Procedure to Obtain Hydromagnesite from a MgO-Containing Residue

The Mg(HCO3)2 solution can be obtained by carbonation of MgO-containing materials such as cyclone dust or byproducts generated during the calcination ...
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Ind. Eng. Chem. Res. 2000, 39, 3653-3658

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KINETICS, CATALYSIS, AND REACTION ENGINEERING Procedure to Obtain Hydromagnesite from a MgO-Containing Residue. Kinetic Study Ana I. Ferna´ ndez, Josep M. Chimenos, Merce` Segarra, Miguel A. Ferna´ ndez, and Ferran Espiell* Department of Chemical Engineering and Metallurgy, University of Barcelona, Martı´ i Franque` s, 1, 08028 Barcelona, Spain

A procedure for obtaining hydromagnesite is described. The process involves the formation of a Mg(HCO3)2 solution by carbonation of a MgO-containing residue slurry and further precipitation of hydromagnesite by addition of pure magnesium oxide as the precipitating agent. The solid residues generated can be landfilled and the water effluents discharged or reused at the plant. The optimum parameters for the precipitation were established, and the kinetic data were examined according to the equation for diffusive control. The activation energy diagram suggests a transition from diffusive to mixed control of the reaction rate. Under certain experimental conditions, an undescribed crystalline structure was obtained, and its stability and structural transformation into hydromagnesite were also investigated. Introduction Hydromagnesite, Mg5(CO3)4(OH)2‚4H2O, is commonly used as a source of magnesium oxide and, when its purity is suitable, for pharmacological use. Other applications such as use as an additive for pigments and papermaking are being developed. Its decomposition with increasing temperature also makes hydromagnesite interesting for use as a flame retardant agent for polymers. The basic magnesium carbonate of general formula xMgO‚yCO2‚zH2O is obtained industrially in different ways. Rosa1 described a procedure for obtaining basic magnesium carbonate by carbonation of calcined dolomite and further heating of the magnesium bicarbonate solution formed after its filtration. Ca´ceres and Attiogbe2 studied the separation of Mg/Ca carbonates from dolomite, finally recovering pure hydromagnesite from the aqueous residue. Manufacture of basic magnesium carbonate by mixing a basic magnesium carbonate suspension and a MgO suspension while bubbling with CO2 is described by Morie et al.3 Another commercial method for obtaining this chemical, described by Prakash and Gupta,4 is the formation of magnesium carbonate trihydrate by carbonation of Mg(OH)2 slurries and then boiling of the pulp. Cosic et al.5 described different stoichiometries of the basic magnesium carbonates obtained by varying the experimental conditions for carrying out the hydrothermal decomposition. The industrial procedures involving the addition of a precipitating agent to a magnesium salt solution using basic reagents such as sodium or potassium carbonates or bicarbonates are described in the bibliography. The magnesium salt solutions have been described as Mg* Author to whom correspondence should be addressed. Telephon: (34)934021316. Fax: (34)934021291. E-mail: [email protected].

(HCO3)26 or other salts such as magnesium nitrate, chloride, acetate, and sulfate. With this last method, the precipitation of basic magnesium carbonate with defined stoichiometry was not possible, and the properties of the magnesium oxide obtained afterward were strongly influenced by the preparation conditions.7 Langmuir8 has published a review of the stability of carbonates in the system MgO-CO2-H2O, and additional data for dypinguite, Mg5(CO3)4(OH)2‚5H2O, were obtained by Raade.9 This paper is focused on a nonpolluting process that involves the revalorization of a residue by obtaining a technical-grade hydromagnesite suitable for many applications. Particularly, it exposes the kinetic aspects of the formation of synthetic hydromagnesite, Mg5(CO3)4(OH)2‚4H2O, starting from a magnesium bicarbonate solution and adding MgO as the precipitating agent. The use of magnesium oxide as the basic precipitating agent, instead of alkaline carbonates, allows the formation of hydromagnesite without the introduction of foreign ions, avoiding a later purification operation. Thus, the water effluents generated do not cause an environmental impact when discharged and can be reused to obtain the magnesium bicarbonate solution. The process is actually being carried out on a pilot scale of 300 kg/day. Experimental Procedure In the described process, the Mg(HCO3)2 solution was obtained by carbonation at atmospheric pressure of a slurry, using as the raw material a calcined magnesite residue generated by the company Magnesitas Navarras (Navarra, Spain) with an average content of MgO of 4245%, the remainder being CaCO3, Fe2O3, and SiO2. The optimal conditions for leaching MgO slurries were previously studied.10

10.1021/ie0003180 CCC: $19.00 © 2000 American Chemical Society Published on Web 09/07/2000

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The kinetic precipitation experiments were performed in a 1-L cylindrical glass vessel reactor at atmospheric pressure, into which the previously obtained Mg(HCO3)2 solution was introduced. To avoid premature precipitation, the Mg2+ concentration in this solution was 4.2 g/L, and the solution was filtered and immediately used. The solution in the reactor was magnetically stirred, and the stirring started for all experiments just before the addition of MgO so as to avoid decarbonation and further precipitation from the solution. Both pH and temperature were measured during the reaction with a microprocessor pH meter. For all of the experiments, the initial pH and temperature were those from the fresh magnesium bicarbonate solution, that is, 7.3 and 30 °C, respectively.10 An increase in pH occurred during the reaction, whereas a slight decrease in the temperature was observed. To study the effect of temperature, a thermostatic water bath was used to keep the reactor temperature constant. The course of the reaction was followed by measuring the amount of CO2 and its related species, HCO3- and CO32-, in solution by titration with standardized 0.025 M HCl and phenolphthalein and methyl orange as indicators.11 At various time intervals, 2-mL aliquots of the slurry were taken, immediately filtered, and titrated with HCl. At the initial pH, most of the CO2 in solution is present as HCO3-; the rest, present as CO2(aq), was negligible after titration with standardized 0.1 M NaOH. Thus, the conversion was calculated as

X)

[HCO3 ]o

- ([CO23 ] + [HCO3 ]o

[HCO3 ])t

(1)

where [CO32-]t and [HCO3-]t were determined by titration at a time t and [HCO3-]o was determined by titration of the magnesium bicarbonate solution before the addition of precipitating MgO. The precipitating MgO used was obtained by calcining natural magnesite at 650 °C for 10 h and was analyzed by ICP-OES (inductive coupled plasma optic emission spectroscopy) and AA (atomic absortion) and found to have a content of 89.3% MgO, 3.8% CaO, 2.6% Fe2O3, and 3.0% SiO2. Commercial grade MgO (98%) was also used, and the results compared with those obtained with calcined magnesite. The MgO samples were ground, and the particle size fractions were achieved by sieving the solid between 200 and 63 µm using standard sieves DIN 4188. The particle size distributions, as well as the mean diameter, were determined with a Microtrac particle analyzer. The solid obtained was analyzed by X-ray diffraction. Elementary analysis of C and H by chromatography and thermogravimetric analysis (TGA) were also used. Results and Discussion The kinetics experiments were performed using calcined magnesite instead of commercial-grade MgO as the precipitating agent for several reasons. First, the effect of particle size can be evaluated by sieving the sample. For the commercial-grade MgO, this was not possible because of its very small particle size. Second, the commercial-grade MgO is strongly reactive, having a specific surface area, measured by the single-point BET method, of 120 m2/g compared with 40 m2/g for the calcined magnesite. Thus, with commercial-grade MgO,

Table 1. Chemical Analysis of the Impurities in the MgO-Containing Raw Material, the Precipitating MgO, and the Final Product Obtaineda impurities CaO (%) Fe2O3 (%) Al2O3 (%) SiO2 (%) a

raw material

precipitating MgO

hydromagnesite

4.5 1.6

1.80 0.05 0.15 0.35

1.34 0.052 0.030 0.18

2

Mg2+

Initial concentration, 4.2 g/L; temperature, 30 °C; pH 7.2; precipitating MgO, 10 g/L; and stirring speed, 500 min-1.

Figure 1. Effect of particle size on bicarbonate conversion (initial Mg2+ concentration, 4.2 g/L; temperature, 30 °C; pH 7.2; precipitating MgO, 8.9 g/L; and stirring speed, 500 min-1).

no appreciable differences were observed when the influences of various parameters were studied. In contrast, when using calcined magnesite that is less reactive, the effects of the different factors studied were measurable. The level of impurities in the magnesium oxide determines the applications of the obtained product. These impurites, however, were not considered in the study of the different parameters affecting the reaction rate. Otherwise, to evaluate the process for further industrial application, the concentration of impurities for both the initial MgO-containing raw material and the commercial-grade precipitating MgO were analyzed, as was the final product obtained. These results are shown in Table 1. Effect of Particle Size. The precipitation experiments were performed by adding 10 g/L of calcined magnesite (8.9 g/L MgO) to a Mg(HCO3)2 solution with an initial Mg2+ concentration of 4.2 g/L, at a temperature of 30 °C, initial pH of 7.2, and stirring speed of 500 min-1. A slight increase in the conversion was observed with the reduction of mean particle size, as is shown in Figure 1. A mean diameter (d50) of 114 µm was used for the rest of the experiments. Effect of Stirring Speed. The stirring speed has a strong influence on the apparent reaction rate. The initial experimental conditions were 4.2 g/L Mg2+, 8.9 g/L of precipitating MgO, 30 °C, and pH 7.2. The experiments performed for stirring speeds of 100-900 min-1 show that conversion increases with increasing stirring speed. The slope of the plot of Napierian logarithm of initial reaction rate vs Napierian logarithm of stirring speed (Figure 2) has been found to be 0.3, suggesting a diffusive or mixed control of the apparent reaction rate.12

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Figure 2. Effect of stirring speed on bicarbonate conversion (initial Mg2+ concentration, 4.2 g/L; temperature, 30 °C; pH 7.2; and precipitating MgO, 8.9 g/L with d50, 114 µm).

Figure 3. Effect of precipitating MgO concentration on the bicarbonate conversion (initial Mg2+ concentration, 4.2 g/L; temperature, 30 °C; pH 7.2; stirring speed, 500 min-1; and precipitating MgO with d50, 114µm).

Effect of the Concentration of Precipitating MgO. Various experiments were performed with the initial amount of MgO added being varied from 1.8 to 10.2 g/L. The initial experimental conditions were 4.2 g/L Mg2+ as Mg(HCO3)2, temperature of 30 °C, pH of 7.2, and stirring speed of 500 min-1. For this series of experiments, the maximum concentration of precipitating MgO was stated when an excess of MgO was observed by X-ray diffraction of the obtained solid. An increase in the conversion is observed with increasing precipitating MgO concentration, as depicted in Figure 3. Linear plots were obtained with the equation for diffusive control 1 - 2/3X - (1 - X)2/3 against time,13 where X is the conversion according to eq 1, as shown in Figure 4. Effect of Initial Concentration of Mg(HCO3)2. The initial concentration of Mg2+ was varied from 1.90 to 5.25 g/L, and 8.9 g/L of precipitating MgO was added for all the experiments, with an initial pH and temperature of 7.2 and 30 °C, respectively. No appreciable effect on the reaction rate was observed when the initial concentration of the Mg(HCO3)2 solution was increased. Although it is possible to obtain supersaturated magnesium bicarbonate solutions with higher magne-

Figure 4. Kinetic model for diffusive control 1 - 2/3X - (1 - X)2/3 against time (conversion data from Figure 3).

Figure 5. Effect of temperature on bicarbonate conversion (initial Mg2+ concentration, 4.2 g/L; pH 7.2; stirring speed, 500 min-1; and precipitating MgO, 8.9 g/L with d50, 114µm).

sium concentrations, these concentrations have not been tested to avoid precipitation of magnesium carbonate from the solution. Effect of Temperature. To study the effect of temperature, the magnesium bicarbonate solution temperature was kept constant by using a thermostatic bath and was varied from 6 to 50 °C. The initial experimental conditions were 4.2 g/L Mg2+, pH 7.2, stirring speed of 500 min-1, and 8.9 g/L of precipitating MgO. An increase in the conversion was observed with increasing temperature, as depicted in Figure 5. When the experimental temperature is 26 °C or lower, the conversion reaches a certain value and remains for during approximately 10 min; after this time, the conversion increases again. This first stage can be attributed to a stabilization of the system, as the measured pH in this period of time was the pH corresponding to the buffer HCO3-/CO32- (8.3). At higher temperatures, this buffer effect was not be observed. Kinetic Aspects. The integrated rate equation

2 1 - X - (1 - X)2/3 ) kt 3

(2)

is usually tested in heterogeneous systems in which the unreacted nucleus moves inward and a reacted diffusion layer of product is formed on the surface of the particle.

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Figure 6. Kinetic model for diffusive control 1 - 2/3X - (1 - X)2/3 against time (conversion data from Figure 5).

Figure 7. Arrenhius diagram for determining the activation energy (k data calculated from Figure 6).

Although it was derived for single spherical particles, it has been applied successfully in heterogeneous systems with isometric particles.14 Plots of this equation are straight for the whole period of time only for those experiments carried out at higher temperatures (40 and 50 °C). For the rest of the experiments, an induction period is observed, after which straight lines are obtained, as shown in Figure 6. The experimental rate constant k was calculated from the slope, and the Napierian logarithm of the rate constant k (min-1) versus 1/T is depicted in Figure 7. As can be observed from this figure, there is a transition from a diffusive to a chemical or mixed control of the process rate. The activation energy at each temperature can be calculated using the following empirical equation

Ea(T) )

6.80 × 104 - 200.75 T

(3)

where T is the temperature (in K) and Ea is the activation energy (in kJ/mol). Thus, the activation energy values for the reaction temperatures 40, 30, and 6 °C are 16.5, 23.7, and 43.0 kJ mol-1 respectively. To propose a sequence of reactions according to these results, variations in the concentrations of bicarbonate and carbonate ions were studied separately. During the first 10 min of reaction, the concentration of bicarbonate ions decreases, whereas the concentration of carbonate ions peaks until precipitation begins. The beginning of precipitation was observed visually at the reactor by the white coloration of the formed product, different from the reddish coloration of the precipitating MgO used (caused by the Fe2O3 impurity). From these results, two separate stages can be described. In the first, no precipitation is observed, the concentration of carbonate ions increases, and the pH increases slightly. This stage can be described by the following reactions

MgO(s) + H2O(l) S Mg(OH)2(s)

(4)

+ + OHMg(OH)2(s) S Mg(OH)(ads) (i)

(5)

OH(i) S OH(aq)

(6)

2+ OH(aq) S CO3(aq) + H2O(l) HCO3(aq)

(7)

where OH(i) is the OH- at the interface.

Figure 8. SEM image of the crystals of basic magnesium carbonate formed on the particle surface (×6000).

The maximum carbonate concentration was achieved under the following experimental conditions: particle size of MgO corresponding to the size fraction 90-63 µm, 8.9 g/L of precipitating MgO, 5.25 g/L [Mg2+] in the initial magnesium bicarbonate solution, 900 min-1 stirring speed, and 6 °C reaction temperature. These results suggest that diffusion of OH(i) , eq 6, through the boundary layer to the bulk solution is the rate-determining step.

The second stage begins when the concentration of carbonate decreases and presumably precipitation starts. Precipitation of basic magnesium carbonate can occur in two different ways: in the bulk solution if the Ksp is achieved (eq 8), 2+ 25Mg(aq) + 2OH(aq) + 5CO3(aq) + 4H2O S Mg5(CO3)4(OH)2‚4H2O(s) (8)

or on the particle surface, resulting in an unreacted nucleus that moves inward and a diffusion layer of product.

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Figure 9. X-ray patterns from the obtained basic magnesium carbonate and its transformation into hydromagnesite. HCO3(aq) S HCO3(i) 2HCO3(i) + OH(i) S CO3(i) + H2O(l)

(9) (10)

+ + nCO2(n + 1)Mg(OH)(ads) 3(i) + mH2O(l) S

nMgCO3‚mH2O(s) + (n - 1)OH(i) (11) OH(i)

S

OH(aq)

(12)

This last equation agrees with the observed increment of pH as the precipitation progresses. Thus, in accordance with the sequence of reactions proposed, the formation of a basic magnesium carbonate layer is limited by the diffusion of HCO3- ions from the bulk solution, as well as the OH- ions formed. An experiment was performed using a less-reactive MgO (calcined at 900 °C) and a reaction time of 10 min. The observation through an scanning electron microscope (SEM) of the recovered solid corroborated the presence of basic magnesium carbonate crystals formed on the particle surface, as can be observed in Figure 8. Product Stability. The temperature of formation determines the stability of natural basic magnesium carbonates8 as well as the synthetic ones. With the experimental conditions 4.2 g/L Mg2+, 8.9 g/L of precipitating MgO, and initial pH and temperature of 7.2 and 30 °C, respectively, the recovered solids presented an X-ray pattern from an undescribed crystalline structure, as shown in Figure 9A. A series of experiments involving changes in the temperature were carried out to evaluate the stability of this crystalline phase. It was transformed into hydromagnesite by heating the final pulp from 40 to 60 °C. X-ray patterns from the solid, at various intermediate temperatures, were obtained after

the sample was dried at 105 °C for 24 h, as shown traces B, C, and D of Figure 9. Using thermogravimetric, chemical, and elementary analysis, an approximate composition of this solid phase was calculated to be 0.39MgO 0.42CO2 0.19H2O. Conclusions The conclusions drawn from these findings are that satisfactory conditions for the production of hydromagnesite are achieved through the addition of MgO to a continuously stirred Mg(HCO3)2 solution, at room temperature, followed by heating of the final pulp at 55 °C. The Mg(HCO3)2 solution can be obtained by carbonation of MgO-containing materials such as cyclone dust or byproducts generated during the calcination of magnesite. This process follows a kinetic model for diffusive control. From the Arrenhius plot, it can be concluded that there is a transition from diffusive to mixed control of the process. The kinetics of the reaction, as well as the purity of the product, are strongly influenced by the magnesium oxide used. Thus, by using a reactive, high-purity MgO, a high-purity hydromagnesite can be obtained. The described process allows the production of hydromagnesite without the introduction of foreign ions, and therefore, the water effluents generated can be discharged or reused to obtain the magnesium bicarbonate solution. Literature Cited (1) Rosa, R. Proce´de´ de production de de´rive´s du magnesium, EP0732304A1, 1996.

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(2) Ca´ceres, P. G.; Attiogbe, E. K. Thermal decomposition of dolomite and extraction of its constituents. Min. Eng. 1997, 10 (10), 1165-1176. (3) Morie, T.; Kuroki, T.; Matsumoto, Y. Manufacture of porous spherical basic magnesium carbonate with narrow particle size. JP86-235240, 1986. (4) Prakash, S.; Gupta, K. N. Commercial production of Magnesium Chemicals From Industrial Waste Dust and Fines. Chem. Eng. World 1987, 22 (1), 39-41. (5) Cosic M.; Burevski, D.; Milkovic, M. Hydrated magnesium carbonates, their transformations and thermal decomposition to give magnesium oxide. Kem. Ind. 1994, 43 (2), 41-8. (6) Pattinson Polytech. J. Dinglers 1873, 209, 467. (7) Choudhary, V. R.; Pataskar, S. G.; Gunjikar, V. G.; Zope, G. B. Influence of preparation conditions of basic magnesium carbonate on its thermal analysis. Thermochim. Acta 1994, 232, 95-110. (8) Langmuir, D. Stability of carbonates in the system MgOCO2-H2O. J. Geol. 1965, 73, 730-754. (9) Raade, G. Dypinguite, a new hydrous basic carbonate of magnesium, from Norway. Am. Mineral. 1973, 55, 1457-1465.

(10) Ferna´ndez, A. I.; Chimenos, J. M.; Segarra, M.; Ferna´ndez, M. A.; Espiell, F. Kinetic Study of Carbonation of MgO slurries. Hydrometallurgy 1999, 53 (2), 155-167. (11) Kolthoff, I. M.; Sandell, E. B.; Meehan, E. J.; Bruckenstein, S. Ana´ lisis quı´mico cuantitativo, Sexta Edicio´n; Librerı´a y Editorial Nigel S.R.L.: Buenos Aires, Argentina, 1979. (12) Barret, P. Cine` tique He´ te´ roge` ne; Gauthier-Villars: Paris, 1973. (13) Sohn, H. Y.; Wadsworth, M. E. Rate Processes of Extractive Metallurgy; Plenum Press: New York, 1979. (14) Nun˜ez, C.; Espiell, F. The shape of the bodies and its consequences on the chemical attack of solids. Chem. Eng. Sci. 1986, 41 (8), 2075-2083.

Received for review March 13, 2000 Revised manuscript received June 27, 2000 Accepted June 29, 2000 IE0003180