ELECTROKIYETIC BEHAVIOUR AND ELECTRODE POTENTIAL BY F. L. USHER
The double layer of ions generally assumed to be responsible for electrokinetic effectsis similar, electrically, to that implied in the generally accepted theory of electrode potential. The question of their identity is one of considerable import,ance. This identity was assumed by Billiterl mho formed an estimate of single electrode potential by determining the concentration of a particular ion required to annul electrokinetic effects, the normal potential being then calculated by means of Nernst’s formula. The value so obtained differed by about half a volt from that based on experiments with a dropping mercury electrode and on the electrocapillary behaviour of the same metal. The reeults obtained by Haber and Klemensiewicz2 and by Freundlich and Rona3 have on the other hand furnished conclusive proof that, at least in the case of the glass surfaces studied by them,electrode potential and electrokinetic potential are distinct. In view of the apparently decisive nature of these later results, it is interesting to find that the identity of the two potentials has quite recently been again assumed4 in the case of metals. If the outer member of the double layer which is the seat of electrode potential is mobile, this and the inner member-and with them the surfaces to which they are respectively attached-must, when exposed to a suitably oriented electric field, undergo a relative displacement which will give rise to electrokinetic effects of the kind generally observed. If, on the other hand, the two members of the “Nernst” double layer are not relatively displaced by the tangential component of an applied electric force, the theory of electrokinetic phenomena due to Helmholtz is by implication deprived of its fundamental assumption, unless we postulate the existence of a second superposed double layer similar to the first in every respect but that of the relative mobility of its components. The former alternative, if true, requires a modification of a somewhat revolutionary character in the current theory of electrode potential, inasmuch as the magnitude and even the sign of the charge on the mobile surface of a metal can be changed by the addition of small quantities of electrolytes that can have little if any influence on the concentration of the metal ion generally regarded as the sole determinant of electrode potential, It raises the question whether electrode potential may in fact be due primarily to an electric charge imparted by an adsorbed layer of ions of any kind, rather than to the existence in the liquid of ions of one particular kind, In order to obtain information which might help in deciding between the two views just referred to, experiZ. Elektrochem., 8, 638 (1902); 15, 439 (1909). Z. physik. Chem., 67, 385 (1909). Sitzungsber. Akad. Wiss. Berlin, 20, 397 (1920). Garrison: J. Am. Chem. Soc., 45, 37 (1923).
ELECTROKINETIC BEHAVIOR AND ELECTRODE POTENTIAL
955
ments on the influence of some uni-univalent electrolytes on the electrode potential of mercury have been carried out, side by side with electrocapillary measurements. The influence of a single electrolyte, potassium chloride, on the electrokinetic properties of a positively and a negatively charged surface has also been examined.
Experimental The mercury used was first purified by spraying through a long column of acidified mercurous nitrate solution, and was afterwards distilled slowly in a vacuum. The electrolytes examined were lithium, potassium, and ammonium chlorides, hydrochloric acid, poiassium hydroxide, and sodium citrate, in four degrees of dilution from millinormal to normal. Two of these substances, ammonium chloride and sodium citrate, are capable of holding finely divided mercury in a stable suspension, positively charged in the first case, negatively in the second. The combination whose E. M. F. was measured consisted of the mercury in contact with the solution under examination, which was connected through a saturated solution of potassium chloride with a decinormal calomel electrode, this pair of elements being connected in series with a standard cadmium cell through a reversing key. Half the difference between the readings obtained before and after the poles of the first-named pair of elements mere reversed mas taken as its E. AI. F. The values at 25' are given in Table I.
TABLE I Potential aaainst N / I calomel ~ electrode at 2 i 0 KCl NHiCl HCl KOH S a Citrate 0.068 0.065 0.111 -0.104 0.125
Dilution
LiCl
N/ I O O Q
0.067
N/IOO
0.052
0.051
K/IO
0.003
0.002
N
-0.062
-0.068
0,049
0.049
-0.150
0.073
-0.003
-0.001
-0.182
0.008
-0.663
-0.068
-0.227
-0.014
Four of t'he electrolytes, potassium chloride, ammonium chloride, potassium hydroxide, and sodium citrate, were then made up with a dilute solution of hydrazine hydrate (0.038 N) instead of water, and the results are given in Table 11.
TABLE I1 Dilution
N/ 1000 N/ I 00 N/IO N
KC1
-0.57 -0.56 -0.47 -0.45
Potential against S / r o calomel electrode at 2;' XHaC1 KOH S a Citrate -0.53
-0.48
-0.50
-0.57
-0.55
-0.44
-0.70
-0.56
-0.27
-0.72
-0.35
-0.55
The electrocapillary curves were determined in the usual manner, and the same solutions (omitting lithium chloride) were used as for the measurements of electrode potential. The voltage corresponding to the maximum interfacial tension was obtained by adjusting the applied voltage in steps of 0.01
956
F. L. USHER
volt; otherwise readings were taken only for steps of 0.1volt. The potential corresponding to the maximum tension (considered to be the natural potential) is given to the nearest hundredth of a volt, and the error is probably not greater than this for the normal and decinormal solutions, but the uncertainty in the case of the millinormal solutions may amount to as much as 0.05 volt, owing lo the flatness of the curve near the maximum a t great dilutions. I t is unnecessary to give the full experimental readings. These were plotted, and from the smoothed curves the constants of the ascending branch of the parabolas were calculated by means of the formula y = Y ~ +CE2, ~ ~C and - E being respectively the capacity and potential of the double layer in electrostatic units. In Table I11 the natural potentials are given, for convenience, in volts, the other electrical quantities in electrostatic units.
TABLE I11 Natural potential (volts)
Maximum interfacial tension (dynes per cm.)
KCl N/~ooo
0.75
IO0
0.70
420 42 6 426 420
2.34 2.32 2.32 2.32
5.9 5.4 4.8 4.1
1.66 2.58 3.36
4.2 5.4
3.8
0.55
414 42 6 423
7.1
1.9
0.63 0.60 0.56
426 427 419
2.2 2.92 3.8
4.6
2.9
5.8
2.2
7 . 1
1.7
0.60 0.59 0.54 0.40
42 1 42 4 417 413
I . 70
.36 I .30 I ,32
3.4 2.7 2.3 I .8
3.8 4.7
XlIOO
0.67
0.57 0.42
423 425 414
2.44 2.94 4.2
5.2 j.6 5.9
2.6
AT/ I O AT
AT-KC1 with hydrazine
0.42
43 4
2.7
3.8
2
Solution
AT/
N/IO N
0.62 0.53
Capacity (e. s. unit
x 10-7)
Surface density of charge (e. s. unit X IO-’)
Thickness of double layer (cm. X 107)
2.7 2.7 2.7
2.7
KH4C1 N/Iooo
0.75
N/IO N
0.63
2.5
HC1 h’/IOO
N/IO
N KOH N/rooo N/Ioo
N,!Io AT
I
5.0
4.8
Na Cit. 2.2
1.5
‘4
ELECTROKINETIC BEHAVIOR A S D ELECTRODE POTERTIAL
957
In Table 111, the interfacial tensions have been taken as proportional to the corrected measured heights of the mercury column, the factor of proportionality being obtained by measuring, in the same apparatus, the height of the column supported by solutions of hydrochloric acid of the strengths used by Lenkewitzl in his determinations of this constant. The following electro-osmotic experiments were carried out in order to ascertain ( I ) whether potassium chloride brings about an initial increase of charge in the case of a positive as well as a negative surface, and ( 2 ) whether this effect persists in the presence of a strongly adsorbed ion having a charge of the same sign as the surface, An “alundum” thimble (made for use in an extraction apparatus) was used as the diaphragm. It was soaked for some days in concentrated hydrochloric acid in order to remove any soluble mineral matter, and then thoroughly washed with water until the electro-osmotic flow became constant. This diaphragm was negative t o water, and when required as a positive diaphragm mas soaked first in dilute alum solution, then in ammonia, and finally washed as before. It was found convenient to measure the pressure balancing the electro-osmotic force in the case of the negative diaphragm, and the rate of flow of liquid in the case of the positive. The inner electrode was a straight piece of stout platinum wire sealed into a glass tube a t its upper end, this tube passing through a hole in the centre of a rubber stopper which fitted tightly into the top of the thimble. The glass tube was provided with distance pieces which ensured the replacement of the electrode in a central position. The outer electrode was a long spiral of platinum wire wound on a glass frame which fitted closely into the containing vessel. The thimble mas replaced after each refilling jn the same relative position with respect t o the outer electrode by taking cross bearings on two fixed pairs of uprights. Measurements mere made of the effect of potassium chloride when the negative diaphragm was in contact with ( I ) water, ( 2 ) 0.2-millinormal NaOH, and (3) millinormal S a O H ; and when the positive diaphragm was in contact with ( I ) water and ( 2 ) 0.5-millinormal Alz(S0J3. The potential difference between the electrodes was 2 2 volts, which produced an easily measurable effect, but even at this voltage the measurements a t concentrations greater than six milli-equivalents per litre were unreliable. The potassium chloride solution was made up with the appropriate liquid, water, NaOH, or Alz(SOJ3, so that the only variable concentration in any particular experiment was that of the potassium chloride. I n Table IF7the figures given under “relative surface charge” are proportional to the readings of pressure or rate of flow of liquid. They denote the surface density of the uncompensated electric charge, and are comparable within each series, although the numbers given in Series I are not comparable with those in Series 11. In all cases the figures given represent the mean of two or more readings. The error depends on the concentration of the electrolyte. I n the pressure measurements the maximum disagreement varied from three per cent. below a concentration of one milli“Untersuchung uber Kapillaritat”. Munster, 1904.
958
F. L. USHER
equivalent per litre to ten per cent. at higher concentrations; and in the measurements of electro-osmotic flow the corresponding errors were roughly twice as great.
TABLE IV Influence of potassium chloride on electro-osmotic effect SERIESI. Diaphragm negative. Relative surface charge in water N/5000 NaOH
Blilli-equivs. KC1 per litre
X/IOOONaOH
0
1
2.71
2.22
0.1
0.5
1.39 1.58 1.44 1.66 1.73
2 ,I2
1.56 2.46 2.39 2.66 2.37
0.6
2.02
0.7 0.75 0.8
1.93
0.2
0.3 0.4
2.58 2.48 2.13
-
-
-
2.44
2.66
1.5
1.83 1.95 I .88 1.97
2 .O
2.05
1.93
3.0 4.0 5 .o 6.0 8.0
2.02
2.17
1.68
2.05
.66 2.13
2.10
0.9 I
.o
2.54
2 . I2
2.20
-
I
-
I
2.17
1.95 1.39
IO, 0
Milli-equivs. KC1 per litre
SERIES11. Diaphragm - -positive. Relative surface charge in water Y/ZOOO Alz(S04)3
0
I
0. I
1.84
0.2
2.20
0.4 0.6 0.8 .o
2.80 2.54 2.92 2.68
1.5 2 .O
2.48
I
3*o
-
4.0
I . 56
6.0
0.56
2.64 2.16 2.04 1.80 I .96 2.32 2.24 2.28
2.08 I .60 0.84
-
.80
ELECTROKINETIC BEHAVIOR AKD ELECTRODE POTEIiTIAL
959
Fig. I shows the relative surface densities of the charge plotted against the concentrations of KC1 so as to indicate the general sense of the effects in the case of the negative diaphragm, and Fig. 2 shows the initial portions of the curves with the concentration scale magnified. Figs. 3 and 4 show the corresponding curves for the positive diaphragm.
I
2
i
4
FIG.I
Fr
I
FIG.2
Discussion of Results If, as it will be assumed, the liquid contact potentials between the solutions examined and the decinormal potassium chloride used in the reference electrode were eliminated by the use of the salt bridge, the measured potentials must be the algebraic sum of the potential differences across all double layers situated a t the interface of the mercury and the solutions used, minus the sum of the corresponding potentials at the mercury-solution interface in the reference electrode. The value of the latter is not known with certainty, but is in any case constant, so that, by subtracting some constant quantity from the experimental values, it might be expected that figures would be obtained bearing some relation to the electrokinetic behaviour of mercury in the corresponding solutions, It is quite evident that no such constant quantity can be chosen so as to reveal even a suggestion of parallelism between the transversely measured potential and the electrokinetic behaviour. Finely divided mercury is known to move towards the cathode in dilute solutions of ammonium chloride, towards the anode in dilute sodium citrate, and not at all in dilute potassium chloride. On comparing the experimental values for these three salts in centinormal and millinormal solution, it is seen that, while there is very little difference between the values for potassium and ammonium chlorides, the mercury is more positive in sodium citrate than in
960
F. L. USHER
either of the others. Again, it is more negative in solutions of potassium hydroxide, although no stable suspension can be obtained in them. There is, in fact, a complete absence of any correlation between the transversely measured potential and the electrokinetic behaviour. One or two points of interest emerge, however, from a consideration of the figures. It may be noticed that the potentials of mercury in decinormal solutions of the chlorides examined are substantially the same as those of the
Y
LL
Milli-equiwak.ntr K C l per l i t r e
= ul er
Milli-equrvalent's KCI per fitre
.z
.4
.6
.8
FIG.4
mercury in the decinormal calomel electrode, and that the influence of change of concentration is, at least qualitatively, the same as in the case of a calomel electrode. This striking similarity of behaviour is not likely to be accidental, and there can be little doubt that the potential of mercury is in both cases determined by the same factor, namely, the activity of the mercurous ions in the liquid. But it was found that mercury could not be detected (by means of hydrogen sulphide) in any of the solutions used, although of course it is easily recognisable in the solution used in the calomel electrode. It would appear therefore that at great dilutions the concentration of mercurous ions a t the mercury surface is almost independent of the bulk concentration, and
ELECTROKINETIC BEHAVIOR AND ELECTRODE POTENTIAL
961
it is reasonable to infer that in all the solutions examined sufficient mercurous ions were formed by chemical action to provide a layer covering the mercury surface. Assuming the chemical action to be due to dissolved oxygen, one finds that only about 2 X IO-^ gram of oxygen mould be required to produce I sq. cm. of a unimolecular film of mercurous ions, It is obviously impracticabie to carry out experiments in which so small a quantity of oxygen could be excluded with certainty, and some measurements were therefore made in presence of hydrazine hydrate which, provided the solution is alkaline, reduces mercury compounds even at very great dilution : mercuric sulphide was indeed the only compound examined which resisted reduction. As is evident from Table 11, the effect of the hydrazine, which was present in the same ccncentration throughout, is remarkable. The potentials are given only to the nearest cen tivolt, since they were not reproducible with greater accuracy, presumably because the electrodes under these conditions are practically irreversible. It will be seen that the effect of changing concentration of the electrolyte is now in the opposite sense, the potential of the mercury becoming more positive with increasing concentration, except with sodium citrate and potassium hydroxide, an observation which suggests that it is no longer determined by mercurous ions. It will also be noticed that there is still little to suggest a connection between the measured potentials and the mobility of mercury in the corresponding solutions.
Electrocupdlary measurements. The constants given in Table I11 are those relating to the parabola y = ymax -$CE2, and were obtained by drawing the curves with the maximum interfacial tension at E = O . The “natural” potentials deduced from the electrocapillary experiments shorn divergencies from those reckoned from the measurements of electrode potential (taking the absolute potentid of the decinormal calomel electrode as 0.61 volt) which are difficult t o explain quantitatively in the present state of our knowledge, but the general results are similar. The figures are recorded here chiefly because we can deduce from them some particulars regarding the properties of the double layer responsible for the natural potential. It has been statedl that the relation between the interfacial tension and potential should be expressed by the equation y = ylnar- CE2,following a suggestion made by Kruger and Krumreich2 that the variation of interfacial tension with surface charge should be doubled in order to include the effect of both sides of the double layer. This suggestion was originally put forward to explain a discrepancy between results obtained by two different experimental methods, but it appears to be contrary to theoretical requirements as deduced, for example, by applying the principle of Virtual Work. Thus, in a double layer of unit area, forming a condenser of capacity C, let a small positive charge, SQ, be transferred from the positive to the negative side. The work terms involved are E6Q+dy/dQ. SQ, which, since the system is in equilibrium and the change takes place a t constant temperature and volume, must be equated to zero. Freundlich: “Kapillarchemie”, 395 (1922). Kruger and Krumreich: Z. Elektrochem., 19, 620 (1913).
962
F. L. USHER
Hence dy/dQ= -E, or dy/dE= -CE. If the capacity is assumed to be constant we obtain on integrating, y = ymax - iCE2, without having made any assumption regarding the effectiveness of one or both sides of the double 12yer. The capacity being thus determined from the electrocapillary curves, it is easy to calculate the charge per unit surface, and also the thickness of the double layer. These quantities are given in Table 111, the thickness of the double layer being calculated on the assumption that the dielectric constant of the medium separating its two components is 8 0 : the values obtained by taking it as unity are so much smaller (about 3 X IO-^ cm.) than the normal spacing of atoms that the former alternative seems preferable. It is interesting to find that, although potassium chloride is probabl. the only one of the electrolytes used which behaves in an approximately normal manner (in not giving rise to any adsorption of its ions at the electrocapillary maximum), the values for the charge per unit surface and for the thickness of the double layer are all of the same order of magnitude. Thus the surface charge (positive) in ammonium chloride is about ;X104 e. s. units, and in sodium citrate it is practically the same, and still positive; yet in regard to its electrokinetic behaviour mercury is positive in the former but negative in the latter. E f e c t of potassium chloride on electrokinetic charge. The curves in Figs. I and 3 indicate that the general effect of potassium chloride on the electrokinetic potential of a negatively charged surface, when no other electrolyte is present, is first to increase the negative charge and afterwards to reduce i t ; while with a positively charged surface under similar conditions the effect is first to increase the positive charge and then .to reduce it. Khen another electrolyte of a kind which strongly enhances the natural charge is alsopresent, the tendency of the potassium chloride to increase it is much less marked. While these statements represent the general tendency of the electrokinetic potential under the conditions named, a closer examination of the figures for very small concentrations reveals a series of subsidiary maxima and minima in the curve, which can hardly he attributed to experimental error: the initial portions of the curves in every case have a sinuous form. It is to be noticed that the first effect of the potassium chloride is to increase the natural charge, whether positive or negative, unless this has already been enhanced by the previous addition of a suitable electrolyte. It has been pointed out elsewhere1 that an initial increase of surface charge may be expected to occur when the free ions produced by the dissociation of surface molecules are replaced by others which have less tendency to combine with the residual sessile ions. Thus, the addition of an alkali metal salt in sufficiently great dilution ia the case of a surface giving rise to free hydrogen ions should always cause such an increase, and this is confirmed by the data so far available. But there are instances in which this explana€ion 1
Trans. Faraday SOC.,21, Part
2
(1925).
ELECTROKINETIC BEHAVIOR AND ELECTRODE POTENTIAL
963
is inapplicable.’ Kruyt, in measuring the electrokinetic effect at a glasswater interface, has found that an initial increase of charge is produced not only by potassium chloride, but also by hydrochloric acid and by barium chloride. This behaviour can hardly be covered by the considerations just referred to, and may best be explained by assuming an adsorption of chloride ions a t the surface of the glass. A similar assumption seems necessary to explain the results shown in Table IT.’ and Figs. 1-4.It is evident that in these experiments two processes are at work whose effects are opposed. In the case of the “alundum” diaphragm, whose natural charge in water is negative, the potassium chloride added furnishes chloride ions which we adsorbed and thus increase the negative charge of the surface, further addition of the electrolyte then actkg in the normal manner and recudipg the dissociation of the surface molecules. These two effects predominate alternately, and would presumably continue to do so until the surface was fully occupied by the negative ions of dissociated and undissociated molecules-a stage which cannot be reached experimentally on account of difficulties of measurement. With an alumina dia,phragm whose natural charge is positive, the cations of the electrolyte are adsorbed, and the further course of events is then much the same as before. This view of what occurs is supported by the observation that when in the first case a strongly adsorbed anion, in the second a strongly adsorbed cation, is initially present, the increase of the initial charge is either much smaller, or does not occur.
Summary I. Measurements of the contact potential and of the electrocapillary behaviour of mercury in solutions of different mi-univalent electrolytes indicate the absence of any direct connection between electrode potential and electrokinetic behaviour in the case of this metal. A comparison of the potentials of mercury in solutions of chlorides 2. with and without the addition of hydrazine, indicates that whenever the metal is brought into contact with such solutions, chemical action results in the production of mercurous ions which are present at the interface in nearly the same degree of concentration as if the solutions had previously been saturated with calomel. 3 . Electro-osmotic experiments on the influence of potassium chloride on the surface charge of a diaphragm point to an initial adsorption of that ion which enhances the original charge, this effect being greatly reduced when a strongly adsorbed ion bearing the same charge as the surface is present. Department of Chemzstry, Cnzcerszty of M y s o r e , Bangalore. Kolloid-Z., 22, 81 (1918).
