T H E PHYSICAL CHEMISTRY OF COLOR LAKE FORMATIO”.
I. GENERAL PRINCIPLES BY HARRY B. WEISER AND EVERETT E , PORTER
The adsorption of many dyes by wool, silk, and cotton is SO weak that they are of value to the practical dyer only when used in connection with mordants. The term mordant (from mordre, to bite or to corrode) was first applied by the French to metallic salts which were supposed to act by biting or opening up a passage into the fibers of the cloth giving access to the color. Thus alum was believed to be effective in fixing certain dyes, owing to the solvent or corrosive action of the sulfuric acid. I t is now known that the real mordant is the hydrous oxide and not the acid derived from the salt. In general, a mordant may be defined as any substance that is taken up strongly by the fiber and in turn takes up the dye strongly. Two classes of mordants are generally recognized; acid, and basic or metallic. The acid mordants are the tannins, the fatty acids, albumin, hydrous silica, arsenic acid and phosphoric acid, while the basic mordants are the hydrous oxides of the heavy metals, the most important of which are the hydrous oxides of chromium, aluminium, iron, tin, and copper in the order named. In dyeing a mordanted cloth it is, in most cases, the mordant rather than the fiber which takes up the dye. When a dye is taken up by a mordant in the absence of a fiber, the complex is called a color lake. This investigation deals with the mechanism of the formation of color lakes of the hydrous oxides of iron, chromium and aluminum as lake bases. For a long time all such color lakes were believed to be definite chemical compounds between the dye and the mordant. This view was questioned by Biltz and Utescher’ who investigated the behavior of alizarin with the hydrous oxides of chromium and iron. With the former, the amount of dye taken up increased continuously with increasing concentration of solution, giving no indication whatsoever of the formation of chromium alizarate. With hydrous ferric oxide, on the other hand, the amount of dye taken up increased so rapidly with relatively small changes in the concentration of the bath, that Biltz was lead to assume the formation of a ferric alizarate. More recently Bull and Adam9 investigated the iron-alizarin lake and concluded that it is a ferric oxide-sodium alizarate complex. Marker and Gordon3 studied the influence of hydrogen ion concentration on the amount of the basic dyes, crystal violet and methylene blue, and the acid dyes, orange I1 and metanil yellow, taken up by ferric oxide, alumina, and silica. In most instances, breaks were obtained in the pH-adsorption curves which were interpreted to mean that chemical compounds were formed between the hydrous oxides and ‘Ber., 38, 4143 (1905).
* J. Phys. Chem., 25, 660 (1921). Ind. Eng. Chem., 16, 1186 (1924); 15, 818 (1923).
I384
HARRY B. WEISER AND EVERETT E. PORTER
the dyes investigated. As we shall see, the value of Marker and Gordon’s paper would seem to lie chiefly in their recognition of the importance of hydrogen ion concentration in color lake formation rather than in the data submitted or the conclusions drawn therefrom. I n the light of our observations( on adsorption of ions from a mixture of electrolytes, it would seem that the concentration of the ions present in the dye bath, including hydrogen and hydroxyl, might be quite as important in determining the composition of a color lake as the concentration of the dye itself. Moreover, the nature, physical character, and purity of the mordant would be expected to influence the composition of the lake. In order to determine the effect of each of the several factors on the lake-formation process, the first investigations were made on systems in which all the variable constituents could be determined quantitatively. T o simulate the conditions of lake formation as nearly as possible, observations were made a t varying hydrogen ion concentrations on the taking up of the inorganic anion, sulfate, and the organic anion, oxalate, both separately and simultaneously, by highly purified hydrous oxides of known composition. The results of these observations are recorded in this communication. In the subsequent papers the observations will be extended to the true color lakes. Experimental Procedures T h e Adsorbent. In most of the earlier investigations on adsorption of the dyes by the hydrous oxides, the latter were precipitated from the respective salt solution with ammonia, washed by decantation, and dried. A weighed amount of the adsorbent prepared in this way was shaken with the dye and the amount taken up was determined from the change in the concentration of the solution. This method of procedure is open to certain objections: first, it is impossible to obtain equilibrium conditions in a reasonable length of time; second, it is impossible to get absolutely uniform samples; and third, the purity of the oxides is uncertain. To obviate these difficulties, the hydrous oxides used in the present investigation were prepared in the form of sols. This mas done by precipitating the gels from the rp -ctive chloride solutions with ammonia, washing by decantation with the aid of the centrifuge until peptization began, and finally, peptizing with the smallest oossible quantity of hydrochloric acid. The sols formed in this way were Cidyzed in boiling solution for approximately one hundred hours according to the method of Xeidle,* using a five-liter balloon flask to replace the open beaker. The rate of flow of water was kept at about two liters per hour. By this procedure, products were obtained having a much higher degree of purity than the precipitated gel. Determinatzon of Adsorptzon. Two methods may be employed for determining the adsorption by the g d thrown down on the addition t o the sol of a given amount of electrolyte: the gel may be washed by the aid of the centri1 2
R’eiser: J. Phys. Chem., 3 0 , 20, I j27 (1926). J. Am. Chern. SOC.,39, 7 1 ( 1 9 1 7 ) .
PHYSICAL CHEMISTRY OF COLOR LAKE FORMATION
I385
fuge and the analysis made on the supernatant liquid with all the wash water; or an aliquot part of the supernatant liquid may be pipetted off and analyzed, calculating the adsorption from the change in concentration. From the results of a number of preliminary experiments, it was found that the latter method gives more reliable and consistent results. The errors eliminated by using this method are the differences in the conditions of washing and the variation in the reversibility during washing. The latter factor is especially important since the adsorption is reversible in all cases, equilibrium between the wash water and the precipitate being established more or less rapidly. The error introduced by the failure to take into account the volume occupied by the solid is kept small, amounting to less than one tenth of one percent in most cases. The fact that this error is constant in any series of experiments, reduces still further any objection to the method. The mixing of the hydrous oxide sol with the electrolyte was done in 250 cc wide-mouth bottles with ground-glass stoppers. The solution to be mixed with the sol was put in the bottle and diluted to exactly 150cc. Fifty cubic centimeters of sol containing the desired amount of hydrous oxide, was placed in a 60 cc bulb with a large neck inserted through a rubber stopper which fitted the bottle. I n order to retain the sol when the bulb was inverted, the neck of the latter was fitted with a thin slice of rubber cut from a stopper. To mix the solutions, the stopper with the bulb was inserted in the bottle which was given a quick jerk thereby dislodging the thin stopper from the neck of the bulb and allowing very rapid and thorough mixing of the solutions. Numerous tests showed that equilibrium between the adsorbent and the adsorbate was established almost at once. However, the mixtures were allowed to stand in the glass-stoppered bottles for about 2 4 hours after which they were centrifuged and an aliquot part withdrawn for analysis. The hydrogen ion concentration was determined in the sJpernatant liquid as well as in a similar mixture in which water was substituted for the sol. The former determinations are designated in the tables as “pH after adsorption” and the latter as “pH before adsorption.” Since the sol is approximately neutral, if the small error introduced by the space occupied by the solid particles of the oxide is neglected, the difference between the two pH values is the change due to the adsorption by the colloidal particles.
Hydrogen Ion Concentratzon. The hydrogen ion concentrations were measured throughout with the hydrogen electrode. The apparatus employed was the Leeds and Northrup “Type K” potentiometer in connection with one of their sensitive galvanometers. A small electrode vessel of special design requiring but 4 cc of solution for a determination, was used. The gas entered the electrode vessel at the bottom through a very small capillary, a pressure of seven to eight centimeters of mercury being required to cause a flow of one cubic centimeter of gas each three seconds. This served to agitate the liquid very effectively without blowing it out of the vessel. The electrodes were prepared by platinizing 4-millimeter glass tubes to a height of I centimeter. These electrodes come to equilibrium more quickly than the platinized
1386
HARRY B. WEISER AND EVERETT E. PORTER
platinum electrodes and are easily and economically renewed. The hydrogen used was purified by conducting it over hot platinized asbestos, bubbling through potassium hydroxide, sulfuric acid, and finally through water at the same temperature as the electrode vessel.
The Adsorption of Sulfate Adsorption Experiments. The adsorption of sulfate by a constant amount of chromic oxide was determined at varying hydrogen ion concentrations, according to the procedure described above. In all the experiments, the initial sulfate concentration was maintained constant at 0.005 normal, while the pH value was varied through wide limits using suitable proportions of sulfuric acid, potassium sulfate, and potassium hydroxide. Obviously, this procedure introduces varying amounts of potassium ion as well as of hydrogen and hydroxyl ions. This effect may be considered as negligible, since it was found in some experiments to be recorded in a subsequent paper, that very little or no alkali ion is adsorbed under similar conditions. If the alkali adsorption were appreciable, it would serve to increase the adsorption of the negative ions. It should be pointed out that the same initial concentration of sulfate leads to a varying equilibrium concentration. Since the final concentration is smaller the greater the adsorption, it follows that the variations in the adsorption values would have been somewhat larger than those observed if the final concentration had been kept constant. However, the final concentration is sufficiently large in all cases to put the adsorption well on the flat of the isotherm. The marked effect of the hydrogen and hydroxyl ions on the adsorption of sulfate is shown by the data recorded in Table I and shown graphically in Fig. I . It will be seen that raising the concentration of the hydrogen ion increases the positive charge on the particles, thereby increasing their capacity TABLE I Adsorption of Sulfate by Hydrous Chromic Oxide at varying pH Va!ues Cc of s o h . mixed with go cc of 0.125grams of CrlOa in a total of 200 cc
sol containing
pH values Adsorption values milli-equiv. Before After per gram Cr203 adsorption adsorption
N/go H I S O l N/go KISOl X/50 KOH
30
20
0
20
30
0
10
40
0
0
SO
0
0
SO
S
0
SO
IO
0
50
20
4.66 4.60 4.52 4.28 3.48 I . 90 1.28 0.48 -0.08
0
50
30
-0.
SO
0
0
40
IO
0
IO
2.42 2.49 2.s9 2 ' 73 3.09 8.68 10.61 10.9s 11.28 11.4-
2.87 2.97 3.1s 3.64 5 . IS 8.29 8.73 8.96 9.39 9.87
PHYSICAL CHEMISTRY O F COLOR L A K E FORMATION
I387
to adsorb sulfate. The adsorption of the sulfate falls off quite rapidly with increasing concentration of the hydroxyl ion, the carrying down of the sulfate being completely nullified a t a pH value of about 9.2. I t is obvious however that both sulfate and hydroxyl are adsorbed in the alkaline range from a pH of 7 to 9. In Loeb’s investigations on the proteins’ it seems probable that he overlooked the specific effect of cations other than hydrogen and of anions other than hydroxyl since he worked with relatively low concentrations of salts.
The pH-adsorption curve is smooth throughout, showing no indication of the formation of a compound between the sulfate and the hydrous oxide even at pH values as low as 2.0.
TABLE I1 Precipitation of Chromic Oxide Sol with Sulfate at varying pH Values Cc of eoln. mixed with 5 cc of sol containing 0.0125grame of Ct,O, in a total of 20 cc X/~O HzSO, N / ~ oK2S04 S / ~ KOH O
pH values Before adsorption
2.50
0.0
0.0
2.50
2.92
3.16
0.2
0.0
2.50
3 .oo
3.30
2 . IO
0.4 0.6 0.8
0.0
2.50
0.0 0.0
2
3.07
3.38
2 .
3.18 9.35 I O .67
3.74 5.61
.85
I . 50 1.25 0.00
0.9 1.1
0.0 0.0
0.00
0.0
I
.o
’
45
2.30
I5
I . IO I
.oo
“Proteins and the Theory of Colloidal Behavior”
(1922).
-
After adsorption
2.30
I
1
Precipitation values milli-equiv. per liter of mix
6.85
1388
HARRY B. WEISER AND EVERETT E. PORTER
Precipitation Erperiments. Some observations on the precipitation value of sulfate ion at varying pH values were next made. 5 cc of the sol was placed in the inside compartment of a mixing apparatus' made entirely of glass, and varying amounts of potassium sulfate, sulfuric acid, and potassium hydroxide diluted to exactly I 5 cc were placed in the outside compartment. These solutions were mixed rapidly, allowed to stand for about 2 4 hours, and centrifuged to determine the presence or absence of complete precipitation. The pH values were determined for the supernatant liquid as well as for like mixtures substituting water for the colloid. The results are given in Table 11 and represented graphically in Fig. 2 . L
c
IO
-1
k c
8
-n z
3 v
-' 3
6PH
2
f : 2
4
421
2
-
$
cf
I t should be noted that the precipitation value for potassium hydroxide alone is 1.0milliequivalent per liter as compared with 1.1 for the potassium sulfate. If we make the usual assumption that the relative amounts required for precipitation are a measure of the adsorbability of the two ions, this would mean that hydroxyl is adsorbed only a little more strongly than sulfate. This is not in accord with the facts, for it was found that no measurable amount of sulfate is adsorbed from a solution in which the concentration of hydroxyl is 0.00016 when the concentration of sulfate is o.ooj. However, the adsorption of sulfate ion or hydroxyl ion is sufficiently strong that but little of either remains in solution at their respective precipitation values. This is especially true in the case of hydroxyl ion, the pH value of the supernatant liquid after precipitation with potassium hydroxide being 6.85 which is approximately that of pure water. It will be noted that the stability is not increased by further addition of acid after the final pH value of the mixture has reached 3 . 4 which brings the hydrogen adsorption to the flat of the isotherm. In Fig. I it will be seen that the sulfate has attained its maximum adsorption at Weiser and Middleton, J. Phys. Chem. 24, 48
(1923).
PHYSICAL CHEMISTRY O F COLOR LAKE FORMATION
I389
the same pH value, illustrating that maximum adsorption of the precipitating ion corresponds with maximum stability of the sol. The course of the curves demonstrates clearly the truth of the time-honored assumption that the stabilizing action on the acid side is due to the adsorption of hydrogen ions, and the precipitating action of the KOH is due to neutralization by adsorption of the hydroxyl ions.
T h e Eflect of Calcium Zon. I n view of the marked effect of hydrogen on anion adsorption, one would expect a similar effect for other cations that ars strongly adsorbed. The adsorptions of sulfate by hydrous chromic oxide from solutions of potassium sulfate and of calcium sulfate of the same equivalent concentrations were compared. From potassium sulfate, the adsorption was 1.86 milliequivalents at a pH of 7.32, while from calcium sulfate the adsorption was 2.23 at a pH of 7.39. This suggests that one could determine the effect of varying calcium ion concentration on the adsorption at constant pH values. In a subsequent paper on the alizarin lakes this behavior will be shown to be of particular significance. We should not espect the calcium ion to have a yery marked effect in the acid range because it is so much less adsorbed than hydrogen. On the other hand, the effect should become large in neutral and basic range on account of the low concentration of the hydrogen ion. Adsorption of Oxalate Adsorption by Hydrous Chromic Oxide. Experiments similar to those described above for sulfate were carried out for oxalate. The results are tabulated in Table I11 and shown graphically in Fig. 3. The form of the oxalate curve is strikingly similar to that of the sulfate curve throughout the entire range. The difference in the pH values before and after precipitation is greater in the case of oxalate than of sulfate due to stronger adsorption of
TABLE I11 Adsorption of Oxalate by Hydrous Chromic Oxide at varying pH Values Cr of aoln. mixed with jo cc of sol containing o 125 grams of
Adsorption values milli-equiv. per gram Cr203
Cr,03 in a total of z o o cc
0
5
0
IO
4.88 4.77 4.47 4. I7 3.48 2.06 1.32 0.65
0
20
0.00
0
30
0.06
50
0
40
0
30
0
20
0
IO
0
0
0
pH values Before adsorption ' 73 2.87 3.16 3.81 4.45 8.56 I O . 58 10.67 11.12 11.23
2
After adsorption
2.89 3.16 3.78 4.56 7.54 8.73 8.99 9.20 9.48 9.99
I390
HARRY B. WEISER AND EVERETT E. PORTER
oxalate ion which is accompanied by a corresponding increase in the adsorption of hydrogen ion. The fact that the quantity of oxalate adsorbed is a little less than of sulfate in the range of total acid around 0.0015 normal is not an exception to the above statement. The primary and secondary ionization constants for the two acids are as follows: H2SOa,Kl = 0.450, Kz = 0.017; H2C20a,K1 = 0.100,KS = o.ooooj. I t is obvious that at the normality under consideration, the concentration of the divalent oxalate is considerably less than that of divalent sulfate. Jloreover, there is a considerable difference in the hydrogen ion concentration of the two solutions as is indicated by the greater pH value of oxalic acid as compared with sulfuric. In the section on the simultaneous adsorption of oxalate and sulfate are graphs which show this
same effect. On the other hand, at the same hydrogen ion concentration, the adsorption of oxalate is greater than that of sulfate throughout the entire range investigated. There is no way of telling just what part of the adsorption is bioxalate or bisulfate but it is evident that a considerable quantity of bioxalate is adsorbed in the acid range. Precipztation Experiments. In Table IT’ and Fig. 4 the relation between the hydrogen ion concentration and the precipitation value of oxalate is shown. It will be noted that the precipitation value for oxalate is less than for sulfate a t p H values between 6 and 9, a behavior which would be considered normal since oxalate is adsorbed more strongly than sulfate. On the other hand, a t pH values less than 5 the precipitation value of oxalate exceeds that of sulfate although the adsorption of oxalate still continues greater than sulfate. This apparent discrepancy is due to the fact that the secondary ionization of oxalic acid is but 1/350 of that of HZSO,. I t is obviously erroneous, therefore, to assume that equal concentrations of sulfate and oxalate give equal concentrations of the strongly adsorbed divalent ions. In the case of oxalate the
1391
PHYSICAL CHENISTRY O F COLOR LAKE FORMATIOS
TABLE IV Precipitation of Chromic Oxide Sol with Oxalate a t varying pH Values Precipitation values milli-equiv. per liter of m i s
Cc. of s o h mixed with 5 cc of so1 containing 0.0125 grams of Cr203in a total of 20 cc
pH values Before adsorption
-
After adsorption
S / j o H2Cs04S i j o IifC:OI S j o IiOH
3.35 3.15 2.80 2.60
0.0
0.0
3.25
3.05
0 . 2
0.0
3.35
3.11
3.19 3.32
0.4
0.0
3.20
0.j
0.0
3.10
3.17
3.40
2.20
0 .i
0.0
2.00
0.8 0.9
0.0
2.90 2.80 2 . jo
3.38
3.5’
0.00
1.0
0.0
I
0.00
0.0
1.0
I , 00
9,40 10.67
5.79 6.8j
I
.60
.2
0.0
A
Conrentration y
K,C,O+
6
a
.oo
1.0
in MiiIioquivdIents
actual concentration of the divalent ion, which is a measure of its effective concentration, is but a small fraction of the total concentration in the acid solution. Hence the equivalent adsorption of oxalate appears t o be greater than that of sulfate since a larger percentage of the former is adsorbed as the univalent bioxalate ion and is expressed as if all of it were adsorbed as the bivalent ion. Attention should be called to Weiser and Middleton’sl observation that the precipitation value of sulfate was less than that of oxalate for a Crum’s alumina sol although the adsorption of oxalate by the precipitated gel was 1
J. Phys. Chem., 24, 630
(1920).
1392
HARRY B. X E I S E R A S D EVERETT E, PORTER
greater than that of sulfate. Since the Crum sol was quite acid this anornalous behavior is readily understood. Thus, the precipitation value in milliequivalents per liter of sulfate in neutral solution is 1 . 1 and that of oxalate is I . This is the reverse of the order of adsorption, as it should be. On the other hand, when the hydrogen ion concentration is as great as 0.0001normal, the order of precipitation values is reversed. Adsolption by Hydrous Ferric Oxide arid A l u m i n i i u i Oxide. The data for the adsorption of oxalate by hydrous ferric oxide and by alumina are given in Tables V and VI and represented graphically by Figs. 5 and 6. As one might expect the effect on adsorption of varying the hydrogen ion and hydroxyl ion concentrations is similar to that observed with the hydrous chromic
oxide. Due to the inherent differences in the colloids, the total quantity i s different for each oxide but the general character of the pH-adsorption curves is the same. TABLE V Adsorption of Oxalate by Hydrous Aluminum Oxide at varying pH Values Cc of soln. mixed with 50 cc of sol containing 0.0845 grams of AlzOs in a total of zoo cc
N/jo HzCzOi N/jo K2CzOr
0.037
__
N KOH
Adsorption values milli-equiv. per gram AlpOs
50
0
0
I.20
30
20
0
20
0
5
30 40 45
0
50
0
0
50
5
0.99 0.94 0.76 0.69 0.43 0.16
0
50
IO
0 .I 1
10
0 0
pH values after adsorption
2.82 3.45 4.35 6.62 7.58 8.92 10.08 I O .58
PHYSICAL CHEMISTRY O F COLOR L A K E FORMZATIOX
1393
TABLE TI Adsorption of Oxalate by Hydrous Ferric Oxide at varying pH Values Cc of soln. mixed with 50 cc of sol containing 0.132 grams of F e 2 0 3in a total of zoo cc H~CZO S /~ j O II?CzOI
S/jO
0.0.037 S
__
Adsorption values milli-equiv. per gram Fen08
pH values after adsorption
KOH
50
0
0
I .32
3.09
30
20
0
4.18
20
30
0
1.31 1.29
IO
40
0
1.25
5
0 0
.08 0.46
5.92 6.32
0
45 50 50
5
0.07
0
50
IO
0.00
0
50
20
peptized
0
4.74
I
8.79 9.84 42
IO.
IO 8
PH 4 2
Simultaneous Adsorption of Sulfate and Oxalate Observations on simultaneous adsorption by hydrous oxides have been niade by IIehrotra and Sen’ and by The former studied the mutual effects of copper and barium i n varying proportions on their relative adsorption by hydrous manganese dioxide. Since no account was t a k m of the hydrogen ion concentration to which hydrous manganese dioxide is quite sensitive, their results are of questionable value. Data in one specific case led them to the general conclusion that “The sum of the adsorption of two ions from a mixture is always greater than the adsorption of either of them present alone in the solution.” Reference to Table TI11 as well as to Table I 9 and 1
J. Indian Chem. Soc., 3. 397 (1926).
2U.S. Public Health Reports, 39, I j O 2 (1924).
‘394
HARRY B. WEISER AND EVERETT E. PORTER
Figs. 7 and 8 will show that this is not true in every case, and it may not be true in any case, provided the precipitations take place under the same conditions. Some of the data obtained by Miller on the simultaneous adsorption of oxalate and sulfate by hydrous alumina are reported in Table VII. On the
TABLE YII Simultaneous Adsorption of Sulfate and Oxalate by A1203(Miller), Before precipitation Molar concentration KzSO4
Molar concentration Ii*C*O1
At equilibrium after precipitation
I
