2814
INDUSTRIAL AND ENGINEERING CHEMISTRY
pendent of flow rate. Therefore, in the lower flow rate region, increasing the flow rate materially increases the yield per unit of equipment per dag, but this increase is less a t higher flow rates and must be balanced against increased cost of gas handling and separation of the desired product which becomes a smaller fraction of the effluent gas a t the high rates of flow. I n these experiments the flow rate range covered a t each temperature extends beyond the point where a further increase in the flow rate will result in a profitable increase in the space-time yield. In plotting these curves of Figure 7 the H , versus flow rate plot serves a useful purpose. Since that is a linear plot, it is easy to draw the straight line through a good average of the experimental points. Then from the average value of If, a t any flow rate a corresponding value for the space-time yield can be computed and smooth curves result. Further, some other value for the depth of the catilyst bed can be selected and similar curves drawn, from which the optimum flow rate can be estimated for the new catalyst bed. ACKNOWLEDGMENT
Vol. 41, No. 12
LITERATURE CITED (1)
Beckman. R. B.. Pufahl. A. E.. and Houxen. - 0. A.. IND.ENG. C H E M . , 35,558
(1943).
Brattain, R. R., Rasmussen, R. S., and Cravath, A. M., J . A p plied Phys., 14, 418 (1943).
Brunauer, S., Emmett, P. H., and Teller, E., J , Am. Chem. SOC., 60, 309 (1938).
Coggeshall, h’. D., and Saier, K.L., J . Applied Phys., 17, 450 (1946).
Coull, James, and Bishop, C. A . , Chem. Eng. Progress, 44, 443 (1948).
Dodd, R. H., and Watson. K. M.,Trans. Am. Inst. Chem. Engrs., 42, 263 (1946).
Hougen, 0. A , , and Watson, K. M., IND.Em+.CHEM.,35, 529.47 (1943).
Hurt, D. M., Ibid., 35, 522-8 (1943). Perry, J. H., ed., “Chemical Engineers’ Handbook,” pp. 790-1. 2nd ed., New York, McGraw-Hill Book Co., 1941. Thacker, C. M., Folkins, H. O., and Miller, E. L., IND.ENG. CHEM.,33, 584-90 (1941). Voge, H. H., and May, N. C., J . Am. Chem. Soc., 68, 550-3 (1946).
The authors are grateful to N. D. Coggeshall and E. L. Saier for spectroscopic analysis of the butene samples. The pelleted alumina catalyst waa kindly furnished by the Harshaw Chemical Company, Cleveland 6, Ohio.
RECEIVED January 1 4 , 1Y49. Abstracted from a portion of the dissertation presented b y Russell G . Hay to the Graduate School of the University of Pit,tsburgh in partial fulfillinent of the requirements for a Ph.D. degree. The work was done under a fellowship from the Gulf Research and Development Company, Pittsburgh 80, Pa.
Carbonation of Aqueous Sus ensions Containing Magnesium Oxides roxides d
ROBERT L. EVANS’ AND HILLARY W. ST. CLAIR’ United S t a t e s Bureuit of Mines, Salt Lake City, U t a h
olutions of magnesium bicarbonate having metastable concentrations more than twice the equilibrium concentration may be consistently prepared by leaching magnesium hydroxide, magnesium oxide, and calcined magnesite or dolomite with carbon dioxide and water. Metastable solutions cannot be prepared from magnesium carbonate or in the presence of precipitated magnesium carbonate or when the leaching temperature is much in excess of 30’ C. Formation of metastable solutions of maximum concentration requires that the solution be kept nearly saturated by carbon dioxide; the solids should not be added at a rate in excess of that at which an equivalent amount of carbon dioxide is absorbed. The magnesium goes into solution as magnesium hydroxide; the hydroxide ion reacts with dissolved carbon dioxide to give bicarbonate ion, thereby allowing dissolution of magnesium hydroxide to continue. The principal reaction involved in the breakdown of solutions of metastable concentration is that between bicarbonate and hydroxyl ions, a reaction that occurs in a very narrow zone at the solidliquid interface.
D
URING a study of carbonic acid leaching of magnesite and dolomite it was observed that, under certain conditions, a metastable solution of magnesium carbonate could be produced in which the concentration of magnesium was more than twice I 2
Present address, University of Minnesota, Minneapolis, Minn. Present address, U. S. Burrail of Lline8. CollPge Park, Md.
the equilibrium concentration. This obseivation led to an extensive investigation to determine the practicability of applying this phenomenon in leaching magnesia from magnesite and dolomite by an improved Pattison process ( 2 ) . Because of the importance of control of the metastable solution to the buccess of the process, an independent study was made t o clarify the mechanism by nhich the metastable solution rias formed and to establish the essential conditions required for producing and maintaining such a solution. BlETASTABLE SOLUTION O F MAGNESIA BICARBONATE
Magnesium oxide or magnesium hydroxide is readily dissolved by an aqueous solution of carbon dioxide; magnesium carbonate is also t>alreninto solution but a t a much slower rate. The solid phase with which the solution is in equilibrium is not crystalline magnesium bicarbonate but may be either magnesium hydroxide or the hydrated carbonate, ICIgC0~.3H~O.The solubility of magnesium bicarbonate increases rapidly as the partial pressure of carbon dioxide is increased up to a pressure of 3.85 X lou4 atmospheres or 0.28 mm. of mercury. Up to this point, the solid phase in equilibrium with the solution is magnesium hydroxide. As the partial pressure is increased further, the solubility of magnesium bicarbonate increases but a t a slower rate, and the stable solid phase is magnesium carbonate trihydrate, known by the mineralogical name, nesquehonite. These solubility relations are shown in Figure 1, ll-hich was drawn from data obtained by Kline ( 8 ) . The metastable solubility of magnesium bicarbonate is repre-
INDUSTRIAL A N D E N G I N E E R I N G C H E M I S T R Y
December 1949
1
i
m .000385
QI
PARTIAL PRESSURE OF
Cq
TIME,MINUTES
10
.I
2815
(ATMOSJ
Figure 2. Course of Carbonations"0f Different Amounts of -200-Mesh Caloined Magnesite
Figure 1. Solubility of Magnesium Oxide at 25 ' C. and Partial Pressure of Carbon Dioxide lo-' to 1 Atmosphere
sented by the dotted projection of the first leg of the solubility curve. Metastable solutions are represented by the area between the dotted line and the solubility curve for MgCO3.3H20. Metastable solutions of this nature were also observed by Terada (6) while studying the relative rates a t which magnesium hydroxide and magnesium carbonate were taken into solution by aqueous solutions of carbon dioxide. He likewise found that magnesia or magnesium hydroxide dissolves rapidly, but magnesium carbonate dissolves very slowly. CARBONATION EXPERIMENTS
L
The carbonation experiments were carried out in a 3.125-inch cylindrical vessel with a T-shaped, 1.875-inch, propeller-type stirring rod driven a t high speed. One-half liter of water was used in each experiment, and 10-ml. samples were removed during carbonation. Excess carbon dioxide was bubbled into the solution direct from the tank, without dilution, and a t a rate of about 2 cubic feet per hour under the prevailing atmospheric pressure of 630 to 650 mm. of mercury. ATo attempt was made to carry out the carbonations isothermally. Unless otherwise stated, the water used was initially a t about 20' C. The powdered sample of magnesium carbonate used in this work was C . P . grade, containing 42% magnesium oxide. The magnesium hydroxide was prepared by hydrating magnesium oxide; it contained 69% magnesium oxide and less than 2% carbon dioxide. The magnesite used in the experiment was from the deposit near Chewelah, Wash. ; the dolomite was from a very large deposit near Sloan, Nev. The analyses of the magnesite and dolomite, calcined a t 700" and 800" C., respectively, are as follows: Analyses, % Chewelah magnesite Sloan dolomite
MgO
CaO
COS
R~03
43.0 21.35
32.8
2.6
28.2 44.8
2.8
4.2
,.,
0.8
Insoluble
Three typical carbonation tests are shown in Figure 2; they were made with samples of high grade calcined magnesite from Chewelah, Wash. The magnesium content of the solution (in equivalent grams of h4gO per liter) is plotted against the time in minutes. Tests 1 and 2 show typical carbonations in which metastabIe solutions are produced. The concentration increases to about 11 grams of MgO per liter in about 30 minutes, after which it decreases to 8 or 9 grams per liter. After the maximum concentration of magnesium has been reached, the residue begins t o t8akeon a bulky appearance as magnesium is reprecipitated.
.I A
15
30
25'
8
30°
9
45
1
60
TIME, MINUTES
Figure 3. Carbonation of 7 Grams of Light Magnesium Oxide at Different Initial Temperatures
Chemical and x-ray analyses of the bulky residues show i h a t they contain the hydrated carbonate, MgC03.3H20. In test 3 the dissolution of magnesia ceased after 15 minutes, and the concentration of magnesium in the solution decreased sharply to a value less than 6 grams of MgO per liter. Dissolution then continued, and the magnesium concentrat,ion approached the final equilibrium value. This minimum in the curve is observed frequently when precipitation of the carbonate occurs prematurely. Tests 1, 2, and 3 were made with increasing amounts of suspended solids. Premature precipitation of the trihydrate, as in test 3, is more likely to occur when an excessive amount of solids is in contact with the solution. The proportion of solids to solution is an important factor influencing the stability of the metastable solution. TEMPERATURE OF CARBONATION. Three carbonation tests on 8-gram portions of magnesium oxide were made a t initial temperatures of 20°, 25', and 30" C. to determine qualitatively the effect of temperature of carbonation. No attempt was made to maintain a constant temperature during the reaction. The temperature rose 5" to 10' during each test because of heat releascd by the reaction. The results are shown in Figure 3. The solubility of MgC08.3HzO a t 20°, 25", and 30" is 10.1, 8.7, and 7.4
I N D U S T R I A L A N D ENGINEERING CHEMISTRY
2816
Vol. 41, No. 12
depending on the particle size) or by adding the solids a t a rate not in excess of t h a t a t which an equivalent amount of carbon dioxide is absorbed. 3. The metastable solution is formed only from the hydroxide or oxide and not from the carbonate. 4. The metastable solution breaks down by the precipitation of magnesium carbonate trihydrate (MgCOa.3Hz0)a t the bolid-liquid interface. Once any appreciable amount of carbonate is formed, the concentration of MgO drops t o the equilibrium concentration and the residue takes on a bulky flocculant appearance. 5 . Metastable solutions are very sensitive t o rise in tcmpera ture; their stability decreases markedly as the temperature r i s e above normal room temperature. 6. Calcium as well as magnesium forms a metastable bicarbonate solution. I n leaching calcined dolomite (calcium magnesium carbonate), the calcium goes into solution first, but it soon reaches a maximum and is reprecipitated. The magnesium usually attains its maximum concentration after the calcium has been reprecipitated. ~
Figure 4.
-1' hlINliTES 6, Carbonation 01 Rlagnesium Hydroxide and Magnesium Carbonate
grams of MgO per liter, respectively; n i t h increasing temperature there is less tendency to form metastable solutions At 3Q" C. the MgO Concentration leveled off a t the equilibrium concentration without passing through a maximum. CARBONATION OF MAGNESIUM HYDROXIDE AND MAGNESIUM CARBONATE. Comparative tests were made using increasing amounts of magnesium hydroxide; these tests are compared with carbonation of a sample of magnesium carbonate in Figure 4. The amount of magnesium hydrovide leached varied from 9 grams in test 10 to 15 grams in test 12. Here again there was the tendency for larger amounts of solids t o prevent attaining metastable concentrations. Furthermore, test 12 repeats the phenomenon observed in test 3 (Figure 2), in which the concentration of magnesium in solution goes successively through a sharp maximum and minimum before coming t o the final equilibrium Pidue. The concentration of magnesium during leaching of magnesium carbonate follows the usual course of a dissolved substance approaching saturation. There is no tendency to form a metastable solution. CARBONATION OF CALCINED DOLONITE SUSPENSION. Leaching tests on suspensions of calcined dolomite (Figure 5 ) show that metastable solutions of calcium bicarbonate, analogous t o those of magnesium bicarbonate, are formed. Thcy decompose and are followed by metastable concmtraiions of magnesium bicarbonate. The final solutions of these carbonations contain less than 0.01 gram of CaO per liter as compared with 0.5 gram of CaO per liter when calcined magnesite was leached. The results of leaching calcined magnesite, with and w t h o u t a small amount of calcined dolomite, are shown in Figure 6 The amount of calcium dissolved from impure calcined magnesite is reduced by addition of about 5% of calrined dolomite to accelerate the decomposition of metastable calcium bicarbonate solution ~
CHE-MICAL REACTIONS DURING CARBON4TIOh
The significant facts derived from experimental tests concerning the formation and dccomposition of metastable solutions of magnesium and cdrium bicarbonate mag be summaiizpd as folloT7~s:
1. Solutions of magnesium bicarbonate having concentrations
of more than 22 grams of RlgO per liter have been prepared by leaching magnesium hydroxide and oxide (or calcined magnesite) in water kept nearly saturated with carbon dioxide a t a partial pressure of about 1 atmosphere. The equilibrium concentration for the same temperature and partial pressure is only 8.6 grams of MgO per liter. 2. Formation of the metastable solution is favored b y a low ratio of solids t o solution (about 20 grams solids per liter or less
Solutions of bicarbonate may l w formed either from the carbon ate or the hydroxide,
Solutions of metastable concentration are formed from the hydroxide, oxide, and calcined magnesite and dolomite, but not from the carbonate. Undoubtedly the surface layers of the oxide or calcined carbonate are hydrated before reaction with carbon dioxide; therefore only the dissolution of Mg(0H)S need be considered,
16
2 /'
i!
,,-e:,,/.. , q 7 ,
Jq-
HYDRATED
17
TIME,MINUTES
Figure 5. Course of Carbonations of Hydrated and Unhydrated Calcined Dolomite
The foregoing observations suggest the folloRing explanation of how bicarbonate solutions of metastable concentrations are formed. The magnesium hydroxide dissolves and ionizcs, forming a saturated solution in contact Rith the solid:
Mg(0H)z = M g + "
+ 2OH-
(3aj
The conditions for saturation are defined by the solubility product:
[RIg++l [OIT-I~= 1.2
x
10-11
(3b)
When dissolved carbon dioxide is present in the solution, it reacts with the hydroxyl ion to form bicarbonate ion: OH-
+ [C0,Ja4= HCOa-
(4)
INDUSTRIAL AND ENGINEERING CHEMISTRY
December 1949
TIME, MINUTES
Figure 6. Course of Carbonation of 10 Grams of Calcined Magnesite, with and without Addition of 0.15 Gram of Calcined Dolomite
The essential function performed by the carbon dioxide is t o conBume hydroxyl ion and permit further dissolution of magnesium hydroxide. The limiting concentration of magnesium that may be attained by this reaction is represented by the dotted extension of the solubility curve for Mg(0H)g shown in Figure 1. As the bicarbonate ion accumulates in the solution another competing reaction must be considered-that is, the reaction between hydroxyl and bicarbonate ion: HCOI -
4-OH-
= HzO
+ cos--
(5)
The carbonate ion combines with the magnesium ion to form the carbonate as soon as the two ions reach saturation concentration as defined by the solubility product of magnesium carbonate: [Mg++] [COS--]
1.66 X loWa
(6)
The solubility curve i n Figure 1 shows that the equilibrium reaction (6) and the subsequent precipitation of magnesium carbonate, MgCOJ.3H20, may become the dominant reactions after the concentration of MgO in solution exceeds 0.55 gram of MgO per liter. However, the fact that metastable solutions are produced indicates that the reaction between hydroxyl and dissolved carbon dioxide under proper conditions can proceed at such a rate that it precludes the competing reactions by which the carbonate is precipitated. Since the reactions occurring during carbonation are heterogeneous reactions involving gaseous, liquid, and solid phases, rates of reaction depend to a great extent on diffusion across the interfaces between gas and liquid and between liquid and solid. Therefore, the reactions may be considered on the basis of the Nernst-Brunner theory of interfacial films as applied by Davis and Crandall (1 ). When the solution is saturated with carbon dioxide, the rate of carbonation is determined by the respective rates of diffusion of OH- and [COP],, through the stationary film of solution immediately in contact with the solid Mg(OH)2. This film may be divided into an acidic zone in which [C02]., is diffusing toward the solid and a basic zone in which OH- ion is diffusing away from the solid. The concentration gradients of each of these components decrease in the direction of diffusion, becoming virtually zero a t the reaction zone where the two components combine t o form bicarbonate ion. Any reaction between hydroxyl and bi-
2817
carbonate ions (Equation 5 ) will be largely restricted to the basic part of the film or that zone between the solid surface and the reaction zone, hence the importance of the relative amount of solid surface in contact with the solution. The concentration of hydroxyl ion is very low, even a t the surface of the solid magnesium hydroxide. T h e concentration for a saturated solution of magnesium hydroxide in pure water is only moles per liter; in a solution containing an equivaabout 3 X lent of 8 grams of MgO per liter, the common ion effect depresses moles per liter. the OH- concentration to only 2.4 X In an aqueous solution saturated with carbon dioxide the con0.3 X 10-2 moles per liter. Accentration of [ C O Z ]is~ about ~ cording to the stationary film theory, if equal rates of diffusion for OH- and [COZ],, are assumed, the ratio of the thickness of the basic zone to that of the acid zone in the film is less than Therefore, it may be concluded that the reaction between hydroxide and bicarbonate ions to form carbonate ion (Equation 5 ) is restricted to a very narrow zone a t the solid-liquid interface. Furthermore, this zone becomes more restricted as the carbonation reaction proceeds and the magnesium in solution increases. This explains why it is desirable to regulate the rate of addition of solids to be less than the rate of absorption of carbon dioxide if leaching is to be carried to metastable concentrations. The preceding discussion applies only t o a solid phase of magnesium hydroxide free from crystals of magnesium carbonate, MgC03.3Hz0. Once an appreciable amount of magnesium carbonate is present to act as nuclei for precipitation of the stable solid phase, carbon dioxide is abstracted from the solution very rapidly by precipitation of the carbonate, When this precipitation occurs a t a relatively low degree of supersaturation, the concentration of dissolved carbon dioxide may locally decrease below the saturation concentration so that the concentration of magnesium temporarily decreases to less than the equilibrium concentration for a saturated solution. This explains the minimum in the concentration curves observed for tests 3 and 10, shown in Figures 2 and 4,respectively. Heretofore the carbon dioxide in the solution has been considered t o occur partly as bicarbonate and partly as a simple solution of carbon dioxide in water, [COz]8q; no mention has been made of carbonic acid. The distinction between [COzIsa and metacarbonic acid was first suggested by McBain (4). On the basis of McBain’s data the true dissociation constant for carbonic acid was calculated by Thiel and Strohecker (6). Only a few per cent of the dissolved carbon dioxide is actually hydrated to form carbonic acid, which is in fact a strong acid. Consequently, although bicarbonate ion may be formed from ionization of metacarbonic acid, this path has been disregarded in favor of the more probable reaction between hydroxyl ion and dissolved COQ. ACKNOWLEDGMENT
The authors acknowledge the assistance of V. E. Bell in carrying out some of the laboratory tests, E. V. Potter for x-ray diffraction analyses, and H. E. Peterson and L. W. Butcher for chemical analyses of the samples. LITERATURE CITED (1) Davis, H. S., and Crandall, G. S., J . Am. Chem. Sac., 52, 3757-84 (1930). (2) Doerner, H. A., Holbrook, W. F., and Fortner, 0. W., U.S.BUT. Mines,Tech. Paper 684 (1946). (3) Kline, Walter, J . Am. Chem. SOC.,51, 2093-7 (1929). (4) McBain, J . Chhem. SOC.,101,814 (1912). ( 5 ) Terada, Kiyomatsu, Bull. Inst. Phys. Chem. Reseawh (Tokyo), 7 ,
452-65 (1928). (6) Thiel, A.,and Strohecker, E. R., Ber., 47,945,1061 (1914). RECEIVED February 18, 1949. Contribution from the Intermountain Experiment Station, Bureau of Mines, U. 9. Department of t h e Interior.
