640
Ind. Eng. Chem. Process Des. Dev. 1981, 20, 640-646
Nomenclature a, b = constants in eq 8 a', b' = constants in eq 17 Bo = (ApgdN2/y),Bond number D = drop diameter dN = nozzle diameter Fl = buoyancy force F2 = interfacial tension force F3 = drag force g = acceleration due to gravity n = flow index in power-law model Q = volumetric dispersed phase flow rate t = time t , = detachment time u = linear velocity of continuous phase V = drop volume V, = volume of force-balance drop u, = velocity of expanding drop u, = dispersed phase velocity through nozzle W e = (u2pDo/y)(Do/dN),Weber number x = distance measured from nozzle tip y = (V/V,) Greek Letters 0 = constants in eq 11 y = interfacial tension Ap = difference in the densities of two phases CY,
p p
= viscosity
$I = angle between the
center of the nozzle and the vertical
+ = Harkins-Brown correction factor
Subscripts 0 = quiescent continuous phase 1 = very low continuous phase flow rate 2 = very high continuous phase flow rate c = continuous phase
d = dispersed phase Literature Cited Acharya, A.; Mashelkar, R. A.; Ulbrecht, J. J. I d . Eng. Chem. Fundem. 1978, 17, 230. Chuang, S. C.; Goldschmlt, V. W. J . Besic Eng. 1970, 92, 705. Davklson. J. F.; Schuler, B. 0. 0. Trans. Inst. Chem. Eng. 1980, 38, 335. Harklns, W. D.; Brown, F. E. J. Am. Chem. Soc. 1019. 41, 499. Heart@, P. M.; de Nle,I. H.; de Vrles. H. J. Chem. Eng. 1971, 26,441. Itoh, T.; Hlrata, U.; Inoue. K.; Kitagawa, Y. Chem. Eng. Jpn. 1979a, 5, 288. Itoh, T.; Hirata, U.; Inoue, K.; Yamamoto, T. Chem. Eng. Jpn. 1979b, 5 ,
a/.
313. Kumar, R.; Kuloor, N. R. Adv. Chem. Eng. 1970. 8 , 255. Sada, E.; Yasunishi, A.; Katoh, S . ; Nlshioka, M. Can. J . Chem. Eng. 1978, 56, 669. Sulllvan, S. L.; Hardy, B. W.; Holland, C. D. A I C E J. 1984, 70, 848. Takahashi, T.; Mlyahara, T. Proceedings of the Meetlng of Society of ChemC cal Engineers, Japan, 1978. Ulbrecht, J. J.; Ranade, V. R. Paper No. 1161%presented at the 72nd Annual AIChE Meeting, San Franclsco, 1979.
Received for review August 11, 1980 Revised manuscript received April 13, 1981 Accepted April 20, 1981 The work was supported in part by NSF Grant No. CPE 7916866.
= density
Thermal Decomposition of Inorganic Sulfates and Their Hydrates Jacob Mu and Daniel D. Perimutter" Depaflment of Chemical Engineering, Universiiy of Pennsylvania, Philadelphia, Pennsylvania 19 104
A study is reported of the controlled decompositions of a series of inorganic sulfates and their common hydrates, carried out in a thermogravimetric analyzer (TGA),a differential scanning calorimeter (DSC), and a differential thennal analyzer (DTA). Various sample sizes, heating rates, and ambient atmospheres were used to demonstrate their influence on the results. The purposes of this study were (1) to reveal intermediate compounds, (2) to determine the stable temperature range of each compound, and (3) to measure reaction kinetics. In addition, several solid additives: carbon, metal oxides, and sodium chloride, were demonstrated to have catalytic effects to varying degrees for the different salts.
The thermal decomposition of common inorganic sulfates has long been an important class of reactions in the chemical industry. Applications may be found, for example, in such diverse areas as ore beneficiation (McWilliams and Hixson, 1976), in metallurgical dead-roasting (Pechkovskii and Ketov, 1957), in the preparation of catalysts and molecular sieves (Wagner, 1963), and in proposed thermochemical cycles for water-splitting (Cox, 1977). Kinetic studies of such decompositions are rarely truly isothermal, for it is very difficult to establish an isothermal condition before a substantial degree of reaction has occurred in the solid. When this is the case, dynamic techniques are preferable since they monitor the change of a selected parameter in a sufficiently large temperature interval continuously. For efficient use of experimental time, dynamic kinetic studies are commonly run at relatively high heating rates 0196-4305/81/1120-0640$01.25/0
(10 OC/min or 20 OC/min). The solid reactants have been assumed to follow closely the programmed temperature increase; however, this may be a rather crude approximation when thermal decomposition is a strong endothermic reaction. Under such conditions temperature inhomogeneities may develop in the solid, as well as temperature differences between the heating phase and the solid reactant. Inconsistencies in reported values for initial decomposition temperatures of some metal sulfates were summarized by Kolta and Askar (1975). The differences in the values are as large as 150 OC. Another disadvantage to the use of high heating rates is the possible by-passing of intermediate compounds, particularly among salt hydrates. In the study reported here a series of ten common sulfates were decomposed in a thermogravimetric analyzer (TGA) with three objectives: (1) to identify intermediate 0 1981 American Chemical Society
Ind. Eng. Chem. Process Des. Dev., Vol. 20, No. 4, 1981 641 Table I. Manufacturers' SDecifications of ComDounds insoluble matter
compound ZnS0;7H20 MnSO;H,O C o s 0 ,.7 H, 0 CuSO;5H,O NiSO; 6Hz0 MgSO,*'7II,O Cr,(SO,),~n-H,O A1,(SO4),~16H,O Fez(SO,),.n-H, 0
100.47 99.5 100.3 99.2 97.0 99.9 n.a. 105.0 75.6 as anhydrous salt
0.002 0.005 0.002 0.005 0.001 0.001 0.002 0.009
chloride
nitrate
0.0005 0.001 0.0005 0.001 0.001 0.0003 0.001 0.0004 0.0005
0.002 0.005 0.001 0.01
compounds, (2) to determine the temperature range over which each compound is stable, and (3)to measure reaction kinetics as a function of conversion and temperature. Where applicable the range of study included the common hydrates of the several salts. Experimental Section The samples used in the decomposition studies reported here were all obtained as Baker Analyzed Reagent grade except for the tin sulfate, which was Fisher Reagent grade. The manufacturer's specifications are listed in Table I, showing purities always above 97% and typically over 99%. The apparatus used was a commercial instrument (Model 990) manufactured by Dupont Instrument Co. with exchangeable modules available to perform thermogravimetric analysis (TGA), differential scanning calorimetry (DSC), and differential thermal analysis (DTA). The TGA module (Model 951) monitors the weight of a material and its rate of change continuously, when heated in any inert or reactive gas, either as a function of increasing temperature or at a preselected temperature. The DSC unit measures the differential heat input between the sample and a reference and the DTA measures the temperature difference between a sample and a reference. In order to conform to common commercial practice all experiments were run on powdered samples with particle sizes between 60 and 80 pm. Preliminary runs on samples of 5,10,15, and 20 mg of each of the metal salts showed no sensitivity to sample size. Accordingly, all subsequent tests used sample weights between 10 and 20 mg as suggested by previous workers, and a nitrogen flow of 80 cm3/min was maintained through the gas space (64 cm3) over the sample to drive off the gas product of reaction. Each of the runs was also repeated in air but demonstrated no effect of this change. Information on the stable temperature ranges of the intermediate compounds and the thermal characteristics (endothermic or exothermic) of the chemical reactions involved were obtained by a series of preliminary decomposition runs on the TGA, DSC, and DTA modules at heating rates of 5,10, or 20 OC/min. After the preliminary tests were finished, TGA runs for the same materials were conducted at the slower heating rates of 1or 1.5 OC/min over the temperature ranges of interest. This time the weight and the derivative of weight were recorded on the TGA as a function of time instead of temperature, in order to identify (1)the initial decomposition temperature, (2) the final products and any possible intermediates, and (3) the stable temperature ranges of these intermediates.
Rssults and Discussion Intermediate Compositions. Each of the ten salt hydrates studied exhibited a series of relatively stable intermediate compositions as it was heated in the TGA. As a prototype for discussion it is useful to examine the results for dehydration and decomposition of zinc sulfate
2ot
oo 2w
LOO
800
600
TEMPERATURE,
OC
Figure 1. Thermogravimetric results for zinc sulfate hydrate obtained at successive heating rates of 1 OC/min (20 to 300 OC), 6 OC/min (300 to 500 "C), and 1.5 OC/min (500 to 950 OC).
_______
100
ZnSO,.
Zn5.0,
i::
5,44H20
L
. H20 _______
-1
A
1.2
w
3
0. 0
'
I
20
LO
60
BO
TEMPERATURE. C '
Figure 2. Thermogravimetricresults for zinc sulfate hydrate at a heating rate of 1 OC/min between room temperature and 80 O C .
hydrate presented in Figure 1,obtained at several heating rates depending on the temperature range: 1 OC/min (20 to 300 "C), 5 OC/min (300to 500 "C), and 1.5 OC/min (500 to 950 "C). In the as received condition, the zinc sulfate was labeled as a heptahydrate; however, because ZnSO4.7H2Ois unstable at room temperature, it was already partially dehydrated to ZnS04.4H40 at the beginning of the run. As evidenced by back-calculations based on several stable intermediate compounds (ZnS04.H20, ZnS04, and ZnO) of known compositions, the initial material corresponded at a mixture of hepta- and tetrahydrates with an effective composition of ZnS04*5.44H20. The enlarged segment of the curves reproduced in Figure 2 shows that the decomposition rate has dropped to zero
Ind. Eng. Chem. Process Des. Dev., Vol. 20, No. 4, 1981
642
0
LOO
200
600
TEMPERATURE,
8W
1000
OC
Figure 5. DTA results for zinc sulfate hydrate obtained at a heating rate of 10 OC/min between room temperature and 950 "C. Table 11. Compositions Identified by TGA temp, "C compound
TEMPERATURE, ' C
Figure 3. Thermogravimetricresults for anhydrous zinc sulfate at a heating rate of 1.5 OC/min between 550 and 890 O C . ,oOm
80
___ ZnSOb _ ' . 5 L_ L H 2 0_ _ _
:bV
0
1
3
.
'
1
1
200
1
,
'
I
LO0
s
t
,
1
600
1
8
8
,
I
1
1
800
TEMPERATURE, C'
Figure 4. Thermogravimetric results for zinc sulfate hydrate obtained at higher heating rates: A, 5 OC/min; B, 10 OC/min between room temperature and 950 OC.
a t 47 "C, indicating that the initial mixture is completely converted to ZnS04.4Hz0. This stable intermediate has never been reported before, possibly because prior studies passed this temperature range too quickly. As the temperature is again increased the ZnS04-4Hz0dehydrates further to ZnSO4.H20, the reaction having reached completion at 72 "C. The monohydrate is stable up to 224 "C, when ZnS04 begins to form. At 590 "C the anhydrous sulfate begins to convert to an oxysulfate, Zn0.2ZnS04, a process complete a t 712 "C. The slight plateau in the temperature interval 712 to 715 "C is enlarged in Figure 3 to show in the derivative curve that the two decomposition steps are completely independent. The oxysulfate intermediate was also reported by Pechkovskii (1957) and Stern and Weise (1966). The reaction to ZnO is complete at 837 "C. The initial decomposition temperature of 590 "C found for ZnS04 in this study a t low heating rate is 85 "C lower than the value reported by Kolta and Askar (1975) and 60 "C lower than that of Ostroff and Sanderson (1959). To reconcile these findings TGA results a t intermediate heating rates of 5 and 10 "C/min are presented in Figures 4 and 5 showing that a t these rates the brief plateaus corresponding to ZnS04.4Hz0 and ZnO-2ZnS04fail to be detected and the maximum rates are delayed. At still higher rates of 10 and 20 "C/min the results from TGA
ZnSO;7H, 0 ZnS0;- 4H; 0 ZnSO,-H,O ZnSO, ZnOs2ZnS0, ZnO CuSO4*5H,O CuSO;2.5H2O CuSO;H,O cuso, cuo*cuso, CUO cu,o MnSO,.H,O MnS0;0.8H2 0 MnSO, Mn,O, NiS04.6H,0 NiSO. NiO CoSO;6HZ 0 COSO,. 3.5H, 0 CoSO,.H, 0
coso,
c030, MgS04.6H,0 MgSO,.H, 0 MgSO, MgO SnSO, SnO, Al,( SO,),. 16H,O Al,( SO4);14H,O A12(S0,),~8H,0 AlZ(S0,),~2.5H,0 A12(S04)3
A1203
Cr,(S04),~6H,0 Cr,(804),~1.5H,0 Cr2(S04)3 Cr203
Fe,( S0,),.9HZO Fe,(60,),.6H,O Fe,(S0,)3.4H,0 Fe2tS04
Fe203
)3'H2
initial decompn
< 25 49 224 590 715 > 900 35 78 178 572 660 840 > 900 115 150 704 > 900 65 637 >850 45 70 180 67 0 850 45 150 780 >1,000 290 >700 25 65 110 252 580 > 900 50 325 395 > 800 < 25 90 150 220 494 > 750
final decompn 47 72 252 712 837 75 100 230 678 704 868 135 238 828 373 762 85 150 265 800 140 290 980 620 70 110 242 350 805 310 348 645 55 160 215 260 635
and DTA both exhibit maximum temperatures in the derivative curves delayed by 50 to 80 "C. Qualitatively similar results were obtained for several other sulfate hydrates, although no intermediates were identifiable in some cases and no hydrates were found in
Ind. Eng. Chem. Process Des. Dev., Vol. 20, No. 4, 1981 643 Table 111. Temperature Ranges for Several Experimental Heatinn Rates cation of sulfate hydrate Fe Cr Ni Mg
co
Mn cu Zn A1 Sn
51-
heating rate 1 "C/min
5 "C/min
1.5 "C/min
20-250 20-350 20-400 20-350 20-270 20-280 20-230 20-300 20-350 20-700
200-500 n.a. 400-650 350-750 270-660 280-700 230-570 300-500 350-560 n.a.
500-750 350-750 650-850 750-1000 660-900 700-900 570-910 500-950 560-900 n.a.
IC
Table IV. Stable Compounds Only Detected at Heating Rates G1.5 "C/min ZnO.PZnS0, ZnSO,+4H2O Fez(S0,),~4H,0 Fez(SO,),.l. 5H,O
CuS0,.2.5HZO cuo~cuso, MnS0;0.8HZO MgSO,. H 0 AIZ(SO,);14H,O Cr,(SO,), .1.5H20
FeZ(S04)3'HZ0
CoS0;3.5HZO CoSO;H20
o,o'l
0.005
Table V. Comparison of Temperatures of Initial Decomposition temperature, "C reference
this compound study
value
source
572 7 04 637 67 0 780 290 580 395 494
625 844 675 7 20 1124 360 630 455 575
Kolta and Askar (1975) Duval(l963) Kolta and Askar (1975) Kolta and A s k s (1975) Weast (1975) Perry (1972) Parazian et al. (1972) Duval(l963) Kolta and Askar (1975)
cuso, MnSO,
others. A summary of all the compositions identified in this study is presented in Table I1 together with the pertinent decomposition temperatures. Heating Rates. As noted above, TGA results are notoriously sensitive to heating rates. The several rates used in this study are therefore listed as Table 111. As was the case with ZnSO4*4H20,many intermediate compositions identified in this study could only be found at low heating rates, 11.5 OC/min. These are summarized in Table IV. It may be noted further that initial decomposition temperatures recorded for a number of compounds are appreciably lower than those reported in the literature from higher heating rate experiments. Several of these are compared in Table V. Reaction Kinetics. I t is not to be expected that any single kinetic expression would be applicable to the wide range of decompositions of this study. Nevertheless, the nth order equation da _ - ko(1 - a)nexp(-E/RT) dt is a convenient basis for comparison since it subsumes most of the prior nucleation and diffusion models (Young, 1966). Equation 1 reduces to the Erofeev form of Mampel's unimolecular law for n = 1; it conforms to the three-dimensional and two-dimensional shrinking core models for n = 1/3 and n = 1/2, respectively, and the equation follows Avrami's nucleation law (constant density of nuclei and one-dimensional nucleus growth) with n = 0. For these
0.001
1.90
1.92
1.94 IT
1.96 OK-'
198
2.00
2.02
x to3
Figure 6. Kinetic correlation test for the dehydration of zinc sulfate monohydrate.
reasons the data from each decomposition run were tested empirically by fitting to eq 1rearranged to the linear form
It should be noted that TGA data are continuous; a t least 20 points at equal temperature intervals were taken from each curve to fit eq 1. The values of the kinetic parameters were computed by standard curve-fitting procedures, using appropriate statistically evaluations for estimating the confidence intervals of the computed slope and intercept of the linearized form based on the Student "t" distribution (Bennett and Franklin, 1954). Again, it is convenient to refer to the zinc sulfate dehydration and decomposition experiments as typical. For kinetic interpretations data points were taken from Figure 1at even intervals for each of the sequential reactions and fit to the prediction of eq 1. Figure 6 shows, for example, the data at 1"C intervals for the dehydration reaction of ZnS04.H20 computed for values of n at each of five levels from 0 to 2. The coordinates were chosen to produce a straight line for the correct choice of n in accordance with eq 2. The data points for n = 2 lie on a straight line with standard error of 5.4%, whereas the data points for other values of n do not behave linearly. The activation energy E and frequency factor ko,determined from the slope and intercept individually, were found to be 111.0 kcal/g-mol and 1.83 X lo4 s-l. The 95% confidence interval for E based on the Student "t" distribution is f0.7 kcal/g-mol; the corresponding interval for k covers the range from 4.65 X 1044to 7.20 X 1046s-'. The spread of the frequency factor over a broader range arises from the great extrapolation of the best-fit line to the ordinate intercept. Figure 7 uses the same coordinates to correlate the decomposition data for ZnS04, where n was again tested at
644
Ind. Eng. Chem. Process Des. Dev., Vol. 20, No. 4, 1981
Table VI. Kinetic Parameters compound initial
final
ZnS0;5.4HZO ZnS0;4HZO
ZnS0,.4Hz0 ZnSO,.H,O
conv range for regression fit
reaction order
energy of activation kcal/g-mol
preexponential parameters, s e l
0.01-0.87 0.01-0.93
1 2
34.0 t 0.3 82.2 t 0.9
1.39 X loz1 5.81 x 1051
std error, % 2.3 7.6
15
SOL
T E M P E R A T U R E , 'C
Figure 8. Thermogravimetricresults for the dehydration of nickel sulfate hexahydrate at a heating rate of 1 OC/min between room temperature and 400 O C . O D O l L ' 103
1
'
'
1 0L
I
1
1
'
I
I
'
108
iii,
OK-'
1
112 x
1
'
I
'
1
116
lo3
Figure 7. Kinetic correlation test for the decomposition of zinc sulfate.
five levels. Twenty-two data points were picked over the temperature range of 613 to 697 "C with equal temperature intervals of 3.75 "C. The data points for n = 2 / 3 closely fit a straight line with standard error of 6.3 % , providing an activation energy of 84.0 kcal/g-mol and frequency factor of 2.46 X 10l6 s-l. A summary of the corresponding results for a variety of the salts studied is given as Table VI, but it should be emphasized that many of the decompositions examined followed rate patterns that were impossible to model by eq 1, possibly due to diffusional resistance in the solid reactant. In Figures 8 and 9, for example, data are reproduced for NiS04-6H20and Fe2(S04)3*6H20, neither of which follows simple nth order kinetics. In the case of stannous sulfate the best fit line changed from n = 2 to n = 1 at 470 "C. This may be caused by a mechanism change, but the data fit does not alone permit a definitive answer to this question. Effect of Solid Additives. The oxides of the transition elements are known to have catalytic activity with respect to the thermal decompositions of potassium chlorate (Rudolff and Freeman, 1970), potassium perchlorate (Udupa, 1975; Furuichi et al., 19741, ammonium perchlorate (Mayer et al., 1970; Korobeinichev et al., 1975), and potassium permanganate (Sugier and Phuong, 1975). Furthermore, McWilliams (1972) used carbon to decrease the decomposition temperature of manganese sulfate and increase its decomposition rate, and a similar effect on
t 0
1
"
'
~ ' 200
"
"
" ' 400
~ 600
1 800
TEMPERATURE, ' C
Figure 9. Thermogravimetric results for ferric sulfate hexahydrate obtained at successive heating rates of 1 OC/min (20 to 250 "C), 5 OC/min (250 to 500 "C), and 1.5 OC/min (500 to 750 "C).
barium carbonate was reported by Satterfield and Feakes (1959). Stimulated by these earlier reports, a series of solid additives were tested in this study in an effort to decrease reaction temperatures and/or enhance reactivity. The additives tested were carbon, sodium chloride, and a series of metal oxides. The carbon (A. B. Thomas Co.) was from animal bone sources, specified to be over 85 w t % carbon, the balance ash. The effects of carbon addition on the decompositions of metal sulfates were measured by TGA and DTA. Both results show enhanced reactivity. Again choosing ZnS04 as an example, TGA results are shown as Figure 10 for a series of experiments run at a heating rate of 20 "C/min with different amounts of carbon addition. McWilliams and Hixson (1976) proposed that this effect results when metal sulfates are pre-reduced to some unstable intermediates by carbon monoxide, which forms from the reaction of carbon and carbon dioxide. The C02 is in turn a by-
Ind. Eng. Chem. Process Des. Dev., Vol. 20, No. 4, 1981 845
20 L
'\ f
30
.-c
25[
E
20
iy
1.0
70 LOO
~
1
500
,
"
600
T E M P E R A T U R E , "C
"
'
7W
TEMPERATURE,
"
I
800
900
1WO
OC
Figure 10. Thermogravimetricresults for zinc sulfate-carbon mixtures of different molar ratios at a heating rate of 20 OC/min.
Figure 12. TGA analyses of ZnS04 and NaCl mixtures: (1)undiluted ZnS04, (2) with NaCl heated at 20 OC/min, (3)with NaCl heated at 10 OC/min. Table VII. Effect of Carbon Addition on the Decomposition Temperatures of Metal Sulfates TGA tests
compound ZnSO, ZnSO, NiSO, NiSO, MgSO, MgSO,
30 r
600-900 475-790 640-880 450-780 800-1090 560-980 700-900 510-800 cuso, 580-820 CUSO, t C 400-820 MnSO, 720-920 MnSO, + C 440-900 A1*(SO,), 540-860 Al,(SO,), + C 450-825 Fe, ( so, 1 3 540-740 Fe,(SO,), + C 350-670
+C +C +C coso, COSO, + C
TEMPERATURE, ' C
Figure 11. TGA analyses for ZnS04 and mixtures: (1)undiluted ZnSO,, (2)NiO added, (3)CrzOaadded, (4)Fe203added.
product of the reaction between sulfate and carbon. The intermediates readily decompose into the final product oxide. This multiple-step model avoids the difficulty of otherwise interpreting direct reaction between solid reactant particles having very small interfacial area per unit volume. Most of the powder in the mixture never contacts the other reactant, except via the intermediate gases. The increased decomposition rates observed at high ratios of sulfate to carbon are thus explained by the higher reaction rates between carbon and carbon dioxide, where the specific interfacial area between carbon and carbon dioxide is increased. The enhanced reaction rates lead to the production of high local concentrations of carbon monoxide which then facilitates the decomposition of the metal sulfate. The catalytic activity of each of a series of metal oxides (CrzO3, Fe2O3, NiO, Mn304,Co304,MgO, Al2O3, and ZnO) was also investigated using the decomposition of zinc sulfate as a test reaction. The TGA results on Mn304, Co304, MgO, Al2O3, and ZnO did not show any effect, but marked differences in rates were detected for samples that were mixed with Cr203,FezOs, and NiO, respectively, as shown in Figure 11. The metal oxide to zinc sulfate ratios were set at 1:l on a molar basis, and samples were heated a t 20 OC/min over the range indicated. The peaks of the derivative traces show maximum rate at 890 O C for the
decompn temp range, "C
DTA tests
temp decompn ofmax temp rate, "C range, "C
770,883 615,725 840 730 1050 880 905 765 780,800 560,800 890 765 810 745 700 625
610-915 505-810 645-885 360-790 810-1100 550-960 705-920 520-800 570-810 410-830 735-940 455-905 550-880 470-835 545-750 360-670
temp of max rate, "C 750,880 600,710 840 725 1065 870 910 720 76,805 540,800 925 740 832 760 705 630
control run, and reductions of 30, 50, and 70 "C for the three mixtures. The temperature of initial reaction similarly decreased. These catalytic effects are in agreement with those reported for the same additives by Udupa (1975)and by Pechkovskii (1957)on the decomposition rates of potassium chlorate and zinc sulfate. Rudolff and Freeman (1970) attributed such catalysis to a charge transfer mechanism in the defect structures of the transition metal oxides. By this argument, MgO, A1203, ZnO, Mn304,and Co304have too stable a structure to effect the decomposition of zinc sulfate. As a practical matter, the measured effects are relatively small; their significance lies in the demonstration of possibly larger effects of other additives. As a final additive test, a ZnS04-NaC1 mixture (molar ratio 1:2)was dried for 1 h at 300 OC to drive off hydrated water and run in the TGA. Figure 12 shows that in the presence of NaCl the decomposition begins at about 500 "C, some 250 OC lower than the control run. With heating rate lowered to 10 O C / m i n the decomposition was complete at 720 OC, a temperature at which the zinc sulfate control had just begun to decompose. This pronounced catalytic effect of sodium chloride can probably be attributed to the great reduction in melting point of ZnS04-NaC1 mixtures, combined with enhanced reactivity in the liquid state. Such an interpretation is consistent with the finding of
846
Ind. Eng. Chem. Process Des. Dev. 1981, 20, 646-651
Rudolff and Freeman (1970). The same additive carbon was also investigated as a possible catalyst for the other sulfates of this study, using a carbon to salt ratio set at 1:l on a molar basis. The thermogravimetric and differential thermal curves were qualitatively the same as obtained from the zinc sulfate. The temperature ranges for the decompositions of the several pure sulfates are compared with those of the mixtures in Table VI1 together with temperatures of maximum rate from the TGA and DTA curves for a heating rate of 10 OC/min. The initial decomposition temperature for each mixture shows a significant decrease in the presence of carbon, and in each case the temperatures of maximum reactivity are substantially reduced, up to 220 “C in the case of copper sulfate. Nomenclature E = activation energy ko = frequency factor n = order of reaction R = gas constant t = reaction time T = reaction temperature a = fractional conversion of the decomposing solid Literature Cited
Duvai, C. “Inorganic Thermal Analysis”, 2nd ed.; Elsevler: New York, 1963. Furuichl, R.; Ishii, T.; Kobayashi, K. J. ThermlAnal. 1974, 6 , 305. Koita, G. A.; Askar, M. H. T h e r m h h . Acta 1875, 1 1 , 65. Korobeinichev, 0. P.; Anisiforov, G. I.; Tereshchenko, A. G. AIAA J . 1975. 13(5),628. Mayer, S.W.: Weinberg, E. K.; Schieler, L. AIAA J. 1970. 8(7), 1328. McWilliams, J. P. Ph.D. Thesis, Unhrerstty of Pennsylvania, 1972. McWiUiams, J. P.; Hlxson, A. N. Ind. Eng. Chem. Process Des. Dev. 1976. 15, 365. Ostroff, A. G.; Sanderson, R. T. J . Inorg. Nwl. Chem. 1958, 9 , 45. Parazian, H. A.; Pizzolato, P. J.; Orreii, R. R. Thermochm. Acta 1972, 6 , 337. Pechkovskii, V. V. J. Inorg. Chem. USSR1857, 2 , 1467. Pechkovskii, V. V.; Ketov. A. N. J. Appl. Chem. USSR, 1957, 30, 1506. Perry, J. H., Ed. “Chemical Engineers Handbook”, 5th ed.; McGraw Hill: New York, 1972. Rudolff, W. F.; Freeman, E. S. J. phys. Chem. 1970, 74, 3317. Satterfieid, C. N.; Feakes, F. AIChE J. 1958, 6 , 122. Stern, K. H.; Weise, E. L. “High Temperature Properties and Decof Inorgank Salts, Part I . Sulfates”, Natlonai Standard Data Reference Serles, NSRDS-NBS-7, 1966. Sugier, H.; Phuong, N. T. H. Radiochem. Radioanel. Lett. 1975, 21, 121. Udupa, M. R. Themwehim. Acta 1975, 12, 165. Wagner, R. “Adsorbents and Catalysts” In “Handbodc of Preparative Inorganic Chemistry”; Brauer, G., Ed.; Academic Press: New York 1963. Weast. R. C.. Ed. “Handbook of Chemistry and Physics”, 56th ed.; CRC Press: Boca Raton, FL, 1975. Young, D. A. “Decomposltlon of Solids”; Pergamon Press: Oxford, 1966.
Bennett, C. A.; Franklin, N. L. “Statistical Analysis in Chemistry and the Chemical Industry”; WHey: New York, 1954. Cox, K. E., “Thermochemical Processes for Hydrosen Production”: Los Alamos Scientific Laboratory Report LA-6970-PR, 1977.
This research was funded by the U.S.Department of Energy, Office of Basic Energy Sciences, under Contract No. EX-76-S-
Received for review August 29, 1980 Revised manuscript received May 20, 1981 Accepted May 20, 1981
02-2747.
Model for Predicting Emissions from Fixed-Roof Storage Tanks James R. Beckman’ and James R. Gllmer Department of Chemical and Bio Engineering, College of Engineering and Applied Sciences, Arizona State Universiiy, Tempe. Arizona 85281
A mathematical model was developed which predicts hydrocarbon breathing losses from fixed-roof storage tanks containing crude oil. The model was able to predict within 10% error the hydrocarbon emission data taken from existing in-fiekl fixed-roof storage tanks by Western Oil and Gas Association in 1977. Due to the success of the model, it was concluded that gas inside a fixed-roof tank is stratified and not mixed. The upper strata near the tank top is most important in regulating emissions while the lower strata near the liquid surface plays a lesser role in determining tank emissions.
Introduction In 1977 the Western Oil and Gas Association (WOGA) appointed a task force to oversee a study of hydrocarbon emissions from fixed-roof storage tanks as reported by Engineering-Science, Inc. (1977). Some of the results showed that hydrocarbon vapor loases from a single storage tank amounted to 32 m3 of liquid/year (200 barreb/year). The hydrocarbon losses occurred from standing storage and were due solely to the sun’s diurnal heating and cooling of the vapor space within the storage tank. Others also have described standing storage hydrocarbon emissions from fixed-roof tanks (see API Bulletin 2513,1959; 2518, 1962; Harrer, 1978; and Danielson, 1973). Not only does this loss deplete our gasoline and crude oil supplies but it also contributes to the production of photochemical smog. Countless thousands of fixed-roof storage tanks are now in existence and are still being built since they are less expensive than the lesser polluting floating-roof type tanks. 0 196-430518 1f 1 120-0646$0 1.25f 0
The American Petroleum Institute in 1956 established the Evaporation Loss Committee to advance the basic knowledge of hydrocarbon evaporation loas from fixed-roof tanks as reported in API Bulletin 2512 (1957). The committee concluded that tank breathing losses were more or less directly proportional to true liquid pressure and the l w rate is probably less than directly proportional to vapor volume and daily atmospheric temperature change. These conclusions were the results of Boardman (1952) and Bridgeman (1955). Their work was based on equilibrium considerations assuming complete saturation of the vapor space with hydrocarbon which led to the following theoretical equation for breathing loss
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American Chemical Society
