MECHANISM OF PRECIPITATION OP ALUMINA PROM SODIUM ALUMINATE
Sept., 1945
sodium hydroxide by a current of air a t a boiling temperature is a clean-cut method of preparing manganosic oxide. 2. Oxidation a t room temperature invariably
[CONTRIBUTION FROM
THE
1449
leaves in the product an excess of oxygen amounting to a fraction of 1%. GOATHILL,LAMBERTVILLE, N. J.
RECEIVED FEBRUARY 14, 1945
CHEMICAL RESEARCH DIVISION, DEPARTMENT OF CHEMICAL ENGINEERING, TENNESSEE VALLEY AUTHORITY, WILSON DAM,ALABAMA]
Mechanism of the Precipitation of Alumina from Sodium Aluminate Solutions BY KELLYL.ELMORE, CHARLES M. MASONAND JOHN D. HATFIELD In several processes for the production of alumina, alumina trihydrate is precipitated by treating sodium aluminate solution with carbon dioxide. The present paper describes an investigation of the PH of sodium aluminate solutions during precipitation of alumina and the determination of the mechanism of precipitation a t 60'.
Theory Alumina in' sodium aluminate solutions is believed to occur as NaA102.' The form precipitated from these solutions in the present work was found by X-ray analysis to be gibbsite, A1203.3H20; this finding is in agreement with the observations of Laubengayer and Weisz.2 The precipitation reaction therefore may be postulated as 2A101-
+ 2Ha0 +AIz0~~3HzO
(I)
This reaction is represented by the rate equation - dC,,o,- / d t = k'Ci10,- Ckso+
(3)
where CAIO, - represents the concentration of aluminate ion in moles per.liter, and t the time in hours. The velocity of the reverse reaction is negligible, since the solubility of gibbsite is small and is practically constant in the pH range of this investigation. At constant pH equation (3) would reduce to - dCuo,-/dt
kCi10,-
(3)
which is the simplest form of the differential equation for a bimolecular reaction. Integration yields 1/CA1o2- = kt const. (4) froin which it is apparent that a plot of 1 / c A I O ~ versus t a t constant PH should be a straight line. The concentration of aluminate ion in moles per liter cannot be calculated exactly from the concentration of alumina found by analysis of the solution, because the analytically determined value represents the sum of the gibbsite, with which the solution may or may not be supersaturated, plus the aluminate ion. It is possible, however, to calculate the change in concentration
+
(1) J. W. Mellor, "A Comprehensive Treatise on Inorganic and Theoretical Chemistry," Vol. 5 , Longmans, Green & Co., London,
..
1929. D. 285.
(2) A. W. Laubengayer and R . S . Weisz, THISJ O U R N A L , 66,'247249 (1943).
of aluminate ion from the carbon dioxide consumption, because the carbon dioxide added is directly responsible for the conversion of aluminate ion to gibbsite. Let A represent the number of moles per liter of aluminate ion initially present in the solution, x the number of moles per liter of carbon dioxide reacting with the aluminate ion, and y the number of moles per liter of carbon dioxide reacting with the carbonate ion to form bicarbonate ion. If the equation for 'the formation of carbonic acid COI
+ 3H20 = 2H10' + cos-
(,-I 1
is added to equation (l),the equation
+
2A1023Hz0 ( A - 2x)
+ COz = A1203~3HzO+( xC01= - y) X
(61
X
is obtained for the carbon dioxide used to precipitate the alumina. For the carbon dioxide used in converting carbonate to bicarbonate, the equilibrium equation may be written COa' (x
+ COz + HzO
- Y)
2HCO1-
y
(7 1
2Y
If the carbon dioxide consumed in converting carbonate to bicarbonate is applied as a correction to the total carbon dioxide consumed, the remainder should represent the amount required for conyersion of aluminate ion to gibbsite. The y, represents the total number of moles sum, x per liter of carbon dioxide added to the solution. For the second ionization of carbonic acid the equations are :
+
HCOa-
+ Hz0 = COi' + H30+
18)
where f is the stoichiometric ionic molar activity coefficient. Substitution of the equalities ccos- = Y - .y LSRros- = 2y
(10) (111
into equation (9) yields the equation
By extrapolation of the values of Harned and Scholes3 for lower temperatures, Kz a t 60' was estimated to be 7.35 X 10-l'. Activity coefficients, f,for carbonate and bicarbonate ions were (8) H. S. Harnrd and S. R . Scholes, i b i d . , 6S, 17013-1709(1941).
KELLY L.
I450
ELMORE, C H A R L E S
M.
obtained from Walker, Bray and J o h n ~ t o n the ,~ ionic strengths of the solutions having been calculated to be in the range O.(i to 0.7. Since the ratio of these ionic activities was employed, small differences in the values selected for the activity coefficients would have little effect on the result. Activity of the water was calculated from Iiaouit's law. T
