tive solubility of benzene and cyclohexane in the 3-methylsulfolene. I n this regard it is interesting to note that the optimum temperature for aromatics extraction with sulfolane is about 120' (Ileal, et al., 1959) and so the relative solubility probably increases with increasing temperature up to a point. Hence the observed increase in separation factor with temperature. The long-term performance of the 3-methylsulfolene membrane has not been determined and probably membrane deterioration resulting in a decrease in permeability and an increase in selectivity due to plasticizer removal could be expected. The plasticizer removal rate does not appear to be extremely rapid, however, because, for example, the same membrane was used to test all of the aromatic-naphthenic systems (Table I) and during this time (about 100 hr) the benzene-cyclohexane flux dropped only about lo%, prob-
ably within experimental error of the measurements. From the standpoint of plasticizer removal possibly another aromatic extraction agent or a heavier sulfone could be used to advantage and this will be the subject of future research. literature Cited
Binninz. R. C., Lee, R. J., Jennings, J. F., Martin, E. C., Znd. Eng.-Chem., 53,46 (1961). Deal, G. H., LCvs.rs, -._ H. D., Oliver, E. D., Papadopoulos, M. N., . . Petrol. ReJinpr Jj - , -3,185 (1969). Fels, M., AZCi1E Svmp. Ser., No. 180,68,49 (1972). Huang.. R. Y . M.. Lin. V. J. C.. J . Avvl. .. Polvm. Sci.., 12., 2616 (1968).
I
,
Kucharski, M., Stelmaszek, J., Znt. Chem. Eng., 7, 618 (1967). Li, N. N., Znd. Eng. Chem., Prod. Res. Develop., 8,281 (1969). RECEIVED for review October 10, 1972 ACCEPTED February 20, 1973
Simultaneous Catalytic Reduction of Nitric Oxiide and Sulfur Dioxide by Carbon Monoxide Charles W. Quinlan," Vurnel C. Okay, and J. R. Kittrell Department of Chemical Engineering, University of Massachusetts, Amherst, Massachusetts 01002
Experimental work is reported on the simultaneous reduction of NO and SO2 by CO over a commercial copper on alumina catalyst. It is shown that the catalyst is less active in promoting the combined reduction than in either case taken separately. A predictive equation for SO2 conversion as a function of temperature and CO level is presented for NO levels from 0 to 1000 ppm. Catalyst selectivity does not appear altered by addition of NO and maximum obtainable removal of sulfur compounds appears to be limited to75-80% in a one-bed process.
Emissions of nitric oxide arid sulfur dioxide from stationary sources have been an increasing problem in the United States over the last several years. It has been estimated that 36.6 X IO6tons of sulfur dioxide were emitted from stationary sources in 1971 (Chilton, 1971) and that sulfur dioxide emissions could quadruple by the year 2000 if adequate controls are not provided. Approximately one-half of this sulfur dioxide emission level is due to power plants burning coal and oil. Nitrogen oxide emissions in 1968 were estimated to be 16 X lo6 tons, about 60% of which originated from stationary sources (Bartok, et al., 1971). Several processes have been suggested for reducing sulfur dioxide and nitrogen oxide emissions from stacks of combustion operations. Such suggestions have included reduction of sulfur or nitrogen (content of the fuel (Chilton, 1971), modifications of the combustion process (Bartok, et al., 1971), and various wet and dry processes for removing sulfur dioxide and nitric oxide from the stack gases. These processes have a variety of widely publicized disadvantages, including cost, by-product sales, and associated water or solid waste problems. Kinetic processes for removing nitric oxide and sulfur dioxide have also been considered, for example, by utilizing a reducing agent to reduce sulfur dioxide to elemental sulfur and nitric oxide to elementary nitrogen.
Carbon monoxide has been one of the most widely studied reducing agents for sulfur dioxide and nitrogen oxide. Sulfur dioxide reduction has been studied by Khalafalla, et al. (1971), over an iron catalyst and by Ryason and Harkins (1967), Querido and Short (1973), and Quinlan, et al. (1973), over copper catalysts. Nitric oxide reduction by carbon monoxide has been extensively studied with reference to automobile emissions, by Shelef and Kummer (1969), Baker and Doerr (1965), and many others. Ryason and Harkins (1967) have also studied the activity of several catalysts for the simultaneous reduction of sulfur dioxide and nitric oxide by carbon monoxide. The purpose of the present study was to evaluate the effect of sulfur dioxide on the activity of a commercial copper on alumina catalyst for nitric oxide reduction, and the effects of nitric oxide on the activity and selectivity of the catalyst for sulfur dioxide reduction. This information, coupled with other catalyst studies in the absence of nitric oxide (Quinlan, et al., 1973; Okay and Short, 1973), can be of value in process design and optimization efforts needed for an evaluation of carbon monoxide reduction as a process scheme for removing sulfur dioxide. The primary reactions that occur in the simultaneous reduction of nitric oxide and sulfur dioxide by carbon monoxide are Ind. Eng. Chem. Process Des. Develop., Vol. 12, No. 3, 1973
359
LEGEND GAS CYLINDER
=
I2 I
N2 NO/ N2
=
COIN2
TI
TO bNALYTIC4L TR4IN
73
Copper level 8 wt % Surface area 140 m2/g Pore volume 0.33 cc/g Data provided by the Harshaw Chemical Co.
7 4 rn C%/Np 1 5 I SD2/N2
FURN4CE-l
Table 1. Catalyst Propertiesaof Harshaw Cu-0803 (T in.)
U
ROTAMETER 8-VALVE LKCTROSTATlC AECIPIT4TOI
TO AN4LYTIC4A TR4IN
I-5
Figure 1 . Schematic flow diagram of experimental equipment
CO
+ KO
2CO
+ so2
+ +
co + '/2SZ 2COS
+ SO1
+ '/2N2 '/2SZ + 2C02
(2)
cos
(3)
+
(4)
COz
+
+
3 / ~ S ~2C02
Equations 1 and 2 are, respectively, the direct reduction of nitric oxide by carbon monoxide and the direct reduction of sulfur dioxide by carbon monoxide. Equation 3 represents the reaction between carbon monoxide and sulfur to produce carbonyl sulfide, a toxic and undesirable side product. In addition carbonyl sulfide can react directly with sulfur dioxide to produce elemental sulfur by eq 4 (Querido and Short, 1973). Numerous other reactions could take place, particularly if water is present as in a flue gas, such as the water-gas shift reaction
HzO
+ CO
-
Hz
+ COa
(5) With the hydrogen which is generated by the water-gas shift reaction several other components can be formed, such as H2S and ammonia. Although the water-gas shift reaction has been reported to occur a t the higher carbon monoxide concentrations present in automobile exhaust gases (Klimisch and Barnes, 1972), the reaction apparently does not occur a t the lower carbon monoxide levels of furnace stack gases with the copper on alumina catalyst of the present study (Okay and Short, 1973). Reactions associated with the shift reaction will therefore not be considered further here. Experimental Section
All experiments were conducted on a plug flow, tubular reactor; the catalyst employed was a 10% copper oxide (8% copper) on y-phase alumina, commercially available as Harshaw CU-0803. The specifications for this catalyst are contained in Table I. For each run, approximately 3 g of this catalyst was placed in the reactor and the temperature of the reactor was raised to 825°F in flowing nitrogen. The catalyst was activated by passing the reaction gases (carbon monoxide, sulfur dioxide, etc.) over the catalyst a t 825'F for a time period of 6-12 hr, during which the sulfur dioxide conversion gradually rose to 100%. The catalyst activity was then checked after each start-up by dropping the catalyst temperature to a predetermined level where the conversion was about 700/0. All data were taken after steady-state reactor operation was achieved. 360
Ind. Eng. Chem. Process Der. Develop,, Vol. 12, No. 3, 1973
The reactor system used in this research is shown schematically in Figure 1. Cylinder gases employed were chemically pure nitrogen, 8% carbon monoxide in nitrogen, 3000 ppm nitric oxide in nitrogen, 1% sulfur dioxide in riitrogen, and 50% carbon dioxide in nitrogen. All gases were obtained from the Matheson Co. Blends of these gases allowed evaluation of reactor performance a t levels of 2000 ppm of SO2, 0-1000 ppm of NO, 14% COZ,and about 6000 ppm of CO. To achieve these blends, the cylinder gases were delivered through individual rotameters to a blending rotameter. The bulk of the mixed gases was directed through the main reactor rotameter to pass over the catalyst bed, while a portion was removed through a sample rotameter and directed to the analytical train. The reactor consisted of a '/B-in. IPS titanium pipe, and was 36 in. in length; the middle 28 in. of the reactor was enclosed in the controlled heating zones of a three-zone Lindberg heavy-duty furnace. Several reactor metals were evaluated for utility in this system. Stainless steel, nickel, and monel were all found to catalyze the nitric oxide-carbon monoxide reaction (Quinlan, 1972) as observed by others (Meguerian and Lang, 1971). Aluminum was satisfactory but tended to soften near the temperature range desired for the present studies. Quartz glass and titanium could withstand the temperature and did not catalyze the nitric oxide reduction reaction. Because of the more robust nature of the titanium pipe, all experiments were conducted with this metal. A catalyst charge was typically 3 g of Harshaw catalyst, ground to 20-30 mesh, and yielding a 4 in. bed height. Stainless steel screen of 40-50 mesh supported the catalyst. Apparently due to the low residence time of the gases over the screen, it did not contribute to the conversion of nitric oxide a t the flow rate used in the present study. The usual experimental tests showed little or no effect of internal or external heat and mass transfer on the rate of reaction. A Lindberg heavy-duty controller, attached to the Lindberg furnace, was used to control the reaction temperature. Temperature measurements were made with chromel-alumel thermocouples attached to the outside wall of the reactor, spaced about 1 in. apart over the bed length. Output was monitored with a Leeds & Northrup potentiometer. Bed conditions remained isothermal to within h2'F. The reactor system was operated essentially at atmospheric pressure with the total pressure drop from inlet t o exhaust never exceeding 60 mm. This magnitude of pressure drop was as expected from calculations made using the Ergun equation (Bird, et al., 1960). After leaving the catalyst bed, the reactor stream was passed through an electrostatic precipitator for removal of elemental sulfur. A 32°F cold trap was also used for sulfur removal after the electrostatic precipitator, as shown in Figure 1. Specific tests were conducted to confirm that no components were lost by reaction or deposition in the downstream lines (Quinlan, et al., 1973). The analytical train utilized a Perkin-Elmer Model 900 gas chromatograph for analysis of carbon monoxide, carbon
t
I
h
1000
b
r; 5
8
100
ul U
3
r l FEED CONTAINS: 2000 ppm SO 140,000 ppm CO 34,000 ppm 0 WITHOUT NO WITH 6OOppm NO
--
Figure 2. Effect of temperature and CO ratio on equilibrium SO2 level
Figure 3. Effect of temperature and CO ratio on equilibrium COS level
dioxide, sulfur dioxide, and carbonyl sulfide. Helium was used as the carrier gas for the chromatograph and a hot wire detector employed for all analyses, with 150" detector temperature and a 300mA detector current. Carbon dioxide, carbonyl sulfide, and rsulfur dioxide were analyzed using a 6 ft X l i s in. stainless steel column, packed with molecular sieve 5A (80-100 mesh). The columns were maintained a t a temperature of approximately 70". Nitric oxide was analyzed separately on a Dynasciences Pollution Monitor NX 130 equipped with a scrulbber to remove sulfur dioxide. This instrument utilizes an electrochemical potential cell to provide a continuous monitoring of the nitric oxide content of the gas stream. Output is read directly from a calibrated meter. Other components of the process do not interfere with the nitric oxide analysis, as claimed by the Dyriasciences Corp. and demonstrated experimentally. ilnalysis was not attempted for N 2 0 based on reports of earlier investigators that formation of that compound was not important a t the temperatures of operation (Shelef, et al., 1968). Furthermore, material balances on CO indicated that the reaction product of KO reduction was N2. Further details on experimental apparatus and procedure are contained elsewhere (Quinlan, 1972).
well as CO, C02, SOa, SO2, CSZ,SZ,Ss,COS, 0 2 , and N2. The JANAF tables were used as the sources of free energy data. All calculations are represented as a function of stoichiometric carbon monoxide ratio (CO ratio). This ratio represents the stoichiometric excess of carbon monoxide for the primary reactions of the process and is defined as
Equilibrium Calculatiorir
Calculations were made to determine the equilibrium composition of the several reactants in this system a t 1 atm pressure and with the templsrature range expected to be important for the simultaneous reduction of sulfur dioxide and nitric oxide. These calculations were performed by a free energy minimization program (White, et al., 1958; Querido, 1971). This program does not require a predetermined series of reactions to be specified, but rather requires only a tabulation of all possible reactants and products. Results have already been reported for feed gases comprising SO2 and CO (Querido and Short, 1973) and SO:rC02-H20 (Okay and Short, 1973). Hence, we report here only the results of the introduction of nitric oxide as a feed component. Possible reaction products which were considered were therefore all nitrogen oxides as
CO ratio =
concn of CO in feed 2(concn of S02) concn of K O in feed
+
(6)
The solid lines of Figure 2 present the equilibrium SO2 concentrations for an initial mixture containing 2000 ppm of SO2 as a function of reactor temperature and CO ratio. It is apparent that low temperatures and high CO ratios favor the conversion of S02. The effect of the addition of 600 ppm of NO to the initial mixture is shown by the dashed lines of Figure 2. Because the nitric oxide reduction reaction of eq 1 consumes CO competitively with the sulfur dioxide reduction reaction of eq 2, it is not surprising that higher levels of sulfur dioxide remain a t equilibrium for a given CO concentration in the initial mixture when nitric oxide is also in the initial mixture than when it is not present. However, when the CO ratio is defined as in eq 6, the deviation of the solid and dashed lines of Figure 2 is not great. Hence, the lines of Figure 2 can be used to describe equilibrium SO2 levels for the entire range of NO levels commonly encountered in stationary source emissions. The conversion of nitric oxide to elemental nitrogen a t equilibrium is 100% for all of the temperatures shown in Figure 2. All other components of the reaction products except those shown in eq 1-4 have negligible equilibrium concentrations. I n Figure 3 is shown the analogous effects of temperature and CO ratio on eqiilibrium concentrations of carbonyl sulfide for an initial mixture containing 2000 ppm of S02. The solid lines again represent equilibrium concentrations when the initial mixture contains no nitric oxide and the dashed lines represent the results for an initial mixture also containing 600 ppm of nitric oxide. For CO ratios exceeding 1.0 (Le., for CO levels exceeding the stoichiometric requireInd. Eng. Chem. Process Des. Develop., Vol. 12, No. 3, 1973
361
FEED 2920 ppm NO 8625 ppm CO BALANCE N2
4c
Table II. Parameter Estimated for Eq
3c
Parameter
Estimate
20
A B (at 732°F) B (at 795°F)
19.8 1.53 1.87 2.16 x 104 O R
'r -r
EIR a
10
I
120
I
125
I
130
I
I
135
140
RECIPROCAL TEMPERATURE I 10','R"
Figure 4. Temperature dependence of first-order NO rate constant for NO reduction
L
p 90 2
- 770-771
0
70
\ \ \
x)L
IAO
260
A
- 754'753
360 460 5bO 660 700 NO CONCENTRATION IN FEED, PPM
sb0
io0
ld00'
Figure 5. Effect of nitric oxide level on SO2 conversion ments for the conversion of SO2 and NO), higher temperatures decrease the equilibrium level of COS formed. Higher CO ratios increase the equilibrium level of COS produced, due to the influence of the formation reaction of eq 3. It is apparent from Figure 3 that, if the CO ratio is defined as in eq 6, equilibrium levels of COS production may be predicted by Figure 3 for all KO levels within the range commonly observed in stationary source emissions. It should further be noted that, since COS is produced in this reaction system, the COS equilibrium curves of Figure 3 will be approached from the bottom in a reaction system. Hence, kinetic restrictions will likely lead to significantly lower COS levels in the reactor effluent than would be predicted by Figure 3. Catalyst Activity Studies
Nitric Oxide Reduction. Several runs were conducted to determine the activity of the Harshaw catalyst for the reduction of nitric oxide by carbon monoxide in the absence of sulfur dioxide. Using the 20-30 mesh particles, complete conversion of nitric oxide was obtained a t temperatures of 500'F and above, and a space velocity range from 35,000 to 50,000 hr-1 (defined as the volumetric flow rate of the reactant gas a t the catalyst inlet conditions of temperature and pressure divided by the volume of the catalyst bed). -1series of four data points were obtained in the temperature range from 280 to 370°F and a t a space 362
Ind. Eng. Chem. Process Des. Develop., Vol. 12, No. 3, 1973
k8o2 in sec-1.
velocity of approximately 40,000 hr-1. The first-order rate constant for the disappearance of carbon monoxide is plotted in Figure 4. For the range of variables covered, the linearity of Figure 4 suggests that a first-order model may adequately describe the system behavior. The activation energy estimated from the slope is approximately 10 kcal/mol. This low value indicates that, despite the small particle size, internal diffusion may be important. However, the inherent catalyst activity for this reaction is excellent. The previous work on the reaction of sulfur dioxide with carbon monoxide over this catalyst (Quinlan, et al., 1973) indicated that temperatures of operation much higher than 500°F would be required for the simultaneous reduction of nitric oxide and sulfur dioxide; a t these conditions the above data on NO reduction would indicate complete conversion of nitric oxide. Hence, no further work was conducted on the kinetic behavior of the nitric oxide-carbon monoxide system in the absence of sulfur dioxide. It would appear that nitric oxide reduction is not the critical reaction in the refinement of this process with the copper-alumina catalyst, unless catalyst deactivation problems for the nitric oxide reduction reaction should appear. Furthermore, even this problem is best evaluated a t the more severe conditions required for simultaneous reduction of sulfur dioxide and nitric oxide. Simultaneous Reduction of Nitric Oxide and Sulfur Dioxide. Data from our laboratories have generally supported the dual site reaction mechanism proposed by Khalafalla and Haas (1972) for the reduction of SO2 by carbon monoxide. However, Quinlan et al. (1973), have shown that an adequate description of the sulfur dioxide conversion as a function of the experimental variables can be obtained by a first-order rate constant with a correction term for CO ratio -In (1 - X S O A
kSOz = _
1nkso2 = A
_ 8
_
- RT E + B(C0 ratio
~
- 1.4)
(7) (8)
This model was shown to be adequate over a temperature range from about 700 to 900°F, contact times from 0.16 to 0.5 sec, and CO ratios from 0.9 to 1.6, in the absence of nitric oxide and water. The parameter values reported for eq 8 are shown in Table 11. Data are reported in Table I11 showing the effect of nitric oxide level in the reactor feed a t three different reactor temperatures in the presence of approximately 14% Con. These data are plotted in Figure 5, showing the effect of increasing nitric oxide levels in the reactor feed for a fixed upstream SOZ concentration of about 2200 ppm and CO concentration of about 6000 ppm. The point a t zero nitric oxide conversion is the normal activity of the catalyst before introduction of nitric oxide as described extensively by Quinlan, et al. (1973). It is apparent that an increase in the nitric oxide level of the feed results in a significant decrease in the rate of sulfur dioxide reduction.
~~
~~~
Table 111. Data for Reduction of SO2 and NO by Carbon Monoxide Upstream compositions, ppm Run no.
Temp,
DN-1 DN-2 D-41 DN-3 DN-4 DN-5 DN-6 DN-9 DN-10 DN-11 DN-12
co
OF
so2 2190 2250 2190 2250 2250 2210 2210 2160 2255 2200 2200
6140 6060 5750 6060 6060 5830 5830 6320 6160 6624 6560
796 796 79'5 770 770 770 77 1 7514 7514 7514 7513
Conversion, %
Contact time,
NO
sac
50,
cos
NO
293 990 0 995 1005 352 0 0 400 908 0
0.178 0.178 0.230 0.228 0.228 0.228 0.228 0.228 0.229 0.229 0.229
72.0 40.9 92.4 36.9 41.7 63.9 73.1 82.9 62.1 43.1 73.7
10.5 9.2 11.6 11.6 12.0 11.5 12.2 11.0 11.3 12.6 11.8
97.7 94.5 83.2 92.0 97.0 100.0 96.6
.
CO Ratio = I 35 H,O 0%
-
4.0-
A
-
I
770.F 704.F
CO RATIO I F NO REACTS
3 0-
(WINLAN ot PI 1972')
1
1
1.0
1.1
1.2 ADJUSTED
1.3 CO
1.4
1.5
RATIO
Figure 6. Dependence of SO2 rate constant on adjust d CO ratio
20-
NITRIC
One possible interpretation of the decreased activity in the presence of NO is that CO normally available for SO2 reduction is preferentially removed a t the top of the catalyst by the more rapid NO reduction reaction. As the NO level of Figure 5 is increased, ithe available CO would be decreased, thereby decreasing the rate of SO2 reduction. If this interpretation was valid, thle SO2 con? ersion should be dependent upon an adjusted CO ratio, defined as adjusted CO ratio
=
(conon of CO - concn of NO) in feed 2(concn of SO2 in feed)
(9)
This dependence should be identical with the dependence of SO2 conversion upon CO ratio determined in the absence of NO as reported by Quinlan, et al. (1973). The data of Figure 5 are plotted in Figure 6 a t two temperature levels, as the first-order rate constant of eq 7 us. the adjusted CO ratio of eq 9. As shown in Figure 6, the effect of nitric oxide is far more severe than simply determined by a carbon monoxide ratio effect a t the entering portion of the bed. Specifically the effect of nitric oxide is five times larger than would be expected simply from the effect of nitric oxide on the carbon monoxide ratio. It appears, therefore, that the nitric oxide is preferentially adsorbed on the surface sites normally available for reduction of SO2, most probably on the copper sites adsorbing CO. Alternatively, the loss of activity could be due to oxidation of the copper catalyst by NO to a less active oxidation state.
OXIDE
LEVEL, ppm
Figure 7. Dependence of SO2 rate constant on nitric oxide level
Further evidence of such a phenomena is seen by the points marked by crosses in Figure 5. These present the activity determinations of sulfur dioxide with no nitric oxide in the feed, but immediately after the catalyst had been exposed to the feed containing nitric oxide. I n this case the resulting activity was considerably lower than observed for a catalyst not exposed to nitric oxide; however, this low activity for the sulfur dioxide reduction was fully reversible. The full activity for sulfur dioxide reduction could be obtained by increasing the catalyst temperature to 825°F and then lowering it to the previous test conditions. Alternatively, after about 12 hr on stream a t the original test conditions, the activity is fully recovered, probably due to the reversible desorption of the nitric oxide from the catalyst. If this is indeed a reversible adsorption of nitric oxide, the behavior shown in Figure 5 will be dependent upon the temperature level a t which the experiment is conducted. Because the equilibrium adsorption constant for nitric oxide will decrease as the temperature increases, the effect of the presence of nitric oxide should be less pronounced a t higher bed temperatures. No extensive investigation was conducted to determine the effect of sulfur dioxide upon the catalytic activity for nitric oxide reduction. Throughout the series of runs shown in Figure 5, the nitric oxide conversion was normally above 90% Ind. Eng. Chem. Process Des. Develop., Vol. 12, No. 3, 1973
363
I 5r
Ep
mi
401
Y
I1 I
SYSTEM VARIABLESt CATALYST: M R W A W CU 0803 PARTICLE SIZES'20/50 MESH ux) 1/8 IN TEMPERATURE 700 -WO.F E ;;T ; iZ$" ,610SECS
-
EU t bg
NOLEVELS: zoo-iryx m WATER LEVELSz.8 12&!
concentration (in ppm) is sufficiently linear to allow modification of eq 8 to In kso, = A
E - RT - + B(C0 ratio - 1.4) - C(N0 level)
(IO)
As determined from Figure 7, the best value for the parameter Cis 1.33 X IO-* (ppm-I). Hence, this simple empirical model is able to account not only for variations in temperature and CO level but for the presence of NO and its subsequent depression of activity. Again, no mechanistic interpretation has been attempted in generating this model; we have merely presented a predicative tool capable of characterizing system behavior relative to the level of SO2 conversion.
A '0
PO
IO
30
40
50
60
70
SOI REDUCED, PERCENT Of
80 90 INLET SOl
100
Figure 8. Dependence of carbonyl sulfide production on sulfur dioxide conversion
t
I
8OC
Selectivity Considerations
Extensive representations of the conversion of sulfur dioxide to carbonyl sulfide as a function of the per cent sulfur dioxide reduced have been presented by Quinlan, et al. (1973), and by Okay and Short (1973). Figure 8 summarizes this relationship. Apparently for a given catalyst system, any condition of temperature, space velocity, and carbon monoxide ratio which leads to a given severity of the sulfur dioxide conversion reaction also provides a comparable severity for the carbonyl sulfide formation reaction. Also demonstrated in Figure 8 is the absence of selectivity alteration when nitric oxide is introduced into the feed. Hence, although the presence of nitric oxide depresses the activity of the catalyst for conversion of sulfur dioxide it does not alter the relative product mix. As shown in Figure 8, Okay and Short (1973) have shown much the same to be true for the presence of HzO in the absence of NO. A single data point involving both HzO and NO indicates no difference for that case. However, there is evidence that the activity depression of H20 dominates that associated with NO but that conclusion, drawn from a single point, is not well substantiated. Figure 9 presents this same selectivity data in terms of the SO2) remaining in the total sulfur compounds (COS effluent. Nitric oxide does not have any effect upon the behavior of this relationship as expected from Figure 8. Hence, as concluded by Quinlan, et al. (1973), single-bed operation is capable of a t best 80% removal of sulfur with the effluent containing both SO2and COS. The efiect of nitric oxide on the carbon monoxide-sulfur dioxide reaction is to require higher temperatures, longer contact times, or higher CO ratios to achieve this optimum than in the NO-free case.
+
10;
lo
4b
io
io
IbO
SO2 REDUCED, PERCENT OF INLET SO,
Figure 9. Dependence of total sulfur compounds remaining on sulfur dioxide conversion
and often above 95%, Although this is considerably lower activity than would be expected from the initial nitric oxide reduction experiments of Figure 4 without the presence of sulfur dioxide, it is acceptable from the point of view of the overall process with criteria of 90% reduction of sulfur dioxide and nitric oxide levels. Hence, no investigation of the dependence of nitric oxide reduction on the sulfur dioxide level was deemed necessary. The empirical model presented in eq 8 can be modified to account for the activity decline attributed to the presence of nitric oxide. I n Figure 7 is shown the relationship between the first-order rate constant for SO2 conversion and the level of nitric oxide in the reactor feed. To obtain this relationship, all points from Table I11 were first corrected to a common temperature (771'F) and a common CO ratio (1.34) via the parameters of eq 8. Note that the CO ratio required for use in eq 8 neglects the presence of NO; hence, for these calculations the NO level in eq 6 was set to zero. As can be seen in the figure, the resulting relationship between In k8o2 and the NO 364 Ind. Eng. Chem. Process Des. Develop., Vol. 1 2, No. 3, 1973
Conclusion
It is thereby possible to obtain simultaneous reduction of SO2 and NO over a copper catalyst. Nitric oxide removal is essentially complete a t all conditions considered while SOZ conversion is depressed from levels reported in the NO-free case. The selectivity and maximum sulfur removal capabilities appear unaltered by NO despite this activity depression. Nomencloture
A
= logarithm of preexponential factor for SO2 rate constant, evaluated a t CO ratio of 1.4, eq 9 and 10 B = empirical slope of plot of logarithm of SO2 rate constant us. CO ratio C = emDirical sloDe of plot of logarithm of SO2 rate constant us. NO conceitratiin (ppm), Ppm-1 E = activation energy, cal/mol kso, = firsborder rate constant for SO2 conversion, sec-I R = universal gas constant,:al/mol "R T = absolutetemperature, R
X8o2 = fractional co:nversionof inlet SO2 = contact time, based on inlet reactor conditions and total catalyst volume, sec
e
literature Cited
Baker, R. A., Doerr, R. C., Ind. Eng. Chem., Process Des. Develop., 4,188 (1965).
Bartok, W., Crawford, A. R., Skopp, A., Chem. Eng. Progr., 67, 64 - - (1971 ,- - . - ’i. ,. Bird, R. B., Stewart, W. E., Lightfoot, E. N., “Transport Phenomena,” Wiley, New York, N.Y., 1960. Chilton, T. H., Chem. Eng. Progr., 67,69 (1971). Khalafalla, S. E., Haas, L. A., J . Catal.,24, 121 (1972). Khalafalla, S. E., Foerstw, E. F., Haas, L. A., Ind. Eng. Chem., Prod. Res. Develop., 10, 133 (1971). Klimisch, R. L., Barnes, G. J., Environ. Sci. Technol., 6, 543 (1079.) \A%,.-,.
Meguerian, G. H., Lang,, C. R., S A E (SOC.Automot. Eng.) J . , IIEC SP-361 (1971). Okay, V. C., Short, W. L., Ind. Eng. Chem., Process Des. Develop., 12,291 (1973).
Querido, R., University of Massachusetts, Ph.D. Thesis, 1971. Querido, R., Short, W. L., Ind. Eng. Chem., Proc. Des. Develop., 12. 10 (1973).
Quidan, ‘C. W’., University of Massachusetts, M.S. Thesis, 1972. Quinlan, C. W., Okay, V. C., Kittrell, J. R., Ind. Eng. Chem., Proc. Des. Develop., 12, 107 (1973). Ryason, P. T., Harkins, J., J. Air Pollut. Contr. Ass., 17, 796 (1967).
Shelef, M., Kummer, J. T., Chem. Eng. Progr., Symp. Ser., 67, 117,74 (1971).
Shelef, M., Otto, K., Gandhi, H., J.Catal., 12,4 (1968). White, W. B., et al., J . Chem. Phys., 28,751 (1958). RECEIVED for review October 16, 1972 ACCEPTEDJanuary 19, 1973 This project has been financed with Federal funds from the Environmental Protection Agency Grant No. APO 1443-01. The contents do not necessarily reflect the views and policies of the Environmental Protection Agency, nor does mention of trade names or commercial products constitute endorsement or recommendation for use.
Acentric Factor. A Valuable Correlating Parameter for the Properties of Hydrocarbons Charles A. Passut and Ronald P. Danner” Department of Chemical Engineering, The Pennsylvania State University, University Park, Pa. 16806
Revised acentric factors, based on the original defining equation of Pitzer, have been determined for 192 hydrocarbons, Critical property and vapor pressure data were collected, evaluated, and then used to obtain new values foir the acentric factors. These values are recommended for use in most available correlations. Correlators are urged to use values for the acentric factor based on the literal definition rather than values determined by requiring optimum agreement with their final equations.
Presently the most useful tool in the prediction of physical properties of hydrocarbons is the theory of corresponding states. This theory hypothesizes that fluids behave similarly when they are compared a t the same reduced temperature (T/T,) and reduced pressure ( P / P c ) . Through statistical mechanics a third parameter, in addition to the critical temperature and pressure, has been introduced, which has greatly improved the accuracy of corresponding state predictions. The parameter was developed by Pitzer (1955a,b) and termed the acentric factor, W. It is defined in terms of the vapor pressure behavior of the material. w = -log (Pr’)Tr-o.7
- 1.00
(1)
where P,’ is the reduced vapor pressure (P’/Pc).The vapor pressure used is the vallue a t a reduced temperature of 0.7. The definition of w was chosen so as to make w = 0 for the heavier rare gases Ar, Kr, and Xe, Le., the simple spherical molecules. Extensive tables for correlating physical properties were developed by Curl and Pitzer (1958) in the following form G
=
+ wG(’)(T,,P,)
G(o)(I’,,P,)
(2)
where G is the particular property being correlated and G(O) and G ( l ) are tabular functions of the reduced temperature
and reduced pressure. The G(O) tables were developed for fluids obeying the simple two-parameter law of corresponding states. The G ( l ) tables were developed to account for deviations from the two-parameter law of corresponding states caused by molecular size and shape (acentricity). These tables are intended for use with normal fluids (essentially nonpolar materials) which were defined in terms of reduced surface tension by Curl and Pitzer (1958). It is clear the value of w as defined by eq 1 r d l depend on the accuracy of available values for the vapor pressure, critical temperature, and critical pressure. As more accurate values have become available for these physical properties the accepted acentric factor values have changed. Since the work of Curl and Pitwr (195S), the acentric factor has also been used in many other correlations which do not take the form of eq 2 (Edmister, et al. (1968), Fisher and Leland (1970), Johnson and Colver (1970), Lee and Edmister (1971), Starling and Han (1972)). Some of these efforts, in fact, have resulted in the workers backing out of their correlations “modified acentric factors,” Le. , empirical values which best fit their correlations. Such an approach tends to lead to a proliferation of w values and increased confusion. A comprehensive list of hydrocarbon acentric factors was presented in 1966 in the API “Technical Data Book-PetroInd. Eng. Chem. Process Des. Develop., Vol. 12, No. 3, 1973
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