THE MECHANISM OF FORMATION OF ORANGE I1 ALUMINA COLOR LAKES ANDRIES VOET Research Department, Sun Chemacal & Color Company, Dauiszon General Printang Ink Corporation, Harrison, New Jersey Received November 6 , 1942 I. INTRODUCTION
Orange 11, a water-soluble dye (Color Index # 151, Schultz tables #86) is the monosodium salt of 1-p-sulfophenylazo-2-naphthol. I t is formed by diazotization of sulfanilic acid and coupling the diazo compound with @-naphthol. When a solution of Orange I1 in water is mixed with an aqueous alumina hydrate gel a color lake may be formed. The mechanism of the formation of this color lake, the Persian Orange lake, has been the subject of several investigations. Reinmuth and Gordon (4) proved that the precipitate formed upon addition of an aqueous solution of an aluminum salt to a solution of Orange I1 in water is a definite salt of the composition A&, where X represents the univalent Orange I1 acid radical. The same authors also proved that upon interaction of the free Orange I1 acid (a very soluble and comparatively strong acid) and alumina hydrate the compound A1X3 is also formed. This observation was confirmed by Bancroft and Farnham (1). These latter authors, however, also studied the interaction between the sodium salt of Orange I1 and alumina hydrate and concluded that the mechanism of the interaction must be considered as a pure adsorption. These conclusions seem to be directly contradictory to the results of investigations of Marker and Gordon (3), who studied the amount of dye taken up by alumina hydrate as a function of the acidity of the solutions. A sharp break was found to exist a t a pH of 3.0, where a tremendous increase in the quantity of dye taken up by the hydrate was observed. The conclusion was drawn that a chemical reaction occurred which was responsible for the lake formation. Weiser and Porter (B), in earlier investigations, questioned the existence of a chemical compound A& and attributed the formation of the Persian Orange lake to pure adsorption. In subsequent publications Weiser ( 5 ) conceded the existence of a definite alumina salt of Orange I1 acid. He still was inclined, however, to believe that the Persian Orange lake formation proper mas a process of adsorption, rightly pointing out that the formation of a definite salt by interaction of aqueous solutions of aluminum chloride and the sodium salt of Orange I1 does not prove that the same type of reaction does necessarily occur by interaction of the sodium salt of Orange I1 and alumina hydrate. The following study is part of an investigation as to the mechanism of the interaction between Orange I1 and alumina hydrate. 191
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AXDRIES VOET
(1) Orange I I : The Orange I1 dyestuff used was a purified commercial product. The ivater content \\-as found to be 2.8 per cent by measuring the loss of weight on drying in vacuum until constant weight at room temperatures. The ash content, \vas the amount theoretically required for a pure product. ?io foreign metals were present. Since Orange TI solutions are readily decomposed on standing, fresh solutions were made up each time. ( 2 ) dlumiim Iiudrate: Alumina hydrate \vas prepared by precipitating a 0.3 N solution of aluminum sulfate with a 0.3 N solution of sodium carbonate at 70°C. The precipitate n a s filtered and washed n.ith distilled water until addition of barium chloride to the filtrate did not cause an opalescence. The filter cake was reslurried and rewashed, and this procedure was repeated again. Finally a suspension was made containing 37.2 g. of A 1 2 0 3 per liter, and was allowed to age for several months at room temperature. A sodium analysis of the alumina hydrate was carried out. A known volume of the suspension vas evaporated and heated to red heat in a quartz crucible to constant weight. The contents of the crucible was treated with sulfuric acid and once more h&ed to red heat. An extraction wa8 carried out xith boiling water and sodium was estimated in the extract as sodium zinc uranyl acetate, according to the method of Barber and Kolthoff (2). An amount of sodium corresponding t,o 0.41 milliequivalent per gram of mas found. Sulfate was determined by dissolving a known volume of the suspension in hydrochloric acid and precipitating the sulfate as barium sulfate. An amount of 1.57 milliequivalents of sulfate per gram of A1203was found. The pH of the alumina hydrate suspension was G.50. The particks of the suspension showed a positive charge, as \vas indicated by t,heir electrophoretic motion in a U-tube. (3) Chemicals: All chemicals used were of C.P. quality.
B. dletkod Dye solutions were made up by neighing the required amount of dye and adding distilled ivnter up to a given volume. Lakes were prepared by adding a measured volumr of the dye solution to a measured 1-olume of the alumina hydrate suspension. The solutions in Pyrex bottles sealed with rubber stoppers were placed in a thermostatic shaker filled with m t e r of 40°C. =!= 0.1". After the period of interaction the solutions were centrifuged in an International clinical centrifuge for 15 min. a t approximately 2000 R.P.U. A measured volume of the top layer was removed carefully with a pipet and recentrifuged, generally after dilution with a measured volume of water. All solutions were examined in a strong light, and analyzed only in the complete absence of suspended pa.rticles. 1)etermination.; of dye concentrations weye carried out by means of a HelligeDuhoscq optical colorimeter which \vas adjuated daily by comparing Polut ions containing known quantities of dye.
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Estimations were made of the concentration of the dye remaining in solution after removal of the lake by centrifuging. The quantity of dye taken up by the alumina hydrate was found by subtracting these estimated quantities from the amount of dye originally added to the solution. The pH of the solutions was measured with a Coleman glass electrode pH meter, which was standardized with a Coleman standard buffer solution. Measurements were made in the liquid containing dyestuff and lake suspension after the end of the period of interaction. Consequently, all reported pH values refer to the final stage of suspension. Adjustments of the p H were made by addition of the required quantities of hydrochloric acid or sodium hydroxide.
1
1000
500t
250c pH-7.16
PH
WE CONCENTRATION IN PER CENT
FIG.1 FIG.2 FIG. 1. Amount of dye taken up by the alumina hydrate plotted against the final concentration. FIG. 2. Amount of dye taken up by the alumina hydrate plotted against the pH. , data of the author; ------------- , data of Marker and Gordon.
C . Data ( 1 ) Influence of the dye concentration: The influence of the dye concentration
was studied first. To the alumina hydrate suspension equal volumes of dye solutions of the different concentrations were added, the pH of which had been previously adjusted to a standard. The amount of dye taken up was estimated after 72 hr. of interaction at 40°C. The result is given in figure 1, where the amount of dye carried down by the alumina hydrate is plotted against the final concentration for solutions of a pH of 7.16,as well as for solutions of a pH of 6.40. It was found that in solutions of-a fina;l’PHof 5.5, as well as in more acid solutions, all the dyestuff present i s a r r i e d down by the hydrate independent of the initial dye concentration. (2) Injluence of the pH: Although the importance of the p H may be observed from the curves of figure 1, the influence of the p H on the amount of dye taken
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ASDRIES
Tom
up by the hydrate is plotted separately in figure 2 for a solution containing 1.90 per cent of Orange 11. It was found that only a slight amount of dye is taken up in neutral and basic solutions, while a tremendous increase is observed in more acid solutions. The data of Marker and Gordon (3) are plotted in figure 2, indicated by a dotted line. Aside from the sharp break in this curve, vhich might be attributed to lack of data in the critical range, the general character of both curves shows a certain similarity, although a shift over a range of 4 pH units is observed. 111. DISCUSSION
When we consider the curve of figure 1 measured at a pH of 7.16, it becomes evident that the shape of the curve is that of a regular adsorption isotherm. Consequently the mechanism of interaction betxeen Orange XI and alumina hydrate at a pH of 7.16 is adsorption. When we consider the curve of figure 1 at a pH of 6.40, a more complicated picture is observed. The quantity of dye taken up by the alumina hydrate, ho\Tever, may be formed by superposition of a constant quantity independent of the concentration, and a variable quant'ity which follow a regular adsorption isotherm. Finally, at still higher acidity, it is observed that all the dye present in the solution is carried down by t,he hydrate. This rather complicated picture becomes comparatively simple when it is borne in mind that precipitated alumina hydrate is readily dissolved in acid. Consequently we may expect that a t a pH beloiv 7.0 the hydrate will dissolve. The speed of this process is greatly increased a t greater hydrogen-ion concentrations. Thus a very substantial increase of the quantity of hydrate dissolved per unit of time occurs on lowering the pH. As a result of the dissolution of alumina hydrate, aluminum ions are formed which are able, as previously mentioned, to form the insoluble aluminum salt of Orange 11. Consequently, a t a pH directly below 7.0 the amount of dye carried down by the hydrate is made up by superposition of a precipitated and an adsorbed part. At a still higher acidity, where a large quantity of aluminum ions is present, practically all the dyestuff carried down will be in the form of the aluminum salt,. This range was found to start at a pH of 5.5 for our alumina hydrate. At a pH above 7, in the absence of precipitating ions, the entire mnount taken up is adsorbed. More proof of these opinions is furnished by further experiments. In a separate blank experiment, under conditions exactly similar to those of the experiments reported in figure 1, an estimation was made of the quantity of alumina hydrate dissolved by the amount of hydrochloric acid added to reach the final pH of 6.40. It was found that this quantity, calculated per gram of Also3,was able to precipitate 560 mg. of Orange I1 dye; this corresponds exactly to the point of the curve extrapolated to zero concentration. Thus the meaning of figure 2 becomes clear. At a high pH, above 7.0, the interaction of hydrate and dye is a pure adsorption. Directly below the neutral point, chemical reac-
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tion and adsorption are both observed, while a t lower pH the chemical combination becomes rapidly dominating, owing to abundance of aluminum ions. Although the basic part of the curve of figure 2 represents a true adsorption equilibrium, the acid part does not. Upon prolonged interaction more hydrate will dissolve and more dye will be precipitated. The theoretical limit is reached only when all the excess acid is used up. On the other hand, if the type of hydrate used is less reactive, it will dissolve only when a higher acid concentration is present and, although the general character of the curve is maintained, a shift in pH towards the acid side must occur. The data of Marker and Gordon (3) given in the dotted line might be explained by this c0nsideration.l Physically and chemically a very distinct difference is observed between “adsorbed” lakes formed a t a pH above 7.0, and “precipitated” lakes formed in the presence of aluminum ions a t a pH lower than 7.0. Adsorption complexes have a yellowish shade, while precipitated lakes have a more reddish shade. On prolonged washing with distilled water of 25OC. the filtrate of a precipitated lake was only very slightly colored and had a pH of 5.0, while under similar conditions the filtrate from the adsorption complex was markedly colored and had a pH of approximately 4.0. Since the solubility of the aluminum salt of Orange I1 is very slight, the precipitated lake does not lose much of its dye on prolonged washing, while the absorbed lake is rapidly decomposed. Finally, a color lake was prepared which was known to contain only the aluminum salt of Orange 11. This was done by first precipitating the aluminum salt of Orange 11, by adding an equivalent quantity of aluminum chloride to a solution of Orange 11, and then mixing the wet precipitate with an alumina hydrate gel. The resulting product could not be distinguished physically or chemically from a color lake made by interaction of Orange I1 and alumina hydrate a t a pH belov 5.5. A further study revealed the type of adsorption occurring a t a pH above 7.0. Analysis of adsorption lakes showed that no sodium mas present in the lake. This observation eliminates a direct adsorption of the dye and leaves as an explanation of the adsorption phenomena either an exchange adsorption or a hydrolytic adsorption. Hydrolytic adsorption, however, must also be disregarded, since no change in pH is observed as result of adsorption. Consequently, exchange adsorption is the only remaining possibility. This is proved directly by analysis, which shoiyed that sulfate ions previously adsorbed by the hydrate are released upon adsorption of Orange I1 acid, in such a way that one sulfate ion is exchanged for every two Orange I1 ions. It is interesting to note that the apparent contradictions mentioned in the literature are non-existent. Marker and Gordon (3) rightly claimed a chemical combination, since their experiments were made in acid solutions, while Bancroft and Farnham (1) were justified in claiming an absorption, since it may be deduced from their reports that they studied neutral or slightly basic solutions. 2 A source of errors is very likely the use of the rather doubtful Bayliss electrode by Marker and Gordon in 1924 to measure pH values. The modern glass electrode is much more reliable.
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M. E. LAIXG MCBAIh' IV. SGYU.4RY
The interaction of Orange I1 and alumina hydrate is essentially an exchange adsorption, where one sulfate ion is exchanged for two dye anions. When aluminum ions are present, as may be the case at a pH below 7.0, the aluminum salt of Orange I1 is precipitated. Orange I1 lakes prepared a t a pH below 5.5 with the type of alumina hydrate described previously are pure aluminum salts, while lakes prepared a t a pH between 7.0 and 5.5 consist of mixtures of an adsorption complex and a chemical combination. " The author wishes to express his thanks to the Sun Chemical PE Color Company, Division General Printing Ink Corporation, for their permission to publish this investigation. REFERENCES (1) BANCROFT ASD F ~ R S H AJ. MPhys. : Chem. 36, 3127 (1932). (2) BARBER AND K O L T H O F F : J. h n . Chem. S O C . 60, 1625 (1928). (3) MARKERASD GORDOS:Ind. Eng. Chem. 16, 1186 (1924). (4) RCINYUTH A N D GORDON: Colloid Symposium Monograph 7, 161 (1930). (5) WEISER:The H y d r o u s Oxides. John Wiley and Sons, Inc., Yew York (1935). (6) WEISERA N D PORTER:J. Phys. Chem. 31, 1704 (1927).
MIGRATION DATA I N SOLUTIONS OF A COLT,OIDAL ELECTROLYTE: LAURYLSULFONIC ACID' hl. E. LAING McBBIN Department of Chemistry, Stanford University, California Received December 7, 1948
Laurylsulfonic acid has proven to be the most satisfactory colloidal electrolyte for exact experimentation. It is a freely soluble, simple, uni-univalent, straight-chain, free sulfonic acid with twelve carbon atoms, and it canr:ot hydrolyze. Its properties are closely similar to those of all straight-chain colloidal electrolytes. Indeed, A. P. Brady (work not yet published) has shown that the osmotic properties of all the membws of all kinds of straight-chain compounds, anionic or cationic, can be superimposed on one single curve. Most of the simple unambiguous quantitative methods of classical physical chemistry have now been applied to laurylsulfonic acid. The chief one lacking has been a study of the electrolytic migration, N hich is here supplied. The results are highly interesting in themselves, in as much as the transport number, n, exhibits a sharp maximum in moderately low concentration (0.055 N ) . The data yield direct information as to the properties of the carriers of the electric current in solutions of colloidal electrolytes. 1 Presented at the Sineteenth Colloid Symposium, which Ras held a t the University of Colorado, Boulder, Colorado, June 18-20, 1942.
