t o have greater effects in the zinc chelate. From this point of view, it is not too surprising t o see the p-CH3 point somewhat further above the line than the p-F and p-C1 points, whose electron-release contributions by conjugation are offset by electron-withdrawing effects by induction. The great differences in the behavior of the pair of methylsubstituted dithizones with nickel and zinc is noteworthy.
I
I1
I t is apparent from Figure 2 that the position of the methyl group has a more profound effect on the formation of the zinc than on the nickel complex. Since, as indicated by molecular models, the phenyl ring in the Ni complex (11) lies out of the plane of the chelate ring, the presence of a methyl substituent in either ortho o r para position would have about the same effect. In contrast, the presence of an ortho substituent on the phenyl ring in the zinc chelate (I) would severely limit the free rotation of this ring. The steric effect shows up as strongly in the rate of formation of the zinc chelates as in their stabilities. At this time it is impossible t o evaluate separately the factors of the rate constant, k , which is a composite of a rapidly formed intermediate complex and a first-order decomposition constant. The behavior of the p-OCH3 derivative should resemble that of the p-CH3 compound. Its failure t o d o so cannot be readily understood. Further work involving rate studies with other metal ions is under way. RECEIVED for review August 26, 1966. Accepted December 15, 1966. The authors are grateful t o the Atomic Energy Commission and the National Science Foundation for financial support in this work.
Effect of pH on Ion and Precipitate Flotation Systems Alan J. Rubinl and J. Donald Johnson Department of Encironmental Sciences and Engineering, Schoo! of Public Health, Unicersity of North Carolina at Chapel Hill, N . C.
The effects of pH on the ion flotation and precipitate flotation of several metal ion-collector systems were examined. The anionic collector, sodium lauryl sulfate, was used to remove both soluble and insoluble copper(l1) and iron(ll1) species. Stearylamine, a cationic collector, removed dissolved copper and copper hydroxide, while soluble iron was not removed and ferric hydroxide was only partially removed by this collector. A weak acid collector was found to be less efficient than a strong acid collector for removing iron by precipitate flotation. A method for predicting precipitate flotation as a function of p H is described.
IONFLOTATION and precipitate flotation are foam separation processes used to remove surface inactive substances from aqueous dispersions. This is accomplished by adding a surface active agent, a collector, that will react to form a surface active product with the component to be removed. With these processes the separation is obtained at the interface of the bulk and foam phases without assistance from the extended phase, having the advantage of producing a dry foam of small volume, and thus allowing the use of compact equipment. Both ion and precipitate flotation are sensitive to bubble size and, therefore, an alcohol frothing agent is added along with the collector. The ion flotation technique was introduced by Sebba in 1959 ( I ) , and the precipitate flotation technique by Baarson I Present address, College of Engineering, University of Cincinnati, Cincinnati, Ohio 45221
__ ( I ) F. Sebba, A ’ a ~ e 184, , 1062 (1959).
298
ANALYTICAL CHEMISTRY
and Ray in 1963 (2). Although precipitate flotation has not been extensively investigated (3), Sebba has published several papers on ion flotation (4-7). A detailed study in which the effects of several variables o n the two processes are compared has recently been described by Rubin et 01. (8). With Sebba’s process an insoluble metal ion-collector product is floated, whereas the ion flotation process described in this paper involves the partition of a soluble ion-pair product. Ion and precipitate flotation differ from one another in that with the latter process the component to be removed is precipitated before the addition of collector. In principle this may be accomplished by adding any substance that forms an insoluble compound, however, this work is restricted to hydroxide precipitates. Since surfactant need only react with ions on the surface of the condensed phase to produce a hydrophobic surface, only small amounts of collector are required. Thus, while ion flotation requires stoichiometric o r greater concentrations of collector, precipitate flotation is effective in some systems in which the metal ion concentra-
__________ ( 2 ) R. E. Baarson and C. L. Ray, “Precipitate Flotation-A New
Metal Extraction and Concentration Technique,” Amer. Inst. of Mining, Metallurgical and Petroleum Engineers Symposium,
Dallas, Texas, 1963. ( 3 ) J. A. Lusher and F. Sebba, J. Appl. Cllem., 16, 129 (1966). (4) Ihid.,15, 577 (1965). (5) N. W. Rice and F. Sebba, Ihid., p. 105. (6) F. Sebba. “Ion Flotation,” Elsevier, New York, 1962. ( 7 ) F. Sebba, Narzrre, 188, 736 (1960). (8) A. J. Rubin, J. D. Johnson, and J. C. Lamb, Ind. Errg. Ckem. Process Design Decelop., 5, 368 (1966).
MERCURY U - TUBE
LOW-FLOW- RATE
I I Figure 1. Schematic of experimental apparatos
tion is 100 times that of the collector (2). In addition, ion flotation is very sensitive to gas flow rate which affects the rate of removal; to co Lector concentration which determines the amount removed but not the rate; and to ionic strength which affects both the rate and total removal. Precipitate flotation on the other hand is relatively insensitive to gas flow rate and collector concentration and completely independent of ionic strength (8). Thus, precipitate flotation is a true flotation process that is similar in many respects to froth flotation. It is unlike Froth flotation, a mineral dressing process, in that the precipitate being rernoved is a flocculent noncrystalline material. With precipitate flotation the suspended material is rapidly removed accumulating on the surface of the foam bed allowing the float to be easily collected. The purpose of the work described in this paper was to investigate the effect of initial pH on several metal icn collector systems. E:XPERIMENTAL
Apparatus. Figure 1 shows a schematic of the apparatus used in this work. Nitrogen gas was passed to the flotation cell through a gas humidifier, glass-wool filter, Manostat rotameter (model FM 1042B), and Moore low flow rate controller (model 63B1J-L). A 600-ml glass Buchner funnel having a diameter of IO-cm with a fine sintered glass frit served as the flotation cell. Line gas pressure was monitored with an open mercury U-tube manometer and maintained at 30-inches upstream of the controller. The rotameter was calibrated at this pressure from data supplied by the manufacturer. All gas f l o ~ rates ! were converted t c average atmospheric pressure and room temperature and are reported as milliliters per minule. A Beckman Zeromatic meter was used to measure pH, and metal ion concentrations were determined colorimetrically using a Bausch and Lomb Spectronic 20. Chemicals. The collectors, sodium Iaiiryl sulfate (NaLS), stearic acid, and stearylamine, were used without further purification. Stearic a:id and NaLS were supplied by Fisher Scientific in reagent and U.S.P. grades, respectively, and practical grade stearylamine was supplied by Eastman. Before use, stearylamine and :,tearic acid were dissolved in reagent grade absolute alcohol. NaLS was dissolved in a 50% by volume ethanol-water solution. This allowed the simultaneous addition of collector and alcohol frother to the flotation cell. A stock iron solution of 4.0mM concentration was prepared from iron wire ‘dissolved in nitric and sulfuric acids, and a 6.0mM stock copper solution was prepaied from reagent grade copper(l‘1) perchlorate. Laboratory distilled water did not contain (detectable amounts of the metals and
was used without further treatment. Reagent grade sodium hydroxide, sulfuric or perchloric acid, and sodium perchlorate were used to adjust the p H and/or ionic strength. Perchlorate reagents and reagents for the iron and copper analyses were supplied by G. Frederick Smith Chemical Co., except 1-hexanol which was supplied by Fisher Scientific. The hexanol was redistilled before use. Procedlire. Experimental solutions, 0.2mM (1 1 mglliter) iron(II1) o r 0.3mM (19 mgiliter) copper(II), were prepared from the stocks, being adjusted to the desired ionic strength and approximate pH, and transferred to the flotation cell. The experiments were batch type using a volume of 400 ml. Gas flow was started and a final adjustment in p H was made, this p H was recorded and the gas flow rate adjusted. Two or three 1-ml samples were taken for analysis. Collectorfrother solution was added with a syringe, the collector dose being constant at 1 ml added to 400 ml. Samples of the bulk were taken with volumetric pipets at predetermined intervals. The data reported in this paper. except for those in Figure 5, are for maximum removaib. The foam was allowed to remain in the cell during the course of the experiment and was not analyzed. The samples were boiled in small flasks with acid, cooled, and then transferred to 250-ml separatory funnels with stopcocks of Teflon. The analytical procedures for the iron and copper were very much the same involving an extraction with 1-hexanol, and using bathophenanthroline and bathocuproine, respectively, as the chromogenic reagents (9-11). After each experiment the flotation cell and separatory funnels were washed with alcohol, acid, and distilled water. Pipets and other glassware were washed in detergent and then treated in a similar manner. RESULTS
Iron(II1) Systems. If it is assumed that the precipitate of a metal ion is completely removed from dispersion as soon as it is formed, say by precipitate flotation, then a “theoretical” precipitate flotation curve can be drawn based on calculations from hydrolysis data. For iron(II1) at room temperature, and an initial concenLration, Co, low enough (ca. 0.2mM) so that soluble polynuclear species are not formed or make u p only an insignificant fraction of the total, the maximum concentration of soluble (aqueous) iron(II1) species, C, is : C,
=
+ [FeOH+*]+ [Fe(OH)2+]
[Fe+3]
or iis a function of pH: C,
=
KdH3 f K d k H 2
+ KdKiK2H
(2)
where Kd is the solubility or dissociation constant, (Fe+3)/ (Hf)3, Ki and K2 are the first and second hydrolysis constan?s for iron(III), respectively, and H is the hydrogen ion concentration calculated from pH (12, 13). Assuming 100% efficiency the per cent removal of the metal precipitate may be calculated by substituting C, in Equation 2 into Equation 3 below : =
lOO(1
- C,jC,)
(31
(9) H. J. Cluley and E. J. Newman, A~iolyst,88, 3 (1963). (10) H. Diehl and G. F. Smith, “The Iron Reagents,” G. Frederick
Smith Chemical Co.. Columbus, Ohio, 1960. (11) H. Diehl and G. F. Smith, “The Copper Reagents,” G. Frederick Smith Chemical Co., Columbus, Ohio, 1958. (12) J. N. Butler, “Ionic Equilibrium. A Mathematical Approach,” Addison-Wesley, Reading, Mass., 1964. (13) A. J. Rubin, Ph.D. thesis, University of North Carolina at Chapel Hill, 1965. VOL. 39, NO. 3 , MARCH 1967
299
PH
.
Figure 3. Low gas flow rate foam separation of 0.2 mM iron(II1) with stearylamine 0
3
2
1
4
5
PH Figure 2. Foam separation of 0.2 m M iron(II1) as a function of collector concentration, C mM, and gas flow rate, G mlimin, us. pH using sodium lauryl sulfate as collector
The calculated precipitate flotation curve, that is, the precipitate formation curve for 0.2mM iron(II1) is shown in Figures 2 and 3. The values of the various constants used to calculate these curves are given in Table I. These curves show that precipitate just begins to form at p H 2.67 (the Precipitation point), the metal being completely soluble at pH values less than this. At pH 2.67 and above, iron(II1) at this concentration will exist in both soluble and insoluble forms. At and above p H 4 only an insignificant fraction of the metal remains in solution. Figure 2 shows the results of several foam separation experiments at various collector concentrations, C mM, and gas flow-rates, G ml/min, using NaLS as the collector. All the removals in this figure fall o n the curve o r to the left of it. The removals at p H values less than the precipitation point must be due to the foam separation of dissolved iron(II1) species, that is, these removals are by ion flotation. Removals on the curve, then, are due to precipitate flotation, while removals between the curve and the precipitation point are due both to ion and precipitate flotation. This latter type of mixed flotation is discussed elsewhere (8). The removals at the higher collector concentrations and gas flow rates are observed to increase with an increase in pH. At the lowest gas rates and collector concentrations the removals are lower but increase sharply at p H values at and above the precipitation point in agreement with the calculated curve.
Table I. Hydrolysis Constants Used to Calculate the Precipitate Flotation Curve Constant Value Source
Ki Kz Kd
2.5 2.6 9.1
x x
x
10-3 10-5 108
(14) (15) (12)
(14) W. C. Bray and A. V. Hersey, J. Am. Chem. SOC.,56, 1889 (1934). (15) W. Stumm and G. F. Lee, Schweiz. 2. Hydrol., 22, 295 (1960).
300
ANALYTICAL CHEMISTRY
Gas flow-rate 10.8 ml/min, collector concentration 0.1 m M , flotation time 10 to 11 minutes
NaLS is a strong acid collector; Figure 2 suggests that this collector is capable of removing dissolved, colloidal, and condensed iron species. Stearic acid, in contrast to NaLS, is a weak acid surfactant, its dissociation being strongly affected by the pH of the solution. Several similar precipitate flotation experiments were run using stearic acid as the collector. The removals were not as high as predicted by the calculated curve, and were not reproducible between experiments. In contrast to NaLS, the removals by this collector were also very dependent upon the collector concentration and gas rate indicating that a weak acid collector is less efficient for removing iron(II1) than a strong acid collector such as NaLS. Soluble iron(II1) species, except at very high pH, are positively charged. Therefore, from electrostatic considerations, a cationic collector such as a surface active amine would not be expected to remove soluble iron. Insoluble iron substances, however, are known to be floated by amine collectors (2, 16). These effects, the nonfloatability of soluble iron and the floatability of condensed iron using stearylamine as the collector, are shown in Figure 3. N o removal occurs in this system until the precipitation point. Just at the precipitation point the removals fall on the curve but as the p H is further increased, they irregularly approach a maximum of only 82 %. During precipitate flotation the condensed iron species were removed within the first few minutes of foaming forming a distinct orange layer on top of the foam bed. With ion flotation no visible product was observed when the p H was below the precipitation point. Copper(I1) Systems. Copper(I1) is similar in many respects to iron(I11) but it precipitates at a higher p H forming a more reproducible precipitate. Soluble copper species, however, are much more complicated in solution and the various hydrolysis constants for many of these copper species are unknown or disputed (17, IS), and therefore a precipitate formation curve for copper was not derived. Instead, ion and precipitate flotation of copper(I1) may be compared to the curve shown in Figure 4. The foam separation of copper by ion and precipitate flotation using stearylamine as the collector is shown in Figure 4. (16) F. N. Belash and A. I. Andreeva, Sb. Nauchn. Tr. Kricorozhski Gornorudn. Znst., 12, 275 (1963); C.A. 59, 136256. (17) C. Bereck-Biedermann, Ark. Kemi.,9, 175 (1957). (18) D. D. Perrin, J. Chem. SOC.,1960, p. 3189.
100
I
I
I
I
I
I
I
00
$ B
2 Y
a.
60
E
40t 20
0
0
2
4
6
a
10
PH
0 PH
Figure 4. Low gas flow rate foam separation of 0.3 m M copper(I1) with stearylamine
Figure 5. Interrelationship of pH and ionic strength in copper(I1tNaLS ion flotation system
Gas flow rate 5.2 or 10.8 ml/min, collector concentration 0.1 or 0.2 m M , ionic strength 0.030, flotation time, 8 to 10 minutes
Gas flow rate 10.8 ml/min, copper(I1) perchlorate concentration 0.3 mM, collector concentration 0.6 mM, flotation time 40 minutes. Ionic strength adjusted with perchloric acid or sodium perchlorate
Precipitated copper was separated very efficiently, the removals forming a srnooth curve and approaching 100%. Below pH 7 soluble copper species at an ionic strength of 0.03 were partially removed by the amine collector. These effects were undoubtedly due to strong coordination between copper and nitrogen. The results of several experiments on the Cu(I1)-NaLS ion flotation system are simmarized in Figure 5 . The removals obtained in 40 minutcs are shown as a function of ionic strength and pH, all 3ther conditions being constant. The removals are compared to the precipitate flotation curve for the Cu(I1)-stearylamire system. Points on the extreme left are for removals at t i e lowest pH possible at each ionic strength, Removals are seen to decrease with ionic strength; it being apparent that ion flotation is extremely sensitive to this parameter. Hurrps in the curves at the higher ionic strengths are not readily explainable except in terms of differences in the ionic compositions of the solutions. Shown also is a single point cin the precipitate flotation curve at an intermediate ionic strength, indicating that the foam separation of condensed species is independent of the concentration of soluble salts, DISCUSSION
Ion Flotation Systems. With the four ion flotation systems just described different patterns of removal were observed. As expected, the anionic collector, sodium lauryl sulfate, removed soluble iron(II1) species, while no removals were obtained using the cationic collector, stearylamine. Removal of iron by NaLS increased as a direct function of pH reaching a maximum coinciding with the precipitate flotation curve. In contrast to iron, copper(I1) was removed by stearylamine as well as NaLS, and at high ionic strength the maximum in removal with these twm3copper systems occurred before the precipitation point, removals decreasing upon approaching the pH of precipitation. These results can be explained by three considerations. In these systems the sodium ion plus hydrogen ion concentrations were kept constant, the sum of the two being approximately equal to the ionic strength and thus sodium ion concentration decreased as the pH was decreased. With the systems exhibiting ion flotation, removal decreased from the maximum with decreasing pH apparently because the metal
ion could not compete against hydrogen ion for the collector as successfully as against sodium ion. This is readily explainable since Hf has a higher charge density than Na+, and is therefore harder to displace from the laurylsulfate anion. At high ionic strength-Le., high Na+ and/or H+ concentrationthe removals decrease (Figure 5 ) because of the increased competition from Naf and H + with the metal ion for the collector. Iron is an oxygen coordinator while copper is primarily a nitrogen coordinator. Thus in comparing the two metal ionNaLS ion flotation systems, copper removal is observed to be much more sensitive to ionic strength. On the other hand, soluble iron is not removed by stearylamine while copper is removed in spite of unfavorable electrostatic effects. Copper removal by stearylamine, however, would be expected to be less sensitive to Na+ competition than the other systems since the main competition is between Hf and copper ion for the electron pair on the nitrogen in the amine. It is concluded that coordination between the metal ion and collector is more important than electrostatic considerations with the copper(I1)-stearylamine system. As the pH is increased toward the precipitation point the concentration of metal polynuclear species also increases. Apparently, soluble copper polynuclear species-e.g., Cu3(OH),+2 and C U ~ ( O H ) ~(17, + * l8)-do not compete as well for collector as does Cu+2 which has a higher charge density and more coordination sites available for collector. Iron polynuclear species, if formed, and of which Fe2(OH)2+4is the most important (19, 20), would be expected to compete very well with Fe+3. Thus because of the differences in the charge densities of the respective polynuclear species, copper removal would decrease while iron removal would increase in the vicinity of their respective precipitation points. In systems where coordination effects are a minimum charge effects predominate. Thus removal in soluble systems involves a competition between collector and the various cations in the solution, this effect being altered to a greater or lesser extent by complex formation. It would be expected (19) B. 0. A. Hedstrom, A r k . Kemi., 6 , l(1953). (20) E. Matijevic and G. E. Janauer, J . Co/loid Interface Sci., 21, 197 (1966). VOL. 39, NO. 3, MARCH 1967
301
then that selectivity of removal could be effected by choosing a collector that has coordination o r chelation properties selective for the component to be removed. Ionic strength also affects the degree of separation by ion flotation by directly altering the ionic competition for the collector as it is being carried to the foam phase. Selectivity then could be improved in some systems by adjusting the concentration of neutral salts as well as the pH. Precipitate Flotation Systems. Iron and copper hydroxide precipitates were removed from dispersion using several types of collectors. As with ion flotation, coordination effects were observed; stearylamine, for example, efficiently removed copper hydroxide but not iron hydroxide. The removals of iron by the nitrogen collector were nonreproducible and reached a limiting removal of about 82%. It is concluded that the lack of reproducibility of the results was largely due to the variable character of iron hydroxide precipitates (19), and that a maximum removal was observed because iron and stearylamine form an unstable complex. NaLS efficiently removed both copper and iron precipitates; most likely differences in coordination were not observed because of favorable charge effects between collector and metal hydroxide. The strong acid collector, NaLS, was more efficient for
removing iron than the weak acid collector, stearic acid. This effect occurs because of competition between metal and protons for the latter collector, lauryl sulfate maintaining its negative charge even in very acid solutions. In general, removals with NaLS were rapid and virtually complete within a relatively short period of time. The experimental data for iron(II1) agreed very well with the calculated precipitate formation curve demonstrating the applicability of hydrolysis data for predicting and comparing removals by precipitate flotation. Because of differences in precipitation points among transition metals, it is very likely that separations can be effected by simple pH adjustment. The separation of metals by this technique should be examined in more detail. RECEIVEDfor review July 28, 1966. Accepted January 9, 1967. Research made possible by the Department of Environmental Sciences and Engineering with financial support provided by U. S. Public Health Service training grant ES 7-04, and fellowship number WPPM-16, 462. The research described in this paper is part of a dissertation submitted by Alan J. Rubin in partial fulfillment of the requirements for the Ph.D. degree, University of North Carolina at Chapel Hill
Some Aspects of Thermoelectric Vapor Pressure Osmometry Arnold Adicoff and Warren J. Murbach Organic Chemistry Branch, Code 5056, Research Department, US.Natal Ordnance Test Station, China Lake, CUI$. 93555 The thermoelectric vapor pressure osmometer has been established as a satisfactory instrument for the determination of the vapor pressure lowering of solutes in solutions. The calibration constant, a,, of the instrument i s demonstrated to be independent of the nature of the solute and can be calculated from a knowledge of the system geometry and measurable thermodynamic and transport parameters. Experimental results have been evaluated in terms of the change in resistance of the thermistor and the concentration of solution. It was necessary to use a second degree equation in concentration plot to fit the data. Evaluation of the coefficient of the second order term has required the incorporation of a heat of mixing t e r m in the form of the Van Laar equation. The concentration independent term p for the heat of mixing was obtained for 14 solutesolvent pairs.
WITHINTHE LAST FEW YEARS, interest in a rapid and precise method for measuring the number-average molecular weight of low molecular weight polymers and prepolymers has increased. In the range of molecular weights of polymers below 10,000 the thermoelectric technique has become popular. Ebullioscopic and cryoscopic methods require difficult differential techniques for the highest sensitivity and suffer markedly from problems that lead to such complications as limited solvent choice, foaming, coprecipitation, solvation, and molecular stability. These methods are often tedious and time consuming. The thermoelectric method on the other hand is rapid, requires small samples, and permits a wide choice of solvents. I n addition, recently, the method is said to have been extended to polymers of 40,000 molecular weight ( I ) . (1) M. J. R. Cantow, R. S. Porter, and J. F. Johnson, J. Polymer Sci., A2, 2547 (1964).
302
ANALYTICAL CHEMISTRY
It therefore became of interest to determine the extent to which a popular commercially available instrument for thermo, electric determinations could be used as a research tool, andif possible, determine the necessary conditions for its proper use as such a tool. A number of authors (1-14) have utilized the method of Hill (15) to measure either vapor pressure differences or, from these differences, molecular weights of solutes. A paper by Tomlinson and coworkers (12) has reviewed vapor phase osmometry and calculated thermodynamic efficiencies for a number of solvent systems and Van Dam (13) has optimized a thermoelectric vapor phase osmotic system using thermocouples and a detailed analysis of the mass and heat transport
(2) E. J. Baldes, Biodynarnica, 46, 1 (1939). (3) E. J. Baldes and A. F. Johnson, Zbid., 47,l (1939). (4) A. P. Brady, H. Huff, and J. W. McBain, J. Phys. Colloid Chem., 55, 304 (1951). (5) D. E. Burge, J . Phys. Chem., 67, 2590 (1963). (6) W. I. Higuchi, M. A. Schwartz, E. G. Rippie, and T. Higuchi, J. Phys. Chem., 63, 996 (1959). (7) S. Kume and H. Kobayashi; Makromol. Chern., 79, 1 (1964). (8) R. H. Muller and H. J. Stolten. ANAL.CHEM., 25, 1103 (1953). (9) J. J. Neumayer, Anal. Chin?.Acfa, 20, 519 (1959). (IO) R. Pasternak, P. Brady, and H. Ehrmantraut, Paper presented at the ACHEMA 1961, 13th Chemical Engineering Congress, June 1961, Frankfurt am Main, Germany. (11) G. B. Taylor and M. B. Hall, ANAL.CHEM., 23, 947 (1951). (12) C. Tomlinson, Ch. Chplewski, and W. Simon, Tetrahedron, 19, 949 (1963). (13) J. Van Dam, Rec. Trac. Chitn., 83, 129 (1964). (14) A. Wilson, L. Bini, and R. Hofstader, ANAL.CHEM.,33, 135 (1961). (15) A. V. Hill, Proc. Roy. SOC.(London),A127, 9 (1930).
