MOLAL HEAT CAPACITIES AND VOLUMES
309
liberated when chloroform is mixed with a polyethylene glycol ether. This difference is explained as due to intermolecular association of the polyamines through K-H+-?; bonds. Because of this solvent association, the locations of the maxima in the heat of mixing curves for the polyamines do not indicate how many of the nitrogen atoms in the polynminrs are available for bonding. REFEREYCES
(1) BUBWELL, RODEBUSH, AND ROY:J. Am. Chem. Sac. 60, 2528 (1938). ZELLHOEFER,AND MARVEL:J. Am. Chem. Sac. 60,2666,2714(1938). (2) COPLEY, (3) GINSBERG,E.:Private communication. (4) GORDYAND STANFORD: J. Chem. Phya. 8, 170 (1940). AND BONNER: J. Chem. Phys. 7, 880 (1939). (5) KIRBY-SMITH (6)MCLEODAND WILSOX:Trans. Faraday Sac. 31,596 (1935). (7) ZELLHOEFERAND COPLEY: J. Am. Chem. SOC.60,1343 (1938). (8) ZELLHOEFER,COPLEY, AND MARVEL: J. Am. Chem. Sac. 60, 1337 (1938).
T H E APPAREST A S D PARTIAL MOLAL HEAT CAPACITIES AXD VOLUMES O F GLYCISE A S D GLYCOLAMIDE. I1
RESULTSFOR CONCENTRATED SOLUTIONS OF GLYCOLAMIDE FRASK T. GUCKER, JR., AND WILLIAM L. FORD
Department of Chemistry, Northwestern University, Evanston, Illinois Received April 26, 1940
In a comparison of the properties of glycine and glycolamide (I), presented before the Symposium on Intermolecular Action, which was held at Brown University in December, 1938,we reported densities and specific heats of glycine over the whole range of concentration, and those of glycolamide up to 2.3 iM. Measurements a t several higher concentrations, necessary to complete this series, could not be made in time for the Symposium. They are given in the present article. MATERIALS AND SOLUTIONS
In preparing the large amount of glycolamide required for this work, we first used the method of Edsall, described in our previous paper (1). Later we found that, by passing dry gaseous ammonia into freshly distilled ethyl glycolate, without the use of alcohol as a solvent, the yield was increased from 50 per cent to over 80 per cent of purified material. The white product, after three crystallizations from 95 per cent alcohol, had electrolytic impurities of only about 0.001 per cent and melted a t the same
310
FRANK T. QUCKER, JR., AND WILLIAM L. FORD
temperature as Dr. Edsali's material. This was 117°C.(corrected), as measured with a calibrated thermometer, and agreed with the value given by Heilbron (3). I n one case, a trace of an orange-colored organic impurity, which persisted in the crystallization from alcohol, waa removed easily by means of animal charcoal. I n order to find out how concentrated a solution could be used for measurements at 5"C., a rough determination of the solubility of glycolamide at 0°C.was made. After 11.66 g. of a 6.04 m solution had stood in an ice bath for 6 hr., the mother liquor was decanted and the crystals were found to weigh 1.84 g. These figures give for a minimum solubility 23 g. of glycolamide in 100 g. of water, corresponding to a 3.0 m solution. A value of 3.3 m for the saturated solution, which probably is more nearly correct, waa obtained by a method which was somewhat more complicatcd. As before, the solutions were made up determinate, in sufficient quantity for the measurement of both density and specific heat. They were 80 concentrated that shaking for a considerable time was required to insure homogeneity. The first solution of 4.4 m concentration, used for experiment No. 7 of the specific heats and experiment No. 17 of the densities, was not adequately stirred. The density of the residual solution in the flask was found to be appreciably greater than that in the calorimeter. Because of the uncertainty in the concentration, these results were discarded. Two other separate solutions, made up to the same concentration intended for the first, gave results for the density and specific heat which checked satisfactorily. EXPERIMENTAL RESULTS
The experimental methods have been described in the previous paper. The present results are numbered chronologically to extend the previous tables. Table 1 includes all of the new results for specific h a t s , except those of experiment No. 7,which were discarded as explained above. In experiment No. 8 the solution was too concentrated to be studied at 5°C. Experiment No. 9 was carried out only at 25'C. These results were combined with those previously published, to yield equations for the apparent molal heat capacity as a function of the molality at each temperature. These are given in table 2. The coefficients were determined by the method of leaat squares. The values of As, the observed specific heats of the concentrated solutions, minus those calculated from the equation, are shown in table 1. All of the results at 5' and 26°C.were reproduced by quadratic equations, with average deviatiom of f 6 and f4 X le6 in the specific heat. One deviation of 10 and one of 15 was found a t the lower temperature, while none exceeded 8 a t 26OC. At 40°C. we had found that the results at low concentrations were best fitted by a quadratic equation with a negative initial slope. The
311
MOLAL HEAT CAPACITIES AND VOLUMES
results a t higher concentrations mere decidedly lower than the extrapolation of this curve. When all of the results were included, the best quadratic equation had a positive initial slope. Since the coefficient of m2 has TABLE 1 Specijic heals and appareibt molal heat capacilzes of concentrated solutions of glycolamide at 5 " , 85", and 40°C. NO.
1
AT 5'C.
z"s%
EXF'C.
8
AT 25'C
1 -105Aa~ WCZ-n
lo5A8
AT
WCP.
8
40'C.
1wA8
~ _ _ - 0.87900 +3 38 29 33 440 87899 +3 38 290.88612 &O /O 84891 $4 39 01,0.85720 4-1
______ 9 3.49524 3691 10 3.49524.36960.863031+10 8 4.52826.1124
1-
8
~
~
~
1
QCpl
I-
40 46
1 40 99
TABLE 2 Equations for the apparent molal heat capacity of glycolamide, and related properties, at 5", 86", and 40°C.
+ +
+
+ +
W p l = 26.66 2.20nz - 0.15m* = 35.69 0.72m - 0.03m2 = 39.15 0.30m Cpi = 26.66 4 . 4 m - 0.45m2 35.69 1 . 4 4 ~-~0.09m2 = 39.15 0.Wm cpy c p =~ 0.0396mz - O.OOMm3 = 0.0130m2- 0.0011mS = 0.0054nb2
+
-
TABLE 3 Densities of concentrated aqueous solutions of glycolamide at 26°C. and corresponding apparent and partial molal volumes -___ c
ExPT.NO.
16a
' i 19a 20a
rn
106 Ad
VI
ovz
(CALCULATED)
OBSERVED)
1
2.8518
3.4096
1.05049 1 ,05049
-9 -9
56.50 56.50
56.85
18 .OM1
1
3.4952
4.3697
1 ,06225 1 ,06224
- 12
56.59 56.59
56.98
18.0406
- 11
3.4952
4.3696
1.06222 1 ,06227
- 14
-9
56.59 56.58
56.98
18.0406
4.5282
6.1124
1 ,08085 1.08066
-9 f10
66.73 56.78
57.19
18.0393
b 18s. b
d
only 0.01, we then computed the best linear equation. This reproduced all the results with an average deviation of *S X 10-5,-actually less than the average deviation of k18.given by the best quadratic equation. The linear equation Iisted in table 2 almost exactly reproduces the observations a t the two highest concentrations reported in this article, and a t
312
FRANK 1'. GUCKER, J R . , AND WILLIAM L. FORD
2.6 m. It yields values too high by 19, 12, and 6 X 10-5 at intermediate concentrations, and too low by 16 and 10 in the dilute range. Equations for the partial molal heat capacities of solute and solvent, calculated by the methods of our previous article, are also included in table 2. Table 3 gives all of the new density measurements, except those of experiment No. 17, which were discarded as explained above. We have also included an unpublished measurement, at 2.85 M , made by Dr. Charles E. Moser in this laboratory. These results were combined with our previous ones, by means of the method of least squares, to give the best linear equation for the apparent molal volume as a function of the molarity, c (moles per liter). This equation, which does not differ much from our previous one, reproduces all of our experimental results satisfactorily, with random deviations in the densities. Table 3 includes the deviations for our new results and also the values of the apparent molal volume of glycolamide and the partial molal volumes of solute and solvent, calculated from the equations given in table 4. TABLE 4 Equations for the density and related properties of aqueous solutions of glycolamide at %'C.
+ + +
d = 0.997074 0.019065~- 0 . 0 0 0 1 2 9 ~ ~ .PV2 = 56.168 0.1294~ P2 = 56.168 0.25% 0.00727~2 PI = 18.0691 - 0.002338~2
-
Whenever the apparent molal volume of a solute can be expressed by the equation @Vz = @Vi ac (1)
+
the general equations for the partial molal volumes of solute and solvent (2) are
and
lOOoP! =
[IO,
+a61
(3)
Expanding these fractions by division gives and
v2= +v;+ 2ac - 10-3u@v:c2 - 2 x 10-3uzc8+ - . . . p, = v!(1 - 1 0 - ~ + ~ *10-6a2C4- . . . .)
(4) (5)
STUDY O F UNIVERSAL INDICATORS
313
In dealing with our results for glycolamide, all terms in c higher than the second can be neglected, even for the most concentrated solutions. This must be true in any case where a is small, and leads to a general simplification in calculations. The last two equations of table 4 are just as satisfactory as the more complicated equations of the type given in table 6 of our previous paper (1). We have measured the specific heats of concentrated aqueous solutions of glycolamide up to 3.5 A I at 5“C., and 4.5 M at 25” and 4OoC., and the densities at 25°C. up to 4.5 M . The new results double the range of our previous experiments. We have derived equations which express the apparent and partial molal heat capacity and volume, and the density of solutions of glycolamide over the whole range of concentration. REFERENCES (1) GUCKER, FORD,AND MOSER:J. Phys. Chem. 43, 153 (1939). (2) GUCKER,GAGE,AND AIOSER:J . Am. Chem. SOC.80, 2582 (1938). (3) HEILBRON: Dictionary of Organic Compounds. Oxford University Press, New York (1934).
A SPECTROPHOTOMETRIC STUDY OF UNIVERSAL INDICATORS1 J. T. WOODS WITE M. G. MELLON Departmnt of Chemistry, Purdue University, West Lafayette, Indiana Received March 8, 1940
The term “universal indicator” is usually applied to a solution which contains two or more acid-base indicators and is used to measure acidity over a wide pH range. Such a mixture acts as a single indicator with several stages of ionization, each associated with a change in hue. A number of these mixtures have been proposed by different investigators (1, 4, 6, 8, 9, 12, 14, 15, 16). I n addition, several commercial products, of unannounced composition, are now available. In general, the hues of the solutions follow those of the spectivm in being red in acids and blue or purple in bases. While the general nature of the absorption spectra of indicator solutions has been known for years, few reliable spectrophotometric transmittancy curves have been reported until recently for individual compounds (2, 5 ) , 1 Abstracted from a thesis presented by J . T. Woods to the Graduate School of Purdue University in partial fulfillment of the requirements for the degree of Master of Science, June, 1939.
