388
H. T. BRISCOE AND THEDFORD P. DIRKSE
(11) KAHLBAUM: 2. physik. Chem. 26, 577 (1898). (12) LANQE:Handbook of Chemistry and Physics, 2nd edition, p. 735. Handbook . Publishers Inc., Sandusky, Ohio (1936). J. Phys. Chem. 40,481 (1936). (13) MASONAND WASHBURN: AND COOMBS: J. Am. Chem. SOC.97, 1656 (1915). (14) RICHARDS (16) RICHARDS AND MATHEWS: J. Am. Chem. SOC.3 0 , s (1908). (16) SMITHAND WOJCIECHOWSKI: Rooaniki Chem. 16, 104 (1937). AND MARTIN:J. chim. phys. 23, 733 (1926). (17) TIMMERMANS J. Chem. Soo. 101, 2438 (1912). (18) WADEAND MERRIMAN:
T H E CONDUCTANCE O F SALTS I N MONOETHANOLAMINE
H.T. BRISCOE
AND
THEDFORD P. DIRKSE
Department of Chemistry, Indiana University, Bloomington, Indiana Received July 94, 1030
Solutions in which monoethanolamine acts as the solvent have received very little theoretical consideration. Monoethanolamine is very viscous, it has a dielectric constant approximately half that of water, and it is a very good solvent for many salts. This study deals with the conductances of solutions of salts in this solvent. The equations to be tested include (a) the equation of Debye-Hiickel for calculating activity coefficients, and ( b ) Onsager's conductance equation. EXPERIMENTAL
The bridge used in this work was similar to the bridge recommended hy Jones and Joseph (8). Every precaution they suggested was taken. All the leads were shielded to reduce capacitance, and a condenser was used to balance out capacitance that could not be eliminated by mechanical means. The average amount of this capacitance was 0.000005 microfarad. The frequency of the alternating current was 1250 cycles per second. Transformer oil was used in the constant-temperature bath to reduce stray capacitance between the leads of the cell. The temperature of the bath was maintained a t 25°C. f 0.015", as measured by a Beckmann thermometer which was calibrated against a U. S.Bureau of Standards' thermometer. The cell was made of Pyrex glass and was fitted with platinum electrodes.. It was constructed according to the directions of Jones and Bollinger (6). The electrodes were not platinized, since the solutions used were dilute and alkaline (5). The cell constant was measured by using 0.01 demal potassium chloride which has a conductance of 0.00140877 mho, as determined by Jones and Bradshaw (7). The salts were prepared according to the directions given by Dasher (2). The ethanolamine was purified by distillation. In the first distillation the middle fraction was directed into another distilling flask. This frac-
TABLE 1 Monomethylammonium picrate Ao e 1724.14; Ki 0.4705 X IO-’
c x 1P 1495 252 59 19 7.6 2.5 0.9
fierp.
A
0.989 0.992 0.994 0.995 0.995 0,997 0.998
10.97 36.90 83.05 189.47 381.58 600.00 888.88
0.873 0.946 0.973 0.985 0.990
0.994 0.997
KO-ar x 101
Kostrald x lo-,
-____ 1546 260 62 21 9.8 3.8 1.8
0.06 0.11 0.14 0.25 0.47 0.46 0.49
K’XlV
KXlV
_ _ _ 0.05 0.06 0.11 0.11 0.14 0.14 0.25 0.25 0.47 0.47 0.46 0.46 0.49 0.49
TABLE 2 Dimethylammonium picrate An = 757.57; K I = 5.611 X 10-o cx
1(P
1394 600 262 224 83 33 11 2
A
KO-,
fi,,
19.08 40.30 55.72 64.98 148.69 252.38 372.62 573.00
0.979 0.980 0.984 0.984 0.985 0.988 0.991 0.995
Xlo-,
0.877 0.917 0.944 0.948 0.968 0.979 0.9% 0.994
KatX1V
K’Xlo”
KXlW
0.91 1.80 1.53 1.82 4.02 5.63 5.32 5.55
0.87 1.73 1.49 1.76 3.90 5.50 5.23 5.50
K’XliP
KX1O-d
-____ 1469 0.90 1.79 645 1.52 286 1.80 248 104 3.99 5.59 50 22 5.48 9 5.51
TABLE 3 Trimethylammonium picrate Ao = 543.48; Ki = 3.584 X 1 0 - 6
cx
1P
A
f*W.
.~
953 368 106 34 14
4
28.85 49.91 91.49 129.46 220.84 410.12
f*&
-___
0.975 0.979 0.985 0.989 0.991 0.993
0.897 0.934 0.964 0,979 0.986
~~
1029 412 128 45 24 17
2.84 3.42 3.61 2.59 4.04 10.22
2.86 3.45 3.63 2.58 4.04 10.2
2.72 3.31 3.53 2.53 3.98 10.01
KOitasld
x lo-,
K’ x lo“
K X 10-0
0.06 0.06 0.06 0.09 0.08 0.10 0.17
0.06 0.06 0.06 0.09 0.08 0.10 0.17
0.06
TAB214 Tetramethylammonium picrate ho = 2222.22; K , 0.09145 X lo-’ c x 10s
1256 401 175 65 23 7 3
A
15.52 27.12 41.09 80.99 131.86 251.85 457 Io4
f*W.
0.989 0.992 0.993 0.994 0.995
f*&
K0-w x lo-,
___0.883 1297 0.932 412 0.954 180 0.972 68 0.983 25
0.996
0.M
0.997
0.993
8 4
0.06 0.06 0.08 0.08 0.10 0.17
~
KOnuser
A c x 101 'Aexp. ~ _ _ ~_ 1707 16.17 0.981 0.984 667 29.20 227 50.97 0.988 87 82.74 0.990 119.96 30 0.993 9 0.994 291.97 0.995 3 602.59
cx
fit&.
~ 0.865 0.913 0.948 0.967 0.980 0.989 0.993
x
_ 17860 7004 2423
954 342 135 55
Konnscr
101
XlW
7.23 12.44 17.79 30.25 44.08 74.33 98.28 112.69
7466 2388 927
290 108 45 16 13
0.946 0.960 0.970 0.978 0.984 0.986
0.738 0.842 0.898 0.941 0.964 0.976 0.985 0.987
0.990 0.990
KOit-ld
x lp
K,X
K
X I P
4.44 5.74 6.10 6.36 4.80 11.27 29.28
4.48 5.78 6.13 6.38 4.81 11.30 29.37
_ 4.31 5.60 5.98 6.26 4.75 11.17 29.08
Koatmld X1P
K ' X 1o-a
IC X 1 0 1
8.42 8.13 6.60 6.36 5.43 7.82 5.81 7.07
7.54 7.49 6.21 6.08 5.27 7.61 5.70 6.94
K ' X 10-4
K X l W
13.6 14.8 12.4 12.9 12.9 24.2
12.4 13.8 11.7 12.4 12.6 23.7
-___ 8236 8.26 2626 8.01 1033 6.53 341 6.33 137 5.414 114 7.79 29 5.79 27 7.07
_
TABLE 7 Dimethylammonium chloride ho 3 373.13; K1 = 12.64 X IO-'
cx
101
25.63 47.48 61.04 97.74 155.96 259.83
2641 787 381 138 42 15
CXIP
4311 1418 508 216 69 24 16 8
KO-r
A
I
firdo. x 10-1 ~-
0.953 0.965 0.972 0.979 0.985 0.988
0.835
2949 923 463 189 74 49
0.906
0.933 0.959 0.977 0.986
13.40 14.6 12.2 12.9 12.9 24.1
Trimethylammonium chloride 216.45; KI = 24.97 X lo-'
Bo A
15.56 27.22 47.17 63.16 88.21 132.79 125.88 208.74
K' X 10-
fhXP.
0.939 0.954 0.963 0.972 0.981 0.986 0.989 0.990
0.794 0.876 0.923 0.949 0.971 0.982 0.985 0.990 390
4889 1673 664 310 117 64 40 242
24.1 25.7 31.1 26.7 19.5 24.4
13.8 242
24.6 24.7 31.4 26.4 19.6 24.4 13.8 225
1
K X 10-4
21.8 22.5 29.2 24.9 18.9 23.8 13.5 221
~
CONDUCTANCE OF SALTS IN MONOETHANOLAMINE
391
tion was redistilled and fractionated again; only the middle fraction was collected. These distillations were conducted in an all-glass apparatus, and a t no time did the ethanolamine come in contact with the air. The TABLE 9 Tetramethylammonium chloride An = 45.05; Ki = 150.3 X
cx
10s
2750
1180 550 260 110 60
20
A
11.27 14.07 17.27 23.46 34.55 33.33 40.00
f*sd
Lala.
0.911 0.934 0.950 0.959 0.968 0.976 0.985
0.832 0.886
0.921 0.945 0.963 0.973 0.984
K X I W
3875 1785 919 556 489 236 184
230 167 130 148 27 1 125 133
240 173 135 151
288 129 145
ux)
151 121 139 270 123 141
TABLE 10 Tetraisoamylammonium iodide A 0 = 617.28; KI = 7.117 X lo-' c x 10s A ______ 1245 11.00 21.31 436 162 38.43 65 62.19 24 96.04 10 189.81
0.983 0.986 0.988 0.991 0.993 0.993
0.883 0.929 0.956 0.971 0.982 0.988
Kower x IW
KOstwsld
12990 4589 1750 737 297 145
4.03 5.39 7.08 7.44 7.17 13.68
x
101
K'
x
10-
KXIW
4.05 5.41 7.10 7.46 7.18 13.70
3.92 5.27 6.94 7.33 7.08 13.53
K ' X 101
K X 101
TABLE 11 Tetraisoamylammonium nitrate = m 2 . 7 ; K~ = 1.01 x 10-7
cx
1(P
761 546 178 64 24 9 5
____
A
22.05 26.55 51.63 84.98 145.51 229.12 352.05
.'+ex*.
0.990
0.991 0.992 0.994 0.995 0.996
0.996
f*&.
0.907 0.921 0.954 0.972 0.982 0.989 0,991
KO-r
Xlol
KOstwdd
-___ 7844 5618 1841 676 265 107 63
Xlol
0.72 0.75 0.94 0.93 1.08 1.08 1.53
-~ 0.72 0.75 0.94 0.94 1.08
0.71 0.74 0.93 0.93 1.07
freezing point of the purified ethanolamine was 10.53'C. compared with 10.5'C., the value reported by the Carbide and Carbon Chemicals Corporation. The refractive index of the ethanolamine a t 20°C. was found to be 1.45342 compared with 1.4539,the value reported by the Carbide
392
H. T. BRISCOE AND THEDFORD P. DIRKSE
and Carbon Chemicals Corporation. The dielectric constant a t 25OC. was measured by V. E. Parker of this laboratory and found to be 37.72. W. C. Davis of this 14boratory measured the viscosity of the ethanolamine a t 25OC. and found it to be 0.19346 poise. The solutions were made up by weight. A portion of the salt was weighed in a special flask. This flask was then filled with ethanolamine and reweighed. The cell was filled directly from this flask. The remaining solution was then diluted and t h e process repeated. Ground-glass joints were used throughout to prevent the solution from coming in direct contact with the air. The cell was wavhed out several times with the solution to be measured before a measurement was taken. TABLE 12 Conductance of triisoamylammonium picrate cx
SOLVIIIT
D.K.= 26.04 79.8 per cent ethanolamine 20.2 per cent dioxane
D.K.= 18.80 59.7 per cent ethanolamine 40.3 per cent dioxane
D.K.= 14.62 49.7 per cent ethanolamine 50.3 per cent dioxane
10s
A
120.8 46.51 22.55
12.42 19.35 31.04
719.7 288.8 111.9 47.81 20.41
7.92 8.56 9.83 12.55 68.59
830.4 334.5 116.9 43.94 17.60
7.23 7.77 8.55 13.65 17.05
RESULTS
Tables 1 to 12 show the experimental results. The concentration of the salt in equivalents per liter is represented by C; A is the equivalent conis the experimentally obtained activity coefficient as ductance; fa calculated by the method of Kraus and Fuoss (9), while ffcalo. is the theoretical activity coefficient of the binary salt as a whole calculated by means of the Debye-Huckel equation (3):
KoneasePis the ionization constant calculated from Onsager’s equation (11):
CONDUCTANCE O F SALTS I N MONOETHANOLAMINE
393
KOatwdd is the ionization constant calculated from Ostwald’s dilution law; K’ is the stoichiometrical ionization constant; K is the thermodynamic ionization constant. The conductance a t infinite dilution, Ao, and K I , the experimental ionization constant, were obtained by using the method of Kraus and Fuoss (9). According to D. Belcher (l), extrapolation may not give correct values of ho for weak electrolytes ( K 5 lo+). On this basis the values of A0 for all the salts except the chlorides may be in error. For this reason the
E FIG. 1. Plot of d c against A. 0, monomethylammonium chloride; a, diethylammonium chloride; 8,trimethylammonium chloride; @, tetramethylammonium chloride.
results for the chlorides only are given in graphical form (figure 1). The theoretical slope of the curve &S given by the Onsager equation is also plotted for the chlorides. DISCUSSION
Figure 1 shows that for each salt the experimental slope is much steeper than the Onsager slope. Only for tetramethylammonium chloride do the experimental and theoretical slopes approach each other. The fact that the conductances of all the solutions vary greatly with the concentration shows that these salts behave as weak electrolytes in monoethanolamine.
394
H. T. BRIBCOE AND THEDFORD P . DIRKBE
Walden and Birr (12) measured the conductances of substituted ammonium picrates and chlorides in acetonitrile. Their results show that in acetonitrile tetramethylammonium picrate behaves as a strong electrolyte in the dilute region. Wright, Murray-Rust, and Hartley (13)found that this salt behaved similarly in nitromethane. In monoethanolamine this same salt behaves as a weak electrolyte except at infinite dilution where its conductance is higher than that of the other salts. Walden and Birr’s results vary considerably from the theoretical Onsager values. But in ethanolamine the same or similar salts all vary still more than in acetonitrile. For the chlorides, the experimental slope in monoethanolamine deviates from the theoretical Onsager slope about four times as much as in acetonitrile. The differences in the behavior of these salts in acetonitrile and monoethanolamine may be explained in one or two ways. First, the differences may arise from differences in the extent to which ions are solvated in the two solvents. For all the salts studied, the ionization constants calculated according to Onsager’s approximate equation decrease as the concentration decreases. This may be explained by solvation, which would cause some of the solvent to be used in solvating the ions, thus raising the concentration of the ions in the remaining solvent. This effect should diminish with increasing dilution. The chemical character and molecular structure of monoethanolamine strongly support the belief that it may strongly solvate ions. I n the second place, these differences may be explained by differences in some physical property of the solvents. The dielectric constant of acetonitrile is 36, and that of nitromethane is 37. These values are practically the same as that for monoethanolamine (37.72). If the apparent degree of dissociation depends solely upon the dielectric constant of the solvent, then the salts studied should behave similarly in these three solvents. The fact that this is not so indicates that the apparent degree of dissociation is not dependent entirely upon the dielectric constant of the solvent. An important physical difference in the three solvents is shown by their viscosities; the viscosity of monoethanolamine is thirty-two times as great m that of nitromethane and sixty-five times that of acetonitrile. Since increming viscosity tends to decrease the mobilities of the ions and thus to decrease the conductance, the high viscosity of monoethanolamine may account for the fact that the salts here studied behave as weak electrolytes in monoethanolamine. Kraus and Fuoss (10) found that tetraisoamylammonium nitrate is a strong electrolyte, but that triisoamylammonium picrate is a weak electrolyte when dissolved in ethylene dichloride and in mixtures of this solvent with benzene and dioxane. In monoethanolamine these two electrolytes show a difference only at infinite dilution. It may be that a t dilutions
CONDUCTANCE OF SALTS IN MONOETHANOLAMINE
395
other than infinite the high viscosity of monoethanolamine slows the mobilities of the ions down to about the same level, since the curves of all the salts studied are approximately the same. The stoichiometrical ionization constants agree fairly well with the experimental ionization constants up to concentrations of about 1 X equivalents per liter. The ionization constants calculated according to Ostwald’s dilution law agree very well with the stoichiometrical ionization constants a t all concentrations. Onsager’s equation does not express accurately the ionization constants in these solutions. The values calculated from it do approach the experimental ionization constants as the
2.8 0
c (3
s
1.60 1.20
0.80
1
8
5
4 LOG
3
2
i
C
FIG.2. Plot of log C against log A. (3, dielectric constant = 37.72; 63, dielectric constant = 213.04; 0 , dielectric constant = 18.80; 0 , dielectric constant = 14.62.
dilution is increased, but only in one or two instances is there any agreement in the vicinity of infinite dilution. At very great dilutions, for some of the salts, the calculated ionization constants are greater than the experimental values. This is perhaps explained by the fact that a t these high dilutions the conductance measurements cannot be made as accurately as in more concentrated solutions. The experimental value is only an approximate value., In some instances the ionization constant increases rather abruptly a t certain concentrations. In monoethanolamine the Debye-Huckel equation for %he activity coefficient of a binary salt as a whole reduces to
396
€I T. . BRISCOE AND THEDFORD P . DIRKSE
where C i s the stoichiometric concentration. This equation assumes complete dissociation, since the stoichiometric concentration is substituted for the ionic concentration. Tables 1 to 11 show that for the solutions studied the experimental activity coefficients agree with the theoretical values up to a concentration of about 1 x 10-4 gram-equivalents per liter. Table 12 and figure 2 show the results of the conductance 'measurements of triisoamylammonium picrate in solvents of different dielectric constants. The solvents were obtained by mixing monoethanolamine and dioxane in varying proportions. The percentages are given by weight. As the dielectric constant of the solvent decreases, the conductance a t a definite concentration also decreases. This is to be expected, since a decrease in dielectric constant means greater electrical attraction between the ions. The curves in figure 2 show a tendency to flatten out as the concentration increases, and seem to approach a minimum. This is explained by means of the concept of triple ions as suggested by Fuoss and Kraus (4). It was hoped that it would be possible to extend the measurements to higher concentrations in order to determine a t just which concentration the minimum appears. However, solutions of greater concentrations changed from an orange-yellow to a rather deep red color upon standing a few hours. This change of color probably indicates a reaction between the picrate and the ethanolamine. Thus, measurements in this range of concentration are, in all probability, in error. With dilute solutions such as those actually used, reproducible results were obtained for solutions over varying periods of time. This color change was also noticed in concentrated solutions of mixed solvents containing more than 60 per cent of ethanoladine. For this reason only measurements of the most dilute solutions are used. SUMMARY
1. The salts studied behave as weak electrolytes in monoethanolamine as judged by the change in conductance with changing concentrations, but their conductances are high. 2. The Onsager equation does not give theoretical values for the conductances of these solutions which agree with the experimental values. The lack of agreement is the result, a t least in part, of the failure of this equation to account fully for the viscosity effect. 3. The Debye-Huckel equation for activity coefficients agrees as well as could be expected with the experimental results for the solutions studied. 4. In dilute solutions the conductance of triisoamylammonium picrate decreases as the dielectric constant decreases and approaches a minimum with increasing concentration.
CONDUCTANCE OF SALT8 I N MONOETHANOLAMINE
397
REFERENCES
(1) (2) (3) (4) (5) (6) (7) (8) (9) (10) (11) (12) (13)
BELCHER,D.: J. Am. Chem. SOC. 80, 2744 (1938). DASHER,P. : Thesis, Indiana University, 1937. DEBYE,P., AND H~JCKEL, E.: Physik. Z. 24, 185 (1923). Fooss, R., AND KRAUS,C.: J. Am. Chem. SOC.66, 2387 (1933). JONES, G., AND BOLLINGER, D.: J. Am. Chem. SOC.67, 280 (1935). JONES, G., AND BOLLINQER, G.: J. Am. Chem. SOC.63, 411 (1931). JONES, G., AND BRADSHAW, B.: J. Am. Chem. SOC.66, 1780 (1933). JONES, G., AND JOSEPH, R.: J. Am. Chem. SOC.60,1049 (1928). KRATJS,C., AND Fooss, R.: J. Am. Chem. SOC. 66,476 (1933); 57,488 (1935). KRAWS,C., AND Fuoss, R.: J. Am. Chem. SOC.66, 22 (1933). ONSACIER, L.: Physik. Z. 28, 277 (1927). WALDEN, P., AND BIRR,E.: 2. physik. Chem. A144,269 (1929). WRIQAT,C., MURRAY-RUST, D., AND HARTLEY, H.: J . Chem. SOC.1931, 199.
NEW BOOKS An Introduction to the Principles of Physical Chemistry. By 0. MAASSAND E. W. R. STEACIE. Second edition. 15.5 x 23.5 cm.; ix 395 pp.; 102 fig. New York: John Wiley and Sons, Inc., 1939. Price: $3.00. The second edition of this text has been materially improved by the addition of new chapters on atomic structure, phase rule, and colloidal systems, and also by the rewriting, revision, and substantial extension of chapters dealing with chemical reactions, electrolytic dissociation, and electrochemistry. As in the first edition, the clarity of presentation is excellent. The inclusion of many new figures also constitutes a commendable addition and increases the value of the book as a text. There can be no serious criticism of the scope of the topics presented. Although the content was limited to that which could be covered in about fifty lectures, the inclusion of a short chapter on crystalline state and crystal structure would probably have been justified. Possibly greater emphasis should have been given to determination of molecular weights in the chapter on the gaseous state. Also i t might have been more logical to give the method for determination of molecular dissociation in this chapter than in the chapter on electrolytic dissociation near the end of the book. While the chapter on electrolytic dissociation has been somewhat expanded and improved in this edition, it would have been further improved by more rigorous and, in some instances, more detailed treatment. On the whole, this new edition is an excellent book and will maintain its high place as a text for the students for whom i t was written, namely, those who have had but limited experience in physics and mat,hematics. F. E. BARTELL.
+
+
Eleitromagnetics: A Discussion of Fundamentals. By ALFREDO'RAHILLY.xii 884 pp. London, New York, and Toronto: Longmans, Green and Company, 1939. Cork, Ireland: Cork University Press, 1939. Price: $12.50. Here is a book written with all the freshness and good-natured humorous combativeness that might be expected from one carrying the name of its author.
