Extractive Crystallization for the Production of Calcium Acetate and

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Extractive Crystallization for the Production of Calcium Acetate and Magnesium Acetate from Carbonate Sources Dionysios Dionysiou,*,† Marina Tsianou,‡ and Gregory Botsaris Department of Chemical Engineering, Tufts University, Medford, Massachusetts 02155

Extractive crystallization was employed for the preparation of deicing compositions of calcium acetate (CA) and magnesium acetate (MA) at room temperature. The design system comprises an aqueous phase as the source of calcium or magnesium ions and an organic phase as the source of acetate ions. The process was designed with the consideration that acetic acid is used in an organic phase as extracted from a fermentation broth. The conditions under which the extractive crystallization method resulted in the formation of single crystals or aggregates of calcium acetate and magnesium acetate were studied. Furthermore, the effect of acetic acid concentration in the organic and aqueous phases on the characteristics of the obtained crystals was investigated. Overall, the extractive crystallization process proved to be feasible at least in the laboratory and resulted in the production of well-formed, nonporous, large single crystals and clusters of either calcium acetate acetic acid hydrate [Ca(CH3COO)2‚CH3COOH‚H2O] or magnesium acetate tetrahydrate [Mg(CH3COO)2‚4H2O, R], depending on the type of salt (CaCO3 or MgCO3) used in each case. Introduction Calcium acetate (CA), magnesium acetate (MA), their double salt calcium magnesium acetate (CMADS), and mixtures of them (collectively designated as CMA) have been extensively reported as promising chemicals for environmental applications.1 For example, CMA is considered as the best alternative material for replacing the corrosive and environmentally unacceptable sodium chloride and calcium chloride deicers,2-4 as a powerful SO2, NOX, and toxic particulate emission control agent in coal combustion processes to reduce acid rain,5-7 and as an effective catalyst for the facilitation of coal combustion.8,9 CMA is also biodegradable, noncorrosive, and nontoxic. Extensive research in the last 10-20 years has indicated that CMA has minimal or no detrimental effect on vegetation, aquatic life, water reservoirs, construction materials, and automobiles.10-19 In an extensive study focusing on finding an alternative deicing material to sodium and calcium chlorides, only calcium acetate, magnesium acetate, and their combination were selected as the most competitive alternatives for replacing chloride deicing salts.2,3,20 Any of the CMA forms would be an effective deicer for most geographical areas. The selection of the chemical form will be dictated by economic criteria and by the calcium or magnesium sources that are available. The Cryotech CMA deicer, a commercial product, consists mostly of CMADS, as the mineral dolomite, a double carbonate salt of calcium and magnesium, has been used in the reaction with the acetic acid. Temperature criteria may also enter into the selection of the chemical form. The solubility of CA increases substantially as the temper* Author to whom correspondence should be addressed. † Current address: Department of Civil and Environmental Engineering, University of Cincinnati, Cincinnati, OH 452210071. ‡ Current Address: Physical Chemistry 1, Center for Chemistry and Chemical Engineering, University of Lund, P.O. Box 124, Lund S-221 00, Sweden.

ature is lowered below approximately 40 °C until -14.7 °C, which is the eutectic point for CA. On the other hand, the solubility of MA declines slowly below 0 °C, and its eutectic point was estimated by extrapolation to be approximately -29 °C.4,14 For this reason, MA is considered a more effective deicer than CA in areas where temperatures well below freezing are anticipated. Cryotech CMA deicer has a eutectic point of about -27 °C, which is lower than those of both NaCl and urea. Sodium chloride has a eutectic point of about -22 °C, whereas urea has a comparatively high eutectic point of approximately -11 °C. Urea has been employed as a deicer in certain cases, but its application is very limited. Because of its nitrogen content, urea exhibits eutrophic, corrosive, and toxic effects, and it can also be converted to nitric acid or ammonia.2,4 Although CMA deicers are commercially available and are used in certain U.S. highways where the use of chloride salts is prohibited because of their detrimental effects on environment and construction materials and utilities, the use of CMA in coal combustion technology has not yet been used in industry. However, the effectiveness of CA and CMA as catalysts for the combustion of coal, as well as their sulfur capture capability in such a process, has been the subject of many studies reported recently.5,6,8,9,21-24 Although CMA combines outstanding chemical and environmental properties, its large-scale application is impeded by its high cost. CMA costs more than $700 per ton, whereas sodium chloride costs only $30-40 per ton. Despite the fact that the real cost of sodium chloride use, including the corrosion effects on automobiles and bridges, is higher than its selling price, complete replacement of sodium chloride by CMA is not expected in the near future as long as the CMA deicer price remains high. The price could only be lowered by a drastic reduction in the price of the acetic acid (production by a biochemical route) and/or by changes in the current industrial process. The feasibility of a new process based on extractive crystallization that has the

10.1021/ie9906823 CCC: $19.00 © 2000 American Chemical Society Published on Web 09/20/2000

Ind. Eng. Chem. Res., Vol. 39, No. 11, 2000 4193

potential of reducing the cost of CMA production and that is more attuned to the biochemical source of acetic acid was the object of this investigation. Characteristics of the Current Industrial Process In the current industrial process, CMA is prepared from the reaction of low-grade acetic acid, produced from petroleum or natural gas, with dolomite lime.25 The process includes the following major steps: preparation of a solution that contains the reaction product (CMA); evaporation of the solution to dryness to produce a powder; and finally, pelletization of the powder to produce large pellets. The drying and pelletization steps employed in the current industrial processes are costly, and they considerably increase CMA production costs. Moreover, the drying process that follows the preparation of the slurry can lead to a product that contains unreacted acetic acid or dolomite lime. The latter can affect the quality of the final product. Finally, some physical properties of the final CMA pellets are not ideal and result in additional problems. These pellets are generally spherical and tend to bounce off roadways and pavements during application. In addition, they are porous, and their density is low. Consequently, they require a larger storage area, special handling during transportation, and more frequent deliveries during application. Features of Future Processes for CMA Production As has been stated above, for the production of a lessexpensive CMA deicer with better handling characteristics, there are at least three parameters that must be modified from the current industrial processes. First, the evaporation to dryness and pelletization steps must be avoided, as such steps further increase the production cost. Second, the method employed should lead to nonspherical and nonporous particles. Finally and most importantly, a less-expensive acetic acid should be used for CMA preparation. The high cost of acetic acid is the major reason for the high production cost of CMA. Such a less-expensive acetic acid can be produced by anticipated biochemical methods. The most important current processes for acetic acid production include carbonylation of methanol, liquidphase oxidation of hydrocarbons, and oxidation of acetaldehyde.26 To reduce the cost of acetic acid currently produced, other methods are under investigation. Currently, there is a renewed effort to develop inexpensive methods for acetic acid similar to those developed earlier, such as fermentation of alcohol and destructive distillation of wood. These inexpensive processes include fermentation of biomass, forestry residues, municipal wastes, and other byproducts (such as whey).26-29 Recovery of acetic acid from fermentation broth is a subsequent consideration for such processes. This process requires technology that is the same or similar to that required for the recovery of acetic acid from dilute aqueous waste or byproduct streams. The latter is mainly due to the low concentrations of acetic acid (e5 wt %) obtained in the fermentation broth. Various separation processes that have been developed for acetic acid recovery include liquid-liquid extraction, azeotropic distillation, and extractive distillation. In an excellent review of acetic acid extraction, King26 points out

that, from a cost point of view, extraction is the best choice for streams containing low to intermediate acetic acid concentrations (up to at least 35 wt %), whereas azeotropic distillation is the most favorable approach for streams containing high acetic acid concentrations. Novel chemically complexing extractants have recently been developed and studied for acetic acid recovery from fermentation broth and dilute aqueous solutions. These extractants are strong, organic Lewis bases, which can form acid-base complexes with acetic acid. Chemically complexing extractants usually have higher selectivity for acetic acid over water; have higher boiling points; are more expensive; and, in general, can obtain higher values of the equilibrium distribution coefficient KD.26,30 As discussed below, in our extractive crystallization process, the acetic acid will be in an organic phase consisting of the acid in extractant systems as removed from fermentation broth. Therefore, an important consideration for this study was the selection of an extractant system that can be widely used for acetic acid recovery from fermentation broth and dilute aqueous solutions. Selection of the most effective extractant system was based on many factors, including relative efficiency for acetic acid recovery, cost, physical interface behavior, toxicity, solubility in water, and tendency for side reactions. Detailed data for the screening procedure are given elsewhere.30,31 The extractant system that was selected comprised a mixture of Alamine 336 and 2-ethyl-1-hexanol in a volume ratio of 50/50. Alamine 336 (Henkel Corporation), a commercial tertiary amine with aliphatic chains with 8-10 carbons each, has been extensively studied and has been successfully shown to be one of the most effective extractants for acetic acid recovery from fermentation broth and dilute aqueous streams.32-36 The system Alamine 336 and 2-ethyl-1-hexanol (50/50 v/v) has also been used in our laboratory for the production of calcium and magnesium acetates from reagent acetate salts.30 The system showed good results, and it was not a limitation for the process. In a recent study for the extraction and sorption of acetic acid for calcium magnesium acetate formation, Reisinger and King37 indicated that Amberlite IRA-2 in octanol would make an even better choice than our selection. Nevertheless, it is worth noting that the current investigation is not limited in the selected extractant system used in the organic phase. Other extractants can also be used in the process, and their effectiveness can be evaluated in a comparative study. The latter was not an objective of the current study. This work is a laboratory investigation of an alternative process that satisfies the three parameters stated in the beginning of this section. The process utilizes a mixture simulating the acetic-acid-containing extractant that resulted from in situ solvent extraction of the acid from a fermentation broth. This loaded organic phase is brought into contact with an aqueous solution of calcium and/or magnesium ions. The acetic acid is reextracted into the aqueous phase, which becomes supersaturated with respect to calcium or magnesium acetate, and crystals of these acetates nucleate and grow. This is, in essence, a process of extractive crystallization. In this paper, only the first phase of the investigation is reported. The objective in this phase was, first, to demonstrate the feasibility of obtaining CA or MA crystals by extractive crystallization. In the industrial

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process, where supersaturation is created by evaporation of water, the acetates precipitate in the form of a gelatinous cloud that cannot be separated, which necessitates the continuation of the evaporation process to dryness. The chemical nature of the crystals produced under the various conditions was determined. The size of the crystals was also important. The study also demonstrated that crystals of a size suitable for the current methods of deicer spreading, i.e., roughly 0.30.7 cm, can be obtained. In addition, the study also identified the solution parameters that determine the precipitation of such crystals and the trends when the values of these parameters increased or decreased. It was not the purpose of the investigation to provide exact data for the design of a large-scale process. Such a process will be an extractive crystallization method using the same two phases in the study. However, a different means for the contact between the two phases will definitely be employed, which provides a larger interfacial contact area between the phases and increased mass-transfer rates. In view of these objectives, this study considers separately the production of either CA or MA. The investigation of the conditions for the preparation and characterization of mixtures of CA/MA and CMADS crystals using this process will be presented in a forthcoming publication. CMA Precursor Salts Calcium carbonate (CaCO3) is generally insoluble in water (0.00153 g of CaCO3, aragonite/100 mL water; T ) 25 °C),38 and it reacts with acetic acid, forming calcium acetate, carbon dioxide, and water, as shown in the following reaction:

CaCO3 + 2CH3COOH S Ca(CH3COO)2 + CO2 v + H2O (1) On the other hand, the solubility of calcium acetate is comparatively high [e.g., 24.8 g of Ca(CH3COO)2‚H2O/ 100 g of saturated solution at T ) 24.7 °C or 34.7 g of Ca(CH3COO)2‚2H2O/100 mL of water at T ) 20 °C) and decreases slightly with the temperature.30,38-40 Magnesium carbonate (MgCO3) is also insoluble in water [0.0106 g of MgCO3, natural magnesite/100 mL of cold water or 0.04 g of 3MgCO3‚Mg(OH)2‚3H2O, natural hydromagnesite/100 mL of cold water]. On the other hand, magnesium acetate, which is the reaction product of magnesium carbonate with acetic acid, has a comparatively high solubility in water [38.8 g of Mg(CH3COO)2/100 g of saturated solution at T ) 20 °C or 120 g of Mg(CH3COO)2‚4H2O/100 mL of cold water at T ) 15 °C].38,39 In contrast to the solubility of calcium acetate, the solubility of magnesium acetate increases with the temperature.30,39 The reaction of magnesium carbonate with acetic acid is shown below.

MgCO3 + 2CH3COOH S Mg(CH3COO)2 + CO2 v + H2O (2) Experimental Procedures As previously explained, the extractive crystallization system comprises a system with two phases: an aqueous phase as the source of calcium and magnesium ions and an organic phase as the source of acetate ions. The aqueous phase comprises a saturated solution of calcium

or magnesium acetates that is prepared by reacting acetic acid with calcium carbonate or magnesium carbonate. It should be noted that an actual process could start with an aqueous suspension of calcium or magnesium carbonate in contact with the organic phase. The acetic acid transferred to the aqueous phase would react with the carbonate to form acetate. The solution would initially become saturated with respect to the acetate and subsequently supersaturated, at which point crystallization would occur. Because our objective was simply to investigate the feasibility of crystal formation, we started with a saturated solution prepared before contact with the organic phase. The organic phase comprised a mixture of acetic acid and organic solvents, simulating a phase removed from a fermentation broth by solvent extraction. The mixture of Alamine 336 and 2-ethyl-1-hexanol (mixed at equal volume proportions), one of the most effective systems for acetic acid extraction from fermentation broth, was employed as the extractant system for our study. Materials. For the current study, reagent-grade calcium carbonate, CaCO3 (>99.0%); magnesium carbonate, MgCO3 (reagent-grade, assay: 40.0-43.5% as MgO); and glacial acetic acid, CH3COOH (99.7%), were purchased from Fisher Scientific, and 2-ethyl-1-hexanol (99.0%) was purchased from Fluka Chemica-Biochemica Co. Alamine 336 (a mixture of tertiary amines consisting of trioctyl- and tridecylamines, 95-97%) was purchased from Henkel Corporation. Equipment and Chemical Characterization. Crystals were photographed with a Minolta XG-M camera with macro lenses (50 mm) and Vivitar extension tubes, depending upon the size of the crystals. Filtration of the supersaturated solution was achieved by using Whatman and Marke Selecta, RFP (Germany) filters. Crystals were analyzed by X-ray diffraction (Nicolet, Model I 2/2000 with Cu KR radiation) in order to identify their chemical composition. Method. Saturated solutions of calcium acetate or magnesium acetate were prepared by reacting calcium carbonate or magnesium carbonate, respectively, with aqueous solutions of acetic acid. For that purpose, small amounts of a mixture of calcium carbonate (CC) or magnesium carbonate (MC) were added in a 400-mL beaker containing 200 mL of aqueous solution of acetic acid until the solution reached the saturation point. The mixture of salts and aqueous solution was mixed with a magnetic stirrer at room temperature (20 °C). The reason for adding small amounts of salts at each time was to avoid overflow of the reaction liquid due to the formation of large amounts of carbon dioxide (CO2). When the carbon dioxide was substantially removed, the beaker was covered with Parafilm to avoid evaporation of both water and acetic acid while the reactants were continually being mixed. The procedure for the addition of a small amount of salt was repeated until supersaturation conditions were achieved wherein a small amount of undissolved material could be seen in the beaker even 2 days after the last addition of salts. The above period of time was considered sufficient after preliminary observations during the experimental procedure. This procedure was followed by filtration of the saturated solutions, which was also conducted at room temperature. The filtration was repeated two or three times to ensure that no undissolved material was contained in the saturated solution.


Figure 1. Crystals obtained from the system containing Ca2+ with the following initial conditions: (a) A ) 20.70, B ) 46.1; (b) A ) 20.70, B ) 46.1; (c) A ) 31.02, B ) 12.5; (d) A ) 31.02, B ) 12.5; (e) A ) 31.02, B ) 56.2; and (f) A ) 31.02, B ) 56.2, where A is the concentration of acetic acid in aqueous phase (% wt/wt) and B is the concentration of acetic acid in organic phase (% wt/wt).

The procedure continued with the formation of the two-phase system. At this step, 5 mL of saturated solution were added in each of two series of 40-mL beakers. To each 5 mL of saturated solution was added 20 mL of organic phase via down-flow on the beaker’s wall, thereby avoiding any mixing of the two phases. The direct addition of organic phase to aqueous phase was avoided because of the observation that the formation of small droplets of one phase in the other phase occurred when the mixing of the two phases was not gentle. The organic phase was composed of a mixture of 2-ethyl-1-hexanol and Amine 336 (50/50 v/v) and glacial acetic acid. At this step, two series (the experi-

ments were performed in duplicate), each containing usually seven 40-mL beakers, were prepared, with each beaker containing 5 mL of saturated solution (aqueous phase) and 20 mL of organic phase in which the added volume fractions of acetic acid varied from 12.5 to 83.7 wt %. Control samples with only the aqueous phase or with the organic phase containing only the extractant system were also prepared. No formation, however, of crystals occurred in the control samples. The obtained crystals were collected after a period of 5 days, cleaned carefully to remove any attached amount of organic phase, and photographed. The crystals, after drying at room temperature in a desiccator,

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Table 1. Characteristics of Crystals Obtained from the System Containing Ca2+ acetic acid in aqueous phase (% wt/wt) acetic acid in organic phase (% wt/wt)

a

20.70

31.02 crystal size (cm)

crystal type

La

Wb

12.5

small needles

0.4-1.0

0.05-0.1

24.3

rodlike crystals & clusters

0.4-1.2

0.1-0.3

35.5


0.5-1.3

0.1-0.3

46.1


0.5-1.5

0.1-0.4

56.2


0.4-1.2

0.2-0.4

65.8

no crystals

-

-

75.0

no crystals

-

-

83.7

no crystals

-

-

crystal type small needles, thin rodlike crystals & clusters large rodlike crystals & clusters large rodlike crystals & clusters thin/thick rodlike crystals & clusters thin/thick rodlike crystals & clusters thin/small rodlike crystals & clusters thin/small rodlike crystals & small clusters thin/small rodlike crystals & small clusters


La

Wb

0.5-2.0

0.05-0.1

0.5-2.5

0.1-0.2

0.4-2.0

0.1-0.2

0.5-1.5

0.1-0.15

0.5-1.5

crystal size (cm)

crystal type

La

Wb

migration crystals

0.2-0.5

0.03-0.05

thin/long rodlike crystals & clusters rodlike crystals & clusters

0.5-2.0

0.1-0.15

1.0-1.5

0.05-0.2


0.2-1.5

0.05-0.2

0.1-0.15

thin rodlike crystals & clusters

0.2-1.5

0.05-0.1

0.3-1.0

0.05-0.1


0.6-0.8

0.1-0.2

0.3-1.0

0.05-0.1

no crystals

-

-

0.1-0.3

0.05-0.1

no system examined

Length. b Width or diameter.

were ground to a powder and analyzed by X-ray diffraction to identify their chemical composition. Results and Discussion (a) Calcium Carbonate as the Source of Ca2+. Nucleation and Growth: Size and Shape of Crystals. The first series of experiments was designed with the aqueous phase having 5.23 wt % of acetic acid for the reaction with the carbonate salt. The organic phase contained between 12.5 and 83.7 wt % of acetic acid. For concentrations greater than 75.0 wt % of acetic acid in the organic phase, formation of a two-phase system was not distinct. In this series of experiments, no crystals were obtained under all conditions. Formation of crystals also was not observed in the second series of experiments in which the acetic acid concentration in the aqueous phase was increased to 10.44 wt %. In the third series of experiments, the proportion of acetic acid in the aqueous phase was increased to 20.78 wt %. In this case, formation of crystals was observed for the systems containing 12.5-56.2 wt % of acetic acid in the organic phase. For the system containing 12.5 wt % of acetic acid in the organic phase, small needles were obtained, and their nucleation started 6-7 h after the two phases were brought into contact. The size of the obtained crystals increased with an increase in the concentration of acetic acid in the organic phase to the level of 56.2 wt %. Nucleation and growth of crystals under these conditions occurred in approximately 8-10

h after the preparation of the two-phase system. Crystals for 24.3-56.2 wt % of acetic acid in the organic phase had a rodlike shape, with sizes between 0.5 and 2.0 cm in length and 0.1 and 0.3 cm in diameter (or width). Clusters of crystals with similar dimensions were also formed under these conditions and are shown in Figure 1a,b. Similar results were obtained at a concentration of acetic acid in the aqueous phase of 31.02 wt %. Clusters of crystals obtained under these conditions are shown in Figure 1c-f. In this case, crystals were formed for all of the concentrations of acetic acid in the organic phase. However, at the very high concentrations (65.883.7 wt %), the obtained crystals were small and few in number. Very similar results were also obtained when the acetic acid concentration in the aqueous phase was increased to 41.16 wt %. In this case and for the system containing 12.5 wt % of acetic acid in the organic phase, formation of small migrate crystals was observed on the beaker wall in the organic-phase region. More data about the type, shape, and size of crystals obtained under all the conditions examined are shown in Table 1. The reproducibility of the experiments was very good, and such results are extensively reported elsewhere.31 In general, formation of crystals under certain conditions, crystal shape, and crystal size were almost the same for the repeated experiments. However, there were some differences with respect to the shape of the various clusters. This characteristic is very difficult to control


Figure 2. (a) Typical X-ray diffraction spectrum of crystals obtained from the system containing Ca2+ with A ) 31.02% wt/wt and B ) 24.3% wt/wt. (b) X-ray diffraction spectrum of calcium acetate acetic acid hydrate, (C6H10CaO6‚H2O); Ca(CH3COO)2‚CH3COOH‚H2O. Note: * ) calcium acetate hydrate, C4H6CaO4‚H2O, x ) calcium acetate hydrate, C4H6CaO4‚5H2O. See Figure 1 for definitions of A and B.

as it depends on the various defects on the initial crystal and the locations of the various crystals at the interface and in the saturated solution. These parameters do not seem to follow any specific trend and are rather random attributes. The data in Table 1 show the same trend with respect to crystal size for all three concentrations of acetic acid in the aqueous solution. Thin rodlike crystals are obtained when the acetic acid concentration in the organic phase is 12.5 wt %. As this concentration increases, the crystal size increases, and the length/ width ratio decreases. However, at high concentrations of acetic acid in the organic phase, the crystals again become small and rodlike. This trend can be explained by the following facts: At low concentrations of acetic acid in the organic phase, the driving force for the transfer of the acetic acid to the aqueous phase is low. Thus, the supersaturation created in the aqueous phase is low, the crystal growth rate will be low, and the resulting crystals will be small. As the concentration of the acetic acid in the organic phase increases, higher transfer rates, higher supersaturations in the aqueous phase, and higher growth rates will result, yielding larger crystals. However, as the amount of the transferred acetic acid increases at high acetic acid concentrations in the organic phase, the volume of the aqueous phase becomes larger. This dilution effect results in a lowering of the supersaturation and in the production of smaller crystals or even no crystals at all in concentrations of acetic acid in the organic phase higher than 75.0 wt %. The formation of clusters is of secondary importance for our investigation. A single crystal or a crystalline aggregate of the same size will be equally effective deicers. However, certain observations about the formation of these clusters should be noted. There are two types of aggregates. The first was an artifact resulting

from the experimental setup. The crystals were formed initially at the interface of the two phases, and then, when they were large enough (0.5 cm × 0.05 cm) to overcome the surface forces keeping them at the interface, they fell to the bottom of the beaker. At this point, some crystals were very close to each other, and as they continued to grow in the supersaturated solution, intergrowth occurred, resulting in cluster formation. The second was a crystal-growth phenomenon. They were actually spherulites resulting from a single nucleation event. Type and Hydration Level of Crystals. X-ray powder diffraction analysis showed that the obtained crystals were mostly calcium acetate acetic acid hydrate, Ca(CH3COO)2‚CH3COOH‚H2O. This conclusion can easily be reached through a comparison of the main peaks in Figure 2a,b. Figure 2a shows the X-ray spectrum of crystals obtained under these conditions, whereas Figure 2b shows the X-ray spectrum for calcium acetate acetic acid hydrate obtained from the JCPDS library (Hanawalt Search Manual and PDF-2 Database, International Center for Diffraction Data, Philadelphia, PA, 1993). In both cases, the main peaks are at ∼7.07, ∼7.30, ∼10.72, ∼14.51, ∼21.45, ∼23.45, and ∼25.95 2Θ (degrees). Small amounts of calcium acetate monohydrate, Ca(CH3COO)2‚H2O, and calcium acetate pentahydrate, Ca(CH3COO)2‚5H2O, were also formed under these conditions. Saury et al.40 found that the solubility of calcium acetate decreases with temperature and that there are three stable calcium acetate compounds in the temperature range from 0 to 180 °C. Calcium acetate monohydrate, Ca(CH3COO)2‚H2O, is crystallized below 58 °C, calcium acetate hemihydrate, Ca(CH3COO)2‚0.5H2O, is crystallized above 58 °C, and an amorphous phase is formed above 140 °C.40

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Figure 3. Solubility of calcium acetate in aqueous solutions of acetic acid at 25 °C. [Data were reported by Dunn and Philip (1934).41]

Figure 4. Crystals obtained from the system containing Mg2+ with the following initial conditions: (a) A ) 41.16, B ) 12.5; (b) A ) 51.20, B ) 12.5; (c) A ) 51.20, B ) 35.5; and (d) A ) 61.15, B ) 24.3. See Figure 1 for definitions of A and B.

Nevertheless, in the current investigation, the existence of acetic acid in the aqueous phase results also in the formation of complex salts with acetic acid. This result is in agreement with the data reported by Dunn and Philip,39,41 who found that calcium acetate acetic acid hydrate crystallizes from calcium acetate solutions when the concentration of acetic acid in the solution is higher than 10.8 g/100 g of saturated solution, whereas calcium acetate monohydrate crystallizes at concentrations of acetic acid below this value, as shown in Figure 3.

The possibility that the main form of CA might be the calcium acetate acetic acid hydrate has not been mentioned or taken into consideration in the investigations of acetates as deicers. This result of our investigation raises the question that remains to be answered: Does the existence of acetic acid in the deicer present a problem in the deicing application? (b) Magnesium Carbonate as the Source of Mg2+. Nucleation and Growth: Size and Shape of Crystals. In these experiments, the concentrations of acetic acid both in the aqueous phase and in organic phase were varied,

Ind. Eng. Chem. Res., Vol. 39, No. 11, 2000 4199 Table 2. Characteristics of Crystals Obtained from the System Containing Mg2+ acetic acid in aqueous phase (% wt/wt) acetic acid in organic phase (% wt/wt) 12.5

24.3

crystal type few big prismatic crystals (MA tetrahydrate) no crystals (in ten days)

La × Wb × Hc 1.0 × 0.5 × 0.3

-

no crystals (in ten days)

-

46.1

no crystals (in ten days) no crystals (in ten days) no crystals (in ten days) no crystals (in ten days)

-

65.8 75.0


35.5

56.2

a

41.16

-

crystal type small prismatic crystals (MA tetrahydrate) small prismatic crystals (MA tetrahydrate) small prismatic crystals (MA tetrahydrate) no crystals (in ten days) no crystals (in ten days) no crystals (in ten days) no crystals (in ten days)


La

Wb

0.3-0.8

0.2-0.4

0.2-0.6

0.2-0.4

0.2-0.4

0.1-0.2

-

-

-

-

-

-

-

-

crystal size (cm)

crystal type small prismatic crystals (MA tetrahydrate) small prismatic crystals (MA tetrahydrate) small prismatic crystals (MA tetrahydrate) no crystals (in ten days) no crystals (in ten days) no crystals (in ten days) no crystals (in ten days)

La

Wb

0.2-0.6

0.2-0.4

0.2-0.5

0.2-0.4

0.1-0.3

0.1-0.2

-

-

-

-

-

-

-

-

L ) Length. b W ) Width or diameter. c H ) Height.

Figure 5. (a) Typical X-ray diffraction spectrum of crystals obtained from the system containing Mg2+ with A ) 41.16% wt/wt and B ) 12.5% wt/wt. (b) X-ray diffraction spectrum of magnesium acetate tetrahydrate (R). Mg(CH3COO)2‚4H2O (R). See Figure 1 for definitions of A and B.

as in the case of the calcium carbonate experiments. For the series of experiments with 10.44, 20.70, and 31.02 wt % of acetic acid in the aqueous phase, no crystals were obtained for the entire range of acetic acid concentration in the organic phase investigated (12.5-75.0 wt %). In the case of 41.16 wt % of acetic acid in the aqueous phase, only a few large crystals were formed for the sample containing 12.5 wt % of acetic acid in the organic phase. These crystals had a prismatic shape and dimensions of about 1.0 cm × 0.5 cm × 0.3 cm. A representa-

tive crystal is shown in Figure 4a. When the concentration of acetic acid in the aqueous phase increased to 51.20 and 61.15 wt %, smaller well-formed prismatic crystals were obtained for the samples containing 12.5, 24.3, and 35.5 wt % of acetic acid in the organic phase. The shape of the crystals remained the same for all conditions. The dimensions of the obtained crystals were varied between 0.2 and 0.8 cm in length, between 0.1 and 0.4 cm in width, and between 0.5 and 0.3 cm in height. Crystals were first observed within 15-30 h after the preparation of the two-phase system. These

4200


Figure 6. The system magnesium oxide-acetic acid-water at 25 °C. [Data were reported by Iwaki (1914) and were obtained from Linke (1958)39.]

nucleation times are longer than those observed for the case of Ca2+ systems, and they are in agreement with the results obtained by Tsianou.30 The largest crystals were obtained in systems having 12.5 and 24.3 wt % of acetic acid in the organic phase. Smaller crystals were obtained for 35.5 wt % of acetic acid in the organic phase, and no crystals were formed in the systems containing higher acetic acid concentrations in the organic phase. Clusters of crystals were not generally formed. It was also noticed that migration crystals were not formed on the beaker’s wall in the organic-phase region. Figure 4 shows representative crystals obtained in these experiments. Detailed results for the type, shape, and size of crystal obtained for all of the conditions examined are shown in Table 2. The reproducibility of these results was even better than that for the case of the crystals obtained from Ca2+ solutions.31 Type and Hydration Level of Crystals. The type of crystals obtained under these conditions was also determined with X-ray powder diffraction. The obtained crystals were found to be magnesium acetate tetrahydrate Mg(CH3COO)2‚4H2O (R). This can easily be determined through comparisons of the X-ray spectra of all of the samples analyzed to the X-ray spectrum for magnesium acetate tetrahydrate (R) obtained from the database of the X-ray diffractometer (JCPDS). Figure 5a shows a representative spectrum of crystals obtained in the current study, whereas Figure 5b shows the spectrum of magnesium acetate tetrahydrate (R). Magnesium acetate tetrahydrate (R) belongs to the monoclinic crystallographic system with R ) γ ) 90°, β ) 95.37°, x ) 8.55 Å, y ) 11.995 Å, and z ) 4.807 Å. Comparison between the main peaks at ∼12.74 (110), ∼18.51 (001), ∼19.98 (011), ∼20.83 (200), ∼21.71 (-111), ∼22.15 (210), and ∼27.68 (031) 2Θ degrees of the spectra shown in Figure 5a,b can easily show that the obtained crystals are magnesium acetate tetrahydrate (R). This is also in agreement with the results reported by Tsianou30 in a similar study. This result is also supported by solubility data of the system magnesium oxide-acetic acid-water reported by Iwaki in 1914 (see ref 39). Iwaki showed that magnesium acetate tetrahydrate crystallizes from such a system when the concentration of acetic acid is higher than 14 g/100 g and lower than 36 g/100 g of saturated solution. These data are graphically presented in Figure 6. In various studies,

it was also reported that the solubility of magnesium acetate in water increases with temperature, and magnesium acetate tetrahydrate, Mg(CH3COO)2‚4H2O, is the only phase that crystallizes in the temperature range from 0.1 to 68 °C.30,39,42 Conclusions This study investigates an alternative approach to the production of CMA deicer. The method of extractive crystallization was employed for the production of, separately, calcium acetate and magnesium acetate from abundant and inexpensive carbonate salts, CaCO3 and MgCO3, respectively, as the first step. This method involves the formation of the desired product through crystallization in a two-phase system consisting of an aqueous phase as the source of Ca and Mg ions and an organic phase as the source of acetate ions. In this manner, the controlled production of well-designed materials can be achieved. Overall, the method of extractive crystallization demonstrated the feasibility of the production of well-formed single nonspherical crystals or aggregates of relatively large size under most of the conditions investigated. In the system in which CaCO3 was used as the reacting material, the process resulted in the formation of relatively large (0.2-1.5 cm) single rodlike crystals or aggregates of mostly calcium magnesium acetate acetic acid hydrate. In the case of MgCO3 as the reacting material, the process resulted in the formation of smaller single prismatic crystals of magnesium acetate tetrahydrate (R). Starting concentrations of acetic acid in the aqueous and organic phases of 20-40 and 2045 wt %, respectively, were found as the best conditions in the Ca2+ system. In the Mg2+ system, the best starting acetic acid concentrations were 50-60 and 2045 wt % in the aqueous and organic phases, respectively. The size and the form of the crystals (angular nonspherical and nonporous particles) indicate that they can be an excellent replacement for the spherical and porous pellets of the current industrial process and are more suitable for direct road application. However, because the calcium acetate crystals obtained in these experiments contain extra acetic acid in their lattice, the question remains to be answered whether that presents a problem in a possible deicing application. The


production of calcium magnesium acetate (CMA) crystals, in any form, from aqueous solutions containing both Ca2+ and Mg2+ ions is the subject of another publication.43 Acknowledgment This work was funded by the Center of Environmental Management (CEM) at Tufts University through grants from the U.S. Environmental Protection Agency. The authors acknowledge the assistance of Mr. Alan Yen and Mr. Tom Dunn in obtaining some of the X-ray diffraction spectra. Literature Cited (1) Calcium Magnesium Acetate: An Emerging Bulk Chemical for Environmental Applications; Wise, D. L., Levendis, Y. A., Metghalchi, M., Eds.; Elsevier Science Publishers B. V.: Amsterdam, The Netherlands, 1991. (2) Chollar, B.; Smith, D.; Zenewitz, J. The Involvement of the Federal Highway Administration with Calcium Magnesium Acetate. In Calcium Magnesium Acetate: An Emerging Bulk Chemical for Environmental Applications; Wise, D. L., Levendis, Y. A., Metghalchi, M., Eds.; Elsevier Science Publishers B. V.: Amsterdam, The Netherlands, 1991; pp 1-20. (3) Schenk, R. Alternative Road Deicer. In Calcium Magnesium Acetate: An Emerging Bulk Chemical for Environmental Applications; Wise, D. L., Levendis, Y. A., Metghalchi, M., Eds.; Elsevier Science Publishers B. V.: Amsterdam, The Netherlands, 1991; pp 37-48. (4) Schenk, R. Physical and Chemical Properties of Calcium Magnesium Acetate. In Calcium Magnesium Acetate: An Emerging Bulk Chemical for Environmental Applications; Wise, D. L., Levendis, Y. A., Metghalchi, M., Eds.; Elsevier Science Publishers B. V.: Amsterdam, The Netherlands, 1991; pp 21-35. (5) Levendis, Y. A.; Zhu, W.; Wise, D. L.; Simins, G. A. Effectiveness of Calcium Magnesium Acetate as an SOx Sorbent in Coal Combustion. AIChE J. 1993, 39 (5), 761. (6) Steciak, J.; Levendis, Y. A.; Wise, D. L. Effectiveness of Calcium Magnesium Acetate as Dual SO2-NOX Emission Control Agent. AIChE J. 1995, 41 (3), 712. (7) Shuckerow, J. I.; Steciak, J. A.; Wise, D. L.; Levendis, Y. A.; Simons, G. A.; Gresser, J. D.; Gutoff, E. B.; Livengood, C. D. Control of Air Toxin Particulate and Vapor Emissions after Coal Combustion Utilizing Calcium Magnesium Acetate. Resour., Conserv. Recycl. 1996, 16, 15. (8) Levendis, Y. A. Catalysis of the Combustion of Carbonaceous Particles (Synthetic Chars and Coal) by Addition of Calcium Acetate. In Calcium Magnesium Acetate: An Emerging Bulk Chemical for Environmental Applications; Wise, D. L., Levendis, Y. A., Metghalchi, M., Eds.; Elsevier Science Publishers B. V.: Amsterdam, The Netherlands, 1991; pp 221-252. (9) Sharma, P. K. Calcium Impregnation of Coals as a Means for Sulfur Emissions Control in Combustion. In Calcium Magnesium Acetate: An Emerging Bulk Chemical for Environmental Applications; Wise, D. L., Levendis, Y. A., Metghalchi, M., Eds.; Elsevier Science Publishers B. V.: Amsterdam, The Netherlands, 1991; pp 273-284. (10) Chollar, B. H.; Virmani, Y. P. Effects of Calcium Magnesium Acetate on Reinforced Steel Concrete. Public Roads 1988, 51 (4), 113. (11) Hiatt, G. F. S.; George, N. A.; Gushman, J. R.; Griffis, L. C.; Rausina, G. A. Calcium Magnesium Acetate: Cooperative Toxicity Tests and an Industrial Hygiene Site Investigation; Transportation Research Record 1157, Deicing Chemicals and Snow Control; Transportation Research Board, National Research Council: Washington, D.C., 1988; pp 20-26. (12) Horner, R. R. Environmental Monitoring and Evaluation of Calcium Magnesium Acetate (CMA); National Cooperative Highway Research Program, Report 305, Transportation Research Board, National Research Council: Washington, D.C., April 1988. (13) Kennelley, K. J.; Locke, C. E., Jr. Electrochemical Behavior of Steel in Calcium Magnesium Acetate. Corrosion 1990, 46 (11), 888. (14) Kennelly, K. J.; Locke, C. E. The Corrosivity and Electrochemical Behavior of a New Deicer, Calcium Magnesium Acetate.

In Calcium Magnesium Acetate: An Emerging Bulk Chemical for Environmental Applications; Wise, D. L., Levendis, Y. A., Metghalchi, M., Eds.; Elsevier Science Publishers B. V.: Amsterdam, The Netherlands, 1991; pp 129-151. (15) Locke, C. E.; Kennelley, K. J. Corrosion of Highway and Bridge Structural Metals by CMA; Report No. FHWA/RD-86-064; Federal Highway Administration: Washington, D.C., June 1986. (16) Pianca, F.; Carter, K.; Sedlak, H. A Comparison of Concrete Scaling Caused by Calcium Magnesium Acetate and Sodium Chloride in Laboratory Tests; Ontario Ministry of Transportation and Communications, MI-108: Downswille, Ontario, Canada, March 1987. (17) Slick, D. S. Effects of Calcium Magnesium Acetate (CMA) on Pavements and Motor Vehicles; Report No. FHWA-RD-87-037; Federal Highway Administration: Washington, D.C., April 1987. (18) Slick, D. Effects of Calcium Magnesium Acetate (CMA) on Pavements and Motor Vehicles. In Calcium Magnesium Acetate: An Emerging Bulk Chemical for Environmental Applications; Wise, D. L., Levendis, Y. A., Metghalchi, M., Eds.; Elsevier Science Publishers B. V.: Amsterdam, The Netherlands, 1991; pp 49-56. (19) Winters, G. R.; Gidley, J. L.; Hunt, H. Environmental Evaluation of Calcium Magnesium Acetate; Report No. FHWARD-84-094; Federal Highway Administration: Washington, D.C., June 1985. (20) Dunn S. D.; Schenk, R. U. Alternative Highway Deicing Chemicals; Report No. FHWA/RD-79-108; Federal Highway Administration: Washington, D.C., March 1980. (21) Ohtsuka, Y.; Tomita, A. Catalytic Gasification of Low Rank Coals with Calcium Acetate. In Calcium Magnesium Acetate: An Emerging Bulk Chemical for Environmental Applications; Wise, D. L., Levendis, Y. A., Metghalchi, M., Eds.; Elsevier Science Publishers B. V.: Amsterdam, The Netherlands, 1991; pp 253271. (22) Rising, B.; Hazard, H. Combustion of COM with Alkaline Acetate Additives. In Calcium Magnesium Acetate: An Emerging Bulk Chemical for Environmental Applications; Wise, D. L., Levendis, Y. A., Metghalchi, M., Eds.; Elsevier Science Publishers B. V.: Amsterdam, The Netherlands, 1991; pp 285-291. (23) Durych, A.; Laszuk, A.; Wiechawski, A. Modified Limestone Method for Removing Sulfur from Flue Gases. In Calcium Magnesium Acetate: An Emerging Bulk Chemical for Environmental Applications; Wise, D. L., Levendis, Y. A., Metghalchi, M., Eds.; Elsevier Science Publishers B. V.: Amsterdam, The Netherlands, 1991; pp 293-296. (24) Manivannan, S.; Wise, D. L. A Preliminary Evaluation of CMA for Sulfur Removal in Coal-Fired Boilers. In Calcium Magnesium Acetate: An Emerging Bulk Chemical for Environmental Applications; Wise, D. L., Levendis, Y. A., Metghalchi, M., Eds.; Elsevier Science Publishers B. V.: Amsterdam, The Netherlands, 1991; pp 297-317. (25) Todd, H. E., Jr. Deicing Compositions Comprising Calcium Magnesium Acetate Double Salt and Process for Their Manufacture. U.S. Patent No. 4,913,831, April 3, 1990. (26) King, C. J. Acetic Acid Extraction. In Handbook of Solvent Extraction; Lo, T. C., Baird, M. H. I., Hanson, C., Eds.; John Wiley & Sons: New York, 1983; pp 567-573. (27) Bungay, H. R.; Hudson, L. R. Calcium Magnesium Acetate from Biomass. In Biomass Conversion Technology: Principles and Practice; Moo-Young, M., Lamptey, J., Glick, B., Bungay, H., Eds.; Pergamon Press: New York, 1987; pp 189-192. (28) Palasantzas, I. A.; Wise, D. L. Preliminary Economic Analysis for Production of Calcium Magnesium Acetate from Organic Residues. Resour., Conserv. Recycl. 1994, 11, 225. (29) Wise, D. L.; Augenstein, D. An Evaluation of the Bioconversion of Woody Biomass to Calcium Acetate Deicing Salt. Sol. Energy 1988, 41 (5), 453. (30) Tsianou, M. Extractive Crystallization in the Production of Calcium and Magnesium Acetates for Environmental Applications. M. S. Thesis, Tufts University, Medford, MA, 1995. (31) Dionysiou, D. D. An Extractive Crystallization Process for the Production of Calcium Magnesium Acetate. M. S. Thesis, Tufts University, Medford, MA, 1995. (32) Althouse, J. W.; Tavlarides, L. L. Analysis of Organic Extractant Systems for Acetic Acid Removal for Calcium Magnesium Acetate Production. Ind. Eng. Chem. Res. 1992, 31, 1971. (33) Louie, E. C. Effect of Organic Solvents for Extraction of Acetic Acid During Fermentation by Clostridium Thermoaceticum. S. M. Dissertation, Massachusetts Institute of Technology, Cambridge, MA, 1985.

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Received for review September 9, 1999 Revised manuscript received July 25, 2000 Accepted August 6, 2000 IE9906823

Extractive Crystallization for the Production of Calcium Acetate and

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