FUNCTION O F C.4RBOKATE IK SYNTHESIS O F GLTCISE
1123
(1)E r ~ z r a ~J., F. : Recherches Refractometriques. (5) (6)
(7) (8) (9)
(10) (11)
(12)
de Erven Loosjes, Haarlem (1919). GROSSE,A. V., A N D EGLOFF,G.: Physicul Comtants o,f Parafin Hydrocarbons. Bulletin 219, Universal Oil Products Company, Chicago (1938). HULST,L. J. K.V A N D E R : Rec. trav. chim. 64, 518 (1935). KRAFFT,F.: Ber. 16, 1687 (1882). MAIR,B. J., A N D STREIFF,A. .J. : ,J. Research T a t l . Bur. Standards 24, 395 (1940j . SHEPARD, A. F., HEKSE, A. L . , ASD MIDGLET, T., JR.: J. .hi, Chem. Soc. 63, 1918 (1931). T I M M E R M A N S , J., A N D l , i A R T I S , F.: J. chim. phys. 26, 111 (1928). TIN\IERMASS, J., A N D HESS.AST-ROL.ASD, hfarz: J. chim. phys. 32, 501 (1935). \vIIIB.%UT, J. P., HOOG, H., LASGEDIJK, s. L., O V E R H O F F , J . , A S D S M I T T E S B E R G , J.: Rec. trav. chim. 68, 329 (1839).
QL-ASTITATIVE ISVESTIGA4TIONS O F AXIhTO ,1CIDS ASD PEPTIDES. VI THE FUNCTIOK OF CARBONATE IN
SYNTHESIS OF GLYCINEFROM HYDROXIDE, AND A M M O N I ~CARBOKATE' M THE
CHLORO.4CETIC ACID, i h M O S I U h l
LIAY S DV:",
A W BUTLER,
AND
EDWARD H F R I E D E S
Department of Chemastry, Cnauerszty of Calafornza, Los Angeles, Culzfornza Recezued February 16, 1041
The first synthesis of glycine from a-halogen acids and ammonia,namely, that reported in 1858 by Perkin and Duppa (12),-was folloned by a number of investigations (4), of nhich the one by Robertson (11) is the most comprehensive. In the early studies glycine was isolated as the copper salt in order to separate it from the mixture of the amino acid and the ammonium salts of hydrochloric (or hydrobromic) acid, iminodiacetic acid (so-called diacid), and trimethyleneaminetricarboxylic acid (so-called triacid) (equations 1 to 3, table 1). Studies by Robertson (11) on the rates of formation of chloride ion and primary amine under varying conditions led to the discovery that the production of secondary products is markedly depressed, and the yield of glycine is increased nearly to the theoretical quantity, by the use of high 1 Presented before the annual meeting of the American Society of Biological Chemists, held in Toronto, Canada, April, 1939 (see reference 5). The authors were aided in this work by a grant from the Univemty of California. For the fifth communication in this series see Dunn and Porush: J Biol Chem. 137, 2G1 (1939)
1124
M. S. D U " ,
A. W. BUTLER AND E. H. FRIEDEN
molar ratios of ammonia to chloroacetic acid. The molar ratio of ammonia to chloroacetic acid of 60 to 1, which was adopted for practical use, was shown by Orten and Hill (10) to give approximately 65 per cent of the theoretical yield of pure glycine when methyl alcohol is used to precipitate the amino acid from the reaction mixture. Ammonium carbonate was employed as the aminating agent for a-halogen acids by Nencki (9), Slimmer (14),and Cheronis (3). When a mixture containing a 4 to 1 molar ratio of ammonium carbonate to chloroacetic acid was heated for 12 hr. at 65OC., Cheronis found that glycine was formed in 60 per cent of the theoretical yield. Because of the unusual interest of these results, it seemed desirable to investigate the influence of temperature, molar ratios of ammonia to carbonate to chloroacetic acid, and partial pressures of ammonia and carbon dioxide on the formation of glycine. The results of these experiments are given in the experimental part. TABLE 1 Postulated reactions occurring i n mixtures of chloroacetic acid, ammonium hydroxide, and ammonium carbonate solutions
+ + + + + + +
+
+
2NHa ClCHzC00- -+ NHd2HzCOONHT C1XH&HnCOOClCHzC00- -+ KH(CH&OO-)S SHa NHa NH(CHzCOO-)n ClCHzC00- -+ N(CH&OO-)a H+ HCOa e H+ cos-COS Ha0 Hz0 :2" Cos-COS 2NHa 2"s NHzCOONH: COz NHzCHzC00- F! -0OCKHCHZCOOH+ COz KHzCOOClCHzC00NHa -0OCNHCH2COO-
+
+
+
+ +
+
+
+
+
+
+
+
+ NH: + C1+ NH: + C1+ iiH: + C1-
EXPERIMENTAL
Fractionally distilled chloroacetic acid, C.P. 15 N ammonium hydroxide, and commercial grade ammonium carbonate were used throughout. The chloroacetic acid and the ammonium carbonate solutions were standardized by titration procedures and the ammonium hydroxide by density determinations. Variation of the carbonate to chloroacetic acid ratio while the ammonia concentration was maintained constant was accomplished by the introduction of potassium carbonate. Measured quantities of the reaction mixtures were placed in special, long-necked, 50-ml. Florence flasks and, at recorded times, the flasks were immersed in a thermostat at the desired temperature. The flasks were removed from the thermostat at stated intervals, and the reaction mixtures were cooled rapidly by immersing the flasks in ice water. The concentrations of chloride ion and of primary amine were determined by Volhard and Van Slyke analyses, respectively. Solutions to be analyzed for primary amine were first made alkaline with sodium hydroxide and boiled to remove
FUNCTION OF CARBONATE I N SYNTHESIS OF GLYCINE
1125
ammonia. All analyses were made in duplicate with 1 per cent or higher precision. Considerable difficulty was experienced in measuring the partial pressures of ammonia and carbon dioxide with satisfactory precision and accuracy. The manometric Van Slyke apparatus, with sulfuric acid in the Hempel pipet for the absorption of ammonia, was first used for this purpose. This method was developed until partial pressure measurements of ammonia over aqueous ammonia solutions could be duplicated to 2 to 3 mm. of mercury, but the results with ammonium carbonate solutions were unsatisfactory. Copper sulfate solution for the absorption of ammonia and sodium hydroxide solution for the absorption of carbon dioxide were tried without success. It was found, also, that the Blacet (1) micro gas
FIG.1. Apparatus for collecting known volumes of carbon dioxide and ammonia measured at atmospheric pressure.
apparatus and procedure, as well as an aspirator method, bvere unsatisfactory for the present purposes. A static vapor pressure method and the calibrated apparatus shown in figure 1 were utilized finally to measure partial pressures of carbon dioxide and ammonia. The following technique was used: Bulb C is filled ivith mercury to stopcock F. About 500 ml. of the experimental solution is placed in bulb A and allowed to run into bulb B until the latter is half filled. Stopcock F is closed. Bulbs A and B are attached at J to a motor and shaken for 15 min. to attain equilibrium between the liquid and gaseous phases. Stopcock F is opened and bulb D lowered sufficiently to permit the gas to flow from bulb B into C down to mark H while the level of the mercury in D is a t H. Stopcock G is closed and F opened to connect with the absorbers. The gas is run into the absorbers a t the rate of about 10 ml. per minute.
1126
M. S. DUNN, A. W. BUTLER AND E. H. FRIEDEN
The absorbers consisted of two Geissler bulbs, connected in series, one containing 3.5 N sulfuric acid and the other saturated barium hydroxide solution. A 3.5 Ai sulfuric acid solution mas used as the absorbent for ammonia because it has approximately the same partial pressure of water vapor as the experimental solutions and dissolves a negligible quantity of carbon dioxide. Although the aqueous tension of the barium hydroxide solution is somewhat higher than that of the experimental solutions, the error introduced was calculated to be small. The quantities of ammonia and carbon dioxide absorbed were determined from the increase in weight of the Geissler bulbs, and the partial pressure of each gas was calculated from the following expression:
P= W X R X T *If x Ti where P is the pressure in atmospheres, W is the grams of absorbed gas, M is the molecular weight of the gas, l7 is the volume in milliliters of bulb E to the mark H, R is 82.06 cc.-atmospheres, and T is the Absolute temperature. The average value for the experimentally determined partial pressure of ammonia over an aqueous ammonia solution of known concentration deviated 6.8 per cent from the best values previously reported (8). This agreement was less close than desired, although vapor pressures measured by this method were considered to be precise enough for the intended purpose. Because of the possibility that some of the chloride might arise from hydrolysis, as uell as amination, of the chloroacetic acid, the degree of hydrolysis under the experimental conditions was measured. The pH of the experimental solutions, one of the significant factors, was found to vary from 9.3 to 11.0, according to measurements made by means of standard buffer solutions and the indicators thymolphthalein, thymol blue, and bromothymol blue. Standard solutions of sodium chloroacetate a t pH 12, prepared by dissolving chloroacetic acid in sodium hydroxide solution, were maintained for 27 hr. in a thermostat a t 45OC. The degree of hydrolysis, calculated from the experimentally determined chloride concentration, n-as found to be only 1.1 per cent. DISCUSSIOX
I t was found that (a) the time required for the production of a given quantity of primary amine a t constant ratio of ammonia to carbonate is inversely proportional to the concentration of chloroacetic acid and ( b ) the maximum percentage yield of primary amine is the same for the two concentrations (0.2 and 0.4 N) of chloroacetic acid investigated. The data from these experiments are omitted to conserve space.
FUSCTIOS O F C.%RBONATE IK SYNTHESIS OF GLYCINE
11s
It is an obvious deduction from the authors’ experiments (compare the curves in figures 2 and 6 , 3 and 7 , 4 and 8, and 5 and 9) that the rate of production, as well as the yield, of primary amine (glycine) in a given time increases with increasing temperature, although this effect is more marked a t ion-, than a t high, molar ratios of ammonia to chloroacetic acid. It wm observed that the mass action influence of ammonia on the production of glycine (compare the curves in figures 2 , 3,4, and 5, and those in figures 6, 5 , 8, and 9) is highly significant, as was first emphasized by Robertson (11). The assumption that ammonia is the principal, if not the sole, aminating
TIME
OF
R E A C T I O N IN
HOURS
FIG.2 . Curves showing the rate of production of chloride and primary amine at 25°C. The d a t a are plotted according to the following notation: 0 , C1-, and 0 , primary amine, in solutions of ammonia and chloroacetic acid in the equivalent ratio 5 to 1; and H , C1-, and A, primary amine, i n solutions of ammonia, carbonate, and chloroacetic acid in the equivalent ratio 5 t o 4 to 1. I n this and all other figures designated quantities of ammonia refer to total ammonia present as SH,and XH4+
agent is supported by the observations of Groggins and Stirton (7), who showed that ammonolysis may be carried out with anhydrous ammonia in the vapor phase or in an inert solvent. The traces of amines which are formed in solutions of ammonium salts were attributed by these authors to result from the action of ammonia formed in small amounts by the hydrolysis of KHZ. The hypothetical substance, NH40H, is assumed to be non-existent. The net effects of carbonate upon the production of primary amine may be deduced from a consideration of the curves in figure 12. When the
d sp
TIME
OF R E A C T I O N IN HOURS
FIG.3. Curves showing the rate of production of chloride and primary amine a t 25°C. The data are plotted according to the following notation: 0 , C1-, and 0 , primary amine, in solutions of ammonia and chloroacetic acid in the equivalent ratio 10 to 1; and., Cl-, and A, primary amine, in solutions of ammonia, carbonate, and chloroacetic acid in the equivalent ratio 10 t o 4 to 1.
FIG.4. Curves showing the rate of production of chloride and primary amine a t 25'C. The d a t a are plotted according to the following notation: 0, C1-, and 0 , primary amine, i n solutions of ammonia and ohloroacetic acid in the equivalent ratio 15 t o 1; and m , C1-, and A , primary amine, i n solutions of ammonia, carbonate, and chloroacetic acid in the equivalent ratio 15 t o 4 t o 1. 1128
0
10
20
30
40 50 60 10 TIME OF REACTION IN HOURS
SO
90
100
FIG.5 . Curves showing the rate of production of chloride and primary amine a t 25°C. The data are plotted according to the following notation: 0, C1-, and 0 , primary amine, in solutions of ammonia and chloroacetic acid in the equivalent ratio 30 to 1;and , GI-,and A, primary amine, in eolutiona of ammonia, carbonate, and chloroacetic acid in the equivalent ratio 30 to 5 to 1.
TIME
OF REACTION IN HOURS
FIG.6. Curves showing the rate of production of chloride and primary amine a t 45OC. The data are plotted according to the following notation: 0, C1-, and 0 , primary amine, in solutions of ammonia and chloroacetic acid in the equivalent ratio 5 to 1; and , C1-, and A, primary amine, in solutions of ammonia, carbonate, and chloroacetic acid in the equivalent ratio 5 to 4 to 1 . 1129
4 E
3 #
I=
2
s9
0
5
10
I5
20 TIME OF
25
30
35
40
45
50
REACTION IN UOURS
FIG.7. Curves showing the rate of production of chloride and primary amine at 45°C. The data are plotted according to the following notation: 0 , C1-, and 0 , primary amine, i n solutions of ammonia and chloroacetic acid in the equivalent ratio 10 t o 1; and H , C1-, and A, primary amine, in solutions of ammonia, carbonate, and chloroacetic acid in the equivalent ratio 10 to 4 to 1.
TIME OF WACTION IN HOURS
FIG.8. Curves showing the rate of production of chloride and primary amine at 45°C. The d a t a are plotted according t o the following notation: 0,C1-, and
0,
primary amine, i n solutions of ammonia and chloroacetic acid in the equivalent ratio 16 t o 1; and H , C1-, and A, primary amine, in solutions of ammonia, carbonate, an% chloroacetic acid in the equivalent ratio 15 t o 440 1.
1130
FUNCTIOS OF CARBONATE IS STSTHESIS
OF GLTCISE
1131
FIG.9. Curves showing the rate of production of chloride and primary amine a t 45°C. The data are plotted according to the following notation: 0 , C1-, and 0 , primary amine, in solutions of ammonia and chloroacetic acid in the equivalent ratio 30 to 1; and , C1-, and A , primary amine, in solutions of ammonia, carbonate, and chloroacetic acid in the equivalent ratio 30 to 5 to 1.
PRESWRE Of
NH, IN M U
OF
HG
FIG.10. Curves showing the partial pressures of ammonia in millimeters of mercury in solutions of ammonia and chloroacetic acid (curve A ) , and in the presence of Sequivdents of carbonate per mole of chloroacetic acid (curve B). The equivalent ratio of ammonia to chloroacetic acid varies from 0 to 1 to 38 to 1.
1132
M. 8. DUNN,
-4.
W. BUTLER AND E. H. FRIEDEN
0
10 20 30 40 : MOLES OF NH, PER
MOLE OF CNOROACETIC ACID
FIG. 11. Curves showing the time for maximum production of primary amine a t 45°C. in solutions of ammonia, carbonate, and chloroacetic acid in the equivalent ratios 5-40 to 4 to 1 (Curve A) and 640 to 0 to 1 (curve B).
FIG.12. Curves showing the percentage of the theoretical primary amine produced a t 45°C. in solutions of ammonia, carbonate, and chloroacetic acid in the equivalent ratios 5-60 to 4 to 1 (curve NHa CO1) and 5-220 to 0 to 1. The points noted a8 A refer to data reported by Robertson (11).
+
FUNCTION OF CARBONATE IN SYNTHESIS OF QLYCINE
1133
equivalent ratio of carbonate to chloroacetic acid in the reaction mixture is 4 to 1, only about a 15 t o 1 equivalent ratio of ammonia to chloroacetic acid is required to give 85 per cent of the theoretical amount of primary amine. When carbonate is absent, the ratio of ammonia to chloroacetic acid required to give a comparable yield of amine rises to about 50 to 1. It may be noted (figure ll), however, that a considerably longer time is required for the production of a given yield of primary amine in the presence, than in the absence, of carbonate, particularly a t low equivalent ratios of ammonia to chloroacetic acid. On the other hand, the yield of primary amine in a given time with a constant equivalent ratio of ammonia to chloroacetic acid first rises to a maximum a t an equivalent ratio of carbonate to chloroacetic acid of about 5 to 1 and then falls off as the eo
9 '$
70
i
55 4
'0
40
i
Y f
30
20
0
5 IO I5 EQUIVALENTS Of K.CO, PER EQUNALWT O f CHLOROACCTIC ACID
20
FIG.13. Curves showing the percentage of the theoretical prmairy amine produced in 43 hr. a t 25°C. and 20 hr. a t 45°C. in solutions of ammonia, carbonate, and chloroacetic acid in the equivalent ratios 10 to 0-20 to 1. Under theseconditions the yield of chloride was approximately 100 per cent (at 45°C.) and 58 per cent (at 25°C.) of the theoretical amount.
equivalent ratio of these constituents is increased to 20 to 1 (figure 13). At the same time the yield of chloride (not shown in figure 13) remains constant. Since the percentage yield of primary amine produced in solutions containing ammonia, carbonate, and chloroacetic acid in an equivalent ratio of 15 to 4 to 1 is as high as in solutions containing only ammonia and chloroacetic acid in an equivalent ratio of 60 to 1, it is evident that relatively small concentrations of carbonate are as effective as high concentrations of ammonia in preventing the formation of diacid and triacid. It is proposed, as a logical explanation for these observations, that some complex of glycine is formed in ammonia-carbonate-chloroacetic acid reaction mixtures which is stable in alkaline solution but is decomposed
1131
M. S. DU",
A. W. BUTLER AND E. H. FRIEDEN
in acid solution during the analysis or isolation of glycine. A natural assumption is that the complex of glycine is ammonium carbaminoglycinate, since Siegfried (13) has prepared the calcium and barium carbamates of glycine and other amino acids and has presented evidence pointing to the existence of carbaminoglycine in solution a t 0°C. It would appear from the investigations of Stadie and O'Brien (15) that glycine carbamate may be formed by the reaction of carbon dioxide and the glycinate anion (equation 7 , table 1). Although it was observed that glycine does not combine with carbonate ion, bicarbonate ion, or carbon dioxide (at pH 6.5)' these authors found that glycine carbamate is formed
FIG.14. Curves showing the initial (curve A) and equilibrium (curve B) vapor pressures of ammonia in solutions of ammonium hydroxide, potassium carbonate, and chloroacetic acid. The molar ratio of ammonium hydroxide to chloroacetic acid is constant a t 10 t o 1. (Equilibrium values were measured 10 hr. after initial values.)
from glycine and carbon dioxide in alkaline media. The proportion of glycine carbamate was shown to increase with increasing alkalinity of the reaction mixture. That carbonate ion per se is not instrumental in forming glycine carbamate may be inferred from a consideration of the curves shown in figures 10, 13, and 14. The increased yield of glycine resulting from the introduction of potassium carbonate into ammonia solutions may be attributed to the increase in ammonia vapor pressure, as shown in table 2. The partial pressure of the carbon dioxide in the authors' reaction mixtures for which the equivalent ratio of ammonia to carbonate t o chloroacetic acid was 5 to 4 to 1 was approximately 110 mm. There was a
1135
FVTNCTION OF CARBONATE IN SYXTHESIS OF GLYCINE
marked effect of the carbonate upon the amine production under these conditions. On the other hand, a definite carbonate effect was observed in mixtures containing 10 or more equivalents of ammonia to 4 to I of the other constituents, even though the partial pressure of the carbon dioxide was too lon- to measure. It may be concluded, therefore, that the carbonate effect is not explained by the reaction of carbod dioxide and glycinate ion to give glycine carbamate. =in alternative explanation would seem to be that ammonium carbaminoglycinate is formed from carbamate ion and chloroacetic acid (equation 8, table 1). Some justification for this hypothesis is found in the investigations of Faurholt (6) and of Stadie and O’Brien (15), who observed that the reactions between carbon dioxide, ammonia, and water come to equilibrium slo~vlga t room temperatures. From investigations by Buch (2), Wegschneider (16), and Faurholt (6) it was shown that independent equilibria are established between ammonia, carbon dioxide, and ammonium carbonate on the one hand and ammonia, carbon dioxide, and ammonium carbamate TABLE 2 Equilibrium pressures of ammonia and maximumpercentage yields of glycine in solutions of ammonia, potassium carbonate, and chloroacetic ocid C~C~,COOH
“ I
mOlaS
mlda
10 10 15
0 10 0
”
1,*,,
1 1
1
MAXIMUM YIELD OF GLYCINE
EQUILIBRIUM PRESS U R E O F AMMONIA
per cent
millimeferr
50 60 60
45 80 80
on the other in any mixture of these component substances. Faurholt observed, also, that the formation of ammonium carbamate is a slow reaction, since he was able to determine carbonate quantitatively as barium carbonate in the presence of carbamate. It may be observed from the curves in figure 2 that, although the rate of production of glycine is greatly diminished, the formation of by-products is reduced practically to zero when carbonate nearly equivalent to the ammonia is present. Since, under these conditions, there is no measurable partial pressure of ammonia (figure lo), it is a logical conclusion that glycine can be formed only by a reaction which involves some aminating agent other than ammonia. If carbamate ion is this reactant, its action on chloroacetic acid must be slower than that of ammonia. It would appear that the addition of carbonate to mixtures of chloroacetic acid and ammonia results in shifting of the equilibria (reactions 5, 6, and 8, table 1) to favor increased production of carbon dioxide, carbamate ion, and carbaminoglycinate. As the result of these reactions the concentra-
1136
>I. 8. D-,
A. W. BUTLER AND E. H. FRIEDES
tion of glycinate ion is decreased, the production of diacid and triacid is diminished, and the yield of glycine is increased. In carbonate zolutions of low ammonia Concentration, the partial pressure of carbon dioxide is high, the partial pressure of ammonia is lon., and there i+ a marked carbonate effect, because the carbamate ion, assumed to be the aminating agent, k formed in relatively high concentration. In carbonate solutions of relatively high ammonia concentration and partial pressure, the carbon dioxide is depressed to negligible proportions. Severtheless. there is a marked carbonate effect, because the equilibrium condition; are such that the carbamate-ion concentration is increased to a relatively high level. St3IMARY
1. The influence of temperature, molar ratios of ammonia to carbonate to chloroacetic acid, and partial pressures of ammonia and carbon dioside on the production of primary amine (glycine) and chloride in reactions of chloroacetic acid with ammonia, ammonium carbonate, and mixtures of these aminating agents ha+ been investigated. 2. It has been shown that the rates of these reactions are sloiver, but the yield of glycine is higher, in the presence than in the absence of carbonate. 3. That relatively small concentrations of carbonate are as effective as high concentrations of ammonia in increasing the yield of glycine is believed to be explained by formation of ammonium carbaminoglycinate, v-hich diminishes the concentration of glycinate ion and depresses the production of diacid and triacid. REFERESCES (1) BLACET, F. E., ASD LEIGHTOS,P. A : Ind. Eng. Chem., rinal. Ed. 3, 266 (1931). ( 2 ) BIXH, K.: Z . physik. Chem. 70, 66 (1910). (3) CHEROSIS,S. D.: Ahstracts of papers presented a t the Chicago meeting (September 10-15, 1933), the Pittsburgh meeting (September 7-10, 1936), and 9-13, 1940) of the American Chemical the Detroit meeting- (September .
Society. (41 D c s s . 11.9 . : I n Chenlistru of the Amino Acids and Proteins, edited by C. L. 1. Thomas, Springfield, Illinois (1938). Schmidt, p. 69. C. \ - I
6.
(5) D n s s , 11.S., BTTLER.A , K . , ASD FRIEDES,E. H.: Proc. hm.SOC.Biol. Chemists 9, S S I I (1939). (6) FACRHOLT, C.: 2. anorg. Chem. 120, S5 (1922). (7) GROGGISS, P. H.. ASD Q T I R T o s . -\. J.: Ind. Eng. Chem. 26,12 (19331. (8) International Critical Tables, Vol. 111, p. 362. JIcGraw-Hill Book Company, Inc., Yew Tork 119251. (9) SESCKI, 11.:Ber. 16, 2 9 i (1Y83). (10) ORTES,J. Il,, ASD HILL.R . 11,:J. Am. Chem. SOC. 63, 2797 (1931); Organic Syntheses. Coll. Vol, I , p. 253. John Wiley and Sons, Inc., S e x f o r k (1932). (11) ROBERTSOS, C. R . : J. Im.Cheni. SOC.49, ZSS9 (1927). (12) PERKIS.IV. H., .\SD DCPPA,B. F.: .4nn. 108, 112 (1858).
COMMUNICATION TO TH’E EDITOR
1137
(13)SIEGFRIED, M.:Z.physiol. Chem. 44, 85 (1905);48,401 (1905). (14)SLIMMER, M.D.:Ber. I,400 (1902). (15) STADIE,W.C., AND O’BRIEN,H. : J. Biol. Chem. 112, 723 (1936). (16) WEGBCHNEIDER, R.:Monstsh. 87,425 (1916). \
Addendum
Since this paper was submitted for publication, Cheronis2has published a detailed description of his investigations in which the experimental approach was similar to that employed by the present authors. The effects of ammonium salts, as well as the influence of pH upon the production of amino acids, were studied by Cheronis and Spitzmueller. The present studies were confined to chloroacetic acid, while chloropropionic acid and seven bromo acids were investigated by Cheronis and Spitzmueller. The investigations of the present authors on the effect of varying concentrations of ammonia and carbonate upon the production of glycine appear to be somewhat more comprehensive than those of Cheronis and Spitzmueller, and our studies upon the relation of the partial vapor pressures of the various constituents above the reaction mixtures remain the only ones available. The conclusions reached by the present authors concerning the function of ammonium salts in increasing the yield of glycine and the mechanism of the ammonolytic reactions are essentially in agreement with those of Cheronis and Spitzmueller. The increased yield of glycine obtained with ammonium salts other than the carbonate may be attributed to the increase in the equilibrium concentration of ammonia in the system. When ammonium carbonate is present, the unstable amino acid carbonate is agreed by both authors to be an easential factor in the production of high yields of glycine. ‘Cheronie, N. D., and Spitzmueller, K. H.: J. Org. Chem. 6, 349 (1941).
COMMUNICATION TO T H E EDITOR T H E MOLECULAR PROPERTIES OF LIGNIN SOLUTIONS’
Dr. H. C. Howard has recently called our attention to an error in our paper on “The Molecular Properties of Lignin Solutions” (J. Phys. Chem. 40, 1117 (1936)). In calculating the shape factor of the lignin molecule from viscosity data, using the Eisenschitz equation, the constant was taken as 1.59 instead of 0.159. The shape factor so calculated was 7.5, 1
Received June 3, 1941.
