THE SOLUBILITY O F UREA I N WATER* B Y LOUIS SHNIDMAN AND ARTHUR A . SUSIER
Introduction Although the solubility of urea in water has been the subject of several researches,'@ the precision of measurement in most cases is poor, and judging from the procedures used, the accuracy of the results may well be questioned. One purpose of the present research was, therefore, to attempt to determine more precisely and accurately the solubility of urea in water. The data are also of some interest from the standpoint of a study of concentrated solutions-their ideality, non-ideality, and the like. Materials I n work of this nature the materials used become of prime importance. Urea from two sources was used. First, the C. P. Grade of Baker's urea which was imported from Germany and made by the Synthetic Ammonia Carbon Dioxide Process4 was used. It is interesting to note that a t the present time apparently no urea is being made in the United States, all being imported from Germany. This sample of urea as received showed a melting point of 132.5', but some solid insoluble material seemed to be present, so the urea was twice re-crystallized from distilled water. During the re-crystallization the temperature was never allowed to exceed 65'. The urea was then carefully dried a t 55'C. in an electric oven for sixteen hours, after which the melting point was run. The melting point of the resulting crystals was 132.7', which agrees exactly with the value of 132.7'recorded in the 1iterature.j Another sample of urea made from Calcium Cyanamid was used. This sample obtained through the courtesy of the Union Carbide Company, was made by them from Calcium Cyanamid in an experimental plant in 1 9 2 j , and showed the following analysis:
Dry Basis
% Total Nitrogen Urea Nitrogen Dicy Nitrogen Guanylurea Nitrogen Cyanamid Nitrogen Ammonia Nitrogen Insoluble Nitrogen Combined H2S04 hloisture as received
43 . g o 40.40
.18 I .85 I8
.87 '
Urea Dicy Guanylurea Cyanamid Ammonia
'
86.62 '27
3.37 '27
I
.06
I3
5.52
7.35
*Communication from the Chemical Laboratory of the University of Rochester.
THE SOLUBILITY OF UREA IN WATER
I233
This sample although impure was re-crystallized carefully four times from distilled water and twice from Baker's C. P. Methanol. During the re-crystallization, the temperature never exceeded 65'. The final crystals from the second methanol re-crystallization were dried in an electric oven a t 52' for 24 hours, The melting point of the resulting crystals was 132.6' which agrees closely with the value recorded in the literature. The melting point was taken as the prime criterion for the purity of the substances used. Thus, two samples of urea from different sources, originally made by different methods, were obtained in a high state of purity.
Procedure and Apparatus The synthetic method of Alexejew6 was employed in making the solubility determinations. This method consisted in heating weighed quantities of solvent and solute in a sealed tube, shaken in a water bath, and noting the temperature at which the solid phase had nearly disappeared. In recent years other investigators7,8*9have found this method to be an accurate and a reliable means for determining the solubility of solids in various solvents. As pointed out by these investigators, care must be taken in attaining true equilibrium conditions at the solubility temperature; this can ordinarily be obtained through slow heating. The apparatus used has been described earlier.g A temperature rise of 0.01' per minute was used in some cases, though in many cases thermostating for a period of time was employed. Sunierg pointed out that with a rate of heating of 0.01' per minute, results well within 0.1' of the true solubility temperature were obtained for naphthalene-aliphatic alcohol systems. The authors feel that this same degree of accuracy would hold for the urea-water system. I n the preparation of the sealed tubes for a run, precautions were taken to insure the presence of small crystals. Other investigators' * . p have shown that the size of the crystal is of importance in attaining true equilibrium conditions. The method was that ordinarily employed and consisted in rapidly heating the tube to a temperature where all the solute dissolved, and then cooling rapidly with vigorous shaking. Thin-wall pyrex tubes of seven millimeters internal diameter and approximately fourteen centimeters long were used. The tubes were cleaned with sulphuric-chromic acid cleaning solution, rinsed with distilled water, and then heated over an open Bunsen burner to dull redness, placed in a desiccator, allowed to cool, and weighed. In these determinations a thermometer certified by the Bureau of Stitndards was employed. The thermometer could be read to *O.OIO with the aid of a magnifying glass. The temperatures recorded should be accurate to *o.oz'. Statements are found in the l1 to the effect that care is necessary in heating solutions of urea because of the danger of decomposition a t higher temperatures. I n order to study this point more clearly, and %tthe same time check the accuracy of the results obtained at higher temperatures, a tube originally showing a corrected solubility temperature of 73.11' was
LOUIS SHNIDMAN AND ARTHUR A. SUNIER
1234
heated to the temperature of boiling water (99') with constant shaking for varying periods of time. The solubility temperature was re-determined after each period. The results are given in Table I, and are plotted in Fig. I,
TABLEI Effect of Heating a t 99' on Solubility Temperature of Urea Hrs. @
Solubility Temperature
99" 0
73.11 73 .os 72.48 71.36 69.67
I
5 IO 20
Difference in degrees
-0.03 -0.63 -1.75
-3.44
As shown above after one hour heating and shaking a t 99', a change of 0.03' was noticed. This change is probably within the degree of accuracy of the method a t that temperature. Hence, i t is concluded that the solubility temperatures in the higher range are accurate and are not affected by any decomposition, because at no time was the tube subjected to its solubility temperature for more than one hour, and then considerably below 99'. If little or no deDl.lLllLllCL Cunoarl% composition took place a t 99' after one hour, it was quite certain that the change FIG.I or decomposition taking place a t temperChange in Solubility Temperature of atures of 73' or below would be negligible Urea on Heating. for periods of one or two hours. Fig. I shows clearly that the rate of decomposition after the first hour and during the next nineteen hours of heating and shaking a t 9 9 O , as indicated by its solubility temperature, is a straight line function. I t is of course quite possible that the presence of the ammonium cyanate produced, affects the solubility temperature to some extent; hence the solubility temperatures do not give a true indication of the amount or rate of decomposition. Walkerlo heated a decinormal urea solution at 100' and found that the transformation of urea into ammonium cyanate had reached equilibrium after one hour. The tube used for the above heating at 99' was approximately a 57.5 normal urea. It may be that the more concentrated solution shows a slower rate of transformation, which would explain the progressive lowering of the solubility temperature during the 20 hour heating period.
1 - 1
I*
THE SOLUBILITY O F UREA I N WATER
I235
Results The results of the various solubility determinations are presented in Tables I1 and 111. Concentrations have been calculated and tabulated on both the mol fraction and weight per cent basis. The data was plotted on a large scale according to the method of Hildebrand and Jenks,I2 as the log Nz vs. rooo/T. The solubilities at rounded temperatures were read off and are given in Table IV.
TABLEI1 Solubility of Urea in Water Urea made by Snythetic NHJ Grams Grams Urea Solvent I . 5140 I . 7891
I .8899 I . 5119
2.3267 I ,8298 2.8030 2.2347 2.1987 2.5702 2.3602
+ Wgt. CO1 Process (re-crystallized from water) % Mol Fract. Solub. Urea
Urea
Temp.
I
51. I O
0.2387 0.2712
18.72 26.83
1.4951 I ,0103
55.37 55.83 59.94 60.87 64. I9 65.39 69.33 70.38 72.59 77.57
0.3498 o ,3618 0.4041 0.4163 0,4428
27.31 35.42 37.36 43.94 46.56 54'77 57.02 61.76
0.5093
73.11
,4486 1.4420
1,4955 I ,0205
,4832 0.9884 0.9250 I
0.9703 0,6823
0.2749 0.3099 0.3182
TABLE I11 Solubility of Urea in Water Urea made from Calcium Cyanamid-re-crystallized from water and methanol Mol Fract. Solub. Grams Grams Wgt. % Solvent Urea Urea Temp. Urea I . 1217
I . 0027
52.80
1.7794 2.0496
I . 5176
53.97 57.51 59.97 62.95 64.31 69 ' 53 7 0 . IO
1.5155
,6631 I ,8291 I , 1159 2,2329 I
1.8579 1,3635 2.4081
I . 5140
,0114 0.9788 I ,0148 0.4888 0.9520 0.7406 I
0.4879 0,7355
71.49 73.64 76.60
0.2513 0,2602 0.2888 0 . 3I 0 2 0.3377 0.3510 0.4065 0,4130 0.4294 0.4561 0.4956
21.59 23.85 30.38 35. I 5 41.11 43.85 54,97 55.88 59.13 63.79 70.49
1236
LOUIS SHNIDMAN AND ARTHUR A. SUNIER
TABLEI V Solubility of Urea in Water a t Rounded Temperatures (expressed in Mol Fractions of Urea) Temp. “C.
Synthetic Process Nz Urea
Cynamid N2 Urea
Mean Nf Urea
20
25
0.2435 0.2642
30
0.2856
0.2447 0.2655 0.2868
0.2441 0 .2649 0,2862
40
0.3308 0,3554
0.3322 0.3569 0.3818 0.4079 0.4349 0.462, 0.4920
0.4069 0,4340 0.4618 0 49=4
45 50
0.3802
55
0.4060 0.4332 0.4609 0,4906
60 65
70
Discussion of Results The results of these determinations were compared with those published by earlier workers. Speyers,’ many years ago, determined the solubility of urea in water. He used Kahlbaum’s urea (no doubt made from cyanamid), re-crystallized from ethyl alcohol, and dried on a steam radiator t o prevent decomposition. Ordinary distilled water was used in these determinations. H e used a thermometer said to be accurate to a tenth of a degree, and kept the solution a t a temperature to within a tenth of a degree for ten minutes, and then analyzed the filtrate by the Kjehldahl distillation method. H e states that the filtration was carried on while the temperature was cooling slightly. Speyers’ results were not very precise, and a study of his method of attaining equilibrium and analyzing his samples leads to conclusion that the results are not very accurate. Krummacher* determined the solubility of urea, re-crystallized from ethyl alcohol, a t the three temperatures, viz. 5 . 5 , 17.1, and 20.92OC. His results at the two higher temperatures are in essential agreement with the results presented in this paper. Comparison at the lower temperature is not possible, except by a considerable extrapolation of the present curve. Pinck and KelleyS more recently determined the solubility of urea in water and found that at higher temperatures there was a marked deviation from the results published by Speyers. Their method consisted in heating from three to four hundred cc. of solution in a water bath in the presence of solid urea a t a temperature a few degrees above that a t which the solubility was determined. When the solution was cooled to the desired temperature, and after being maintained a t this point for about ten minutes, a sample of twenty-five to thirty-five grams was taken. The dissolved urea was determined by the urease method of Fox and Geldard.” These authors used a sample of synthetic urea purified by two re-crystallieations from water. Their procedure in obtaining equilibrium was not of
THE SOLUBILITY OF UREA I N WATER
1237
extreme accuracy. Some of their results agree closely with the authors’ determinations, but some of them, e.g. a t the lower and higher temperatures, show some variations. The results of the above mentioned investigators are compared in graphic form with the authors’ work in Fig. 2. When Speyers’ data are plotted according to the method of Hildebrand and Jenks,l* no smooth or straight line curve results. When Pinck and Kelley’s data are plotted in a similar manner, the curve also is not smooth. This deviation from the straight-line function was undoubtedly partly due to their method of obtaining equilibrium. The present data on the solubility
x 0
.R*C*
& WbbK
-AUTHOR6 ( U R U
B - k r * o R a wucn
LOO
MOL
rRACTlOP4
URU
FIG.2 Solubility of Urea in Water.
of urea made by two independent methods, plotted according to the method of Hildebrand and Jenks, yield a smooth and straight line curve. The results on the two samples of urea are so close that they overlap in many cases. A study of Tables I1 and I11 shows the deviation obtained relative to the solubility of urea and water in two samples of urea made from different sources, and re-crystallized from water and methanol respectively. The fact that urea made by the synthetic method and urea made from calcium cyanamid shows such close agreement in regard to their solubility in water, leads to the conclusion that these determinations are more accurate than those of previously published work. This appears more plausible when it is recalled that the urea made by the synthetic process was re-crystallized from water, while that made from calcium cyanamid was re-crystallized from methanol. Hence, samples of urea made by two independent methods, and re-crystallized from water and methanol respectively, show close agreement as regards their solubility in water. These relations are brought out very clearly in Fig. 2 . It is felt that the foregoing results are accurate to well
1238
LOUIS SHNIDMAN AND ARTHUR A. SUNIER
within + 0.1j o of the true solubility temperature; this figure representing the maximum deviation, whereas most of the determinations deviate much less. The mean deviation in the solubility temperature for the synthetic urea sample was O.IO', while for urea from calcium cyanamid it was 0.07". It is believed that the solubility results obtained with the sample of synthetic urea are nearer the true value, than those of urea from calcium cyanamid; first, because the former showed a slightly higher melting point, and secondly the product before re-crystallization was in a higher state of purity. However, in the final results (see Table IV) a mean value of the solubility of urea a t rounded temperatures is presented. This mean value, representing the average of the solubilities of the respective samples of urea, was read off from the large plot, previously referred to. From this plot the equation of the mean straight line was determined and found to be log,,N = - 609.8 (1/T) 1.468 and is valid over the temperature range 20' to 70' studied. It gives results to within one part per thousand of the values as obtained from the plot. Some nine preliminary determinations by Mr. E. Doell in this laboratory, using urea twice re-crystallized from absolute ethanol, has given results which lead to the same conclusion. The fact that the log N vs. I/T curve is a straight line, leads one to inquire whether or not ideal solutions are encountered in this range of concentration and temperature. When the curve is extrapolated to log N = 0 , the intercept on the temperature axis gives a value of 142' whereas the melting point of pure urea is agreed to be 132.7'. This discrepancy rather definitely points to the fact that the solutions are non-ideal, perhaps yielding a reverse S form of curve discussed by Mortimer.14 No data seems to be available concerning
+
TABLE V Vapor Pressure of Urea Solutions at 60.28' (Perman and Lovett's Data) Mols Urea per Mol H2 0
-
,023I ,0445 ,0690 .os13 . I339 ,1763 .2537 ,3409 ,4160 .4712 .j188
,6482 .8127
Mol. Fract. Urea
WP.)
-
,0226 .0426 .0645 .0752
.I181
. I499
,2024 ,2542 ,2938 ,3203 ,3416 ,3933 ,4483
Observed Vap. Press. (ma.)
151.42 I49 ' 1 146.8 144.7 142.2 136. I 131.3 123.8 118.0 111.7 107 .o
104 . o 95.4 85.5
Mol Fract. (Calc.)
% Diff. or
-
47.6 39.7 45.4 23.5 16.7 12.8 10.9
,0153 ,0305 '0444 ,0609
. I012 ,1329 ,1824
Error
,2207
15.2
,2623 '2934 .3132 ,3700 .4353
I2
.o
9.2 9.1 6.3 3.0
THE SOLUBILITY OF UREA IN WATER
I239
the latent heat of fusion of urea (no doubt because of the decomposition of urea at its melting point); hence a comparison of the experimental and ideal slope of the line is not possible. I t may be said that when the slope, 6d9.8, is multiplied by 4.583,a value of approximately 2800 calories is obtained; if the solutions were ideal, this value would represent the molal latent heat of fusion of urea. At least two sets of data are to be found in the literature dealing with the vapor pressures of urea solutions. That of Perman and LovettIs covers quite a range of concentrations at several temperatures. Since these authors did not present any calculations of deviations from Raoult's law (Bancroft'c in a recent article mentions briefly the variation of these data from Raoult's law) calculations have been made using their data. The pertinent data and the results of the calculations at only one temperature-60.28~ would appear to be of sufficient interest to be included in this paper. In Table V columns one and three are taken directly from the paper of Perman and Lovett. I n column two concentrations are recorded on the mol fraction basis. Column four gives the mol fractions calculated by substituting the proper values in the well known equation (po - p)/po = x of Raoult. The difference between the experimental and calculated values of the mol fraction is seen to vary from nearly j o per cent to 3 per cent. Similar calculations using Perman and Lovett's data a t 40.02' where the mol fraction of urea ranged from 0.059 j to 0.402I showed a deviation from Raoult's Law decreasing from nearly 43 per cent to 8 per cent; the data a t 80.10" where the mol fraction of urea ranged from 0.043 I to 0.5547 showed deviations from Raoult's Law decreasing from nearly 180per cent to I per cent. One peculiar point is worthy of mention, viz. the deviations are greatest in the most dilute solutions in all cases. This may in part be explained by the relatively small pressure differences encountered a t the lower concentrations, but it is hardly possible that this explanation holds throughout the range of concentrations studied. It would appear that with solutions of most other substances studied, the deviations from Raoult's Law are greater in the more concentrated solutions. An example of this is found in the work of Berkeley, Hartley, and Burton17 on the vapor pressures of sugar solutions. These workers have found that on increasing the mol fraction of sugar from 0.0175 to 0.1025, the deviation from Raoult's Law increases from I O per cent to 37 per cent. Such results, along with others which are not presented a t this time, lead one to the view that any attempt to express quantitatively the deviations from Raoult's Law for urea solutions should be postponed until further data are accumulated. The data of R. Fricke18 on the vapor pressure of urea solutions were obtained a t temperatures near zero degrees centigrade and are confined to only three concentrations a t two different temperatures, and thus do not materially assist in deciding the question just touched upon. In connection with the work of Perman and Lovett, it may be well to mention further that some of the more concentrated solutions were super-saturated, a fact apparently known to the authors when the work was under way.
I240
LOUIS SHNIDhlAN AND ARTHUR A. SUNIER
As a result of their study of the heat of solution and heat of dilution curves of urea solutions, Perman and LovettI5have suggested a new method of determining solubilities, viz. “If a number of points of the Heat of Solution curve were determined and afterwards a number on the Heat of Dilution curve, it is obvious that these curves would intersect at the saturation point and would indicate the solubility of the substance.” The results so obtained for urea solutions differ by about 1.4 per cent by weight at j o o . It may be said that Perman and Lovett did not claim a high precision for the solubility determined in this way. Summary Samples of urea from two different sources have been carefully purified. Some twenty-two determinations of the solubility of urea in water have been made using the synthetic method in the temperature interval 20’ to 70’; the precision of measurement in these runs is much higher than any previously published; it is believed that this is true of the accuracy also. The data may be accurately represented by the following equation, log,,N = - 609.8 (I/T) 1.468. 3 . Preliminary data are presented showing the effect of varying periods of heating on the solubility temperature. I.
2.
+
Literature cited Speyers: Am. J. Sci., (4) 14, 294 (1902). Krummacher: 2. Bioi., 46,302 (1905). 3 Pinck and Kelley: J. Am. Chem. SOC., 47, 2170 (1925). Bosch and Merser: U. S. Patent-1,429,483 (Sept. 19, 1922). 6 Inter. Crit. Tables, 1, 177. 6 Wied. Ann., 28, 305 (1886). 7 Ward: J. Phys. Chem., 30, 1316 (1926). 8 Sunier and Rosenblum: J. Phys. Chem., 32, 1049 (1928). 9 Sunier: J. Phys. Chem., 34, 2582 (1930). 10 Walker: J. Chern. SOC.,67, 746 (1895). 11 Thorpe: “Dictionary of Applied Chemistry,” 7, 271 (1927). ’ZHildebrand and Jenks: J. Am. Chem. SOC.,42,2180 (1920). 13 Fox and Geldard: Ind. Eng. Chem., 15, 743 (1923). 14 Mortimer: J. Am. Chem. SOC., 44, 1416(1922); 45,633 (1923). 15 Perman and Lovett: Trans. Faraday SOC.,22, I (1926). 18 Bancroft: J. Phys. Chem., 35, 3160 (1931). “Phil. Trans. Roy. SOC.,218A, 295 (1919). 1s R. Fricke: 2. Elektrochemie, 35, 631-40 (1929). I
