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The isotopic composition of ammonia, nitrogen dioxide and nitrate in the atmosphere
Atmosphpric Enuirontwnt Vol. 11, pp. 1239-1243. Pergamon Press 1977. Pnnted in
Great Britain.
THE ISOTOPIC COMPOSITION OF AMMONIA, NITROGEN DIOXIDE AND NITRATE IN THE ATMOSPHERE HOWARD MOORE Department of Physical Sciences, Florida International University, Miami, FL 33199, U.S.A. (First received 25 February 1977 and in jnal form 27 April 1977)
Abstract-Values of nitrogen isotope ratios in gaseous ammonia and nitrogen dioxide and in ammonium ion and nitrate ion in aerosols and precipitation are presented. These values are shown to vary with source. Ammonia in clean air samples had a Sr5N of - 10.0 2 2.60/, compared to f24.9 + 3.4%, in barnyard samples. Nitrogen dioxide in clean air is -9.3 F 3.5#& compared to auto exhaust of +3.7 * 0.3% Ammonium ion in particulates and in rain had 6”N values of 5.6 t 5.5 and -1.4 + 3.5%,, respectively. Nitrate ion in particulates and in rain had 615N values of 5.0 &- 5.7 and - 6.6 + 3.9%, respectively. Several possible isotope fractionation mechanisms which might be operative in the atmosphere are discussed. The most plausible fractionation mechanism appears to involve isotopic equilibrium between gaseous and condensed phases with subsequent removal of the condensed phase.
atmosphere. Both equilib~um The mechanism by which gases such as NH,, NO2 and SC& react with water droplets or are converted to particulate matter and are eventually scavenged from the atmosphere is a complex problem which is
not well understood. Evidence for the existence of ammonium sulfate aerosols in the atmosphere is extensive (Twomey, 1971; Heard and Wiffen, 1969; Charlson et al., 1974; ~~derpol et al., 1975). While evidence for nitrate aerosols is less extensive, compounds such as NH,NOJ undoubtedly exist and must be formed in reactions analogous to those which result in the formation of ammonium sulfate (Robinson and Robbins, 1970). Another problem associated with the atmospheric nitrogen cycle is the possible conversion of NH3 to NO,. Several authors have given argument supporting such a conversion (Hoering, 1957; Georgii, 1963; McConnell, 1973). Robinson and Robbins (1968, 1970) argued both for and against such a conversion and finally concluded that no satisfactory reactions are known for converting NH3 to NO; in the atmosphere. Determinations of D/H ratios (Dansgaard, 1964; Friedman et al., 1964; Knight et al., 1975), i3C;“C ratios (Stevens, 1972) and 34S/3zS ratios (Jensen and Nakai, 1961; Grey and Jensen, 1972; Holt et al., 1972; Dequasie and Grey, 1970) have been helpful in assessing the importance of sources of reactive atmospheric gases and of various physical and chemical processes occurring in the atmosphere. In comparison, few determinations of 15N/r4N isotope ratios in atmospheric compounds have been made (Hoering, 1957; Moore, 1974; Wada et ai., 1975). The isotope ratio in a particular chemical species in the atmosphere is due to the source or combination of sources from which it originates and is altered by chemical and physical processes which occur in the
and unidirectional
reactions such as
reactions such as, i4NH3:
14N02
(2a)
“NH,-II:
iSNO,
(2b)
may alter isotope ratios. It is the purpose of this paper to present and discuss data on the 15N/i4N isotope ratios in atmospheric nitrogen containing compounds. The i5N/i4N isotope ratios are discussed in terms of possible isotope fractionation mechanisms such as those above. EXPERIMENTS ‘PROCEDURES AND
RESULTS
The experimental methods utilized have been described previously with the exception of the NH, collection procedures and NO, collections from auto exhaust (Moore, 1974). The nitrogen species in air were collected by passing air through a series of five filters. The first filter is a very efficient polystyrene filter (Delbag-Luftfilter, GMBH, No. 99/98) which removes particulate NH: and NO;. The next two filters are Whatman 541 cellulose filters which are chemj~lIy impregnate with triethanoia~ne and act as efficient colkctors of NOz (Moore, 1974). Similar filters have been used for the collection of SO2 in the atmosphere by Forrest and Newman (1973). The last two filters are Whatman 541 cellulose filters impregnated with phosphoric acid. Various other acids and salts were tried for
NH3 collection but were unsatisfactory. Analysis of dual filters in series indicates that the collection efficiency of the phosphoric acid impregnated filter is approximately 95% at air flow rates of 0.85 m” min-’ through a 25 cm square filter and at NH, concentrati~s existing in the un~lluted atmosphere. Each of the filters is placed in a modified Kieldaht tvne distilling apparatus; concentrated base is added to each filter and the NH, produced is distilled into a standard acid solution. Zinc which has a fresh coating of finely divided copper is then added to reduce the NO; to NHs
1239
HOWARD
1240
MOORE The precision for individual sample measurements is estimated to be +0.2%,. Errors due to instability of the mass spectrometer are smaller, but the conversion of the various compounds to N, introduces an additional small error. The overall range of values for the various nitrogen species in the atmosphere is undoubtedly due to various atmospheric processes discussed below. The concentration of gaseous ammonia collected at the laboratory ranged from 0.22 to 0.56 ppbv with a mean of 0.38 ppbv. The concentrations in the barnyard samples ranged from 70.2 to 82.0 ppbv with a mean of 77.7 ppbv. The greatest source of error in determining concentrations is the error in determining air sample volumes (+lO%). This is the estimated error in the values given above. The concentrations in samples collected at the laboratory are lower than concentrations cited by Robinson and Robbins (1970) but are in the range of values found by Georgii and Muller (1974) for the most ammonia-free regions in the atmosphere. Typical concentrations for the other nitrogen species were reported earlier (Moore, 1974).
and a second distillation is performed. The distillates are reacted with NaOBr in a vacuum system in order to produce N, suitable for mass spectrometric analysis (Hoering, 1955; Hoering and Moore, 1958). The amount of N, produced is measured manometrically in order to determine the original concentrations of NH: and NO.7 in the filters. Ammonium and nitrate ions in rainwater are separated bv a orocedure similar to that described bv Hoerina (1957).
Atmospheric N, is used as the standard (Moore, 1974; Wada et al., 1975). Most of the samples were collected on the NCAR laboratory rooftop or near the laboratory in Boulder, CO. The collection site is characterized by predominantly westerly winds and the immediate upwind area is dominated by the forested Rocky Mountain region. One sample was collected in a local greenhouse and the nitrogen was found to have a similar isotopic composition to’ the above samples. The isotopic ratios for all of these samples are shown in Fig. 1. Three samples were collected in a barnyard in which sheep are kept. The isotopic composition of the NH, nitrogen samples was 21.5, 25.8 and 27.55,,, compared to a mean of - lO.O%,,for the samples collected near the laboratory. The concentration of NH, was about 200 times higher in the barnyard samples. Whether this is all NH3 or perhaps some urea and other amines is not certain, as the analysis procedure will not distinguish between these compounds. The amounts and the isotopic composition of the other nitrogen species were similar to those in the air samples collected near the laboratory. Three samples of auto exhaust were taken from the same automobile. The isotopic ratios for the N, in the NO, were 3.9, 3.4 and 3.9&. No difference is apparent between samples at idle or at a higher engine speed.
DISCUSSION The 6” values for NH, in the barnyard and in the clean air near the laboratory demonstrate the possible differences due to variation in sources. Smaller differences are shown by the 6’ 5N values for NOz in auto exhaust compared to the clean air samples. A consideration of the sampling location and the similarity of 6i5N values for the greenhouse samples and the clean air samples lead to the conclusion that the 615N values shown to the left side in Fig. 1 are primarily due to natural soil emanations. The similarity of the 61sN values for NH: and NO; collected on filters to the 615N values for similar species in soil (Cheng et al., 1964) suggested that these species might be due to the entrainment of soil particles (Moore, 1974). In one sample collected near the laboratory, the aerosol concentration was 14.4 pg me3 air. Ammonium ion accounted for 4.3% and nitrate ion 1.2% of the total concentration. If
6 15N
-15
-10
I
I
NH:-
-5
Solid
.
NH,+- Rain
..
NHs -Gas ..w
.
.“g
IO
15
20
25
I
I
I
I
I
A
.
”
“..
.
.
.
Solid
.
. .. I
NO;- Rain . . .“.I - Gas . . .. . A
.
A
. .
.. .
A-
l
.
.
A
NO, .
.
5
. .
A
NO;-
0
I
.
I
Figure 1. rsN/N14N ratios for various atmospheric nitrogen containing compounds. A represents the mean value. The NH,-gas values at the extreme right are for samples collected in a barnyard.
L241
Isotopic compositions in the atmosphere
these are assumed to be in the form of ammonium nitrate and ammonium sulfate, 16% of the total aerosol concentration is accounted for. Comparing the percentage of ammonium ion in the filter to that in soils of 5 x lo-“/, (Richardson, 1938) shows that bulk soil entrainm~t into the atmosphere cannot account for the ammonium ion in aerosol particles. Postulation of a surface soil particle enriched in ammonium and nitrate is invalid due to the high solubility of these compounds and the extreme enrichment factor necessary. These facts point to gaseous precursors in the atmosphere. Ammonia does not appear to be produced by reactions occurring in the atmosphere. Whether it is destroyed by oxidation to nitrogen oxides is not certain. McConnell (1973) has considered the reaction of NH3 with OH radicals in the atmosphere. The product of this reaction is NH, and it is not certain whether NH2 reacts further to give NO or, in fact, consumes NO to produce N> Based on an aerosol residue time of 30 days, McConnell calculates that 4lO/, of the NH, is converted to NO, rather than being removed by rainout and dry deposition. If an aerosol residence time of 5 days is assumed (Martell and Moore, 19743, the conversion would be 10%. Even this percentage may be high in view of the uncertainties which exist in the rate equations. Miyaka and Wada (1971) studied the bacteria1 oxidation of NH3 using a marine nitrifier (N~~~o~~ys~~s oceanus) and found the NO; depleted in “N by 5.4 to 21.1?& While it is possible that dit%rent bacterial nitrifiers might produce different fractionation effects, the similarity between 615N for the NH3-NO2 and NH:-NO; couples [Fig. 1) suaests that bacterial oxidation of NH3 in the atmosphere is unimportant, Other oxidation mechanisms may be possible, but the evidence thus far appears to support the conclusion of Robinson and Robbins (1970) that the ammonia and nitrate cycles are not coupled in the troposphere. This conclusion requires a rather Iarge natural flux af nitrogen oxide and/or nitrogen dioxide (NO,) into the atmosphere to balance the measured nitrate deposition by rainfall. It has been known for some time that freshly filled silos produce NO, and persons who enter such silos are subject to “silo filler’s disease” (Lowry and Schuman, 1956). In one case, brown fumes (NO,) were so dense that color photo~aphs were obtained (Peterson et al,, 1958). A sample of the gas was found to be approximately 10% nitrogen dioxide., Other analyses have shown varying ratios of NO and NOz. Experiments by Peterson et al. (1958) indicate that bacteria are responsible for the formation of nitrogen oxides including NO, NO* and NzO from nitrate. Thus it is reasonable that these oxides are produced in soils by similar mechanisms and can escape to the atmosphere. If large amounts of NO, were produced in tbunderstorms, a correlation should be found between thunderstorm activity and the nitrate content of rain. Georgii (1963) found no correlation, while Reiter
(1970) did find four cases of enrichment of NO2 at a mountain station.‘Atmospheric nitrogen fixation by lightning has been reported by Noxon (1976) and fixation by high-voltage discharges which simulate lightning in the laboratory has been reported by Zipf and Dubin (1976) Recently, GrifIing (1977) developed a model which predicts the amount of nitrogen oxides produced by a lightning flash in a model thunderstorm, While clearly fortuitous, the amount predicted by Griffing agrees with the amount found by Noxon (1976). Ingerson (1953) suggested that nitrates fixed by lightning should be depleted in i5N with respect to atmospheric nitrogen if the N 2 and NO are in equilibrium at a low temperature. At 25°C for the reaction 14N2 + 2’5NOr-?:‘SN,
+ 2”NO,
(4)
the equilibrium isotope fractionation factor, CI,equals 1.015. This suggests that nitrogen in NO would be about - 157&, with respect to atmospheric Nz (a = 1 + ziis/loOa). If NO is rapidly oxidized to NOz with little or no fractionation then the NOz nitrogen should also be about -15x0. The data in Fig. 1 appear to support Ingerson’s suggestion. Hoering (1957) attempted to test this idea by passing air through a spark between electrodes of tungsten and platinum but could find no isotope fractionation. He therefore concluded that this is not a feasible production mechanism. Thus, it appears certain that NO, is produced, but the relative importance of this source is not certain. If it is assumed that the 6r5N variations for the ammonia species shown in Fig. 1 are representative of a single source in the atmosphere and not due to oxidation of ammonia to nitrogen oxides, the variations can be explained on the basis of heterogeneous gas-liquid or gas-solid reactions. Healy et al. (X970) reviewed the evidence for possible mechanisms of ~rnoni~ sulfate formation and concluded that the most probabIe mechanism is by reaction of NHs with SOz and an oxidant such as O2 in mist or cloud droplets. They also concluded that photochemical reactions, direct reactions of NH% and SOz and reactions on particles are unimportant. Junge and Ryan (1958) have shown that the oxidation of SO1 in water proceeds fairly rapidly providing that the sulfuric acid formed is neutralized so that the pH remains greater than 2.5. The sequence of events leading to the formation of ammonium sulfate particles in the atmosphere appear8 to involve the reaction of SO2 with Hz0 to form H&SO,, which is then oxidized to H,SO,. The H$G4 is then partially neutralized by NH, to form ammonium sulfate. The nitrogen isotopes might be fractionated in several ways in the above sequence. If the fractionation was due to diffusion of NH3 to the mist droplet the “*‘NH, would be enriched in the resulting rain and aerosol samples relative to 15NH3. The opposite iS in fact true. The rates of diffusion are much faster
HOWARDMOORE
1242
than the rate of droplet removal (Cadle and Robbihs, 1960) hence eq~iib~um must be established and the diffusion process must be relatively unim~r~t. If the ammonia system represents a closed box equilibrium system, the following equilibria must be satisfied: I 4NH3 (sol)f ’ +@b (gas) *
F
15NH~wj + 14NHzw 14NH,+w, + “NH,
(3
=
&[H’l
‘JN‘H~,s,, + 14NHs(sasj (6) NH3 + H,O=:NH:OH-
(7)
H+ + OH-.
(8)
For reactions (5) and (6), values of c( the equilibrium isotope fractionation factors are 1.005 and 1.036 at 25°C respectively, .Fc, for reactions (7) and K, for reaction (8) are 1.77 x lo-’ and 1 x 10ei4, respectively. Kirshenbaum et al. (1947) have shown that for this equilibrium system a0 = M(Q - c(5)+ ag
(9)
where tie is the measured fractionation factor and A4 is the ratio of ammonia to ammonia plus ammonium ion in solution. In addition, tIo = 1 + &il~ 1 + S,/lOcMl’
(10)
where S, and 6, represent the 6r5N values for the condensed and vapor phases, respectively. The values of 6, and 6, are shown in Fig. 1. Using these values, A4 can be determined from equation (9). The pH of the rain is then given by
b
h-1 . i-
I
(11)
Solving these equations for the pH of the rain yields a value of 10. Since precipitation usually has pH value between 4 and 6 (Junge, 1963), it appears that the ammonia system cannot be a closed-box equilibrium system. If it is assumed that the a~onia system is an open system in which the ammonia di5N is the ratio in the system prior to any removal mechanism and the rain or aerosol 6rSN value represents the ratio in the fraction removed then (Knight et al., 1975) A0 = A, + A, GoA0 = &A, + 6,A,
(12)
’
(16)
&-+l-[l* cc*(c!- 1)’
(17)
and L is the liquid water content in the atmosphere. For the gas-rain system F [equation (1611 is 0.76 and the product of [H’].L is 1.97 x 10e6, where L is in g rnT3. This value corresponds to liquid water contents of 0.2, 1.97 and 19.7 gme3 at pH values of 5, 6, and 7, respectively. For the gas-aerosol system F is 0.57 and corresponding liquid water contents are 0.01, 0.78 and 7.8 gmm3 at the above pH values. These values of L are reasonable and support the assumption that the fractionation of the nitrogen isotopes in the ammonia system is due to heterogeneous reactions with removal of the condensed phase. The nitrogen oxide system is complicated by the fact that several different oxides exist and that NO2 reacts with water to form HN03 and NO upon solution. However, if equation (1) is assumed to represent the dominant reaction (a = 1.036), the fraction of NO* in the condensed phase is 0.91 for the gas-rain system and 0.60 for the gassaerosol system. These values show the same trend as the ammonia system values as might be expected from the data in Fig.
REMARKS
The nitrogen isotope ratio in a particular nitrogen compound found in the atmosphere depends in part upon its source. This source dependent variation should be fully explored, as it will be useful in ascertaining the importance of various natural and anthropogenic sources within the atmosphere. The nitrogen isotope ratios are altered by reactions occurring in the atmosphere. In particular, heterogeneous reactions in which the liquid or solid phase is removed from the atmosphere appear to be operative. Further studies of nitrogen isotope fractionation factors for well-characterized atmospheric reactions are needed to aid an understanding of the entire atmospheric nitrogen cycle.
(13) (14)
0
L
NH,,,, + HA+,, = NH4OH~,~
CONCLUDING
pH=-log;?
.
K;K
where I< is the equilibrium constant for the reaction
(gas) e
H20=:
in each of these fractions. The value of c1 is 1.036 [equation (6)] as the eq~librium between ammonia gas and a~o~ia in solution is not imprint at less than pH 7. The fraction of ammonia in the condensed phase can also be shown to be given by (Junge, 1963; Junge and Ryan, 1958)
Acknowle~e~e~~-This work was performed at the National Center for Atmospheric Research which is sponsored by the National Science Foundation.
(15)
where &, A, and A, represent the total ammonia, and the amount in the condensed phase and vapor phase, respectively. 6,,, S, and 6, refer to the 615N
REFERENCES Cadle R. D. and Robbins R. C. (1960) Kinetics of atmospheric chemical reactions involving aerosols. Discuss. ~u~~~~ Sot. 30, 155-161.
Isotopic compositions in the atmosphere Charlson R. J., Vanderpol A. H., Covert D. S., Waggoner A. P. and Ahlquist N. C. (1973) Sulfuric aoid-ammonium sulfate aerosol: optical detection in the St. Louis region. Science X&Q,156-l 58. Cheng H. II., Bremner J. M. and Edwards A. P. (1964) Variations of rtitrugen-15 abundance in soils Science I& f574-1575. Dansgaard W. f1964j Stabfe isotopes in precipitation. T&us X6, 4X-468. Dequasie H. L. and Grey D. C. (1970) Stable isotopes applied to pollution studies. Am. Lab. 2 (12) 19-27. Forrest .I. and Newman L. (1973) Sampling and analysis of atmospheric sulfur compounds for isotope ratio studies. A~mo$@iericEnvironment 7, I-13. Friedman I,, Redfield A. C., Schoen B. and Harris .I. (1964) The variation of the deuterium con&tentof naturati waters in tke kydrolo~G cycle. Rea ~~o~~ys. 2, $77~224~ Georgii II. W. (1963) Oxides of nitrogen and ammonia in the atmospkere. d. geophys. Res. 6&, 3963-3970. Georgii II. W, and Muller W. _I.(1974) On the distribution of ammonia in the middle and lower troposphere. Tellus 26, 180-I 84. Grey D. C. aud Jensen .I. L. (1972) Bacteriogenic sulfur in air pollution. Science 177, 1099-l 100. Griffing G. W. (1977) Ozone and oxides of nitrogen production during tkunderstorms. 3. ~~~~~~1s.J&X g& 943-9X Healy T. V., McKay H. A. C1 FiIbeam A. and SarugiII D. {I970) Ammonia and ammonium sulfate in tke troposphere over the United Kingdom. J. aaophys. Res. 75, 2317-2321. Heard M. J. and Wiffen R. D. (1969) Electron microscopy of natural aerosols and the identification of particulate ammonium sulphate. Atmospheric Enziironment 3, 337-340. Hoering T. C. (lass) Variations of nitrogen-15 abundance in natura~Iy occurring substances. Science 132, 1233-1234. Hoering T. C. ff957) The isotopic composition of the ammonia and the nitrate ion in rain. C%chim. COSBWckim. Acta 12, 97-102. Hoering T. C. and Moore H. E. (19%) The isotopic composition of the nitrogen in natural gases and associated crude oils. Genchim. cosmochim. Acta 13, 2X-232. Holt B. D., ~~~~l~erneir A. G. and Venters A. (1972)Variations of sulfnr isotope ratios in samples of water and air near C&ago. &tinm. S&. TechnuE, 6, 338-341. Ingerson E. (1953) Nonrad~og~~~ isotopes in geoIogy_ I&& @Oi. sot. Am 64, 301-374. Jensen M. I.,. and Nakai N. (1961) Sources and isotopic composition of atmospheric sulfur, Science 134, 2102-2104. Junge C. E. (1963)Air Ckemistry and Radioactivity, p. 338. Academic Press, New York. Junge C. E. and Ryan T. G. (1958) Study of the SO, oxi-
1243
dation in solution and its role 3inatmospheric chemistry. Q< Jl R. met. Sot. 84, 46-55. Kicshenbaum I., Smith J. S., Crowell T., Graff J. and McKee R. (1947) Separation of the nitrogen isotopes by the exchange reaction between ammonia and solutions uf ~~~~~i~~ nitrate. d. ckem. P&s. 35, 44%446. Knight C, A., Ekkalt D. H., Roper N. and ISmgkt N. C, (1975) Radial and tangential variation of deuterlum in hail stones. J, Atmos. SC!. 32, 199@2000. Lowry T. and Schuman I.. M. (1956) Silo-filler’s disease-a syndrome caused by nitrogen diaxide. J. Am. med. Ass. X62, 153-160. Martell E. A. and Moore H. E. (1974) Tropospheric aerosol residence times: a critical review. 1. Reck. Atmos. 8, 903-910. McConnell .I. C. fl973j Atmo~he~~ ammonia: J_ geophgis. Res. 78,7g12-7g21. Miyake Y. and Wada E. (1971) The isoiope effect on the nitrogen in biochemical oxidation-reduction reactions, Rec. oceanogr. Wks Japan 11, 1-6. Moor@ H. E. (1974) Isotopic measurement of atmospheric nitrogen compounds. Tklllns 26, X69-174. _ Noxon J. F. (1976) Atmosnheric nitroaen fixation bv, Ii&tning. Geophys. ies. Let;. 3, 46346% Peterson W. I-I., Burris R. I-I., Sam R. and Little II. N, {I%!) Prducti5n of toxic gas (nitrogen oxides) in silage m&q. Agrr Fd C&m. 6% K&-l26 Reiter E. fl970) 0n tke casual relation between nitrogen oxygen compounds in the troppspkere and tke atmospheric electricity. TeNus 22$ 122--l 36. Richardson H. L. (1938) The nitrogen cycle in grassland soils with especial reference to the Rothamsted Park grass experiment. J. Agr. SeE, 28, 73-121. Robinson E. and Robbins R. C. (1968) Sources, abundances and fate of atmospheric pdlufants. Am petrol ”
Inst. Prr& NG. 4035. K3binson E. and Robbms R, C. /I970) Gaseous nittogen
compound po%ttants from urban and natural sources. Air, PoItut. Control Ass. J. 20, 303-306. Stevens C. M. (1972) The Isotopic Camposition of Atmospheric Carbon Monoxide. Coordinating Research Council, New York. Twomey S. (1971) The composition OF cloud nuclei. J. Atmos. Sci. 28, 277-281.
Vaaderpol A, H., Carsey F. D., Covert D. S., Ckarlson R. f. and Waggoner A. I?, (197s) Aerosol chemical parameters and air mass character in the St. Louis region.
Scicnrce1%&57&571. Wada E., Kadonaya T. and Matsuo S. (1975) lsN abtmdance in nitrogen of naturally occurring substances and global assessment of denitriflcation from isotopic viewpoint. Ceochem. J. 9, 139-148. ZipE E. C. and Dubin M. (1976) Laboratory studies of the synthesis of N,O and NG, and on the destruction of CF,Cl,? by 1igktning. EOS cS7, 156.
Great Britain.
THE ISOTOPIC COMPOSITION OF AMMONIA, NITROGEN DIOXIDE AND NITRATE IN THE ATMOSPHERE HOWARD MOORE Department of Physical Sciences, Florida International University, Miami, FL 33199, U.S.A. (First received 25 February 1977 and in jnal form 27 April 1977)
Abstract-Values of nitrogen isotope ratios in gaseous ammonia and nitrogen dioxide and in ammonium ion and nitrate ion in aerosols and precipitation are presented. These values are shown to vary with source. Ammonia in clean air samples had a Sr5N of - 10.0 2 2.60/, compared to f24.9 + 3.4%, in barnyard samples. Nitrogen dioxide in clean air is -9.3 F 3.5#& compared to auto exhaust of +3.7 * 0.3% Ammonium ion in particulates and in rain had 6”N values of 5.6 t 5.5 and -1.4 + 3.5%,, respectively. Nitrate ion in particulates and in rain had 615N values of 5.0 &- 5.7 and - 6.6 + 3.9%, respectively. Several possible isotope fractionation mechanisms which might be operative in the atmosphere are discussed. The most plausible fractionation mechanism appears to involve isotopic equilibrium between gaseous and condensed phases with subsequent removal of the condensed phase.
atmosphere. Both equilib~um The mechanism by which gases such as NH,, NO2 and SC& react with water droplets or are converted to particulate matter and are eventually scavenged from the atmosphere is a complex problem which is
not well understood. Evidence for the existence of ammonium sulfate aerosols in the atmosphere is extensive (Twomey, 1971; Heard and Wiffen, 1969; Charlson et al., 1974; ~~derpol et al., 1975). While evidence for nitrate aerosols is less extensive, compounds such as NH,NOJ undoubtedly exist and must be formed in reactions analogous to those which result in the formation of ammonium sulfate (Robinson and Robbins, 1970). Another problem associated with the atmospheric nitrogen cycle is the possible conversion of NH3 to NO,. Several authors have given argument supporting such a conversion (Hoering, 1957; Georgii, 1963; McConnell, 1973). Robinson and Robbins (1968, 1970) argued both for and against such a conversion and finally concluded that no satisfactory reactions are known for converting NH3 to NO; in the atmosphere. Determinations of D/H ratios (Dansgaard, 1964; Friedman et al., 1964; Knight et al., 1975), i3C;“C ratios (Stevens, 1972) and 34S/3zS ratios (Jensen and Nakai, 1961; Grey and Jensen, 1972; Holt et al., 1972; Dequasie and Grey, 1970) have been helpful in assessing the importance of sources of reactive atmospheric gases and of various physical and chemical processes occurring in the atmosphere. In comparison, few determinations of 15N/r4N isotope ratios in atmospheric compounds have been made (Hoering, 1957; Moore, 1974; Wada et ai., 1975). The isotope ratio in a particular chemical species in the atmosphere is due to the source or combination of sources from which it originates and is altered by chemical and physical processes which occur in the
and unidirectional
reactions such as
reactions such as, i4NH3:
14N02
(2a)
“NH,-II:
iSNO,
(2b)
may alter isotope ratios. It is the purpose of this paper to present and discuss data on the 15N/i4N isotope ratios in atmospheric nitrogen containing compounds. The i5N/i4N isotope ratios are discussed in terms of possible isotope fractionation mechanisms such as those above. EXPERIMENTS ‘PROCEDURES AND
RESULTS
The experimental methods utilized have been described previously with the exception of the NH, collection procedures and NO, collections from auto exhaust (Moore, 1974). The nitrogen species in air were collected by passing air through a series of five filters. The first filter is a very efficient polystyrene filter (Delbag-Luftfilter, GMBH, No. 99/98) which removes particulate NH: and NO;. The next two filters are Whatman 541 cellulose filters which are chemj~lIy impregnate with triethanoia~ne and act as efficient colkctors of NOz (Moore, 1974). Similar filters have been used for the collection of SO2 in the atmosphere by Forrest and Newman (1973). The last two filters are Whatman 541 cellulose filters impregnated with phosphoric acid. Various other acids and salts were tried for
NH3 collection but were unsatisfactory. Analysis of dual filters in series indicates that the collection efficiency of the phosphoric acid impregnated filter is approximately 95% at air flow rates of 0.85 m” min-’ through a 25 cm square filter and at NH, concentrati~s existing in the un~lluted atmosphere. Each of the filters is placed in a modified Kieldaht tvne distilling apparatus; concentrated base is added to each filter and the NH, produced is distilled into a standard acid solution. Zinc which has a fresh coating of finely divided copper is then added to reduce the NO; to NHs
1239
HOWARD
1240
MOORE The precision for individual sample measurements is estimated to be +0.2%,. Errors due to instability of the mass spectrometer are smaller, but the conversion of the various compounds to N, introduces an additional small error. The overall range of values for the various nitrogen species in the atmosphere is undoubtedly due to various atmospheric processes discussed below. The concentration of gaseous ammonia collected at the laboratory ranged from 0.22 to 0.56 ppbv with a mean of 0.38 ppbv. The concentrations in the barnyard samples ranged from 70.2 to 82.0 ppbv with a mean of 77.7 ppbv. The greatest source of error in determining concentrations is the error in determining air sample volumes (+lO%). This is the estimated error in the values given above. The concentrations in samples collected at the laboratory are lower than concentrations cited by Robinson and Robbins (1970) but are in the range of values found by Georgii and Muller (1974) for the most ammonia-free regions in the atmosphere. Typical concentrations for the other nitrogen species were reported earlier (Moore, 1974).
and a second distillation is performed. The distillates are reacted with NaOBr in a vacuum system in order to produce N, suitable for mass spectrometric analysis (Hoering, 1955; Hoering and Moore, 1958). The amount of N, produced is measured manometrically in order to determine the original concentrations of NH: and NO.7 in the filters. Ammonium and nitrate ions in rainwater are separated bv a orocedure similar to that described bv Hoerina (1957).
Atmospheric N, is used as the standard (Moore, 1974; Wada et al., 1975). Most of the samples were collected on the NCAR laboratory rooftop or near the laboratory in Boulder, CO. The collection site is characterized by predominantly westerly winds and the immediate upwind area is dominated by the forested Rocky Mountain region. One sample was collected in a local greenhouse and the nitrogen was found to have a similar isotopic composition to’ the above samples. The isotopic ratios for all of these samples are shown in Fig. 1. Three samples were collected in a barnyard in which sheep are kept. The isotopic composition of the NH, nitrogen samples was 21.5, 25.8 and 27.55,,, compared to a mean of - lO.O%,,for the samples collected near the laboratory. The concentration of NH, was about 200 times higher in the barnyard samples. Whether this is all NH3 or perhaps some urea and other amines is not certain, as the analysis procedure will not distinguish between these compounds. The amounts and the isotopic composition of the other nitrogen species were similar to those in the air samples collected near the laboratory. Three samples of auto exhaust were taken from the same automobile. The isotopic ratios for the N, in the NO, were 3.9, 3.4 and 3.9&. No difference is apparent between samples at idle or at a higher engine speed.
DISCUSSION The 6” values for NH, in the barnyard and in the clean air near the laboratory demonstrate the possible differences due to variation in sources. Smaller differences are shown by the 6’ 5N values for NOz in auto exhaust compared to the clean air samples. A consideration of the sampling location and the similarity of 6i5N values for the greenhouse samples and the clean air samples lead to the conclusion that the 615N values shown to the left side in Fig. 1 are primarily due to natural soil emanations. The similarity of the 61sN values for NH: and NO; collected on filters to the 615N values for similar species in soil (Cheng et al., 1964) suggested that these species might be due to the entrainment of soil particles (Moore, 1974). In one sample collected near the laboratory, the aerosol concentration was 14.4 pg me3 air. Ammonium ion accounted for 4.3% and nitrate ion 1.2% of the total concentration. If
6 15N
-15
-10
I
I
NH:-
-5
Solid
.
NH,+- Rain
..
NHs -Gas ..w
.
.“g
IO
15
20
25
I
I
I
I
I
A
.
”
“..
.
.
.
Solid
.
. .. I
NO;- Rain . . .“.I - Gas . . .. . A
.
A
. .
.. .
A-
l
.
.
A
NO, .
.
5
. .
A
NO;-
0
I
.
I
Figure 1. rsN/N14N ratios for various atmospheric nitrogen containing compounds. A represents the mean value. The NH,-gas values at the extreme right are for samples collected in a barnyard.
L241
Isotopic compositions in the atmosphere
these are assumed to be in the form of ammonium nitrate and ammonium sulfate, 16% of the total aerosol concentration is accounted for. Comparing the percentage of ammonium ion in the filter to that in soils of 5 x lo-“/, (Richardson, 1938) shows that bulk soil entrainm~t into the atmosphere cannot account for the ammonium ion in aerosol particles. Postulation of a surface soil particle enriched in ammonium and nitrate is invalid due to the high solubility of these compounds and the extreme enrichment factor necessary. These facts point to gaseous precursors in the atmosphere. Ammonia does not appear to be produced by reactions occurring in the atmosphere. Whether it is destroyed by oxidation to nitrogen oxides is not certain. McConnell (1973) has considered the reaction of NH3 with OH radicals in the atmosphere. The product of this reaction is NH, and it is not certain whether NH2 reacts further to give NO or, in fact, consumes NO to produce N> Based on an aerosol residue time of 30 days, McConnell calculates that 4lO/, of the NH, is converted to NO, rather than being removed by rainout and dry deposition. If an aerosol residence time of 5 days is assumed (Martell and Moore, 19743, the conversion would be 10%. Even this percentage may be high in view of the uncertainties which exist in the rate equations. Miyaka and Wada (1971) studied the bacteria1 oxidation of NH3 using a marine nitrifier (N~~~o~~ys~~s oceanus) and found the NO; depleted in “N by 5.4 to 21.1?& While it is possible that dit%rent bacterial nitrifiers might produce different fractionation effects, the similarity between 615N for the NH3-NO2 and NH:-NO; couples [Fig. 1) suaests that bacterial oxidation of NH3 in the atmosphere is unimportant, Other oxidation mechanisms may be possible, but the evidence thus far appears to support the conclusion of Robinson and Robbins (1970) that the ammonia and nitrate cycles are not coupled in the troposphere. This conclusion requires a rather Iarge natural flux af nitrogen oxide and/or nitrogen dioxide (NO,) into the atmosphere to balance the measured nitrate deposition by rainfall. It has been known for some time that freshly filled silos produce NO, and persons who enter such silos are subject to “silo filler’s disease” (Lowry and Schuman, 1956). In one case, brown fumes (NO,) were so dense that color photo~aphs were obtained (Peterson et al,, 1958). A sample of the gas was found to be approximately 10% nitrogen dioxide., Other analyses have shown varying ratios of NO and NOz. Experiments by Peterson et al. (1958) indicate that bacteria are responsible for the formation of nitrogen oxides including NO, NO* and NzO from nitrate. Thus it is reasonable that these oxides are produced in soils by similar mechanisms and can escape to the atmosphere. If large amounts of NO, were produced in tbunderstorms, a correlation should be found between thunderstorm activity and the nitrate content of rain. Georgii (1963) found no correlation, while Reiter
(1970) did find four cases of enrichment of NO2 at a mountain station.‘Atmospheric nitrogen fixation by lightning has been reported by Noxon (1976) and fixation by high-voltage discharges which simulate lightning in the laboratory has been reported by Zipf and Dubin (1976) Recently, GrifIing (1977) developed a model which predicts the amount of nitrogen oxides produced by a lightning flash in a model thunderstorm, While clearly fortuitous, the amount predicted by Griffing agrees with the amount found by Noxon (1976). Ingerson (1953) suggested that nitrates fixed by lightning should be depleted in i5N with respect to atmospheric nitrogen if the N 2 and NO are in equilibrium at a low temperature. At 25°C for the reaction 14N2 + 2’5NOr-?:‘SN,
+ 2”NO,
(4)
the equilibrium isotope fractionation factor, CI,equals 1.015. This suggests that nitrogen in NO would be about - 157&, with respect to atmospheric Nz (a = 1 + ziis/loOa). If NO is rapidly oxidized to NOz with little or no fractionation then the NOz nitrogen should also be about -15x0. The data in Fig. 1 appear to support Ingerson’s suggestion. Hoering (1957) attempted to test this idea by passing air through a spark between electrodes of tungsten and platinum but could find no isotope fractionation. He therefore concluded that this is not a feasible production mechanism. Thus, it appears certain that NO, is produced, but the relative importance of this source is not certain. If it is assumed that the 6r5N variations for the ammonia species shown in Fig. 1 are representative of a single source in the atmosphere and not due to oxidation of ammonia to nitrogen oxides, the variations can be explained on the basis of heterogeneous gas-liquid or gas-solid reactions. Healy et al. (X970) reviewed the evidence for possible mechanisms of ~rnoni~ sulfate formation and concluded that the most probabIe mechanism is by reaction of NHs with SOz and an oxidant such as O2 in mist or cloud droplets. They also concluded that photochemical reactions, direct reactions of NH% and SOz and reactions on particles are unimportant. Junge and Ryan (1958) have shown that the oxidation of SO1 in water proceeds fairly rapidly providing that the sulfuric acid formed is neutralized so that the pH remains greater than 2.5. The sequence of events leading to the formation of ammonium sulfate particles in the atmosphere appear8 to involve the reaction of SO2 with Hz0 to form H&SO,, which is then oxidized to H,SO,. The H$G4 is then partially neutralized by NH, to form ammonium sulfate. The nitrogen isotopes might be fractionated in several ways in the above sequence. If the fractionation was due to diffusion of NH3 to the mist droplet the “*‘NH, would be enriched in the resulting rain and aerosol samples relative to 15NH3. The opposite iS in fact true. The rates of diffusion are much faster
HOWARDMOORE
1242
than the rate of droplet removal (Cadle and Robbihs, 1960) hence eq~iib~um must be established and the diffusion process must be relatively unim~r~t. If the ammonia system represents a closed box equilibrium system, the following equilibria must be satisfied: I 4NH3 (sol)f ’ +@b (gas) *
F
15NH~wj + 14NHzw 14NH,+w, + “NH,
(3
=
&[H’l
‘JN‘H~,s,, + 14NHs(sasj (6) NH3 + H,O=:NH:OH-
(7)
H+ + OH-.
(8)
For reactions (5) and (6), values of c( the equilibrium isotope fractionation factors are 1.005 and 1.036 at 25°C respectively, .Fc, for reactions (7) and K, for reaction (8) are 1.77 x lo-’ and 1 x 10ei4, respectively. Kirshenbaum et al. (1947) have shown that for this equilibrium system a0 = M(Q - c(5)+ ag
(9)
where tie is the measured fractionation factor and A4 is the ratio of ammonia to ammonia plus ammonium ion in solution. In addition, tIo = 1 + &il~ 1 + S,/lOcMl’
(10)
where S, and 6, represent the 6r5N values for the condensed and vapor phases, respectively. The values of 6, and 6, are shown in Fig. 1. Using these values, A4 can be determined from equation (9). The pH of the rain is then given by
b
h-1 . i-
I
(11)
Solving these equations for the pH of the rain yields a value of 10. Since precipitation usually has pH value between 4 and 6 (Junge, 1963), it appears that the ammonia system cannot be a closed-box equilibrium system. If it is assumed that the a~onia system is an open system in which the ammonia di5N is the ratio in the system prior to any removal mechanism and the rain or aerosol 6rSN value represents the ratio in the fraction removed then (Knight et al., 1975) A0 = A, + A, GoA0 = &A, + 6,A,
(12)
’
(16)
&-+l-[l* cc*(c!- 1)’
(17)
and L is the liquid water content in the atmosphere. For the gas-rain system F [equation (1611 is 0.76 and the product of [H’].L is 1.97 x 10e6, where L is in g rnT3. This value corresponds to liquid water contents of 0.2, 1.97 and 19.7 gme3 at pH values of 5, 6, and 7, respectively. For the gas-aerosol system F is 0.57 and corresponding liquid water contents are 0.01, 0.78 and 7.8 gmm3 at the above pH values. These values of L are reasonable and support the assumption that the fractionation of the nitrogen isotopes in the ammonia system is due to heterogeneous reactions with removal of the condensed phase. The nitrogen oxide system is complicated by the fact that several different oxides exist and that NO2 reacts with water to form HN03 and NO upon solution. However, if equation (1) is assumed to represent the dominant reaction (a = 1.036), the fraction of NO* in the condensed phase is 0.91 for the gas-rain system and 0.60 for the gassaerosol system. These values show the same trend as the ammonia system values as might be expected from the data in Fig.
REMARKS
The nitrogen isotope ratio in a particular nitrogen compound found in the atmosphere depends in part upon its source. This source dependent variation should be fully explored, as it will be useful in ascertaining the importance of various natural and anthropogenic sources within the atmosphere. The nitrogen isotope ratios are altered by reactions occurring in the atmosphere. In particular, heterogeneous reactions in which the liquid or solid phase is removed from the atmosphere appear to be operative. Further studies of nitrogen isotope fractionation factors for well-characterized atmospheric reactions are needed to aid an understanding of the entire atmospheric nitrogen cycle.
(13) (14)
0
L
NH,,,, + HA+,, = NH4OH~,~
CONCLUDING
pH=-log;?
.
K;K
where I< is the equilibrium constant for the reaction
(gas) e
H20=:
in each of these fractions. The value of c1 is 1.036 [equation (6)] as the eq~librium between ammonia gas and a~o~ia in solution is not imprint at less than pH 7. The fraction of ammonia in the condensed phase can also be shown to be given by (Junge, 1963; Junge and Ryan, 1958)
Acknowle~e~e~~-This work was performed at the National Center for Atmospheric Research which is sponsored by the National Science Foundation.
(15)
where &, A, and A, represent the total ammonia, and the amount in the condensed phase and vapor phase, respectively. 6,,, S, and 6, refer to the 615N
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1243
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