P L E N A R Y ACCOUNT
EXTRACTION BY PHASE SEPARATION WITH MIXED ONIC SOLVENTS R.Grintad
S. C. Doris
Descriptive Chemistry R O B E R T R. G R I N S T E A D J A M E S C . D A V I S SCOTT LYNN' R O B E R T K. C H A R L E S W O R T H
Western Division Laboratories, The Dow Chemical CO., Walnut Creek, Calif. 94598
THEseparation of mixtures of substances into the individ5. Lynn
R.K.Charlenworth
R. R. GRINSTEMIholds a B. S. degree in Chemistry from the Uniuemity of California at Berkeley (1946), and a Ph. D. in Chemistry from California Institute of Technology (1950). He participated in the early deuelopments of salwnt extraction and ion exehange processes for the recouery of uranium and uanadium from low gmde ores, and mare recently has been studying other applications of solwnt extraction to inorganic problems, including the processing of brines for recowry of baron, magnesium chloride, and halogens. He is (I member of Phi Beta Kappa, the American Chemical Society, and the Ameriron Associotionfor the Adoancement of Science. 3. C. DAVISholds a B. S. degree in Chemistry from the Uniuersity of California at Berkeley (1966). He hm worked in the fields of industrial inorganic sepamtions by means of soluent extmtian and membrane methods. He ia a member of the American ChemicalSociety. SCOTTLYNN,B. S. (1950). M. S. (1951), and Ph. D. (1954) in Chemical Engineering, California Institute of Technology, 1953-4, Technical Uniwrsity, Delft, Netherlands, is currently Professor of Chemical Engineering at the Uniwrsity of California, Berkeley, where he has taught since 1967. He U M S formerly (195441966) with the Dow Chemical Company. His field of interest is pmcess synthpsis and deuelopment, particularly in the area of inorganic chemicals, industrial electrolysis, heat transfer in polymerizing systems, andgas absorption. He is a member of the American Chemical Society, the American Institute of Chemical Engineers, and the Electrochemical Society. ROBERTK. CHARLESWORTH has Chmicul Engineering degreesfmm the Uniuersity of Washington (B. S.), the Uniwrsity of Wisconsin (M. S.), and Purdue Uniuersity (Ph. D., 1951). He semd as an engineering officer (USNR) during World War I I and taught Chemistry at Idaho State College (1947-8). He is currently leader of a group of Chemical Engineers engaged in p m s dewlopment. His ezperience ot Dow has been in process research, design, and dewlopment in the arem of polymerization, high polymer handling systems, fibers,organic synthesis, and solwnt extmction. He is a member of the American Chemical Society, the American Institute of Chemical Engineers, and Sigma Xi.
218
I & E C P R O D U C T RESEARCH A N D D E V E L O P M E N T
ual components is one of the major tasks of the chemist or chemical engineer. The mixture may he a homogeneous one of gases or of liquids, or i t may he composed of various solid phases, or even of one type of phase in another-e.g., a finely divided solid suspended in a liquid or gas. For a number of reasons, the problem of separation of the components of a liquid mixture (usually a solution) is encountered as often as any other, perhaps more often in modern chemical operations. As a result, a large number of unit operations have heen developed as tools for the chemist and engineer faced with a separation problem' involving liquids. Distillation, precipitation, adsorption, ion exchange, solvent extraction, and dialysis are some of the more common and useful operations of this type. Solvent extraction has enjoyed increasing utility in the last ,several years. Industrially speaking, it is not a new operation, having heen in widespread use in petroleum processing almost from the beginning of that industry. There the twin factors of a liquid raw material and immiscibility with water, a common solvent for many chemicals, have made solvent extraction a useful tool in the processing of hydrocarbons. Various acidic and alkaline aqueous solutions, as well as other hydrocarbon-immiscible solvents, are used to remove impurities from petroleum, as well as to separate the major fractions of the raw material. Recent reviews of the applications of solvent extraction in the petroleum and other organic chemical industries have been published hy Treyhal (1963) and by Vashist and Beckmanu (1968). Solvent Extraction in Inorganic Industries
A more recent development has heen the introduction of solvent extraction into the processing of inorganic compounds. In a sense this is not a new application, either, since analytical chemists have known for years how t o 'Present address, Department of Chemical Engineering, UNversity of California, Berkeley, Calif.
The extraction of inorganic salts from aqueous solution by organic salts dissolved in toluene has been studied. A typical system is shown by the reaction:
R4N' R'COO-
+ Na' + CI-
R4N' CI-
+ R'COO- Na'
where the barred species are in the organic phase. Using O.5M alkylammonium carboxylates in toluene, the effectiveness of different amine types as extractants is in the order: quaternary primary secondary > tertiary. The effectiveness of organic acids is in the order: carboxylic > alkylphosphoric > arylsulfonic. The order of extraction of inorganic salts by carboxylic acid systems is: Ca2' Mg'- > Na' > K', and NO; BrCISO:-. Because the equilibrium is readily reversible, the extracted species are stripped from the organic phase simply by contact with water. Since no chemical energy is used, the system functions generally as a separation, rather than a concentration system. In certain cases, however, where a common ion exists in the feed solution, or where a two-temperature cycle is used, a substantial concentration of the extracted species may also be achieved.
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separate halogens, ferric chloride, and uranium nitrate from aqueous solutions with appropriate organic solvents. However, industrial applications hardly existed prior to the 1940's, because of a number of problems which rendered most proposed or potential processes economically unattractive. .4s the petroleum industry was the cradle of organic applications of solvent extraction, so the nuclear industry was the cradle of inorganic applications. The need for uranium of extremely high purity for nuclear uses, coupled with the fact that costs were a secondary consideration, led to the scaling up of the classical method of purifying uranium by extraction of the nitrate into diethyl ether. A vigorous research program sponsored by the U. S. Atomic Energy Commission expanded this development into the raw materials processing of uranium as well as other nuclear raw materials. Solvent extraction in uranium processing has been reviewed by Ellis (1958), and more recently by Gittus (1963), and a number of recent developments in the use of extraction in the nuclear industry have been presented by McKay et al. (1965). Inevitably, further developments have occurred beyond the nuclear field as this new-found technology has diffused into other chemical and engineering areas. At the present time a variety of solvent extraction processes is in commercial operation. Cuthbert (1958) has discussed the extraction technology of thorium, which has developed closely behind that of uranium. The recovery of both vanadium (Swanson et al., 1961) and beryllium (Chemical and Engineering News, 1965) has been described. Closely related metals are also being separated commercially by solvent extraction. The separation of zirconium and hafnium is described by Hudswell and Hutcheon (1957); the separation of niobium and tantalum by Gustison and Pilloton (1963). The rare earths, which have been among the most idifficult groups to fractionate into individual components, are also separated by solvent extraction processes described by Johnston (1966) and Weaver (1968). The above separations involve relatively valuable species, but a number of applications to relatively low value products have been reported. I n particular, the recovery of boric acid from natural brines has been reported to be in operation (Havighorst, 1963), while the separation of phosphoric acid from calcium after reaction of phosphate rock with hydrochloric acid has been reported by Baniel et al. (1959) among others.
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An extensive review of these developments is beyond the scope of this paper, and the reader is referred to recent general reviews by Hanson (1968), Zakarias and Cahalan (1966), and Smithson (1966). Characteristics of Inorganic Extraction Processes
The attractive features of extraction processes are twofold. First, they involve only liquids, which are relatively simple to handle and readily adaptable to continuous operations. A second feature is that a good deal of flexibility exists in the selection of an extractant system. While extractants must usually meet the requirement of low miscibility with water, they need not be highly insoluble in it, nor be liquids themselves. If they can be dissolved in a suitable water-immiscible diluent, and "prefer" that diluent to water, they are potentially suitable components of an inorganic extraction system. The hurdle which solvent extraction has had to surmount in finding applications in inorganic chemical processing is, basically, the relatively high polarity of most inorganic compounds. Many are either salts, or, if un-ionized, have high dipole moments, and are accordingly characterized generally by solubility in water and extensive hydration or interaction with that solvent. Accordingly, most potential inorganic applications involve an aqueous solution as one phase of the system. Because of the ability of water to dissolve many compounds, both inorganic and organic, the concept of a solvent extraction process implies a relatively nonpolar liquid as the immiscible phase. Such liquids are poor solvents for most inorganic compounds, for the same reasons that they are immiscible with water: Their interaction with polar materials is low. As a result, the extraction of inorganic compounds, particularly salt-like compounds, by water-immiscible solvents is usually more difficult than the extraction of neutral molecules. T o the industrial chemist or engineer, this difficulty manifests itself in two ways. Both stem from the fact that the organic phase must contain a compound capable of interacting significantly with the inorganic species to be extracted. Unfortunately, the functional groupings needed for this purpose also tend to increase the interaction of the extractant compound with water. One result is a higher tendency of the organic compound to distribute into the aqueous phase and be lost. The second problem is that its bipolar nature tends also to confer upon the extractant the properties of an emulsifying agent. This VOL. 8 NO. 3 S E P T E M B E R 1 9 6 9
219
is, of course, an undesirable situation, since the rate of throughput in an extraction system is a direct function of the rate a t which emulsified phases coalesce after mixing. Inorganic Extraction Processes
Some Examples. The extraction of an inorganic species from an aqueous solution by a solvent is generally the result of one of two types of interaction. If the species extracted is a neutral compound, solvation or complex formation with some component of the organic phase is usually involved, accompanied by at least partial loss of the hydration shell of the compound. I n fact, the aqueous phase in such systems is frequently a highly concentrated solution with a relatively low water activity. The latter characteristic assists in driving or “salting” the extracted species into the organic phase. Examples of this type of system, as shown in the equations below, are the extraction of uranyl nitrate from strong nitratenitric acid systems by tributyl phosphate, or the extraction of phosphoric acid from calcium chloride solutions using butanol or tributyl phosphate (TBP).
+ 2 NOs-mHzO + 2TBP 2 UOz(N03)z(TBP)z + (2m + n)HzO H3P04’nHiO + ~ B u O H Z H~PO~*~BU + OnH2O H
UO:2*nH20
(1) (2)
In this and other equations, bars denote species in the organic phase. Equation 1 is the basis of current methods for purifying uranium. I n this process tributyl phosphate, usually dissolved in a hydrocarbon diluent, has replaced the relatively hazardous diethyl ether. Equation 2 is the basis for the production of industrial phosphoric acid after dissolution of phosphate rock in aqueous hydrochloric acid. The second common type of extraction is exactly analogous to conventional ion exchange. An organic acid in the solvent, in either free or salt form, exchanges a proton or other cation for the preferred cation in the aqueous phase:
2RS03Na + Ca” 2 (RSOS)&a
(3)
+ CU+*2
(4)
SRCOOH
+ 2Na’ (RCOO)~CU + 2H-
Similar reactions involving the exchange of anions by solvents containing organic bases may also be written:
+ C12 RNH, + NO; + H 2 0
RNH3C1+ I - 2 R”,I RNH3N03+ OH-
(5) (6)
I n the first type of extraction, where a low water activity is often an important component of the driving force for extraction of a neutral compound, the reaction is readily reversed with water alone. Water is effective in two ways: I t interacts specifically with the inorganic compound, as shown in Equations 1 and 2; and it dilutes the inorganic species, shifting the equilibrium toward the aqueous phase by simple mass action. I n this type of extraction system the concentration of the product is generally somewhat lower than the original aqueous solution, since without the application of chemical energy the extraction can only proceed “downhill”-that is, toward more dilute solution. 220
I & E C PRODUCT RESEARCH A N D DEVELOPMENT
When cations or anions alone are exchanged, however, as in the second type of extraction, it is usually necessary to strip them from the solvent with acid, base, or some other regenerant, typically a t a cost of one equivalent of regenerant for each equivalent of extracted ion. I n return for this expenditure of reagents, however, it is frequently possible to recover the extracted ion in an aqueous solution which is much more concentrated than the original. The recovery of copper from leaches of ore by acidic extractants, as indicated in Equation 4, is an example of this type of process, as is the recovery of uranium by similar methods. Uranium recovery has been an industrial process for many years, and the application to copper is currently under development by a number of companies. Although Equations 3 and 4 involve relatively simple acidic reagents, more complex ones can be used. Included in this group are chelating agents possessing more than one ionizable proton, and often containing a number of uncharged coordinating groups as well. As the value of the extracted species becomes lower, the importance of reagent costs increases. T o avoid the cost of regenerating acidic or basic solvents, it is preferable in some instances to extract inorganic ions as neutral compounds in systems which can be regenerated with water. However, the number of water-immiscible organic liquids which exhibit appreciable solvating capacity for very many inorganic salts is limited. When one adds the requirement of reversible extraction of the salts from aqueous solution, the number of both solvents and the salts which they can extract becomes smaller still. Mixed Ionic Extraction Systems
We describe here a new type of solvent-extraction system, which has elements common to both types of mechanisms discussed above-that is, they extract the ionic constituents of inorganic salts as separate species, but in stoichiometric ratio, so that the result is extraction of a “neutral” compound. However, unlike the ion exchange systems, these new systems can be stripped with water; and unlike the neutral extraction systems, a net gain in concentration of the extracted species through the system can often be realized as a result of common ion effects. Finally, because these systems use conventional acidic and basic components, they form a basis for further development of selective extraction systems, simply by the modification of the structure of the individual components of the extractant. Because of these features, we believe this type of system may have considerable promise in extending solvent extraction further into the field of industrial processing of high volume, low cost materials. The operation of these new systems, which we have termed “mixed ionic extraction” systems, is analogous to the “snake-in-cage” or ion-retardation resins reported by Hatch et al. (1957). I n those materials, prepared by polymerizing acrylic acid absorbed in a quaternary amine anion exchange resin, both cationic and anionic sites are present. When in contact with an aqueous salt solution, salt is removed from solution according to the reaction:
R,N+R’COO-
+ Na+ + C1-
2 RdNfC1- +
(Ebrs denote species in the resin phase.) R’COO-Na+(7)
By replacing the +itial aqueous solution with water, the absorbed salt can be removed from the resin. An identical reaction applies to the mixed ionic extraction systems, differing only in that the water-immiscible phase is a liquid rather than a solid. Experimental Methods
Methods. 2-ETHYLUNDECANOIC ACID (EUD). Diethyl ethylnonylmalonate was prepared by a malonic ester synthesis. The diethyl ethylnonylmalonate in ethanol solution was saponified with KOH and decarboxylated by heating in sulfuric acid. The resulting product was distilled at 128°C. a t 0.5 mm. Yield: 63.5%; 99% pure by VPC. SODIUM DINONYLNAPHTHALENE SULFONATE. A commercial compound, Nasul BSN (barium salt), was obtained from the Vanderbilt Chemical Co. as a heptane solution. That compound in solution was brought in contact with excess 2.ON sulfuric acid, separated, and theh evaporated; the residue was titrated to determine the equivalent weight. BIS-(2-ETHYLHEXY1,)HYDROGEN PHOSPHATE was used as obtained from the Virginia-Carolina Chemical Co. ALIQUAT3368 (AQ) was obtained from General Mills. This is said to be a mixture of methyl trialkyl (mainly octyl and decyl) ammonium chloride. A typical analysis wasC 70.33%; N 3.08%; H 7.81%; C1 18.75%. Equivalent weight by chloride titration: 445. PRIMENE JM-T, a tert-alkyl primary amine, was obtained from the Rohm and Haas Chemical Co. I t had an equivalent weight by acid titration of 342. Other amines were prepared as described by Grinstead and Davis (1968, 1969b). RADIOACTIVE TRACERS. Chlorine 36, a P-emitter, was obtained from the New England Nuclear Corp. as an aqueous hydrochloric acid solution and used without further purification. Counting times were adjusted to produce a maximum standard deviation of 1.2% in the net count. Sodium 22, a y emitter, was obtained from the New England Nuclear Corp. as a sodium chloride solution and used without further purification. Counting times were adjusted to produce a =t2% maximum standard deviation in the net count. Procedures. Mixed ionic solvent systems were prepared in most cases simply by mixing the desired amine and acid, and diluting with the diluent of the desired concentration. I n the case of quaternary systems, the quaternary ammonium chloride and the acid were mixed, diluted with diluent, and brought in contact with a slight excess of dilute sodium hydroxide to convert the acid to the corresponding anion. The organic phase was usually washed with water once or twice to remove the sodium chloride formed in this conditioning step. Samples were shaken for one hour and then separated in a centrifuge. Both organic and aqueous samples were subsequently analyzed. Where tracer elements were not used, metallic cations were determined by atomic absorption spectroscopy and chloride was determined by potentiometric titration. Water content was determined by a standard Karl Fischer titration. The distribution of the organic extractants into the aqueous phases was examined in a number of cases, in order to make certain that losses into the aqueous phases were negligible. Back extraction and spectrophotometric
0.05K
Aqueous (MgClp),
Figure 1. Effect of acid type on distribution of magnesium chloride between water and mixed ionic systems Organic phase. 0.48M Aliquot 336s salts of 2-Ethylundecanoic acid A Bis-(2-ethylhexyl) phosphoric acid 0 Dinonylnaphthalenesulfonic acid All in toluene
0
-u
0.00029
I:
0,0001
2
4
Aqueous (NaCI), M
Figure 2. Distribution of sodium chloride between water and 0.50M alkylammonium-2-ethylundecanoates in toluene
0 Aliquot 336s (quaternary) (0.48M) & N-Dodecyl-2-ethylhexylamine (secondary)
V Trioctylamine
(tertiary) Q2,2-Dihexyl- 1 -aminooctane (primary)
methods were used (Grinstead and Davis, 1969a). With all of the systems studied, the distribution loss of organic compounds was less than 10-3M, which is negligible. Experimental Results
This work has been centered around systems of carboxylic acids and either quaternary or primary amines, since these systems extract salts more strongly. Figure 1 illustrates the effect of the acid type on magnesium chloride extraction. The observed order is carboxylic > phosphoric > sulfonic. The data in Figure 2 represent the extraction of sodium chloride with a variety of amines in conjunction with VOL. 8 N O . 3 SEPTEMBER 1969
221
a reference acid and solvent. I t is obvious that the amine component also has a great effect on the extent of extraction, ranging over as many as four orders of magnitude in organic sodium chloride concentration at a given aqueous sodium chloride concentration. The observed order of extraction is quaternary > primary > secondary > tertiary. The same order of extraction is found with mag-
nesium chloride for the same systems, as shown in Figure 3. However, the variation in the magnesium chloride extraction is only about two orders of magnitude. Figures 4 and 5 show the variation of distribution with extractant concentration with a quaternary amine and a carboxylic acid for sodium chloride and magnesium chloride, respectively. In both cases the distribution coefficient is almost directly proportional to the extractant concentration. Because of the relatively high loading of the organic phase, the curves have a slightly concave downward shape. In contrast, Figure 6 shows the same data for a primary I
I
A
Aqueous ( Mg CI2 1 ,M
Figure 3. Distribution of magnesium chloride between water and 0.50M alkylammonium-2-ethylundecanoates in toluene
A
Aliquat 3365 (quaternary) (0.48M)
0 2,2-Dihexyl-1 -aminooctane (primary)
0.5
0 Prirnene JM-T (primary)
1.0
1.5
Aqueous ( Mg Clz), M
0 N-Dodecyl-2-ethylhexylamine
(secondary)
V Trioctylamine (tertiary)
Figure 5. Distribution of MgClz between water and Aliquat 336S:EUD in toluene Organic concentrations shown on curves; analyses by flame photometry, except A by tracer CI"
I
1
P 0.0015
I
0.00102
:I
0.1
0.0005
I
2
3
4
AqUeOU8 (NoCI),M
Figure 4. Distribution of N a C l between water and Aliquat 336S:EUD in toluene at room temperature Extractant concentrations given in figure
222
I&EC PRODUCT RESEARCH A N D DEVELOPMENT
5 Aqueous ( N a C I ) , M
Figure 6. Distribution of N a C l between water and dihexyl-1 -amino octane:EUD in toluene Extractant concentrations given on curves
2,2-
I
1 -1
I
I
0.41
0.05
Aqueous ( MX 1, M
0
0 -5
0
Figure 7. Distribution of alkali salts between water and 0.48M Aliquat 336S:EUD in toluene
1.5
1.0
Aqueous ( M CI,),M
Figure 9. Distribution of MgCh and CaClz between water and alkyl-ammonium 2-ethylundecanoates in toluene 0 , A 0.48M Ali4uot 3365 system
0,V 0.50M
l-omino-2,2-dihexyloctone system
0.0I
0 0
2
4
Aqueous ( Na X 1, M
0 Figure 8. Distribution of sodium salts between water and 0.50M 2,2-dihexyl-1 -amino octane:EUD in toluene
amine system, where the loading is relatively low. I n this case the curves are distinctly concave upward. The data in Figures 7 and 8 illustrate the variation in extraction of various sodium salts in the quaternary and primary amine mixed ionic solvents. The data for potassium chloride distribution are included in Figure 7 for comparison. The data in Figure 9 show the extraction of calcium chloride and magnesium chloride for both the quaternary and primary amine mixed ionic solvents. In both cases the calcium chloride extraction is greater than the magnesium chloride extra.ction. However, the difference is greater for the primary amine system. As can be seen, the over-all extraction is much greater for the quaternary system. Figure 10 shows the effect of a common ion (in this case chloride as sodiuin chloride) on magnesium chloride
I
I
I
0.5
1 1.0
I 1.5
Aqueous ( M g Cl,), M
Figure 10. Effect of sodium chloride on distribution of magnesium chloride between water and 0.48M Aliquat 336S:EUD in toluene O 4 . O M NoCl present
A No NoCl present
extraction. The effect of the 4.OM sodium chloride on magnesium chloride distribution is dramatic, particularly a t lower concentrations of magnesium chloride. The data in Figure 11 show the variations in the sodium chloride distribution as a function of diluent. The chloroform and toluene diluents are similar in effect, but the amyl acetate shows a marked increase in extraction of sodium chloride. Figure 12 illustrates the effect of the same diluents on magnesium chloride distribution. Here toluene and amyl acetate are similar in behavior, and, VOL. 8 N O . 3 S E P T E M B E R 1 9 6 9
223
0.2
-
10.2 3
IE 0,I
Aqueous ( NaCI1, M
Figure 11. Effect of diluent on sodium chloride distribution between water and 0.48M Aliquat 336S:EUD
Aqueous ( Na CI) ,M
0 Toluene
Figure 13. Effect of temperature on the distribution of sodium chloride between water and various extractants in toluene
0 Amyl ocetote A Chloroform
I
I
I
0 , 0 23" C.;A,A60°C.
Open points, 1.OM Primene JM-T:naphthenic acid EEE; filled points, 0.48M Aliquot 336S:EUD
I
0.2
z
go,I
I-P 0 Aqueous ( MgCIZ) , M
Figure 12. Effect of diluent on magnesium chloride distribution between water and 0.48M AQ-EUD in toluene
Figure 14. Effect of temperature on distribution of magnesium chloride between water and various extractants in toluene .,O23"C.;A,A6OoC. Open points, 1.OM Primene JM-T:naphthenic acid EEE; filled points, 0.48M Aliquot 336S:EUD
0Toluene A Chloroform 0 Amyl acetate in some cases, the magnesium chloride level in the organic phase exceeds the theoretical limit of 0.25M. Chloroform shows a decrease in magnesium chloride distribution relative to the other diluents. Figures 13 and 14 show the variations in distribution with temperature for sodium chloride and magnesium chloride, respectively. Both a primary and a quaternary amine system were studied a t 23" and 60°C., using the same reference acid and diluent. In Figure 13 the data show that the distribution of sodium chloride is an inverse function of temperature. The quaternary amine system is more sensitive to temperature than the primary amine system. Figure 14 shows similar data for magnesium chloride distribution; again the quaternary amine system is the more sensitive. However, magnesium chloride distribu224
I 6 E C PRODUCT RESEARCH A N D DEVELOPMENT
tion is less sensitive to temperature than that for sodium chloride. Extraction of water by mixed ionic solvent systems is shown in Figure 15. The difference in behavior between the quaternary and primary amine systems is striking, the former taking up about 8 moles of water per mole of extractant when in contact with water, and the latter only about 0.5 mole. Increasing concentrations of sodium chloride cause little change, largely because the extraction of sodium chloride is relatively low. Magnesium chloride, on the other hand, extracts appreciably even at low aqueous phase concentrations, and at least two effects are evident in the primary amine system. The initial decrease in water extraction with magnesium chloride concentration is probably the result of decreased water activity in the aqueous phase. As magnesium extraction increases, the water coordinated to the extracted magnesium becomes
1 I l
I I I 2 3 4 (NaCI) or(MgCI,), Nondity
I 5
I 6
Figure 15. Water extraction by mixed ionic extraction systems in toluene A. B.
0.40M Aliquat 336S:EUD 1 .O Primene JM-T: naphthenic acid EEE
0
NaCl
0
MgClz
appreciable and causes a rise in the organic water concentration. The later decrease is probably a combination of the reduced activity of free water and the reduced amount of water coordinated to the extracted magnesium. Because of the large amount of water already present in the quaternary system, these effects are not clearly discernible in this system.
Discussion The present paper is intended primarily to describe the chemical behavior of the mixed ionic extraction system. A more detailed analysis of the mechanism of extraction and the nature of the organic phase will be published subsequently. Certain unique features of the system merit some discussion a t this point, however. The operation of the mixed ionic solvent system clearly involves an organic salt as the extractant. I n the case of quaternary ammonium systems, this is obviously the case, but even with other amine types the reaction between amine and carboxylic: acid to produce a salt is fairly extensive. Preliminary study of the infrared spectra of these systems shows clearly the presence of frequencies characteristic of the cationic and anionic species, although it also appears that reaction to form the salt is probably not complete. This is in general agreement with the results of Barrow and Yerger (Barrow and Yerger, 1954; Yerger and Barrow, 1955), who found that acetic acid and various aliphatic amines react fairly completely in carbon tetrachloride. I n nonpolar diluents such as toluene it is fairly certain that the ions cannot exist separately, but are necessarily present as ion pairs. The extraction of inorganic salts by these systems appears to involve the formation of two organic salts, a metal salt of the acid constituent, and a halide or other salt of the amine constituent. This conclusion is based primarily on the behavior of the system toward various cations, which extract in the order divalent > monovalent, and in the latter category, sodium >
potassium. These are the normal order of complex formation with simple ligands. Further evidence is provided by the behavior of the three acid types used, which extract in the order carboxylate > phosphate > sulfonate. While no good comparison of the relative chelating ability of these three functional groups was found, the above order is also the order of affinity of the group for the proton. Martell and Calvin (1952) have pointed out this relationship and shown it to be rather general, and one might expect therefore that the carboxylate would form the strongest chelate, the sulfonate the weakest. Behavior of the amine component of the system also suggests the formation of an amine salt during the extraction. The order of extraction of the sodium salts of various anions was shown to be nitrate > bromide > chloride, the same as that reported by Marcus (1963) for the relative extractabilities of these anions by simple alkylammonium ion extraction systems. The relative positions of nitrate and bromide are uncertain a t higher concentrations, but at low levels where activity coefficients approach 1 the stated order is clearly followed. Diluent effects in the mixed ionic solvent system were studied briefly by using two diluents which provided an example of an electron donor (amyl acetate) and a proton donor (chloroform) solvent. Toluene was the reference solvent. The data are inconclusive, but we believe the rather substantial increase in sodium chloride extraction caused by replacing toluene by amyl acetate is a result of solvation of the sodium ion by the diluent. The importance of having hydrophobic donors in the coordination sphere of extracting cations is well known, and has been discussed a t some length by Stary (1964). The coordination sphere of the sodium is only partially filled by the carboxylate ion, and further coordination by amyl acetate molecules would be expected to stabilize the sodium species greatly in the organic phase. We expect the interaction of amyl acetate with the original extractant to be low, since the quaternary ammonium ion is not susceptible to coordination in the same manner as the sodium ion. On the other hand, chloroform should exert an effect through hydrogen bonding to negative sites, which in this system could involve either the carboxylate ion in the original extractant, or the chloride ion in the extracted species. These two effects are in opposite directions, and could account for the relatively small difference between chloroform and toluene. I n the case of magnesium chloride extraction, amyl acetate behaves the same as toluene, although the former is generally a considerably better solvating agent. We believe this is a result of the reduced need of the magnesium ion for solvation, since its coordination sphere is more nearly filled (as compared with sodium) by the two carboxylate groups. Another interesting phenomenon is the relatively high levels of magnesium in the organic phase when toluene or amyl acetate is the diluent. These levels exceed the stoichiometric limit for a divalent ion, and the most reasonable explanation is the existence of the monovalent ion MgCl+ in the organic phase. The two possible forms of magnesium may exist in equilibrium in the organic phase, as follows: M g 2 + ( R C O O - ) 2 + O c 3 M e N f C 1 2 MgC1' RCOO-
+ RCOO-
Oc3MeN+
(8)
(Oc = octyl, M e = methyl) VOL. 8 N O . 3 S E P T E M B E R 1 9 6 9
225
This excess loading phenomenon is not observed in chloroform’, and in addition, the extraction of magnesium chloride in this solvent is poorer than either of the other two. We suggest that the likely explanation is that chloroform can affect the equilibrium in Equation 4 by hydrogen bonding to the chloride ion, thereby suppressing the formation of the MgC1’ species. One of the unique features of the mixed ionic systems is the common ion effect shown in Figure 10. The extraction of both anion and cation components of magnesium chloride in the manner of Equation 7 results in the equilibrium expression
(9) where CA represents the extractant molecule (CA = cation.anion). Although the extractant is not monomeric as the equation indicates, neither this fact nor the nature of the extracted species affects the dependence of extraction on both aqueous chloride and magnesium concentrations. The dependence is more precisely on the activity of the aqueous species, although attempts to correlate extraction of magnesium chloride from single and mixed salt systems on the basis of salt activity alone have been unsuccessful. We believe this to be due to the importance of water in the extraction, which is appreciable, particularly for the quaternary systems. Any correlation of extraction of inorganic salts from mixed salt systems will have to take into account the variations in water activity. Applications. I t is useful to consider the mixed ionic extraction systems as analogs of semipermeable membranes-that is, the solvent phase provides a transfer medium which allows the desired species to be absorbed and transferred to the aqueous strip or product stream. Since no chemical energy is involved in the operation, the activity of the extracted species in the product can approach but not exceed that in the feed. I n most cases, therefore, the system is of interest as a separation method rather than a concentration method. However, the existence of common ions can provide some concentration through the system, an outstanding case being the extraction of small concentrations of divalent cations from large concentrations of monovalent ones. Consequently, the separation of divalent cations from monovalent cations constitutes the most important class of application visible at this stage of development. Additional studies have been made of the use of mixed ionic systems to separate magnesium chloride from sea water concentrates. This study, and a discussion of the possibilities of application of the process to chemical recovery or sea water softening, are presented in a separate paper (Grinstead and Davis, 1969a). The preference of the system for bromide over chloride raises the possibility of recovery of bromine, and perhaps iodine. However, because the selectivity is not high, and because the methods currently in use are simple and effective, this application appears less attractive, except perhaps as a by-product operation in combination with a divalent ion separation. A further possible application lies in the removal of nitrate ion from waste water. While the extraction data given here for the individual systems shows only about a threefold higher distribution for nitrate than chloride, unpublished data indicate that the selectivity in the mixed 226
I & E C PRODUCT RESEARCH A N D DEVELOPMENT
ionic systems may be higher by as much as 10- to 20-fold. Studies of this possible application are currently in progress. Although less promising, another possibility is the concentration of salt solutions by use of a temperature cycle. For example, Figure 13 demonstrates that the extraction of sodium chloride a t room temperature is about a factor of 2 better than a t 60°C. I t would therefore be possible to extract sodium chloride more or less completely from a brine a t the lower temperature and recover it at about twice the concentration a t 60”, perhaps even more concentrated a t a still higher temperature. It is also possible to remove the extracted salts chemically. Unless a relatively high loading can be obtained, the use of sufficient acid and base to strip the resulting organic phase completely would probably be out of the question. A restriction on the application of mixed ionic solvent systems which we have not examined in detail is provided by the operable pH limits in the aqueous phase. I t can readily be understood that a t very high or very low pH either deprotonation of the ammonium ion or protonation of the acid ion can occur, destroying the dual ionic nature of the system. The quaternary system, of course, should be operable in even strongly alkaline solutions. The lower pH limit will depend upon the acid group involved, and will be lower for phosphoric acids, for example, than for carboxylic acids. Consequently, the most immediate potential applications of these systems are for the processing of sea water concentrgtes, and other brine systems where the pH is within a few units of neutrality. Since the pH limits can be broadened, however, by proper choice of the components of the system, further development can be expected to provide additional applications in more acidic and alkaline solutions. literature Cited
Baniel, A., Blumberg, R., Alon, A., Brit. Chem. Eng. 1959, 223. Barrow, G. M., Yerger, E. A., J. Am. Chem. SOC.76, 5211 (1954). Chem. Eng. News 43, 70-1, April 19, 1965. Cuthbert, F. L., “Thorium Production Technology,” Addison-Wesley, Reading, Mass., 1958. Ellis, D. A., in “Uranium Ore Processing,” J. W. Clegg, D. D. Foley, Eds., Addison-Wesley, Reading, Mass., 1958. Gittus, J. H., “Uranium,” Butterworths, London, 1963. Grinstead, R. R., Davis, J. C., IND.ENG. CHEM.PROD. RES. DEVELOP. 8, 000 (1969a). Grinstead, R. R., Davis, J. C., J. Phys. Chem. 72, 1630 (1968); U.S. Office of Saline Water, Research and Development Rept. 320 (1968), 406 (1969b). Gustison, R. A., Pilloton, R. L., in “Columbium and Tantalum,” F. T. Sisco, E. Epremian, Eds., Wiley, New York, 1963. Hanson, C., Chem. Erg. 75 (18), 76 (1968). Hatch, M. J., Dillon, J. H., Smith, M. B., Ind. Eng. Chem. 49, 1812 (1957). Havighorst, C. R., Chem. Eng. 70, 228, (1963). Hudswell, F., Hutcheon, J. M., in “Extraction and Refining of the Rarer Metals,” Institution of Mining and Metallurgy, London, 1957. Johnston, J., Chem. Age 96, 71 (1966). McKay, H. A. C., Healy, T. V., Jenkins, I. L., Naylor, A,, “Solvent Extraction Chemistry of Metals,” MacMillan, London, 1965. Marcus, Y. A., Chem. Reu. 63, 139 (1963).
Martell, A. E., Calvin, M., “Chemistry of the Metal Chelate Compounds,” I’rentice-Hall, Englewood Cliffs, N. J., 1952. Smithson, G. R., Jr., Shea, J. F., Tewksbury, T. L., J. Metals 1966, 1037. Stary, J., “Solvent Extraction of Metal Chelates,” Macmillan, New York, 1964. Swanson, R. R., Dunning, H. N., House, J. E., Eng. Mining J . 162, No. 10, 110 (1961). Treybal, R. E., “Liquid Extraction,” 2nd ed., McGrawHill, New York, 1963. Vashist, P. N., Beckmann, R. B., Ind. Eng. Chem. 60, 43 (1968).
Weaver, B., Progr. Sci. Technol. Rare Earths 3, 129 (1968). Yerger, E. A., Barrow, G. M., J. A m . Chem. SOC.77, 6206 (1955). Zakarias, M. J., Cahalan, M. J., Trans. Inst. Mining Met. 75, C 245 (1966).
RECEIVED for review February 6, 1969 ACCEPTED May 12, 1969 Research supported by the Office of Saline Water, U. S. Department of the Interior, under Contract 14-01-0001-1134. 98th Annual Meeting, A.I.M.E., Washington, D. C., February 16-20, 1969.
EXPLORATORY PROCESS STUDY Base-Catalyzed Reaction of Organic Chlorides with Sodium Acetate I - D E R
H U A N G ’
A N D
L E O N A R D
D A U E R M A N ’
Department of Chemical Engineering, New York University, New York, N . Y . 10453 The heterophase acetylation reactions of organic chlorides with sodium acetate in the presence of various catalysts were investigated. There is a maximum catalyst concentration beyond which the reaction rate becomes independent of the catalyst concentration, the rate of reaction increases with the dielectric constant of the solvent, the rate of reaction is increased when bromide is added as cocatalyst, and when the chloride is conjugated the isomer ratio in the products is affected by the degree of branching on the carbon atom adjacent to the amine group. Several kinetic models have been analyzed as they relate to the results in this study and the observations of other investigators.
THEcatalytic acetylation of organic chlorides in heterophase reactions with sodium acetate is a class of reactions of considerable industrial importance. T o date, the kinetics is still controversial (Hennis et al., 1968; Merker and Scott, 1961; Ruggeberg et al., 1946; Yamashita and Shimamira, 1957). Work in this area has been reviewed in recent publications (Hennis et al., 1967, 1968). This study reports data which are of importance in process development and are pertinent to the kinetics. Experimental
Material. BENZYLCHLORIDE.Commercial grade benzyl chloride from Velsico was used throughout (except for runs A-70 to A-99, in which reagent grade benzyl chloride was used). Purity, 98.9%, moisture content, less than 0.1%. The reagent grade, from Matheson Coleman & Bell, has a moisture content of less than 0.06%. MYRCENE 85. [(CH&C =CH-(CH&-C(=CH~)C H =CHz]. Commercial grade myrcene was distilled a t 20 mm. of Hg, and Myrcene 85 was collected a t a pot temperature up to 90°C. AMYLCHLORIDE(1-chloropentane) , reagent grade from Eastman Kodak, b.p. 105”-07” C., purity 99.9+%. SODIUM ACETATE(anhydrous), commercial grade from Celanese Co., moisture content 1.0%. For runs A-70 to
’ Present address, Givaudan Corp., Clifton, N. J.
07014 address, Department of Chemical Engineering and Chemistry, Newark College of Engineering, Newark, N. J. 07102
* Present
A-99 reagent grade anhydrous sodium acetate was used (Matheson, Coleman & Bell). The moisture content was less than 0.1% by the Karl Fischer method. CATALYSTS, all C.P. grades from Eastman Kodak. Solvents. Xylene, b.p. 137-39, from Hess Oil & Chemical. Toluene, b.p. 110-11”, from Hess Oil & Chemical. Chlorobenzene, b.p. 130-2”, from Eastman Organic Chemicals. Nitrobenzene, congealing point 5.8“C., from American Cyanamid. Ethylene glycol, R.I. 1.4309; H 2 0 solubility loo%, from Union Carbide. Procedure. CONVERSIONOF BENZYLCHLORIDETO BENZYLACETATE.I n a 500-ml., round-bottomed, threenecked, jacketed flask equipped with a condenser, stirrer, and thermometer, were charged 1 mole of sodium acetate, 1 mole of benzyl chloride, and 0.01 mole of various types of amines with or without solvent. With stirring, the mixture was heated to the desired temperature in about 10 minutes by circulating hot oil through the jacket. A Lauda Circulator, Model K-2, constant temperature bath (0.1”C. precision) was used to control the oil temperature. The liquid layer of the reaction product was analyzed by gas chromatography (F & M Model 720 dual-column, 6-foot x %-inch i.d. Apiezon columns, column temperature 150”C.) and in some cases was analyzed by chlorine value. T h e agreement was within 2%. VOL. 8 N O . 3 S E P T E M B E R 1969
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