Formation and chemical reactions of atmospheric particles

Formation and Chemical Reactions of Atmospheric Particles R. D. CADLE National Center for Atmospheric Research,~ Boulder, Colorado 80303 Received February 26, 1971; accepted March 3, 1971 Particles may enter the atmosphere preformed, for example, as a result of dust storms or industrial operations. They may also be formed by chemical reactions in the atmosphere and may be changed physieMly and chemically as a result of such reactions. The particles of photochemical smog are formed when nitrogen dioxide is photolyzed and the resulting atomic oxygen reacts with certain of the organic vapors present in such smog to trigger a series of complex reactions. There is considerable evidence that such reactions also occur in the ambient atmosphere involving essential oils emitted by vegetation. Hydrogen sulfide and sulfur dioxide are oxidized both in smog and in the ambient atmosphere by several different processes, the end products being sulfuric acid droplets and various sulfates. Sea salt particles are prevMen~ over the oceans and they undergo a number of chemical reactions including reactions with oxides of nitrogen and sulfuric acid droplets to liberate hydrogen chloride and produee nitrates and sulfates. Sulfuric acid droplets produced by the oxidation of H2S and SOs react with atmospheric ammonia to form ammonium sulfate. INTRODUCTION

chemicM reactions that produce "photochemical smog" (1, 2). Such smog results from the action of sunlight on a mixture of air, oxides of nitrogen (especially nitric oxide, NO, and nitrogen dioxide, NO~), and organic vapors such as the unburned hydrocarbons in automobile exhaust gases. Although several types of initiating reactions m a y occur, the most important is probably the photolysis of nitrogen dioxide to form nitric oxide and atomic oxygen. The atomic oxygen reacts with molecular oxygen to form ozone (03) but ozone reacts with nitric oxide to form nitrogen dioxide and molecular oxygen. The amount of Oa produced by this sequence is much smaller than that often occurring in photochemical smog. However, atomic oxygen reacts with hydrocarbons to form free radicals and these undergo a large sequence of reactions, some of which result in the oxidation of NO to NO2. Other reactions m a y produce 03 and some produce unpleasant products such as formaldehyde, acrolein, and peroxyacetyl nitrate. In simplified form, the reactions are:

The atmosphere is a huge aerosol system and the particles in it are produced in many different ways. Some enter the atmosphere preformed, for example as soil dispersed by winds, as volcanic "ash," and as micrometeorites. Others are produced in situ by chemical reactions of gases and condensation of vapors. Most such particles, however they are introduced into the atmosphere, m a y undergo chemical changes while they are in the air. The formation of airborne particles in situ by chemical reactions, and the chemical reactions airborne particles m a y undergo once they are formed, are discussed in this article. PARTICLE FORMATION Photochemical Smog A city can introduce great quantities of particles into the atmosphere, and a major mechanism of production is the sequence of The National Center for Atmospheric Research is sponsored by the National Science Foundation. Copyright @ 1972 by Academic Press, Inc.

NO2 + h~ --~ NO + O,

[11

Journal of Colloid ancl Interfitce Science, Vol. 39, No. 1, April 1972

25

26

CADLE O + O~+M--*O3+M, O, + NO --~ NO2 + O2 , O + olefins --+ R'O, or O / \ R --R',

[21 [3]

[4]

O3 ~- olefins -~ products,

[5]

R + O2 --* RO2 ,

[6]

RO2 + O~ --* RO + O~ .

[7]

Reactions [4] and [5] initiate a series of reactions producing various organic acids, aidehydes, ketones, and nitrogen-containing compounds. Reaction [7] may be the one mainly responsible for ozone production, but this is questionable. Ozone may accumulate as a result of RO2 reacting with NO in competition with reaction [3]. The direct excitation of various organic molecules by sunlight may lead to other reaction chains. Only a small part, probably about 5 % of the organic vapor in photochemical smog is converted to particles. Research with artificial smogs indicates that aerosols are formed when many, if not most, 6-carbon and larger straight chain, branch chain, and cyclic olefins, and various aromatic hydrocarbons mixed with nitrogen dioxide in air are irradinted.

Oxidation of Essential Oils Reactions similar to those occurring in photochemical smog probably also occur in the natural atmosphere. In this case, the organic vapors (essential oils) are given off by plants. Many species of plants emit such vapors, and hundreds of organic compounds have been identified in essential oils. Terpenes are especially common, a-pinene occurring in oils from over 400 species. Oxides of nitrogen are produced naturally by nitrogen fixation in forest and grass fires. Furthermore, many of the essential oils absorb solar radiation and the resulting electronic excitation may lead to reaction chains producing particles. Such photochemical reactions involving naturally occurring organic vapors may be responsible for much natural haze, such as the very intense haze that forms over the jungles of Colombia (3, 4). Rasmussen and Went (5) have estiJournal of Colloid and Interface Science, Vol. 39, No. 1, April 1972

mated that the worldwide emission of organic vapors to the atmosphere by plants is 4.4 X 10s tons/year, while Ripperton et al. (6) suggest that this estimate may be too low by a factor of 2 to 10. Some of these organic compounds, those that are highly unsaturated, oxidize rapidly even in the absence of sunlight by attack of oxygen on the carbon atom in an a-position relative to a double bond: --C--C=C--C--+ O~ -~ --C--C~C--C--

I o--o

[8l

The peroxide radical thus formed may undergo various reactions leading to particle formation. The reactions are similar to those that occur during the autoxidation of cooking fats and of drying oils.

Sulyates Sulfates and sulfuric acid droplets are a major fraction of the aerosol particles over much of the world, and in both the troposphere and stratosphere. Part of this sulfate is from the oceans, incorporated in dried-out droplets of sea water, but more is produced by the oxidation in air of hydrogen sulfide (H2S) and sulfur dioxide (SO2). The former is produced largely by the decay of biological materials, while the latter is formed by the oxidation of H2S, the combustion of sulfurcontaining fuels, and volcanoes (7). Hydrogen sulfide does not react with 02 at an appreciable rate in the gas phase. It does react slowly with 03, but at rates probably too slow to be important (8). Furthermore, hydrogen sulfide does not absorb solar radiation; thus it does not undergo photolysis or react photoehemically with molecular oxyten. Hydrogen sulfide, 03, and 03 are soluble in water. The oxidation of H2S by these oxidants in aqueous solutions, for example in fog or cloud droplets, may be very fast, but these reactions have not been studied. Hydrogen sulfide reacts rapidly with atomic oxygen : H~S + O -+ OH + HS,

[9]

with a rate constant at 300°K of 4 × 10-14 cm3 molecule-1 see-1 (9). Reaction [9] is followed by a chain reaction leading to products such as SOs, SOs, H2SO4, H2, and

ATMOSPHERIC PARTICLE FORMATION

27

TABLE I ANALYSESOF PARTICLESCOLLECTEDON CELLULOSE("IPC") FILTERSIN THE STRATOSPHEt~,E OVEn T~IECnNTRALUNITED STATES Conc (~g/m~ ambient)

At start of sampling Latitude (N)

Longitude (W)

(°)

(min)

(°)

(rain)

37 33 31

00 40 45

109 112 106

05 45 25

SO~-

0.20 0.27 0.17

Si

0.040 0.036 0.10

Na

0,00 0.012 0.035

CI

NOa-

NH4+

0.068 0.043 0.094

0.22 0.26 0.21

0 0 0

All samplings were 30 min long at 18 km altitude. The samples were collected on Sept.. 9, 1969. H20. This reaction must occur to a slight extent in photochemical smog, and will oxidize any H~S reaching the stratosphere. Sulfur dioxide in its electronic ground state does not react at an appreciable rate with either O~ or 03. I t does undergo a third-order reaction with atomic oxygen: SO~ + 0 + M --* SOs + M.

[10]

This reaction, like the corresponding reaction of H2S, must occur to some extent in smog. I t is probably a very important, reaction in the stratosphere, contributing to the formation of a worldwide layer of particles having a peak concentration at 16-20 km altitude (10). We have been malting a study of the composition of this layer, collecting samples on filters and impactors flown on RB-57F aircraft (11). The results obtained on one such flight for a series of cellulosefiber filter samples is shown in Table I. The predominant species were SO~-, probably present largely as sulfuric acid droplets, and NO3- which was probably largely present in the stratosphere as HNO3 vapor and was absorbed by the cellulose fibers. No NIl4 + was found in these samples, although small amounts had been found on other flights. Sulfur trioxide never exists to an appreciable extent in the open atmosphere. I t reacts extremely rapidly with water vapor to produce sulfuric acid (12) which may undergo chemical reactions (for example, to form ammonium sulfate), condense to form droplets, or be adsorbed on solid particles. SOs -}- H20 --+ H2SO4 .

[11]

Sulfur dioxide absorbs near-ultraviolet radiation, becoming electronically excited. Numerous studies have shown that the

excited species reacts at an appreciable rate with 02, but reported quantum yields (4) vary by several orders of magnitude. Recent studies in our laboratories suggest that • for the near-ultraviolet region of sunlight is about 2 ,~ 10-3, too small for this reaction to be significant in the atmosphere. Electronically excited sulfur dioxide may react with ozone, but this reaction has not been studied. When sulfur dioxide is included in an artificial photochemical smog (an irradiated mixture of air, NO, NO2, and an appropriate hydrocarbon), the SO2 is oxidized and hydrated to H2SO4 droplets much more rapidly than in clean irradiated air (13, 14). Although the mechanism of the oxidation is not known, it m a y result from reaction with the alkyl peroxide radicals produced by reaction 6: ROO -b SO2 --+ SO3 + RO.

[121

This, of course, would be followed by reaction [11]. As mentioned earlier, photochemical reactions similar to those occurring in smog can occur in the ambient atmosphere. I t is possible that. reactions such as [12] are responsible for most of the oxidation of sulfur dioxide in the troposphere. Volcanoes at times erupt tremendous quantities of sulfur compounds into the troposphere and stratosphere. Thermodynamic considerations suggest that little sulfur in the gas phase in the depths of a volcano exists as SO3, most being present as S02, S, and H2S (15). Recent studies by Cadle et al. (16) at the volcano Kilauea in Hawaii indicate that the sulfur compounds may be largely oxidized to S03 immediately after ejection into the atmosphere where the gases are mixed Journal of Colloia and Interface Science, Vol. 39, No. I, April 1972

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CADLE

with air but are still very hot. Finely divided lava (tephra) may serve as a catalyst for this oxidation. Sulfur dioxide also undergoes considerable oxidation in power plant plumes and the amount oxidized may be 20 % or more in a few minutes. However, in this case, the oxidation probably occurs in solution in aqueous droplets condensing in the plumes (17). Similar oxidation must occur in fogs and clouds in the ambient atmosphere. The SO2 dissolves in water droplets to form sulfurous acid (H2S03) :

fraction of the chloride in the atmosphere is combined in hydrogen chloride. Sulfuric acid droplets and particles of limestone can react in similar fashion, carbon dioxide being evolved as a result. The theory of this type of reaction in aerosols is, of course, the theory of coagulation, which can occur by several mechanisms, such as Brownian diffusion, electrostatic interaction, and differential rates of sedimentation. The theory of these mechanisms has often been reviewed (20-22). Gas-Particle Interaction

S02 -b H 2 0 - ~ H2SO~ .

[13]

The H2SO3 is then oxidized by dissolved oxygen from the air: 2H2SO3 ~- O2 --* 2tI2SO4 .

[14]

The rate and extent of the oxidation is markedly affected by the acidity of the droplets, decreasing with increasing acidity (18). Catalysts such as iron and manganese salts dissolved in the droplets greatly increase the rate of oxidation. Johnstone and Coughanowr (19) investigated this reaction. They estimated the possible rate of acid formation in a fog nucleated with manganese salts to be about 1%/min, orders of magnitude faster than the photochemical reaction in the absence of other pollutants. Urone and his co-workers (14) found that, even in the absence of sunlight, S02 in air is oxidized very rapidly in the presence of powdered oxides of various metals such as aluminum, calcium, and iron. Airborne particles may have a similar effect, but since the work was only qualitative, the magnitude of this oxidation cannot now be estimated.

-diAl/dr

= 3/~[A]0[B]o/r0,

[16]

and

AEROSOL REACTIONS

k = PZ exp(-E/RT),

Coalescence

The atmosphere often contains particles of different composition that can react on contact. An example is the particles of sea salt formed by the evaporation of droplets of sea water, and droplets of sulfuric acid. In the latter case, coalescence may result in the liberation of hydrogen chloride: 2 N a C 1 -t- I-I2SO4 --~ Na~SO4 ~- 2 H C I ~.

Many gases in air can and do react with the aerosol particles. Examples are ammonia (NH3) with H2SO¢ droplets, nitric acid vapor (HNO@ with sea salt particles, and ozone with organic particles. The rate of reaction of a particle after it has been introduced into a gaseous mixture containing a reactive gas may be controlled by gas-phase diffusion of the reactive gas, by the rate of reaction at the surface of the particle, or by diffusion of reactants and products in the particle. The upper limit to the initial reaction rate can be calculated from the kinetic theory of gases by estimating the number of collisions per second of the gaseous reactant per unit area of particle surface and assuming that each collision results in reaction. Cadle and Robbins (23) pointed out that this is the equivalent of calculating the rate constant k from the collision theory of reaction rates, assuming a steric factor P of unity and a zero activation energy E:

[15]

There is considerable evidence that a large Journal of Colloid and Interface Science, Vcl. 39, No. 1, April 1972

[17]

where [A] is the concentration of the reactive gas, [B] is the concentration of spherical aerosol particles of radices r expressed in weight or volume units, the subscript 0 refers to initial conditions, and Z is the collision number. :\lany reactions involving gases and airborne particles maintain a steady state throughout much of their course, gashevsky (24) derived equations for steady-state diffu-

ATNIOSPHEI~IC PAI:~TICLE FOI~MATION sion and chemical reaction in living cells, which are applicable with slight modification (19, 23). The differential equations applying to such eases can be integrated for various types of steady-state reactions. When gas-phase diffusion is rapid relative to the rate of diffusion and reaction in the particle, and the rate of reaction of the dissolved gas with the material of the particle q is a first-order reaction: R~ -

ciD~

r0

[18]

• [ ( k r / D ~ ) m r o coth (k~/D~)I/2ro -- 1]

where R~ is the steady-state rate ( - d [ A ] / d t ) per unit area of particle surface, c~ is the concentration of the dissolved reactant gas within the particle, D~ is the diffusion coetficient of the reactant gas within the particle, and ]c~ is the conventional rate constant for reaction. For large values of ( l ~ / D ~ ) l / 2 r o , Eq. [18] reduces to R~ = c~(DJc~) ~/2.

[19]

If q is a zero-order reaction and the penetration of the reactant gas into the particle is small, 1,~ = (2ciDJc~) 1/2.

[20]

When gas-phase diffusion is slow relative to the absorption by the particle, and the concentration of the gaseous reactant in the air is sufficiently high that it remains nearly constant, R, -

ceD~ __ pD~ ro RT~ '

[21]

where R is the gas constant, T the absolute temperature, D~ the diffusion coefficient of the reactant gas outside the particle, and p is the partial pressure of the reactant gas. Equation [21] is valid only when there is essentially no movement of the particle relatire to the air. If P is very small or E is large (Eqs. [16] and [17]), the reaction rate may be controlled by the reaction at the surface of the particle, [A] and [B] may be nearly constant for long periods of time, and Eq. [16] will quantitatively describe this steady-state condition.

29

Such equations are useful for analyzing the data from experiments with steady-state reactions involving aerosols. For example, Johnstone and Coughanowr (19) found that the rate of absorption of sulfur dioxide by stationary droplets of dilute hydrogen peroxide could be described by Eq. [21], suggesting that the rate was controlled by gas-phase diffusion of the sulfur dioxide. They also found that the reaction of sulfur dioxide with droplets of aqueous solutions of manganese sulfate followed the steady-state rate equations for reactions controlled by liquid-phase diffusion and zero-order reaction (Eq. [20]). Very few quantitative studies have been made of chemical reactions between gases and aerosol particles that occur in the atmosphere. Since ammonium sulfate is an important constituent of atmospheric aerosols, and is probably formed in the atmosphere by the reaction between I-I2SO4 droplets and ammonia, the kinetics of the latter reaction was studied by Cadle and Robbins (23) and Robbins and Cadle (25). They found that when the droplets were concentrated H2SO4, the initial reaction rates could be represented by Eq. [16], in which the average secondorder rate constant/c is 1.5 X 10~em4 mole-1 see-I at 300°K. The overall reaction could be represented by the empirical equation d[A__]] = k 3[A]0 dt

r

[22] • ( 0 . 1 8 0 . 1 1 ~ 0 . X 2 X ) [NH3], where X is the fraction of the droplet reacted. The sulfuric acid droplets in a city smog or fog are generally quite dilute. Cadle and Robbins (23) found this reaction too fast to measure by their technique and suggested that every collision of an ammonia molecule with a droplet resulted in reaction. These investigators also studied the reactions in the system air-H20 vapor-NO2NaC1 (26). When NO2 and an aerosol of NaC1 particles were added to moist air, an equilibrium mixture containing NaNO3 and HC1 (gas) was produced so rapidly that, in most cases, the rates could not be deterJournal of Colloid and Interface Science, Vol. 39, No. 1, April 1972

30

CADLE

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mole ratio NO2/NaCI Fza. 1. P e r c e n t a g e conversion of NaC1 aerosol to NaNO3 aerosol by NO2 at different relative humidities. Journal of Colloid and Interface Science, Vot. 39,' No. l, April 1972

ATMOSPHERIC PARTICLE FORMATION mined (Fig. 1). The first step is probably the reaction 3NO2 + H20 --~ 2HNO3 + NO.

[23]

The equilibrium constant for this reaction is 0.004 a t m -1 at 300°K, but in air, where there is much more water vapor t h a n NO2, about 5 % of the NO2 is converted to nitric acid vapor. The second step is either adsorption of nitric acid v a p o r on the relatively dry sodium chloride particles or, when the relative humidity exceeded a b o u t 75 %, solution in aqueous droplets containing NaC1. This was followed b y reaction of the nitric acid with NaC1 and desorption of hydrogen chloride, either immediately following the renttion or during subsequent evaporation of the droplets. This sequence of reactions m a y be of particular importance in the air near seacoast cities, the cities supplying the NO2 and the ocean the NaC1. In fact, particles collected from air near such cities have been found to contain high concentrations of nitrate b u t little or no nitrite (27). Goetz and Pueschel (28, 29) have produced artificial photochemical smog in a flow system and have concluded t h a t particle formation occurs largely by chemical reactions at the surfaces of previously existing airborne particles, which t h e y called reaction centers. T h e y found t h a t other conditions being equal, the size, type, and concentration of the reaction centers determined the rate of the formation of the products which contribute to aerosol formation. T h e y also studied the effects of relative humidity, sulfur dioxide concentration, and order of mixing. Large rather complicated effects t h a t were difficult to explain were observed. REFERENCES 1. LEIGHTON, P. A., "Photochemistry of Ai~ Pollution." Academic Press, New York, 1961. 2. ALTSttULLER, A. P., AND BUFALINI, J. J., Environ. Sci. Technol. 5, 39 (1971). 3. WENT, F. W., Sei. Amer. 192(5), 63 (1955); Proc. Nat. Acad. Sci. U.S.A. 46,212 (1960). 4. CADLE, R. D., "Particles in the Atmosphere and Space." Reinhold, New York, 1966. 5. RASMUSSEN, R. A., AND WENT, F. W., Proc. Nat. Acad. Sci. U.S.A. 53,215 (1965). 6. RIPPERTONj L. A., WHITE, O., AND JEFFRIES,

H. E., presented: Div. Water, Air, and

31

Waste Chem., 154th Meet. Amer. Chem. Soc., Chicago, IL, 1967. 7. ROBINSON, E., AND ROBBINS, I~. C., Final report, Stanford Res. Inst. Project PR-6755, Menlo Park, CA, 1968. 8. CADLE,l~. D., AND LEDFORD, M., Int. J. Air Water Pollut. 10, 25 (1966); HALES, J. M., WILKES, J. O., AND YORK, J. L., Atmos. Environ. 3, 657 (1969). 9. LIUTI, G., DONDES, S., AND HARTECK, P., J. Amer. Chem. Soc. 88, 3212 (1966). 10. CADLE, R. D., AND POWERS, J. W., TeUus

18, 176 (1966). 11. CADLE,R. D., LAznus, A. L., POLLOCK,W. t{. ANDSHEDLOVSKY,J. P., in "Proceedings of the Symposium on Tropical Meteorology," Univ. of Hawaii, Honolulu, 1970. (C. S. Ramane, Ed.), American Meteorological Society, Boston, Mass., 1970. 12. GOODEVE,C. F., EASTMAN,A. S., ANDDOOLEY, A., Trans. Faraday Soc. 30, 1127 (1934). 13. RENZETTI, N. A., AND DOYLE, G'. Z., J. A i r Pollut. Contr. Ass. 8, 293 (1959); Int. J. Air Pollut. 2, 327 (1960). 14. URONE, P., LUTSEP, H., NOYES, C. M., AND PARCHER, J. F., Environ. Sci. Technol. 2,611

(1968). 15. HEALD,E. F., NAUGHTON,J. J., AND BARNES, I. L., JR., J. Geophy.s. Res. 68,545 (1963). 16. CAD~E,R. D., WARTB~mG,A. F., ANDGRAHEK, F. E., Geochim. Cosmochim. Aeta 35, 503 (1971). 17. FOSTER, P. M., Atmos. Environ. 3, 157 (1969). 18. JUNGE, C. E., AND RYAN, T. G., Quart. J. Roy. Meteorol. Soc. 84, 46 (1958). 19. JOHNSTONE, H. F., AND COUGHANowR, D.R., Ind. Eng. Chem. 50, 1169 (1958).

20. CADLE, R. D., "Particle Size. Theory and Industrial Applications." Reinhold, New York, 1965. 21. GREEN, H. L., AND LANE, W. R., "Particulate Clouds. Dusts, Smokes and Mists," 2nd ed., Van Nostrand, Princeton, N.J., 1964. 22. DAVIES, C. N., ed., "Aerosol Science." Academic Press, New York, 1966. 23. CADLE, ]:~. D., AND ROBBINS, R. C., Discuss. Faraday Soc. 30, 155 (1961).

24. RASItEVSI4Y,N., "Mathematical Biophysics," Univ. of Chicago Press, Chicago, 1938. 25. ROBBINS, R. C., AND CADLE, R. D., J . Phys. Chem. 62, 469 (1958). 26. ROBBINS,R. C., CADLE,R. ]~., ANDECKHARDT, D. L., J. Meteorol. 16, 53 (1959). 27. JUNGE, C. E., "Air Chemistry and Radioac-

tivity." Academic Press, New York, 1963. 28. GOETZ, A., AND PUESCHEL, R., J . A i r Pollut. Contr. Ass. 15, 90 (1965). 29. GOETZ,A., ANDPUESCttEL, R., Atmos. Environ.

1, 287 (1967). Journal of Colloid and Interface Science, Vol. 39, No. 1, April 1972