Test of a Permeation Sampling Device for Integrated Sulfur Dioxide Concentrations in Ambient Air Colin M. Killick Warren Spring Laboratory, Stevenage, Hertfordshire, England
Mean ambient and artificially prepared sulfur dioxide concentrations were measured over periods varying from 1 h to 19 days using a new sampling permeation device supplied by P. W. West, Louisiana State University. The analysis was carried out by the West-Gaeke method. Concentrations obtained in this way were compared with those recorded by a continuous conductimetric instrument. For 75% of the tests, the ratio of concentrations measured by the device to those recorded by the instrument ranged from 0.80-1.17. An attempt is made to explain the results of the remaining tests, which lie outside this range. Statistical tests show that agreement between the two methods is significant.
concentrations of sulfur dioxide. As a conductimetric instrument, the Littlewood-Nash has the disadvantage of interference from other pollutants, but on the other hand, the design of the electronics has led to high sensitivity and stability ( 3 ) . In a series of unreported comparison tests with commercial instruments in a major urban area it has been shown that the Littlewood-Nash compares favorably with the Wosthoff U3S SO2 monitor (conductimetric)and the Philips PW9700 SO2 monitor (coulometric). The following regression lines were obtained over a range of ambient sulfur dioxide concentrations between 100 and 700 pg/m3. Nash = 0.886 Philips Nash = 0.922 Wosthoff
A number of manual and instrumental methods have been developed for the estimation of sulfur dioxide in the air. Considerable ingenuity has been shown in the automation of simple conductivity methods, flame photometry, and the application of specific coulometry, particularly to provide short-period or instantaneous pollution measurements. This is a far cry from the lead dioxide candle previously widely employed for the assessment of sulfur dioxide pollution over periods of a month or so. A number of factors have contributed to the candle’s decline in popularity-sampling period, nonspecificity, variable reactivity, and the dependence of results on wind speed and humidity, as well as on sulfur dioxide concentration. Many of these disadvantages apply also to the use of zinc plates or cans. Nonetheless there remains a use for measurement devices that can be sited throughout a locality to give a picture of mean distribution of sulfur dioxide concentrations over periods of a week or more, particularly devices which need no power or attention during long sampling periods. A new device in this category has been proposed by West ( I ) . Briefly, it consists of a glass cylinder closed a t one end by a silicone rubber membrane permeable to sulfur dioxide. Ten milliliters of sodium tetrachloromercurate solution are exposed in the device, in contact with the membrane. Sulfur dioxide permeating this membrane is trapped as the sulfitomercurate ion (2) and subsequently is determined by reaction with formaldehyde and bleached pararosaniline hydrochloride solutions. One of these permeation devices was made available to this laboratory. I t was decided to check its performance against that of a continuous measuring instrument normally used here to monitor ambient sulfur dioxide concentrations. Experimental Apparatus. The permeation device provided by West had been calibrated by K. D. Reiszner on 22 August 1972. The calibration factor and temperature were stated. The device was tested against a Littlewood-Nash sulfur dioxide monitor designed a t this laboratory (3, 4) working on the conductivity principle. The monitor provided 15-min mean
- 7 (correlation coefficient 0.977)
+ 11 (correlation coefficient 0.971)
Reagents. All reagents were prepared as described by West ( I ) . The pararosaniline hydrochloride stock solution was prepared from the solid (Fisher Scientific Co., No. p389) as the concentrate described by West was not available. All other reagents were prepared from AnalaR-grade chemicals supplied by Hopkin and Williams Ltd. Sampling. The permeation device was prepared by pipetting into it 10 ml of 1 M sodium tetrachloromercurate solution. The bung was replaced in the open end and the device supported with the silicone rubber diaphragm facing downward. Variations in the temperature of the air surrounding the membrane affect the rate of diffusion of sulfur dioxide through it ( I ) . During the first few experiments the ambient temperature was recorded every few hours, using a mercury-in-glass thermometer suspended nearby. The arithmetic mean of these readings was used to correct the diffusion constant. For the first five experiments, the device was tested outof-doors-in practice its most likely use. A louvered box of the type normally used to expose a lead dioxide candle ( 5 ) was modified to sit on a metal tripod. The permeation device was taped to a small piece of angled metal fitted inside the box (Figure 1). The tripod was erected about 1.5 m from the north-facing wall of a small wooden hut, such that the device was about 1.5 m above the grass. The Littlewood-Nash recorder was operated inside the hut. Outside air was drawn into the instrument through 6-mm bore P T F E tubing terminating in a funnel about 1 m from the louvered box (Figure 2). The remaining experiments were carried out inside the Laboratory. For two of these experiments the device was exposed to room air alongside the sulfur dioxide recorder. Other experiments used an artificially prepared atmosphere containing a known amount of sulfur dioxide ( 4 , 6). For these tests the permeation device was modified slightly so that it could be placed inside a desiccator through which the artificial atmosphere passed (Figure 3). The modification consisted of taping three short lengths of copper wire to the device to act as legs to hold the membrane away from the base of the desiccator. The sulfur dioxide recorder also sampled the atmosphere of the desiccator throughout the experiments. For the last two experiments outside air was drawn through the desiccator. Volume IO, Number 5,May 1976
473
PERMEATION DEVICE MEMBRANE
_
ANGLED SUPPORT
_
_
_
I I
BASE
~
TRIPOD
Figure 1. Mounting device for outside sampling
NFUNNEL FUNNEL INTERIOR O F HUT
TRIPOD
GRASS
MONITOR
This figure was used to obtain the sulfur dioxide content of the exposed solutions. Exposed Solutions. At the end of each exposure period, the louvered box was taken into the laboratory for the removal of the permeation device. The volume of solution remaining was determined in a 10-ml measuring cylinder and brought back to 10 ml, if necessary, with deionized water. The solution was transferred to a conical flask for analysis; a second flask was used for the reagent blank of 10 ml of 1 M sodium tetrachloromercurate solution. To each flask was added in turn 1 ml of pararosaniline hydrochloride working solution and 1 ml of 0.2% formaldehyde solution as with the calibration solutions. The solutions were placed in the dark for 40 min and then read against distilled water in 1-cm cells a t 575 nm. The absorbance of the exposed solution, corrected for the absorbance of the reagent blank, was multiplied by the gradient of the calibration line to calculate the sulfur dioxide content of the exposed solution. Calculation of Results. The mean concentration of sulfur dioxide in the air during the exposure time is calculated from the formula derived by West ( I ) :
WEIGHT FOR STABILITY
e =wk -
Figure 2. Experimental layout for outdoor tests
EXCESS G A S V E N T
S U L P H U R DICXIDE FROM DILUTER
--~,
- T ! c p,--9
q\
T O I h L E T OF LITTLECOOD/NASH SULFLiUR CIOXIDE R E C C ? D E R
L--4
//
-DESICCATOR
I
'LEGS"
M E ~ R A ~ E
t
where C = mean sulfur dioxide concentration, pg/m3 w = pg sulfur dioxide found in the exposed solution k = diffusion constant t = time of exposure, h For the permeation device provided by West, k was given as 1027 a t 23 "C and the temperature coefficient for the membrane described as "a permeability decrease of about 5 per cent for a 10 "C increase" ( I ) . During the outdoor experiments (tests 1-5) temperatures were recorded from time to time and the mean of these readings was used to correct k . The remaining experiments were carried out a t temperatures close to 23 "C.
Figure 3. Experimental layout for laboratory tests
Results a n d Discussion Analysis
Calibration of Sampling Reagent. Before the sulfur dioxide content of the exposed sodium tetrachloromercurate solutions could be determined, the reagent was calibrated against a standardized sodium sulfite solution. A solution of sodium sulfite containing approximately 8 g/l. was standardized against 0.1 N solutions of iodine and sodium thiosulfate. The sulfite solution was then diluted 500 times in two stages, the second stage using 1 M sodium tetrachloromercurate solution as the diluent. The final solution thus contained the equivalent of approximately 3 pg of sulfur dioxide per ml. From it, calibration standards were prepared by mixing x ml of the standard sulfitomercurate solution with (10 - x ) ml of sodium tetrachloromercurate solution, where 1 < x < 6 ml. Ten milliliters of 1 M sodium tetrachloromercurate solution was used as reagent blank. One milliliter of pararosaniline hydrochloride working solution was added to the blank and each standard in turn, followed by 1 ml of 0.2% formaldehyde solution. The solutions were shaken to ensure good mixing and left to stand in the dark for 40 min to allow full color development. They were then read against distilled water in 1-cm cells a t 575 nm, using a Pye-Unicam SP500 series 2 spectrophotometer. The reagent blank was very low, usually less than an absorbance of 0.01. The readings were corrected for the blank reading and plotted on a graph of absorbance against p g sulfur dioxide/lO ml reagent. The gradient of the resulting straight line passing through the origin was calculated. 474
Environmental Science & Technology
The results of the tests used to check the working of the permeation device are shown in Table I. As a preliminary examination, the ratios of the concentrations measured by the device and the continuous monitor were calculated. If tests 6, 7, 16, and 17 are excluded, the ratio varies, 0.801.17. This compares favorably with the short-term tests (mostly one-day exposures) used in Reiszner and West's field trials ( 1 ) . They obtained ratios of 0.63-1.08 in four tests against a coulometric monitor and 0.83-1.09 in 9 out of 10 tests using the normal West-Gaeke procedure. These present tests utilized exposure times varying from 1 h to 19 days. The ratios obtained in tests 1-5 on ambient air indicate that the device apparently recorded higher concentrations than those obtained with the continuous monitor. This could be due to a number of reasons. The correction made to h for temperature variations was approximate. On occasion the temperature was as low as 0 "C (Table 11);then the permeability of the membrane would be higher than average. If the ambient sulfur dioxide concentration was high a t such times a disproportionate amount of sulfur dioxide could have been absorbed. The continuous monitor is also temperature sensitive. It was not operated in a constant temperature environment, neither was any correction made for temperature when the mean concentrations were calculated. However, it was situated in a heated hut for these tests, hence temperature variations will have been smaller than those recorded by the device outside. During the remaining tests (nos. 6-17), the monitor and the device were at the same temperature.
Table I. Comparison Between Sulfur Dioxide Concentrations Measured by Permeation Device and Littlewood-Nash Continuous Recorder Concentration, @g/m3 Test no.
Duration
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17
7 days 7 days 7 days 19 days 2 days 7 days 7 days 2 days
3h 3h 2h l h 2% h 3% h l h 4 days 10 days
Mean temp, 10.5 5.7 4.5 5.8 9.7 22.4 20.7 21.2 21.0 B
a a a B
a
a
O C
Reagent remaining, ml
Device
Mean Nash
Ratio (devlce/Nash)
9.9 9.6 9.3 9.6 Not recorded 6.4 6.1 Not recorded Not recorded Not recorded Not recorded Not recorded Not recorded Not recorded Not recorded 8.4 6.2
35.3 48.1 85.9 38.2 97.9 10.9 8.0 105.7 1832 2620 4570 3670 1630 3340 7140 21.5 43.2
35 41 76 45 84 72 58 132 1730 2620 4090 4150 1900 3825 7370 49 70
1.0 1.17 1.13 0.84 1.17 0.15 0.14 0.80 1.06 1.o 1.12 0.88 0.86 0.87 0.97 0.44 0.62
Temperature variations not recorded, but probably within the range 20-25 OC. Note, Tests 1-5, outside, in the grounds of the laboratory; 6-7, inside, sampling room air; 8-15, inside, sampling a constant sulfur dioxide concentration: 16-17, inside, sampling outside air through the desiccator. a
The lower levels of sulfur dioxide recorded could also have been the result of absorption of ammonia by the continuous monitor. The presence of ammonium ion would lower the conductivity of the reagent and result in an apparently lower sulfur dioxide concentration being recorded. No measure of the level of ammonia in the air was available, but as negative readings have never been obtained with the Littlewood-Nash monitor a t this site, the concentrations of ammonia are known not to exceed those of sulfur dioxide. For the tests carried out inside the laboratory, neither temperature nor interferences play any part in the deviations from unity. The air used to produce the sulfur dioxide concentrations was cleaned and, dried before use and the laboratory temperatures were reasonably stable for long periods between 20" and 25 "C. West derives the formula used to calculate sulfur dioxide concentrations from basic considerations of permeability through a membrane. The formula above ( I ) shows that the amount of sulfur dioxide fixed by the sodium tetrachloromercurate solution is proportional to the sulfur dioxide concentration and inversely proportional to the time of exposure. In their consideration of the lead dioxide candle, Liang e t al. (7) calculated a mathematical model based on the diffusion resistance of the lead dioxide paste as the rate-limiting step. In this model, the amount of sulfur dioxide fixed by the lead dioxide is proportional to the square root of the sulfur dioxide concentration and initial lead dioxide concentration, and inversely proportional to the square root of the time of exposure. In neither formula does a wind speed term appear. During part of the period of outdoor experiments a t Warren Spring Laboratory, wind speed results were available from a Meteorological Office-type anemometer mounted on a 30-m mast 30 m from the hut. Mean speeds were calculated. The results are given in Table 111. No correlation can be found between the mean wind speed and the ratio of the sulfur dioxide concentrations recorded by the permeation device and the Littlewood-Nash monitor. During tests 6 and 7 on room air inside the main laboratory, the main reason for the low ratios may be the lack of
Table II. Temperature Variations During Each Test Temperature Test no.
Duration
Highest, 'C
Lowest, O C
Mean, 'C
1 7 days 14.6 6.0 2 7 11.0 0.0 3 7 8.0 0.0 4 19 13.5 -1.0 5 2 12.4 5.0 6 7 24.7 19.0 7 7 22.0 19.5 8 2 Not recorded 9 3h Not recorded No temperature data was recorded for the remaining tests
10.5 5.7 4.5 5.8 9.7 22.0 20.7 21.2 21.0
Table 111. Comparison Between Mean Wind Speeds and Ratio of Sulfur Dioxide Concentration Recorded by Device and Monitor Test no. 1 2 3 4 5
Date Nov. 1-8 Nov. 8-15 Nov. 17-24 Nov. 24-Dec. 13 Dec. 13-15
Mean wind speed, m/s
Ratio (device/Nash)
Not available Not available 6.9 6.9 5.5
1.0 1.17 1.13 0.84 1.17
movement of air past the membrane as there was no significant movement of air in the room.. This would point to one possible defect in this method of measurement, although the same effect applies to the lead dioxide candle. The results of tests 16 and 17 are unexplained. There is no apparent reason why such low results should have been recorded, as the rate of air movement through the desiccator was the same as that obtained from the artifically prepared sulfur dioxide concentrations. A plot of concentrations obtained by the device (y-axis) against those recorded by the Littlewood-Nash monitor (x-axis) indicated a linear relationship between the two. Using a Hewlett Packard Model 9100A Calculator and two prerecorded programs STAT-PAC IV-10 and 09100-70803, Volume IO, Number 5. May 1976
475
the regression line of the device on the continuous monitor was calculated. The results of these statistical tests for all results, and for all results less those obtained in tests 6 and 7 , are shown in Table IV. The significance tests on the regression analysis indicate that the value of the gradient obtained is very significant, and that of the intercept not significant a t all. As a further check on the significance of the results, a t test was carried out on all the results, using the difference between each pair of readings and the null hypothesis that the differences do not differ significantly from zero. A value of 1.15was found for t . Entering the t table a t u = 16, P is found to lie between 0.25 and 0.30 and thus our null hypothesis is upheld.
Table IV. Statistical Comparison of Results Parameter Correlation coefficient Gradient 95% confidence limits t statistic Probability U
Intercept 95% confidence limits t statistic Probability U
All results
Omitting 6 and 7
0.995 1.02 f0.06 39 -65 % 16 i-30 f141 0.453
0.995 1.02 f0.06 35 -75% 14 4-25 f169 0.313
((0.1 %
((0.1 %
16
14
Conclusions
Acknowledgment
Reiszner and West claim that their permeation device will provide reasonable mean concentrations of sulfur dioxide without the use of power for sampling, a t low cost and with the minimum of analysis. The results of the tests described above generally support this claim. The device offers not only an alternative to the lead dioxide candle as a power-free sulfur dioxide monitor but can be expected in most circumstances to provide results that are better than merely indicative. Furthermore, depending on the level of sulfur dioxide in the air surrounding the device, it has been shown that results can be obtained over periods as short as 1 h or as long as 19 days; the upper time limit is dependent only on the loss of water from the device by evaporation. The chemistry is practically specific for sulfur dioxide. The toxicity of the reagent could be a drawback in use, particularly at unattended measuring sites.
The author is indebted to Dr. Egan, the Government Chemist, London, for the opportunity to test one of Professor West’s permeation devices. L i t e r a t u r e Cited
Reiszner, K. D., West, P. W., Enuiron. Sei. Technol., 7, 526-32 (1973). (2) West, P. W., Gaeke, G. C., Anal. Chem., 28,1816-19 (1956). (3) Littlewood, A,, J . Sei. Instrum., 44,878-80 (1967). (4) Killick, C. M., ibid., ( J . Phys. E ) , Series 2,2, 1017-20 (1969). (5) British Standards Institution, London, “Methods for the Measurement of Air Pollution: The Lead Dioxide Method”, BS1747: Part 4, 1969. (6) Saltzman, B. E., Anal. Chem.. 33. 1100-12 (1961). (7) Liang, S.F., Sternling, C. V., Galloway, T. R., J.A.P.C.A., 23, 605-7 (1973). (1)
Received for review February 25, 1975. Accepted January 5 , 1976.
Reactivity of Zinc Oxide Fume with Sulfur Dioxide in Air William L. Dyson and James E. Ouon* Environmental Health Engineering, Northwestern University, Evanston, 111. 6020 1
The reactions of metal oxide particles with sulfur dioxide have many implications for understanding the air pollution phenomenon and its control. The formation of acids and sulfates under certain conditions are well known. In urban atmospheres, zinc is as prevalent as lead and the average concentration of zinc is 0.7 llg/m3 ( I ) . Zinc ammonium sulfate was considered to be partly responsible for the irritant properties of the fog in the Donora episode ( 2 ) . Amdur and Corn ( 3 ) reported that zinc sulfate and zinc ammonium sulfate aerosols were respiratory irritants capa’ ble of producing significant increases in pulmonary flow resistance in guinea pigs. The synergistic health effects of some combinations of gaseous and aerosol pollutants are generally recognized. In studying the reaction between sulfur dioxide and zinc oxide as a possible air cleaning process, Gressingh et al. ( 4 ) found that up to 50% by weight of sulfur dioxide could be absorbed by zinc oxide. They implied that zinc sulfite was the main reaction product. The health implications of zinc sulfite are unknown. Hence, depending on the nature of the reaction product and the reactivity, the sulfur dioxide-zinc oxide reaction may be a secondary source of an irritant in the atmosphere or an atmospheric sink for sulfur dioxide. The focus of this 476
Environmental Science & Technology
study was to evaluate the reactivity or capacity of zinc oxide fume for reaction with sulfur dioxide in air. The concentrations of sulfur dioxide used were 4.0-17.6 ppm. The reactivity was evaluated a t temperatures of 15, 2 5 , and 35 “C and for relative humidities of 2-95%. Differentiation was made between sulfite and sulfate reaction products. Some information on the order of the reaction and the initial reaction rate was also obtained. S h r i n k i n g Core M o d e l If, as suggested by Gressingh et al. ( 4 ) ,zinc sulfite is the main product for the reaction of zinc oxide with sulfur dioxide in air, and the reaction does not proceed in the absence of water vapor, then the reaction falls within the general class of noncatalytic gas solid reactions. Relationships for the conversion of zinc oxide as a function of exposure time when any one interaction step is rate controlling have been derived by Wen ( 5 ) for a constant diffusing gaseous reactant concentration in the gas mixture. For the case where gas film diffusion controls, the relationship is
