TITRATION CURTrES FOR ALUJIINUM SALTS ... - ACS Publications

TITRATION CURTrES FOR ALUJIINUM SALTS WITH ALKALIES B Y H E R B E R T L. DAVIS AYD ESTHER C. FARNHAM

During some work now in progress on the preparation of color lakes with alumina it was observed that the sodium salt of an acid dye is capable of doing the work of far more than an equivalent amount of alkali in precipitat’ing alumina from solutions of aluminum salts. In similar fashion, far less alkali is needed to produce the first permanent turbidity of alumina from solutions of aluminum sulphate than from solutions of aluminum chloride, and the alkali needed in both cases increases with dilution of the salt. The explanation in these cases is that t,he sulphate ions and the dye anions are adsorbed very strongly and cause the neutralization of the positively charged colloidal alumina much inore easily than does the less strongly adsorbed chloride ion. These observations raised questions concerning the titration of such aluminum salts and alkalies as have been given as proof of the formation of such salts as NaAlOz in the alkaline solutions. The present paper reports a repetition of these t,itrations under varying conditions of concentration, and it is shown that in these titration curves there are no inflection point,s corresponding to the amount of alkali needed to begin the precipitation of alumina while in the more concentrated solutions and especially in the sulphate solutions there are inflection points which correspond roughly to the disappearance of the visible particles of alumina at pH 10.5-1 I . These points are achieved in every case with less than four mols of alkali per atomic weight of aluminum and are affected by dilution and by t,he nature of the anion present. They do not, therefore, prove the compound formation but are strong evidence for the peptization of alumina by hydroxyl ions. The earliest work on this subject is contained in the classical paper by Joel Hildebrand’ on “The Hydrogen Electrode in Analysis, Research, and Teaching.” In this paper Hildebrand included the curve for the titration of aluminurn sulphate (history, purity and concentration are not stated) with sodium hydroxide. The curve thus obtained is closely similar to that soon after published by Blum? for the titration of aluminum chloride, “The aluminum chloride solution, prepared from the recrystallized salt, was about decimolar (for AlCl,) and contained some free hydrochloric acid: the sodium and potassium hydroxide solutions prepared from the met,als were about fifth-normal. The initial volume of the titration was about 50 cc.” Blum knew of the work of Hildebrand and says: “Subsequent to the appearance of * This work is part of the programme now being carried out at Cornell University under a grant to Professor Bancroft from the Heckscher Foundation for the Advancement of Research established by August Heckscher at Cornell University. Hildebrand: J. Am. Chem. SOC.,35, 863 (1913). * Blum: :. .4m.Chern. Soc., 35, 1499 (1913).

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HERBERT L. DAVIS A S D ESTHER C. FARNHAM

+

the paper by Hildebrand, the experiment shown in curve B (AICla NaOH) was repeated and continued until considerable excess of alkali was present, No evidence was found, however, of another point of inflection corresponding with the second hydrogen of aluminic acid, the possibility of which was noted by Hildebrand.”

FIG.I

The titration curves of these two aluminum salts, sulphate and chloride, are found to be closely similar, each possessing the same three break-points in the curve of pH against alkali added. Both authors took as quite significant the points a t which the slope of the curve changes, these being analogous to the inflection points which mark the end of titration acid by a base. These curves are shown in Fig. I and the pH values of the inflection points are shown in

TITRATION C C R ~ E SFOR ALUMISUM SALTS WITH ALKALIES

10j9

Table I together with the alkali required to bring the solutions to those pH values. Only the ratios of the volumes used are significant since the concentrations are uncertain. TABLE I Break Points in Titration of Aluminum Salts 1 5 3 7 IO j I1 pH of solutions at break points 0 5 cc 14 18 Alkali used (Hildebrand &(S04)3) Alkali used (Blum AlCl,) 0 9 8 34 42 For the interpretation of these curves we turn first to Blum who was more explicit. "The curve A for the neutralization of hydrochloric acid with sodium hydroxide is shown simply to indicate the normal course of such a reaction in the absence of metals precipitable by hydroxides. I t also furnishes a means for determining the points. of inflection in the precipitation curves, L e . , the points of departure from a normal neutralizat,ion curve. From the curve B for the action of sodium hydroxide upon aluminum chloride, it niay be noted that precipitation of aluminum hydroxide begins when (H') is about IO^, and is complete before (H') is IO-^. The abscissa between t'hese two points, i.e., from 9.8 to 34 cc, or 2 4 . 2 cc, represents the volume of sodium hydroxide required to precipitate the aluminum present. The portion of the curve from (H') = IO-' to (H') = 1 0 - l O . j represents the dissolvingof aluminum hydroxide in sodium hydroxide, the solution being almost entirely clear at' the latter point. That the solution then contains a definite compound consisting of one atom of sodium to one of aluminum is indicated by t'he fact that the volume of alkali from the neutral point to this point of inflection, or from 34 to .p cc is 8.0 cc., i.e., almost exactly one third of t'he volume ( 2 4 . 2 cc.) required to precipitate all the aluminum. This confirms the observation of Prescottl hat one molecule of freshly precipitated aluminum hydroxide dissolves in exactly one molecule of sodium or potassium hydroxide. The same relation holds true in curve C for aluminum chloride and potassium hydroxide where the volumes of alkali required to precipitate and t o redissolve the aluminum hydroxide are respectively I j . j cc. and 6. j. cc. The determination of the exact point of inflection when the aluminum hydroxide is all dissolved is rendered difficult by the partial conversion of the colloidal form as first precipitated, to the crystalline variety, which is difficulty soluble even in great excess of alkali. On standing for several hours some crystalline aluminum hydroxide always separated from the alkaline solutions. The difficult solubility of the last traces of aluminum hydroxide, even with considerable excess of alkali, was more marked with potassium than with sodium hydroxide. No reason is advanced for this difference. Its effect upon the curve for potassium hydroxide and aluminum chloride was practically eliminated by first determining the precipitation portion of the curve, i.e., to 2 0 cc. of potassium hydroxide, and then starting with a fresh portion of solution, making a single addition of 2 0 cc. potassium hydroxide and completing the curve in I cc intervals. By I

Prescott: J. Am. Chem. SOC., 2,

27

(1890).

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HERBERT L. DAVIS AND ESTHER C. FARNHAM

this means the time in which the aluminum hydroxide could change to the crystalline difficulty soluble variety was reduced to such an extent that an almost perfectly clear solution was obtained a t the point of inflection, i.e., (H) = 10-11.’’ Hildebrand is more nearly correct in the interpretation of his sulphate curve and his observation of the precipitation for he says: “Aluminum hydroxide is precipitated while the solution of aluminum sulphate is still strongly acid, the hydrogen ion concentration during the precipitation varying roughly between IO-^ and 10-5.” He does, however, see in his curve the formation of aluminate. “The solution of aluminum hydroxide in an excess of alkali is indicated by the final portions of the curve. It will be seen that the proportion of alkali used corresponds to the formula of NaAlOz.nHz0. There seems also to be a slight break in the curve corresponding to the addition of another equivalent of base. This cannot be regarded as certain until confirmed by more accurate measurements. We believe, however, that the curve here given supports the theory of the solution of aluminum hydroxide as an acid rather than as a colloid, as claimed by Mahin, Ingraham, and Stewart.‘ The ultramicroscope, which these investigators seem not to have applied and which should give rather decisive evidence for or against a colloid theory fails to show the presence of a colloid in this solution, although the ordinary solutions of aluminum hydroxide produced by dialysis show submicrons very plainly in the ultramicroscope.” Blum also asserts that his titration shows the existence of NaAlOz and KAIOz and continues: “If this process of solution were due entirely, or even principally, to the colloid properties of aluminum hydroxide there would probably be a reduction in the alkalinity of the solution, but it is improbable that it would be of such magnitude as has been shown here, much less that the maximum reduction in alkalinity would occur when the alkali was chemically equivalent to the aluminum hydroxide.” The present paper demonstrates that this statement is generally not in accord with the facts. These titrations of Hildebrand and of Blum have been accepted by Weiser2 as the best available evidence to prove the existence of aluminates in alkaline solutions as against the colloidal peptization of alumina by alkalies. I t appears wise to review the whole problem in the light of some similar experiments done more recently which are more in accord with what we might expect of such systems. I n the first place the essential identity of these curves is their undoing and shows either that one or both of these curves are in error or that such titrations show no unique points related to the precipitation of alumina. The close similarity of the titration curves for these salts is confirmed by our own experiments but it becomes evident that this is because they represent a common reaction, the neutralization of the acid freed by hydrolysis of the salts. This is the view later expressed by Blum3 in discussing the estimation Mahin, Ingraham, and Stewart: J. Am. Chem. SOC.,35, 30 (1913). Hydrous Oxides,” rt7 (1926). Blum: Bur. Standards Sci. Paper, 286 (1917).

* Weiser: ”The

TITRATION CURVES FOR ALUMINUM SALTS WITH ALKALIES

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of aluminum. He said that the precipitation of alumina “may be considered as a progressive hydrolysis brought about by the neutralization of the acid continuously set free.” The later investigations of 1LIiller’ will throw some light on this question although, unfortunately, he employed very dilute solutions so that his conclusions might be applied to water purification. Thus the titration of 0.005 molar potassium alum solut,ion and of 0.01molar aluminum chloride solution gives two curves which are very similar. With the first addition of alkali there is a slow increase in pH and then with the addit’ion of the last half of an equivalent of sodium hydroxide the pH of the solutions increases sharply from 5 to 9. There is no sign of any inflection of the curve with low additions of alkali such as at pH 3 which could be interpreted as the beginning of the precipitation of alumina, nor is there any sign of an inflection point up to the four mols of alkali added to the at’om of aluminum. Especially interesting are the experiments of Miller that concern the precipitation of alumina directly. Portions of 500 cc of the aluminum salt solutions were precipitated by the slow addition of 500 cc of sodium hydroxide solution with mechanical stirring. The final concentration was 0.005 molar with respect to aluminum. After the addition of the reagent had been completed, the solution was permitted to st,and a half hour. The hydrogen ion concentration was determined colorimetrically and the precipitated alumina determined. This experiment gave the data shown in Table 11.

TABLE I1 Anion Effect on the Precipitation of Alumina First Precipitate

Potassium alum pH 4.1 required 1 . 2 NaOH/Al Aluminum chloride 7.1 2.9

Complete Precipitation

pH

j.2-2.4

NaOH/Al

7.8-2.9;

These data are in part extrapolated from the curves and data of the paper but they are essentially correct. The smallest actual amount of alkali added to the sulphate solution was 1 . 2 mols NaOH per atom of aluminum and this gave a precipitate of 3 2 7 0 of all the alumina at pH 4.3, while when 2.4 mols of alkali were added, y8YOof the aluminum precipitated at pH j . 2 , and 2 . 7 mols of alkali gave complete precipitation at pH 6.7. In the case of aluminum chloride the addition of 2.9 XaOH gave no precipitate even though the pH was 7 . 1 while the use of 2.9 j mols of alkali brought down 99% of the available alumina at pH 7.8. A further study of these data reveals that the range for complete precipitation of the alumina from chloride solution is a narrow one included within the limits of pH 7.8 to 8.5, while from sulphate solutions practically complete precipitation is achieved over the wide range from pH 5 . 2 to 8.0. This work of Miller taken in connection with the earlier work of Theriault and Clark? demonstrates that the widt,h of these precipitation areas as I

Miller: U. S. Public Health Reports, 40, 351 (1925). Theriault and Clark: U. S. Public Health Reports, 38, 181 (1923).

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well as their actual position depends on the concentration of the salts and upon the anions present. By the addition of ammonium chloride, Miller found the precipitation area for chloride alumina t o be extended from 6.2 to 9.5 and a titration of aluminum chloride with potassium carbonate gave a curve very like that obtained from the sulphate and alkali except that the carbonate curve is displaced about one unit to a higher pH value. One other set of experiments by Miller is interesting. He found it possible to coagulate the aluminum chloride mixtures containing deficient amounts of alkali if he added sodium sulphate. His table is reproduced as Table 111. “In Table I11 are given data showing the smallest quant,ities of sodium sulphate which, when added to one liter quantities of 0.00 j molar aluminum chloride-sodium hydroxide mixtures cause complete flocculation of the colloidal material, leaving the supernatant liquid clear and sparkling after the floc has settled.” TABLE 111 Flocculation of Aluminum Chloride-NaOH Mixtures by XazSO4 PH 4.6 4.9 6.8 8.4 9.0 Mols NaOH/Al 2.0 2.40 2.75 2.90 3.1 Equivalents of Ka2S04per mol A1C130 . 7 0.7 0.3 0 .2 0.0 These data show clearly the rBle of the strongly adsorbed sulphate anion in effecting the precipitation of the alumina at a much lower pH and with considerably less alkali than would be required in the presence of the chloride ion only. A strongly adsorbed dye anion would have a comparable effect. As a first approximation, then, we have that the addition of 3 XaOH per atom of aluminum to a chloride or sulphate solution will result in the complete precipitation of the alumina and will give a solution of about pH 8 to 9. Kolthoff’ found quite empirically that he could titrate aluminum salts with alkali with phenolphthalein as indicator. “Finally I have investigated whether aluminum salts may be titrated with alkali using phenolphthalein as indicator. “Ten cc 0.1 molar alum neutralized i 2 5 cc 0.1 N XaOH (calculated 30 cc) (Kolthoff is using KAl(SOa)2as alum and his solution is o.3N withrespect to aluminum). The end-point was vague. In the presence of an excess of barium nitrate 29.1, 2 9 . 2 , 29.2 cc 0.1 N alkali was used. When the titration in presence in an excess of barium nitrate was carried out at the boiling point, 30.4,. 30.3, 30.3, 30.4 cc 0.1N alkali was neutralized. “On cooling, the solution became red again and it required 0.3-0.4 cc 0.1N HCl to decolorize the phenolphthalein. This method leads to excellent results. The mixture of the alum solut,ionand excess barium nitrate is titrated a t the boiling point until the color of phenolphthalein is pale pink. The solution is cooled and the small excess of alkali is titrated with acid. “The experiment was repeated with aluminum chloride solution which was 0.151 N (standardized by silver nitrate according to Mohr, with an excess of ‘Kolthoff: 2. anorg. Chem., 112, 172 (1920).

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magnesium oxide). Ten cc a t room temperature required 11.75,14.70 cc 0.1N alkali with phenolphthalein (calculated 15.1 cc). At the boiling point IO cc required 15.20, I j.22 cc 0.1N alkali; after cooling it was back-titrated with 0.12, 0.15cc 0.1 N HCI. The results show that the method described is a good one.” Our first attempt to duplicate these results of Kolthoff showed that he had not emphasized sufficiently the necessity of considerable boiling of the sulphate systems. Thus, following his directions, I O cc of 0.3 N alum (with respect to aluminum) required 29.1 cc 0.1N alkali and it was only when the solution was heated for about five minutes before addition of alkali and boiled for five minutes after the faint pink of phenolphthalein has been obtained that 29.9 and 30.3 cc were required. I t is obvious that the method is not immediately useful for the determination of the amount of aluminum present unless there is some assurance that the aluminum salt used is a neutral salt Excess acid will be accounted as aluminum and a basic salt will give low values for the aluminum present. It seemed desirable to have a clearer idea of the reason for this titration and the effect of the variables upon it. Ten cc portions of aluminum chloride and of aluminum sulphate solutions were titrated with 1.19 N NaOH and back-titrated with 0.136 K H2S04. The apparent concentrations of these solutions are shown in Table IV found by titration under various conditions.

TABLE IV Apparent Kormality of Aluminum Salt Solutions A1C13

Cold, no salt Same after boiling and cooling Cold 2 5 cc 0.5 N Ba(X03)2 Same, boiled and cooled 12.5 cc N Na2S04 Cold Same, boiled and cooled

+

+

975 K 0 997 o 988 I 008 0 940 0 997 0

hl (soi)3 0

940 N

I 003 0

991

I

023

0

910

I 00;

Some very interesting observations of these data can be made. In the first place, with both solutions the blank and the system with sodium sulphate give results which are identical while the addition of barium nitrate requires the use of more alkali. It was shown that in the case of the chloride solution the value obtained by titrating with silver nitrate gave 1.001N for the confirming the lower titration value and showing that the barium salt titration is in error. Kolthoff was not clear as to whether barium nitrate should be used in the chloride titration but it follows from this that it should not. A glance a t the table will show exactly the same relations in the sulphate solutions except the error introduced by the barium nitrate is even larger. If he had tried it, Kolthoff would have found that the sulphate solution too could be titrated without salt addition. The essential difference is that the solutions containing sulphate require much longer heating and Kolthoff merely made the plus

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error of barium nitrate addition balance the minus error of insufficient heating so as to arrive a t the proper result. We are now prepared to arrive a t an understanding of this titration. The important process is the liberation of the acid produced by hydrolysis and then the separation of this acid from the strongly adsorbing hydrous alumina. Even the original solutions contain colloidal alumina and heating and alkali addition increase the amount of this form until it exceeds the specific solubility under the circumstances, the most important factor being the anions present. This colloidal alumina as formed adsorbs aluminum ions and hydrogen ions and in the presence of sulphate ions these too are adsorbed and tend to neutralize the charge on the alumina and precipitate it. I t is clear, then, that the alumina will carry down acid and thus prevent its neutralization until a high pH is reached. The table shows this effect to be greater in the presence of sulphate ions than of chloride ions. Thus the first persistent (not permanent) pink of phenolphthalein makes the chloride solution cold appear to be 0.975 N while the sulphate solution which finally requires the same amount of alkali shows this pink as of 0.94 N, the explanation being that the sulphate is markedly more effective in increasing hydrogen ion adsorption on the alumina. The addition of sodium sulphate to the aluminum chloride has the same effect. Although the sodium sulphate is barely acid to phenolphthalein, its addition to the aluminum chloride makes necessary the addition of 0.3 cc N alkali less in order to obtain the pink color. This is undoubtedly the explanation of the recent work by Thomas and Whitehead.' They reported that sulphates when added to solutions of aluminum chloride or sulphate were more effective in raising the pH than chlorides are. Sulphates are more effective in increasing the adsorption of hydrogen ions on alumina and thus raising the pH of the solution so that in the present case it requires less alkali to titrate and therefore appears less concentrated than it really is. Since the sulphate ion is notorious for its effect on alumina it was logical for Kolthoff to add barium salts to remove that effect. The difficulty is that an excess overcorrects the error and actually no salt is needed if the boiling be s;fficiently vigorous and for a long enough time. Thus even excess sulphate added has little effect on the titration under these conditions. But the excess of barium salt appears to act by being adsorbed and in the alkaline solutions increasing the adsorption of hydroxyl ions leading to high alkali requirements. Other strongly absorbed cations should behave similarly and the addition of 2 5 cc N Ca(NO& to an aluminum sulphate titration gave 1.021 N while an equivalent amount of sodium nitrate gave 1.003 N as the concentration of the aluminum sulphate solution. A similar behaviour was shown for barium sulphate when I O cc N sodium sulphate and 2 5 cc 0.5 N barium nitrate were made very faint pink to phenolphthalein and then mixed. On boiling the excess barium ion was adsorbed on the barium sulphate and carried down increasing amounts of hydroxyl ion until 0.5 cc 0.I N alkali had been added to keep the color a t the faint pink of the indicator. On cooling one drop of 0.136 'Thomas and Whitehead. J. Am. Leather Chem. Assoc., 25, 127 (1930);Chem. Ab8. 24, 2659 (1930).


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N H z S 0 4decolorized the suspension so that this experiment shows that adsorption by barium sulphate accounted for about one fourth of the error observed in the tit,ration of the aluminum sulphate in the presence of excess barium sulphate. The high adsorptive power of the alumina acc0unt.s for the remainder as it is due likewise to the adsorption of barium and hydroxyl ions. As the table shows, barium ions are effective in reducing hydrogen ion adsorption and in t,hus freeing a larger amount of the acid for titration cold but on heating in the alkaline solutions they do increase the error. Similarly sulphate ions on the acid side increase acid adsorption but in the alkaline solutions oppose adsorption of hydroxyl ions. Since the titrat’ion is to be finished on the alkaline side, the use of sodium sulphate might be indicated. But the table shows that’ no salt is needed and that salts of the type of barium nitrate are distinctly not to be used. The conclusions of this,matter, then, is that the addition of alkali to cold solutions of aluminum salts will give a pink color of phenolphthalein before three equivalents of alkali have been added and that this error is greater in the sulphate solution than in the chloride solution. This is the definite proof that Hildebrand and Blum by adding alkali to such solutions did not titrat,e the salt solutions as they believed, but that a pH of 7 was reached with considerably less than the equivalent,amount of alkali. I n the curves to be shown later the magnitude of this error will be demonstrated and its variat,ion with concentration and anions present. An investigation of 0.1 N AlCla and 0.1 N A12(SOa)3 showed that comparable effects are found also in more dilute solutions. The Explanation of the Curves of Blum and of Hildebrand We are now prepared to study these curves of Hildebrand and of Blum in more detail and to point out that not only do t’hey include free acid but, that they cannot mean the things that their observers interpret’ed them to mean. The first portion of their curves from pH I . j to pH 3 is as Blum intimated due to the presence of excess acid, since this curve is quite like the final portion of the curve of free acid. Further neither Clark nor Uiller had any such curve for the salts in question and they were using more refined methods. One comment’ on the earlier work is of interest. “If these curves are compared with those published by Hildebrand (1913) and by Blum (1913-IA,), there will be found a general agreement in main feat,ures. However the measurements by Hildebrand and by Blum were made with comparatively crude instruments, and for this reason the observers probably hesitated to call attention to the detailed features in the titration curve which must have appeared to them very peculiar. One notable feature is the distinct slope of the curve between pH 5 and 8. The fact that the steeper part. of the curve should occur so distinctly ahead of three equivalents of alkali is also food for thought. The flatness of the curve at the start is, of course, accounted for by the throwing out of one or more constituents of the equilibrium as the titration proceeds.” Theriault and Clark: U. S. Public Health Reports, 38, 181 (1923)

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Our own titrations on salts which were known to be free from excess acid or base shows no first portion of curve from 1.5 to 3 but start from p H values between 2.3 and 3.4 and are very slow to change their pH on addition of alkali. It is probable, but not proved, that excess acid if present would be titrated sharply to pH 3 as these early authors believe. The next point that should be made is that the essential identity between the curves for the sulphate and chloride titrations is the best evidence that neither tells anything about the start of the precipitation of alumina. The experiments of Clark and of Miller show that alumina begins to be coagulated into particles caught by the filter under quite different conditions of alkali addition and pH. If the alkali be added to 0.1N aluminum salt solutions and they be heated a little to accelerate equilibrium, it is found that from 80 to 90% of the theoretical alkali may be added to the chloride without producing a permanent precipitate of alumina while to the sulphate solution not more than about 20% can be added. These are systems in equilibrium while the observations in titrations are distinctly not in equilibrium. If the titrations are carried out slowly enough the alumina precipitated by local alkali concentrations on first additions will be dispersed and the chloride solution will stay clear much longer than the sulphate solution. Alkali addition must increase greatly the proportion of aluminum present as alumina before it appears as visible aggregates and this last step will be aided greatly by sulphate adsorption. I t is clear therefore that a t pH 3 Hildebrand had a rough index of the beginning of titration of the aluminum salts after the neutralization of the free acid and that the formation of visible permanent precipitates of alumina would lag behind this point and will occur for the sulphate before it does for the chloride. The practical identity of the curves can be related only to the fact that the essential reaction taking place is the titration of the hydrochloric and sulphuric acids set free by hydrolysis. The lag in this hydrolysis causes the p H during the addition of alkali to rise more slowly than if all the acid were free originally. This slow setting-free by hydrolysis makes the strong acids titrate as though they were weak acids merely because of the relatively low free acid concentration in equilibrium with the hydrolyzing solution. The formation of the alumina is then given by the equation: 2A1C13 3HOH ?=? A1203.xH20 6HC1.

+

+

and the displacement of this equilibrium to the right by the neutralization of the acid results in the accumulation of the alumina, first as colloidal positively charged alumina. As its concentration increases, the alumina begins to precipitate but a very important factor is the presence of strongly adsorbed anions capable of neutralizing the original positive charge of the alumina below its critical value and thus bring about the precipitation. I n this picture the hydrogen plays only a secondary rBle and it is quite impossible to speak intelligently of the isoelectric point of alumina unless one specifies the other constituents of the system. The center of Miller’s precipitation ranges is of the nature of an isoelectric point for alumina but he found this point to be


1067

great,ly affected by the anions present. A strongly adsorbed anion will neutralize the positive charge impressed on alumina by a solution of low pH or high hydrogen ion concentration and this low pH will be the isoelectric point for that anion and for that concentration. Similarly alumina precipitated in t,he presence of the weakly adsorbed monovalent chloride ion will adsorb so little of this anion that the hydrogen adsorbed from a solution at pH 7 or 8 is sufficient to give an aggregate whose charge is below the critical value, and 7 or 8 is the isoelectric point then. I t is, therefore, highly improbable that. aluminum chloride should start to precipitate at pH 3 and complete its precipitation at pH 7 ; and the volume of alkali used between these points can have no such meaning as Blum ascribes. The use of one third of this amount to redisperse the alumina is sheer coincidence and it did not, favor Blum when he used potassium hydroxide, the titration being at least 10% in error. The explanation of this is that the excess alkali is dispersing the alumina as a colloid and the difference between sodium and potassium hydroxides is to be ascribed to a different adsorpt'ion for the potassium and sodium ions. It has already been shown that it is quite impossible for aluminum sulphate to give a curve identical with that of the chloride if the break points were unique for the alumina produced. We may conclude then that this first break point, at pH 3 is an approximate measure of the free acid present. Aluminum chloride and sulphate are notorious for forming acid of basic salt's but potassium alum is crystallized in better form and Miller with this salt observed no free acid. Our runs on neutral salts show none. This point on the early curves drops out of consideration. The second point of interest is at pH 7 which they selected as the eyuivalent point. Our experiments show that this is below the amount of alkali required for real titration and that it varies with the anion. Further, the curve as it' passes through pH 7 has a distinct slope and is not vertical as it is for free mineral acid. By far the most interesting points on these curves are the break-points at pH IO.j to 1 1 . I t has been shown that the relative volumes of alkali required for these points was nearly one third more than that required to reach pH 7 with sodium hydroxide. I t must now be emphasized that neither Hildebrand nor Blum showed that these were in the ratio of 3 NaOH/Xl and 4NaOH/Al. Our own experiments agree with this as to the relative amounts of alkali used; but the actual figures come out about 2 . 6 NaOH and 3.7 NaOH for the sulphate and about 2.7 NaOH and 3 . 4 NaOH for the chloride titration. There is in these figures no support for sny conclusion as to the sharp formation of sodium aluminate.

The Titration Runs In order to study these curves in more detail, solutions were prepared to be N, 0.3 N and 0.1 N with respect to aluminum chloride and aluminum sulphate and sodium hydroxide. The salts used were of the best grade and had been previously studied by the precipitation of alumina and sulphate and by the silver nitrate titration. They were thus established as neutral

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aluminum salts without free acid or base, and the solutions prepared were standardized by comparing them with the alkali solution according to the method of Kolthoff. I n each case thirty cc portions of the aluminum salt

T i a t i o n Cutves

N Solutions .3N Solutions o.lN Solutions eElectPode left 0

Cubic Centimeters NaOH FIG.2

solutions were titrated in an open beaker stirred by hydrogen bubbling over a strip of platinized platinum electrode. Hydrogen from a commercial cylinder was purified by passage through alkaline pyrogallol, sulphuric acid and water. The alkali of the same concentration as the salt solution was added in 1-2 cc portions. This procedure provides a constant volume in


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each titration although the concentration changes, and in each case 30 cc of alkali is equivalent to 3 NaOHjAl. The saturated calomel half-cell was used with a potassium chloride bridge, the value for this being taken as

FIQ.3 0.245 volts a t

25'. All determinations were made at room temperature and corrections made for z 5 O - 3 0 ~ . Although neither Blum nor Hildebrand mentions removing the electrode during the alkali addition, it appears that both must have done so. Permitting the electrode to remain in the solution during the addition of the

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alkali causes the deposition of a thick coating of adherent alumina on the electrode and this gives very unsatisfactory results. Typical curves for such titrations are shown in Figs. 2 and 3 for 0.3 N AIC13 and 0.3 N Alz(S04)s, approximately the concentration used by Hildebrand and Blum. I n these curves it is observed that on adding alkali to the aluminum salt solutions, the apparent hydrogen ion concentration of the solution actually increases. This may bring about a reversal of the cell and the calculations show that in the case of the sulphate the apparent hydrogen ion concentration of the solution is of the order of IOO N, a perfectly fantastic figure. This represents striking poisoning of the electrode by the alumina which is coagulated upon it. Originally this alumina is highly adsorptive and as formed carries on to the electrode with it adsorbed hydrogen ions in the hydrous sheath. This however would not account for a continued increase in the apparent hydrogen ion concentration of this adsorbed layer and furthermore it is, in the present state of our knowledge, questionable whether adsorbed hydrogen ions are capable of affecting a hydrogen electrode. For the present no full explanation can be given of this phenomenon and it appears to be comparable to the ordinary poisoning processes already known about. This may have in this case a mechanical action by which the continuously increasing film of alumina prevents the access of the hydrogen gas to the electrode and therefore the setting up of the equilibrium between molecular and atomic hydrogen upon which the action of the electrode depends. It may be merely that the alumina blocks off this activating adsorption of the hydrogen on the platinum. Whatever the mechanism may be it is clear that the observed poisoning effect will be calculated as an apparent high hydrogen ion concentration in such systems. Actually of course the free solution is much more alkaline than the electrode shows, for phenolphthalein changes color when about 80-9oyo of the theoretical alkali has been added although the apparent pH by the electrode is still far on the acid side of pH 7. The continued addition of alkali has a peptizing action on the alumina layer so that in the case of the chloride a pH value of 7 is shown by a little more than 4 NaOH/Al although phenolphthalein turns pink a t about 2 . 7 NaOH/Al. I n the case of the sulphate solution the peptization of this layer appears more difficult and even 5 NaOH fails to produce a solution which is alkaline according to the electrode. In view of these gross abnormalities the procedure was adopted of withdrawing the electrode from the solution during the addition of the alkali although stirring was continued by means of the hydrogen. Two to five minutes were allowed for the precipitation of the alumina, the electrode was replaced and the equilibrium reading obtained. This method is a great improvement over the first one and is apparently the one used by Hildebrand and by Blum since it gives curves similar to theirs. It is, however, not a full solution to the problem for even thus a film of solution adhering to the electrode contains some alumina which tends to form a light film an the electrode when the latter is reimmersed in the solution which has in the meantime become more alkaline. This light film should have some such

log1


effect as the heavier one has been shown to have so that the curves given may be a t a somewhat lower pH than the solution itself. That this error is a very slight one is shown by the fact that the readings of the electrode agree in most of the cases with the first appearance of the pink color of phenolpht,halein which was added to all titrations. I t is quite probable that the deviations increase with departure from neutralit'y and that a better way to study this system would be by portions made up separately and permitted to come to equilibrium before the electrode is used. The values recorded are approximately correct and are comparable among themselves.

Titration pf

TABLEV Aluminum Chloride with Sodium Hydroxide N A1C13

Original solution Phenolphthalein showed pink At 3 NaOH/Al Break points At 5 NaOH/Al

XaOH

pH

o

2.3 6 .4 8.4 10.7

25 . 2

26.5 30 34.5 50

0.3 W AlCh

NaOH 0

pH

2.9 5.8 8.4 9.7

11.0

26.3 27.4 30 33

I O ,I

12.7

50

12.2

0.1

N AlCla

NaOH 0

27.0

27.75 30 50

pH

3 .o 7'5 8.5 9.3 11.7

TABLE VI Titration of Aluminum Sulphate with Sodium Hydroxide NA~~(SOI)I XaOH pH

Original solution Phenolphthalein showed pink At 3 NaOH/Al Break points At 5 NaOH/Al

o 26 27

30 37.5 50

2.8 7.4 8.4 10,4 11.2

12.6

0.3 N A12(S04)3 NaOH pH 0

25.2 26.3 30 37.5 50

3 .o 6.8 7'9 10.1

10.9 12.4

0.1 N Al*(SO,) NaOH pH

0

26.3 27.4

3.4 7'5 8.6 9.6

30 37

10.2

50

11.8

In Tables V and VI are shown those dat'a from the two titrations which are significant, including under each concentration of salt the volume of alkali added (of the same normality as the salt) and the pH value observed. The pH of the chloride solutions is lower than that of the corresponding sulphate solutions and in both the acidity increases with concentration. Next are shown the volumes of alkali and t,he pH readings before and after the phenolphthalein showed its first pink color. I n all cases save one these pH values are on opposite sides of pH 8.3 the lower end of the phenolphthalein range. This checks the correctness of the electrode reading a t this point and since in no case is more than 27.7 j cc of alkali required to pass the pH 8.3 this shows that the Hildebrand and Blum assumptions that all the alumina was titrated a t pH 7 is a t least ten per cent in error. This is confirmed by the pH values observed when 3 NaOH is actually added in each system,

1072


the lowest pH being 9.3, which is quite different from 7 . As break-points are shown those volumes of alkali and the corresponding p H at which the curves appears to change direction in the irregular portion and finally are shown the pH values obtained on addition of 5 NaOH/Al. The tables should be compared with the curves in Figs. 2 and 3 where the complete systems are shown. The first item of interest is that there is no break in these curves corresponding to those found by Hildebrand and by Blum at pH 3 and which we have explained as being due to the free acid present. Further there is no irregularity in these curves which might be taken as the beginning of precipitation of alumina. This is difficult to observe under the conditions of the experiment but ppproximately four to five cc of alkali gave precipitates in the sulphate solutions which did not clear while the electrode was coming to equilibrium while in the chloride solutions three times this amount of alkali gave a temporary precipitate which cleared before the next addition of alkali. This is in accord with what would be expected and with the 20 and 80% additions when time for equilibrium was allowed. It is, therefore, somewhat hazardous to state definitely just what amount of alkali gave a visible precipitate of the hydrous alumina; it is clear that there is a marked difference between chloride and sulphate in this repect. At any rate the agglomeration of the alumina to particles large enough to see is probably not a unique point but merely a stage in a continuous process. The courses of the two sets of curves are quite similar, both being fairly flat between pH 3 and 4 until 2 NaOH has been added when there is a sharp but not a vertical ascent ceasing shortly before 3 NaOH/Al is reached. If any point were to be taken as marking the end of the precipitation of the alumina, it might be a t 2 NaOH for there the flat ceases. But common experience in determining alumina precludes the acceptance of this point. The gradual increase in pH with alkali addition on the steep portions of these curves confirms the observations of Clark and distinguishes these curves from a true titration curve. I n all these aluminum salt curves the change from pH 4 to pH I O requires approximately I O cc of alkali while a similar change in the titration of hydrochloric acid requires only a drop or two. This is, of course, a type of buffer action in which the precipitated alumina, having adsorbed the hydrogen ion gives it up slowly as the solution becomes alkaline. This buffer action is particularly marked in the acid range and if one did not mind the alumina present, aluminum chloridesodium hydroxide buffers might be made for the range from pH 3 to 4. The adsorption of the acids by alumina makes the systems titrate as relatively weak acids instead of the strong acids HC1 and HzS04 actually are. The curves also show graphically that it is impossible t o titrate these systems cold using phenolphthalein or any other indicator since there is no vertical portion of the curves a t 3 NaOH/Al. Heating swings the lower end of the steep portion of the curves to the right and makes the curves approximately vertical a t the pH 8.3 of phenolphthalein, since heating reduces the acid adsorption. This is all in agreement with the Kolthoff method of titrating these salts.

TITRATION CURVES FOR ALUMIKUM SALTS WITH ALKALIES

I073

The most interesting point in connection with these curves is that they show break-points such as were reported by the earlier authors at pH I O to pH 11, and further that these points do coincide roughly with the apparent final disappearance of visible particles of alumina. There is some reason to believe that this may be an error in observation and this point will be studied further. The data show, however, that these points come at about 34 cc of alkali for the chloride systems and about 3 7 cc for the sulphate solutions. These break-points differ in pH wit,h the concentration of the salt and with the anion present, although the volumes of alkali needed for the various concentrations are approximately the same for each of the two salts separately. The volumes of alkali used are about 34 cc for the chloride solutions and 37 cc for the sulphate solutions and by no stretch of the imagination could these be taken as 4 NaOH/Al. In these curves too the volumes of alkali needed to bring the systems to pH 7 and then to this second break are roughly in the 3:1 ratio but the actual values are for the chloride system 2 . 6 : 0 . 8 NaOH/Al and for the sulphate 2.6:1.1 NaOH/Al. The failure of these points to come out in simple stoichiometric proportions removes the last support for the theory of aluminate formation in these alkaline solutions. The explanation then is to be sought in a peptization of the alumina by the hydroxyl ions. The tendency to a flattening of the curve in the alkaline solution shows that the added alkali is disappearing from the solution by being adsorbed on the alumina until the latter is dispersed and saturated with hydroxyl ion. Then further alkali addition builds up a free hydroxyl ion concentration in the solution comparable to that obtained in the presence of sodium chloride only and the curves assume the shape of the HC1 curve. If there were merely adsorption of the alkali on the alumina and one measured the pH of the intermicellar liquid, the pH of the systems containing alumina should be lower than that of the HC1-NaOH titration in the alkaline range. The more alumina present should adsorb more alkali so that the curve for the normal solution should come lowest of all. The observed order is the reverse of this and the normal solution is as strongly alkaline as the system containing no alumina. This is merely another result of the deposition of a film of alumina on the electrode. In the alkaline solution this film by adsorption of hydroxyl ion makes the electrode appear to be in an abnormally alkaline solution and this error is greater in the more concentrated solutions. The thinner layers of alumina are more easily peptized from the electrode. Up to the break-points the hydroxyl ions are peptizing the outer layers of alumina on the electrode and then the sharp upturn denotes the adsorption of these ions on the lower layers until these become saturated. The marked difference in t'he shapes of these upper irregular portions of the curves is further evidence that we are dealing with peptization and not compound formation. The sulphate alumina is easier to precipitate and harder to peptize than the chloride alumina and this is mirrored in the greater change of direction of the sulphate curve. The gradual disappearance of this irregular portion with dilution is significant and explains why Clark and Miller did not observe it in their very dilute solutions.

I074


summary

The alkali titration curves of Hildebrand for aluminum sulphate and of Blum for aluminum chloride have been considered as the best available evidence for the formation of aluminates in aklaline solutions. 2 . These titrations have been repeated and the general features of the curves confirmed except in so far as the earlier work was done on salt which contained free acid. 3. Such titrations have been made also for more dilute and more concentrated solutions and it becomes evident that the interpretations previously offered are not valid. The breaks in the curve are not unique for alumina alone but are influenced by the conditions. 4. There is no break or irregularity in the titration curves corresponding to the beginning of the precipitation of alumina. Some colloidal alumina is present in the original solutions and its quantity and aggregation increases steadily with alkali addition. Visible and filterable particles of alumina appear a t markedly different alkali additions which depend on the concentration, the temperature, and the anions present. 5 . I t is shown that aluminum salts can be titrated with alkali if boiled, and that salt addition is not necessary and is in some cases actually bad practice. 6. I t is shown that aluminum salts cannot be titrated cold by adding sufficient alkali to bring the system t o pH 7 as was believed by Hildebrand and Blum. The volumes of alkali used for this and for the break in the curves at pH I O to I I are not in the simple ratio of 3 and 4 NaOH per atomic weight of aluminum and therefore lose all value as a proof of compound formation in the excess alkali. 7. Since the beginning or the end of the precipitation of alumina produces no inflection points in these curves their similarity of form is to be ascribed to the common reaction occurring, namely the neutralization of the hydrochloric or sulphuric acids as this process is affected by their adsorption on the alumina. 8. I t is shown that alumina deposited on the hydrogen electrode can cause very large errors in the readings obtained. 9. This paper shows that the action of alkalies on alumina is one of dispersion and that this process is affected by such factors as concentration and particularly by the difference between sulphate and chloride ions which should affect the adsorption of the peptizing hydroxyl ion. I.

CbTndl Universzty.

TITRATION CURTrES FOR ALUJIINUM SALTS ... - ACS Publications

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