T H E ELECTROCAPILLARY CURVE AND ITS DISPLACEMENT WITH CONCENTRATION AND TEMPERATURE' LESTER A. HANSEN AND J. W. WILLIAMS Laboratory of Colloid Chemistry, University of Wisconsin, Madison, Wisconsin Received June 1.4, 193.4 INTRODUCTION
A change in the electrical potential aoross the interface formed by an aqueous solution in contact with mercury is accompanied by a change in interfacial tension which can be measured directly. The relationship between interfacial tension and applied potential may be represented graphically, forming the so-called electrocapillary curve. If we plot some function of interfacial tension as ordinate with applied potential as abscissa a curve of parabolic shape is usually obtained. Before a difference of potential is applied across the interface the mercury is positively charged and anions are preferentially adsorbed in the solution side of the interface. Usually the anions are capillary-active. In this event the more capillary-active the anion is, the greater will be the depression in the rising branch of the curve. As the applied potential difference is increased the charge of the mercury surface is lowered, thus increasing the interfacial tension and causing the electrocapillary curve to rise. A maximum in the curve is reached when the charge on the mercury surface has been, neutralized. Further increase in applied potential causes the mercury to become negatively charged, and the curve descends. If the cations are capillary-active, the descending branch of the curve will be depressed because of the increased concentration of these ions a t the interface. The general mathematical theories of the electrocapillary curve have been presented and discussed in a number of places in the literature. Reference may be made a t this time t o articles by Lippmann (ll),Gouy (5), Frumkin (3), Koenig (€9,and Craxford, Gatty, and Philpot (1). The electrochemical study of the interfacial potential difference between mercury and a solution containing electrolytes, or electrolytes and neutral 1 Presented before the Eleventh Colloid Symposium, held a t Madison, Wisconsin, June 14-16, 1934. More complete details of this work may be found in the thesis of L. A. Hansen, presented to the Faculty of t h e University of Wisconsin in partial fulfillment of the requirements for the degree of Doctor of Philosophy and filed in the Library of the University of Wisconsin, June, 1934. 439
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LESTER A. HANSEN AND J. W. WILLIAMS
molecules, is important because a knowledge of the effects produced gives information about the properties and arrangement of the ions and molecules present in the interface. I n this report it will be our purpose to give the results of a study of the effect of changes in temperature and concentration of ammonium nitrate in aqueous solution on the position and shape of the electrocapillary curve. In a later report there will be presented the results of similar electrocapillary studies with solutions of sodium sulfate and of certain nonelectrolytes, PREPARATION O F MATERIALS
Mallinckrodt’s “Reagent Quality” sodium sulfate was used. Mallinckrodt’s “c.P.” quality ammonium nitrate was recrystallized from water and dried in a desiccator over concentrated sulfuric acid for use. The mercury was purified by placing concentrated sulfuric acid over it and passing air through the liquids. This process was continued for a period of a week, the sulfuric acid being changed several times. After this treatment the mercury was thoroughly washed to remove all traces of acid; then it was dried and distilled. Mercurous nitrate was prepared by allowing concentrated nitric acid to react with an excess of the purified mercury. The mercurous nitrate was removed from the solution by filtration through a sintered glass filter and was dried over concentrated sulfuric acid in a desiccator. All of the solutions were prepared by dissolving the proper weights of the materials in conductance water. The final distillation of the conductance water was carried out in quartz. APPARATUS AND EXPERIMENTAL PROCEDURE
The apparatus employed in this research was constructed of Pyrex glass. It is shown diagrammatically in figure 1. The liquid to be studied was placed in the cell T, which was connected t o the electrometer proper by means of the ground glass joint S. The side arm U was joined with K, a reflux condenser. Because of the fact that temperatures of from 25” to 75°C.were to be used, it was necessary to employ this reflux condenser to keep the pressure in the cell equal t o the atmospheric pressure without changing the composition of the solution. The column of mercury P ends in the capillary L. The inside radius of the capillary was approximately 6 X 10-4 centimeter. Because the capillary was so small, the technique ordinarily used in sealing capillary tubing to larger tubing could not be employed. Eight-millimeter Pyrex tubing was used, because the outside diameter of the capillary tubing was just slightly less than the inside diameter of tubing of this size.. The capillary tube was inserted to a distance of about one-fourth of an inch into the 8-
ELECTROCAPILLARY CURVE AND ITS DISPLACEMENT
441
mm. tubing. The two tubes were then clamped in this position on a ring stand. The glass tubing which extended over the capillary tubing was heated gently with a hand blast torch until a small portion of it became soft. This pliable portion was pressed against the capillary tubing by rolling a carbon rod over the surface. This process was continued until the capillary tubing was entirely sealed into position. By following this method it was possible to seal capillary tubes without collapsing their inner walls. The stopcock J and the ground gIass joint R made it possible to take the apparatus apart whenever it was necessary to replace the capillary tip L. Two brass rings were clamped on the tube; one above, and
FIG.1. APPARATU~~
one below, the joint R. Metal springs were suspended between the two brass rings, thus holding the ground glass joints securely in position. The height of the mercury in column P could be changed by applying pressure from a nitrogen tank through the tube F. The negative electrode, located a t B, consisted of a tungsten wire sealed into the glass tube. The tungsten wires A and C were sealed into the glass vessel coming in contact with the duplicate anodes of mercury, N and M respectively. The stopcock H allowed the pressure in the reservoir 0 to be reduced by allowing nitrogen gas to escape to the atmosphere through it. When the
442
LESTER A. HANSEN AND J. W. WILLIAMS
stopcock I was opened, the pressure was moderated by having another reservoir connected to the outlet D. The entire electrometer was placed in a well-insulated air thermostat. The heating elements consisted of nichrome wire wound in spiral form. The air was thoroughly circulated by means of a fan, the blades of which were placed inside the air bath, while the motor which drove the fan was located outside the thermostat. The temperature was regulated by means of relay and deKhotinsky regulator. The difference of potential to be applied across the electrodes A and B was obtained by means of an ordinary potentiometric circuit. The mercury in the capillary tip L was illuminated by means of an automobile spotlight and a lens which focused the light on the capillary. The source of light was placed at right angles to the telescope of the comparator. It was found that by having the distance from the objective of the comparator t o the mercury capillary very slightly greater than the focal length of the telescope, the mercury in the capillary took on the appearance of a wide ribbon with well-defined tip. If the comparator was so arranged that it was sharply in focus, it was observed that the mercury in the capillary would take on the appearance of a fine thread with a diffuse or ill-defined tip. Instead of measuring the change of position of the mercury in the capillary, it was more convenient to keep the mercury at a constant level by increasing or decreasing the height of the mercury in the column P. The position of the mercury in the capillary L was kept constantly at a point 0.05 cm. from the lower tip of the capillary. It was found by Koenig (9) that if the mercury was kept a t a position such that it was more than 0.1 cm. from the end of the capillary the slight amount of hydrogen which was slowly liberated a t the capillary cathode could not diffuse out fast enough but would form bubbles of hydrogen inside. The height of the mercury in column P was measured by means of a traveling telescope suspended on a 3-in. cold rolled steel rod and a steel scale graduated t o millimeters. The distance from the mercury meniscus in the capillary to the bottom of the steel scale was a constant for each capillary, so the height actually measured with the traveling telescope is the difference between the bottom of the scale and position of the mercury meniscus in the tube P. The distance from the mercury meniscus in the capillary to the bottom of the steel scale was measured by means of a cathetometer. The solution in the cell T extended a short distance above the bottom of the capillary tip L. In order to make the correction for the hydrostatic pressure of this column of solution from the bottom of the mercury meniscus to the top of the solution it was necessary to measure this height accurately. This measurement was made by means of the comparator. To obtain the values of the interfacial tension in absolute units from the
ELECTROCAPILLARY CURVE AND ITS DISPLACEMENT
443
corresponding heights of the mercury column the following procedure was adopted. The value of the interfacial tension at the maximum of the electrocapillary curve for a half-molar aqueous solution of sodium sulfate has been determined by Gouy ( 6 ) . The apparatus was calibrated by means of this solution for each capillary used. EXPERIMENTAL RESULTS
Electrocapillary curves were obtained a t 25", 50", and 75°C. for the following solutions of ammonium nitrate: (1) 1.0 M ammonium nitrate in a solution 0.002 M with respect to mercurous nitrate and nitric acid. (2) 0.2 M ammonium nitrate in a solution 0.002 M with respect to mercurous nitrate and nitric acid. (3) 0.02 M ammonium nitrate in a solution 0.002 M with respect to mercurous nitrate and nitric acid. The electrocapillary data are collected t o form table 1. I n the first column are found the values of 4 , the applied potential difference. The second, third, and fourth columns contain values of interfacial tension corresponding to the several applied potential differences a t 25", 50", and 75"C., respectively, for 1.0 M ammonium nitrate solution; the fifth, sixth, and seventh columns are made up of similar data for the 0.2 M solution; and the eighth, ninth, and tenth columns contain the results for the 0.02 M solution. Representative electrocapillary curves which show the effect of change in concentration and temperature, plotted from these data, are shown in figures 2 and 3. I n order to determine accurately the position of the maximum of an electrocapillary curve, an equation for a portion of the curve extending to a short distance on either side of the apparent maximum was obtained by du the method of the least squares. The first derivative -, when equated to d4 zero, gave the exact position of the curve. This method of obtaining the maximum was used in all cases. In table 2 there are given representative results obtained a t 25°C. The heading for each column is self-explanatory. Such results are shown graphically in figure 4. The curves of this figure are from values obtained by using the mathematical equations, but the points on the curves are experimental values. au (max.) The results of calculations for 4 (max.) and for the coefficients
aT
and
W (max*) for the ammonium nitrate solutions are summarized in table aT
3. I n the third column are given the values of the applied potential which correspond to the maxima of the electrocapillary curves. The fourth column presents the differential coefficient of change of u (max.) with temperature, and the last column gives the corresponding coefficient for change of 4 (max.) with temperature.
444
LESTER A. HANSEN AND J. W. WILLIAMS
TABLE 1 Electroca llaru data .for ammonium nitrate solutions 1.0 M NHrNOr 25°C.
1.0 M 1.0 M 0.2 M NHrNO: NHiNOs NHiNO; 50°C. 75°C. 25°C. dunes
votts
dynes
cm.-l
cm.3
0.000 0.025 0.050 0.075 0.100 0.125 0.150 0.175 0.200 0.250 0.300 0.350 0.400 0.450 0.500 0.550 0.600 0.650 0.675 0.700 0.725 0.750 0.775 0.800 0.825 0.850 0.875 0.900 0.925 0.950 0.975
312.57 319.26 325.41 331.43 337.11 342.47 347.48 352.33 356,95 365.75 373.85 380.94 387.71 393.85 399.45 404.34 408.65 412.48 414.16 415.65 416.82 418 .OO 419,05 420.07 420,66 421.09 421.21 421.21 421.05 420.62 419.92 419.09 418.15 417.14 416.16 414.40 412.68 410.91 408.92 406.88 404.65 402,27 399.76 397.37 391.93 385.87 379,45 372.52 365.52 357.81
314.78 321.52 327.72 333.21 338.82 343.96 348.99 353.62 358.22 366.52 374.43 381.64 388.10 393.91 399.28 403.88 407,74 411.48 412.68 414.40 415.37 416.35 417.24 417.71 417.98 418.10 418.10 418.02 417.75 417.36 416.39 415.37 414.52 413.23 411.59 410.31 408.40 406.53 404.66 402.91 401.23 398.62 396.13 393.56 388.88 382.65 376.49 369.71 362.78 355,03
1.000
1.025 1.050 1.075 1.100 1.125 1.150 1.175 1.200 1.225 1.250 1.275 1.300 1.350 1.400 1.450 1.500 1.550 1.600
dynes
cm.-l
315.61 322.20 328.17 333.80 339.11 344.31 349.12 353.89 358.42 366.64 374.17 381.07 387.32 392.94 397.98 402,44 405.93 409.23 410.27 411.40 412.37 413.11 413.53 414.08 414.39 414.42 414.27 413.80 413,34 412.72 411.90 410.78 409.62 408.53 407.13 405.58 404.22 402.56 400.77 398.91 396.82 394.57 392.32 389.68 384.33 378.63 372.97 365.99 359.20 351.83
dynes
cm.-1
306.27 312.73 319,18 325.37 331.47 337 I34 342.27 347.13 351.94 361,41 369.90 377.57 384.77 391.19 397.45 403.08 407.82 412.48 414.63 416.24 417.80 419.25 420,54 421.72 422.81 423.59 424,06 424.30 424.49 424.49 424.34 424.02 423.28 422.41 421.17 419.99 418.70 417.18 415.65 414.08 412,32 410.33 408.25 406.26 401.52 396.55 391.58 385.13 378.59 371,86
0.2 M NHiNOs
50T. dynes
cm.-l
306.99 313.93 319.93 325.92 331.77 337.22 342.52 347.39 352.22 360.99 369.32 376.96 383.89 390.32 396.17 401.43 405.94 410.62 412,22 413.85 415.45 416.81 417.75 418.53 419.35 420.01 420.36 420.52 420.59 420,52 420.36 419.97 419.11 418.06 416.58 415.65 414.13 412.68 411.17 409.30 407.66 406.06 404.31 401.78 397.69 392,70 387.75 382,22 375.95 369.21
1 1
0.2 M 0.02 M 0.02 M 0.02 M NHINOa YHdNOa NHiNOa N H ~ N O I 75°C. 25°C. 50°C. 75°C. dynes
cm.-'
307.74 313.98 320.22 326.08 331.86 337.02 342.10 347.18 351.95 360.67 368.47 375.84 382.70 388.48 394.41 399,30 403.60 1107,17 1109.03 410.58 411.71 412.68 413.49 414,42 414.93 415.39 415.63 415.78 415.78 415.70 415.55 415 .OO 414.23 413.45 412.49 411.44 410.35 408.88 407.64 405.85 404,65 402.79 401.01 399.18 394.76 389.95 385.03 379.13 373.20 367,ll
dynes cm.-l
1 1 I ! $~t!
dynes
cm.-1
310.46 311.32 310.64 317.03 317.59 316.97 322,67 323.98 322.59 329,12 329.39 328.41 334.72 334.89 333.80 339.73 340.07 339.07 344.821 345.131 344.19 350.41 349.92 349.15 355.30 354.56 353.96 363.95 363.29 361.91 372.29 371.47 369.79 379.88 378.83 376.81 386.96 385,88 383.36 393 381 392.041 389.68 399.331 404.54 397,721 402.44 394.84 399.53 409.31 406.92 403.64 413.65 415..57/ 410 412:411 85 406,QO 408.49 419.17 413.931 417.651 415.33 409.77 411.01 420.78 416 50 411.98 412.91 414,OO 414.58 415.39 415.98 416.40 425.391 421,181 416.52 416.52 416.44 416.25 413.05 415.70 415.04 414.35 413.53 412.37 411.13 418.43 414.28 409.93 416.94 413.00 408,80 415.38 411.67 407.68 413.42 410.23 406.47 411.58 408.36 404.88 408.14 404.39 401.94 403.67 400.69 397.44 399.17 396.05 392.86 393.62 390.99 388.21 388.02 385.65 383.48 381.68 383.74 377.58
FIG. 2.
AMMONIUM
FIG.3. ELECTROCAPILLARY CURVESFOR 0.2 M AMMONIUM NITRATESOLUTION AT 25", 50°, AND 75°C. 445
446
LESTER A. HANSEN AND J. W. WILLIAMS
TABLE 2 Data .for 0.8 M ammonium nitrate solution at 86°C. a
0
OBBERVED
CALCULATED
volts
dynes cm.-1
dynes cm.-l
0.825 0.850 0.875 0.900 0.925 0.950 0.975 1.000 1.025 1.050
422.81 423.59 424.06 424.30 424.49 424.49 424.34 424.02 423.28 422.41
422.81 423.52 424.06 424.39 424.54 424.51 424.29 423.88 423.28 422.48
u
(max.) = 424.49 dynes cm.-* u = 293.52 $ 280.966,
DEVIATION
DEVIATION SQUARED
0.00 -0.07 0.00 +0.09 +O .05 $0.02 -0.05 -0.14 0.00 $0.07 -0.03
0.0000 0.0049 0.0000 0.0081 0.0025 0.0004 0.0025 0,0196 0.0000 0.0049 0.0429
6,
- 150.606,*
(max.)= 0.933 voIt
FIG.4. ELECTROCAPILLARY CURVES FOR 0.2 &fAMMONIUM NITRATESOLUTtON IN THE
REGION OF
THE
MAXIMAOF
THE
CURVESAT 25", 50",
AND
75°C.
447
ELECTROCAPILLARY CURVE AND I T S DISPLACEMENT
In measurements carried out a t 27.5"C., the potential of the 1.0 M ammonium nitrate solution in contact with mercury with respect to the normal calomel electrode was found to be 0.338 volt, that of the 0.2 M solution 0.385 volt, and that of the 0.02 M solution 0.414 volt. TABLE 3 Summary of electrocapillary data SOLUTION
TEMPBRATUBE
$(max.)
ar(max.) 8T
ad(max.) aT
voltv
dynes em.-' degree-'
volts degree-1
"0.
1.0 M NH4N03 1.0 M NHaNOa 1.0 M NH4NOa 0.2 M NH4N03 0.2 M NH4NOs 0.2 M NHrNOj 0.02 M NH4NOs 0.02 M NH,NOs 0.02 M NH4NOa
25 50 75 25 50 75 25 50 75
0.887 0.862 0.846 0.933 0.921 0.914 0.961 0.961 0.953
-0.12 -0.15
-0.00100 -0.00064
-0.16 -0.19
-0.00048 -0.00028
-0.16 -0.19
-0.0000 -0.00032
TABLE 4 Electrical capacity data for 0.8 M ammonium nitrate solution at constant potential values dr IN COULOMBS OM.* x 10' dd
d
0.825 0.850 0.875 0.900 0.925 0.950 0.975 1.000 1.025 1.050
A t 25°C.
At 60°C.
At 75°C.
32.47 24.94 17.41 .9.88 2.35 -5.18 -12.71 -20.24 -27.77 -35.30
28.61 21.21 13.74 6.27 -1.20 -8.67 -16.14 -23.61 -31 .OS -38.55
22.71 16.34 9.97 3.60 -2.77 -9.13 -15.50 -21.87 -28.24 -34.61
da The Lippmann equation for the electrocapillary curve is - = e, where d$ e is the electrical capacity of the surface at constant potential. The value of e can be calculated from the slope of the electrocapillary curve. If the b4 - c42, are differentiated with equations for these curves, a = a respect to 4, values of E are obtained. Data of this sort are presented in table 4.
+
448
LESTER A. HANSEN AND J. W. WILLIAMS DISCUSSION
For discussion it is convenient to divide the electrocapillary curves into three parts, an ascending branch, a maximum, and a descending branch. The ascending branch Along the ascending branch the mercury is positively charged and, as a result, there is an excess of anions in the water side of the interface. From an inspection of figure 3 it would appear that a change in temperature has
FIG. 5. ELECTROCAPILLARY CURVESFOR 0.2 M AMMONIUMNITRATESOLUTION IN THE REGION OF THE ORIGIN OF THE CURVES AT 25", 50°, AND 75°C.
no effect on the position of the first portions of the ascending branch of the curve. However, if these portions of the ascending branches are plotted on a larger scale, as shown in figure 5, it is observed that a t the higher temperatures the interfacial tension values are higher. As the applied potential difference is increased the curves cross each other in such a manner that, at the higher temperatures, the interfacial tensions eventually become lower. This can be explained in the following manner (10). There are two factors operative. An increase in thermal agitation serves t o lower the interfacial tension, and an increase in temperature decreases
ELECTROCAPILLARY CURVE AND ITS DISPLACEMENT
449
the adsorption of mercury salt and acts to increase the interfacial tension. It is apparent that the effect of the adsorption is the more important one near the origin of the curves, but as the curves rise the effect of thermal agitation becomes more and more important and eventually causes reversal in the position of the curves. The maximum At the maximum of the electrocapillary curve the charge on the mercury should be zero, but Gouy (7) considered that additional ions will be present on the surface if they are capillary-active, so that a t the maximum of the electrocapillary curve there may be still a considerable potential difference between the solution and the mercury due to this layer of capillary adsorbed ions. According to Freundlich (Z),if the anions are capillary-active, the mercury must carry a negative charge a t the maximum in order to balance the tendency of anion adsorption. In the same way, with capillary-active cations present the mercury must have a positive charge a t the maximum. It will be observed from an inspection of table 3 that the values of 4 (max.) for the ammonium nitrate solutions become smaller as the temperature is increased. This is consistent with expectation and in agreement with the results published by Koenig (10) for potassium nitrate solutions. Formally, the situation may be discussed with the use of the equation,
4 =
- 4~
(1)
in which 4 r e f . is the potential of reference, and 4 may be described as the potential of the double layer at the polarized electrode. For the temperature variation of the maximum we have,
Contributions to q5p result from the presence in the interface of ions and dipole molecules; in addition, the distribution of electrons and mercury ions in the charged surface of the metal itself is a factor, At the present time it is impossible to determine 4 p and a4 -' with any aT degree of accuracy, because the available methods for the estimation of a4 &f. and a4re.P ' are not entirely satisfactory. But even if c $ and ~ aT aT could be determined, it would be a matter of extreme difficulty to separate the amounts contributed by the three factors,-ions, dipoles, and metal. For solutions of ammonium nitrate there are good reasons to believe the contributions of both ions and dipoles to be finite. In the maximum region of the curve the double layer at the interface contains some adsorbed
450
LESTER A. HANSEN A N D J. W. WILLIAMS
anions. It also contains water molecules which are partially oriented in such a way that the negative end of the dipole is turned toward the mercury side of the interface. Now, since increasing temperature favors z1 decrease in the adsorption of anions and tends t o increase the randomness of any effective orientation, a shift of the position of 4 (max.) toward smaller values with increase in temperature is explained. As would be expected, the actual value of the interfacial tension decreases with increasing temperature and the entire curve is flattened out. From the observed vertical shift of the maximum with concentration of added ammonium nitrate, calculations can be made for the surface densities at the maximum, I' (max.), in equivalents per unit area (cM.~),by using TABLE 5 Surface density data for ammonium nitrate solutions TIN
I
"e.
1.0 M
r(max.1 NHrN03 NHrNOa
- 0.2 M
r(max.1
0.2 M
5.3 x 10-11 3.7 x 10-11 1.9 x 10-11
25 50 75
NHrNOs - 0.02 M NHiNOa
1.0 x 10-11
0.7 x lo-" 0.6 X 10-11
an equation of Gibbs. The chemical potential of the electrolyte, is defined by the expression, P N H ~ N O=~ Const.
+ RT In
CNHdt ' fNH4+
+ RT In
eNor
,UNH,NO~,
' fNO3-
Therefore, ~CLNH~NO =~
2 RT d In
CNHdNOs fNHdN03
= 2 RT d In ~ N H I N O I
The equation of Gibbs, in the form assigned it by Koenig (8), is
The numerical results of table 5 for r$j;&&,have been obtained from graphs in which logarithms of ammonium nitrate activity were plotted as a function of the maximum interfacial tensions. The surface densities are seen to decrease with an increase in temperature and with a decrease in concentration. Also, since aa
(max.)
aCLNH4NOa
is negative, we may conclude that ammonium nitrate is a capillary-active electrolyte.
ELECTROCAPILLARY CURVE AND ITS DISPLACEMENT
451
The descending branch Along the descending branch the mercury side of the interface acquires a greater and greater negative charge and the anions are replaced by cations at the interface. In figure 3 it is observed that the descending branches of the electrocapillary curves are more or less uniformly displaced with a change in temperature. Higher temperatures lower the interfacial tension. The effect of a change in concentration of ammonium nitrate on the position of the descending branch of the curve is shown in figure 2. As the concentration of the ammonium nitrate is lowered, the interfacial tension for equal values of applied potential increases. It was shown by Frumkin (4), and later by Koenig (8), that one can predict the effect of a change in conchtration of the electrolyte on the shape of the electrocapillary curve. According to Gibbs there exists a simple relationship between the adsorbed quantity of a substance and the change of surface tension with concentration, which is expressed by the equation
In this expression I’ is the excess of the solute in gram-moles per square centimeter of the dividing surface, a is the effective concentration or activity of the solute, and u is the surface tension. Frumkin has shown that the equation for the electrocapillary curve can be written in the form,
where pi is the chemical potential of the ithspecies present a t the interface. It can be shown from equations 3 and 4 that
Here the subscripts A and K refer to the anion and cation, respectively, and n represents the valence of the ion. Along the descending branch of the curve only the cations are adsorbed a t the interface, therefore the terms referring to the anions can be neglected, and we obtain the equation
RT When integrated between limits this equation gives
,
452
LESTER A . HANSEN AND J. W . WILLIAMS
In a similar manner one obtains the following expression for the ascending branch.
Equation 7 has not been applied with any degree of success because anions are usually capillary-active. The application of equation 6 to the data for the ammonium nitrate solutions now can be illustrated. There are given in table 6 values calcuRT lated for the quantity - In 2 a t 25", 50", and 75"C., together with n d ai observed values of $2 - $1. The data may be considered to confir& the identity required by the equation. TABLE 6 Electrocapiltary data for 0.09 and 0.8 M ammonium nitrate solutions AT
25°C.
C
410 405 400 395 3 90 385 380
Average =
--RT In az - = F at
AT
0%- 01
a
0.041 0.039 0.043 0.047 0.048 0.044 0.043
410 405 400 395 390 385
50°C.
0.044
Average =
0.053
RT a2 -- In - = F a1
I
AT
76°C.
0%- 01
0.051 (1.053 0.053 0.055 0.051 0.053
0.053 0.058
02
40 5 400 395 390 385 380 375
e1
0.054 0.054 0.054 0.051 0.053 0.053 0.057
Average = az --RT InF a1
-
=
0,054 0.062
Activity coefficient data of Scatchard and Prentiss (12), obtained by the freezing point method, have been used in our calculations. It was necessary to assume these coefficients to be constant over the entire temperature range. SUMMARY
1. A modified Lippmann electrometer was designed and built. With this apparatus electrocapillary curves were obtained for 1.0 molar, 0.2 molar, and 0.02 molar aqueous ammonium nitrate solutions at the temperatures 25", 50", and 75°C. The position of the maximum in each curve was located by means of an analytical procedure. 2. An explanation based upon the capillary activity of the ammonium
ELECTROCAPILLARY CURVE AND ITS DISPLACmMENT
453
nitrate has been presented for the displacement of these curves toward lower values of interfacial tension with increase in temperature and with increase in concentration. 3. A displacement of the maxima of the curves toward lower values of applied potential a t the higher temperatures has been partially accounted for on the basis of a decreased adsorption of anions and an increased tendency toward random distribution of solvent dipoles caused by the rise in temperature. REFERENCES CRAXFORD, GATTY,AND PHILPOT: Phil. Mag. 171 16, 849 (1933). FREUNDLICR: Kapillarchemie, 4th edition, Vol. 1, p. 416. Leipzig 11930). FRUMKIN: Z. physik. Chem. 103, 43 (1922). FRUMKIN: Phil. Mag. [6] 40, 375 (1920). GOUY:Ann. Physik 7, 129 (1917). GOUY:Ann. Physik 6, 5 (1916). (7) GOUY:Compt. rend. 131, 939 (1900). (8) KOENIQ:J. Phys. Chem. 38, 111, 339 (1934). (9) KOENIQ:Z. physik. Chem. 166A, 38 (1931). (10) KOENIQ:Z. physik. Chem. 167A, 96 (1931). (11) LIPPMANN: Ann. chim. phys. [5] 6, 494 (1875). (12) SCATCHARD AND PRENTISS: J. Am. Chem. SOC.64, 2696 (1932). (1) (2) (3) (4) (5) (6)
