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Ind. Eng. Chem. Res. 2008, 47, 1002-1004
APPLIED CHEMISTRY Simple Method for Synthesis of Magnesite (MgCO3) Kristian Sandengen,*,† Leif O. Jøsang, and Baard Kaasa Department of Materials Science and Technology, Norwegian UniVersity of Science and Technology (NTNU), 7491 Trondheim, Norway
A simple method for synthesizing magnesite (MgCO3) from hydromagnesite (3MgCO3‚Mg(OH)2‚3H2O) at atmospheric conditions has been suggested. Monoethylene glycol (MEG) was used to decrease the water activity and avoid the hydrate phases of magnesium carbonate. An aqueous solution containing 95 wt % MEG was refluxed at 150 °C with continuous CO2(g) flow. This aided transformation of the metastable hydromagnesite into the thermodynamically stable magnesite. Introduction Transport of hydrocarbons and water in long subsea flow lines results in new challenges in the control of hydrate, corrosion, and scale formation. As the fluids cool, water condenses and gas hydrates may form. Hydrate formation can be inhibited by an antifreeze agent such as monoethylene glycol (1,2-ethanediol or MEG). MEG is completely miscible with water, but its presence lowers the solubility of most salts.1,2 Mineral precipitation may therefore occur, and such deposition of inorganic minerals from brine is called “scale”. During development of a scale model,2 the need to measure magnesite (MgCO3) solubility in water + MEG solutions arose. Magnesium carbonate powder, claimed to be magnesite, was consequently purchased from a commercial supplier. An X-ray diffraction (XRD) analysis of the material, however, showed that the salt was hydromagnesite (3MgCO3‚Mg(OH)2‚3H2O). It was therefore necessary to synthesize magnesite in the laboratory. The goal of the present work was to produce magnesite (MgCO3) of good purity in common laboratory glassware, i.e., without expensive steel bombs or other high-pressure vessels. Theory Solid magnesium carbonate can exist in several modifications, where the most common are lansfordite, MgCO3‚5H2O; nesquehonite, MgCO3‚3H2O; hydromagnesite, 3MgCO3‚Mg(OH)2‚ 3H2O; and magnesite, MgCO3. Ko¨nigsberger et al.3 used the notation 4MgCO3‚Mg(OH)2‚ 4H2O for hydromagnesite, while in this work the 3MgCO3‚Mg(OH)2‚3H2O notation of Marion4 was arbitrarily chosen. It is noted that the stoichiometry of the solid phase is not important for the discussion in the present work. The relative solubility3-5 in water is expected to change with temperature according to Figure 1. Magnesite has the lowest solubility in the whole temperature interval; i.e., it is the thermodynamically stable phase. Lansfordite and nesquehonite do not become more stable than * To whom correspondence should be addressed. Tel.: +47 74 86 20 00. E-mail:
[email protected]. † Present address: STATOIL Stjørdal, Strandveien 4, 7501 Stjørdal, Norway.
Figure 1. Schematic of relative solubility for various magnesium carbonate phases in an aqueous solution (PCO2 ∼ 1 atm) from literature data.3-5 2, Nesquehonite; [, magnesite; 0, lansfordite; s, hydromagnesite.
hydromagnesite for 0 < T (°C) < 80. It has been, however, stated3,4 that both lansfordite and nesquehonite are kinetically stabilized. Thus transformation to hydromagnesite does not occur at room temperature. In this work transformation denotes simultaneous dissolution of the less stable and precipitation of the more stable solid phase. Kinetics6 of dissolution/precipitation is generally faster at high temperature; hence both nesquehonite and lansfordite rapidly convert to hydromagnesite at temperatures above 5060 °C. Precipitation of magnesite is similarly inhibited; thus hydromagnesite will be stable in aqueous solution for a long period, even at temperatures approaching 100 °C. The solubility of magnesite at room temperature has been found to be about 2-2.5 mmol/kg of water.7 This is actually a much lower solubility than that of CaCO3, and comparable to that of FeCO3. Seawater is normally slightly supersaturated with CaCO3 but contains approximately 5 times more Mg2+ than Ca2+. Thus it should have a large magnesite supersaturation. Since MgCO3
10.1021/ie0706360 CCC: $40.75 © 2008 American Chemical Society Published on Web 01/16/2008
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is normally not formed, this supports the statement that the reaction is in some way inhibited. Investigation of the actual mechanism leading to such kinetic inhibition is, however, beyond the scope of this work. Magnesite Synthesis. The reaction from hydromagnesite to magnesite in an aqueous solution under CO2 gas bubbling can be written as
3MgCO3‚Mg(OH)2‚3H2O(s) + CO2(g) f 4MgCO3 + 4H2O(l) hydromagnesite + CO2 f 4magnesite + 4H2O
(1)
There are generally three parameters that can be varied to produce magnesite: CO2 partial pressure, water activity, and temperature. Both a higher CO2 partial pressure and a lower water activity are advantageous, since eq 1 will be driven to the right. Temperature is likely the most important factor, as an increase in temperature will increase the reaction rate. Langmuir7 reviewed carbonate stability in the MgO-CO2H2O system, and also summarized methods for magnesite synthesis. A more recent work by Deelman8 provides an interesting discussion concerning the mechanism of magnesite formation in nature, as well as a thorough summary of magnesite research history. Syles and Fyfe9 synthesized magnesite by reacting hydromagnesite and CO2 in closed glass ampules at 126 °C. They concluded that increased Mg2+ concentration tended to inhibit the reaction, while ionic strength and CO2 pressure acted as positive catalysts. From the literature it is evident that magnesite is easily produced from an aqueous solution, provided that a high temperature (>100 °C) is used. An elaborate study of salt solubility in water + MEG solutions2 has been performed in our laboratory. This experience with mixed-solvent solutions consequently gave rise to the idea of using MEG to lower the water activity and avoid the hydrate phases of magnesium carbonate. Intuitively the presence of MEG is advantageous for two reasons. First, it lowers the water activity, which should favor magnesite formation relative to the magnesium carbonate phases containing crystal water. Second, MEG increases the boiling point of the solution and a higher temperature can be reached at atmospheric conditions. Experimental Section A 95 wt % MEG solution was made by weighing MEG (p.a. 99.9%, Acros, 950 g) and distilled water (50 g) into a 1 L screwcap bottle. NaCl (1 mol/kg of solvent) was thereafter added because it increases the solubility of hydromagnesite and also reduces both the vapor pressure and the water activity. Solid hydromagnesite (Merck, z.a., 10 g) was added to a three-neck round flask containing the MEG + water + NaCl solution (250 g). Water-cooled reflux was utilized to ensure that the solvent did not evaporate. The mixture was magnetically stirred for 3 days under continuous CO2 bubbling (20-50 mL/min), with the flask being open to atmospheric pressure. A thermostat was used to keep the temperature at 150 ( 5 °C. The resulting precipitate was filtered using (Schleicher & Schuell, Black Ribbon paper) a Bu¨chner funnel, and washed with distilled water and absolute ethanol. Phase analysis was performed using X-ray diffraction (XRD). Results and Discussion Magnesite was successfully synthesized from hydromagnesite by refluxing a solution of 95 wt % MEG at 150 ( 5 °C, as
Figure 2. X-ray diffraction pattern of the hydromagnesite starting material with the intensity (counts) from Cu KR radiation plotted as a function of the 2θ angle (deg).
Figure 3. X-ray diffraction of the synthesized magnesite. The intensity (counts) from Cu KR radiation is shown as a function of the 2θ angle (deg). Literature data12,13 are indicated by bars.
described above. Figure 2 shows an X-ray diffraction (XRD) pattern of the hydromagnesite starting material, while Figure 3 shows the X-ray diffraction pattern of the synthesized magnesite. The latter showed no sign of hydromagnesite, which should appear with a strong peak at a 2θ angle of about 31°. At the bottom of Figure 3 there are given bars corresponding to magnesite in the International Centre for Diffraction Data databank,12 showing that the produced solid is clearly magnesite. For reference it is noted that the particular databank file shown in Figure 3 has been derived from the work of Catti et al.13 Temperature. It was desirable to keep the reaction vessel at a temperature slightly lower than the boiling point (∼163 °C) to maintain a certain CO2 pressure. CO2 (or CO32-) is needed to transform the hydroxide-containing hydromagnesite into magnesite. The total pressure always equals the atmospheric pressure, and the CO2 pressure (PCO2) is therefore given as the difference between the atmospheric pressure (Ptot) and the vapor pressure of the solvent (Psolvent). When approaching the boiling point, Psolvent f Ptot; thus PCO2 approaches zero. Because CO2 is continuously bubbled through the solution, there must be some gas present that gives a certain partial pressure. This will depend on the bubbling rate and cause some “steady state” condition, but PCO2 will certainly have the lowest value at the boiling point of the solvent. Raoult’s law provides a good estimate for the boiling point of a water + MEG solution.10,11 Significant amounts of NaCl present in the solution will increase the boiling point slightly. NaCl and MEG Concentrations. The solubility limit of NaCl in 95 wt % MEG is about 1.1 m (mol/kg of solvent) at room temperature. It is undesirable to precipitate NaCl if the
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solution is cooled; hence a NaCl concentration of 1 m was chosen in the present work. An increased MEG concentration yields a higher boiling point and a higher possible reaction temperature. The reaction was tried carried out also in pure MEG, but the resulting solid was gel-like, not easily filtered, and gave a typical amorphous pattern during XRD analysis. It is likely that some water has to be present in the solvent to convert the hydroxide-containing hydromagnesite into magnesite, as H2O participates in the CO2 dissociation reactions yielding the necessary carbonate. Any further discussion concerning this topic is beyond the scope of this work. A thorough discussion of CO2 equilibria and thermodynamics in water + MEG mixed solvent solutions can rather be found in Sandengen.2 Optimization. It is not necessary to wait until all the hydromagnesite has reacted at the elevated temperature. A simple second step could be to collect the solid material and stir it in saline water under CO2 bubbling. Any remaining hydromagnesite should then rapidly dissolve, leaving low soluble magnesite as the only solid. Conclusions Magnesite (MgCO3) was successfully synthesized in a water + MEG solvent at 150 °C with a CO2-containing atmosphere. Conventionally magnesite has been produced in aqueous solutions by applying an autoclave to reach temperatures higher than 100 °C. In this work monoethylene glycol (MEG) was used as cosolvent together with water to reduce the vapor pressure of the solvent. Hence a reaction temperature of 150 °C was achieved at atmospheric conditions and the reaction could be carried out in common laboratory glassware. Acknowledgment We thank STATOIL and Norsk Hydro for financial support.
Literature Cited (1) Kan, A. T.; Fu, G.; Tomson, M. B. Effect of Methanol and Ethylene Glycol on Sulfates and Halite Scale Formation. Ind. Eng. Chem. Res. 2003, 42, 2399-2408. (2) Sandengen, K. Prediction of mineral scale formation in wet gas condensate pipelines and in MEG (monoethylene glycol) regeneration plants. Ph.D. Dissertation, Norwegian University of Science and Technology, 2006; ISBN 82-471-8036-7. (3) Ko¨nigsberger, E.; Ko¨nigsberger, L.; Gamsja¨ger, H.; Low-temperature thermodynamic model for the system Na2CO3-MgCO3-CaCO3-H2O. Geochim. Cosmochim. Acta 1999, 63 (19/20), 3105-3119. (4) Marion, M. G.; Carbonate mineral solubility at low temperatures in the Na-K-Mg-Ca-H-Cl-SO4-OH-HCO3-CO3-CO2-H2O system. Geochim. Cosmochim. Acta 2001, 65 (12), 1883-1896. (5) Yanat’eva, O. K.; The metastable equilibrium in the system CaCO3MgCO3-H2O. IzV. Akad. Nauk SSSR 1960, 1, 180-182 (Engl. transl.) (6) Mullin, J. W. Crystallization, 4th ed.; Butterworth-Heinemann: Woburn, MA, 2001. (7) Langmuir, D. Stability of carbonates in the system MgO-CO2-H2O. J. Geol. 1965, 73, 730-754. (8) Deelman, J. C. Low-temperature formation of dolomite and magnesite, version 2.1; 2005; http://www.jcdeelman.demon.nl/dolomite/bookprospectus.html (accessed September 2007). (9) Sayles, F. L.; Fyfe, W. S. The crystallization of magnesite from aqueous solution. Geochim. Cosmochim. Acta 1973, 37, 87-99. (10) Trimble, H. M.; Potts, W. Glycol-water mixtures. Vapor pressureboiling point-composition relations. Ind. Eng. Chem. 1935, 27, 66-68. (11) Nath, A.; Bender, E. Isothermal vapor-liquid equilibriums of binary and ternary mixtures containing alcohol, alkanolamine, and water with a new static device. J. Chem. Eng. Data 1983, 28 (4), 370-375. (12) International Centre for Diffraction Data, Powder Diffraction File, Release 1999. (13) Catti, M.; Pavese, A.; Dovesi, R.; Saunders, V. R. Static lattice and electron properties of MgCO3 (magnesite) calculated by ab initio periodic Hartree-Fock methods. Phys. ReV. B 1993, 47 (15), 9189-9198.
ReceiVed for reView May 6, 2007 ReVised manuscript receiVed October 18, 2007 Accepted November 25, 2007 IE0706360
