Treat precipitation reactions as reactive crystallizations

Crystallization concepts can provide crucial insights for improving precipitate filterability.

By Tim Frank, Wayne Fort and Chris Jones, The Dow Chemical Company

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In disposing of phosphoric acid waste, simply neutralizing the waste with caustic and sending it to a wastewater treatment plant (WWTP) often is not an option because the resulting phosphate salts are soluble in water and the total amount of phosphate that can be discharged from the WWTP is strictly controlled for environmental reasons. Incineration methods are sometimes used, but they can be expensive and generally do not allow recovery of the phosphate content.

A number of years ago we implemented an alternative method involving reaction of phosphoric acid with hydrated or “slaked” lime (Ca(OH)2 particles slurried in water) to precipitate calcium phosphate solids. The processing scheme involved adding slaked lime to a batch of phosphoric acid in a stirred tank, and filtration of calcium phosphate solids from the resulting slurry; we opted for batch operation as it fit well into the specialty chemicals plant where the phosphoric acid waste was generated. Because calcium phosphate solids are only sparingly soluble in water at neutral pH, the filtrate from this process could be sent to the WWTP. The low product solubility not only allowed removal of more than 95% of the phosphate present in the feed, but also made the task of building large, filterable particles challenging. This article explains why and illustrates the importance of understanding and controlling solute supersaturation levels in a crystallization operation.

Chemistry and product solubility

The mix of calcium phosphate reaction products generated depends upon the ion speciation — that is, the type of ions present in solution, whether Ca2+, H2PO4-, HPO42-, or PO43-, among others. This depends upon pH and also upon whether the products of the various possible calcium phosphate reactions are fairly soluble and able to undergo further reaction in solution, or only sparingly soluble. A review of the literature [1, 2] indicates that hydroxyapatite, Ca10 (PO4)6(OH)2, is the primary product obtained at alkaline conditions, while dicalcium phosphate dihydrate, CaHPO4•2H2O, is the primary product at pH 2-4 and temperatures below 80°C. We focused our efforts on crystallizing CaHPO4•2H2O. Understanding the reaction product’s solubility behavior in the mother liquor is critical to understanding supersaturation, the driving force for crystallization. Figure 1 illustrates the solubility/dissociation behavior of CaHPO4•2H2O in water at pH 2-5 and 37°C, as estimated from solubility isotherms reported by Elliott [2]. Solubility decreases rapidly above a pH of about 2.5. This corresponds to the pKd value of the dissociation constant for that compound. At pH 7, the solubility of CaHPO4•2H2O is about 150 ppm.

Promoting crystal growth

The relationship between nucleation and growth provides the basis for understanding the conditions needed to promote crystal growth. For dicalcium phosphate, the supersaturation level can be expressed as:

 

where g denotes the activity coefficients for the ions in solution, [Ca2+] and [HPO42-] are the concentrations of ions in the supersaturated solution, and Ksp is the solubility product, which equals [Ca2+][HPO42-] at equilibrium at the given solution conditions. The numerator is the ionic activity in the supersaturated solution and the denominator represents thermodynamic equilibrium.

The nucleation rate often is defined as the number of crystals produced per unit time per unit volume of crystallizer. It generally follows a power-law relationship such that Nucleation Rate µsn(2).

The rate of crystal growth, often expressed as a change in a characteristic crystal dimension per unit time, is generally a more-linear function of supersaturation. For this reason, the growth of existing crystals will dominate nucleation at supersaturation levels below a certain critical level [3-6].

General control strategy. For the purpose of producing large, easily filtered crystals, control of a batch crystallization operation can be viewed in terms of initiating a nucleation event followed by controlling a growth period. The control strategy involves manipulating the operating variables that affect the rate at which supersaturation is generated — so as to first nucleate a small population of crystals and then grow those crystals with minimal additional nucleation. The approach to building large crystals is also relevant to producing higher purity crystals. How a given crystallization process is operated to generate supersaturation will depend upon the particular application and may include slow change in temperature, evaporation of solvent, addition of an anti-solvent, or reaction of reagents to generate a product concentration in excess of the product’s solubility in the mother liquor, as in the present example. After nucleation, supersaturation should be kept at a level that minimizes or at least reduces to an acceptable extent the nucleation of additional crystals.

Techniques to avoid high supersaturation. In carrying out precipitation reactions, special care is needed to limit as much as possible the generation of high localized supersaturation levels at the reagent feed point. In the example, the base is added to the reactor as a slurry of Ca(OH)2 particles suspended in water, so Ca(OH)2 particle size, the available solid surface area, and mass transfer to and from the solid surface — as well as the chemical kinetics and product solubility — play a role in generating supersaturation. (For discussions of the general kinetic aspects of solid/liquid reactive crystallization processes, see Refs. 7 and 8.) In general, the supersaturation level at the feed point may be minimized by improving circulation (without generating too much shear), using dilute reagents, slowing down the reagent addition rate, and operating under conditions that increase product solubility.

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