AN IIWESTIGATIOT\; O F T H E DEGREES O F HYDRATION O F T H E ALKYL AMINES 11; AQL-EOVS SOLUTION BY
\vILLIARI
CAMPBELL SOJIERVILLE
Introduction The state of combination of ammonia in aqueous solution has been the subject of much controversy in recent years, as may be seen from a series of papers by R. M. Caven, T. S.Moore, and A. E. C. Smith.’ The point on which these writers disagreed was whether a hydroxide KH,OH actually existed, which ionised directly into SH,’and OH-, or whether the observed ionisation was due t o the fundamental reactions
+
H 2 0 e - H ~ OH-
NH3
+H+SXHl+
I t has been shown that a hydrate of ammonia, SH,,&O, does exist at low temperatures by Rupert,? and Moore and \\7inmil13 obtained results by means of distribution coefficient determinations which indicated that ammonia is hydrated in solution. I t was shown by Caven‘ however that the ammonium ion is similar in some of its properties to the ions of the alkali metals. On these grounds “,OH would be expected to be a strong base, whereas it is a very weak one. Caven suggests therefore that there exists in solution a hydrate, say NH3,nH20 which is not directly ionisable, and that the ionising process might be represented by X H ~ , E H ~f OH+FfXH4+,nH*O Similar equations might be applied t o the alkyl amines. Thus the hydrate formation on Caven’s theory takes place as a side reaction, and the extent to which the amines are hydrated may have little connection with their relative tendency t o give positive ions in solution. On Moore’s hypothesis, while we cannot postulate definitely that the order of hydration will parallel the order in which the amines are ionised, there is a greater possibility of ita doing so, since combination with water is a preliminary to ionisation. The results obtained by Moore and \Tinmill for the degrees of hydration of the various amines show very marked irregularity. The pres.ent research was undertaken in an attempt to obtain results from various physical properties which might be correlated to establish the degrees of hydration of the amines. The first line of attack was through the l
J. SOC. Chem. Ind., 42 (1923).
* J. Am.
Chem. Soc., 31, 866 (1909); 3 2 , 148 J. Chem. SOC., 101, 163j (1912). e.g. J. Chem. Soc., 121, 1406 (1922’1.
(1910)
DEGREES O F HTDRATION O F THE ALKYL AMINES
2413
freezing point curves. These have been examined by Pickering.' It was felt advisable to check the compositions of the hydrates he obtained, and also t o determine more accurately the part of each curve where the crystallising substance is ice. For this part of the curve, deviations are apparent from the ideul curve given by Washburn,*
+
ATr = 103.20(2 0 . 4 2 8 ~ ~ ) where ATr = depression of freezing point z = mole fraction of solute.
If the solute combines with water, less water will be available in the free state. Thus the effective concentration of the solute, and hence the freezing point depression, will be increased. Conductivity measurements (and with them density and viscosity measurements), were undertaken to correct for the ionisation of the amines in dilute solution. It was also thought it might be possible to deduce from the conductivity results directly, approximate values for the degree of hydration of each amine a t o°C and a t 25°C. This is treated fully in the discussion. Conductivities have been determined by Bredig3 at zs0, Bruni and Sandonnini' a t 2 jo,Moore and Winmill$ a t IS', 2 5 O , and 32.35'. None of these used solutions more concentrated than N/8. Further it was necessary to determine the degrees of ionisation a t o°C, as that temperature corresponds most nearly to the temperatures encountered in the freezing point determinations. Densities and viscosities of solutions of the methylamines have been determined at 2 j" by Kanitq6 but it was thought advisable t,o check his values. Materials Monomethylamine. The hydrochloride was freed from ammonium chloride by adding to the solution twice as much sodium hydroxide as was equivalent t o the ammonia impurity (6Tc)and boiling until no more base came off. More caustic soda was added, and the amine distilled into water, precautions being taken to exclude carbon dioxide. On testing the amine as hydrochloride by titrating a weighed amount with silver nitrate solution, the percentage of ammonia was found to be zero. K h e n a sample was required to be used in very concentrated solution, the amine solution was boiled in a flask under a reflux condenser, the amine being passed down a spiral immersed in a cooling mixture of ether and carbon dioxide snow. It was collected directly in the freezing point tube, only the amount required in the experiment being distilled over. J. Chem. Soc., 63, 181 (1893). "Principles of Physical Chemistry," equation 54 (1915) 2. physik. Chem., 13, 289 (1894). ' 2. Electrochemie, 16, 223 (1910). loc. cit. 2. physik. Chem., 22, 236 (1897).
WILLIAM CAMPBELL SOUERVILLE
2414
D i and t r i methylamines, mono, d i and tri ethylamines. These were prepared by the usual methods,' or bought from Hopkins and Williams, and purified. Standard solutions of hydrochloric acid and sodium hydroxide. These were standardised using pure dry anhydrous sodium carbonate, or ealcspar.
Determination of Concentration The mole fraction of amine in a solution was found by weighing a flask containing a pipetted volume of standard hydrochloric acid ( N or N / I O according to the concentration of the amine solution), running in a quantity of the solution, and reweighing the flask. The quantities of acid and of amine were taken so as to leave a slight excess of acid. This excess was titrated with standard sodium hydroxide. Provided the volume of acid was at least 2 5 cc. the titration was accurate to one in one thousand. Knowing the density of the solution, the normality could also be calculated. Methyl orange was used as indicator throughout.
Densities of Amiie Solutions The pyknometers were of the usual Ostwald type, of 1.1-1.8cc. capacity. The accuracy of the results was of the order of 3 in 10,000. Density data a t o°C xa = mole fraction of amine in solution DE = relative density of solution a t o°C. M e N H 2 . .. . . . . .x3 DE Me2".
. . . . . . .x8
,05218
,02793
,01503
,00789
,00423
,9782
,9871
,9930
'
9960
,9980
,03817
,02426
.01z99
,00829
,00543
,9852
.991o
'
9942
,9962
DE Mes>-. . . . . . . . .xa
EtNHz..
,02385 ,9804
.02318 ,980~
,01008 ,9902
,0041j
D:
. . . . . . .za
,02984
.0142j
.00724
,00362
.9851
,9914
'9959
'9983
,01340
,00684
.00380
.00206
,991
,994i
'9974
,9989
.02530
.01258
.007jz
,00397
.oozo8
.98I O
.992O
.9944
.9974
'
Dg EtzNH.. . . . . . ..xa
D," E t & . . . . . . . . . . xa
D::
Ber., 12, 523, 38, 882; J. Chem. Soc., 109, I j 4 .
'9963
9989
DEGREES OF HTDRATIOS OF THE ALKYL A N I S E S
241,;
Density data at 25OC M e N H n .. . . . . . . r a
Di: M e 2 S H . .. . . . . . r a
D:.:
.ozj92
,01352
.006j2
,00344
,987;
,9932
,9970
,9987
.018jo
,00837 ,9939
,00426
,00180 ,9989
,9871
,9969
Me3??. . . . . . . . .x D::
,04623 ,961~
. 01j16 ,984;
,00864 ,991;
E t S H z . , . . . . . . .ra
,01837 ,9890
.0079j ,9947
,00348 ,9971
.OISII
,9986
,00609 .995O
,00293 ,9980
,00998 ,9912
.00;18 ,9951
,00261
D:: E t z T H . . . . . . . . .x,
Dl: E t 3 S . . . . . . . . . .x3.
D::
.OOI~Z
,9990
,9980
Viscosities of Amine Solutions The method used was that given by Findlay,' using a viscometer of the type shown in Fig. I , taking 2 cc. of liquid. The thermostat a t 2 j c was a large glass jar, stirred and regulated to 2 j . 0 0 =t o.01"C. At ocC a large vacuum flask filled with distilled water, and ice prepared from distilled water, was used. Accuracy obtainable. The stop watch was graduated in fifths of a second, and read to 0.1sec. After some practice, and by repeating the timing until concordant results were obtained, the error in starting and stopping the watch was about 0.1 sec., equivalent t o an error of slightly less than 2 in 1000,since the time of outflow a t 25' was I niin. 7 sec. A t o°C this error was less. Some deviation from Poiseuille's law due to the construction of the viscometer was found when a 2 0 7 ~solution of sucrose Frc, I was employed (compare Washburn and MacInnes2). It was Viscosity noticed a t oo that the surface tension of pure water differed Apparatus so much from that of a solution that an error was being introduced in the time taken by the liquid surface to sink down the part of the capillary between the bulb and the mark B. A new mark was accordingly etched a t C. It was now found that the results obtained with the 2 0 7 ~sucrose solution obeyed Poiseuille's law to within I in 1000, less than the error due to the stop watch.
' "Practical Physical Chemistry," 4th ed., p. 3j. J. Am. Chem. SOC., 33, 1692 ( 1 9 1 1 ) .
2416
TPILLIAhl CAMPBELL SOMERVILLE
Viscosity data at 0°C.
x.
=
mole fraction of amine
q / a w = viscosity of solution over viscosity of pure
water. ,04549 1.525
...
.0225j
1.239
. 0 4 2 j ~
I . j34
.00900
1.095
,003667
.001800
1.019
1.040
,01835 ,00962 ,004647 .002247 1.373 1.186 1.084 1,039
,02395 ,00971 ,00693 ,003747 ,001771 1.884 1.294 1.204 r . 1 0 8 1.048 ,02269 .00836 ,002958 ,001538 1.427 1.138 1,043 1 . 0 2 1
...
,01888 .oogg3 ,003854 .001701 1.818 1.374 1.133 1.055
2
,02530 , 0 1 1 1 4 ,005395 , 0 0 1 5 7 2 696 1.601 1.248 1.064
Viscosity data at 26°C MeNH z.......x, ,02058 q/qw
hle~?;H... . . . .x8 q/aw
iMesr\f.........x, q/qw
EtXH2.. , , , . , .x8 o/qw
EtZKH.. . . . . , ,xa q/qw
E t & . . . . . . . . .x. q/qw
1.153
,01219 .oog4j ,00465 1.082 1.061 1.030
.00202o
1.013
,02617 . ~ 1 3 1 3 .0041;1 ,001576 1,365 I 165 1.046 1.018 ,04672 ,02537 .01oj9 ,004685. . 0 0 2 2 2 5 1.523 1.200 1.081 1.034
2.074
.02318 ,01114 .00503o ,002574 1.128 1.oj1 1 . 0 2 2
1.281
,02654 1.727
,01171
1.278
.0009j1 1.016
,001064 1.00;
,00590 ,00264 .00140 1.049 1.023
1.126
,01349 ,005299 ,002347 ,001087 1.402 1.141 1.053 r.015
Conductivities of Amine Solutions Using a valve oscillator, headphones, and a three metre Kohlrausch slide wire bridge (Leeds and Northrup), an accuracy of 4 in IO,OOO in the ratio of one side of the bridge to the other was obtained. A variable capacity was connected in parallel with the resistance box, bahncing the capacity of the cell, giving a silent sound minimum. Platinum black electrodes were used. It is well known that the amines are catalytically decomposed by platinum black, but before each determination, the solution to be examined, a pipette,
DEGREES O F HYDRATION O F THE ALKYL AMISES
2417
and the cell (previously dried with steam), were alloweti to attain thi. thermostat temperature. Sonic solution m s then transferred to the cc.11, and the conductivity determined niter four minutes. On leaving the cell another four minutes, only w r y sinall changes of conductivity were recorded, so that no error was being introduced. The cell constnnts were det~errnined :tftcr each run, and rcmainctl unchanged throughout. Good distilled water haring a conductivity ncrer greater than 2 . 0 X IO& mhos. a t 25' was employed. .Is no high dilutions were being dealt with, this was satisfactory. Themusticfa. .It 2 5 ' a large copper hath was used, protected by asbtstos o.oI'('. ht 0°C a large Ilewnr sheets, stirred and regulated to 2j.00 vcwel T ~ eniployccl S as in the viscwity J-iork. The eyiiirnicnf eondudii,itl'cs of the n ~ t ~ l ' n ents irt-fiJtite dilittioti. Tlle mohilitic:: of the nminc ions hnve been determined accurately at 2 io by X o o r e and ll-iniiiilljlwho also found temperature coefficients for these mobilities. In thc following table
-1z6 = ralue at =
.lo
=
25' obtained by Moore and Kinmill value at oo calculnted from A 2 5 using the temperature coefficicnts given value at oo calculated from . l Z 5 by means of the equation givcn by Iiohlrausch' &I, dT19 = 0.01341 0.640/A18- 4.94/(A11s)?
+
Ion JIeSHBlIe2SH2+ lIe3SH+ EtSH3EtJH2EtBSH+
'
.I25
A '0
A0
60.6
30,8
30.4
'- 5 9
26.0
26.0
'
-
3
49.5
22.8
23.4
47. T
19.6
22.3
38.4 33.7
15.8
1L9
I j . 1
14.9
The values finally adopted at oo were those given by the Iiohlrausch equation. The values for the mobility of the OH- ion n-ere taken to be 1 9 4 at 2 jo and 118 at oo as in the Internat,ional Critical Tables. Corrections. 1 correction was applied for the conductivity of the amine carbonate formed by the C'On in the water used. The data for the solubility of carbon dioxide in water were taken from Iienda11.3 As shown by Moore and lVinmill,4 practically all the carbon dioxide originally present in the water will be in the form of carbonate ion. At the concentrations considered here, it can be shown that this may be assumed without introducing error. The mobility of the carbonate ion was taken from the International Critical Tables.
' loc. C l t 2. Elektrochernie, 14, 125 (1508'. J. Am. Chem. SOC., 38, 2460 (1516) IOC. C l t .
2418
WILLIAM CAMPBELL SOMERVILLE
Amine
Conductivity correction for amine carbonate a t z jo at o'
3 . o X -"mhos.
Monomethylamine Dimethylamine Trimethylamine Monoethylamine Diethylamine Triethylamine
2.7
3 . 3 X ro+mhos. 3.0 2.9
2.6
2.8
2.4 2.3
2 . 5
2.8
2.4
The corrected conductivity of a solution was given by multiplying the experimental conductivity by q/vw, adding the conductivity of carbon dioxide in pure water, and subtracting the above correction for the amine carbonate, and the conductivity of the water. A correction was also applied to the experimental normality for the amine withdrawn in forming amine carbonate. This was equal to 2 . 8 X I O + moles /litre a t 2 j o l and to 5.88X IO-+ molesllitre a t 0'.
Conductivity data.
mole fraction of amine normality of amine a = degree of ionisation K = ionisation constant of the amine. 2. =
N
iMonomethylamine Xa
,04549 .02255
.00900 .003667 .001800
=
2 So
O0
x
a
2.398 1.219 0.4940 0.2026 0.0997
,00965
K X
104
2.25 3.28 ,02797 3.97 ,01625
,04406
4.11
,06224 4 . 1 2
0.06172
,07867
0.03086 0.01 543
.1o79
4.14 4.02
.1464
3.56
Xa
N 0.8914
,01643 0.7250 :o1331 .00999 0.5467 . 0 0 5 7 ; 5 0.3164 ,002678 0.1479 0.0997 0.04986 0.02493
4
a
K X I O
.0210o
4.02
,02380 ,02793 ,03752 ,05498
4.21 4.39 4.63 4.73
,0667;
4.76
,0924 ,1269
4.69 4.63
Dimethy lamine ,04255
2.168
,01835 .00962 ,004647 .ooz247
0.977 0.5231
,00987 2.13 ,01899 3.59 ,02789 4 . 1 8
0.2555
,04082 4 . 4 4
0.1240
,05863 4.52 ,06200 4 . 5 4
0 .I
108
0.05540 0.02770
,08602 ,1176
4.48
4.33
,01910
1.014
0.4639
,02160 ,03537
4.83
,00854
,004235
0.2322
,05080
6.31 6.45
,001758
0.1196 0.0970 0.05980 0.02990 0.01495
,07080
6.02
,0983 , I342
6.46 6.41 6.21
.1825
6.08
,07834
DEGREES OF HYDRATIOS OF THE ALKYL AMINES
Trimethylamine .0239j 1.235 ,02395 1.228 ,00971 o . j 2 2 1 .00693 0.3766 ,003747 0.2056 .ooij;r 0.0980 0,08268 0,04134 0.02067 Monoethylamine Za N ,02269 1 . 2 0 4 ,00836 c ) . 4560 ,002958 0.1632 0,09173 0.04586 0.02620 0 .
or263
Diethylamine ,01888 0,978 ,00993 0.5308 ,003854 0.2109 . O O I ~ O I 0.0938 0.05142 0.02
-5 j1
0.01286
Triethylamine .02j30 1.233 ,011xq 0.5834 .ooj39j 0.2910 ,001jj2 0.0866 0.06910 0.03455 0.01728
,003jj3 0 . I i j
,0037jj 0.176 .006602 0.229 ,00796 0 . 2 4 0 ,01111
1.093 ,008jI 0.4690 ,004256 0.2323 .0022j2 0 1237
,02112
. ooj942 o 388
,01058 0.j31 . O I j69 0.~j81 . 0 2 2 0j o 61j 0 , 0 5 4 1 I ,03375 0 637 0.0~706 ,04763 0 644 0.01353 ,06630 0 635 o 631 0.0067j ,0922
0 . 2 j ;
,01622 0.262 . 0 1 j j 8 0.266 ,02492 0.263 ,03482 0.259
2 50
O0
a
2119
K X
IOI
,01709 3.58 ,03060 4.40 ,05227
4.70
.06930 ,0967 ,1234 ,1706
4.73 4.74 4.54
4.41
N ,01837 o 978 ,00795 0.4325 ,003476 0.1910 ,001522 o 0839 0.08196 0.04098 0.02049 Za
a
K
x
IO*
.019jo 3.87 03228 4.65 ,04939 4.92 .oi4oj 4.97 ,07501 4.98 ,1038 4.92 ,1409 4.73
0.7902 .03255 8.65 , 0 5 4 4 1 10.32 .006091 0.3292 .00~924 0.1601 ,07880 IO 79 0.0j1j2 ,1167 1 1 oj 0.02586 ,1600 IO 92 0,01793 ,2162 I O 68 0.0089; ,2860 IO 01
,02702
7.34 ,03860 8.23 ,06225 8 . 7 1 ,0917 8.67 ,122j 8.78 ,1663 8.51 ,2201 7.95
.OISII
,01233 1.90 ,02160 2 . 7 8
,00998
,02801
.00j180
,04166 5 . 0 4 ,02990 5 . 4 4 .0823j 5 . 5 7
,03137 2.95
3.05 ,06443 3.06 ,08969 3.05 ,1231 2.98 .05;62
0.5231 0.~j8; .002612 0.1426 0.07534 0.03767
,1138
.1jj9 0.00942 ,2090
0,01884
4.22
5.50 5.42
5.19
The observed values of the ionisation constant fall away as the concentration decreases, after attaining a constant value. This cannot be explained by the presence of traces of ammonia, since these would h a r e caused a low value of K in concentrated solution, rising on dilution. Absorption of carbon dioxide previous to the experiment would have given a diminishing K, but the initial value would have been very high. A test was carried out with
TTILLISU CAMPBELL SOMERVILLE
2420
dimethylamine, runs being done with a solution which had been kept some months in a bottle from which portions were withdrawn from time to time, and also with a freshly distilled solution. S o change in K was found. The most likely explanation is that the amine is adsorbed on the platinum black of the electrodes. The effect on the conductivity of a concentrated solution would be negligible, but would increase as the solution was diluted. I n the preceding tables, values are given down to S I 6 4 solution only, as with still more dilute solutions the results become too uncertain. Freezing Point Diagrams of the Amines and Water
Section 1 . T h e crystallising substance i s a hydrate. The apparatus used is shown in Fig. 2 . The tube containing the solution was cooled in an air jacket surrounded b y a mixture of ether and carbon dioxide snow, not shown in the draiying, and inoculated with a little previously frozen hydrate if possible. If not, it was further cooled and rubbed, until some solid separated. It was ~$;2$ then placed as shown in the figure. Air was bubbled through the freezing mixture, the rate being adjusted by the screw clip so that the solution warmed up very gradually. Hand stirring was employed, and the temperature kept under observation. Around the maxima on the curves, t,he freezing point could be determined most conveniently by noting the temperature a t which a sudden change in the value of 6T:/6t (rat'e of change of temperature with time) occurred, since this indicated that practically all the hydrate had melted which had separated dr o~ out. At points more remote from the maxima however, ??!!PT+ :o2sm* the change in composition in the liquid phase produced by the separation of the hydrate caused this value to be inaccurate. The method used in such cases, was to take U the temperature a t which the last traces of the solid dlsFIG.2 appeared. This was not generally easy, as Pickering Freezing-point found also, solutions of amines a t such concentrations Apparatus and a t such temperatures being very viscous, while the hydrates are often gelatinous in character. Section d . The crystallising substance i s ice. The same apparatus was used as in section I . The method however was different. Several were tried and the following was adopted. The solution was cooled below the freezing point as before, inoculated with a little hoar-frost, and allowed to warm up gradually, observing the temperature every quarter or half minute. It was found that the temperature rose slowly a t first. Then an increase in the rate of rise occurred, small a t first, followed by a much more rapid increase. The point a t which this second acceleration of the rise came was quite definite, was repeatable to 0.08' even when the freezing point was a low as -31OC, and was found to agree
1
DEGREES O F HYDRATION O F THE ALKYL hXINES
2121
well with the correct freezing point when a few determinations were made using sodium chloride solutions made up in accordance with data given in the Int,ernational Critical Tables. Thermometers. Three thermometers were w e d in sections I and 2 , a toluene one graduated in degrees, reading down t o - I jo', and two mercury thermometers. The first of these read from o°C to -40°C in tenths of a degree, and was used whenever possible. The second vas graduated in twentieths of a degree, from + I j"C to - j"C. It was used only for part of the trimethylamine curve. A11 three were standardised at o°C, and the first t w o at the melting point of mercury. Each solution was n u d e up by diluting the previous one. The concentration v a s determined after each freezing point had been taken.
Re5mltssobtained in scctions 1 cind 2 .Ifononicth!/lanii,.ie ra = mole fraction of amine Ti = freezing point of Polution ("Cj ,0312
,0437
)>--Tij . j 9
4.94
j's.
,0583 6.96
. o ; I ~ ,0818
10.6;
8.97
, 0 9 7 4 ,1056 ,1198 ,1379 13.58 1 j . 3 6 18.69 23.68
1 'i:?:
,1623 ice ,2040 ,2387 . 30;s ,3732 MeSH2,3H20 \,-Ti 31.61 cryst. 39.8 3 i . 8 40.5 4 j . 6 cryst. The w l u e found by Piekering' for the melting point of the hydrate was - 36'. ).r.
IXmd h u h n z n e j.r,
,0473 .0;8j 5 . 6 1 ; jo
,0352 3.94
\,-TI
.iogi> ,1186
,j.~:*
,120;
\*-TI 16.98 1 6 . 5 3 16.j 9
. 0 6 jj 0.14
,0834 ,084-t
. 1 0 0 2 ice 1 ; . j 5 cryst.
12.62
12.90
,1295
,1336
16.49
1 6 . 7 ~ 16.82
. i 3 ( > j ,1153 ,1644 , 1 7 8 4 17.18 18.50 2 0 . 0 9
J.rtL . r 8 6 ~ . z o j j . 2 ; j 9 hIenSH,7FIP0 \-T3 2 1 . 0 4 37 '$4 cryst. 1IPlting point of hydrate, Pickering - 16.6'
Tr i ti7 ct h !/lani in e
J
X u 11
, 0 2 3 8 ice
*A
'(i-~,-z.j~
,2039
f:{
- 2 . 0
11
cryst. 11 ,2663 -9.4
1 ,3384 +a.
I
.ojoo
,0681
,0918
. 1249
,1336
,1631
2 . 8
4.8
5.1
4.8
4.1
1.7
, 3 2 6 6 , 4 0 7 0 hydrate - 1 7 . 3 -33 cryst.
Run -3.
fT: {
,0588 3.77
.0829
,0877
j.09
j.16
j.27
,0991 j.18
,1139 ,1713 hydrate 4.78 1.69 cryst.
,0706 4.74
.08j1
,0906
,0950
,1096
, 1 2 8 ; hydrate
5.13
j.23
5.39
4.97
,0902
Ritn 3. {?Ti
_______
' loc. cit.
4.30
cryst.
2422
WILLIAM CAMPBELL SOMERVILLE
Run 4.
. I 23 5' hydrate
,0639
,0808
,0867
,0882
,0908
,0949
,0790 5.94
,0829 6.03
,0889 6.06
,0957 6.06
,0975 6.04
.oggo hydrate 6 . 0 3 cryst.
Run 5. (4Ti
Runs I , 2 , 3 , 4, were done using the material first prepared. Run 5 was using material obtained from Hopkins and Williams. This sample was used in the density, viscosity and conductivity determinations. The results from runs 2 , 3, 4, and 5 , when plotted, show that the formula of the hydrate is M e 3 N , ~ o H 2 0and , not M e 3 h T , I ~ H ZasOfound by Pickering.
FIG.3 The freezing-point curves of the three methylamines. ( I ) Monomethylamine, (2) Dimethylamine, (3) Trimethylamine, (4) Ideal.
2423
DEGREES OF HYDRATION OF THE ALKYL AMINES
Monoethylamine ,0153 ,0299 {T :i 1.58 3.27 5,
-TI
,0440 5.12
.086j .0870 ice 12.64 1 2 . 7 0 cryst.
,0533 6.44
,0604 7.52
,0687 ,0761 ,0808 , 0 8 5 5 9.04 10.36 11.42 12.49
,0907 ,0931 .IO74 13.30 12.74 10.33
,1423 7.83
,1583 7.73
,1671 7.85
,1862 ,2192 ,2512 ,2787 , 2 8 5 1 zF,t?LTHZ,IIH20 -Ti 8 . 5 4 10.38 13.02 17.5 18.5 cryst. Melting point of hydrate, Pickering, - 7. j o 5%
Diethylamine ,0071
.?&
-Ti xa -Ti
0.68
,0147 1.48
,0574 ,0589 8 . 4 4 8.99
i
xp ,1367 -Ti 7 . 5 8
,0162 .ozo8 1.69 2.19
,0278
.0614 ice 9.51 cryst.
,0653 9.12
1
3.15
,0309 3.54
,0368 4.52
.o42j 5.49
,0490 6.77
,0700
,0816 7.73
,0933 7.44
,1076 7.32
8.43
,1789 ,2296 ,2365 EtzNH,8HzO 9.14 11.83 12.33 cryst.
Indications of a hydrate unsuspected by Pickering were found in more concentrated solutions, the following points being determined,
{T :i
,2450 9.7
,2452 10.0
,2460 9.8
,2480 10.1
,2677 I4
,2838 ,2858 ,2931 ,3503 15 16 17 26
This hydrate was so gelatinous as t o make the points obtained indefinite except for those a t the maximum of the curve. The nearest simple composition for the hydrate is E t z " H ~ H 2 0 . The following values were also determined, .41j 3 ,4522 ,4954 ,5827 ,6318 ,6335 ,6855 ,866 25.2 -TI 29.6 2 7 . 0 21.76 2 0 . 7 5 2 0 . 8 7 2 0 . 4 5 24.9
za
5, ,964 zEtzNH,HZO I . ooo pure amine 50.0 cryst. {--Ti 38 cryst. Melting point of EtzNH,8HzO, Pickering, - 7 . 3 -6.5 1 Triethylamine xa ,0083 .0143 . 0 2 7 0 .0354 ,0409 ,0470 -Tr 0 . 8 5 1.48 2.94 3.42 3.65 3.81
,0542 3.95
,0679 4.26
.88r 30
,0855
4.64
The peculiar character of this curve is accounted for since the amine and water become only partially miscible a t temperatures above I 8' (Rothmundl). J. Am. Chem. Soc., 38, 14p (1916).
2424
XILLIAM CAMPBELL SOMERVILLE
Two substance5 whicli are only part'ially miscible tend t o give 1:trge positive deviations from Itaoul?'~IAJY(Hildebrandl). At very low concentrations no immisiciLility occuw, and the curve mag be normal.
FIG.4 The freezing-point curves of the three ethvlamines. : I ! .\Ionoethylarninc. ( 2 ) Diethylarnine, ' 3 ) Triefhylsinine, (4 Ideal. ~
Section 3. Deterniincitions using (: Reckmc~nrithernorneter. The part.%of ' to - 4' were required more accurately, the freezing point diagrams froin 0 since in that region the frrezing points of all solutions are very near t o the ideul values, and the differences which came up for consideration are relatively small. On this account the method had to be one in which the conditions could be repeated. That finally adopted was to supercool the tube and its contents about a degree below the freezing point, inocuhte with a little hoor frost, and place in an air jacket which in turn was jacketed with ice and water. The temperature was read every quarter, half, or whole minutr, according to the rate of rise. Regular mechanical stirring by an cllectric motor was wed. For a space of at least five minutes, if the supercooling had been sufficient, these readings varied by equal increments. dfter some time the rate increased, a little a t first, and then more rapidly. The teniperature taken with each solution was that a t which the rate of rise of temperature became double its original Yalue. If there wai a choice of two readings, the average was taken. The absolute values of the freezing points were found by preparing a scrips of potassium chloride solutions from data in the International (-'riticnl Tables, and standardising the thermometer and the method thereby. The freezing points of these solut,ions were repeatable t(J 0.01'at -+', the error decreasing to about 0.003" at -1'. 1
J. rim. Chem. Soc., 39,
2303 (1917)
DEGREES O F HTDRATIOS O F THE ALKYL AMINES
2425
K i t h solutions giving a freezing point depression of less than one degree, the temperature rose very slonly. I n such cases the inoculated solution was stirred for two minutes, the temperature was read, a ~ i da sample pipetted out from nhich the mole fraction could be determined I n such cases only the zero correction \\as applied
Thcrtnornelcr The thermometer uwd mas of the usual Beckmann type, .c+itha range of six dppreee, graduated in hundredths of a degree, and read to o 001' b> means of a small lens I t n a y kept a t oo for a meek before use, and not d 1 o v . d t o warni up throughout the experiments S t e m correction Variations in room temperature mere allowed for by taking the Aero after each run The alteration in the zero from an arbitrary standard value g a i e the stern correction for the whole length of stem exposed, from x5hich the coirection for any fraction of the stem could be obtained. arbitrary %:due taken 71 as that found .c\ hen the potassium chloride solutions n ere investigated. Results obtainud in section 3. rc,
=
ATr
=
2,
=
mole fraction of amine corrected for ionisation freezing point depression ("C) mole fraction calculated from the form of the ideal equation z,= o . o 0 9 6 9 o ( A ' ~-~ o.oo4zj.AT;) (Kashburnl)
Dimetiqla T u ine
I\lo~ionaetilylarriine ATr
xt
ti
0 . 0 0 3 4 . ~ o 346 0 . 0 0 3 3 4 .oo.qoo 0.40j ,00393
.oa6jj o.joo .006;6 , 0 1 0 4 4 1.o;6 .01038
,01138 1 . 1 7 1 .o1129 ,01390 1.436 ,01383 , 0 1 4 5 1 1 . 5 0 2 ,01446 .01;,58
I
838
.orjh;
xt -xi
loc. cit
ATi
xi
.00006
0.235 0 . 0 0 2 2 8 0.342 ,00331 ,00587 0.602 ,00581 ,00730 .ooj34 0 . 7 j ;
.00009
,00926 0 . 9 ; ;
.0092z
.OOOO;
.01239 1.287
,01240
.OOOO~
. O I Z ~1 ~ .321
,01272
,01691 1 . 7 6 2
.01jo5
O.OOOIO
.ooooj -.OOOOI
-,00009
.o.orig8 I 883 0.01810 -.OOOIZ ,02148 2 261 ,02169 - . 0 0 0 2 1 ,02148 2 . 2 6 3 , 0 2 1 7 1 - . 0 0 0 2 3 ,02498 z.6jz .ozj41 -.00043 .028j2 3.066 , 0 2 9 3 2 -.OOO~O ,03697 3,992 , 0 3 8 0 2 -.OOIO; 1
Zt
o.00210 ,00336
xt-xi 0.00012 .OOOO~
.00006 .oooo4 .OOOO~
-,OOOOI .OOOOZ - . O O O I ~
0.01810 1.892 0.01819 -.00009 . o z z i 9 2.416 ,02318 -.00039 . 0 2 j i j 2 . 7 4 7 ,02631 - - . o o o ~ ~ ,03144 3.381 , 0 3 2 2 8 - . o o o 8 4 .03j96 4 . 1 j 3 ,039j4 -.00158
2426
WI'LLIAM CAMPBELL SOMERVILLE
Trimethylamine Zt
0.00424
ATf
Monoethylamine Zi
0 . 4 1 5 0.00401
Zt-xi
Zt
o.00023
,00666 0.658 ,00636 .00030 ,00032 ,00790 0.785 ,00758 ,01018 1.017 ,00981 .00037 ,01364 1.368 ,01318 ,00046 , 0 1 5 4 2 1.j61 ,01503 ,00039 ,00023 ,01688 1.731 ,01665 .00002 .02108 2.193 .on106 ,02451 2.578 , 0 2 4 7 1 - . 0 0 0 2 0 ,02680 2.856 ,02733 - ,000 j3 ,02694 2 . 8 7 0 ,02747 - ,00053
Diethylamine 0.00237 0.240 0.00233 ,00430 0.437 ,00423 .00;36 0.74; ,00722 .o1027 1.045 ,01008 ,01054 1 . 0 7 2 ,01034 ,01173 1.193 . 0 1 1 5 0 ,01674 1.734 ,01667 ,01824 1.918 ,01843 ,01955 2.057 .019;6 ,02254 2.406 ,02308 ,02388 2 , 5 7 9 ,02472 . o z j 8 1 2.817 ,02696 .02892 3.241 ,03096 ,03066 3.461 .03305 ,03350 3.875 ,03693 ,03911 4.694 ,04457
o.00004
.oooo j ,000 I 4 ,000
I9
.00020
.00023 .00007
- .00001 - ,0002I - ,00054 - ,00084 - .OOIIj - .00204 - . CsO 2 39
ATr
zi
0.00405 0 . 4 0 5 0.00391 .ooj51 0.762 ,00736 ,01034 1.065 ,01027 ,01126 1.161 ,01119 ,01595 1,653 ,01590 ,01688 1.758 .01690 .02077 2 . 1 8 0 .02094 ,02330 2.454 .0235z .02jo6 2.646 ,02535 .oz900 3.084 .oz9jo ,03141 3.360 ,03209 ,03329 3.565 ,03402 ,03964 4.298 .04087 ,04213 4.598 ,04368
Zt--zi
0.00014 ,000 I 5 .00007
.00007
.o o o o j
- .00002 - ,00017 - ,00022 - .oooz9 - .00050 - ,00068 - ,00073
- .OOI23 - ,00155
Triethylamine 0.00276 ,00123 .OOj56 .01203
0.278 0.00269 0 . 0 0 0 0 j .00008 0,429 ,00415 .000I9 0.763 ,00737 .0000 I 1,247 . O I Z O Z .00009 1 . 5 4 2 ,01484 .014;5 ,000 I5 .or569 ,01554 1.630 ,01837 1.962 .01885 - ,00048 ,02107 2.271 .oz179 - .ooo.i2 ,02205 2.381 ,02284 - . o o o j 9 ,02582 2.750 ,02634 - . 0 0 0 5 2 ,02764 2.928 . 0 2 8 0 2 - ,00038
- ,00343 - .00546
Values of (zt-z,) are shown plotted against values of z,instead of the actual freezing points of the solutions in order to greatly magnify deviations from the ideal curve.
DEGREES OF HYDRATION OF THE ALKYL AMINES
2427
+0.0005
O?
$i
8 -0.0005
FIG.6 Values of (zt -zi) plotted against values of zi. T h e ideal curve is given by the horizontal axis. FIG.5 . ( I ) Monomethylamine, (2) Dimethylamine, (3) Trimethylamine. FIG.6. ( I ) Monoethylamine, (2) Diethylamine, (3) Triethylamine.
Discussion Part 1 , T h e relative hydration of the methyl and ethyl amines. The formulae of the hydrates of each amine which contain the greatest proportion of water, and which give a definite maximum on the freezing point diagram are, MeNH2,3H20 MezT\",7HzO Mes?j,IoHzO P E ~ N H Z IHZO ,I Et0NH,8Hz0 (Et3X,ZHzO) This suggests that the order of hydration in the series of methylamines is Mono(Di(Tri. I n the series of ethylamines it would seem that diethylamine is more highly hydrated than mono. The position of triethylamine is uncertain, as the curve does not follow the usual course. When the part of each curve given when ice is the crystallising substance is examined, it is seen that the ideal curve is followed for some distance, after which each experimental curve lies below the ideal. Two main influences are a t work, internal pressure differences tending to raise the experimental curve, and hydration of the amine tending to lower it (see Hildebrand', and Kendall'). In more concentrated solutions the latter effect more than counter-
' J. Am. C x m . Eoc., 38, 14j2 (1916).
* J. Am.
Chem. Soc., 39,2303 (1917).
WILLIAM CAMPBELL SOYERVILLE
2428
balances the former, and we can calculate approximately the relative degrees of hydration of the amines as follows, let x, = mole fraction of amine by titration, neglecting ionisation, which is negligible at high concentrations zi = value of z corresponding to the observed value of ATf from the ideal curve as before. If the deviation from the ideal curve were entirely due to hydration, and if molecules of water were comhined with one of amine, the mole fraction of water would be (I - x3 - n.2,)
So that
x,
=
x,+ I -xaxi-x, xi.xn
n Values of
i
=
ii
? I . Ia
= x%-,i.Ia
//
obtained in this way are as follows,
.Ilonomethylamine x,o.o6oo 0.1056 0.1623 n I . 8j 2.28 2.39 compare SleSH2,3H20 Ilimethylamine
I,
n
0,0600 2.9j
0.1002
3 .j6 compare Xfe2SH,;H20
Trimethylamzne. The part of the curve in which ice crystallises is in this case too short to allow this calculated to be used. It is seen from the curve ohtained with the Beckmann thermometer that if prolonged, it riould very quickly come beloiv those obtained nith mono and di methylamines (Fig. j ) . Monoethylamine
xn 0.0600 0.08j o n 2.34 2 . 9 0 compare EtSH2,j . sH20
Diethylamine
x, 0.0600 0,06I 1 n 4.88 4 . 9 9 compare EtlS€I,8H20
Triethylamine. The calculation is again impossible as the ice curve quickly deviates from a normal course. I t does however shoivalarge depression compared with the two other ethylamines at concentrations lessthan 0.0220 (Fig. 6) The values of n increase as the influence of the internal pressure decreases. The increase is probably due in large measure to increasing hydrate stability as t h t tpmperature becomes lower, since the viscosity results show that a rise of temperature from oo to 2 jo does significantly affect the relative viscosities. The viscosities at x, = 0.0130are found to be, Me">
via, ?/?-.
ato"C a t 2j°C
1.138 1.088
Me2NH
MeJ
EtSH2
EtrNH
EtlX
1.225
1,410
1.228
1.j10
1.7zj
1.163
1.248
1.1jr
1.31j
1.38:
The results with mono and di methylamines, mono and di ethylamines, therefore suggest that in dilute solution a t 0°C each of the amines is hydrated t o the same extent as in the first hydrate shown by a later maximum on the
DEGREES O F HYDRATION O F THE ALKYL AXISES
2429
freezing point diagram, the stabilit’y of the hydrate in solution increasing as the temperature is lowered. From all the above considerations, the order of hydration in the methylamines and ethylamines appears to be Mono ( D i ( T r i . This agrees not at all with the result,s obtained from the partition coefficients by Moore and JT7inmill,1and seems to show that the method is unsound. The assumptions made by them which appear most dubious are I - the hydroxide (or hydrate) is insoluble in the nonaqueous layer, and 2 - the solubilities in each layer are unaffected by the presence of a small amount of the second solrent (see KendalP) T h e efect of hydration on the ionisation constant. At high concentrations the water withdrawn by the amine in forming hydrates will alter the effective concentration of the base. I n the case of the weak acids, KendalP showed that when account was taken of the rvater associated with the hydrion, constant values of the ionisation constant could be obtained from I i / z solution downwards. In the case of the amines, the chemical equilibrium may be assumed to be either 1\43N,nH20 H2O S ,143S,nH20 H+ OH- $ hf3?1”+,nHzO OH-
+
+
+
+
. . . , . . . . . . . . . . . . .(A)
+
+
or &43T\’,nHzO HzO @ M3KHOH,nHz0F! >f3?;H+,nH20 OH-. . . (B) The same expression for the ionisation constant is obtained in either case, i.e.
If all concentrations are expressed as moles per litre, using fraction ionised M, = molecular weight of water normality of solute D: = specific gravity of solution M = molecular weight of solute we have, from equation (A) [M3NH+,nHzO]= CYX [lf3N,nH20] = ( I - a:).N
a: X
= =
[OH-]
=
a:K
If equation (B) were employed, the last term in the expression for [HzO]would be PN, where /? = fraction of solute as hydroxide. But CY < P < I , and may be neglected, compared with n. Putting v’ = moles solvent dissolving one mole 100oo.D: - N M then K,
solute =
NM,
,
’ loc. cit. * Proc.
Roy. Soc., 85A, 204 (191I).
J. Am. Chem. SOC.,36, 1079 (1914).
=
2430
T I L L I A M CAMPBELL SOMERVILLE
If we now calculate values of K, for different values of n, there may be a value of n which will give constant values of K,, from S / z solution downwards. The equation used finally, bringing the values of K, to the same order of magnitude as the usual K, was Kn = d X j
j . ~ 1-0)
(v'-n)
The results obtained with two amines follow in detail, a summary is given of the others. In every case except that of trimethylamine at 2 j o , a constant was found.
Monomethylamine. 0°C N
K(usua1)
2.398
2.25
1.219
0.4939
3.28 3.79
0.2025
4.11
0.0996 0.06166
4.12 4.14
KO
Kt
K3
2.49 3.44
2.75
2.91
3.61 4 . I3
3.70 4.17 4 . I9
4.05 4. ' 5 4. I3 4 . '5
4.18 4.1.5
K4
3.08 3 ' 79 4.21 4.21
4.16
4.16 4.17
4.16
4.17
n = 3 gives the most constant' values of K, Trimethylamine. o°C
s
K(usua1)
1.23j 1.228 0.5220
1.75 1.76 2.29
2.39
0.3765
2.40
2.48
0.2055
2.57
0.0979 0.0826
2.62 2.66
n =
12
KO I .92 1'
93
2.61 2.64 2.68
2.39
2.55
Kii 2.64
Ku 2.73
2.40
2.56 2.65 2.66
2.64 2.68 2.68
2.73
2.71
2.72
2.70 2.70 2.73
2.69
2.70
2.70
2.72
2.72
2.73
K*
K
2.59 2.62 2.69 2.68 2.71
O
gives best agreement
Summary of hydration values obtained MelZ"2 Mer" MesN EtSH2 EtzNH Et3N
n 0°C
n zj"C
3 7
6 7
I2
6 or 7 4 3 or 4
no const. value 6 or 7 9 or I O 23
EtSH2,j . 5 H 2 0 E t zNH,8Hz0 (Et aN,2 Hz0)
These results show that this method, while giving values of the correct order in several cases, involves a very large error, especially a t 2 so, due to the
DEGREES O F HYDRATION O F THE ALKYL AMINES
2431
interference of other factors. In the case of the methylamines, the order of hydration agrees with that deduced from other considerations. In the case of the ethylamines, no conclusion can be drawn. Part 2 . T h e znternal pressures of the amines. If there is a large difference between the internal pressures of two substances, positive deviations from Raoult's Law are general in the properties of mixtures of the two. Thus since the internal pressure of water is large compared with that of most substances, we would expect the amine with the lowest internal pressure to give the largest deviation.'
FIG.7 Values of izC-zi) plotted against values of Zi. The ideal curve is given by the horizontal axis. ( I ) Monomethylamine, ( 2 ) Dimethylamine, (3) Trimethylamine, (4) Monethylamine, (j) Diethylamine.
The following approximate relative values for the internal pressures of the amines at o°C were obtained by the method of van Laar.2 MeSH2
Me2NH
5850
4550
hleJ
3750
EtNHt
4750
Et>"
3900
Et&
3650
It is evident that a t low concentrations the effects of hydration and internal pressure largely annul each other. For this to happen in all cases, the compound with the highest internal pressure must be the least hydrated, and vice versa. If we could correct for the hydration of the amines, curves could be drawn showing the internal pressure effect much more markedly. If we allow it to be qualitatively true that in dilute solution a t oo, the amines are hydrated to almost the same extent as in their first hydrates, except in the case of triethylamine, tables may be drawn up in which the concentration of amine has been corrected for the amount of water withdrawn to form hydrate. Hildebrand: J. Am. Chem. Soc., 41, 1067 (1919). Chim. phys., 14, 3 (1916); (also Hildebrand: ibid.)
* J.
2432
WILLIAM CAMPBELL SOMERVILLE
The equation which may be used is xt xo = I - ?ZXt where xc = corrected mole fraction xt = uncorrected mole fraction n = number of molecules of water in one molecule of hydrate. In the following table r , = mole fraction calculated from the ideal equation corresponding to each value of xt as before. ~
Jlonoiizefhylai,zine as l f e S H 2 , 3 H z 0 xt
50
Xi
Dimethylnmz‘ne as b f e a S H , 7 H z 0 xt
Jc-Xj
O . O O ~ O O0.0040j
0.00393 O . O O O I ~ ,01044 .o10;8 ,01038 . O O O ~ O ,01451 , 0 1 5 1 7 ,01446 , 0 0 0 7 1 ,02148 ,02296 . O Z I ~ I .oo12j ,02498 . 0 2 j O O ,02541 .OOIj9
ZC
0.00j8j
Xi
Xc-Zi
0.00612 0.0oj81 0.00031
,01239 ,013jj , 0 1 2 4 0 . O O I I j ,01691 .o1919 ,0170; ,00214 ,02279 , 0 2 7 1 0 ,02318 ,00392 .ozj7j ,03142 ,02631 , 0 0 5 1 1
Trimethylamine as hfe,,S,~oHzO
Xonoethylamine as Et?r”n,5. jHzO
0.00666 0.00j1q 0.00636 0.00oj8 01018 .o1133 ,00981 . o o ~ j z .01j42 .01823 .01;03 .00320 .o2108 .026jo ,02106 ,00564 ,32451 ,03248 ,02471 , 0 0 j j 7 ,02680 ,03661 ,02733 ,00928
0 . 0 0 7 5 1 0.00783 0.00736 0.0004j ,01126 .01200 ,01119 ,00081 ,01688 ,01861 ,01690 . O O I ~ I ,02330 ,02673 .023jz .00321 ,02506 ,02906 ,02535 ,00371
Diethylamine as EtzSH78H2O xt
0.00j36
20
Xi
0 . 0 0 j 8 ~ 0.00i22
xo-21
0.00060
,01294 ,01150 ,00144 ,01674 ,01933 ,01667 ,00266 ,01842 ,02160 ,01843 ,00317 ,01173
Triethylamine. In this case the depression from the ideal curve is greatest of all down to - 2’. The internal pressure appears to be a little less than that of trimethylamine. Since the curves obtained by correcting for the hydration of the other five amines are found to lie in the order of internal pressures, an approximate value for the degree of hydration of triethylamine in dilute solution should be obtainable by calculating how many molecules of water must be combined with one of base to bring the curve up to the trimethylamine curve. For each value of 2, corresponding to a freezing point obtained with triethylamine, the value of x c - x , was found from the trimethylamine curve.
DEGREES O F HYDRATION OF THE ALKYL AMINES
2433
The equation used was as before n = ( x , - x ~ ) / J , . x ~ Zt
o.ooij6
Zi 0.00737
.01203
.OI202
. O Ij j 4
. O I569
,01837
,01885
,02107
,02179
. 0 2j82
,02634 ,02802
,02764
zc-zi
0.0009j .002 I 7 ,00344 ,00470 ,00601
zc
0.00832 . 01419 ,01913
n 12.1 12.7 12.1
,02355
12.0
,02780
11.5
,00864
,03498
10.2
,00976
,03778
9.7
n falls off a t greater concentrations as the curve deviates from a normal course. The most probable value of n appears to be 12, that is, the degree of hydration of trimethylamine in dilute solution a t 0°C appears to be not less than 1 2 .
Conclusions On repeating the work of Pickering on the freezing point curves of the amines, the hydrate of trimethylamine is found to be Me,N,IoHzO and not hfe3N,I 1H2O. The existence of another hydrate of diethylamine, possibly Et2KH,3Hz0, is indicated. The degrees of hydration calculated from the curves obtained where 2. ice is the crystallising substance agree as well as might be expected with the criterion that in dilute solution at o°C, each amine is hydrated to approximately the same extent as in the first hydrate shown by a maximum in the freezing point diagram, except in the case of triethylamine. I t appears that the order of hydration in the series of methyl and ethyl amines is hlono < Di < Tri. 3. A method of deducing the degree of hydration of a solute directly from the conductivities of its solutions is suggested. While in a few cases the values agree with those from the freezing point data, in others discrepancies are found. The effect is thus liable to be masked by other factors, especially at 2 5°C. 4. When the results obtained with the Beckmann thermometer are corrected for hydration to the extent indicated, the positive deviation of the curve given by each amine from the ideal curve increases as the internal pressure decreases. From this consideration, the degree of hydration of triethylamine in dilute solution a t 0°C is deduced. The order of hydration of the amines is not the same as the order of basic strength, but appears to correspond exactly with the order of internal pressure. The greater the difference between the internal pressure of an amine and water, the greater is the degree of hydration. I wish to acknowledge my indebtedness to the hloray Fund of the University for the purchase of the Kohlrausch slide wire used in this work. I wish also to record my thanks to Professor James Kendall for his helpful interest and encouragement. I.
Chemistry Department, T h e I-niuersity, Edinburgh. March 30. 1931.
