EDGAR R. STEPHENS, PHILIP L. HANST, ROBERT C. DOERR, and WILLIAM E. SCOTT Franklin Institute Laboratories for Research and Development, Philadelphia, Pa.
Reactions of Nitrogen Dioxide and Organic Compounds in Air A mechanism is proposed for the transitory formation of ozone observed when pure nitrogen dioxide in oxygen is irradiated with artificial sunlight
DURING
recent years, the growing problems of air pollution in major cities have stimulated considerable research. I n particular, the smog problem in Los Angeles is being investigated intensively a t a number of laboratories. This research has demonstrated that photochemical reactions occurring in a polluted atmosphere play a n important part in smog development. The most striking feature of the Los Angeles smog is the presence of abnormally large concentrations of a n oxidizing agent which studies have indicated to be largely ozone. Its concentration usually parallels sunlight intensity, suggesting that it is formed by a photochemical reaction in the atmosphere. Although ozone alone probably is not responsible for all the unpleasant features of smog, the close correlation between ozone concentration and smog severity has made the formation and reactions of this compound the object of considerable investigation (77). There are two important questions concerning ozone-how is it formed in polluted atmospheres, and what role does it play in smog formation? HaagenSmit (4,6) and coworkers have presented evidence that ozone is formed by the photochemical reaction of nitrogen dioxide with hydrocarbons. Both Haagen-Smit (3, 5) and Middleton (70) report that products of reactions between
~
SPECTROMETER
/SOURCE UNIT FOR SHORT PATH SPECTROMETRY
Figure 1. Long-path infrared absorption cell
1498
ozone and various olefins can cause crop damage similar to that observed in the Los Angeles basin. In view of this evidence, it appeared desirable to make a comprehensive investigation of the photochemical reactions of ozone and nitrogen dioxide with hydrocarbons, their oxidation products, and other organic compounds known to be present in polluted atmospheres. Knowledge of the rates of reaction, products formed, and reaction mechanisms would be valuable in seeking a solution to the smog problem. I n May 1953, the Franklin Institute Laboratories undertook to study these factors under the sponsorship of the Smoke and Fumes Committee of the American Petroleum Institute. Experiments were first carried out a t fairly high concentrations of reactants (about 1 mm. of mercury). The photochemical reactions of nitrogen dioxide with hydrocarbons and other organic compounds were studied, as were the reactions of ozone with these compounds. Infrared absorption spectroscopy was used to analyze the products. Since the longest absorption path-length available was 1 meter, products in concentrations smaller than a few tenths of a millimeter of mercury could not be detected. Results of these experiments have been reported (7).
STAINLESS STEEL TUBE CONTAINING RE ACTANTS
INDUSTRIAL AND ENGINEERING CHEMISTRY
Such concentrations are far higher than those encountered in any polluted atmosphere, and estrapolating this information to realistic concentrations might be dangerous. While there is no a priori reason for believing that reactions a t low concentration must be different from those occurring a t high concentrations, there is on the other hand, no reason why different products cannot become important a t low concentration. Reactions Are Conducted in a longPath Infrared Absorption Cell
To study reactions a t lower concentrations, a special combination reaction vessel and multiple-reflection absorption cell was constructed. The design of this cell was based on that first described by White (73) and later modified by Bernstein and Herzberg (7). This multiple-reflection system uses three spherical mirrors (Figure 1)-two large ones a t the right end of the cell and one smaller mirror a t the left. Although only four passes of the light beam through the cell are shown in Figure 1, any integral multiple of four passes can be obtained by adjustment of the angle between the two large mirrors shown in Figure 3. The path length is of course limited by the energy losses encountered after many reflections from the mirrors. The three mirrors inside the cell were coated with gold to provide a high reflection coefficient combined with chemical resistance. A fan mounted inside the cell mixes the reactants. I n the first experiments a Perkin-Elmer model 12-C infrared spectrometer was used to record the spectra; later a Perkin-Elmer model 99 double-pass monochromator was used. Both instruments were equipped with sodium chloride prisms. T h e reaction vessel contained a borosilicate-glass window through which the reaction mixture could be exposed to radiation from a carbon arc placed as shown in Figure 2. This system was similar to that used for providing artificial sunlight in the earlier work a t higher concentrations, except that the average distance between the arc and the react-
A I R POLLUTION ants was much larger in the new vessel. The average intensity of irradiation was correspondingly smaller and reactions slower. T o duplicate the intensity of sunlight more closely, a n AH-6 mercury arc was substituted for the carbon arc. This arc was placed inside a large borosilicateglass tube (70 mm. in outside diameter) which was in turn inserted into the cell through a steel plate replacing the borosilicate-glass window. This resulted in a much higher average intensity of irradiation. The spectral distribution of the radiation from the mercury arc was of course not so smooth as that from the carbon arc but, nevertheless, it provided a fair comparison with natural sunlight (Figure 4). T h e data used in Figure 4 were taken from the National Carbon Co. catalog and the General Electric Co. Data Book. Mercury arc data applies to a point 50 cm. from a 1000-watt AH-6 arc in a Corning 774 jacket. The intensity of irradiation was estimated a t four points inside the cell. using samples of uranyl oxalate solution contained in borosilicate-glass flasks. This system is sensitive from the short wave length cut-off of borosilicate glass. or about 3000 A. to about 4700 A. A comparison with sunlight is shown in Figure 5. As expected, radiation intensity varied markedly along the length of the cell, but the average was of the same order as sunlight. T h e first program to be carried out with this new cell, a study of the photochemical reactions of nitrogen dioxide with various organic compounds, is reported here. The objective was to determine the products formed by these reactions a t realistically low concentrations. Most of the experiments were carried out in extra dry oxygen supplied by the Matheson Co. The prepurified nitrogen and air used in some experiments, and the nitrogen dioxide were cylinder gases obtained from the Matheson Co. The nitrogen dioxide was used without further purification. T h e sample of 3-methylheptane was obtained from the .4merican Petroleum Institute Project 44 a t Carnegie Institute of Technology. T h e l pentene was Phillip’s Petroleum Co.’s research grade hydrocarbon. The other compounds were obtained from Matheson, Coleman and Bell, or Eastman, and were used without further purification. Experiments were carried out as follows: the reaction vessel was first evacuated to a residual pressure of 1 mm. of mercury or less. I t was then filled with oxygen from a cylinder, and a blank spectrum recorded. T h e reactants were next introduced by filling a 1-liter bulb to the required pressure with the
Figure 2.
desired reactant and then flushing the contents into the reaction vessel with a few liters of oxygen. The spectrum of the reactants was then recorded, after which the arc was turned on and scanning of the infrared spectrum was begun. In some experiments the complete spectrum was scanned; in others, selected portions were repeated to obtain a detailed record of concentration versus time for compounds of particular interest.
Figure 3.
Long-path cell
“Blank” infrared spectra taken after filling the cell with oxygen showed water vapor, carbon dioxide, small amounts of hydrocarbon (chiefly methane), and traces of nitrous oxide. The most troublesome of these was the water band which begins a t about 5 microns and extends beyond 7 microns. Although this band was quite strong, even in the driest atmosphere attainable, it was still possible to estimate nitrogen dioxide
Long-path cell with tube removed VOL. 48, NO. 9
SEPTEMBER 1956
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1
1400 I c
e
0
I LOCATION OF Hg A R C
MERCURY ARC
1200 1000
/
n
T H R O T PYREX
N
E
\
800
NATURAL SUNLIGHT
3
q 600
* t
Y
E
v)
z w
400
8 6 k’
W
ga.
200 I
0 3000
4 500
4000
3500
WAVELENGTH
Figure 4.
(
3
5
6
100
DISTANCE
Figure 5.
Spectral distribution of radiant energy
4
50
)
from its band a t 6.2 microns near the center of this water band. I t is possible that some of this water and carbon dioxide leaked in from the atmosphere during the six or seven minutes required to fill the cell. For some experiments, a path length of about 430 meters was used while for others the path was shortened to 240 meters. While concentration sensitivity is, of course, greater a t the longer path length. the infrared energy is smaller so that the signal-to-noise ratio is smaller and records are less satisfactory. The strength of the water band was another reason for shortening the path to 240 meters. Furthermore, a t the longer path length the infrared signal tended to decrease after several hours of irradiation. This effect was attributed to a slight misalignment of the internal optics caused by the heating of the mirror
2
0
5000
mounts by the mercury arc. This effect was not observed a t the shorter path length. Reaction of Nitrogen Dioxide with 3-Methylheptane Produces Ozone
Shown in Figure 6 are the spectra of reactants and products in the photochemical reaction between 3-methylheptane and nitrogen dioxide in oxygen. This hydrocarbon \vas chosen as a typical branched-chain saturated hydrocarbon so that the results could be compared with the data on ozone formation which were reported by Haagen-Smit and Fox (6). Figure 6, A, represents the spectrum of the reactants only-10 p.p.m. of 3-methylheptane and 5 p.p.m. of nitrogen dioxide in 1 atmosphere of oxygen with a 240meter path length. Figure 6, B. is a spectrum recorded while the gases were being
WAVELENGTH (micron) 8 9 IO II
12
7
13
14
15
16
195
30
SO
150
200
250
F R O M E N D OF C E L L
300
(cm)
Intensity of radiation in reaction cell
irradiated by an AH-6 mercury arc. Marked under each of the important bands is the number of minutes between the time the arc was first turned on and the time at which the band was recorded. The ozone band a t 9.6 microns is readily observed. Bands of other reaction products are marked on the spectrum. There is a substantial amount of nitrogen dioxide present even after 13.5 minutes of irradiation. Bands of alkyl nitrate are evident. Bands of an unknown, “compound X,” are indicated on the spectrum, and will be discussed later. Direct tracings of the 9.6-micron ozone band from the laboratory records are shown in Figures 7 and 8. This is by far the strongest absorption band in the ozone infrared absorption spectrum. The infrared spectrum given by Herzberg (8) is somewhat misleading because it shows bands a t 4.75 and 13.9 microns nearly as strong as the 9.6-micron band. This probably results from the high concentrations used for this reference spectrum. When the spectrum is recorded a t an ozone concentration and path length which give less absorption at 9.6 microns, the other two bands are seen to be much weaker. Figure 7 shoivs the ozone formation in
40 c
W c
u
20
a 0 IO0
80
’ 60
‘
V
‘‘
k
OZONE
OD
40 20
0
I
N I T R A T E 11341
I ’
I
I
I
I
(I
.COMPD.X
I
1
I
I
I
I
I
i
I 1
Figure 6. Infrared spectra-1 0 p.p.m. 3-methylheptane and 5 p.p.m. nitrogen dioxide in 1 atmosphere of oxygen. (A) Before irradiation. (B) After irradiation by AH-6 mercury arc. Path length, 240 meters
1 500
INDUSTRIAL~ANDENGINEERING CHEMISTRY
940
97c
WAVELENGTH
1000
I
r n i ~ l o n sI
Figure 7. Ozone formation by irradiation through borosilicate glass by an AH-6 rnercurv arc of 3 13.D.m. 3methylheptane and 2 p.p.m. of nitrogen dioxide in 1 atmosphere of oxygen. Infrared light path, 430 meters .
I
A I R POLLUTION a reaction between 3 p.p.m. of 3-methylheptane and 2 p.p.m. of nitrogen dioxide in 1 atmosphere of oxygen. T h e light source \vas the glass-enclosed AH-6 mercury arc inside the reaction vessel, and the infrared path length was 430 meters. Figure 8 shows ozone formation in a reaction benveen 20 p.p.m. of 3-methylheptane and 4 p.p.m. of nitrogen dioxide in 1 atmosphere of oxygen with a carbon arc light source and an infrared path length of 324 meters. This carbon arc was outside the reaction chamber, the radiation entering the chamber through a borosilicate-glass window. The radiation was much less intense than either the mercury arc or natural sunlight. H o ~ v ever, when carbon arc radiation passes through borosilicate glass, its spectral distribution in the ultraviolet is similar to sunlight; it is ivorth noting that ozone actually \vas formed. even though in small amount. Figure 8 also sho:vs that the ozone disappears after the arc is turned Off.
Ozone Formation Depends on Concentration of Reactants
Ozone formation was studied a t two concentrations of 3-methylheptane and various concentrations of nitrogen dioxide, including some in the range expected in polluted air (Figure 9). With 3 p.p.m. of 3-methylheptane. increasing nitrogen dioxide concentration caused the maximum ozone concentration to be higher and to occur later. Similar behavior was observed with 40 p.p.m. of 3-methylheptane: except that, in duplicate experiments with 8 p.p.m. of nitrogen dioxide, the peak ozone concentration occurred earlier than expected. Ozone formation in two humidified mixtures (\as also studied. Three parts per million of 3-methylheptane and 1 p.p.m. of nitrogen dioxide reacted in 1
atmosphere of oxygen or air containing 10 mm. of mercury water vapor pressure (Figure 9). The ozone formed faster than in the similar reaction in dry oxygen. The maximum amount of ozone was roughly the same. In the experiment starting w4th 4 p.p.m. of nitrogen dioxide and 3 p.p.m. of 3-methylheptane. nitrogen dioxide was undetectable after about 5 hours' irradiation and did not reappear, even though one half the ozone disappeared in the first hour after turning off the arc. Thus the ozone did not disappear by reaction \vith nitric oxide to form nitrogen dioxide.
A Transitory Concentration of Ozone I s Formed when Nitrogen Dioxide in Oxygen I s Irradiated The small initial rise in ozone concentration evident in most of the plots in Figure 9 does not result from the presence of hydrocarbon. \Vhen 8 p.p.m. of nitrogen dioxide in 1 atmosphere of oxygen without hydrocarbon was irradiated \vith the AH-6 arc. the initial surge observed could not be repeated by turning off the arc and then turning it on
again. Even when the gas mixture was left overnight. the surge could not be repeated the next morning. Three reactions can account for formation of a small steady-state concentration of ozone in irradiated mivtures of nitrogen dioxide Tvith oxygen : hv NO2
0
3
+
NO
+0
0 + 02+Oa + so NO2 + O? +
HoIvever, they do not account for the decline of the ozone concentration because the concentrations of ozone and nitric acid would alivays be equal. This decline can be explained by assuming a reaction that \vould allow niiric oxide to build u p in excess of ozone. such as :
0
+ NOy
+
NO
+ 0:
If the nitric oxide forms slightly faster than ozone: the ozone concentration will rise and then fall, just as observed. This scheme accounts not only for the ozone surge but also for the fact that i t cannot be repeated. The first irradiation pro-
3 p p m 3-METHYLHEPTANE
4 0 ppm 3 - M E T H Y L H E P T A N E
kT--AFTER
ARC ON
I I_ 9.40
2;
-L '
1 -
9.70
WAVELLNGT'I
HOURS ~
10 00
l m i ~ r o n Ii
Figure 8. Ozone formation by irradiation through borosilicate glass with a carbon arc of 20 p.p.m. 3methylheptane and 4 p.p.m. of nitrogen dioxide in 1 atmosphere of oxygen. Infrared light path, 324 meters; slit, 400 microns
"
( IRRADIATION
TIME
1
2
3
4
5
6
( hr.)
Figure 9 . Ozone formation in mixtures of 2-methylheptane, nitrogen dioxide, and oxygen by irradiation with an AH-6 mercury arc. Path length, 430 meters VOL. 48,
NO. 9
SEPTEMBER 1956
1501
. -
E
3.0
,"
2.5
Q
- ARC
A DO 2 5 PPm
2.0 0 2
NO
1.5
1.0
g
0.5
0
0
N
0
7 14
8
0
T I M E (rnin) Figure 10. Ozone formation by photolysis of 50 p.p.m. of nitrogen dioxide in 1 atmosphere of oxygen. Ozone formation measured by absorption at 9.5 microns. Path length, 192 meters
duces nitric oxide which prevents the formation of ozone when the arc is turned on a second time. T o test this theory, a n experiment using 50 p.p.m. of nitrogen dioxide in dry oxygen was run with the spectrometer set a t the wave length of maximum absorption for ozone (Figure 10). When the arc was turned on, ozone formed immediately and disappeared rapidly. This surge could be repeated only after 25 p.p.m. of ozone had been added to the mixture to remove nitric oxide.
p.p.m. ofnitrogen dioxide in 1 atmosphere of oxygen was prepared by adding a small amount of ozone to a mixture of nitrogen dioxide in oxygen. After the mercury arc was turned on. the ozone concentration was followed by recording its absorption a t 9.6 microns. In Figure 11. this experiment is compared with one in which 5 p.p.m. of nitrogen dioxide in oxygen was irradiated. The more rapid build-up of ozone in the mixture containing nitrogen pentoxide confirmed the predictions \vhich had been made.
o z o n e Formation IS Enhanced by Nitrogen Pentoxide
Ozone Formation and Nitrogen Dioxide Disappearance in Reactions involving Various Organic
I t is also apparent that a wide variety of organic compounds can produce ozone. Olefins, though they form ozone very quickly, produce a peak ozone concentration which is no higher than that attained by some of the other compounds which produce ozone much more slowly--e.g., 3-methylheptane and butyraldehyde. T h e disappearance rates of nitrogen dioxide when photolyzed in nitrogen and in oxygen are shown for comparison in Figure 13. In nitrogen, nitrogen dioxide could not regenerate by reaction of ozone with nitric oxide; therefore, it disappeared more rapidly than in any other experiment. Presumably, the ultimate product is nitric oxide. Unfortunately this compound has only one absorption band in the 2- to 15-micron range; this band is so weak that small amounts cannot be detected.
Other Reaction Products Are Observed
Infrared spectra for methyl ethyl ketone, 1-pentene, 2-pentene, and butyraldehyde with 5 p.p.m. of nitrogen dioxide in 1 atmosphere of oxygen were observed both before and after irradiation. (Figure 6 shows the products of the 3methylheptane-nitrogen dioxide reac-
I
These results suggested that ozone would accumulate if a substance were added to the nitrogen dioxide-oxygen mixture which would react rapidly with nitric oxide to remove it from the mixture. After a number of unsuccessful attempts, nitrogen pentoxide was proposed as a substance to fulfill the requirement. A mixture containing about 4 p.p.m. of nitrogen pentoxide and 5 1.6
I
I
1.4 L
0
2-PEHlENE
0
METHYL
I
n - BUTTRALQEHYOE
I
n - BUTYRIC
E'HYL
a- BUTYL
T h e ability of various organic compounds to produce ozone when photochemically reacted with nitrogen dioxide, was studied a t concentrations of 10 p.p.m. of organic compound and 5 p , p , m , of nitrogen dioxide in oxygen. T h e rate of disappearance of nitrogen dioxide was also followed (Figures 1 2 and 13). Data in the following tabulation reveals a crude parallel between the rate a t which ozone forms and the rate a t which nitrogen dioxide disappears:
KETOHE
ACID ALCOHOL
i ~
0
,
2 IRRAQIATION
3 TIME
4
(hl)
Figure 12. Ozone formation in photochemical reactions of 5 p.p.m. of nitrogen dioxide with organic compounds, each at a concentration of 10 p.p.m., in 1 atmosphere of oxygen
0.4
z
0.2
0 0
1
2
3
IRRADIATION
4
5
6
7
Organic Compd., 10 p . p . m ~ , Reacted with NOz. 5 p.p.m, 2-Pentene 1-Pentene Biacetyl n-Butyraldehyde 3-Methylheptane 2,2,3-Trimethylbutane Methyl ethyl ketone n-Butvl alcohol n-Butyric acid Isopropylbenzene Benzene Propane
TIME (min)
Figure 11. Ozone formation by irradiation of nitrogen oxides by AH-6 mercury arc. Infrared path, 276 meters
1 502
Aooroximate Time Reauired for: _. NO2 to fall to 2 Ozone to reach 1 p.p,m. of initial conc. 2 min. 7 min. 20 min. 10 min. 10 min. 10 min. 40 min. 30 min. 1 hr. 40 min. 2 hr. 30 min. 2 hr. 1 hr. 30 min. 2 hr. 30 min. 1 hr. 40 min. Only small amount 45 min. formed 45 min None detecteda 1 hr. None detecteda 1 hr. None detected 2 hr. None detected & .
1.0
0.6
0 N 0
I-PEHTENE
D
1.2
0.0
W
3-METHYLHEPTAHE
1
A very small amount of ozone could have gone undetected because of interferences in the ozone spectral region.
INDUSTRIAL AND ENGINEERING CHEMISTRY
A I R POLLUTION
I
r
X
A L L ORGANIC C O M P O U N D S AT A C O N C E N T R A T I O N O F IOppm I N I ATM O F 0 2
~
o
X
A
0
A
I
acid reacts with nitrogen dioxide, it seems likely that reduction of alkyl nitrate to a minor product results from the fact that a t parts-per-million concentration, little or no peroxyacid is formed.
3-METHYLHEPTANE I - P E N T E-N Ez-PENTENE KETONE
I
n- BUTYRALDEHYDE n-BUTYRIC
0
n-BUTYL
A
NO;! ALONE I N I ATM OF NITROGEN NOp ALONE IN I ATM OF OXYGEN
TIM€
What Is the Mechanism of Ozone Formation?
ACID ALCOHOL
4
3
IRRADIATION
ETHYL
V
V
2
METHYL
The peculiar shapes of the plots of ozone concentrarion versus time, and their strange dependence on concentrations of hydrocarbon and nitrogen dioxide (Figure 9) are not yet understood. These data confirm, however, several features of this reaction which were reported by Haagen-Smit and Fox ( 6 ) , namely :
5
(hr)
Figure 1 3. Disappearance of nitrogen dioxide in photochemical reactions with organic compounds
tion.) Carbon dioxide and ozone are products common to all of these experiments. Carbonyl bands also appear strongly in all the product spectra, although in the experiments with butyraldehyde and methyl ethyl ketone, part of them may have been due to the reactant. Both 1-pentene and 2-pentene produced aldehyde. Alkyl nitrate, a prominent product in the earlier work a t high concentration, was also formed in several of these low-concentration experiments. Methyl ethyl ketone did not produce alkyl nitrate while butyraldehyde, both pentenes, and 3-methylheptane did. Methyl ethyl ketone and 3-methylheptane failed to produce carbon monoxide. Both 1-pentene and 2-pentene gave indications of formic acid. S o evidence was obtained for the formation of peroxyacid, a product that was found in the earlier work a t higher pressure. The bands identified as ozonidr bands in the spectra of the products of the nitrogen dioxide-olefin photochemical reactions need some comment. Briner, Susz, and Dallwigk ( 2 ) have studied spectra of ozonides, and report their strongest band to be in the carbonyl region (about 5.8 microns) and weaker bands in the region of 8 to 9 microns. T o check this, ozone and olefins were reacted a t parts-per-million concentrations in the long-path infrared absorption cell. The same type of spectra were obtaineda strong band in the carbonyl region and weaker bands near 9 microns. These are the bands that were tentatively identified as ozonide bands on the spectra. Further work needs to be done on these reactions.
Several bands, more or less prominent in all of these spectra, have not yet been identified; most are believed to belong to one compound (or a class of compounds) since they appear always to increase and decrease together. On several occasions the reaction products were left in the cell overnight with the arc off. I n these cases, bands a t 5.45, 7.7, 9.6, and 12.65 microns were observed to be weaker the following morning even though other bands were unchanged. This suggests that the compound responsible for these four bands (referred to as compound X) is unstable. T h e band a t 12.6 microns is a t a wave length characteristic of alkyl nitrites. However, alkyl nitrites (9) also have a strong band a t about 6 microns which is not present in these spectra. Further work is in progress to identify this substance and to isolate pure samples.
Products Formed Depend on Reactant Concentration
I t is worthwhile to compare some of the results obtained in the present work a t parts per million with previously reported results obtained in the study of similar reactions a t concentrations of the order, parts per thousand (7). Ozone is now detected; it was not detected previously. Alkyl nitrates are now relatively minor products; they previously were major products. Peroxyacids are not now detected as products of the nitrogen dioxide-olefin and nitrogen dioxide-aldehyde reactions as they were before. Since it was previously shown that alkyl nitrate is formed when peroxy-
1. There is an upper limit for the nitrogen dioxide concentration above which detectable amounts of ozone are not produced within a limited irradiation time. 2. This limit is increased by increasing the hydrocarbon concentration. 3. Continued ultraviolet irradiation reduces the nitrogen dioxide concentration and makes the formation of ozone possible, which means that the upper limit concentration of nitrogen dioxide is somewhat artificial. 4. Ozone disappears a t a moderate rate after the ultraviolet irradiation is stopped. The experiments on photolysis of nitrogen dioxide in oxygen show that it is possible to produce small amounts of ozone without adding organic compound. The amounts, however, are smaller than those observed in the presence of a n organic compound. HOW then does the organic compound participate in the reaction? Two theories have been proposed. I n one, it is postulated that the ozone is formed by the mechanism outlined earlier, and that the organic compound in some way reduces the concentration of nitric oxide, thus allowing the ozone to accumulate. It might be assumed that the organic compound, or more probably free radicals formed during its oxidation, reacted with the nitric oxide in much the same way that nitrogen pentoxide reacted with it to allow the ozone to accumulate. But then, the ozone concentration could never exceed the initial concentration of nitrogen dioxide because each molecule of the latter would produce no more than one molecule of ozone. I n some of the experiments at low concentrations (Figure 9), as well as in many of the experiments reported by Haagen-Smit and Fox (O),the maximum ozone concentration does exceed the initial nitrogen dioxide concentration. This may be resolved by postulating that the organic compound regenerates nitrogen dioxide from nitric oxide by a reVOL. 48,
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action like that suggested by Stanford Research Institute (12): Ron
+ NO
+ RO
+ NO2
tvhere ROz represents a n alkyl peroxide radical formed i n the oxidation of the organic compound. If nitrogen dioxide is regenerated. then each molecule initially present could produce many ozone molecules a n d the concentration of the latter could exceed the initial nitrogen dioxide concenrration. There is the follolving objection of this scheme of nitrogen dioxide regeneration. If the function of the organic compound is to oxidize the nitric oxide to nirrogen dioxide, then it would be expected that compounds most effective for oxidizing nitric oxide would be most effective in promoting ozone formation, and, rherefore, most effective in retarding the disappearance of nitrogen dioxide. But the reverse is true; i n rhe experiments where ozone is generated most rapidly. nitrogen dioxide disappears most rapidly. This does not completely rule out this mechanism, however. It may be that the organic material reacts ivith nitric oxide in two ways-combining with part of it a n d reoxidizing part of it to nitrogen dioxide. Perhaps the relative rates of these tivo reactions are dependent on concentration so that regeneration of nitrogen dioxide is favored a t the very low concentration (belo\\- 3 p.p,m.) \
