were found to emit about 2 g of gaseous HCl for every kg of refuse fired. Under typical operating conditions, the daily mass emissions of HC1 from two sources were calculated to be 2900 and 2200 kg. Using a point source dispersion model, maximum surface concentrations of HC1 were estimated t o range between 10 and 43 pg/m3 a t distances of up to 5 km from t h e emissions sources. Acknowledgment T h e authors wish to thank Mike Pleasant, Northrop Services, Inc., and Larry Cottone, Engineering Sciences, Inc., for their efforts in carrying out many of the arduous tasks involving field measurements for this program. Literature Cited ( 1 ) St. Clair, C. LV., F’ollut. Eng. (Sept 1978).
(2) Freeman, H., lintsiron. Sei. Technol., 12, 1252-6 (1978). (:i) Corey, R. C., in “Air Pollution, Vol. IV, Engineering Control of Air Pollution”. Stern, A. C., Ed., Academic Press, New York, 1977, pp 5:31-9:3,
( 4 ) U.S. Environmental Protection Agency, Fed Reg., 36(247), 24876-95 (1971). (5) Robertson, C. A. M., Solid Wastes Management, 64(3), 139-54 (1974). (6) Carotti, A. A,, Kaiser, E. R., J . Air Pollut. Control Assoc , 22(4), 248-53 (1972). ( 7 ) .Jackson, F. R., ”Energy from Solid Waste”, Noyes Data Corporation, Park Ridge, N.J., 1974, p 6. (8) Federal Republic of Germany Federal Laws, Air Pollution Control, par. 3.2.1.1, “Facilities LVhich Are Designed Primarily to E n tirely or Partially Eliminate Refuse from Households and Similar Materials by Combustion”, Bonn, Aug 28, 1974. (9) dahnke, .J. A,, Cheney, J . L., Fortune, C. R., J . 4ir. Pollut. Control Assoc., 2 7 ( 8 ) ,747-53 (1977). (10) Cheney, .J. L., Fortune, C. R., Sci. Total Ent’iron,, in press. (11) See ref 7, p 21. (121 Farmer. S.B., “Workbook of Atmospheric Dispersion Estimates”, Office of Air Programs Publication No. AP-26, Environmental Protection Agency, Research Triangle Park, N.C., 1970. (13) Hosler, C. R., H u l / . Am. M e t . SOC., 56, 1261-70 (1975). (11) Rubel. F. N., “Incineration of Solid LVaste”, Noyes Data Corporation. Park Ridge, N.J., 1974, pp 108-37.
Receiced for recieic. M a y 7, 1979. AcceptPd J u l y 19, 1979
Oxidative Control of Organosulfur Pollutants David C. Ayres” and Catherine M. Scott Chemistry Department, Westfield College, Hampstead, London, NW3 7ST, England
Ruthenium tetroxide is a powerful primary oxidant and may he generated economically using chlorine or hypochlorite as a secondary oxidant. I n this initial study of its possible application to pollution control we determined its effectiveness against a group of thiophenes and related odorants. Rates of oxidation of these substances in saturated potassium permanganate solution have been determined at p H 1 2 and 22 “C. These results showed t h a t a considerable residue would survive wet scrubbing with permanganate since half-life times were in the range 12.5-1215 min. With aqueous ruthenium tetroxide solutioiis, the oxidation rates were at least 100 times those found for permanganate. When control of air pollutants is attempted by means of wet scrubbing, then sodium hypochlorite ( I ) or potassium permanganate is the preferred reagent. Nucleophilic substrates may become stabilized t o oxidation by hypochlorite through their capture of chlorine; the formation from thiophene of chlorothiophenes and their chlorine adducts (2) affords an example of this. Permanganate, although more expensive, is generally more effective (3);nevertheless, its application t o odor control commonly requires better than 99% destruction of a n odorant. This optimum level is difficult to achieve, as Anderson and Adolf have shown ( 4 ) for a range of organosulfur and other odorants from rendering and food processing. T h e efficiency of wet scrubbers is limited by short contact times. We have therefore investigated t h e rates of oxidation of some common odorants by permanganate, as compared t o t h e rates achieved with hypochlorite in t h e presence of ruthenium salts. Salts such as the commercially available hydrate of ruthenium trichloride ( 5 ) are oxidized (6) by hypochlorite to ruthenium tetroxide:
+ +
+
-
[ R U ( O H ) ~ C I ~ . H ~ O 2C10]20H- --* R u 0 4 5C1- 3H20 whence R u 0 4 RuOp
+
Ruthenium tetroxide has been shown ( 7 )t o be both rapid and of wide application in its oxidation of organic compounds; numerous applications of the reagent to pollution control are implicit in the literature. We selected organosulfur compounds and thiophenes in particular for this initial study because these odorants are currently emitted by a variety of large-scale industrial processes. Thus, post-vulcanization gases contain 33% of organosulfur compounds with benzothiophene forming 12°C of the trapped oils from tire production ( 8 ) ,while it has been established (9)t h a t organosulfur release occurs during the reducing stages of brick production from Oxford clays. An extensive investigation of shale oils by Pailer and Grunhaus ( 1 0 ) showed t h a t some fractions contain 15%of sulfur comhined as thiophene and its alkyl and phenyl derivatives. Thirteen substituted thiophenes were identified (11) in t h e products of the hydrogenation of coal, and these compounds are known (12) to form part of the volatile organosulfur products from carbonization of coal. Difficulties in the treatment of organosulfur odorants emitted from coke ovens a t a concentration of 300 mg.m-3 have been described (13) where an oxidative procedure gave only 30-6070 removal. T h e problem posed by these substances in coal conversion processes has also attracted comment recently (14).
+ 2[0]
This compound is the primary oxidant and it persists at a n effective concentration as it is continuously regenerated by the action of hypochlorite (2 equiv) on t h e dioxide. 0013-936X/79/0913-1383$01.00/0
@ 1979 American Chemical Society
Experimental
Source of Materials. 2-Phenylthiophene ( m p 40 “C, ref 2 5 ) was prepared by adaption of a n existing procedure (16). Steam-distilled product was further separated by alumina chromatography using first benzene and then 5% ethanol in benzene as eluant. Final purification of these fractions was effected by preparative GLC using a 3-m column having 10% SE30as the stationary phase on Chromosorb W (60/80) a t 180 “C. T h e same procedure was used a t t h e appropriate column temperature for t h e purification of commercial samples of thiophene (110 “C), dimethyl sulfide (80 “C), 2-n-butylthiophene (130 “C), and 2-ethylthiophene (110 “C). Diphenyl disulfide (1 7 ) was prepared by concentration of an ammoniacal solution of thiophenol in ethanol; the product ( m p 60 “C) was pure by GLC standards on recrystallization Volume 13, Number 11, November 1979 1383
Table 1. Experimental Conditions and Rate Data for Potassium Permanganate Oxidations at pH 12 and 22 OC app rate consiantb Initial concn in water, g L-I
compd thiophene
1.49
2-ethylthiophene
1.82 X
2-n-butylthiophene 3-phenylthiophene 2-methyl-3-phenylthiophene
9.2 10-3 1.64 x 10-3 1.24 x 10-3
( k [ B ] o ) / 2 . 3 0 3X refa std
x
x
3-methyl-2-phenylthiophene
1.13
diphenyl disulfide
suspended in water
10-3
by expi, 1112, mln
min-1 6.4
toluene
lo-'
io3,
toluene pdichlorobenzene
108
53
12.5
time for oxidain, mln
90%
34 1 41
naphthalene naphthalene
21 5.0 (3.8) 15.8 (11.1)
33 138 44
105 524 190
naphthalene
14.7 (13.3)
47 1215
159 2295
0.50
triphenylmethane
[KMn04] = 2.1 X lo-' mol L-'. The figures given were evaluated from the experimentally determined a This was contained in methylene chloride (2.00 mL). half-life times from the relation -k[Bl0 = in 2 / t l I 2 .The constants may also be obtained from the slope of the log ( a / ( a- x ) ) to time plot = k[B]0/2.303. Those figures given in parentheses were calculated in this way for the phenylthiophenes; they differ from the experimentally based figures because the straight-line plot for these three compounds intercepts the x axis before the start. In these experiments, there was apparently a rapid initial reduction in the concentration of substrate, possibly arising from uptake on the walls of the reaction flask.
(petroleum ether, b p 40-60 "C). T h e remaining substances had been synthesized (18)and we are indebted to Professor N. B. Chapman for gifts of 3-phenylthiophene, 2-methyl-3phenylthiophene, and 3-methyl-2-phenylthiophene. Oxidations with Potassium Perr ianganate. Stoichiometry. Aliquots of thiophene (25 mL of a solution containing ca. 0.5 g in 250 mL of water) were treated with a n excess of potassium permanganate (100 mL of a solution of 7.5 g of permanganate in 1L of borate buffer, p H 12.1). After 4 days, the solution was acidified and back-titrated with iron(I1) ammonium sulfate. T h e results were consistent at p H 1 2 with the relationship ( 3 ) : 26Mn04-
+ 3CdH4S
-
26Mn02
+ 2C02
+ 10c032- + 6H20 -t 3s04'-
a demand of 13 0 atoms/molecule of thiophene. This figure was used for general guidance in the oxidation of thiophene derivatives, although clearly the pattern of their reaction products would differ in detail from t h a t of thiophene. Measurement o f Reaction Rates. With the exception of thiophene and thiophenol, it was not possible t o prepare aqueous solutions of these substrates in reproducible concentration owing to the suspension of microdroplets (19). Solutions were therefore prepared indirectly by streaming nitrogen over the organosulfur compound in a closed flask, warmed as necessary, and bubbling the issuing gas through a column of water a t ambient temperature. These initial concentrations and also those obtained sequentially during the kinetic runs were determined by extraction from water into an organic solvent (2 mL by pipet) containing a standard substance (Table I) at a known concentration, followed by GLC analysis. In a typical experiment 3-phenylthiophene was extracted from water (25 mL) with methylene chloride, and injection of a sample (1pL) gave a GLC peak (area ratio 1.23) relative to a naphthalene standard (c = 1.12 X g L-l). T h e relative response factor of 0.67 for 3-phenylthiophenenaphthalene was linear in this range; hence the concentration of 3-phenylthiophene in the aqueous solution was: (1.12 X X 2X X 103)/25 X 1.23/0.67 = 1.64 X g L-l. In the subsequent study, the rate of oxidation of 3-phenylthiophene in this solution (250 mL) was determined after addition of potassium permanganate (1.6602 g) in borate buffer (250 mL). T h e resulting solution ( p H 12.0) was stirred rapidly at 22 "C and samples (50 mL) were pipetted out a t intervals and quenched by extraction from the aqueous permanganate with methylene chloride (2 mL by pipet) containing the naphthag). T h e residual concentration of lene standard (2.247 X :%phenylthiophene a t 150 "C was then determined by GLC as above. This method depends upon the organic substrates being partitioned in a reproducible manner between accurate 1384
Environmental Science & Technology
volumes of the buffered permanganate and the extracting solvent. This was checked by diluting aqueous solutions (25 mL) of thiophene, 2-ethylthiophene, and 2-n-butylthiophene with an equal volume of the buffer and partitioning them with methylene chloride (2 mL). On dilution of the methylene chloride layer (to 50 mL) reproducible ratios were found for the extinctions (at X -240 n m ) in the organic and aqueous phases. Assuming that the extinction coefficients were the same in both phases, approximate values for the partition coefficients (CC14-H20) were calculated from the extinction ratios:
thiophene 2-ethylthiophene 2-n -butylthiophene
partition coeff 1.06 X lo2 8.8 X lo2 26.0 X lo2
proportion extracted
81% 97% >99%
There was a considerable excess of oxidant during these experiments, since thiophene a t an initial concentration of 4 X mol/L was treated with 4.6 X mol/L of permanganate. Typical of the homologues was 3-phenylthiophene a t a concentration of ca. 5 x 10-6 mol/L, equivalent to 5 x 10-5 mol/L of oxidant, in contact with 2 X mol/L of permanganate. Under these conditions all the oxidations (Table I) followed a pseudo-first-order rate law: h [Blot = log ([A]o/[A]). As the relative response plots for the substrates and the standards (naphthalene and toluene) were linear, the apparent rate constants ( h[B]o) were obtained as the slope of the plots (X2.303) of log (peak area initially/peak area a t time t ) to time. T h e records for thiophene and 3-methyl-2-phenylthiophene are shown (Figure 1).
logia/a -4 1.2
/ /
Figure 1. Typical rate plots for thiophene (-)and ylthiophene ( - - - )
3-methyl-2-phen-
Discussion T h e practice of dissolving t h e substrate by passage of a n enriched gas stream through a column of water does not ensure saturation. However, t h e resulting solutions have similar concentrations to those treated in scrubbing towers. T h e potassium permanganate solution used was essentially saturated a t ambient temperature and is capable of considerable reduction in the levels in air of thiols and alkyl sulfides (3);in contrast, the thiophenes we have examined react slowly under these conditions. T h e reaction rates we have determined are of comparable order to those found for other reactive aromatic compounds ( I ) . They are, however, much lower than those based (20) upon 90% removal from t h e vapor phase as evaluated by the organoleptic method. There are inherent errors in this technique ( 2 1 ) ,but fundamentally it relates to the initial oxidation step, whereas t h e GLC method relates to the ultimate removal of the odorants from solution. T h e enhancement of the rate of oxidation on alkylation is exemplified by 2-ethylthiophene, which was t h e most rapidly oxidized of t h e group, but nevertheless it had a half-life time of 12.5 min at 22 "C, while a contact time of 41 min was required to reduce its concentration by 90%. T h e stabilizing effect of a substituent phenyl group was evident in 2- and 3-phenylthiophene, where the rate constants were more than a n order of magnitude less than t h a t of ethylthiophene. A similar stabilizing effect is to be expected, if the nucleophilic strength of the thiophene ring is reduced by insertion of chlorine substituents a t any stage in treatment. So far, Re have not investigated t h e end products of oxidation of these thiophene derivatives, although it has long been known (3, 22) t h a t thiophene itself undergoes total cleavage when it reacts with permanganate. T h e GLC analysis of the products of t h e reactions of 2-ethyl- and 2-butylthiophene showed that intermediates of longer retention could be isolated. T h e rates of oxidation of these intermediates were slower than those of the starting materials, and they could be isolated by extraction (methylene chloride) provided t h a t permanganate was not present in excess; final purification was by preparative GLC (see Experimental). lH N M R spectra of the pure fractions showed t h a t an oxo reaction had occurred for 2-ethylthiophene, which gave rise t o 2-acetylthiophene (23) with characteristic low-field absorption (6 7.02, 1 H; 6 7.53,2 H ) and a singlet (6 2.52,3 H). T h e butyl analogue was converted into 2-butyroylthiophene (24)with comparable N M R absorption a t low field: a sextet due t o t h e central methylene group (6 1.37); one triplet due t o the terminal methyl group (6 0.97) and another due to the deshielded methylene group ( 6 2.38). Both these products were further characterized by their mass spectra, when acetyl- and butyroylthiophene gave rise to molecular ions of mass 126 and 154, respectively, together with the ion C I H ~ S +and others typical of partial cleavage of the side chains. This type of oxygenation of 2-alkylthiophenes has been described (%), and thiophene-2-carboxylic acid is formed on aeration of 2methylthiophene in the presence of cobalt salts (26).We did not detect any intermediates from methyl group oxygenation; if formed in the alkaline buffer they would give ultimate oxidation fragments more rapidly than the ketones. A direct correlation of the rates of oxidation (Table I) with the structures of the various substrates is not possible, as the balance of side-chain oxidation and direct nuclear cleavage will vary with individual compounds. Oxidations with Ruthenium Tetroxide-Hypochlorite. Stoithiometrj. Owing to t h e speed of these reactions, t h e experiments were designed to measure relative rates by competition between pairs of substrates. It was therefore necessary to take a deficit of the oxidant t o ensure t h a t a residue of each compound survived.
The oxygen demand of 2-n-butylthiophene was determined mol) in carbon tetby stirring a sample (0.260 g, 1.86 X rachloride (10 m L ) with sodium hypochlorite (10 mL, 0.77 mol) followed by ruthenium trichloride solution (10 mL, 1.25 X lo-? mol). GLC analysis showed that there was a deficiency of hypochlorite, since almost two-thirds of the butylthiophene was destroyed within 30 min at 20 "C with no further change in its concentration thereafter. Given that 2[ClO-] [Ru04] 2[O], the oxygen demand/mole of butylthiophene was: 0.77 X 10-"/1.86 X lop3 X 0.66 2 6[O] mol-'. This figure was used as a guide to quantities for t h e partial oxidation of binary mixtures. Measurement of Relative Reaction Rates. Solutions of pairs of substrates in water were prepared as described above by the passage of nitrogen over a warm mixture of the two. Alternatively, appropriate volumes of the individual solutions were mixed t o ensure that the substrate concentrations were comparable. An illustration is afforded by t h e competitive reaction between 2-n-butylthiophene and 2-ethylthiophene where equal volumes (250 mL, containing ca. 9 X l o + g L-l) were mixed. T h e ratio of their initial concentrations was obtained by taking samples (25 mL) and extracting with methylene chloride (2 mL) containing the reference substance. (In this experiment p-dichlorobenzene was used, as its t , value (X.37 min) is intermediate between that of 2-n-butylthiophene (4.55 min) and that of 2-ethylthiophene (1.41 min). T h e column was operated here a t 130 "C.) T h e oxidations were then run on further aliquots (25 m L ) to which a deficit of an aqueous ruthenium tetroxide solution (1.12 X 10-2 mol) was added. Since [RuOd] 2 [R.C4H3S]/3 and these solutions contain ca. 9 X 10-1 g L-' of the more reactive butylthiophene (M= 140). then complete oxidation of this component in the total volume (500 mL) would require:
1 10:' 9 x 10+3 X 3 mL of a molar RuOl solution 140 4 ~
Therefore the approximate volume of a 1.12 X lo-' molar solution required in a competitive oxidation of a n aliquot (25 m L ) is: 9 x 3 10' 1 - 0.25 m L 140 X 4 1.12 20 The ratios of the rate constants given for the individual runs were calculated from the relation (27): -kJ-k2 = log [unreacted fraction of l]/Iog [unreacted fraction of 21 where the declining [Ru04] term is common to both substrates, which are assumed themselves to make a first-order contribution. T h e initial concentration a1 = (initial peak areal/mlM1) X c, X M,, where c, and M , are the concentration and molecular weight of the reference, respectively. m l and M1 are the relative response factor of substrate to reference and molecular weight of the substrate, respectively. Similarly, (al - xl) = (final peak area ratiolmlM1) X c, X M,. T h e unreacted fraction of compound 1, (a1 - x l ) / a 1, then simplifies to final peak area ratio (F1)linitial peak area ratio (ZI), since t h e terms M,, cr, and M I are cancelled, as is also the term ml because linear response plots were determined for all substrates. Hence,
T h e ratios calculated for individual runs given in Table I1 show a typical spread; the mean values obtained for pairs of compounds are summarized in Table IIIA. T h e relative rates given in parentheses in the table were calculated as a cross check on the ratios obtained directly from t h e relation: k(compd A) = -kA - kc k(compd R ) k c k B Volume 13, Number 11, November 1979
1385
Table 111
Table II. Determination of the Relative Rates of Oxidation of 2-n-Butyl- and 2-Ethylthiophene by Ruthenium Tetroxide in Water at 25 O C
A. Relative First-Order Rates lor Oxidation of
Pairs of Compounds of Ruthenium Tetroxide compound
GLC peak area ratlo 2-n-Buhef 2-tthef
kdkz
""'} ""'} 0.227
0.475
initial composition by solvent extraction
0.227
0.472
after oxidation (0.5 rnL of Ru04 soln)
0.060
0.194
1.48
0.1 17 0.121 0.105
0.334 0.325 0.315
1.89 1.67 1.a7
(0.25 rnL of RuOI soln)
mean
1.73
For example, when A = 2-nlethyl-3-phenylthiophene, B = 3-phenylthiophene, and C = 2-butylthiophene, then k A / k B = 3.41h.89 = 1.80. As the figures obtained were concordant with experimental results, the same method was used to derive t h e rates relative to thiophene (Table IIIB). Discussion and Conclusions T h e rate of the oxidation of 2-n-butylthiophene in carbon tetrachloride solution could be monitored during the experiment that determined its oxygen demand. Although it is generated by the action of hypochlorite in the aqueous phase, ruthenium tetroxide is partitioned favorably into organic solvents (CCl*-H20 as 58:1), which have always been the preferred media for t h e oxidation of organic substances ( 7 ) . T h e rate of the butylthiophene reaction was therefore controlled by the decline in the concentration of t h e oxidant to zero in a second-order process. Although reaction rates will rise in the presence of excess aqueous hypochlorite, other features which are undesirable in pollution control have been identified for reactions in the organic phase. These are that oxidations of aromatic ethers, furans, and quinolines (28) may be delayed by an induction period and that radical fragments from the solvent can be trapped and incorporated in daughter compounds (29). For these reasons, and also because of its relation with existing industrial practice, we chose to study oxidation by ruthenium tetroxide in water without the addition of a buffer. While it is true t h a t t h e rates of reaction of ruthenium tetroxide with the stronger bases are p H dependent, this is not expected t o have a significant effect on the thiophenes. Any variations would be balanced out by the competitive pairing of very similar molecules of this groupall are very weak bases and a very strongly acidic medium would be needed to sensibly alter the concentration of the free substrates. Ruthenium tetroxide is insoluble in concentrated sulfuric acid, so one can exclude catalysis by its interaction with protons in water. Hence, t h e use of water as a common solvent permitted a practical comparison of t h e rates of reaction with ruthenium tetroxide and those with permanganate a t a p H typical of normal scrubbing procedures. T h e aqueous solubility of the tetroxide a t 25 "C (0.13 mol L-l or -4 X lo-* mol equiv/L) is considerably greater than that of the thiophenes, and hence pseudo-first-order reactions occur at rates that were too fast for direct measurement. Thus, 93% of the sample of 2-n-butylthiophene was oxidized within 2.5 min a t 25 "C and it was completely destroyed within 5 min; 2-ethylthiophene was totally oxidized within 2 min a t 25 "C. Based upon the relative rate order for the reactions with permanganate, it was expected that thiophene itself would be more resistant to oxidation than its derivatives. However, a sequence of measurements still could not be obtained, for thiophene was completely oxidized within 1.5 min a t 0 "C. These results indicate that the rates of pseudo-first-order 1386
Environmental Science 8 Technology
re1 rate
thiophene/anisole
1.91
2-ethylthiopheneianisole 2-n-butylthiophenelanisole 2-n-butyl-/2-ethylthiophene 2-n-butyl-/3-phenylthiophene 2-n-butyl-/2-methyl-3-phenylthiophene 2-n-butyl-/3-methyl-2-phenylthiophene 2-rnethyl-3-phenyl-/3-phenylthiophene 2-methyl-3-phenyl-/2-phenylthiophene 3-methyl-2-phenyl-/-2-phenylthiophene thiophenol/2-n-butylthiophene dimethyl sulfide/2-ethylthiophene 2-methyl-3-phenyl/diphenyl disulfide
6.90 11.2 ( 1 1.9.) 1.73 3.4 (3.38') 1.89 (1.91* 1.70 1.79 (1.80' 1.27 1.39 3.67 3.37 1.03
B. Oxidation Rates Relative to That of Thiophene as Unity compound
rate rei to thiophene (1.0)
3-phenylthiophene 2-phenylthiophene
1.8 2.6
diphenyl disulfide 2-methyl-3-phenylthiophene 2-ethylthiophene
3.2 3.3
3-methyl-2-phenylthiophene 2-n-butylthiophene dimethyl sulfide thiophenol
3.6 3.7 6.2 12 23
reactions with ruthenium tetroxide exceed those with permanganate by a t least two orders of magnitude. Very little is known about the detailed mechanism of ruthenium tetroxide oxidations, but we have shown in two instances (29,30) t h a t single-electron donation t o give radical-cation intermediates is operative. Radical cations are known (31)to feature in the oxidation of thiophene, and t h e observed rate enhancement in water as compared to organic solvents is consistent with these facts. T h e substrate concentrations are similar to those found in scrubbing towers, and the range of rates (Table 111) for thiophenes is small, implying t h a t they could be controlled by a common procedure. T h e observation that dimethyl sulfide is destroyed more rapidly than the thiophenes indicates that the tetroxide will be more effective than is permanganate for the control of sulfides and related aliphatic odorants. T h e thiophenol result is of particular interest, since permanganate is ineffective here, as the daughter compound diphenyl disulfide is extremely resistant to further attack (tllz = 1215) min; release of thiophenols from the disulfides would also occur in the environment under anaerobic conditions. In contrast, thiophenol is rapidly attacked by the tetroxide and there was no evidence of the formation of diphenyl disulfide as a n intermediate: also, its degradation rate under these conditions is of the same order as that of the thiophenes (Table IIIA). Owing to the similarity in their rates of reaction, thiophene was compared (Table IIIA) with anisole, a n observation t h a t makes it possible to compare the results reported here for thiophenes with other benzenoid derivatives, including phenols. One predicts that phenols would react rapidly under these aqueous conditions. Whether or not useful control of polluting aromatic hydrocarbons can be achieved is a matter for investigation; however, it is known that these hydrocarbons are degraded by the tetroxide although stable to permanganate under normal wet scrubbing procedures.
Literature C i t e d “Odours. Part 2”. Department of the Environment Report, Warren Spring Laboratory, Stevenage, Hertfordshire, U.K., 1975, p 112, and papers cited therein. ( 2 ) Suschitzky, H., “Polychloroaromatic Compounds”, Plenum, I~indon,1974, p 214. ( 3 ) Anderson, F. J., Posselt. H. S., Proc. Air Pollut. Control Assoc., Goth meeting (1967). (1)
Anderson, C. E.. Adolf, H. R., Proc. Air Pollut. Control Assoc., 64th meeting (1971). ( 5 ) I’antani, F., J . Less Common Metals, 4, 116 (1962). (6)Avtokratova, T. D.. “Analytical Chemistry of Ruthenium”, Israel Scientific Translations, Jerusalem, 1963, p 137. (7) Lee. D. G., Van den Engh, M., in “Oxidation in Organic Chemistry :iH”. Trahanovsky, W.S., Ed., Academic, New York, 1973, p 177. (8) Shirokov, Yu G., Drugov, Yu S., Petukhova, N. E., Chem. Abstr., (4)
83, 151592 (1976). ( 9 ) Howler. G., London Brick Co., personal communication; Brough, A , . Parry, M. A , , h’hittingham, C. P., J . Soc. Chem. Ind., 51 ( 1978). ( 1 0 ) Pailer. M., Grunhaus, H., Monatsh. Chem., 104,312 (1973). ( 1 1) Akhtar. S., Sharkey, A. G., Shultz, J. L., Yavorsky, P. M., A m . ( ‘ h ~ mSoc., . D ~ L Fuel ,. Chem. Prepr., 19, 207 (1974). ( 1 2 ) hledvedev, K.P., Petropolskaya, V. M., Khim.Ti>erd.Topl., 27-:32 (1972); Chem. Abstr., 77,91053 (1972). (I:{) Yoigt, G.. Kurapkat, H., Kloss, G., Bracht, G., Goetza, G., Ger. Off‘en. 2 205 796; Chem. Abstr., 79,138,407 (1973). (14) Koppenaal, D. W., Manahan, S. E., Enuiron. Sci. Technol., 10, 1104 (1976);Tally, .J. T., Enciron. Sci. Technol., 12, 890 (1978).
(15) (16) (17) (18)
Kues, W., Paal, C., Chem. Ber., 19,3142 (1886). Ayres, D. C., Smith, J. R., J . Chem. SOC.C , 2737 (1968). See Liu, K-T., Tong, Y-C., Synthesis, 669 (1978). Reaton, C. M., Chapman, N. B., Clarke,’ K., Willis, J. bf.,J .
’ Chem. Soc., Perhin Trans. I , 2355 (1976). (19) Furer, R., Geiger, M., Pestic. Sci., 8,337 (1977). ( 2 0 ) Posselt, H. S., Reidies, A. H., Ind. Eng. Chem. Prod. Res. Deu., 4, 48 (1965). (21) Wilby, F. V., J . Air Pollut. Control Assoc., 19,96 (1969). ( 2 2 ) Challenger, F., Ashworth, J. R., Haslam, J . , J .Chem. Soc., 127, l6;i (1926). ( 2 3 ) Reilstein’s Handbuch Org. Chem., 17, 287 (1933). (24) Steinkopf. W., Schubart, I., Justus Liebigs A n n . Chem., 424, 10 (1921). Shchedrinskaya, T. V., Konstantinov, P. A,, Litvinov, V. P.,
Ostapenko, E. G., Zakharov, I. V., Volkov, M. N., Zh. Obshch. Khim., 44,837 (1974). ( 2 6 ) Konstantinov, P. A,, Shchedrinskaya, T. V., Zakharov, I. V., I’olkov, M. N., J . Org. Chem. USSR, 8, 2639 (1972). ( 2 7 ) Ayres, D. C., Kirk, D. N.. Sawdaye, R., J . Chem. Soc. B, 1133 (1970). 128) Ayres, D. C., Hossain, A. M. M., J . Chem. Soc. Perhin Trans. I , 707 (1975). (29) Ayres, D. C., Gopalan, R., J . Chem. Soc., Chem. Commun., 890 (1976). ( : I O ) Ayres, D. C., J . Chem. Soc., Perkin Trans. 1, 585 (1978). ( : { I ) Griffin, C. E., Martin, K. R., J . Chem. Soc., Chem. Commun., 154 (1965).
R c c c i c d f o r recieu. January 22, 1575. Accepted J u l y 20, 1975.
A Five-Stage Cyclone System for in Situ Sampling Wallace B. Smith” and Rufus R. Wilson, Jr. Southern Research Institute, 2000 Ninth Avenue South, Birmingham, Ala. 35205
D. Bruce Harris Industrial Environmental Research Laboratory, U.S. Environmental Protection Agency, Research Triangle Park,
This paper describes the development and calibration of a five-stage cyclone system for in situ sampling of process streams. Cyclones may be used to advantage for collecting large samples and in sampling aerosols of high particulate concentration. T h e cyclone system was calibrated using a vibrating orifice aerosol generator and a pressurized Collison nebulizer to disperse monodisperse latex particles. At 25 “C, 28.3 L/min, and fur a particle density of 1.0 g/cm3, the Dso cut points of the cyclone system were 5.42.1, 1.4,0.65, and 0.32 jm for cyclones numbered I-V, respectively. Results from calibrating the cyclones a t several conditions of flow rate, temperature, and particle density suggest t h a t the Dzo cut points are: (a) proportional to the flow rate of the gas raised to a negative exponent t h a t is between -0.63 and -1.11, (b) linearly proportional to the viscosity of the gas, and (c) proportional t o t h e reciprocal of t h e square root of the particle density. T h e majority of measurements to determine the particlesize distribution in process streams are made with cascade impactors. Impactors, however, have several limitations: When the mass concentration is high, the sampling time may he undesirably short. Considerable experience and skill are required to obtain valid data. Particle bounce and reentrainment cause an unpredictable, but significant, error in t h e stage and backup filter catches ( I ) .
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Frequently there is not enough mass collected on some itages to be weighed accurately. Impactors are used with lightweight collection substrates, which are often unstable in mass when exposed to the process itream (2). There is not enough mass collected for chemical analysis of the particles in each size fraction. A series of cyclones with progressively decreasing cut points will perform similarly to impactors, b u t without many of t h e associated problems. Cyclones, however, also have limitations t o their applicability ( 3 ) : There is no general theory to describe the performafice of small cyclones under field test conditions. Sampling times may be undesirably long a t sources whe?e the mass concentration is low. Thus, depending on t h e nature of t h e test aerosol, either impactors or cyclones may be preferred for a certain apprication. This study was undertaken t o develop and evaluate a system containing five cyclones and a backup filter in series. T h e cyclones were calibrated, using monodisperse aerosols, over ranges in temperature, flow rate, and particle density similar to those expected for field sampling. Background
Several theories have been suggested t h a t attempt to predict cyclone behavior. Typically, the theories are based on the classical equation for centripetal force and include additional terms to describe effects such as viscous drag and turbulent
0013-936X/79/0913-1387$01.00/0 @ 1979 American Chemical Society
Volume 13, Number 11, November 1979
1387
