smallness of these effects. Nevertheless, the results shown should be interpreted in terms of indicative trends rather than in terms of specific numbers.
Literature Cited Air Pollution Control District (County of Los Angeles), “Profile of Air Pollution Control in Los Angeles County,” 1969. Air Resources Board (Calif.), “Air Pollution Control in California: 1970 Annual Report,” 1971. Air Resources Board (Calif.), Los Angeles County Air Pollution Control District, and Western Oil and Gas Association, “Gasoline Modification: Its Potential as Air Pollution Control Measure in Los Angeles County,” p 28, 1969. Air Resources Board (Calif.), “Test Procedures, Emission Standards Revised,” Calif. Air Res. B d . Bull., Vol. 2, No. 9, p 4, Nov. Dec., 1970. “DHEW Urban Dynamometer Driving Schedule,” Fed. Regist., 35 (219),Part I1 App. A, 17311 (1970). “EPA Regulations,” ibid., 36,22456 (1971). Eschenroeder, A. Q , , Martinez, J . R., “A Modeling Study to Characterize Photochemical Atmospheric Reactions in the Los Angeles Basin Area,” General Research Corp. CR-1-152, 1969. Eschenroeder, A. Q., Martinez, J . R., Aduan. C h e m . S e r . , 113, 101-68 (1972). Eschenroeder, A. Q., Martinez, J. R., “Analysis of Los Angeles Atmospheric Reaction Data from 1968 and 1969,” General Research Corp. CR-1-170, 1970. Eschenroeder, A. Q., Martinez, J . R., “Further Development of the Photochemical Smog Model for the Los Angeles Basin,” General Research Corp. CR-1-191, 1971. Available from Nat. Tech. Inform. Serv. Springfield, Va., Document No. PB 201737.
Gay, B. W., Bufalini, J. J., Enuiron. Sci. Technol., 5,422 (1971). Gordon, R., Mayrsohn, H., Ingels, R., ibicl., 2, 1117 (1968). Hocker, A. J., California Air Resources Board, 9528 Telstar Ave., El Monte, Calif. 91731, private communication, 1971. Huls, T., Environmental Protection Agency, Air Pollution Control Office, 2929 Plymouth Road, Ann Arbor, Mich. 48105, private communication, 1971. Kearin, D. H., Lamoureaux, R. L., “A Survey of Average Driving Patterns in the Los Angeles Urban Area,” System Development Corp. TM-(L)-4119/000/01 (1969). Martinez, J . R., Nordsieck, R. A,, Eschenroeder, A. Q., “Morning Vehicle-Start Effects on Photochemical Smog,” General Research Corp. CR-2-191, June 1971. Available from Nat. Tech. Inform. Serv., Springfield, Va., Document No. PB203872. McGraw, M. J., Duprey, R. L., “Air Pollutant Emission Factors” (preliminary document), p 24, U.S.Environmental Protection Agency, 1971. Neiburger, M., Edinger, J. G., “Meteorology of the Los Angeles Basin with Particular Respect to the ‘Smog’ Problem,” Air Pollution Foundation Report So. l , 1954. Roberts, P. J. W., Roth, P. M., Nelson, C. L., “Development of a Simulation Model for Estimating Ground Level Concentrations of Photochemical Pollutants,” Appendix A, Systems Applications, Inc., Report 71-SAI-6, 1971. Scott Research Laboratories, “Final Report on Phase I, Atmospheric Reaction Studies in the Los Angeles Basin,” Vols. I and 11, 1969. Receiced for reuieu March 7, 1972. Accepted June 15. 1973 Work supported by the Air Pollution Control Office of the Encironrnental Protection Agency under Contract N o . EHSD-71-22.
Kinetics of Sulfation of Reduced Alunite N. Gopala Krishnan and Robert W.
Bartlett’
Process Metallurgy Group, Department of Applied Earth Sciences, Stanford University, Stanford, Calif. 94305.
The United States contains more than 500 million tons of the mineral alunite, K~S04.3Alz(SO4)3.6HzO, which is a potential regenerable absorbent for removing SO2 from dilute effluent gas streams. The alunite must be desulfurized to Kz0.3A1203 prior to sorption in a process step identical with absorbent regeneration. The sulfation kinetics of hydrogen desulfurized alunite were determined between 100” and 500°C in SO2 concentrations from 0.10.5%. Sulfation below 300°C is limited to the reaction with KzO. Above this temperature, A1203 is also sulfated, resulting in the formation of potassium aluminum sulfate, and the rates can be expressed as parallel processes. Each sulfation rate is proportional to the SO2 concentration and the remaining unreacted oxide, KzO and A1&, respectively. The experimental activation energy is low. The sulfation of K 2 0 is independent of oxygen pressure and the rate-controlling step appears to be formation of potassium sulfite which is subsequently oxidized to sulfate. The sulfation rate of A1203 is proportional top(02)’Iz. H
Fuel combustion is the chief source of air pollution by sulfur dioxide, and about 18 million tons of sulfur dioxide are being emitted annually by power generating plants in the Cnited States (Bienstock and Field, 1960). Both liquid scrubbing techniques and sorption by dry solids are
* To whom correspondence should be addressed.
being considered for the removal of sulfur dioxide that occurs in dilute concentration in power plant gases. Dry, elevated temperature sorption processes retain the buoyancy of the stack gases for maximum dispersion in the atmosphere after treatment. Among the dry processes are absorption by metal oxides to form sulfites or sulfates, with subsequent regeneration of the oxide and recovery of sulfur in a concentrated form. Alkalized alumina (NazO.Al203) developed by the U.S. Bureau of Mines is an excellent regenerable sorbent for sulfur dioxide (Bienstock et al., 1967). The mineral alunite, a hydrated potassium aluminum sulfate (Kz0.3A1203-4S03.6HzO) is chemically similar to the final sulfated product of alkalized alumina. Heating alunite a t 700°C decomposes the aluminum sulfate to oxide (Fink et al., 1931). Hydrogen reduction converts all the sulfate to hydrogen sulfide and oxides. The oxides thus generated react with sulfur dioxide and oxygen t o form sulfates. Alunite is found in several locations in the western United States and the estimated reserves recoverable by inexpensive surface mining are over 500 million tons ( E n g Mining J , 1972). The purpose of the present study was to investigate the sulfation kinetics of hydrogen desulfurized alunite.
Materials and Experimental Method Pelletized alunite from Marysvale, Utah, was used in this study. The mineral was crushed t o -100 mesh, mixed Volume 7, Number 10, October 1973
923
with 5% bentonite and 10% water, and pressed a t 1000 psi. The M-in. disc pellets were then dried a t 525°C and stored. Pellets were thermally decomposed or reduced with hydrogen in small batches in a horizontal tube furnace as required for the sulfation kinetic studies. ' The absorption of sulfur dioxide by hydrogen-reduced alunite pellets was investigated under various isothermal and constant atmospheric conditions using a quartz helix balance from which the specimens were suspended. The thermogravimetric analysis was supplemented by X-ray diffraction, chemical analyses, and electron beam microprobe analyses. Surface area measurements were made on the pellets before and after reaction, and scanning electron microscopy was used to observe the fractured surfaces. Optimum Hydrogen Reduction of Alunite and Sorbent Regeneration Chemical analyses showed that the sulfur content of alunite reduced in hydrogen a t 800°C for 60 min was about 2% while that of calcined alunite was about 9%. The heating decomposes only the aluminum sulfate, while the presence of hydrogen is necessary for the reduction of potassium sulfate. This was also confirmed by X-ray diffraction of the pellets which exhibited strong peaks of K2S04 for calcined pellets while the reduced alunite was amorphous. Since the SO2 absorption capacity of reduced alunite is 2.5 times t h a t of the calcined alunite, the remaining study emphasized reduced alunite. A series of SO2 absorption experiments using alunite previously reduced at various temperatures showed that the SO2 absorption rate increases with increasing reduction temperature to 800°C. There was no apparent increase between 800" and 900°C and reduction a t 1000°C reduces the absorption capacity. Although the rate of reduction increases with the temperature, the rate of volatilization of K z S 0 4 and permanent loss of potassium also increases and 800°C was the optimum reduction temperature. There was no increase in the reduction and absorbent effectiveness if the reduction period a t 800°C was extended beyond 60 min and these conditions were adopted as the standard procedure for further pellet preparation. The chemical and physical properties of this hydrogen reduced alunite are summarized in Table I. Sulfation Experimental Results The kinetics of SO2 absorption by hydrogen-reduced alunite were investigated from 100-500°C using sulfur dioxide partial pressures from 0.001-0.025 a t m and oxygen partial pressures from 0.001-1 a t m . Both dry gases and gas mixtures containing up to 0.084 atm of water vapor were used. The dependence of the rate of sorption with various pellet sizes was also investigsted. Temperature. The rate of absorption increases with temperature. The absorption of SO2 by reduced alunite a t various temperatures is shown in Figure 1. Below 300°C the rate becomes negligibly small when the weight increase corresponding to the sulfation of all K2O in the pellets has been reached. This limit is indicated by the horizontal dashed line in Figure 1. X-ray diffraction shows K2S04 to be the only phase present. Hence, below 300°C sulfation of K2O accounts for essentially all of the absorption. Above 300°C a new phase, K3A1(S04)3, appears. There is no X-ray diffraction evidence for aluminum sulfate or aluminum sulfite. This evidence suggests that above 300"C, alumina also begins to sulfate, but the aluminum 924
Environmental Science & Technology
Table I. Properties of Alunite Reduced in Hydrogen (8OO0C, 60 rnin)
Chemical analysis Potassium oxide ( K 2 0 ) Aluminum oxide ( A l 2 0 3 ) Silica (SiO2) Sulfite (S03*-) Physical properties Surface area, A0 Porosity Crushing strength
5
1
i
g
62.0 8.4 4.8 24 f 1 m 2 / g - ' 64% 14.1 kg cm-2
B. . ..... . .... ...............
5
21.0 wt %
300c
................-. .......-. ...
200c
8 0 1
IOOC
,
TIME (minl
Figure 1. Sulfation of reduced alunite at various temperatures with 0.003 atm SO2 and 0.024 atm water vapor
sulfate combines immediately with potassium sulfate to form the mixed compound. As the sorption of SO2 nears completion another mixed compound, KA1(SO4j2, appears as the final product in the X-ray diffraction spectra. This compound is identical with dehydrated alunite. The rate of sulfation becomes negligibly small when the weight increase corresponding to the stoichiometric formation of KAl(S04)z in the pellets has been reached. The excess alumina present at this stage, as in alunite, is not sulfated. Sulfur Dioxide Concentration. The amounts of SO2 absorbed at 200" and 5OO"C, for various partial pressures of S02, are shown in Figures 2 and 3, respectively. Oxygen and 0.024 atm of water vapor were also present. The initial absorption rates, determined from the slopes of the weight gain curves a t zero time, appear to increase linearly with the SO2 concentration. The rate decreases with increased sulfation of the solid and becomes zero asymptotically as the absorption capacity of reduced alunite is reached. Oxygen Concentration. The role of oxygen partial pressure was investigated in the absence of water vapor. At 200°C the oxygen concentration has little effect on the rate of absorption. A t 5OO"C, the initial rate is independent of p ( O z ) , but in the later stages of absorption, the rate becomes a function of the oxygen concentration. See Figure 1. Since the early sulfation primarily involves K20,
I-
?
160
I
I
1
I
1
IO0
200
300
400
500
_1
I20
BO
40
200 c
0
600
700
TIME ( m n )
Figure 2. Sulfation at 200°C with 0.024 atm water vapor and various partial pressures of SO2
0
100
200
300
TIME 0 0 2 5 a t m SO2
-
100
200
300
400
500
600
700
TIME ( m i n )
Figure 3. Sulfation at 500°C with 0.024 atm water vapor and various partial pressures of SO2
-
I
I
I
I
1
I
I
A 0 997 o t m
"00w
E
I
c
2
160-
PL
b
5
120-
W
9
2
80-
t-
o
'
40-
0
: i / , %
0
IO0
j 200
300
400
500
600
700
TIME ( m i n )
Figure
Effect of oxygen pressure on sulfation with 0.003 atm SO2 ( d r y ) at 500°C 4.
500
600
700
imin)
Figure 5. Effect of water vapor on sulfation with 0.003 atm SO2
-
0
400
these results indicate t h a t the formation of KzS04 is independent of oxygen partial pressure, while the rate of formation of potassium aluminum sulfate (or aluminum sulfate) depends on the oxygen partial pressure. The sulfation of potassium oxide may proceed through an intermediate sulfite formation step which is rate controlling and independent of oxygen. Aluminum sulfite is thermodynamically unstable a t the experimental conditions used. Water Vapor. The effect of water vapor from 0 to 0.024 atm p(H2O) on the reaction kinetics was investigated. The rates were independent of water vapor a t 5OO0C, as shown in Figure 5 . A t low temperatures, 200°C, the initial rates were also independent of water vapor, but as the reaction proceeds, the rate decreases markedly for reaction in dry gases. Above p(H2O) = 0.012 atm, water vapor has little effect on the reaction kinetics a t any temperature. I t should be noted t h a t water vapor is not a reactant in the sulfation chemical process but acts as a promoter. Nitrous oxide gas is also a promoter. These gases also catalyze the sulfation of alkalized alumina (Town et al., 1969). Surface area measurements on the pellets, reacted with dry and wet gases, showed that there was a reduction in the surface area after reaction with dry SO2 and 0 2 a t low temperatures. Reaction with wet gases caused a smaller surface area reduction. Blocking of pores by the sulfation products probably causes the reduction in surface area and this may account for the marked decrease in the extensive sulfation rate as reaction proceeds in dry SO2 a t low temperatures. Water vapor may affect crystallization of the reaction products but there was no apparent difference in the scanning electron micrographs of reduced alunite after sulfation in wet and dry gases. Pellet Size. The effect of particle size was also investigated using crushed pellets. The standard pellets were 12.7 mm in diameter and about 1.9 mm in thickness. After reduction they were crushed and 0.5, 0.25, and 0.125 mm average diameter grains were sulfated a t 500"C, where the fastest rates were obtained. The particle size had little effect on the sulfation rate. Since the rates remain constant with decreasing particle size, it can be concluded t h a t large-scale pore diffusion within the pellet is not a n important factor in the reaction kinetics. This was also confirmed by microprobe sulfur analyses of cross-secVolume 7 , Number 10, October 1973 925
-2 0
I
I
I
I
1
-2.0
,
I
I
I
085
0 85
0 8
07
0, LY
06
05 04
02
0
TIME
TIME ( m i n )
Figure 6. Logarithmic rate plots for sulfation of K 2 0 at various temperatures with 0.003 atm SO2 and 0.024 atm water vapor
(mini
Figure 7. Logarithmic rate plots for sulfation of K 2 0 at various SO2 pressures with 0.024 atm water vapor at 200°C
tioned pellets. Although there were random point-to-point fluctuations, the average sulfur content traversing the specimen was uniform.
Discussion o f Rate Mechanism Belou 30PCSulfation of K20 I t was noted earlier t h a t the sulfation of KzO is the main reaction below 300”C, while there is sulfation of both KzO and A1203 above 300°C. It is convenient to divide the discussion of rate mechanisms accordingly. A term F ( K 2 0 ) is defined as the fraction of K 2 0 that has already reacted to form K z S 0 4 . FK,O
=
ZIWK
(1)
where u: is the weight change a t any time and WK is the maximum weight gain corresponding to the stoichiometric sulfation of KzO in the pellets. When log, [I - F(K20)] was plotted against reaction time, we obtained a straight line for all wet sulfation kinetic experiments a t 300°C and lower temperatures. Three typical experiments in 0.003 a t m p ( S O 2 ) a t different temperatures are shown in Figure 6. Differentiating the integral expression, lOge(1
-
FK,O)
= --k’K,Ot.
(2)
yields the following rate expression with the rate proportional to the unreacted volume fraction dFK‘O
~dt
-
h’K,O(l
-
FK20)
(3)
where k’(K20) is determined from the slope of the experimental plot of log, [l - F ( K 2 0 ) ]vs. time. The rate increases with the partial pressure of SO2 as shown for five runs a t 200°C in Figure 7. When the experimental rate constant h’(KZ0) is plotted against p ( S 0 ~ a) straight line was obtained, Figure 8, indicating a firstorder dependence on p ( S O z ) :
926
Environmental Science & Technology
a 4 . Y 1 “Ow
1
1
Figure 8. Logarithmic rate constant for sulfation of K 2 0 dependency on SO2 pressure
Other experimental studies of gaseous reactions with porous solids at slow reaction rates, where pore diffusion is not rate controlling, have followed a logarithmic law (Turkdogan and Vinters, 1971; Van Hecke and Bartlett, 1973; Krishnan and Bartlett, 1973) which is superficially analogous with a first-order homogeneous batch reaction. A slowly reacting porous solid can be reasonably described as a pellet consisting of a collection of micrograins often of variable size, a t least partially surrounded by interconnecting pores. The rate-controlling process may be a chemical reaction a t the surface of the micrograins or solid state diffusion through the solid reaction product layer surrounding the micrograin a t a rate much slower than the pore diffusion limiting rate. Pore diffusion through the micrograin reaction product shell can generally be eliminated since the diffusion distances involved are usually very much less than pore diffusion distances through the bulk of the pellet. For each micrograin there is a shrinking core of unreacted material which may be surrounded by the solid reaction product. Reaction proceeds until the micrograin
TEMPERATURE
270
- 3 ,,3:0
-3 4
270
lo
4 01 17
I70
l:5
A E = 8 0 0 CAL
-
I
I 21
19
IT
-I
MOLE‘^
I
1 2 5
I
23
x 103
,
OC
O :l
2 9
27
(OK-')
Temperature dependence of the experimental rate constant for sulfation of K 2 0
Figure 9.
is consumed and smaller micrograins are consumed before larger ones. Consequently, the extensive rate of reaction of the porous solid decreases with time. In this connection, Bartlett et al. (1973) recently showed t h a t rates calculated for a log-normal size distribution of micrograins, within a wide range of size distribution standard deviations, approximates the logarithmic law within the usual experimental error when a surface chemical reaction on the micrograin was rate controlling. The experimental rate constant is a function of the chemical reaction rate constant and micrograin size distribution. In the present experimental study, values of the experimental rate constant, k(K2O), were determined for different temperatures and plotted against 1/T in Figure 9. The apparent activation energy from this Arrhenius plot is only 800 cal mol-', and
0.0725
kK20 =
-
(min-' atm-')
(5)
The following heterogenous chemical reaction steps are suggested for the sulfation of K2O. The sulfite formation step is assumed to be the rate-controlling step, because of the zero-order dependency on oxygen. Formation of K2S04 requires oxygen. -06
Kz01,
4
+
up
K20i,*SOz
-04 m
-
KzOI,.SOz equilibrium c h e m i s o r p t i o n
-
100
0
200
300
400
600
500
700
TIME ( m l n )
Logarithmic rate plots for sulfation of A 1 2 0 3 at various temperatures with 0.003 atm SO2 and 0.024 atm water vapor Figure 10.
r a t e - c o n t r o l l i n g slow s t e p
0 0 2 4 o t m HO ,
-
c1 0
-I 0
U
-0 6
-0 4
0 2
0
100
200
300
400
500
600
700
TIME (min)
Figure 11. Logarithmic rate plots for sulfation of A t 2 0 3 at various SO2 pressures with 0.024 atm water vapor at 500°C
+
%Ok)
+
KzSO,
(7)
(8)
Discussion of Rate Mechanism Above 300'CSulfation of K2O and A1203 Above 300°C, potassium sulfate is produced by the same process as below 300°C, but potassium aluminum sulfates are also produced. From our investigations it was not possible to determine whether potassium aluminum sulfate forms in a single-step reaction [K2S04 + Y3A1203 + SO2 + l/202 2/3K3A1(S04)3]or proceeds through the formation of an amorphous aluminum sulfate intermediate [%A103 + SO2 + Y3AMSO)413; [Y3A12(S04)3 + K2SO4 -* 2/3K3A1(SO4)3].Chao and Sun (1967) report that the sulfation of aluminum oxide is extremely slow even a t 600°C unless alkali metal sulfates are present, which catalyze the sulfation of alumina above 300°C. Extrapolation of the rate expression for the formation of K2SO4 to 500°C yields the experimental rate constant h'(K2O) a t 500°C. The weight increase corresponding to the formation of K2S04 can be computed from the rate expression. Subtracting this weight increase from the observed total weight increase yields the weight increase contributed by alumina sulfation. The fraction of alumina, F(A1203), that has been sulfated was calculated on the basis of conversion to KAl(S04)z a t F(A1203) = 1. When log, [l - F(A1203)I was plotted against reaction time, straight lines were obtained for the runs above 300°C, suggesting a rate mechanism for sulfation of A1203 similar to t h a t of K20, with the rate proportional to the unreacted fraction of the solid. The effect of different temperatures and SO2 partial pressure on the rate of alumina sulfation are shown in Figures 10 and 11. The dependence of the rate m p(S02) is first order as shown in Figure 12. The rate of sulfation of A1203 is proportional to p ( 0 2 ) 1 ' 2 . The Arrhenius plot for the sulfation of alumina in hydrogen-reduced alunite is shown in Figure 13. The apparent activation energy is 9600 cal mol-'. The alumina sulfation extensive rate may be expressed as
-
: :
(6)
K,SO,
KSO,
-02
$
e
SO,(g)
-
Volume 7 , Number IO, October 1973 927
The postulated sequence of reactions is:
'/3A12031
+ SO2 1-1/3A12031
-
where 1/2
(10)
k'A1203 = h.41203PS02P02
'/3A1203 Is.SO,
and k41,0, =
1.16 e-9600
'RT(
(11)
min-' atm+*)
T h e rate-controlling chemical reaction for sulfation of alumina in reduced alunite is expected to be relatively simple and irreversible, and probably involves the sulfation of aluminum oxide by reaction with SO2 and 0 2 .
I
1
I
I
I
I (
7
5
5
E 4
0 X
10
0 -,
(12)
i/ZO? '/3Al~(SO4)~ r a t e - c o n t r o l l i n g slow s t e p
(13)
+ '/jAl,(SO,)J
-
--t
+ 2/3Al(S04)3
2/3K3A1(S04)3
'/,K3Al(SO,),
(14)
ZKAl(SO,)? (15)
Both postulated sets of reactions for sulfation of K2O and A1203 involve chemisorption prior to a rate-controlling step and t h e apparent activation energy will be decreased by the magnitude of the chemisorption free energy, which is a negative quantity. This may explain the low values of the apparent activation energies observed in this study and similar sulfation reactions in porous solids where the rate is much too slow to be controlled by gaseous diffusion in pores.
3
-Y a
W A I h - ~ X P ( - ~ ~ ~ ~ ~ , P S ~ (16) ~P~,"*~ 2
I
Conclusions
-
&300c I
0
0 4
I
I
I
I
I
08
12
16
20
24
PCT
SO,
Logarithmic rate constant for sulfation of pendency on SO2 pressure Figure 12.
-5
e
I
I
I
AI203 de-
I
-6 0
The rate of sorption a t 0.003 a t m SO2 and 300°C is 5.6 x 10-4 g/g of solid per min, and the rate is only slightly sensitive to temperature. The maximum loading is 0.05 g/g of solid below 300°C, and above 300°C it is increased to 0.19 g/g of solid. These values are comparable to those for alkalized alumina (Bienstock e t al., 1967). The kinetic investigations showed that the rate is proportional to the unreacted fraction of the solid sorbent, proportional to the partial pressure of sulfur dioxide, and only slightly sensitive to temperature. Initial rates (K2O) are independent of oxygen partial pressures and only slightly sensitive to temperature. At low temperatures, absence of water vapor causes more extensive pore blocking and decreases the rate. At high temperatures and larger sulfation periods, the sulfation of A1203 becomes important and the rate becomes a function of oxygen partial pressure.
-6 5
Nomenclature -7 0 m
U 1
-75 0 -
-80
-a 5
9 0 VI0
I 012
I 014
I/T x IO2
I
I
016
018 (OK-')
Figure 13. Temperature dependence of the experimental rate constant for sulfation of A I 2 0 3 928
'
Extrapolation of k'(A1203) to lower temperatures using Figure 13 shows t h a t for reaction periods used in our study, the weight gain due to alumina sulfation becomes insignificant below 300°C. We have already seen t h a t a t high temperatures, the rate can be written as the combination of rate expressions for K 2 0 and A1203 sulfation, the resulting weight increase is W = WK[l - exp(-kti,op~~,t)lf
I
10
SO,
Com bined Rate Equation
6
--
K2SO.j
+
*
equilibrium c h e m i s o r p t i o n
Environmental Science & Technology
F(A1203) = fraction of alumina t h a t has reacted to form KAl(S0412 F(K20) = fraction of potassium oxide t h a t has reacted to form K2S04 k ( A l ~ O 3 )= experimental rate constant for alumina sulfation, min-1 atm-3l2 k ( K 2 0 ) = experimental rate constant for sulfation of potassium oxide, m i n - l a t m - l h"(A1203) = logarithmic rate constant for alumina sulfation, min-' k'(K20) = logarithmic rate constant for potassium oxide sulfation, min-1 p ( S 0 2 ) = partial pressure of S 0 2 , a t m p ( 0 2 ) = partial pressure of 0 2 , a t m R = gas constant, 1.98cal " K - l mol-' t = reaction time, min
T
= reaction temperature, "K
w = observed weight gain, g/g of solid-1 WAI = maximum weight gain corresponding to t h e stoi-
chiometric sulfation of alumina, g/g of solid-1 W , = maximum weight gain corresponding t o t h e stoichiometric sulfation of potassium oxide, g/g of solid-1
Literature Cited Bartlett, R. W., Krishnan, N. G., Van Hecke, M. C., Chem. Eng. Sci., 28, in press (1973). Bienstock, D., Field, J. H., J . Air Pollut. Contr. A s s . , 10, 121-5 ( 1960). Bienstock, D., Field, J . H., Myers, J . G., Rept. of Inves. 7021, U.S. Bur. Mines (1967).
Chao, T., Sun, S. C., Trans. S O C . Mining Eng., 238, 420-29 (1967). Eng. Mining J . , 173, N o . 9, "U.S. May Get Aluminum from Alunite," June 1972. Fink, W. L., Van Horn, K. R., Pazour, H. A , , Ind. Eng. Chem., 23, 1248 (1931). Krishnan, N. G., Bartlett, R. W., Atmos. Enairon., 7, 575-86 (1973). Town, J. W., Sanker, P. E., Kelly, H. J., Rept. of Inues. 7275, US.Bur. Mines, 1969. Turkdogan, E. T . , Vinters, J. V., M e t . Trans., 2,3175-88 (1971). Van Hecke, M. C., Bartlett, R. W., ibid., 4,941-7 (1973). Receiued f o r reuieu: September 27. 1972. .4ccepted June 18, 1973. Work was supported by Enuironmental Protection .4gency under Grant A P 00876.
Occurrence of Hexachlorophene and Pentachlorophenol in Sewage and Water Donald R. Buhler,' M. E. Rasmusson, and H. S. Nakaue Department of Agricultural Chemistry and Environmental Health Sciences Center, Oregon State University, Corvallis, Ore. 97331
Concentrations of hexachlorophene (HCP) and pentachlorophenol ( P C P ) in sewage and water samples have been analyzed by gas chromatography. H C P and P C P levels in 24-hr composite samples of sewage influent collected simultaneously from three Oregon cities ranged between 20-31 ppb and 1-5 ppb, respectively. Composite effluent values from these same sewage treatment plants were 6-12 ppb H C P and 1-4 ppb PCP, reflecting a 60-70R removal of H C P and a 4-28% removal of PCP. Analyses of daily and hourly water samples from t h e Willamette River collected just upstream from the city of Corvallis, Ore., showed t h a t H C P and P C P were present in river water in concentrations varying between 0.01-0.1 ppb and 0.100.70 ppb. respectively. Conventional processing of raw Willamette River water a t t h e Corvallis Taylor water treatment plant removed about 60% of the H C P and P C P originally present in the water leaving about 40% of these chlorophenols in t h e finished drinking water. Identifications of H C P and P C P in sewage effluent a n d influent, Willamette River water and treated drinking water were confirmed by mass spectrometry.
ics, and germicidal preparations, whereas P C P is used extensively in industry as a preservative and for slime control. P C P has been found in streams a n d accumulated in fish tissues following industrial discharge (Stark, 1969; Rudling, 1970), a n d H C P is thought to persist similarly in surface waters for long periods (Bandt and Nehring, 1962). Because H C P and P C P are highly toxic chemicals ( N a kaue et al., 1973; Deichmann et al., 1942), we have made use of a sensitive gas chromatographic method to determine the levels of H C P and P C P in municipal sewage from three Oregon cities (Corvallis, Eugene, and Salem); in the Willamette River t h a t receives the sewage discharge from the three plants; and in a public water supply t h a t employs treated Willamette River water. The presence of H C P and P C P in sewage and in raw and treated river water has also been confirmed by mass spectrographic analysis.
Materials and M e t h o d s Chemicals. Hexachlorophene (U.S.P. grade) was a gift of the Givaudan Corp.. Clifton, S.J., and was twice re-
crystallized from isopropranol-water to yield a purified product with m p a t 165-6°C. Reagent grade pentachlorophenol was purchased from Eastman Organic Chemicals, Rochester. N.Y. All solvents or chemicals were reagent Increasing amounts of domestic and industrial waste are grade obtained from commercial sources and were further being discharged into surface waters. Some of these purified as necessary. Ethereal diazomethane (0.5 mmol/ wastes contain toxic chemicals such as chlorophenols, ml) was prepared from Diazald (Aldrich Chemical Co., which resist degradation and tend to persist in t h e enviMilwaukee. Wis.) according to the directions provided by ronment where they may create serious public health and the manufacturer and stored at -4°C until used. Methywater quality problems. lene-14C-hexachlorophene with a specific activity of 2.77 Hexachlorophene [HCP, 2.2'-methylene-bis(3,4,6-tri- mCi/mmol was obtained from New England Nuclear chlorophenol)] and pentachlorophenol ( P C P ) are two such Corp.. Boston, Mass. chlorophenols, widely used in our society in a variety of Apparatus. Gas chromatography (gc) was carried out in fungicidal or germicidal applications. H C P is employed a Varian Aerograph Hy-Fi 600D gas chromatograph primarily in consumer products including soaps, cosmetequipped with a tritium foil electron capture detector. Mass spectrometry was performed with a Varian Atlas CH-7 mass spectrometer. A majority of the samples were introduced into the instrument via the inlet probe. In 'To whom correspondence should be addressed. Volume 7, Number 10, October 1973
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