Adsorption by Precipitates., VI

Chem. , 1924, 28 (3), pp 232–244. DOI: 10.1021/j150237a004. Publication Date: January 1923. ACS Legacy Archive. Cite this:J. Phys. Chem. 28, 3, 232-...
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Adsorption from Mixtures of Electrolytes; Ionic Antagonism In a paper published in the Autumn of ~ g Iawas reported the results of a series of experiments dealing with adsorption during the precipitation of colloids by mixtures of electrolytes1. The experiments indicated that two factors influenced the precipitating action of mixtures of electrolytes : the effect of the presence of each precipitating ion on the adsorption of the other; and the stabilizing action of the ions having the samc charge as the colloid. If the influence of both of these factors was small the precipit,ation value of mixtures was approximately additive while if the influence of both factors was large the precipitation value was greater than would be expected from the critical values for the separate salts. A year after the publication of my observations, Freundlich and Schola2 reported on the results of similar experiments in a paper entitled “Coagnlation by Mixtures of Electrolytes”. A study was made of the precipitation with various mixtures of electrolytes, of colloidal gold prepared by Donau’s m e t h ~ dof, ~colloidal As& and of colloidal sulphur prepared by both von Weimarn’s method4 and Oden’s method6. The results seemed to show that with gold sol and von Weimarn’s sulphur sol the action of the precipitating cations showed an Additive relationship, while the precipitating action was not additive with As2S3sol and Oden’s sulphur sol particularly with mixtures of ions having widely varying precipitating power that are said to be highly hydrated. Since von Weimarn’s sulphur sol is not hydrous while Oden’s sol is hydrous, Freundlich and Scholz conclude that the hydration of the colloid and of the precipitating ions is of primary importance in producing iohic antagonism and so in determining whether the precipitation values of mixtures shall be additive or above the additive value. They are thus led to believe that As2& is a hydrophile sol although it is not usually so considered; and finally, they suggest that the behavior of colloids with mixtures is a suitable means of determining to what extent the stability is influenced by hydration. “We believe with Neuscholsz,” say Freundlich and Scholz, “that the ionic antagonism observed in a test tube is closely related to the phenomenon which has been recognized as such in the biological action of electrolytes and has been followed by J. Loelo, S. Lillie and others. AS an example of this kind of biological action of electrolytes may be mentioned R series of investigations by S. T.illie6. The cilia of the larva of a ringworm, Weiser: J. Phys. Chem., 25, 665 (1921) Kolloidchem. Beihefte, 16, 267 (1922) 3 Donau: Monatsh., 26, 525 (1905) 4Von Weimarn and Malyschew: Kolloid-Z., 8,214 (1911) 5 Oden: ‘‘Der kolloide Schwefel”, (1912) *Lillie: Am. J. Physiol., 10,433 (1904) I




Arenicola, is liquefied by a solution containing sodium ion: the addition of a small amount of a divalent cation stops this process. At the first glance, oiir observations seem to be different in certain respects from this kind of biological action of electrolytes and from the ionic antagonism observed by Neuscholsxl using lecithin sol. In both of the latter cases it was frequently observed that the action of univalent cations could be nullified by the addition of divalent cations while with sulphur sol the action of divalent cations was decreased under the influence of univalent cations. This is probably only a superficial difference. The coagulation of the sulphur sol is realized only at such concentrations of univalent cations that small concentrations of divalent cations can not annul their influence by displacing their adsorption. If we had studied phenomena like the biological action of electrolytes which could be observed at smaller concentrations of univalent cations, we could have nullified their effect by divalent cations. According to the view arrived at in this paper we believe that in these biological processes as well as in the coagulation of hydrophile sulphur sol, pure electrical influences of pronounced ionic antagonism can not be produced; for ionic antagonism a hydration influence is always necessary, such as appears with increasing strength when we go from gold sol or Weimarn sulphur sol to AszSe sol and finally to hydrophile sulphur sol”. From this account it is evident that Freundlich and Scholz are convinced that the extent of hydration of a colloid is the most important factor in causing the ionic antagonism which results in precipitation values for certain mixtures of electrolytes that are considerably above the additive values. It is a great pity that this conclusion was jumped at from a study of but one sol that appears to belong definitely to the class of hydrophile or hydrous colloids. The conclusions of Freundlich and Scholz would have been more convincing if they had extended their observations to the sols of the hydrous oxides which are quite as representative of the class of hydrophile colloids as is colloidal sulphur. Moreover the preliminary experiments reported in my paper, previously referred to, seemed to indicate that the general conclusions of Freundlich and Scholz were open to question. Accordingly a series of experiments bearing on these points was carried out, with the results recorded in the following section. EXPERIMENTAL Experiments with Hydrous Chromic Oxide Since chromic oxide is very highly hydrous, the colloidal solution of this substance was used in the initial experiments. The positively charged colloid was prepared according to the method of Neidle2by dialysis in the hot of the hydrous oxide peptized by chromic chloride. The dialysis was continued for several days until the dialysate contained but a trace of chloride and the pH of the solution was slightly less than that of pure water. The concentration was adjusted to exactly 2 g Cr2O3per liter. The colloid was stored in a pyrex flask and pyrex vessels were used in all experiments. Pfluger’sArch., 181, 17 (1920) J. Am. Chem. SOC.,39, 71 (1907)



Precipitation experiments. The precipitation value of the colloid wm determined in the following manner: test tube was placed I O cc of sol and in a second test tube, a suitable amount of electrolyte diluted to I O cc. The sol was then added rapidly to the electrolyte and the mixture was poured back and forth a number of times. The tube containing the mixture was stoppered and was set aside for exactly 2 4 hours, after which it was shaken vigorously and centrifuged for 2 minutes at 3000 r.p.m. The presence of complete precipitation was evidenced by the entire absence of color in the supernatant liquid. After getting the approximate precipitation value the exact value was determined by carrying out a series of experiments with electrolyte concentrations that varied not more than 0.01to 0.02 cc. The sharpness of the value obtained was increased by the long period of standing before examining for complete precipitation. The precipitation value of mixtures was obtained in the same manner as for single electrolytes, the total volume of the mixture being made up to IO cc before adding to the colloid. The results of experiments using different pairs of electrolytes of widely varying and of similar precipitating power are given in Tables I, I1 and 111. Instead of expressing the precipitation value in the usual way, I have given simply the number of cubic centimeters of the standard solutions necessary for precipitation in a total volume of 2 0 cc. I n other respects the tables are self-explanatory.

TABLE I Preci itation of Colloidal CrzOawith Mixtures of KC1 and KzS04


N/2KCl taken cc

taken cc

N/IOOKzS04to complete coagulation I calculated I diffe I cc I cc

ice percent



0.85 0.65

I .os 0.75






0 .I 3



0.5 I .o 2

- I9 - I3 -I8 -I8

-0.20 -0. IO

TABLE I1 Precipitation of Colloidal Cr2O3with Mixtures of KC1 and K2C204 N / I O OK G O I to complete coagulation difference cc I percent

Nj2 KCl taken cc



I .SO 0.5 I .o


I . IO





0.40 0.14

0.47 0.17

-0.14 -0.07




- 16 -I 8 -1.5

- 18

23 5


1.35 I

0.50 0.675 I .oo








-0.02 0.0 -0.01

-2 0


The above results with the highly hydrous chromic oxide are clearly not in accord with the conclusions of Freundlich and Scholz. It will be noted that mixtures of electrolytes having widely different precipitating power such as KCl and K2S04do not give values considerably above the additive value such as Freundlich and Scholz would predict. On the contrary the values for such mixtures are actually less than additive by a quite appreciable amount. Such a result is altogether in accord with what one might expect if the antagonistic action of the precipitating ions was slight at concentrations below the precipitation value of each. From the slope of the usual adsorption isotherm it is clear that adsorption is relatively greater at lower concentrations. The adsorption of C1' ion is relatively greater at concentrations below its precipitation value so that relatively less sulphate or oxalate is necessary to bring the combined adsorption above the critical value necessary for neutralization and coagulatjon. As before stated such a result would follow only in case there is little or no antagonistic action between chloride ion and sulphate or oxalate ion in the sense that the presence of each decreases the adsorption of the other at concentrations below the precipitation value. That such is the case is evident from the results given in the subsequent paragraph. Adsorpfion experiments. Adsorption of oxalate ion during the coagulation of colloidal hydrous chromic oxide by mixtures of KC1 and KzCz04was determined both below and above the precipitation value. The procedure was as follows: To 30 cc portions of colloid in a 6 0 cc bottle were added various mixtures of KCI and KzCz04 made up to 30 cc. after allowing the mixture to stand until the precipitate started to settle out, it was centrifuged for I O minutes thereby matting the precipitate in the bottom of the bottle. The supernatant liquid was poured off and. a 50 cc portion was acidified with HzS04, heated to 70' and titrated with N/so K M n 0 4in the usual way using a 2 cc pipette graduated in tenths of a cc. The results of a series of experiments are given in Table IV. The first three experiments in Table I V show that a relatively large amount of chloride has no appreciable effect on the adsorption of oxalate below the precipitation value of the latter. A very large excess of chloride decreases measurably the adsorption of oxalate above the precipitation value but it will be noted the adsorption of oxalate is decreased less than 4 percent by the




TABLE IV Adsorption by Hydrous Chromic Oxide of Oxalate in the Presence of Chloride ____ Mixtures added to 30 cc colloid containing 0.06 gms. Cr90. N/2 RCl N/IOOK2C20,

4.50 3 .oo

Oxal e adsorbed cc N/Ioo 2ms per mole CrtOs

I . 50

1.35 1.95 2.85

24.I 5 25 .os 25.65

1.35 1.95 2.85








8.50 8.50

21 .50 0.00

7.34 6.66

8.182 7.421

I 2 .oo I 2 .oo


9.42 8.83

I 5 .oo



I O .68

12.373 1 1 .go5

12.46 12.35

13.889 13.667

21.50 0.00

18.00 0.00

I 5 .oo



0.00 10.00

20.00 20.00



10.00 0.00

1.505 2. I74



presence of 50 times its concentration of chloride (Exp. 7) and less than I per cent by 25 times its concentration of chloride (Exp. 8). In the light of these observations it is evident that the relatively high precipitation value of KCl is due to weak adsorption of chloride ion associated with but slight adsorption of the stabilizing potassium ion.

Experiments with Colloidal Stannic Oxide While the results with colloidal chromic oxide furnish fairly conclusive evidence that a concentration far above the additive value is not necessarily required to precipitate a highly hydrous colloid with mixtures of electrolytes having widely varying precipitatini power, still it seemed desirable to make additional observations on a negative colloid using electrolytes with cations that are supposedly highly hydrated, as Freundlich did. For this purpose the highly hydrous colloidal stannic oxide prepared by Zsigmondy’s method’ was employed and precipitations were made with mixtures of the chlorides of lithium and barium and of lithiunl and magnesium. The method of procedure was identical with that described above for colloidal chromic oxide. After allowing to stand 24 hours and centrifuging, the supernatant liquid was tested for complete precipitation by adding an excess of BaC12. The absence of a trace of gelatinous precipitate on standing indicated complete coagulation. The results are given in Tables V and VI. The colloid contained 7.6 g SnOz per liter. Zsigmondy: Ann., 301,361 (1898)

23 7


TABLE V Precipitation of Colloidal Sn02with Mixtures of LiCl and BaCL N / I O LiCl taken cc


taken cc

N/IOOBaClz to comr calculated cc

5 . IO

3 . IO .oo


2 .so


1.35 0.55

1.58 0.67




-25 -23

-15 - 18


TABLE VI Precipitation of Colloidal SnOz with Mixtures of LiCl and MgClz


N/IOOMgClz to complete coagulation calculated cc

I_ I


difference I percent

5 . IO 2



3,35 I .98 I .oo

2.04 I .06

The observations with hydrous stannic oxide merely confirm the results with chromic oxide. Here again the precipitation values of mixtures of univalent and bivalent precipitating ions of varying degrees of hydration are appreciably less than the additive values. The results are so nearly in line with those obtained with hydrous chromic oxide that there is doubtless little if any antagonistic action between the precipitating cations below the precipitation value of each. Experiments with Colloidal Arsenious Sulphide From our earlier observations with arsenious sulphide sol it was known that there is an appreciable antagonistic action bewteen electrolytes containing univalent and bivalent precipitating cations in the sense that relatively more of each must be used to cause precipitation in case the other is also present in the solution. Since arsenious sulphide sol is usually classed as a hydrophobe colloid and since the pairs of ions employed (K with Ba or Sr) are said to be but slightly hydrated, it is difficult to reconcile our observations with the view that ionic antagonism is a result chiefly of high hydration of the colloid and the precipitating ions. This view seems all the more improbable since the results with highly hydrated colloids using hydrated precipitating ions, show a tendency opposite to those with arsenious sulphide. To explain our earlier results with As& we assumed that the adsorption of Ba' ' ion, say, was decreased appreciably by the presence of K ' ion below the precipitation value so that a



higher concentrat,ion of the former was required to cause the adsorption necessary for neutralization when the latter was present. Freundlich would reject this explanation on the ground that KC1 precipitates only in high concentrations indicating such a low adsorption of K ' ion that the presence of the latter in concentrations below the precipitation value can have no appreciable influence on the adsorption of Ba' ' ion. The weak point in this argument is that a high precipitation value for an electrolyte does not necessarily indicate very low adsorption of the precipitating ionl. This is true only in case the adsorption of the stabilizing ion is relatively slight. In the case under consideration the adsorption of K ' ion may be fairly large, but in spite of this, a high concentration of KC1 may be necessary since the adsorption of C1' ion at certain concentrations below the precipitation concentration may be comparable to that of K ' ion. It is thus a question of fact, whether the adsorption of K ' ion by colloidal AszSa below the precipitation value is or is not sufficiently great to influence appreciably the adsorption of Ba' . ion. This problem is dealt with in the aubsequent experiments.

Precipitation experiments. Colloidal As2Ss was prepared by dropping slowly a solution of As203 saturated a t 30' into H2S water through which was bubbled continuously a stream of HzS,purified by passing through three washbottles containing water. The excess HzS was washed out by bubbling a stream of hydrogen through the solution after which the latter was filtered into a dark glass bottle. Approximately four liters of colloid containing 2 2 . 3 g per liter were prepared. It was necessary to have a colloid that was fairly strong in order to get sufficient adsorbent in a reasonable volume. The colloid was kept in an atmosphere of hydrogen in a dark bottle. To do this the space above the liquid in the bottle was filled with hydrogen and after inserting the stopper the bottle was inverted thus preventing the escape of the gas. Precipitation values were carried out in much the same manner as for the hydrous colloids except that the time of standing after mixing with electrolyte was decreased to 2.5 houre. The mixture was shaken after the first and second hour and again at the conclusion of the 2 . 5 hour interval. Since the precipitation value of mixtures was the same whether the electrolytes were added separately or together, the latter process was followed as in the previous cases2. To test whether precipitation was complete, the supernatant liquid was filtered into a test tube and examined for the presence or absence of the greenish opalescence that characterizes the very dilute colloid. Since KCI with BaC12or SrClz was used in our earlier experiments it seemed desirable to use different pairs of electrolytes in this work. The results in Tables VII, VIII, I X and X were obtained with mixtures of LiCl and BaCL, LiCl and MgC12, BaClz and MgC12, HC1 and MgC12. Weiser: J. Phys. Chem., 25, 680 (1921) Cf. Freundlich and Schols: Kolloidchem. Beihefte., 16,281 (1922)



TABLE VI1 Precipitation of Colloidal As2S3 with Mixtures of LiCl and BaClz N/2 LiCl taken cc

N/IOOBaC12to corn taken


4.03 4.50 4.25 3.76

calculated cc

3.54 3.03 2.03 I .03





38 84 I I8

1.73 I .22

TABLE VI11 Precipitation of Colloidal AS&%with Mixtures of LiCl and MgC12 N/2 LiCl taken cc



N/roo MgCl? to complete coagulation calculated difference cc cc 1 percent

4.05 4.35




5.20 4.25 2.25





2 .o

3.81 3.27

1.34 1.93





35 59 93 126






.oo 3 .oo 2


4.35 2.20 I . I5









4.35 4.65 4.45 3.35 2 .oo

3.83 2.92 2 .oo I .08


4.75 I .o 2 .o

3 .o 4.0



1 .53

52 68

1.35 0.92


2 40


As observed in the case of hydrous chromic and stannic oxides the precipitation value of mixtures is almost additive in case the precipitating power of each is similar (Table IX). On the other hand, with the mixtures containing both a univalent and bivalent precipitating ion there is a wide variation from the additive value as indicated in Tables VII, VI11 and X.

Adsorption Experiments. Since the quantitative estimation of barium can be made wit,h a high degree of accuracy, a study was made of the effect of lithium on the adsorption of barium at the precipitation concentration of various mixtures of the chlorides of the metals. The procedure was as follows: 1 2 5 cc of colloid containing 2.84 g As&a was precipitated in a wide mouth bottle with the precipitation concentration of electrolyte or electrolytes as given in Table VII. The total volume was 2 5 0 cc. After allowing to stand 2.5 hours the mixture was centrifuged I O minutes at 3000 r.p.m. and 2 0 0 cc of the supernatant liquid was removed for analysis. This solution was first acidified with HC1 and allowed to stand for 2 4 hours. As the concentration of electrolytes was so near the precipitation value, there was always a trace of As2& that settled out after the acidified solution had stood over night. This was filtered off, after which the solution was heated to boiling and the barium precipitated with an excess of H2S04. To render the results of a series of experiments comparable it was necessary to add the required amount of LiCl or LiCl and MgClz to the solution that was free from it originally. It goes without saying that the volume of solution and the amount of precipitant were kept constant.

TABLE XI Adsorption by Arsenious Sulphide of Barium in the Presence of Lithium sorbed msper mole AszSs

RaSO4 remaining in 200 cc

_Average __ 0.0288 0.0280




68.75 68.75




I .310

62.50 62.50







I .250

0 .OOg7



1. I45



74 * 2 0




56.25 56.25 56.25 56.25 53.10 53.10 53 ‘ 10 53.10

71.90 71.90

47.00* 47.00 47.00* 47.00

73 .oo


o 6.25 6.25 0


12.5 12.5 0


25 . o 25 .o



59.40 59.40 73 eo0

48.00 48 .oo


0.0375 0 .OS04

o .0301 0.0363 0.0365 0.0368 0.0259 0.0260


0.0358 0.0357 0.0355






Experiment I gives the adsorption of Ba' ' ion at the precipitation concentration for BaC12; experiments 2 , 3 and 4 give (a) the adsorption of Ba' ' ion from BaC12 alone at the concentration necessary to cause precipitation from the mixture with LiCl and (b) the adsorption of Ba' ' ion in the presence of Li' ion. In experiment 4 the concentration of BaC12is below the precipitation value, necessitating the addition of some MgC12. Since the precipitating power of these two salts is very nearly the same, the presence of the small amount of Mg' ' ion can have little or no effect on the adsorption of Ba' ' ion a t this concentration. Attention should be called to the fact that the concentrations of LiCl and BaClz used in experiments 2, 3 and 4 correspond to those for 0 . 5 , 1.0and 2.0 cc of LiCl, respectively, in Table VII. The reeults are given in detail in Table XI. The observations recorded in Table XI show conclusively that concentrations of lithium below the precipitation value have a marked effect on the adsorption of barium. Thus at the precipitation concentration of a mixture containing one-eighth the precipitation value of LiCl alone the adsorption of Ba' ion is lowered more than 2 5 percent; while from a mixture containing one-half the precipitation value of LiCl alone, the adsorption of Ba' ' is decreased 53 percent. This marked effect of Li' ion in concentrations below the precipitation value on the adsorption by As2& of the strongly adsorbed Ba' ' ion is thus a factor of primary importance in raising the precipitation value of BaC12in the presence of LiCI. The high precipitation value of LiCl cannot be due to very low adsorption of Li' ion which would displace but little Ba' ' ion a t concentrations below the precipitation value; but is due to fairly marked adsorption of Li ion associated with appreciable adsorption of the stabilizing C1' ion within the concentration limits investigated. While the adsorption of the stabilizing ion of an electrolyte is usually disregarded since in many cases the effect is small compared to that of the precipitating ion, the time-honored statement is certainly incorrect that only the anions of an electrolyte determine its precipitation value for a positive colloid and only the cations for a negative colloid. Indeed it is not improbable in the case st hand that at certain concentrations of LiCl or KC1 below the precipitation value, the adsorption of C1' ion is actually greater than that of the cation, particularly in the presence of a divalent ion such as Ba' ' ion'. Another possible interpretation of the high precipitation value of LiCl is that neutralization is compensated for in part by a hydration influence which tends to stabilize the colloid. This hypothesis is untenable since the precipitation value of NaCl is almost as high as of LiCl and that of KCl is of the same order of magnitude although Na' and K ' ions are said to be hydrated much less than Li' ion. A'further objection to this hypothesis is that the relative degree of hydration of the ions is still largely a matter of conjecture. I n discussing this paper a t the Milwaukee meeting Professor Bancroft called attention to the results of some recent investigations of Baboroveky2which seemed '



Cf. Weiser: LOC.cit.

* Rec. trav. rhim. (4) 4, 229 (1923)





to reverse the usually accepted order of hydration of the ions. Until we have more definite information on this point it would seem to be a waste of time to attempt to establish some relationship between the stability of a colloid and the hydration of an adsorbed ion. Returning to Table VI1 and VI11 it will be noted that LiCl has a more marked effect on the precipitation value of MgClz than of BaC12. This ip exactly what one should expect in view of the fact that the adsorption of Mg‘ ion is not so great as that of Ba‘ ion as evidenced by the higher precipitation value of MgC12. However, the variation for mixtures is less marked than that observed by Freundlich and Scholz. Itelerring to their observations the latter investigators say: “From the standpoint of the effect of hydration it is not a t all surprising that the peptizing action of Li’ ion stabilizes the solution to L different degree against cations of the same valence. The coagulation value of Mg’ ’ ion is much greater than that of Ba’ ’ ion in a solution containing LiCl although Mg’ ’ and Ba’ ’ ions have almost the same coagulating power in pure solution.” Unfortunately Freundlich and Scholz seem to have overlooked the fact that small differences in precipitation value do not necessarily mean small percentage differences. As a matter of fact they found the precipitation value of MgClz to be 20 percent greater than that of BaClz a difference that is certainly not negligible. Indeed the percentage difference between the precipitation values of LiCl and KC1 was found to be no greater; yet this was clearly recognized as a difference because the actual figures representing the precipitation values were larger. As our figures show, the difference between the precipitation values of MgC12and BaClz is but seven percent as compared with twenty percent observed by Freundlich and Scholz. This is doubtless the chief cause of the variation between Freundlich’P values and our own in the precipitation with mixtures. The percentage difference in precipitation value of pure electrolytes is not the same on account of the wide difference in the concentration of the sols. Freundlich and Scholz used a very dilute sol, 0.4 g AS& per liter, for some unknown reaeon, while we employed a fairly concentrated sol for the reason already given. That the curves representing the precipitation values of ions diverge more markedly the greater the dilution has been demonstrated and an explanation offered in a recent paper from this laboratory1. I n case adsorption studies are not contemplated one would ordinarily not employ such a strong AszS3sol as we have used for obtaining precipitation data. On the other hand there seems no good reason for working with an As& sol as dilute as 0.04 percent, particularly since it is difficult to get accurate precipitation data with a sol of this concentration. This may, account for the approximate values frequently given in Freundlich and Scholz’s tables. Comparing the results in Tables VI11 and X it will be seen that slightly more MgClz is necessary to complete coagulation of the sol in the presence of LiCl than of HCl. This is readily explained when we consider that two factors enter into a comparison of this sort; the neutralizing action of the adsorbed ‘


Weiser and Nicholas: J. Phys. Chem., 25, 742




Li ion 01 13‘ ion and the antagonistic effect of the presence of these ions on the adsorption of Mg’ ’ ion. On account of its greater adsorption, the antagonistic action is greater than that of Li’ but so is its precipitating power: the net result is that a somewhat smaller concentration of Mg’ ’ ion is necessary to bring about precipitation in the presence of H ‘ ion than of Li’ ion. The results given in Table X are interesting from another point of view. Tarta,r’ and his pupils have recently put forth the view that the precipitation value of all electrolytes for colloids is approximately the same at the same hydrogen ion concentration. This can not be even approximately true in the case cited. The precipitation value of MgClz may be the same at relatively enormous differences in the hydrogen ion concentration. As already pointed out the effect of difference in the H ’ ion concentration on the precipitation values of another cation for colloidal As& is determined by the relative adsorbability of the two ions and by the effect of the presence of variable amounts of each on the adsorption of the other. ’

Summary and Conclusions I. The precipitation values of mixtures of pairs of electrolytes of similar and widely varying precipitating power have been found for positive hydrous chromic oxide sol and for negative hydrous stannic oxide sol and arsenious sulphide sol; and the adsorption of certain ions from mixtures of electrolytes during the precipitation of sols has been determined. 2. The precipitation value of an electrolyte for a colloid is that concentration which r e d t s in sufficient adsorption of the precipitating ion to neutralize the combined adsorption of the original stabi‘izing ion and the stabilizing ion added with the electrolyte. The adsorption of the stabilizing ion varies widely but is never negligible if the electrolytes precipitate only in high concentration as is usual with electrolytes having univalent precipitating ions. 3 . The precipitation values of mixtures of electrolytes of similar precipitating power show an approximately additive relationship because the “ionic antagonism” between the precipitating ions is not marked. 4. The precipitating values of mixtures of pairs of electrolytes of widely varying precipitating power and of different degrees of hydration may be less than the additive values. Thus mixtures of KC1 and either KzS04 or K2C204 precipitate hydrous chromic oxide at concentrations less than the additive values; and the same is true for mixtures of LiCl and either BaClz or MgCL with hydrous stannic oxide. 5 . A study of the effect of chloride ion on the adsorption of oxalate ion both above and below the precipitation concentration of the latter discloses the absence of any marked ionic antagonism. The high precipitation value of KC1 for hydrous chromic oxide is thus due to weak adsorption of chloride ion associated with but slight adsorption of the stabilizing potassium ion. The

Tartar and Gailey: J. Am. Chem. SOC.,44,2212 (1922); A supplementary report was read a t the Milwaukee meeting of the American Chemical Society.



absence of appreciable ionic antagonism accounts for the precipitation concentrations of mixtures not rising above the additive values in the several cases referred to in (4). That the precipitation values in these cases are lower than the additive values is accounted for by the well known fact that adsorption is proportionately greater a t lower concentrations of the adsorbed ion. Thus the adsorption of chloride ion is relatively greater at concentrations well below its precipitation value so that relatively less sulphate or oxalate ion is necessary to bring the combined adsorption above the critical value necessary for neutralization and coagulation. 6. The results given in (4)and ( 5 ) are not in accord with the conclusions of Freundlich and Schols that the hydration of a colloid and of the precipitating ions is of primary importance in producing ionic antagonism and that the behavior of a colloid with mixtures may furnish a suitable means of distinguishing a hydrophobe from a hydrophile colloid. 7. The precipitation values of mixtures of electrolytes that vary widely in their precipitating power may be much greater than the additive values. This was observed on As2& sol using mixtures of LiCl with either BaClz or MgClz and mixtures of HC1 and MgC12. 8. The adsorption of barium ion is decreased to a marked degree by the presence of lithium below the precipitation concentration of the chlorides of the metals. The adsorption of lithium ion is also influenced by the presence of barium ion. This cationic antagonism is the important factor in raising the precipitation concentrations of certain mixtures above the additive values. 9. The high precipitation value of alkali chlorides for colloidal As& is not due to very weak adsorption of the cation but results from fairly strong adsorption of the cation associated with appreciable adsorption of the anion. This accounts at once for the high precipitation values of alkali chlorides and for the displacing power of alkali cations. IO The precipitation value of magnesium chloride is increased more than that of barium chloride in the presence of the same amount of lithium chloride. This is not due primarily to the difference in the hydration of the precipitating ions but results for the most part from stronger adsorption of barium ion than of magnesium ion. I I, The influence of the hydrogen ion concentration on the precipitation values of another cation for colloidal As& is determined by the relative adsorbability of the two ions and by the effect of the presence of variable amount of each on the adsorption of the other. It is thus possible to have the same precipitation value for a given electrolyte at widely different hydrogen ion concentrations. Department of Chemistru The Rice Institute Houston, Texas