,heir coefficients tabulated (see Table II).
The most reliable cryoscopic dat,a for Z-aspartic acid are those of Hoskins, Randall, and Schmidt (24). Their “overall” activity coefficients, which compensate for ionization and aggregation, being used, the following expression for ln y as a function of the mola1it.y was obtained: In y = -0.549 ln m - 4.891. From this and the solubility equation for ln m (see Table II) are obtained values for b In r/in m and b In m/bT. According to Equation 3, since the activity coefficients were calculated upon that assump tion, AH = 2 RT2 (0.03499) (1 - 0.549). At 25”, AHzsa = 5580 calories.
The true value for the differential heat of solution probably lies bet.ween this value and AHZSS = 6190 calories which is obtained by
assuming that the perfect solution laws are valid for this sub- stance.
The apparent molal volume of the solute was calculated from the density of an approximately saturated solution of Laspartic acid at 25’ (see Table VI). The usual assumption was made at this concentration, p = 02. The value of 76.9 f 0.3 cc. thus obtained agrees roughly with 74.7 cc. calculat.ed on the basis of Traube’s atomic volumes.
d$-Aspartic Acid-Such dat,a as are reported in the literature relative to the solubility of d&spar& acid are more self-consistent than similar data for I-aspartic acid. This includes work by Pas- teur (58), Engel (55), and Michael and Wing (62).
The dl-aspartic acid used was prepared by heating a solution of I-aspartic acid with barium hydroxide in an autoclave. Thirty- two solubility measurements were conducted at nine evenly spaced intervals between O-65”. The usual equations relating the solubility to the temperature were computed and the coefficients summarized in Table II..
Due to the lack of data concerning t,he activity coefficients of dl-aspartic acid, the values for Z-aspartic acid were used to obtain a rough estimate of the heat of solution. Although this is ques- tionable practise, the freezing point depression of the active and the racemic forms of the amino acids are usually approximately the same. Although fully cognizant that the assumption may not be strictly true, the usual calculations were made.
AH = 2 R!P (1.093 X 10-l - 2.302 X 10-J !I’) (1 - 0.549). At 25”, AHzes = 6500 calories.
The true value probably lies between this value and that ob- tained by assuming that the perfect solution laws are obeyed. Thus AH298 = 6500 - 7200 calories per mole at. saturation.
Berthelot (63) measured the integral heat of solution of aspartic acid at 16”, the final concentration being 0.02 M. He made no mention of the optical properties of his material, but the trend of the paper leads to the belief that his material was dl-aspartic acid. He found AHBe = 7250 f 70 calories. From the solubility measurements the heats of solution at this temperature should be AH2as = 5250 - 5820 calories for Laspartic acid and AHSse = 6400 - 7100 calories for d&aspartic acid.
It was noted that dl-aspartic acid has a higher solubiity and its temperature coefficient is greater than the corresponding values
for 1-aspartic acid. According to Meyerhoffer (35) it is question able whether dl-aspartic acid is a racemic compound or a racemic mixture, although t,he higher heat of solution of the dl form would indicate that, it is a racemic compound.
The density of an approximately saturated solution of dZ-aspartic! acid was determined at 25” (see Table VI). From this the appar- ent molal volume was calculated. As in the other cases the assumption was made that the apparent molal volume, (p, and the partial molal volume, ti2, are identical. The value 75.2 f 0.3 cc. t,hus obtained agrees reasonably well with 74.7 cc. calculated from Traube’s empirical atomic volumes.
It will be noted that this value of CZ = 75.2 cc. for dkspartic acid is definitely smaller t,han & = 76.9 cc. observed for Z-aspartic acid. This fact, according to LMeyerhoffer (35), also should indi- cate that dl-aspartic acid is a racemic compound.
d-Glutamic Acid-Although d-glutamic acid has been well characterized from many standpoints, its solubility in water has been but poorly studied. Various isolated determinations have been reported by Pfeiffer and Wiirgler (l), Pfeiffer and Angern (2), Bayliss (5), Schulze (53), Scheibler (64), and Pertzoff (65). The same correlation between the optical rotation and the solubil- ity was noted as in the case of l-aspartic acid.
The optical rotation when 1 mole of the d-glutamic acid, which was used in the present studies, was dissolved in 4.53 moles of hydrochloric acid was found to be [a]z5 = +32.2O which agrees well with [a]: = +32.36” reported by Wood (66). Thirty-five solubility determinations on this compound were carried out at nine evenly spaced temperatures between 0-6OO”. In order to express the solubility as a function of the temperature the usual equations were calculated, the coefficients of which are summarized in Table II.
The most reliable activity coefficient data for d-glutamic acid are those of Hoskins, Randall, and Schmidt (24). The following expression relating their “overall” activity coefficient to the molality was devised: In y = -0.539 ln m - 5.020. From this expression and the solubility equation for In m (see Table II) b In r/d In m -0.539 and b In m/bT = 0.03714 were obtained.
Since the activity coefficients were calculated with that point in view, upon substitution in Equation 3 for dissociating electrolytes,
it is found that AH = 2 RT2 0.03714 (1 - 0.539). At 25’, AH298 = 6050 calories per mole.
The true value for the differential heat of solut,ion probably lies between this value and that estimated by assuming that the perfect solution laws are valid. Thus AH~M = 6050 - 6550 calories per mole at saturation.
Pertzoff (65) reported AH298 = 9600 calories. This value was obtained by combining his solubility value with those of Pfeiffer and Wtirgler (1) and Bayliss (5). The value cited by Bayliss is quite definitely high due to partial racemization. This makes the slope of the curve much sharper with the result that the heat of solution obtained from these values is too high.
The density of an approximately saturated solution of d-glu- tamic acid was determined at 25” (see Table VI). From this the apparent molal volume was calculated in the usual manner. The customary assumption was made that at these concentrations ~0 and & are ident,ical. The value of 90.8 f 0.2 cc. agrees well with 90.7 cc. calculated from the atomic volumes cited by Traube.
dl-Glutamic A&d-The only available solubility data relating to dZ-glutamic acid are the widely divergent values reported by Lii (67) and by Schulze (53).
The present material was prepared by racemizing d-glutamic acid with barium hydroxide in an autoclave. Twenty-four solu- bility measurements were conducted at nine different temperatures between O-65”. The usual solubility-temperature ,equations were computed and their coefficients tabulated (see Table II).
As there are no available data from which the activity coeffi- cients of dLglutamic acid can be calculated, the values may be assumed to be similar to those found for d-glutamic acid by Hos- kins, Bandall, and Schmidt (24) inasmuch as the freezing point lowerings of the active and racemic forms are usually approxi- mately the same. Using these data we obtain b ln r/d ln m= -0.539 and b ln m/dT = 0.03507 (from Equation 3 in Table II). Then by substitution in Equation 3 we obtain AH = 2 RT2 (0.03507) (1 - 0.539).
The true value of the heat of solution probably lies between the value indicated by this equation and that indicated by assum- ing that the perfect solution laws are obeyed. Therefore, AHzss = 5710 - 6180 calories per mole.
It was noted by comparing the solubility data of d- and dl- glutamic acid that the solubility of the dE form is greater and the temperature coefficient is less than the corresponding values for the d form. The heats of solution are also different. According to the rules outlined by Meyerhoffer (35), dl-glutamic acid may be either a racemic compound or a racemic mixture.
The apparent molal volumes were calculated from the density data of two solutions of dl-glutamic acid which were determined at 25” (see Table VI). The average value of 90.5 f 0.2 cc. agrees well with 90.7 cc. calculated from the atomic volumes given by Traube. It was noted in this case that the partial molal volumes of d- and dl-glutamic acid are identical within the errors of the determination. This indicates that dl-glutamic acid is a racemic mixture. Further data on the point should, however, be obtained by mixed fusion point det~erminations, crystal examination, or some other suitable means.
dl-PhenyluIunke-The data in the literature concerning the solubility of dl-phenylalanine are also fragmentary. Those re- ported by Pfeiffer and Angern (2), Bayliss (5) and Schryver (68) when recalculated to the present basis were found to be from 12 to 28 per cent higher than the values interpolated from the present data.
Twenty-nine solubility measurements upon the present material were carried out at ten different temperatures between O-75”. The solubility4emperature equations were computed in the usual manner, the coefficients of which are summarized in Table II.
With the absence of any data concerning the activity coefficients for aqueous solutions of dZ-phenylalanine, the calculation of the differential heat of solution becomes only approximate. With the equation for In NZ (see Table II), differentiation, and substitu- tion of the proper terms in Equation 1 we obtain AH = RT2 (-2.495 x 10-t - 1.362 X 10”’ T). Therefore at 25”, AHtg8 = 2770 calories. For the reasons outlined for dl-isoleucine, 2770 calories may be regarded as a minimum value for the differential heat of solution at this temperature.
The density of an aqueous solution of dl-phenylalanine was determined at 25’ (see Table VI). From these data the value of 123.3 cc. was obtained for the value of the apparent molal vol- ume, ‘p. The value of 130.6 cc. calculat.ed from Traube’s atomic volumes agrees roughly with the present value of 123.3 cc.
dGNorZeueins--The solubility value report,ed by Kudielka (69) for dl-norleucine is the only one previously reported for t.his amino acid.
Thirty-two measurements were carried out at ten different tem- peratures between O-75’. Equations were derived in t,he usual manner to express these values as functions of the temperature. The coefficients of these equat,ions are summarized in Table II.
Since no data have been reported in the literature concerning t,he activity coefficients for dl-norleucine, only a rough approxima- tion can be made for its differential heat of solution. Following t,he reasoning outlined for dl-isoleucine, from the equation for In NZ (Table II) we find that AH = RT2 (-2.941 X 10m2 + 1.468 X lo4 T). The minimum value for the heat of solution at. 25” then becomes AHzss = 2540 calories.
The density of a solution of this compound was determined at a concentration near saturation (see Table VI). From this the apparent molal volume was calculat.ed and assumed to be equiva- lent to t’he true partial molal volume. The value, fiZ = 110.0 cc., compares roughly wit,h 107.1 cc. calculat.ed from Traube’s mean atomic volumes.
I-Tyrosine-The solubility data for I-tyrosine reported by Schulze ((53) p. 99), von Euler and Rudberg (3), PfeifYer and Angern (2), Ando (38), Sano (46), Hitchcock (70), and Erlenmeyer and Lipp (71) are limited solely to isolated determinations, all of which ranged from 1 to 36 per cent higher than the values int,erpo- lated from the present work. The rotatory powers, when reported by t.hese authors, were lower t*han that found for the present material. The I-tyrosine used showed [CY] 2 = - 16.1” for a 5 per cent solution of tyrosine in 4 per cent hydrochloric acid. This value agrees well with [a]: = -16.2’ reported by Schulze and Winterstein (72) under the same conditions.
Thirty-two solubility measurements were made at nine different temperatures between 0-60’1 The solubility-temperature equa- tions were derived in the usual manner and their coefficients summarized in Table II.
For the calculation of the differential heat of solution, the perfect solution laws were assumed to hold. From the equation for In Nz (Table II) is obtained AH = 0.0337 RT2. At 25’, AH288 = 5960 calories per mole.
Diiodo-l-Tyrosine-The solubility data for diiodo-l-tyrosine have been reported elsewhere (73, 74). They were recalculated and included here to complete the present series. Oswald (75) reported a solubility value between 6 and 7 times greater than the present figures indicate. His material was isolated from natural sources while the present product was synt,hetically prepared from active tyrosine. Any doubt that this synthetic product might be
TABLE VII
Di~erential Heats of Solution of Certain Amino Acids
Ez (289”) 7830 6050 5710 3370 3430 (289O) 1790 830
2070 2540 2770 6960 1590
2140 rt 30 2020 (2sP)
7250 f 70 (289’)
3.530 * 30 3530 f 40 (289°)
1590 f 30
* Corrected for the activity of the amino acid. For a discussion of this subject and the assumptions made consult the text.
racemized was removed by Mr. Winnek of this laboratory when he reduced a sample of the diiodotyrosine to tyrosine by the pro- cedure outlined by Harington (76). No essential change in the optical activity of the tyrosine was noted.
Twelve solubility measurements were carried out at seven differ- ent t,emperatures between O-47”. Higher temperatures were not used as diiodo-l-tyrosine shows some tendency to decompose. The
usual equations were devised and their coefficients t,abulated (see Table II)..
By assuming that the perfect solution laws were obeyed, from the equation for ln NS, AH = 0.0443 RT2 is obt,ained. At 25’, AHzss = 7830 calories per mole.
For the sake of convenience a summary of the differential heats of solution of the amino acids studied in this paper is given in Table VII.
The original data upon which the calculations presented in this paper are based are on file in the University of California library.
Work on the determination of the solubilities of amino acids other than those reported in this paper is in progress.
SUMMARY
1. A review of the literature shows not only a paucity but a con- flict in the solubility and densit.y data of the amino acids. Reasons to account for this have been advanced.
2. The solubilit,ies of t,he following amino acids have been determined over various given temperature ranges: d- and dl- alanine, l- and dLaspartic acid, d- and dl-glutamic acid, l- and dl- leucine, diiodo-l-tyrosine, glycine, dl-isoleucine, dl-norleucine, dl-phenylalanine, l-tyrosine, and dl-valine.
3. Equations have been devised for each amino acid to express the solubility as a function of the temperature. The values calculated from these expressions usually deviate from the ob- served values less than 1.5 per cent.
4. The heats of solution of the amino acids mentioned above were calculated and in some instances were compared with direct calorimetric measurements.
5. The molecular volumes of the amino acids in solution have been calculated from Traube’s empirical atomic volumes. A com- parison of the values so obtained with those found by the estima- tion of the partial molal volumes from the density measurements has been made.
6. Meyerhoffer’s criteria concerning the densities of solutions containing optical isomers were applied to solutions of the amino acids studied. The results substant,iated the conclusions based on the solubility measurements.
7. It was concluded from these criteria that dl-alanine and dL
leucine are racemic compounds, that dl-aspartic acid is probably a racemic compound, and that dl-glutamic acid is probably a racemic mixture.
8. A table giving the solubilities of the amino acids named above at 5” intervals from O-75’ and at 100’ has been included (Table I).
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