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SUMMARY REPORT ON ATMOSPHERIC NITRATES
U. S. ENVIRONMENTAL PROTECTION AGENCY
NATIONAL ENVIRONMENTAL RESEARCH CENTER
RESEARCH TRIANGLE PARK, NORTH CAROLINA
31, 1974
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TABLE OF CONTENTS
Page Number
PREFACE
1. SUMMARY, CONCLUSIONS, AND RECOMMENDATIONS 1
1.1 SUMMARY 1
1.2 CONCLUSIONS- — — - 12
1.3 RECOMMENDATIONS 16
2. INTRODUCTION 19
2.1 GENERAL 19
2.2 REFERENCES .— 20
3. CHEMICAL AND PHYSICAL PROPERTIES 21
3.1 ATMOSPHERIC NITRATES 21
3.1.1 Peroxy Acetyl Nitrate —-21
3.1.2 Nitrogen Pentoxide 23
3.1.3 Nitric Acid-— 24
31 A. Mi trnuc Ar irl — ___— —-—______ —J.*M.^^—»—_—«^_«—-».-»fc-»—^ ?R
• I • H n 1 Ui UUo rtw iu**~- ""• "^ •»——••—— -—»_ h**.*™!*™-**^- ^^.M..*.^ ~—~- *. _^
3.1.5 Nitrate Salts 25
3.1.6 Alkyl Nitrates • 25
3.2 REFERENCES - 26
4. MEASUREMENT TECHNIQUES 27
4.1 ENVIRONMENTAL 27
4.1.1 Sampling, Preparation, and Analysis 27
4.1.1.1 Inorganic Nitrates 27
4.1.1.1.1 Sampling and Sample Preparation -29
4Jj!2' Organic Nitrates « 32
4.1.2.1 Participate Nitrate 33
4.1.1.2.2 Alkyl Nitrates 33
4.1.1.2.3 Miscellaneous Nitrates '• 33
4.1.1.3 Conclusions 40
4.2 SOURCE MEASUREMENTS 41
4.2.1 Stationary 41
4.3 REFERENCES —— — 43
5. ENVIRONMENTAL APPRAISAL 47
5.1 ORIGIN AND ABUNDANCE 47
5.1.1 Natural Sources 47
5.1.1.1 Origin of Atmospheric Nitrates 47
5.1.2 Man-made Sources 50
5.1.2.1 Stationary Sources 50
5.1.2.2 Mobi 1 e Sources - 51
5.2 CONCENTRATIONS 53
5.2.1 Inorganic Nitrates in Air 53
5 2.2 Other Nitrate Concentration Measurements 51
5.3 TRANSFORMATION AND TRANSPORT MECHANISMS— 54
5.3.1 Natural Mechanisms 54
5.3.1.1 Chemical 54
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5.3.1.2 Transport and Removal 70>
5.4 REFERENCES 73
6. EFFECTS - - 75
6.1 HEALTH EFFECTS OF NITRATES AND RELATED COMPOUNDS IN THE AIR- ;|
6.1.1 Route of Entry of Nitrates into the Body '5
6.1.2 Epidemiologic Studies of Health Effects of Airborne
Nitrates 79
6.1.2.1 Airborne inorganic nitrates : 79
6.1.2.1.1 Respiratory effects 7^
6.1.2.1.2 Effects Other Than Respiratory 30
6.1.2.2 Airborne Organic Nitrates • 31
6.1.2.2.1 From Ambient Sources- 31
6.1.2.2.2 From Individual and Occupational Sources 32
6.1.3 Inhalation Toxicology of Nitrates ; 33
6.1.4 Nitrosamines: Possible Human Health Hazard • 35
6.1.4.1 Nitrates and Nitrites: Their Role in the -;
Nitrosation Process 35
6.1.4.2 Factors that Influence Nitrosation 36
6.1.4.3 Nitrosamines in Human and Animal Food Commodities-t —38
6.1.4.4 Biological Activity of N-nitrosocompounds : -39
6.1.4.5 Carcinogenesis: N-nitrosocompounds
6.1.4.6 Human Health Hazard
6.1.5 Research Needs '81
6.2 ECOLOGICAL EFFECTS — 391
6.2.1 Nitrates as Fertilizers — - 1.34
6.2.2 Nitrates in Aquatic Habitats 137
6.2.3 Nitrate Accumulation in Plants 138
6.2.4 Exposure to Airborne Nitrates— • 111
6.3 MATERIALS >—— •—315^
6.3.1 Laboratory -— -115;
6.3.2 Field 116,
6.4 VISIBILITY 118i
6.5 REFERENCES— - : 119
7. CONTROL TECHNOLOGY AND REMEDIAL ACTIONS 132,
7.1 STATIONARY SOURCES- - - 132.
7.1.1 Combustion Processes • 132
7.1.1.1 Source Categories 132:
7.1.1.2 Formation Mechanisms— J32:
7.1.1.3 Control Categories 133'
7.1.1.4 Control Limitations 134
7.1.1.5 Control Application Experience 134
7.1.2 Industrial Processes 136'
7.2 MOBTLE' SOURCES— ——~___________________ •,«
7.2.1 Control Strategies - iS
7.3 REFERENCES
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8. SUMMARY OF NERC/RTP CURRENT RESEARCH ACTIVITIES RELATING
TO NITRATES — 141
8.1 MEASUREMENT AND ANALYSIS 141
8.2 HEALTH EFFECTS STUDIES —141
8.3 FORMATION AND DECAY OF POLLUTANTS 141
8.4 CONTROL TECHNOLOGY - -144
8.4.1 Stationary Sources 144
8.4.1.1 Process Research and Development 145
8.4.1.2 Fuels Research Development— 145
8.4.1.3 Fundamental Combustion Research • 145
8.4.1.4 Combustion Flue-Gas Treatment Processes 146
8.4.1.5 Nitric Acid Manufacturing 146
8.4.1.6 Planned R&D 146
8.4.1.7 Probl em Areas • 147
8.4.2 Mobile Sources 147
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PREFACE
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This document was prepared by a Task Force convened under the
direction of Dr. John Finklea, Director, National Environmental
Research Center (NERC). The objective was to review and evaluate
the current knowledge on nitrates in the environment, as related to
possible deleterious effects upon human health and welfare, with a
view toward evaluating the need for control of emissions from man-made sources
under the provisions of existing statutes.
The following members served on the Task Force:
NERC/RTP
James R. Smith.Chairman
Robert Horton
J.H.B. Garner
F.P. Scaringelli
G. Fairchild
P.L. Hanst
J.G. French
K. Bridbord
J. Baugh
F. Jaye
C. Sawicki
OAQPS
M. Berry
B. Bauman
J. Upham
E. Tabor
T. Ripberger
J.F. Walling
T. Waddell
R. Baron
Gene Sawicki
James Stebbings
Choudari Kommineni
Dale Denney
R.A. Rhoden
M. Jones
NERC/CORVALLIS
Lawrence Raniere
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1. SUMMARY, CONCLUSIONS, AND RECOMMENDATIONS
1.1 SUMMARY
Nitrogen is a principal constituent in the earth's atmosphere.
The most common forms of nitrogen found in nature are nitrogen
gas, oxides of nitrogen, organic nitrogenous compounds, ammonia,
nitrites, and nitrates. Nitrate salts and oxides represent the final
stage in the atmospheric oxidation of gaseous oxides of nitrogen,
starting with the formation of nitric oxide. Nitrogen oxides play
a principal role in secondary aerosol formation, and hence urban
smog.
There is an abundant amount of nitric oxide (NO) and nitrogen
dioxide (N02) of natural origin in the atmosphere. In the lower
atmosphere, NO is created thermally in flames, explosions, and electric
discharges. In the upper atmosphere, NO is formed through the
photodissociation of N~ and Op, followed by combination of N and 0.
For the purpose of this report, the vertical transport of NO from the
upper atmosphere to the lower troposphere can be considered negligible.
Nitrates observed in urban areas result principally from the conversion
of nitrogen oxide emitted from man-made sources. Only a small fraction
is emitted directly as nitrates.
Nitrogen oxides emissions in the U. S. result primarily from the
combustion of fossil fuels in boilers, furnaces, and internal combustion
engines. Practically all man-made emissions from non-combustion
sources are from the manufacture and use of nitric acid. Currently, the
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stationary and mobile sources each contribute approximately 50% of
the oxides of nitrogen emissions from man-made sources.
Nitrate is formed in urban area atmospheres mainly as peroxy-
acetyl nitrate (PAN) vapor, with a lesser amount formed as inorganic
nitrates salts in the aerosols. In outlying areas, the formation of
nitric acid and nitrate salts may well exceed the formation of
PAN-type compounds. Nitrites also are formed in the atmosphere, but
the concentrations are much lower than that for nitrates. The
chemical path from NO to nitrates in the atmosphere may be extremely
complex, and is not well understood, although the photochemical
production of N02 has been studied extensively. A wide variety of
organic and inorganic nitrate compounds might be formed, including
intermediate species which may be unstable and highly reactive. A
number of these compounds have been found in laboratory experiments,
but few have been detected in the ambient atmosphere. Organic
nitrates are formed by the combination of N02 with oxygenated radicals.
Peroxy radicals yield PAN, which measurements show to be the pre-
dominant nitrate in urban atmospheres. Inorganic nitrate begins with
the reaction of NOg with ozone, and is thought to proceed through an
intermediate nitric acid species which, in the presence of other
solubles, will be neutralized to yield various inorganic nitrates.
It seems probable that PAN in the air is slowly taken up by aerosol
droplets and hydrolyzed to nitrite in solution, which is then oxidized
to nitrate. Thus, the ratio of inorganic nitrate to organic nitrate
in a polluted air mass should increase slowly as the air moves away
from the urban generation area. One might expect also that the
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chemical characterization of nitrate aerosols might vary from one
urban area to another, depending upon precursor emission and their
rate of emission.
The majority of the nitrate concentration data available in
the U. S. has been acquired by analysis of particulate matter samples
obtained using the Hi Vol sampler with glass fiber filters. Nitrates
in this incidence are defined as those nitrates that are removed from
suspended particulate matter, having a size range of 0.3 to 10ym in
diameter, collected on glass fiber filters. Such a technique provides
only a limited indication of the nature of the air sample. Fully
satisfactory methods for routine use in measuring particulate nitrate,
PAN, nitric acid, nitrous acid, nitrogen dioxide, and nitrogen pentoxide
are not yet available.
The arithmetic mean for nitrates as measured at the NASN stations
with the Hi Vol sampler for a 5-year period (1966-1970) for those urban
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sites where all data were available ranged from 0.57 to 7.57 yg/m --
the highest being in Los Angeles and the lowest in Concord, New Hampshire.
Nitrate measurements made at CHESS stations show large temporal
and spatial variations. In Chattanooga, Tennessee, observed 24-hour
average concentrations varied from below the detectable level to
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107.3 yg/m, and the monthly arithmetic mean varied from 0.3 to 24.9 yg/m .
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The overall monthly urban averages varied from 1.3 to 7.2 yg/m .
Available data are not sufficient to describe diurnal variations;
however, limited data and theoretical considerations suggest that
significant diurnal cycles may occur, similar to that of ozone, par-
ticularly in the Los Angeles area. If this were the case, the
daily maximum concentration should be significantly higher than
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the annual average—possibly orders of magnitude higher. Since
nitrates in the atmosphere are primarily secondary aerosols resulting
from chemical transformation processes, their distribution repre-
sents the conjugate of all sources in the area, and would extend
beyond the actual source boundaries. Vertical distribution and
dispersion would be directly influenced by meteorological conditions.
Data on the size distribution of nitrates in the atmosphere are
limited. Available data indicate that the nitrate aerosol particles
are normally 2 \tn or less in diameter. Particle sizes in coastal
areas are probably larger.
Atmospheric analyses have generally revealed oxidation products
equivalent to only about half of the suspected emissions. Available
data suggest that N02> N2$, HN03, HN02, PAN, and particulate nitrates
are taken up by such surfaces as leaves and soil at a greater rate
than had heretofore been recognized. These removal mechanisms, coupled
with the transformation processes would help to explain a diurnal
cycle. The other principal removal mechanism is precipitation. In
areas with little rainfall, such as Los Angeles in the summer months,
dry removal processes must play a primary role. Millions of tons
of nitrate containing aerosols reach the earth's surface each year
from these removal mechanisms. The nitrogen content in runoff is
highly variable depending upon the duration and intensity of rainfall,
land use, topography, physical characteristics, and antecedent
conditions. The rough average is 2-3 mg/1. In some cases, runoff
adds enough nitrogen to surface waters to produce algal growth.
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The residence time for nitrates in the atmosphere will vary
depending upon the predominant removal mechanisms. Actual data
on residence times are not available. Since the particle sizes
are primarily sub-micrometer, they will remain suspended for periods
on the order of days, and hence will be transported several hundred
kilometers, unless contained by terrain and meteorological
features—such as is the case with the Los Angeles Basin. Con-
sideration of removal mechanisms and the physical properties of
the aerosol suggest residence times of 2 to 10 days.
Nitrates and nitrites are found in varying concentrations
in most foods, and in water. Tobacco smoke contains higher con-
centrations of nitrogen compounds than does polluted air. The
highest amounts are in the form of total oxidizable nitrogen oxides
and methyl nitrite. A dimethylnitrosamine has been found in cigarette
smoke. Nitrosamines have been reported in a number of foods, but
at much lower concentrations than nitrate and nitrite. Nitrates
and nitrites are used in curing meats; the Food and Drug Administration
has established tolerance limits of 200 ppm for nitrites and 500 ppm
for nitrates for this purpose. The Public Health Service has set
10 mg NO, per liter as the maximum permissible level for drinking
water.
The ubiquity of nitrate in man's environment make possible his
exposure to these substances through a variety of routes of entry
into the body. Water and foods have long been recognized as the
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principal routes of entry. It is difficult to estimate the combined
level of exposure of man to nitrate-nitrite. The question of
exogenous sources versus endogenous production further complicates
the problem. Nitrate and nitrite salts are readily adsorbed through
the stomach. A rough estimate of nitrate in the average daily diet
is 200 mg. Since nitrate and nitrite salts are readily soluble in
water and serum containing systems, and are rapidly adsorbed when
ingested, it seems likely that they also would be rapidly absorbed
through the respiratory tract after deposition in the alveolar
region of the lung. The question thus arises concerning the amount
of nitrate absorbed in this manner as compared to that via ingestion.
Using a standard respiratory volume per day, and assuming exposure
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to nitrate concentrations of 3-5 yg/m for 12 hours per day, the
calculated retained inhaled nitrate would be 21.5 yg. This is sub-
stantially less than 1 percent of the body burden if water alone is
considered. It therefore appears unlikley that inhaled nitrates
would contribute substantially to the body burden of nitrate-nitrite,
and hence to the toxicity resulting from methemoglobinemia and its
consequent effects. However, if the preliminary correlations found
between nitrate and respiratory symptoms are correct, and they are
causally associated, then airborne nitrate may be an important
respiratory irritant.
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The extent to which nitrate may be harmful to human health is
dependent upon its reduction to nitrite.or the possible conversion to
nitrosamines. Nitrate itself is relatively nontoxic, being
rapidly absorbed and rapidly excreted. Nitrite when reaching the
blood stream reacts with hemoglobin to form methemoglobin
resulting in a decrease in the oxygen-carrying capacity of the blood.
Infants, less than 3 months old, are particularly susceptible
to methemoglo&i'nenria. A number of cases, resulting in death, have occurred
from nitrite formed from nitrate in water. A lesser number of cases
have occurred from ingestion of vegetables having a high
concentration of nitrites.
Epidemiological studies in Czechlosovakia and Poland implicated
atmospheric nitrate as the etiologic agent in methemoglobinemia;
however, these studies were not conclusive since the nitrate concentration
in drinking water was not considered. Studies in the U. S. have
associated airborne nitrates with aggravated asthma attacks.
These preliminary data suggest that certain nitrate compounds
found in urban atmosphere may be either direct or indirect
respiratory irritants; however, more information is needed for
confirmation.
Nitrosamines have been shown to be potent carcinogens in
animal experiments. N-nitrosamines are formed by reaction between
nitrous acid and tertiary and secondary amines. The reaction is
dependent upon the concentration of nitrite, the pH of the medium
(acidic), and the reacting amine, and cofactors such as formaldehyde,
thiocyanate, and chloral. Laboratory experiments with
animals suggests rather strongly that such reactions can take
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place -jn the mammalian stomach. Nitrosamines have been shown to
be effective carcinogens in animals by all routes of administration.
The formation of a carcinogen in one tissue may lead to tumor
formation in an entirely different tissue. A slight change in the
molecular structure of the nitroso compounds may cause a
significant change in the carcinogenic activity. Nitrosamines
are highly unstable. They have not been found, and it is unlikely
that they will be found in the atmosphere; however, the
possibility cannot be ruled out at this time. There is much that is
yet to be learned concerning the chemical characteristics of
atmospheric aerosols.
There is substantial evidence that some human cancers, particularly
those of the respiratory tract, are more frequent in urban than in
non-urban areas. There is some evidence to suggest that this
"Urban factor" may in part be due to air pollution. The polycyclic
aromatic hydrocarbons known to be present in urban air have long
been suspect as the responsible agents, but a possible role for
nitrosamines should be investigated when and if they are detected
in the atmosphere.
The most significant effect of the alteration of the nitrogen
cycle by man, through fertilization and waste disposal practices,
is the eutrophication of lakes, rivers, and esturaries, and the
contamination of drinking water supplies. Eutrophication may
have beneficial results by increasing the Productivity 1n the receiving
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bodies of water. However, 1n many areas, the excess of nitrates
{or other nutrients) results 1n excessive growth of algae and
other aquatic plants which reduce the oxygen supply in the deeper
water. This condition may significantly affect marine life in
the water systems. Nitrate in rainfall may contribute approximately
10 percent of the nitrate loading in surfaces waters.
High levels of inorganic nitrate found in-, ambient air do not have adverse
effects on plant life; however, plants will accumulate high
concentrations of nitrate under certain conditions. This has
resulted in cases of acute nitrate poisoning of domestic
livestock. Extensive damage to crops and ornamental plants
has resulted from exposure to PAN.
Nitrates contribute to the initial corrosion of metals, but are
not considered a major problem. High nitrate levels in Los Angeles
have resulted in failure of electronic components through stress corrosion
cracking.
Information concerning primary emissions of nitrates from stationary sources
is lacking. The major source of nitrates in urban atmospheres is the
conversion by atmospheric reactions of oxides of nitrogen emitted from
the combustion of fossil fuels in stationary and mobile sources. Oxides
of nitrogen (primarily NO) are produced in the combustion process by two
mechanisms. The first is the fixation of atmospheric nitrogen,
which has an exponential dependence on temperature and a lesser
dependence upon the oxygen supply. The second mechanism is the
oxidation of chemically bound nitrogen in the fuel itself in which
the rate of oxidation is nearly temperature independent, but is strongly
dependent upon the availability of oxygen. It has been estimated that
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approximately 50 percent of the NO produced by the combustion of
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heavy oils and coal is created by the conversion of the bound nitrogen.
The formation mechanisms provide the basis for control strategies.
The general control techniques are known, however, application to
practical combustion systems is uncertain. Experience to date is
almost exclusively with utility boilers.
Federal emission standards have been promulgated for new utility
boilers greater than 25 MW for all fuels. Boiler manufacturers are
currently selling units guaranteed to meet these standards. Federal
emission standards are being proposed for stationary gas turbines fired
with natural gas or fuel oil. Emission standards have not been proposed
for other stationary sources. Standards for oxides of nitrogen emissions
from mobile sources have been established.
The primary source of nitrates from industrial processes is
fertilizer manufacturing. Data are limited as to the actual nitrate
emitted. Nitric acid plants emit up to 10 ppm of HN03 mist, but the
major contribution to the atmospheric nitrate loading is probably from
the conversion of NO to nitrate. No proven control process is
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available. Advanced waste treatment processes are being developed which
include nitrogen-removal by ammonia-stripping, ion-exchange, biological
nitrification and break-point chlorination. The effectiveness of
these processes vary depending upon the quality of the water involved,
where they are used, and their location in the treatment sequence.
Pilot studies have indicated that a substantial increment of organic
nitrogen in the effluent from an activated sludge plant could be
removed by filtration. The problems associated with the handling
of runoff and solid waste are equally as complex—not the least of
which is economics.
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Although data are limited, it does not appear that significant
amounts of nitrate are emitted directly from mobile sources.
Atmospheric loading from mobile sources results primarily from the
emission of oxides of nitrogen (principally NO) which is then
converted to nitrates in the ambient atmosphere. Control strategies
consist of modification of conventional internal combustion engines
with oxidation catalyst, and the introduction ,of alternative com-
bustion engines. Statuatory emission standards for mobile sources
have been established at 0.4 g/mile to be achieved by 1977.
Nitrate and oxides of nitrogen concentrations in the atmosphere
are not linearly related. An acceptable level of nitrate in the
atmosphere has not been established, and the data to do so are not
adequate at this time. Further, it is not known as to what extent
it would be necessary to reduce oxides of nitrogen emissions in order
to achieve an acceptable level of nitrate.
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1.2 CONCLUSIONS
The total nitrogen cycle of the earth is still not well understood.
The distribution and rate of transfer within and between the lithosphere,
atmosphere, and hydrosphere can be estimated only within very broad
limits. The only quantities thought to be known reasonably well are
the amount of nitrogen in the atmosphere and the rate of industrial
fixation. With the very limited available information, it is difficult
to determine precisely the extent to which man's activities may
influence the total nitrogen cycle. There is no clear evidence to
indicate that these activities tend to deplete the earth's nitrogen
supply in the same manner that has been argued in the case of oxygen.
There is, however, rather conclusive evidence that man has significantly
altered the distribution and chemical characteristics of nitrogen
compounds in some of those regions in which he lives. Some obvious
and undesirable effects of this alteration have been observed, and
speculations made on more subtle effects that may be even more sig-
nificant. Hence assessment can only be partially complete. Based
upon available data, a number of conclusions can be reasonably drawn:
(1) There are three major activities by which man has
significantly altered the nitrogen cycle on a local scale:
combustion of fossil fuels, industrial fixation of nitrogen,
and the generation and disposal of waste materials.
(2) All of the above activities at present are essential to
man's existence and well-being; therefore feasible corrective
actions must be based on management practices which in most
cases are technology dependent.
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(3) The atmospheric nitrate concentration in certain
urban regions in the U. S. has been increased by orders
of magnitude on a temporal basis as a result of the
combustion of fossil fuels. The areas affected may extend
several hundred kilometers beyond the urban boundaries.
(4) Nitrates in the urbar, atmospheres are due primarily
to atmospheric chemical transformations of oxides of
nitrogen. The secondary nitrate-containing aerosols are
normally found in a size range of less than 1 ytn in
diameter, except possibly in maritime areas; therefore,
they are respirable particles, of a size range which
affects visibility, remain suspended in the atmosphere
for several days, and are involved in precipitation
processes.
(5) Nitrates in urban atmospheres consist of a wide range
of inorganic and organic compounds. The present monitoring
system is inadequate to characterize these compounds.
(6) Studies have shown a statistical association between
nitrates, as measured in the atmosphere, and the aggravation
of asthma. These results suggest that some nitrate compounds
may be a direct respiratory irritant, although this has not
been shown conclusively. Specific nitrate compounds have
not been identified as causal agents.
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(7) Certain species of nitrates found in urban atmosphere
are known to cause damage to sensitive agricultural crops
and ornamental plants.
(8) The nitrate concentration in precipitation in urban areas
is generally higher than in non-urban areas.
(9) The nitrate concentrations in some surface waters in the
U. S. have increased significantly as a result of discharges
of municipal waste water; runoff from agricultural land and
urban areas; animal wastes, refuse disposal and precipitation.
Precipitation nitrate accounts for approximately 10 percent
of man-made contributions to surface waters. This has resulted
in eutrophication of lakes and streams, with associated major
shifts in aquatic life and production of algal blooms. Nitrates
in well water in many parts of the U. S. exceed the PHS standard
of 10 mg/1.
(10) The uptake efficiency of nitrogen by plants is less than
50 percent. The probability of loss of nitrogen from the soil
increases with intensified agriculture. Portions of the
natural organic nitrogen, as well as industrial nitrogen applied
as fertilizer, will be lost. The amount and rate of loss varies
depending upon soil conditions, biotic species, rainfall,
agricultural practices, etc. The contribution of nitrate in
precipitation to nitrate concentration in plants is small
compared to that from application of nitrogen containing
fertilizers. The changes in the nitrogen cycle resulting from
agricultural activities do not appear to be permanent or irreversible.
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(11) The nitrate-nitrite content in foods used by man and
domestic animals varied widely. Some of the variability may
be attributed to fertilization practices; and to processing,
handling, storage, and preparation practices. Certain foods
and tobacco contain small concentrations of nitroso compounds
which have been shown to be carcinogens in animals.
(12) Methemoglobinemia in humans (principally infants) and
domestic animals has been associated with excess nitrate in
water and food, but not in "the atmosphere. The contribution
from inhalation is small except in cases of occupational
exposure.
(13) Nitrosamines have been shown to be potent carcinogens
in animals. Ingestion of nitrite with certain secondary amines
can lead to the formation of the carcinogenic nitrosamines in
animals. This information suggest that these compounds in
foodstuffs may contribute to the occurrence of human cancer,
although this has not been shown.
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1.3 RECOMMENDATIONS
Nitrate is one of the end products of urban aerosol and smog
formation, although it is not known what extent the oxides of
nitrogen precursors serve as the universal controlling factor.
The bulk of the available data only provides an estimate of the
water soluble nitrates collected on a glass fiber filter. Based
upon laboratory experiments and theoretical considerations, it is
known that a wide variety of nitrate compounds may occur in the
ambient atmosphere of urban areas and that these may be trans-
ported for hundreds of kilometers from the source area. The exact
chemical nature of these nitrate compounds has not been determined.
The controlling rate factors for urban aerosol formation are not
well known. It is known that the rates of reaction are concentration
dependent, but they are not linearly related. Knowledge concerning
these factors is particularly important to the question of control
strategies. Present capability is inadequate to provide the necessary
data to characterize and understand the nature and extent of aerosol
formation in ambient urban atmospheres in general, and the role of
nitrates in particular. Available data indicate that atmospheric
nitrates may act as direct respiratory irritants, and theoretical
considerations suggest that they may contribute to the increased
incidence of human cancer observed in urban areas; although the
specific compounds which may serve as the causal agents, and their
threshold values, are not known. Concentrations of these specific
compounds have not been measured in the ambient atmosphere. The
above are important gaps in our knowledge concerning nitrates in
the atmosphere. Insight into these problem areas is essential if
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prudent decisions concerning the control of urban smog formation
and its undesirable consequences are to be made. The following
research and interim action programs are recommended.
(1) The present oxides of nitrogen air quality and emissions
standards for stationary and mobile sources should be achieved
within the shortest time frame feasible consistent with economic
and social constraints. Based on possible health effects, specific
emphasis should be given to those geographical areas which currently
experience annual average nitrate concentrations on the order of
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3 to 5 yg/m or greater.
(2) An adequate routine monitoring system for nitrates should
be developed and implemented. This will require the development and
evaluation of suitable nitrate measurement techniques. Information
concerning size distribution and chemical characterization should be
either implicit in the concentration measurement or these parameters
should be studied directly. Standard siting, sampling, preparation and
analysis procedures should be established for the monitoring network.
(3) An extensive research effort is recommended to obtain an
understanding of the nitrate aerosol formation, transport and removal
processes in urban atmospheres. Particular emphasis should be given
to determining the controlling rates of reaction as a function of
precursor concentrations, and the transfer mechanisms (atmospheric
removal) between the atmosphere, water, and soil. An integral part
of this effort should be the chemical and physical characterization
of nitrate aerosols in the different urban areas. Specific emphasis
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should be given to carcinogenic compounds. It should also include
the extent to which atmospheric removal processes contribute to the
eutrophication of our surface waters.
(4) The question of the contribution of urban air pollution
to the prevalence of human cancer and respiratory diseases is
paramount. An integrated clinical, epidemiological, and toxicological
research program is recommended to resolve the,question of the
"urban factor" in general, and the cause and effect relationship of
inhalation of specific compounds in particular. This program should
address specifically the questions of whether or not certain nitrate
compounds observed in the atmosphere serve as direct respiratory
irritants, and whether or not they contribute to human carcinogenicity,
mutagenicity, and teratogenicity.
(5) Suitable strategies and technology for the control of primary
nitrate particulate emissions, and the processes of secondary nitrate
aerosols, should be developed. The program should include the develop-
ment and implementation of interim measures to control the oxides of
nitrogen emissions from the combustion of fossil fuels. Knowledge
obtained from the research effort recommended above will be necessary
to select the appropriate control strategy and to determine the level
at which emissions should be controlled. The ultimate solution must
address the problem of alternative energy sources. Nitrates in the
atmosphere resulting from industrial product sources can best be
controlled on the basis of emissions, equipment performance, and
disposal standards.
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2. INTRODUCTION DO NOT QUOTE OR
2.1 GENERAL
The purpose of this document is to summarize the current
knowledge regarding nitrates in the atmosphere, and their effects
upon human health and welfare; and to assess this knowledge base
with respect to the need for control of the activities of man
which impact upon the distribution of nitrates in the atmosphere
and some aspects of its relation to the total nitrogen cycle.
It is not intended that the document constitute an in-depth
scientific summary. The references cited do not represent a
complete bibliography. Primary emphasis is placed upon those
aspects of the problem which are considered most important rela-
tive to the decision-making processes which are the responsibility
of the Environmental Protection Agency.
The many forms of nitrogen are an integral part of our natural
environment and are intricately related to the complex life cycle ,
details of which are beyond the scope of this document. The
distribution of nitrates on a regional basis in the atmosphere,
soil, water, and food is significantly influenced by the combustion
of fuels, and agricultural and waste disposal practices. Man's
contribution to the nitrate loading in soils, water, and food has
been reviewed recently by the National Academy of Sciences, the
2
Hazardous Materials Advisory Committee for EPA, and the Food and
3
Drug Administration. Although the man-made atmospheric nitrate
contribution to the total nitrogen cycle is small compared to that
from agricultural and waste disposal practices, its effect upon
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urban regions, and hence human health and welfare, is of primary
significance. Here we are concerned with the impact of man's activi-
ties on the transformation processes and distribution of nitrates
in the atmosphere. Particular attention is given to those interface
problems which hopefully will provide the insight necessary for
prudent decisions regarding management and control. These include
measurement and analytical techniques necessary for monitoring
environmental loading and man's contribution thereto, transformations
and behavior within the nitrogen cycle, mechanisms and risks of
exposure and response, undesirable effects, and control technology.
2.2 REFERENCES
1. Accumulation of Nitrate, Committee on Nitrate Accumulation,
Agricultural Board, Division of Biology and Agriculture,
National Research Council, National Academy of Sciences,
Washington, D. C. 1972.
2. Nitrogenous Compounds in the Environment, Hazardous Materials
Advisory Committee, U. S. Environmental Protection Agency,
Washington, December 1973.
3. Review on the Chemistry and Toxicology of Nitrites, Nitrates,
and Nitroso Compounds (Nitrosamines), Bureau of Foods and
Pesticides, Bureau of Drugs, Food and Drug Administration,
Department of Health, Education and Welfare, Washington, D. C.,
August 1970.
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3. CHEMICAL AND PHYSICAL PROPERTIES <
3.1 ATMOSPHERIC NITRATES
3.1.1 Peroxyacetyl Nitrate
Peroxyacetyl nitrate (PAN) is by far t$e most abundant
.»•' 9-
nitrate in the polluted air of major cities. This fact was revealed
1 3
in studies of the atmospheric photochemistry in the 1950's,
It was a surprising result, because the compound is unknown in any
other aspect of chemistry; the material is indigenous only to
photochemical air pollution. PAN is the principal member of a
family of compounds, four of which have been synthesized and studied
to date. The four are peroxyacetyl nitrate, peroxypropionyl ni-
4
trate, peroxybutyryl nitrate and peroxybenzoyl nitrate. The
one-carbon member of the family, peroxyformyl nitrate, has never
been synthesized or isolated, in spite of many attempts. At present
it must be assumed that the formyl compound is too unstable for
existence at normal temperatures.
The family of compounds was discovered in an infrared study of
photochemical reactions of hydrocarbons and nitrogen oxides in air. A
strange set of infrared bands appeared in the spectra, apparently
due to an unknown compound. It was immediately suspected that the
compound was a nitrogen containing species because nitrogen oxides
were being consumed in the reactions. From known infrared spectra
it could be stated that the new compound was not a known type of
organic nitrate, nitrite, nitro compound, or nitroso compound.
For some time the new molecule was merely designated compound X.
Through study of its physical and chemical properties, and through
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the synthesis route, the material was finally identified as peroxy
acetyl nitrate, with the structural formula CH,C (O)OONOo. The
presence of the material in the Los Angeles atmosphere was verified
p
by means of the infrared absorption spectrum in 1957. Shortly
thereafter it was shown that the material was damaging to vegetation
and irritating.to the eyes.
PAN cannot be purchased commercially. For laboratory study, it
'r ' . • , •
must be specially synthesized. There are three methods of synthesis
which involve reactions in the gas phase. First is the photolysis
of diacetyl and NOg in air. Second is the photolysis of ethyl
nitrite in air; and third is the dark reaction of acetaldehyde,
ozone and nitrogen dioxide in air. Each of these reactions follows
the path of reaction in the polluted atmosphere. Acetyl radicals
are produced; they add oxygen; the resulting peroxyacetate radicals
then add !N09 to form PAN. The PAN must then be condensed out of the
ft
gaseous mixture and purified by passing it through a chromatographic
column.
Handlinn liquid PAN is dangerous. A few drops in a test tube
have exploded with great violence. This is not surprising, since
the elemental compositions of PAN and nitroglycerine are almost the
same. PAN may actually release more energy per gram in explosions
than nitroglycerine, and PAN may be easier to detonate.
In the gas phase, PAN has a fairly high degree of thermal stability.
In clean glassware pure PAN will last many days and even weeks. In
the presence of aerosols, metal surfaces, or other gaseous
pollutants, .the PAN may decompose in unpredictable ways.
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It disappears overnight in polluted air, while other more inert
pollutants, such as carbon monoxide, remain. The reactions of
PAN with other substances is a subject much in need of further
study.
PAN is apparently not photolyzed by sunlight at any appreciable
rate. The ultraviolet spectrum reported by Stephens shows no
absorption in1 the solar region.
A most important property of PAN, distinguishing it from
other organic nitrates, is its behavior when in contact with an
4
aqueous medium. The compound appears to have some solubility
in water. The aqueous solution is capable of oxidizing organic
materials, and in so doing it leaves nitrous acid and acetic
acid in solution. If the solution is basic, these acids are of
course neutralized.
This oxidizing behavior differentiates PAN from alkyl nitrates
and nitrate salts which are not easily reduced to nitrites. The
known damaging effects of PAN on plant and animal tissues is
undoubtedly a result of the oxygen evolution. The nitrite ion left
behind after the oxidation undoubtedly has a further potential
for harming the organism.
3.1.2 Nitrogen Pentoxide
There is no doubt that nitrogen pentoxide exists in the
atmosphere, even though it has not yet been measured. It should
be found to have much higher concentrations in the cities
than in the rural areas. The pentoxide is formed when NC^ reacts
with ozone. This is one of the major paths for removal of N02
from the air. Initially, NO^ is formed; but this reacts quickly
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with N0£ to yield the ^Og. This is shown in equations (19) and
(20) of Section 5.3.
Nitrogen pentoxide is a poisonous gas with great oxidizing
power. It will hydrolyze at the surface of fine particles in
the atmosphere, yielding nitric acid. As mentioned previoulsy,
the hydrolysis may be accompanied by neutralization and displace-
ment reactions which yield nitrate salts in solution. The hy-
drolysis of ^Og is the principal source of the nitrates in the
atmospheric aerosol particles.
3.1.3 Nitric Acid
Nitric acid in the gaseous state is colorless and photochemically
stable. It is a volatile acid; for example, a solution which is
90 percent nitric acid and 10 percent water has a nitric acid
vapor pressure of 20 Torr at 20°C. This equilibrium vapor pressure
goes down with increasing dilution of the acid, but for the equili-
brium partial pressure of acid to be only a few parts-per-hun-
dred-mi 11 ion, as might be present in the atmosphere, the solution
would have to be extremely dilute. It is reasonable to conclude
that in the atmosphere, the vapor will not be taken into droplets
and retained unless the droplets contain reactants to neutralize
the acid. Such neutralization reactions are probably responsible
for holding the nitric acid concentrations down to very low levels
in the photochemically reacting polluted air. Low concentrations
of the acid vapor are readily adsorbed on surfaces, especially
metal surfaces. The acid can react with the surface material pro-
ducing nitrate salts.
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3.1.4 Nitrous Acid
DRAFT
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Nitrous acid is not formed in such great quantities as nitric
acid, but it is important because of its photochemical reactivity.
Nitrous acid absorbs sunlight and dissociates into OH and NO
radicals, equation (12), Section 5.3. The photolysis is only about
one-tenth as fast as the photolysis of N02, but it undoubtedly has an im-
portant role in photochemical pollution because of the high degree of
reactivity of the OH radicals.
When NO and N0? are present in the atmosphere, nitrous acid
will be formed as a result of the equilibrium in equation (11),
Section 5.3 The equilibrium amount of HN02 in the morning air of a city
can be comparable to the amount of NOp- When photolysis sets in, the
nitrous acid concentration will be driven below the equilibrium level.
3.1.5 Nitrate Salts
Nitrate salts are all highly soluble in water. The alkali
metal nitrates give a nearly neutral solution, but ammonium
nitrate gives an acidic solution. These salts have many uses
in the chemical industry, in agriculture, and even in medicines
and foods. They are not generally considered poisonous or dan-
gerous to handle.
3.1.6 Alkyl Nitrates
The alkyl nitrates have a more truly organic character than the
peroxyacetyl nitrates. They are more volatile and less soluble
in aqueous media. They do not have a great tendency to hydrolyze.
The alkyl nitrates are not photo-dissociated by sunlight. This,
of course, is the reason that small quantities of them may
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accumulate in the atmosphere as they are formed from the reaction
of alkoxy radicals and N02 (Equation (25), Section 5.3).
3.2 REFERENCES
1. Stephens, E.R., W.E. Scott, P.L. Hanst, and R.C. Doerr. Recent
Developments in the Study of the Organic Chemistry of the
Atmosphere. J. Air Pollution Control Assn. 6;159, 1969.
2. Scott, W.E., E.R. Stephens, P.L. Hanst, and R.C. Doerr. Further
Developments in the Chemistry of the Atmosphere. Proc. A.P.I.
37, (III), 171, 1957.
3. Leighton, P.A. Photochemistry of Air Pollution. Academic Press,
New York, 1961. p. 158.
4. Stephens, E.R. The Formation, Reactions, and Properties of
Peroxyacyl Nitrates (PANs) in Photochemical Air Pollution,
in Advances in Environmental Science and Technology. Vol. I.
J.N. Pitts,(ed.). Wiley, New York, 1969.
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DO A/07-r-' 'l3"i
4. MEASUREMENT TECHNIQUES ' feu7~ O-s
^ W CITE
4.1 ENVIRONMENTAL
4.1.1 Sampling, Preparation, and Analysis
A variety of nitrogen compounds are found in the atmosphere (Table 4^1).
Some are found in the gas phase, such as the nitrogen oxides, ammonia,
lower molecular weight amines (e.g. aniline), nitro compounds (e.g. nitro-
benzene), aza arenes (e.g. pyridines), amino arenes (e.g. N-alkylcarbazoles)
and organic nitrates (e.g. the PAN family). Nitrogen compounds found in
the particulate phase include ammonium salts, large aromatic amines, amino
acids, proteins and other large organic amines, polynuclear aza arenes,
polynuclear imino arenes, inorganic nitrites and inorganic nitrates. In
terms of mass balance for atmospheric nitrogen, organic nitrogen accounted
for 87 percent of the total nitrogen in a coastal Oregon forest opening, with
the remainder coming entirely from nitrate nitrogen, ammonia nitrogen not
being present at detectable levels. Much more of this type of work needs
to be done to delineate the percentage of the various nitrogen compounds
found in the atmosphere.
The purpose of this section is to discuss the measurement of atmospheric
nitrates found in the gas, vapor and particulate phases.
4.1.1.1 laorgairjc Nitrates--The variety of inorganic nitrates available
include alkali nitrates (sodium, potassium), alkaline earth nitrates (calcium,
barium) and many other metallic nitrates, as well as ammonium nitrate and
nitric acid. All of the inorganic nitrates are soluble in water and the
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TABLE 4.1
TYPES OF N COMPOUNDS FOUND IN THE ATMOSPHERE.
.
NOT QL'OT£ OR CITE
1. Nitrogen oxides
a. Nitrous oxide ^0
b. Nitric oxide NO
c. Nitrogen dioxide N02
d. Miscellaneous nitrogen compounds of lesser importance, e.g., N03
N2°3> N2°4 and N2°5-
2. Ammonia and Ammonium Salts
3. Amines
- aliphatic, aromatic, amino acids, proteins and other organic
amines.
4. Nitro compounds
- nitroarenes (e.g., nitrobenzene) and nitroalkanes.
5. Aza arenes
- pyridines, quinolines, benzoquinolines, benzacridines, etc.
6. Imino arenes
- pyrrole, carbazole, etc.
7. Inorganic nitrites
8. Inorganic nitrates
9. Organic nitrates
a. PAN family
1. Peroxyacetyl nitrate
2. Peroxypropionyl nitrate
3. Peroxybenzoyl nitrate?
b. Alkyl nitrates
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thermal decomposition temperatures range from that of 210° C for
NH4N03 to > 590° C for Ba(N03)2.
4.1.1.1.1 Sampling an_d Sample'Preparation--Normal procedure for participate
sampling is to employ the high volume sampler with an 8" X 10" glass fiber
filter that is operated over a 24 hr. period. The average of the flow
rates at the beginning and end of the sampling period is taken as the
rate for the entire sampling period. This permits collection of
3
nitrate from some 2,000 m of air. Usually only a small fraction of the
sample is taken for analysis by the nitration methods, which require from
10-100 yg NOg"* Membrane filters at a sampling rate of 10 liters/min.
and electrostatic precipators at a sampling rate of about 100 liters/min.
2
may also be used. However, the high volume sampler is recommended.
To facilitate storage and transportation, the hi-vol sample filter
is folded upon itself along the 10-inch axis. This fold may result in a
nonhomogeneous area in the sample, so all sample aliquoting is made across
the fold. Using a template, a 3/4 inch wide strip is cut across the 8-inch
dimension of the exposed portion of the filter and the nitrates eluted with
water. The liquid is transferred quantitatively to a graduated cylinder
and diluted to a predetermined mark.
An analytical procedure that has been used to determine particulate
nitrates and that does not require sample preparation.is computer con-
3
trolled high resolution mass spectrometry. A single stage impactor
collects particulate matter of diameters greater than 1-2 ym and the re-
mainder of the particulate matter is collected on a glass fiber paper.
The particulate samples are introduced directly into the mass spectro-
meter utilizing a temperature programmed insertion probe.
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Collection of water samples from ocean, lake, river, pond and well
is straightforward, as is nitrate recovery from soiUs, which can be
19
done by ultrasonic extraction using solvents as copper sulfate, #10 KCI
20
or water. Nitrate can be extracted from powdered plant material with
20
an acid buffer containing ammonium sulfamate. Urine, water samples and
hot water extracts of plant material can be used without further pre-
paration. Whole blood, serum, plasma and milk must undergo protein
21
precipitation. For whole blood, this is accomplished by taking blood
with triple distilled water and precipitating with Ba (OH^ and zinc
sulfates. Meat samples have to be homogenized, treated with buffer,
charcoal, mixed well, treated with zinc acetate and potassium ferri-
22
cyanide and filtered.
4.1.1.1.2 Analysis—The variety of methods available for the analysis
of environmental and biological inorganic nitrates is shown in Table 4.2.
All the methods listed measure total inorganic nitrate while method
4b, thermal decomposition - mass spectrometry, measures some individual
nitrates. However, the effect of heat on the relative composition of
particulate nitrates is not known.
A number of methods have been used for analysis of inorganic nitrate
in atmospheric particulate matter, but essentially they are water analysis
methods applied to the analysis of airborne particles. They include the
23
colorimetric methods using the reagents brucine (2), 2,4-xylenol and
sul fanil amide-N-(l-naphthyl)ethylenediamine. The main difficulties
with these methods are inadequate sensitivity, complexity and interferences.
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TABLE 4.2
AVAILABLE METHODS FOR NITRATE
Ref.
1. Nitration methods
a. Colorimetry (4,5)
b. Fluorimetry
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u
Reduction to nitrite and assay
a. Colorimetry
b . Fluorimetry
3. Reduction to ammonia and assay
a. Colorimetry (9^
b. Gas chromatography
c. Titration 00)'
4. Thermal decomposition
a. Chemi luminescence
b. Mass spectrometry (3)
5. Ion selective electrode
6. Enzymic
a. Colorimetry 02)
b. Fluorimetry
7, Chemical decomposition
a . Manometry Q 3 )"
b. Chemi luminescence
c. Gas chromatography
8. Oxidation methods
a. Colorimetry 04)
b. Fluorimetry
9. Complex formation and extraction into organic solvents.
a. Colorimetry 05)
b. Atomic absorption spectrometry
10. Direct
a. Spectrophotometry (17)
11. Extraction (of nitrate salt)
a. Spectrophotometry
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The high resolution mass spectrometricmethod using thermal decomposition
has been used to identify ammonium and sodium nitrate 1n airborne parti -
culate, but the effect of heat on the relative composition of particulate
nitrates needs more investigation.
Methods used to measure nitrate in water are either direct or
indirect. The direct procedures include the phenoldlsulfonic acid
method, the brucine method, the chromotroplc acid method and the ultra-
25
violet spectrophotometric method. Interferences result from strong
oxidizing and reducing agents, residual chlorine, nitrite and dissolved
organic matter. The indirect procedures involve reduction of nitrate to
nitrite using zinc, cadmium or hydrazine and determination of the nitrite
thus formed with sulfanil amide-N-(l-naphthyl)ethylenediamine. Oxidizing
and reducing agents interfere as well as heavy metals.
Analogous methods are used for the analysis of inorganic nitrates in
soil, foods, body fluids, and plant and animal tissue. In some cases
methods are available for the determination of both nitrate and nitrite
and have been applied to the determination of these entities in biological
fluids.21
4.1.1.2 Organic Nitrates—Only a little information is available on the
organic nitrates present in polluted atmospheres, but even less is known
26
about the organic nitrates in water, food, body fluids and tissue.
Future research efforts on organic nitrates should entail development
of sampling and analytical techniques for the various RN02 and RON02 vapors
present in the atmosphere, e.g., aryl nitrates, alkyl nitrates, nitro-
alkanes, nitroarenes, organic nitrites, etc. In Addition, work should be
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done on the differentiation and identification of organic nitrates present
in the particulate and vapor phases.
4.1.1.2.1 Parti cul ate Ni trate--Very little is known about participate
nitrate: Characterization and analytical procedures have not as yet
been developed. Methods could be readily developed for particulate
organic nitrate utilizing the fact that inorganic nitrates are soluble
in water while organic nitrates are not. A routine survey method for
the determination of total organic particulate nitrates is badly needed.
4.1.1.2.2 Alky! Nitrates—These compounds would be found in the vapor
phase. Alkyl nitrates have been detected in smog chamber studies, but
procedures for their collection from polluted atmospheres and subsequent
analysis have never been developed. More research needs to be done in this
area.
4.1.1.2.3 Miscellaneous Nitrates—Air and water near plants which manu-
facture nitrate containing explosives would be expected to contain TNT,
picric acid, nitrophenola and various organic nitrates. Satisfactory
methods of sampling, identification and analysis have not yet been developed
for these organic nitrates. Organic nitrates have been reported in the
27
urban atmosphere. and have been formed under simulated atmospheric con-
28
ditions in a smog chamber. Thus, from the reaction of cyclohexene and
NO in the presence of light the nitrate of 5-hydroxypentanoic acid has
/\
been tentatively identified. Similarly, from the reaction of toluene and
NO tentative identification with the help of GC-MS has been obtained for
^
nitro-o-cresols, nitrohydroxybenzyl alcohol and 1,3-hexanedione-5-nitrate.
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4.1.1.2.4—PAN Fatiiily—The peroxyacyl njtrates (PANs) are a family of un-
stable highly oxidized organic nitrogen compounds which are formed in
polluted air by the photochemical action of sunlight on hydrocarbons and
29
nitrogen oxides. They can be considered to be acylating, nitrating and
oxidizing agents. The general formula is
0
II
R C 0 0 N00
The first member of the family, in which R is a methyl group is called
peroxyacetyl nitrate (PAN), and has received most of the research effort.
A few tens of parts per billion (by volume) of PAN are present in photo-
chemical smog, concentrations which can injure sensitive plants in a few
hours exposure. Higher homologs, with R = ethyl (peroxypropionyl
nitrate, PPN), R = propyl (peroxy-n-butyryl nitrate, PnBN), and R = phenyl
(peroxybenzoyl nitrate, PBzN) are also toxic but they have been studied
to a lesser extent. The PANs not only damage plants but are powerful eye
irritants as well. PBxN is reported to cause eye irritation at a con-
centration of about 5 ppb. This makes it about 100 times as irritating
as PAN.32
PAN was first detected in the atmosphere, and partially characterized
33-35
with long-path infrared spectrometry. This method, at one time, was
the only means available for measuring PAN at concentrations approaching
those found in polluted atmospheres, and even this method has limitations
because of expense of instruments, the large volume of sample required,
and a threshold of detectability of about 50 ppb, which is somewhat above
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that required to induce plant damage. x
Gas chromatography with electron capture detection evolved as a method
on oc "3Q
for measuring PANs JU>ODr"00 because electron capture detectors are very
sensitive to nitrates and mych less sensitive to hydrocarbons and simple
oxygenates that would be found in atmospheric samples. In one particular
investigation the column that was used was a 3 foot x 3 mm o.d. glass
column, packed with 5% Carbowax 400 to 100 to 200 mesh Chromosorb W, and
was operated at 35°C with a nitrogen carrier gas flow of 25 standard ml
per minute at 10 p.s.i.g. A chromatogram resulting from the injection of
a 2 ml sample containing methyl and ethyl nitrate, PAN and PPN is shown
in Figure 4-1. Identification of the PAN and PPN was confirmed by warming
the sample flasks whereby the PAN and PPN peaks decreased while the
methyl and ethyl nitrate peaks increased; the latter two compounds are
the main decomposition products of the PANs. Since PAN is very unstable
and the concentration in the air is in parts per billion, the operating •
variables must be carefully chosen. Using this system, atmospheric samples
taken on an afternoon of heavy air pollution in Riverside, California,
showed peaks indicating 50 ppb of PAN and 6 ppb of PPN. These results
constitute the first detection of PPN in polluted atmosphere. Further,
the results showed that the important PANs could be measured in polluted
atmospheres at less that 10 ppb with an electron capture detector using
very small untreated samples. Maximum sensitivity permits detection
-9 3
of concentrations of less than 1 ppb (10 by volume or 4.95 yg/M
at 25°C and 760 Torr) using this technique.38
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At the University of California at Riverside an automated gas
chromatograph with electron capture detection has been sampling the
atmosphere every fifteen minutes for several years and providing a PAN
measurement. For calibration purposes, small quantities of PAN are
generated by photolyzing ethyl nitrite vapor in an atmosphere of oxygen.
Ambient air samples can be analyzed by filling a 100 ml syringe outdoors
and then attaching it to the gas sample valve or by using an automated
system consisting of a timing unit, valve and sample loop and solenoid
lab!
38
38
switch. In terms of effect of storage, no reliable means of retaining
PAN-containing air samples is known at this time.'
Another method used to assay PAN involves the measurement of the
chemiluminescence produced from the gas phase reaction of PAN with tri-
39
ethylamine vapor, Figure 4-2. Concentrations of PAN as low as 6 ppb
were detectable with this method. Improvement of the light detection
system may permit measurement of PAN concentrations as low as 1 ppb or
less. Application of this method could result in a technique for measuring
atmospheric concentrations of PAN. In later work done on this technique,
it was found that chemiluminescence generated by bubbling PAN through
a solution of triethylamine in acetone (1% V/V) was approximately twice
as intense as that produced 1n the vapor phase reaction and the results
39
were more reproducible. Therefore, the liquid phase chemiluminescnece
seems to be more suitable for atmospheric monitoring applications.
The existence of the powerful lachrymator peroxybenzoyl nitrate (PBzN)
as a laboratory-generated photochemical smog product was first established
31
in 1968. Although earlier research suggested that all mono-alkyl
benzenes having at least two benzylic hydrogens form PBzN, its formation
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D0 NOT QUOTE OR
o
fe
SAMPLE = 2 Rif
¥ 9 + ,°;
-C-C-O-O-N; !
vo
H-C
H
PAN-250fig/ni3
I
PPN -150
I
I
1234
MINUTES
Figure 4-1. Chromatogram of mixture of peroxyacyl
nitrates and alkyl nitrates. GC(ecd) - 5% CARBOWAX
400 on CHROMOSORB W (100-200 MESH).
-37-
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'ITE
1.0
0.8
0.6
0.4
0.2
PAN
400
500
600
X(nm)
700
BOO
Figure 4-2. Emission spectra of chemiluminescent
reaction of triethylamine with PAN.
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from toluene-NOv mixtures could not be established in smog chambers
A
due to analytical interferences with the Gas-Liquid Chromatography-
Electron Capture (GLC-EC) method. A new and more sensitive procedure has
been proposed which involves quantitative conversion of PBzN to methyl
benzoate (MeOBz) by trapping this compound in basic methanol solution
followed by Gas Liquid Chromatography-Flame lonization Detection (GLC-FID)
determination of the MeOBx. Using this method, the formation of PBzN
from laboratory irradiation of toluene - NO in air was demonstrated and
A
quantitatively determined. In preliminary atmospheric sampling trials
in the San Francisco Bay Area on a day of light photochemical smog
(i.e., oxidant max. =0.1 ppm), PBzN was sampled by means of fritted glass
bubblers containing methanol and sodium methoxide and maintained at 0°C.
The tests demonstrated that, if PBzN were present, its concentration
was £ 0.07 ppb, and a recommendation was made to repeat the tests on days
of high oxidant level. One of the limitations of this analytical technique
for PBzN is that imposed by the need for microliter-size GLC injections
of methyl benzoate in methanol. From 5 ml of collection solution, a
2 yl GLC sample measures only 0.04% of the methyl benzoate formed. Although
this method is the most sensitive for any PAN, it could be considerably im-
proved by the use of wet chemistry and high performance liquid chromatography
(HPLC); sensitivity could be increased about 100 to 1000 times and with
proper modification of this sampling-analytical procedure the whole family
of PANs could be determined. Research effort in this area of investigation
should be supported and a method developed. A more sensitive screening
method for the PANs is badly needed.
-39-
-------
DRAFT
no PfHT (":< ;"'~"~ /"».n n'T
DO KUi QUUJL UK bli
4.1.1.3 Conclusions—There are a number of areas that deserve Investigation
for the characterization and assay of atmospheric N compounds, many of
which have been mentioned in the preceding sections.
One such area is that of sampling for inorganic nitrate, to determine
if N02 is converted to NOZ under certain sampling conditions. Other
variables that need investigating are filter media, composition of
particulates, composition of the atmospheric gas phase, temperature
effects, humidity effects, artifact formation, etc.
Sensitive methods of sampling and analysis should be developed for
the PANs and organic POIM^ vapors.
Methods for the differentiation and identification of organic and
of inorganic nitrates in particulate should be developed, as well as a
thermal decomposition method for total particulate nitrate.
A thorough evaluation of the three most popular methods for parti -
culate nitrate, including a ruggedness test, should be initiated.
An investigation of the nitrogen balance in the particulate phase
and in the vapor phase should be started to quantitatively determine the
composition of the air in terms of the nitrogen compounds.
-40-
-------
DRAFT
I "7"
4.2 SOURCE MEASUREMENTS J)Q fjQJ Qy.J;:- GFi *,.*
4.2.1 Stationary
Since stationary sources emit essentially no nitrates, no specific
measurement technology has been developed for nitrate emissions. The
commonly used measurement techniques all use adaptations of the popular
particulate sampling trains, which collect the particulate on a filter,
thimble, or bag.
Since all inorganic nitrates are water soluble, the methods are
generally adaptations of classical water quality techniques. Most of
them are reviewed in the ambient air section of this document. The most
widely used analysis is the phenol disulfonic acid procedure developed
by Chamot in 1906.
The major precursors to atmospheric nitrate which is present in
stationary source emissions are nitric oxide (NO) and nitrogen
dioxide (N02). The measurement of both of these emissions is generally
done using the PDS method as published by ASTM (DI608-60) or by EPA
(Test Method 7, CFR, 40_, Part 60, Appendix).
These methods involve sampling a portion of the source effluent
o
into an evacuated flask with a weak hydrogen peroxide solution. The
residual oxygen in the flask oxidizes the NO to N02 over a period of
16 to 30 hours. The N02 then dissolves in the peroxide solution and is
oxidized to NOZ. The nitrate in water solution is then analyzed by the
classical PDS method.
This analytical method, although slow and cumbersome, is adequate
-41-
-------
DO NOT QUOTE OR CITE
for stationary source levels of NOX emissions,. The majpr known inter-
ference to this method 1s the chlorine ion, which causes a 10-20%
low result if present at the 50-rlOO ppm level. Fortunately, none of the
major NO emitting sources have chloride emissions high enough to cause
X
appreciable interference.
The only major research work needed 1n this area is to develop
a sampling procedure which 1s better suited to field use. An analytical
procedure which is less time consuming would be helpful from a cost
standpoint.
The Chemistry and Physics Laboratory 1s currently eveluating a
solid sorbent sampling scheme and a rapid analytical procedure using a
43
selective ion electrode.
-42-
-------
D0 NOT QUOTE OR C/TE
4.3 REFERENCES M U|'t
1. Hoeft, R. G., D. R. Keeney and L. M. Walsh, Nitrogen and Sulfur
in Precipitation and Sulfur Dioxide in the Atmosphere in Wisconsin.
J. Envipon. Quality, 1^, 203-208 (1972).
2. Kothny, E. L., et. al., Tentative Methods of Analysis for Nitrate
in Atmospheric Particulate Matter (Brucine Method). Health Lab.
Sci., £, 324-326 (1972).
3. Schuetzle, D., A. L. Crittenden and R. J. Charlson, Application of
Computer Controlled High Resolution Mass Spectrometry to the Analysis
of Air Pollutants. J. Air Poll. Control Assn., 23. 704-709 (1973).
4. Holty, J. G. and H. S. Potworowski, Brucine Analysis for High Nitrate
Concentrations, Environ. Sci. Techno1., £, 835-837 (1972).
5. Sawicki, E., H. Johnson and T. W. Stanley, Determination of Nitrate
or Nitrate plus Nitrite with 1-Aminopyrene. Anal. Chem., 35, 1934
(1963).
6. Axelrod, H. D., J. E. Bonelli and J. P. Lodge, Jr., Fluorimetric
Determination of Trace Nitrates. Anal. Chim. Acta, 51, 21-24
(1970).
7. Sawicki, C. R., and F. P. Scaringelli, Colorimetric Determination
of Nitrate After Hydrazine Reduction to Nitrite. Microchem. J.,
16^, 657-672 (1971). ~
8. Sawicki, C. R., Fluorimetric Determination of Nitrate. Anal. Lttrs.
4_, 761-775 (1971).
9. Keay, J., and P. M. A. Menage, Automated Determination of
Ammonium and Nitrate in Soil Extracts by Distillation, Analyst,
95_, 379-382 (1970).
1Q,. Stanford, G., J. N. Carter, E. C. Simpson, Jr. and D. E.
Schcwaninger, Nitrate Determination by a Modified Conway Micro-
diffusion Method, JAOAC, 56, 1365-1368 (1973).
11. Qureshi, G. A., and J. Lindquist, A Liquid Ion-Exchange Nitrate-
Selective Electrode Based on Carbon Paste. Anal. Chim. Acta., 67,
243-245 (1973).
12. McNamara, A. L., G. B. Meeker, P. D. Shaw and R. H. Hageman, Use
of a Dissimilatory Nitrate Reductase from Escherichia Coli and
Formate as a Reductive System for Nitrate Assays. J. Agr. Food
Chem., _19_, 229-231 (1971).
13. Van Slyke, D± D., and A. F. LoMonte, Manometric Determination of
Nitrate and Nitrite. Microchem. J_., 14, 608-626 (1969).
-43-
-------
i/wr
DO NOT QUOTE OR CITE
14. Williams, R. J. B., The Rapid Determination of Nitrate in Crop£,
Soils, Drainage and Rainwater by a Simple Field Method Using Diphenyl-
amine or Diphenylbenzidine with Glass Fibre Paper. Chem. Ind.,
1735-1736 (1969).
]5. Yamamoto, Y., N. Okamoto and E. Tao, Spectrophotometric Determination
of Anions by Solvent Extraction with Cuproin- or Neocuproin-Copper(I)
Chelate Cations. Anal. Chim. Acta, 47, 127-137 (1969).
16. Houser, M. E., and M. I. Fauth, Indirect Determination of Nitrate,
Nitrite, and Nitro Groups by Atomic Absorption Spectrophotometry.
Microchem. J_., 15, 399-408 (1970).
I/,* Mertens, J., and D. L. Massart, The Determination of Nitrate in
Mineral Waters by Potentiometry and U. V. Spectrophotometry. Bull.
Soc. Chim. Belg., 82_, 179-190 (1973).
•jg . Burns, D. T., A. G. Fogg and A. Willcox, The Estimation of Nitrate
by Extraction with Tetraphenylphosphonium Chloride. Mikrochim.
Acta, 205 (1971).
1|9. $ien, A., and A. R. Selmer-Olsen, Nitrate Determination in
Soil Extracts with the Nitrate Electrode, Analyst, 94,
888-894 (1969).
20. Milham, P. J., A. S. Ward, R. E. Paull and J. H. Bull; Analysis
of Plants, Soils and Waters for Nitrate by Using an Ion-
selective Electrode, Analyst, 95, 751-757 (1970).
21. Wegner, T. N., Simple, Sensitive Procedure in Determining
Nitrate and Nitrite in Mixtures in Biological Fluids, J_.
Dairy Sci., 55_, 642-644 (1972). ~
22. Fudge, R., and R. W. Truman, The Nitrate and Nitrite Contents
of Meat Products - A Survey by Public Analysts' Laboratories
in South Wales and the South West of England, J. Assoc.
Public Analysts, _U, 19 (1973).
^3. Saltzman, B. E., et.al., Tentative Method of Analysis for Nitrate
in Atmospheric Particulate Matter (2,4-Xylenol Method), in Methods
of Air Sampling and Analysis, Intersociety Committee, American
Public Health Association, Washington, D. C. (1972), pp. 322-324.
24. Morgan, G. B., E. C. Tabor, C. Golden and H. Clements, Automated
Laboratory Procedures for the Analysis of Air Pollutants. Technicon
Corporation, flrdsley, New York (1967), pp. 534-541.
A
-44-
-------
25- Standard Methods for the Examination of Water and Wastewater.
Thirteenth edition, American Pubic Health Association,
Washington, D. C. (1971), pp 233-239, 454-467.
26. Faust, F. D., and J. V. Hunter, Eds. Organic Compounds in
Aquatic EnvironmentSjMarcel Dekker, Inc., N. Y., 1971.
'27- Schuetzle, D., A. L. Crittenden and R. J. Charlson, Presented
at the Symposium on Surface and Colloid Chemistry in Air Pollution
during the 166th AC$ National Meeting, Chicago, Illinois, August,
1973.
,,,. Schwartz, W. E., P. W. Jones, C. J. Riggle and D. F. Miller,
c° The Organic Composition of Model Aerosols, Presented at the
Eastern Analytical Symposium, New York City, November 16, 1973.
29. Stephens, E. R. The Formation^.Reactions and Properties of Peroxyacyl
Nitrates (PAflg) in Photochemical Air Pollution, Adv. Env. Sci., 1^,
119 (1969).
3Q. Stephens, E. R. and M. A. Price, Analysis of an Important Air Pollutant;
Peroxyacetyl Nitrate. J. Chem. Educ., 50, 351-354 (1973).
31. Heuss, J. M., and W. A. Glasson, Hydrocarbon Reactivity and Eye
Irritation. Environ. Sci. Techno 1., 2^, 1109 (1968).
32- Stephens, E. R., E. F. Darley, 0. C. Taylor and W. E. Scott, Photo-
chemical Reaction Products in Air Pollution. Proc. Amer. Pet. Inst.,
40_ (III) 325-338 (1960).
33. Scott, W. E., E. R. Stephens, P. L. Hanst and R. C. Doerr, Proc. Am.
Petrol. Inst. Sect. Ill, 37, 171 (1957).
34. Stephens, E. R.}P. L. Hanst, R. C. Doerr and W. E. Scott, Ind. Eng.
Chem., 48, 1498^ (1956) .
35. Stephens, E. R., W. E. Scott, P. L. Hanst and R. C. Doerr, J. Air
Poll. Control Assoc., 6_, 159 (1956) .
35. Darley, E. F., K. A.i Kejtner and E. R. Stephens, Analysis of Peroxyacyl
Nitrates by Gas Chromatography with Electron Capture Detection. Anal.
Chem., 35_, 589-591 (1963).
37. Izumikawa, T., M. Hayafuku, K. Nakano, K. Asakino and T. Odaira,
Continuous Measurement and Determination of PAN in the Air, Tokyo-to
Kogai Kenkyusho Nenpo, 4_, 41-49 (1973) .
3g. Smith, R. G., et. al., Tentative Method of Analysis for Peroxyacetyl
fllitrate (PAN) in the Atmosphere in Methods of Air Sampling and
Analysis, Intersociety Committee, American Public Health Association,
Washington, D. C. (1972), pp. 215-219.
-45-
-------
39. Pitts, J. N., Jr., H. Fuhr, J. S. Caffney and J. W. Peters, Chemi
nescent Reactions of Peroxyacetyl Nitrate and Ozone with Triethylamine.
Environ. Sci. Technol., 7_, 550-552 (1973).
40. Appel, B., A New and More Sensitive Procedure for Analysis of
Peroxybenzoyl Nitrate. J. Air Poll. Control Assn., 23, 1042-1044
(1973).
41. "Improved Chemical Methods for Sampling and Analysis of Gaseous
Pollutants from the Combustion of Fossil Fuels," Vol. II, Nltroqrn
Oxides, APTD 1291, July 1971.
42. Method 7, 40 CFR, Part 60, Appendix, December 1971.
43. "An Improved Manual Method for NO Emission Measurement."
EPA R2-72-067, October 1972.
-46-
-------
5. ENVIRONMENTAL APPRAISAL
5.1 ORIGIN AND ABUNDANCE
5.1.1 Natural Sources-
Atmospheric nitrates result
from the combination of atmospheric nitrogen and oxygen, starting
with the formation of nitric oxide. In the lower atmosphere,
nitric oxide is created thermally in flames, explosions and electric
discharges. In the upper atmosphere, the compound is formed through
the photo-dissociation of Np and Op, followed by combination of N and
0. Nitric oxide reacts further in the atmosphere and proceeds up the
scale of oxidation in a series of chemical changes, finally reaching
its fully oxidized state as a nitrate salt or acid. Some micro-
organisms return nitrogen directly to the atmosphere as inert nitrogen
gas. Other biological actions release large amounts of nitrogen
compounds (N02, NH3, N20). Estimated global emissions of nitrogen com-
pounds are shown in Table 5.1.
The nitrogen-containing compound wh-ich enters the air fn largest
amount is ammonia, which is generated in animal and plant metabolic
processes. The background concentrations of NhU gas probably range
between three and eight parts per billion. This NhU can be
oxidized up to nitrate, but the extent 6f such oxidation in the air
is apparently not known. The most likely fate of the NH^ is to form
ammonium ion in particulates which are washed to the ground.
-47-
-------
L/i\rti i
DO NOT QUOTE OR CITE
Nitrites are a highly significant part of the nitrogen family of
pollutants, and are an integral part of this review. The
biological effects of the nitrite portion of the family may in fact
be more significant than the biological effects of the nitrates.
There are two main points at which nitrites come into the picture.
The first is the equilibrium among NO, NOg, HgO and nitrous acid,
HNOg. The second is the formation of nitrites on the hydrolysis
of the peroxyacyl nitrates.
5.1.2 Man-made Sources
5.1.2.1 Stationary Sources—Nitrates resulting from stationary source
emissions are primarily formed after conversion of emitted NO and N02 to
acid. Literature documenting primary emissions of nitrates from
stationary sources is scarce. NO emitted by stationary sources will
A-
convert in the atmosphere to nitrates, and since approximately 50 percent
of all man-made NO emissions is from stationary sources, it 1S concluded
A
that stationary sources contribute approximately 50 percent of the man-made
atmospheric nitrates.
NO emissions in the U. S. result primarily from the combustion
A
of fossil fuels in boilers, furnaces, and . internal-
combustion engines. Most emissions from industrial sources are from
steam boilers and process heaters, with smaller amounts from internal -
combustion engines, boilers burning waste-fuel gases, catalytic
cracking regenerators, metallurgical ovens, furnaces, and kilns.
Emissions from pipeline and gas plant operations result primarily
from internal combustion engines used to drive pumps and compressors.
Domestic and commercial sources include incinerators, space heaters,
water heaters, ranges and clothes dryers. Essentially all NO
A
emitted by non-combustion sources are from the manufacture or use of
-60-
nitric acid.
-------
The most abundant form of fixed nitrogen is nitrous oxide, N?0.
2
This gas exists in the air world-wide at 0.25 parts-per-million.
The N20, however, is almost as inert chemically as the N?, and it is
not coupled into the atmosphere cycles of NO, N02, nitrites, and
nitrates. It can be neglected in considering the nitrate
problem
The generally recognized rural background concentration of NO
/\
2
(NO plus N02) is about four parts-per-billion. Some of this may
come from ammonia oxidation, and some from direct emission.
The proportions of NO and NOp will shift diurnally because of
photochemical action. Many millions of tons of NO and N02 of
natural origin are present in the atmosphere at any moment. In
the urban areas, however, the background concentration of ni-
trogen oxides and the nitrates resulting from them is trivial
compared to the concentrations of man-made nitrogen oxides and
nitrates. The human exposure to nitrates is almost entirely the
result of the transformation of nitrogen oxides formed in combustion processes
Lightning produces nitrogen oxides, and hence nitrates.
However, this appears to be a small source, even on a world
O
wide basis.
The nitric oxide formed in the upper atmosphere can also be
neglected in considering human exposure, because it does not
propagate down into the lower troposphere. Some of the upper
atmospheric nitric oxide isconverted to nitric acid. Transport
of the acid vapor to the lower atmosphere also appears to be
negligible.
-49-
-------
„
Compound
NOa
Total
NOz
NH3
NOa
NH3
Table 5.11 '
Estimated annual global emissions of nitrogen compounds
Source
Coal combustion
Petroleum refining
Gasoline combustion
Other oil combustion
Natural gas
combustion
Other combustion
Combustion
Biological action
Biological action
Biological action
Source
magnitude
(tons/yr)
3074 x 10s
11 317 x I0*(bbl)
379 x 10*
894 x 10"
20.56 x I011
1290 x 10*
Estimated
emissions
(tons/yr)
26.9 x 10*
0.7 x 10«
7.5 x 10«
14.1 x 10"
2.1 xlO8
1.6x10*
52.9 x 10«
4.2 x 108
500xlO«
5900 x I0«
650 xlO6
Emissions
as nitrogen
(tons/yr)
8.2 x 10s
0.2xlO«
2.3 x 10"
4.3 x 10°
0.6 x I0«
0.5 x 10"
16.1 xlO«
3.5 x 10«
150 x 10"
4900 x 10s
410 x 10"
-48-
-------
NOT O.'nrr . ,,,
V w '-> I il I ; i /^
""' '
5.1.2.2 Mobile Sources—The chemical species, nitr'ateV'
end-product of a chain of chemical reactions which begins in the
combustion chamber. At the high temperatures of combustion, elemental
nitrogen is disassociated and combined with available oxygen to form
nitric oxide. It is the gas, nitric oxide, that is primarily emitted
from the exhaust pipe. Nitric oxide is oxidized to nitrogen dioxide in
the ambient air. Further, photochemical reactions convert the nitrogen
dioxide to acid and nitric acid, which subsequently react to
form various nitrate salts. Both inorganic and organic nitrates are
formed, with peroxyacetyl nitrate being the most abundant of the latter
and one of the main eye irritants of photochemical smog.
There is not much evidence for direct emission of nitrate in
vehicle exhaust in either particulate or gaseous forms. In one case
of measurement of nitrate in raw, undiluted vehicle exhaust by Lee,
C Q
et. al., nitrate was found in concentrations up to 700 yg/m . This
nitrate was found to exist as 99 percent in particles of less than 0.5 ym
diameter. They also found that in chamber studies of vehicle exhaust,
upon irradiation with ultra-violet light, there was a many-fold increase
in the nitrate present with a corresponding decrease in NO and increase
in NOp over the amounts of nonirradiated raw exhaust.
The large mobile source portion of the nitric oxide is emitted
in a city-wide pattern. The NO emission levels in urban traffic areas
3\
have been found to range from 800 to 3000 ppm. . This produces an
urban plume of nitrogen oxides and nitrates which extends out into
the rural areas. The power plants discharge their nitrogen oxides into
-51-
-------
a plume which is more concentrated but smaller in extent. The power
plant plume may eventually merge with the urban area plume and become
indistinguishable from it. The time and distance required for merging
will be a function of such factors as height of stacks, location of
plants, and prevailing winds.
Although the nitrogen compounds are initially gaseous, they end
up being removed in particles and droplets, mainly as nitrate solutions.
The tonnage of nitrate ion (N03~) produced annually in metropolitan
areas can be estimated by multiplying the tonnage of emitted nitric
oxide by a factor of two—the weight increase on oxidation. Emissions
of nitric oxide in Los Angeles County, for example, will yield about
Q
seven hundred thousand tons of nitrate ion per year. This far outweighs
any nitrates emitted directly to the air. The nitrate is formed in the
air mainly as peroxyacetyl nitrate vapor, with a somewhat lesser amount,
not fully measured as yet, formed as inorganic nitrate salts in the
aerosols.
In large cities about two-thirds of the nitric oxide comes from
mobile sources and one-third from stationary sources, mainly
power plants. The proportions of mobile source and stationary source
nitrogen oxides may differ slightly from city to city, depending on
the location of power generating facilities.
-52-
-------
HKFT
5.2 CONCENTRATIONS
5.2.1 Inorganic Nitrates in Air
As one of the routine activities of the National Air Surveillance
Network (NASN), suspended particulate matter is removed from ambient
air by filtering through glass fiber filters. The material
filtered out of the air by this technique has been analyzed for its
content of various substances; among them, nitrate ion, thus
providing data on "concentration of nitrates" in the air.
Table 5.2 contains the annual average values for those urban NASN
q
sites where all data were available from 1966 to 1970. The sites have
been arranged in order of decreasing 5 year average concentration.
Table 5.3 similarly arranged, presents data for nonurban NASN sites
where all 5 years' data were available.
Simple inspection of these tables will reveal the following:
(1) All entries substantially exceed the minimum detectable
value as listed in the data bank. (0.06 yg/m ).
(2) Smallest NASN nonurban values occur in the western third of
the nation.
(3) Urban NASN values are generally higher than nonurban. The
higher urban values are widely scattered over the nation.
(4) These data are inadequate for demonstrating long-term
temporal trends.
-53-
-------
UKMM
DO NOT QUOTE OR CiTI
TABLE 5. 2 URBAN NASN NITRATE ANNUAL AVERAGES, 1966-1970
(ug/m3)
15-year
h966 1967 1968 1969 1970 Average
Los Angeles, CA 5.20 4.25 9.22 6.64 12.55 7.57
Portland, OR 3.16 3.56 4.66 3.23 4.81 3.88
San Diego, CA 2.72 3.09 4.01 3.92 5.35 3.82
Detroit, MI 3.71 2.75 4.69 3.32 4.44 3.78
Oakland, CA 2.27 4.74 3.02 3.36 3.92 3.46
Houston, TX 2.64 3.20 3.62 3.33 3.55 3.27
New Orleans, LA 2.19 3.07 4.09 2.68 3.85 3.12
E. Chicago, IN 2.93 2.21 3.20 2.55 4.39 3.06
Milwaukee, WI 2.18 2.42 4.65 2.74 3.32 3.06
Chattanooga, TN 2.03 2.09 4.12 3.98 3.04 3.05
San Francisco, CA 1.96 3.47 3.90 2.21 2.75 2.86
Columbus, OH 2.07 1.87 2.55 3.76 3.58 2.77
Cincinnati, OH 2.13 1.74 2.51 3.55 3.47 2.68
Youngstown, OH 2.35 2.10 2.79 2.18 3.68 2.62
Dayton, OH 2.18 1.66 2.38 3.61 3.07 2.58
Baltimore, MD 2.40 2.13 2.82 1.81 3.71 2.57
Indianapolis, IN 2.13 2.54 2.52 1.68 3.86 2.55
Denver, CO 1.44 1.91 2.30 3.42 3.09 2.43
Toledo, OH 1.74 2.28 2.85 1.70 3.55 2.42
San Antonio, TX 1.87 2.71 2.95 1.78 2.02 2.27
Cleveland, OH 1.34 1.61 2.64 2.69 2.93 2.24
Memphis, TN 1.94 1.88 2.84 1.16 2.78 2.12
-54-
-------
(TABLE 5.2 CON'T)
DRAFT
00 NOT QUOTE OR CITE
1966
Nashville, TN
Chicago, IL
Des Moines, I A
Atlanta, 6A
Omaha, NE
Norfolk, VA
Tulsa, OK
Salt Lake City, UT
Wilmington, DE
Little Rock, AR
Oklahoma City, OK
Wichita, KS
Tucson, AZ
St. Paul, MN
Newark, NJ
Charlotte, NC
Covington, KY
Minneapolis, MN
Jersey City, NJ
Burlington Co, NJ
Seattle, WA
Albuquerque, NM
Charleston, WV
2.
1.
1.
1.
1.
1.
1.
1.
1.
1.
1.
1.
1.
1.
1.
2.
1.
1.
1.
1.
1.
1.
1.
15
69
72
74
83
66
48
36
54
83
83
62
31
49
39
48
40
00
26
36
41
29
09
1967
1.
1.
1.
2.
1.
1.
1.
1.
2.
1.
1.
1.
1.
1.
1.
1.
1.
1.
1.
0.
1.
1.
1.
95
51
99
27
70
93
36
04
05
81
65
37
67
90
87
81
29
68
44
92
73
81
48
1968
2.93
2.68
1.77
2.09
2.15
2.24
1.73
2.54
2.20
2.23
1.61
1.76
1.64
1.95
1.65
1.61
1.60
1.40
2.00
1.28
1.69
1.53
1.31
1969
1.
1.
1.
1.
1.
1.
2.
2.
1.
1.
1.
1.
2.
1.
1.
1.
1.
1.
1.
2.
1.
1.
1.
40
99
83
59
96
91
88
39
26
38
61
21
18
10
26
28
41
49
30
08
29
69
31
1
2
2
3
2
2
2
2
2
2
2
2
3
2
2
2
1
3
2
2
2
1
1
1
970
.19
.65
.15
.61
.50
.34
.39
.31
.46
.27
.69
.39
.46
.56
.76
.67
.02
.76
.13
.48
.87
.44
.81
5-year
Average
2.
2.
2.
2.
2.
2.
1.
1.
1.
1.
1.
1.
1.
1.
1.
1.
1.
1.
1.
1.
1.
1.
1.
12
10
09
06
03
02
97
93
90
89
88
87
85
80
79
77
74
67
63
62
60
55
40
-55-
-------
DRAFT
DO NOT QUOTE OR CITE
(TABLE 5.2 CON'T) UU " X
1966
Hartford, CO
Glassboro, NJ
Providence, RI
Boise City, ID
New Haven, CN
Ponce, PR
Cheyenne, WY
Guayanilla, PR
Honolulu, HI
Concord, NH
1.
1.
1.
0.
1.
0.
0.
0.
0.
0.
21
12
28
99
06
39
67
28
97
26
1967
2.71
0.96
1.13
0.97
1.00
0.78
0.73
0.55
0.71
0.46
1968
1.
1.
2.
1.
1.
1.
0.
0.
0.
0.
11
57
26
05
29
67
77
93
63
73
1969
0.
1.
1.
1.
0.
0.
1.
0.
0.
0.
67
12
02
19
58
81
04
72
70
37
1970
1.
2.
1.
1.
1.
1.
0.
1.
0.
1.
23
16
20
33
38
00
76
18
51
04
5-year
Average
1.
1.
1.
1.
1.
0.
0.
0.
0.
0.
39
39
38
11
06
93
79 '
73
70
57
Overall Average
1.759 1.918 2.435 2.023 2.826 2.191
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QUOTE
TABLE 5.3 NONURBAN NASN NITRATE ANNUAL AVERAGES, 1966-1970
DO NOT QUOTE OR CITE
(ug/m3)
1966
Cape Hatteras, NC
Matagorda Co. , TX
Cherokee, Co. , OK
Acadia Nat. Park, ME
Jefferson Co. , NY
Monroe Co. , IN
Shannon Co. , MO
Park Co. , IN
Montgomery Co. , AR
Clarion Co. , PA
Thomas Co. , NE
Grand Canyon, AZ
Humboldt Co. , CA
Orange CO. , VT
Coos Co. , NH
Curry Co. , OR
White Pine Co., NV
1.
0.
0.
0.
0.
0.
0.
0.
0.
0.
0.
0.
0.
0.
0.
0.
0.
09
60
63
97
61
29
54
82
56
24
55
42
33
44
52
16
27
1967
1.
0.
0.
0.
0.
0.
0.
0.
0.
0.
0.
0.
0.
0.
0.
0.
0.
23
97
61
93
46
44
80
49
47
37
35
53
54
32
20
28
20
1968
1.23
0.98
0.83
0.81
1.09
1.22
1.13
0.74
0.86
1.00
0.41
0.56
0.59
0.43
0.50
0.43
0.32
1969
1.
1.
0.
0.
0.
0.
0.
0.
0.
0.
0.
0.
0.
0.
0.
0.
0.
16
42
83
64
64
46
28
36
68
39
78
50
54
37
23
23
47
1970
3.
1.
1.
1.
1.
1.
0.
1.
0.
1.
0.
0.
0.
0.
0.
0.
0.
34
39
98
40
60
48
96
22
95
13
74
65
61
71
76
58
39
5-year
Average
1.
1.
0.
0.
0.
0.
0.
0.
0.
0.
0.
0.
0.
0.
0.
0.
0.
61
07
98
95
88
78
74
73
70
63
57
53
52
45
44
34
33
Overall Average 0.533 0.532 0.760 0.619 1.109 0.718
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A word of caution 1s appropriate. The overall averages in the
table apply to the table entries which were chosen for data completeness
over the 5 year period. Therefore, the overall averages are not identi-
cal with the "best" estimate of national annual averages. However, the
"best" estimate would not be expected to be very different.
Temporal variations in the material collected can be assessed on a
national basis but only for certain minimum time scales. In NASN opera-
tions one sample is the amount of material collected on a filter after
air is drawn through that filter for 24 hours. One filter is
exposed thusly at each site in each 12 day period.
As a consequence, it is fundamentally impossible to derive any
information about diurnal variations from NASN data. Moreover, the
frequency with which samples are collected was designed to provide
reasonably characterized quarterly and long-term averages not short-
term averages. Thus for any given site, values for times shorter than
a calendar quarter are not well characterized by NASN data.
In an attempt to provide information on a finer time scale,
monthly national averages have been plotted for urban NASN stations in
Figure 5.1 and for nonurban stations in Figure 5.2. It is obvious in
both figures that fluctuations are sizable. This is due in part at
least to the infrequent sampling schedule (once every 12 days).
The solid line in these figures smooths the data so that a month-
to-month or seasonal change in these averages can be seen more easily.
The fluctuations remain sizable. It should be pointed out, however,
that the method of analysis was changed after all samples from calendar
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o. 2.:
V
o
o
"57
r 61
60
66 "; 67 bS
YEAR
5.1. MONTHLY NATIONAL VALUES FOR NITRATE URBAN NASN SITES
co
E
cr/
UJ
o
•z.
o
2-i
,A A IV
'; \- ' V
5g—bT^~65 ' 66 ' *•' 0« ' -
YEAR
5.2 MONTHLY NATIONAL VALUES FOR NITRATE NONURBAN NASN SITES
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year 1965 were analyzed. This change in method is documented in the
9
National Aerometric Data Bank. It is not clear at this time however,
if any of the differences between pre and post 1966 values can be
attributed to this fact.
It may be of some interest to attempt to infer something about
the nitrate content of collected particles as a function of the
particle size. Data are very limited. However, analysis of material collected
10
by impactors in the Cincinnati area in 1965' and in Riverside,
California in 1968 indicates that most of the nitrate occurs in the
smaller particles -- two micrometers and smaller.
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5.3.2 Other Nitrate Concentration Measurements
The measurement of atmospheric nitrates has been fraught
with error. Satisfactory methods for routine use on particulate
nitrate, PAN, nitric acid, nitrous acid, nitrogen dioxide, and
nitrogen pentoxide are not yet available. Gas chromatographic
and infrared measurements of PAN are the most reliable of the
various nitrate measurements that have been reported.
Stephens and Burleson reported 0.034 parts-per-million (183
micrograms per cubic meter) of PAN in Riverside, California at a
12
time when ozone was about 0.5 parts-per-million. Hanst, et.
al., reported 0.050 parts-per-million (270 micrograms per cubic
meter) of PAN in Pasadena, California, when ozone was about 0.6
23
parts-per-million. Oh irradiation of Los Angeles morning air
in plastic bags, Kopczynski, et_ al^ , produced about 0.050 parts-
per-million (270 micrograms per cubic meter) of PAN along with
U
about 0.4 parts-per-million of ozone. In 1965, Mayrsohn and
Brooks reported 0.214 parts-per-million (1200 micro-
14.
grams per cubic meter) of PAN in Los Angeles. Tingey and Hill
found 0.054 parts-per-million (290 micrograms per cubic meter)
1&
of PAN in the air near Salt Lake City, Utah.;
Nitric acid vapor has not yet been measured in the polluted air of
the lower atmosphere, although it has been detected in the 5 ra-
4
tosphere. EPA investigators using long path infrared spectro-
scopy at Pasadena, California, reported no evidence of nitric
acid vapor at concentrations higher than 0.010 parts-per-million
on two days when ozone concentrations reached maxima of 0.59 parts-
per-million and 0.68 parts-per-million. From consideration of
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the chemistry, it is almost certain that nitric acid is formed
in the air. It is obvious that attempts to measure the nitric
acid concentration must be pursued further.
Detection of alkyl nitrates in the polluted air has been re-
ported. The detected amounts fall !in the range of one to two
parts-per-billion (3 to 6 microgramsrper cubic meter). Since
these alkyl nitrates are formed concurrently with PAN but at
much lower concentrations, and since they are much less reactive
than PAN, they need not be given any independent detailed con-
sideration.
Inorganic nitrates are captured and measured on filters. The
amounts detected and reported have had wide variation. The
Environmental Protection Agency CHESS program has reported 24-
hour average values in the vicinity of 5 micrograms per cubic
I -I I Q
meter. Novakov et aj_., from a four-hour average measurement
made in Los Angeles smog, have reported nitrate at less than one
19
microgram per cubic meter. In contrast, Holmes, et al., of the
California Air Resources Board measured between 100 and 200 micro-
grams of nitrate per cubic meter of air on a smoggy summer day
in Los Angeles. From these wide discrepancies, it is apparent
that future program emphasis must be placed on obtaining more
accurate measurement of atmospheric particulate nitrates.
Nitrate measurements made at CHESS stations show large temporal
20
and spatial variations. In Chattanogga, Tennessee, observed 24-
hour average concentrations varied from below the detectable level to 107.3
3 3
yg/m , and the monthly arithmetic mean varied from 0.3 to 24.9 yg/m .
The overall urban averages varied from 1.3 to 7.2 yg/m .
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IP
TH ™ DO NOT QUOTE OS CI71
Table 5.4
SUMMARY OF NITRATE CONCENTRATIONS BASED ON ChtsS DATA
(1969-1972) yg/m3 .
24 Hr Obs
Monthly mean
Max. Min. Max. Min.
No
OBS
Overall
(nontniy
mean
North Carolina 28.0
(Charlotte and
Greensboro)
Alabama
(Birmingham)
New York
(Riverhead, Queens,
and Bronx)
Utah 27.1
(Ogden, Salt
Lake Cityj Kearns
and Magna)
Tennessee 107.3
(Chattanooga)
New Jersey 11.7
(Ridgewood, Fairlawn,
Matawan, Canteret,
and Elizabeth)
California 87.5
(Vista, Santa Monica,
Thousand Oaks, Anaheim,
Garden Grove, 61en-
dora, and Covina).
4.0 0.1
13.2
23.5
*
*
4.3
5.2
0.6
0.7
6.9 0.8
24.9 0.3
5.1 0.4
28.7 1.9
5784
3660
1978
2572
1732
1.3
1.8
2.4
2.7
4736 3.0
1727 2.5
7.2
*
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5.3 TRANSFORMATION AND TRANSPORT MECHANISMS AT rtHftTF OR CITE
5.3.1 Natural Mechanisms
5.3.1.1 ffienrical —The nitric oxide is converted to nitrogen dioxide by reaction
with various forms of atmospheric oxygen or oxygenated free
21
radicals. The reaction with atomic oxygen:
NO + 0 (+M) = N02 (+M) (1)
is fast and evolves energy, but it is not a major path for N0?
formation. As the oxygen atoms are formed in the air, they
react mainly with molecular oxygen to form ozone.
The reaction of NO with molecular oxygen is a third order
reaction:
2ND + 02 = 2N02 (2^
and as such proceeds only very slowly when the NO is at low
concentration. The reaction is too slow to be of importance in
the ambient polluted air, but it may take place during the dilution
of NO-containing effluents. It is partly responsible for the
starting concentration of N02 in the morning air of cities.
The reaction of NO with ozone:
NO + 03 = N02 + 02 (3)
is extremely fast. It is probably the major reaction of NO when
it is emitted into ozone-containing atmospheres.
N02 in the atmosphere absorbs sunlight and dissociates into
NO and oxygen atoms. The oxygen atoms combine with oxygen molecules
forming ozone. The ozone then reacts with NO, returning it to N02.
This is summarized in the following three reactions:
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N02 + hv = NO + 0 (4)
0 + 0- (+M) -»• 0. OM) (5)
Z o
0 + NO -»• N02 + 02 (6)
Although the photodissociation takes place within one or two
minutes after the exposure of the NO- to the sunlight, the steady
state concentration of NO- is very little changed by the cycle
because the reverse reactions (5) and (6) are so fast. If the
air sample does not contain impurities other than NO-, the photolysis
may go on for an extremely long time with no net change in the com-
position. This is described by the equilibrium
N02 + 02 £v NO + 03 (7)
When other compounds are present, the composition is changed,
but not in the direction one might normally expect. In the pre-
sence of impurities, one might logically expect the photodis-
sociation to cause a decrease in the concentration of nitrogen
dioxide, but instead it causes an increase.
The photochemical production of N02 has been studied extensively
and is accounted for by the intervention of hydrocarbons and other
pollutants in the NO- photolysis cycle.
The hydrocarbons exert their influence by way of free radicals,
R, which are formed by interactions with oxygen atoms, ozone, OH
radicals and other energetic atmospheric species. Hydrogen
abstraction by oxygen atoms is an example.
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RH + 0 = R + OH (8)
The radical R will then add oxygen: ;
R + 02 = R02 (9)
The radical RO.- is capable of oxidizing NO to N0?:
R02 + NO = RO + N02 (10)
The sequence of reactions in the example has regenerated the N0?
which was photo-dissociated, but in so doing has created the
free radicals RO and OH. Since these active species can also
attack hydrocarbons, a branching reaction sequence has been set
up which allows photolysis of N0_ to create additional N0« out of
the reservoir of NO.
Water vapor enters the nitrogen oxide reactions through the
formation of nitrous acid:
NO + N02 + H2<3 = 2HN02 (11)
The nitrous acid molecules are photolyzed by sunlight:
HN02 + hv = OH + NO (12)
The OH radicals may attack hydrocarbons with resultant NO
oxidation:
OH + RH = R + H20 (13)
R + 02 = R02 (14)
R02 + NO = N02 + RO (15)
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LS|\f II I
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The result of the above sequence is conversion of NO to N0? with
additional production of the active species, RO.
Carbon monoxide can also influence the NCL photolysis cycle.
The degree of involvement of carbon monoxide is not nearly as
great as the degree of involvement of hydrocarbons, and the
CO does not play an important overall role in atmospheric nitrate
formation. However, the carbon monoxide involvement in the
atmospheric chemistry does control the steady-state carbon monoxide
concentration in the atmosphere. The key reaction is the oxidation
of carbon monoxide by OH radicals:
OH + CO = H + C02 (16)
The hydrogen combines with oxygen:
H + 02 = H02 (17)
The H02 radicals oxidize NO:
H02 + NO = OH + N02 (18)
When the overall photochemical conversion of NO to NO- in the
atmosphere approaches its end point, the fast reaction between
NO and ozone dies out, and ozone is able to begin to accumulate
in the gaseous mixture. At that point the reaction mixture enters
the stage of nitrate formation. Two types of nitrates begin to form:
inorganic nitrates, derived from the reaction of NO- with ozone,
and organic nitrates, derived from the reaction of N02 with peroxide
free radicals.
Inorganic nitrate formation begins with the reaction of N0_
and ozone, yielding NO,:
N02 + 03 = N03 + 02 (19)
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The nitrogen trioxide couples with NO- to yield nitrogen pentoxide:
N03 + N02 = N205 (20)
At an ozone level of 0.3 parts-per-million, which is typical
of photochemical smog, the half- life of an NO. molecule will be
about thirty minutes. N-O- is the anhydride of nitric acid;
it hydrolyzes readily, but apparently only at 'the surface of
atmospheric particles .
- 2HN03 (£) (21)
If there are other solutes, the nitric acid may be immediately
neutralized, perhaps with the release of volatile products.
Sodium chloride will release HC1 gas:
HN03 + NaCl = NaN03 + HC1 (g) (22)
Calcium carbonate will release CO- gas:
CaC03 + 2HN03 = Ca(N03)2 + H20 + C02 (g) (23)
Ammonium hydroxide will yield ammonium nitrate:
NH4OH + HN03 = NH4N03 + H20 (24)
There are two types of evidence for these neutralization reactions.
First, there are the analyses of collected aerosols, which show the
22
presence of nitrate salts, especially ammonium nitrate. Second,
there are the repeated findings that nitric acid vapor does not
23
exist in the polluted air of cities at detectable levels.
Organic nitrates are formed by the combination of nitrogen di-
oxide with oxygenated radicals. Methoxy radical yields methyl ni-
trate :
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CH30 + N02 = CH3N03 (25)
Peroxyacetate radicals yield peroxyacetyl nitrate (PAN):
CH3C (0)00 + N02 = CH3C (0)OON02 (26)
PAN is the notorious "compound X" of photochemical smog, which
was discovered in the laboratory in 1955 and first measured in
24,25
the atmosphere in 195b. Measurements show it to be the pre-
dominant nitrate in the air. It has been shown to be eye-irritating,
plant-damaging and poisonous to animals.
Laboratory simulations of the atmospheric photochemistry have
shown how the yields of the organic nitrate products vary with
concentration. When the pollutant concentrations are on the order
of a few parts-per-mi11ion the reactions produce comparable amounts
of methyl nitrate and peroxyacetyl nitrate. When the pollutant
concentrations are reduced to the level of a few tenths of a part-
per-million, typical of the real atmosphere, the alkyl nitrate
product drops to a very low value, and the peroxyacetyl nitrate
becomes by far the major product. This behavior is confirmed by
the atmospheric measurements which show PAN to exceed alkyl nitrate
by at least a factor of ten.
It seems probable that PAN in the air is slowly taken up by aerosol
droplets and hydrolyzed to nitrite in solution, which will then be
oxidized to nitrate. Thus, the ratio of inorganic nitrate to organic
nitrate in a polluted air mass would be expected to increase slowly
during transport of the air mass away from an urban center.
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5.3.1.2 Transport and Removal—Formation, transport, and removal of
nitrate air pollutants all proceed simultaneously. A consequence of
the complexity of the many concurrent processes is that atmospheric
measurements have never succeeded in tracing the path of all of the
oxidized nitrogen known to be emitted into the air. Atmospheric
analyses have generally revealed oxidation products equivalent to only
about half of the suspected emissions. The recent EPA spectroscopic
study of the air in Pasadena again resulted only in the detection of
some of the nitrogen compounds. The conclusion must be drawn
from the non-detection of the oxidized nitrogen is that major removal
processes are taking place at ground level. It seems likely that
NOp, NpO,- , HNOo, HNOo. PAN, and particulate nitrates are taken up
by such surfaces as leaves, soil, and man made materials at a
greater rate than has been heretofore recognized. It is necessary
that this subject be explored further in laboratory and field tests.
With appropriate atmospheric conditions, photochemical reaction
in urban centers can reach the stage of rapid peroxyacyl nitrate
formation by midday. By mid-afternoon, the PANs will substantially
exceed the inorganic nitrates in concentration. At night, however,
PAN formation is at a standstill and PAN decomposition is proceeding.
At the same time inorganic nitrate will have assumed the major proportion
of the total nitrate mass. This mass will of course be much reduced from
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the previous afternoon because of removal at surfaces.
At suburban locations downwind from urban centers, the di-
urnal cycle should be less pronounced, the total mass of nitrate
should be smaller, and the proportion of inorganic nitrate should
be greater. Comparison of Measurements at Riverside, California,
with measurements at Los Angeles indicate this to be the case.
Very few of such measurements are available, however, and the
conclusions must be considered tentative. A larger field measure-
ment program on the organic and inorganic nitrates is very much
needed.
Rain-out and wash-out are two pollutant removal processes
associated with precipitation. Rain-out is the process in-
volving growth of rain drops on nuclei containing particulate
pollutants. When the rain drops fall, the pollutants go with them.
Wash-out is the process by which the rain drops pick-up additional
pollutants on the way down. The drops moving through the air scrub
out pollution as they go. These processes will cleanse the air
anywhere, at any time of the year providing that it rains. It
is therefore important to keep in mind that photochemical smog
and its high incidence of nitrate pollution usually occur at dry,
hot times of the year when rain-out and wash-out are not at work.
In Los Angeles, it rains very little in the smoggy summer months.
Dry removal processes must therefore play a primary role. Two
such dry processes are recognized. They can be termed fall-out
and comb-out.
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Fall-out is simply the downward migration of particles in the
gravitational field. Particles grow by accretion and agglomeration;
as they get heavier, their acceleration under gravity can no longer
be countered by their Brownian motion, and the particles settle out.
Comb-out is the removal of particles and gases by leaves, grass,
soil, buildings, and other surfaces. This removal results from
the air sweeping across and through the surfaces. Comb-out pro-
cesses are sometimes simply referred to as dry deposition. The
natural turbulance of the lower atmosphere keeps the air in a
well-stirred condition and maintains a macroscopically uniform
concentration of both gaseous and particulate pollutants. At
the same time the turbulence is continually bringing the pollutants
and the surfaces into contact. Aqueous fine particles and acid
vapors have a high probability of sticking to the surface at the
first collision. Inert gases will bounce off. The "sticking-factor"
of the various pollutant species needs to be established in lab-
oratory tests.
The net result of the turbulence and a high sticking-factor
is that comb-out can cleanse a one centimeter thick layer of air
every second, or a 36 meter thick layer every hour. These figures
illustrate the major role of comb-out in atmospheric cleansing and they
emphasize the need for further study of the dry deposition phenomena.
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5.4 REFERENCES DO NOT QUOTE OR CiTE
.;!• Robinson, E., and R. C. Robbins, "Gaseous Nitrogen Compounds
Pollutants from Urban and Natural Sources", Journal of the
Air Pollution Control Association 20, 303 (1970).
2. Junge, C., and J. Hahn, "N.,0 Measurements in the North Atlantic",
J. Geophys. Res., 76, 8143^(1971).
3. Commoner, B., "Threats to the Integrity of the Nitrogen Cycle:
Nitrogen Compounds in Soil, Water, Atmosphere and Precipitation1,1,
in "Global Effects of Environmental Pollution", S. F. Singer,
Editor, Springer-Verlag, New York, 1970, p. 72.
4. Murcray, D. G., T. G. Kyle, F. H. Murcray, and W. J. Williams,
J. Opt Soc. Amer., 59, 1131 (1969) .
5,OAQPS Data File of Nationwide Emissions 1971, National Air Data
Branch, Monitoring and Data Analysis Division, May 1973.
6. "Concentration and Particle Size Distribution of Particalate
Emissions in Automobile Exhaust," Robert E. Lee, Jr., Ronald
Patterson, Walter Crider, and Jack Wagman, Atmospheric
Environment. Vol. 5_, pp. 225-237, 1971.
7. Air Pollution. Vol. III. A. C. Stern, Editor, Chapter 33,
Mobile Combustion Sources, R. W. Hum, 1968.
8. Los Angeles Air Pollution Control District - "Air Quality
Profile of Air Contaminant Emissions" Los Angeles County -
January, 1971.
g; Data stored in the National Aerometric Data Bank, U.S. Environmental
Protection Agency, Research Triangle Park, North Carolina, 1974.
QK-: Lee, R.E., Jr., R.K. Patterson. Size determination of Atmospheric
Phosphate, Nitrate, Chloride, and Ammonium Particulate in Several
Urban Areas. Atmospheric Environment. 3:249-255, March 1969.
H.; Lundgren, D.A. Atmospheric Aerosol Composition and Concentration as
a Function of Particle Size and of Time. J. Air Poll. Contr Assoc
20:603-608, September 1970.
12- Stephens, E.R., and F.R. Burleson. Distribution of Light Hydro-
carbons in Ambient Air. J. Air Pollution Control Assn. 19:929. 1969.
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13. Kopczynski, S. L., W. A. Lonneman, F. D. Sutterfield, and
P. E. Darley, "Photochemistry of Atmospheric Samples in
Los Angeles", Environmental Science and Technology 6, 343 (1972).
•V$? Mayrsohn, H. and C. Brooks, "The Analysis of PAN by Electron
Capture Gas Chromatography", Western Regional Meeting of the
ACS (Nov. 18, 1965).
15. Tingey, D. T., and A. C. Hill, "The Occurrence of Photochemical
Phytotoxicants in the Salt Lake Valley", UtahiAcad. Proc. 44
(1), 387 (1967).
16. Battelle Memorial Institute, Interim Report to EPA on the
Study of Nitrogenous Compounds in the Atmosphere, 1974.
17. Health Consequences of Sulfur Oxides: A Report from CHESS,
1970-1971. EPA-650/1-74-004. May 1974.
18. Novakov, T., P. K. Mueller, A. E. Alcocer, and J. W. Otvos,
"Chemical States of Nitrogen and Sulfur by Photoelectron Spectro-
scopy", in Aerosols and Atmospheric Chemistry, G. M. Hidy,
Ed., Academic Press, New York, 1972, P. 285.
19. Data not yet published; private communication from John Holmes
to Philip Hanst, April, 1974.
20. Unpublished Data. Health Effects Laboratory, Environmental
Protection Agency, 1974.
21. Leighton, P. A., Photochemistry of Air Pollution, Academic
Press, New York (1961) p. 158.
22. Stephens, E. R. and M. Price, "Comparison of Synthetic and
Smog Aerosols", in Aerosols and Atmospheric Chemistry G. M.
Hidy, Ed., Academic Press, New York (1972) p. 167.
23. Hanst, P. L., W. E. Wilson, R. K. Patterson, B. W. Gay, Jr.,
L. W. Chancy, and C. S. Burton, "A Spectroscopic Study of
Pasadena Smog", EPA, NERC, preprint, January 1974.
24. Stephens, E. R., W. E. Scott, P. L. Hanst, and R. C. Doerr,
"Recent Developments in the Study of the Organic Chemistry of
the Atmosphere", J. Air Pollution Control Assn., £, 159 (1966).
25. scott, E. R. Stephens, P. L. Hanst, and R. C. Doerr, "Further
Developments in the Chemistry of the Atmosphere", Proc. A. P. I.,
37, (III), 171 (1957).
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6. EFFECTS j|g OR CITE
6.1 HEALTH AFFECTS OF NITRATES AND RELATED COMPOUNDS IN THE AIR -
The organic nitrate peroxyacetyl nitrate has been shown to be a
potent eye irritant and may be largely responsible for this effect in
Los Angeles. There is limited evidence that exposure to inorganic
nitrates may induce asthma attacks in some asthmatics under certain
climatic conditions. Except in persons exposed to very high density
automobile traffic or to unusual point sources of an industrial nature,
airborne nitrates probably play an insignificant role in even subclinical
methemoglobinemia. No role for airborne nitrates in carcinogenesis is
known at present. The contribution of airborne nitrates to total
intake of nitrates is very small.
What is known about the health effects of nitrates is described
' below; the extent of what is not known is so large as to be
indescribable.
6.1.1 Route Of. Entry of_ Nitrates into the Body
The ubitiquity of nitrates in man's environment makes
possible his exposure to these substances through a variety of routes of
entry into the body. Water and a variety of foods have long been
recognized as principal routes of entry of nitrates and nitrites into the
body. The subject of accumulation of nitrates in food and water and
1 3
health implications has been extensively reviewed.
The recent report of positive correlations between airborne suspended
nitrates and acute respiratory symptoms in urban populations has- stimulated
interest in the contribution of inhaled nitrates to the total nitrate body-
burden and their potential hazard to health. A number of relevant
questions are posed by these preliminary results and insights into their
explanation are limited by the paucity of research into the inhalation
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toxicology of nitrates.
Nitrate and nitrite salts are readily [soluble in water and serum
containing systems. These salts are rapidly absorbed when ingested. '
It seems likely that they would be readily absorbed through the
respiratory tract aftfcp deposition in the alveolar region of the
lung, through the epithelium of the airways or through the stomach
after nitrate particles are cleared from the airway epithelium
and swallowed. An important question arises about the amount of nitrate
that might be absorbed through the respiratory tract compaEed with
ingestion. While the nitrate dose acquired through the respiratory
route of entry may contribute to methemoglobinemia after absorption
into the blood, no rationale appears to exist which indicates that
the respiratory symptoms which have recently been correlated with
airborne nitrates are related to syndromes which have methemoglobinemia
as their underlying pathophysiological basis.
A calculation of an inhaled dose of nitrate is made on the basis of the
following assumptions: (1) The positive correlation between airborne
nitrates and increased incidence of respiratory disease appears at about
3-5 micrograms per cubic meter of air. (2) It is assumed the particle
size of this suspended nitrate is submicronic metric (0.5-1.0).
The total respiratory deposition rage of particulates with an
8
aerodynamic diameter of 0.5-1.0 pm is approximately 50 to 70 percent.
All of this material will be absorbed through either the respiratory
o
tract or the stomach. The respiratory volume of an adult male at
g
rest may range between 5 and 7 liters per minute. With 6 liters as
a ventilatory estimate, approximately 4320 liters Cor 4.3 m y of air
are inhaled during a 12-hour day of exposure.
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The inhaled nitrate dose for this exposure is calculated to
be approximately 21.5 micrograms. A portion of the inhaled dose would
not, in fact, be deposited (perhaps 30 to 50 percent). Thus, the inhaled nitrate
dose would be substantially less than 1 percent of the body burden if water alone
is considered as the only other source of nitrate. This- percentage
is substantially less when total dietary nitrate intake fs- considered.
It appears unlikely that inhaled nitrates would contribute substantially
to nitrate-nitrite body burdens and toxicity resulting from
methemoqlobinemia and its consequent effects. However, if the preliminary
correlations between airborne nitrates and respiratory symptoms
are correct, and are causally associated, these symptoms appear to
reflect direct irritation to the respiratory tract and suggest that
airborne nitrates may be an important respiratory irritant.
The principle sources of suspended nitrate particulates in ambient
air are i automobile exhausts*combustion of fossil fuels in
stationary.sources, and the manufacture and use of chemical
fertilizers.
10
Studies conducted by Lee, et al, on the distribution of particulate
emissions in automobile exhaust showed large quantities of particulate
nitrate were formed by irradiation of diluted exhaust. The addition
of 0.5 ppm sulfur dioxide to diluted exhaust increased appreciably
both the concentration of nitrate and the sizes of nitrate containing
particles. Irradiation produced a shift of nitrate to smaller particle
sizes with and without sulfur dioxide. The size distribution of nitrate
showed a shift to smaller particle sizes with increasing exhaust
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temperature and operating speed. At 50 mph, 99 percent of the nitrate
was contained in particles of less than 0.5 ym diameters. The
predominantly submicron particle size of these components in
automobile exhaust is consistent with data obtained in air sampling,
indicating mass median diameters in the range of
0.2 - 0.6 ym. The nitrate particulate component is in the respirable
range which can enhance deep lung penetration and deposition. Those
populations who work or reside in areas of high auto traffic density
would be at particular risk to exposure to suspended nitrates.
Cigarette smoke may also be a source of NOx, including nitrate
and nitrosamine exposure. The question of the role of nitrosamines
in tobacco smoke in the induction of lung or other cancer has recently
been reviewed. Beyond the fact that nitrosamines exist in
cigarette smoke, nothing is known. There is a wide variety of
other known and suspected carcinogens in tobacco smoke also. Organic
nitrates, such as methyl nitrite and the nitrohydrocarbons occur in
12-14
cigarette smoke in relatively large quantities. In the case of
nitrohydrocarbons, the amounts may reach microgram per
cigarette-
Since tobacco smoking is an important covariate in air pollution
studies, the effects of nitrogenous compounds derived from smoke must
also be considered.
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6.1.2. Epidemiologic Studies of Health Effects of Airborne Nitrates
6.1.2.1 Airborne Inorganic Nitrates
6.1.2.1.1 Respiratory effects—Epidemiologic data on respiratory
or other effects of exposure to
inorganic suspended nitrates in ambient air are very limited.
In a recent unpublished study conducted as part of the EPA CHESS
program in the New York-New Jersey Metropolitan area.increased asthmatic
attacks were significantly associated with elevated levels of suspended
15
nitrates in six of seven communities. The observed excess risk of
asthmatic attack with elevated levels of suspended nitrates occurred
only when temperatures were 50°F or above. The estimated threshold for these
•3 O
effects ranged from 2.16 yg/m to 7.63 yg/m .
In another study conducted by the same group in two southeast
16
communities, there was some evidence of excess risk of asthmatic attacks
associated with elevated levels of suspended nitrates accompanied by
temperatures of 50°F or above but the findings were less consistent than
those observed in the New York-New Jersey area. The authors stress the
limitation of interpreting these findings in light of the present lack
of knowledge of the chemical and physical characteristics of these
compounds as well as lack of knowledge of their precursors and co-varying
pollutants.
17
In a study by Shy at al. in Chattanooga, Tennessee, an increase in
acute respiratory disease was observed in family groups when the mean
24 hour NCL concentration measured over a 6 month period was between
- o
117 and 205 yg/m-3 and the mean suspended nitrate level was 3.8 yg/nr
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re largely attributed to
or greater* Although the observed effects were largely
NOg the authors pose the question of possible adverse health effects from
suspended nitrates.
6.1.2.1.2 Non-Respiratory Effects—Studies of Czechoslovakian
children who lived near a factory that
emitted nitrate and sulfur dioxide showed that the children in the exposed
community had significantly higher methemoglobin levels than children
living in the control community, but large amounts of nitrates in the
18
drinking water confounded the problem. The study was repeated after the
source of atmospheric nitrates had been controlled and they found no
abnormal levels of methemoglobin. The authors implicated the atmospheric
nitrate as an etiologic agent in methemoglobinemia.
19
Szponar et at. evaluated the level of methemoglobin in 283 subjects
in a Polish village lying within the area exposed to air pollution caused
by a nitrogen fertilizer plant. The study included 88 percent.of the population
and the normal value accepted in the study was 0.13g/100 ml of methemoglobin
in blood. Forty-one percent of children and 37.7 percent of adults above 10 years
had elevated methemoglobin levels. A limitation of this study was failure
to take into consideration the ingestion of nitrates from other sources
such as drinking water.
In 39 tunnel workers* exposed to high concentrations of automotive
emissions»methemoglobin averaged 0.43 _ gm/100 ml, significantly higher
than the level of 0.34 _ gm/100 ml in a control series of maintenance
workers. Total hemoglobin averaged 15.1 gm/100 ml in this population,
suggesting percentage methemoglobin to be about 2.8 oercent. These methemoglobin
levels were in addition to 3 to 5% carboxyhemoglobin levels in the tunnel
workers.
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In general it may be assumed that the contribution of airborne
nitrates to the normal low (1-2% of hemoglobin) levels of methemoglobin
in normal individuals is insignificant relative to that of water, food,
active or passive cigarette smoking, or gaseous NO or M^ (which form
methemoglobin) in the atmosphere. However, studies of subclinical
methemoglobinemia and its chronic effects are needed, especially if
so-called normal adult values do not adequately describe all segments
of the adult population.
Hypertensive heart disease and nephritis mortality rates have recently
been reported to be strongly associated with annual mean N02 levels, but
21
not airborne nitrates per se. Since N02 hydrolyses to nitrous and nitric
acid in the lungs, a lack of observation of a similar effect of airborne nitrates
may imply a threshold level for an effect of airborne nitrates rather than
a complete lack of such an effect. As this study was based on the
correlation of metropolitan mortality rates with the corresponding air
pollution levels for the cities, it obviously cannot be taken as evidence
for the existance of relationships, but only as evidence that these
questions deserve study.
6.1.2.2 Airborne Organic Nitrates
6.1.2.2.1 Ambient Sources—-Several of the organic nitrates have been
incriminated as powerful
eye irritants. Peroxyacetyl nitrate has been purified by gas chromatographic
techniques and its chemical, physical and physiological properties have
been examined. At concentrations in the 1 ppm range»PAN is a strong eye
irritant and is > along with acrolein and
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formaldehyde, probably responsible for eye irritation in ohotochemiral air pollution.
23
In a study by Huiss and Glasson , a new and extremely potent eye
irritant, peroxybenzoyl nitrate, a lachrymator 200 times as potent as
formaldehyde ,was identified as a product from the irradiation of benzylic
hydrocarbons and aromatic olefins. Although this compound has not yet
been found in ambient air, the potential for its presence exists. Benzylic
hydrocarbons are the most common aromatics in gasoline, and as a result
are also in auto exhaust and common solvents. Hence it is probable
that PB N is formed in the atmosphere under smoggy conditions. Only the
development of very sensitive analytical techniques and careful research
will answer the question of the relative contribution of PB N to the eye
irritation observed in photochemical smog.
At present there are no research findings linking organic nitrates
to respiratory disease in man.
6.1.2.2.2 Individual .and Occupational Sour£&s-^A wide
variety of nitrates, nitrites, and nitrohydrocarbon comDounds
used in industry have known physiologic and toxic effects upon inhalation
or absorption.24 These are of interest here only insofar as they suggest
health effects which might be detected in the general population from
widely distributed compounds at much lower concentrations.
A miscellany of acute symptoms may be detected in the industrial
24
setting, and a large number of compounds used form methemoglobin.
Of^particular interest are the chronic cardiovascular toxic effects
which have been reported among workers exposed to organic nitrates in
the explosives industry. Besides acute throbbing headaches these
include elevated diastolic blood pressure, lowered pulse pressure, and
oc on
increased risk of angina pectoris and/or sudden death. ""*"
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The health effects in cigarette smokers of organic substances such
as methyl nitrite and the nitrohydrocarbons are not known. These substances
do cause methemoglobin formation as do NO and N0?. it has recently been
reported that the enzyme D-amino acid oxidase which is concentrated in
the liver and kidneys.will catalyse the oxidation of nitromethane in vitro
30
to yield formaldehyde, nitrite, and H^Og. While one might hypothesize
that kidney damage followed by hypertension might follow the oxidation
of nitro-hydrocarbons in vivo, it is fortunately well "known that cigarette
smokers have not shown increased blood pressure levels in any of numerous
population studies.
6.1.3 Inhalation Toxicology of Nitrates
Very little information is known about the inhalation toxicity of
nitrates. The available information relates to nitric acid aerosols
and peroxyacetyl nitrate.
31
Dautrebande et al. conducted pioneering studies of the effects of
various aerosols, thought to be components of smog, on eye irritation
in mah. Small particle nitric acid aerosol was shown to be an
eye irritant by itself but the irritation was more intense when the
nitric acid vapor was mixed with sodium chloride or used oil aerosols.
3
The lowest nitric acid concentration tested was +_2 mg/m for a 3 minute
exposure.
32 33
Gray et al. conducted some early experimental investigations of the
inhalation toxicity for ratss mice and guinea pigs of N02 generated from
fuming nitric acid. - The particle size of the vapors generated by their
technique were not measured. While the authors describe their studies as
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evidence of NOo toxicity the experimental methods could not adequately
distinguish the difference between NC^ and HN03 effects. Concentrations
of 9 and 14 ppm "NC^" administered 4 hours per day, 5 days per week,
for 6 weeks produced lung pathology. No lesions were demonstrable
when the exposure dose was reduced to 5 ppm.
The acute inhalation toxicity of peroxyacetyl nitrate (PAN) for mice
34
has been reported by Campbell et al. . The median lethal concentration
(LC,-Q) of PAN at 70°F. was estimated to be 106 ppm with most deaths
occurring during the 2nd and 3rd week of continuous exposure. The lethal
concentration was less in ofrder mice, and at higher ambient
temperatures (90°F.) rather than 70°F. The authors did not determine
the pathophysiologic meehanism of action of PAN exposure to
explain why the animals died.
No reference has been found to experimental studies of the inhalation
35
toxicity of any inorganic nitrate. However, Knott and Malanchuk have
called attention to the development of ammonium nitrate aerosol in an
animal exposure chamber where beagle dogs were being exposed to a mixture
3
of feYric oxide (1 mg/m - particle size 0.5 microns) aerosol and nitrogen
dioxide (20 ppm). This report did not describe any Biologic effect
of the nitrate aerosol per se but it is noteworthy as a reminder of the
possibility of new compound formation in animal exposure chambers when
gas-particle mixtures are being investigated concurrently.
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6.1.4 Nitrosamines: Possible Human Health Hazard
N-nitrosocompounds manifest multipotent biological activity in many
species of animals but have seldom been detected in the environment in
significant quantities. The biological actions of nitrosocompounds include
the induction of acute and chronic toxicity, carcinogenicity, mutagenicity
and teratogenicity. The N-nitrosocompounds include nitrosamines, nitrosamides,
nitrosoamidines and nitrosocyanamides. Much of the concern is due to the
widespread occurrence and use of nitrogen compounds that become a part of
our "way of life" and their possible interactions to form the N-nitrosocompounds,
The nitrosamine precursors include secondary and tertiary amines, ureas,
carbamates, guidelines amino acids with secondary amino groups, nitrites,
nitrates,and nitric oxides (NOX).
c i A -iNitrates and Nitrites: Their Role in the Nitrosation Process—The
D • 1 • T- • I ' ' ™ • —. __—__ —__ — -. . . - --"
occurrence and and distribution of nitrates and nitrites in our environ-
ment has been discussed elsewhere in this report (Section 5).Concern centers
on the consumption of products particularly high in nitrate content [e.g., some
well waters, vegetables, cheese and cured meat ] as these nitrates
are reduced to nitrites by microbial agents. » The release of oxides of
n £9
nitrogen from ensiled forages may be harmful to man and animals.*1**'1 Nitrate
that occurs naturally in tobacco is converted into ammonia during the smoking
43
process . This indicates that much of the smoke chemistry of nitrate may
be the chemistry of intermediary ammonia. This ammonia has been shown to
yield amides and N-heterocyclics which in all probability contribute to the
reported presence of nitrosamines in tobacco ' and tobacco
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The use and the presence of nitrite in processed (cured) meat and fish is a
coimon knowledge. With the presence of precursor amines (e.g. trimethylamine,
trimethylamine oxide, and dimethylamine) and nitrite in such preparations
there is a potential for the formation of nitrosamines. Indeed, nitrosamine
48-53
formation has been shown in some such preparations- Nitrite is
present in the human saliva apparently derived from reduction of nitrate by
59
bacteria in the mouth- Nitrite interacts with phenolic compounds
readilyresulting in the formation of nitrosophenols whose biological activity
is not known .
6.1.4.2 Factors That Influence Nitrosation --It has been suggested
that "the protonation of nitrous acid appears
necessary for initiating all nitrosation reactions' and that
carcinogenic N-nitrosocompounds in quantities considered to be potentially
hazardous can not be produced unless the interaction of nitrite and amine
occurs in acidic media. Thus, the acidity prevailing in the mammalian
stomach probably presents favorable conditions for the formation of
jry cp
nitrosamines from amines and nitrites. Sander and his colleagues *
showed the formation of the corresponding N-nitrosocompound by incubating
several secondary amines (e.g., diphenylamine, N-methylaniline and N-methyl-
benzylamine) with nitrite in the presence of gastric juice under various
59
conditions. Sen et al. demonstrated the in vitro formation of
diethylnitrosamine by incubating diethylamine and sodium nitrite in the
gastric juice from rats, rabbits, cats, dogs and man. The basicity of the
amine (weakly basic amines yield more than the strongly basic amines) and
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pH (optimum 1-3) determine the total nitrosamine yield in such reactions.
60
Mirvish showed in vitro that the rate of reaction was maximal at pH 3.4
when formation of dimethylnitrosamine was proportional to dimethylamine con-
centration and to the square of the concentration of nitrite. Lijinsky et
61
al. demonstrated that tertiary amines are optimally reactive towards
nitrous acid in the pH range 3-6 and the reaction proceeds at a measurable rate
even in
nearly neutral solution. The same authors suggest that trisubstituted amines
can compete favorably with secondary amines for available nitrite. The
success of this competition depends strongly on hydrogen ion concentration.
This strongly suggests that tertiary amines are more important than secondary
amines as environmental precursors of nitrosamines in the limit of mildly
62
acidic conditions. Recently, Keefer and Roller have added a new
dimension showing significant synthesis of N-m'trosocompounds using amines,
nitrite and formaldehyde or chloral in the pH range 6.4-11.0. The role of
intestinal bacteria in-the enzymatic synthesis of nitrosamines has been well
63-65
documented- Of considerable relevance under physiologic conditions
56
are the potential catalysts in the reaction medium. Ridd demonstrated
that such nucleophilic anions as chloride and acetate exert a reaction-promoting
64
influence on nitrosation. Catalytic effects on the nitrosation by sulfate*
phosphate66 . and thiocyanate6^ have been shown. The thiocyanate is a
normal constituent of saliva and smokers show higher amounts in the saliva.
Ascorbic acid prevents the formation of N-nitrosocompounds from the interaction
of some drugs and secondary amines both in vitro6** -and in vivo69»70
However, it should be cautioned, that the usefulness of this technique in
practical terms to diminish the nitrosation associated with concurrent
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administration of secondary amines and nitrosatable compounds is limited due
to the high amounts of ascorbic acid required for blocking the nitrosation.
6.1.4.3 Nitrosamines in Human and Animal Food Commodities — Nitro-
samine research started with the indications of nitrosamine
toxicity to humans in an industrial context- The search for nitrosamines
in food commodities was triggered by the dramatic outbreak of serious liver
necrosis in sheep in Norway 72 fed fish meal preserved with nitrite. The
48
presence of dimethylnitrosamine (30-100 ppm) in this meal is not
surprising as fish meal contains relatively large amounts of secondary and
tertiary amines. Sheep given pure dimethylnitrosamine manifested closely
similar liver lesions as seen in the sheep that received the toxic fish meal.
The presence of very small amounts of nitrosamines in smoked herring, kippers,
51 52
smoked haddock, smoked sausage, bacon and smoked ham has been reported- .'
The presence of nitrosamines has been shown in the food and drink
73
consumed by the Bantu population of South Africa. . A significant spatial
correlation between the geographical patterns of incidence of esophageal
74
carcinoma and the drinking of sugar-based alcoholic spirit was noted • The
alcoholic extracts of the fruit of the plant Solanum incanum yielded
74
dimethylnitrosamine'' . The juice from this fruit is used to curdle
milk in South Africa. Lijinsky and his colleagues *. have pointed out
that many drugs and pesticides contain tertiary amino groups and could there-
fore be expected to undergo nitrosation in the body. Examples of such drugs
which react with nitrite in vitro to give varying yields of carcinogenic
nitrosamines are oxytetracycline (antibiotic) and aminopyrene (an analgesic)
which yield dimethylnitrosamine, disulphiram (antialconolic) and nikethamide
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• [;,?-.
r»A MT1" r.'?'v.~r ro psT'r
UJ PsJi y^-ih,: ^K ^'5^
(respiratory stimulant) which yield diethylnitrosamine and tolazamide (oral
hypoglycaemic) which gives nitrosohexamethyleneimine. Aminopyrene produced
enough dimethylnitrosamine in vivo in the rat to cause the typical acute
8n
liver necrosis u and to induce malignant tumors of the liver on continued
81
concurrent administration with nitrite. Combined concentrations
were 250 ppm. Also, the conversion of a morpholine derivative,
phenmetrazine, into its N-nitroso derivative in the rabbit and rat
82
stomach iji vivo was shown. The simultaneous oral administration
of sodium nitrite and dimethyl amine to mice gave rise to
83
typical hepatic necrosis. The relevance of these observations to human
welfare (toxicity and cancer) is very difficult to judge. The evidence
showing similar in vitro metabolism of dimethylnitrosamine in man and
84
rat makes it unwise to assume that man is resistant to the multipotent
biological activities of the nitrosamines. The amounts of preformed
nitrosamines in foods appear to be very small (dimethyl nitrosamine:
0*01-0.08 ppm) but the quantities that may be formed endogenously
are unknown and must depend, among other factors, on the amounts of
nitrites and nitrosatable amines and amides present in the body.
6.1.4.4 Biological Activity of N-nitrosocompounds—The biological
activity of the N-nitrosocompounds is dependent largely upon their
chemical stability. The chemically stable dialykl nitrosamines
(e.g., dimethylnitrosamine) apparently require, as the first metabolic step,
an enzymatic alpha hydroxylation for their activation into carcinogens. In
contrast the chemically less stable nitrosamines (e.g., N-nitrosomethylurea)
have marked local cytopathic effects and do not seem to require enzymatic
activation. Their decomposition is, in some cases, catalysed by sulphydryl
OC
compounds. °
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Several excellent papers review acute toxicity and carcinogenidty of
Q/-_OQ
the nitrosamines. The nitrosamines, with some exceptions, are
selectively hepatotoxic while the nitrosamides damage predominantly organs of
rapid cell turnover, such as the gastrointestinal tract, the blood forming
organs and the lymphoid system. Manifestations of nitrosamine hepatotoxicity
include "blood cysts" areas of destruction of the parenchyma filled with
recently extravasated erythrocytes, and necrosis of the endothelium of the
central and sublobular veins with extrusion of necrotic hepatic parenchymal
cells into the lumen. The kidney lesions are limited exclusively to the
convoluted renal tubules. Rats treated with sublethal doses of
90-
dimethylnitrosamine have veno-occlusive lesions in the liver- Similar
91
veno-occlusive lesions occur in mink treated with dimethylnitrosamine •
Typical acute liver changes have been induced in rats by the heterocyclic
42. 93
nitrosamines. ' Electron microscopy of rat liver treated with
diethyl or dimethyl nitrosamine showed partial separation of the fibrillar
and granular components and the formation of electron dense plaques at the
periphery of the neucleous—microsegregation. Interruptions in the plasma
membrane and increase in the number of microbodies in the cytoplasm were also
94
evident Recent studies indicate that autophagy and disturbance of
lysosomes, begin within 35 minutes after treatment with dimethylnitrosamine
95
and reach a peak at 12 hours- . In contrast, some of the nitrosamides
are extremely irritating locally, for example N-methyl-N-nitrosourethane
OT...
causes severe necrotic stomach lesions after gavage.01 N-methyl-N-nitrosourea
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hamsters treated with N-methyl-N-nitrosourea may provide an experimental
97
model for the study of retinal damage. Evidence indicates that nitrosamines
98
mediate their carcinogenic activity through the formation of carbonium ion
99
and also through alkylation of guanine- The amount of 7-methylguanine
formed in various organs is frequently related to the susceptibility of that
organ to the development of neoplasms, but such a correlation is not always
A 86
observed•
6.1.4.5 Carcinogenesis: N-ni trosocompounds—The
nitrosamines are highly potent (a single dose may suffice to induce tumors)
and versatile (cosmopolitan in species susceptibility, active by various
O/- OQ
routes of administration, multiorgan system involvement). The reviews
quoted earlier contain information on nitrosamine carcinogenesis. Newer
data dealing with the carcinogenesis primarily due to cyclic nitrosamines,
concurrent feeding of amines and nitrites (to determine that the carcinogenesis
was caused by the endogenous formation of respective N-nitrosocompounds), and
possible synergistic effects of other agents will be presented in the following
paragraphs.
Diethylnitrosamine (DEN) has an organotropic toxic and carcinogenic
activity on the respiratory system including the nasal tissues of many species
\
104
of animals. Some neoplasms of the nasal tissues of animals induced
by DEN are histologically and biologically similar to those seen in humans
Nitrosamine inhalation has produced tumors of the upper respiratory tract of
hamsters. -|U*» "eoplasms have developed in the offspring of female mice
ing. 110
and hamsters'. vv^ treated during pregnancy with several nitrosocompounds•
An example of modification of the neoplastic expression of a known carcinogen
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urw I
_ ,,f.~. *•
c ui\
v\)
caused by respirable particulate material occurs in hamsters. They develop
neuroepithelial neoplasms (esthesoneuroepitheliomas) of the nasal tissues
when dimethylnitrosamine (DMN) is given subcutaneously and ferric oxide is
instilled intratracheally •' The intragastric administration of
methylcholanthrene and intraperitoneal administration of dimethyl nitrosamine
to Swiss mice resulted in an increased incidence and decreased latency period
112
of neoplasms, compared with the mice treated with either compound •
Several cyclic nitrosamines induce neoplasms of the liver and other
I 10
organs A high incidence of neoplasms of the liver, tongue and
esophagus was seen in rats given N-nitrosohexamethyleneimine in drinking
93
water. N-nitrosoheptamethyleneimine (N-6-M1) produced squamous neoplasms
114
of the lung and esophagus in rats- . Subsequently, Lijinsky and his
81
colleagues induced high incidence of squamous neoplasms of the lung of
rats fed an amine-heptamethyleneimine together with nitrite. The total
dose administered over a 22 week period was 140 mg. These findings suggest
that lung cancer in cigarette smokers might be caused by carcinogenic
nitrosamines formed from nitrite in the food and secondary or tertiary amines
in cigarette smoke condensed in the mouth and swallowed in the saliva. While
some nitrosocompounds are equally carcinogenic independent of the route of
115
application, the N-6-M1 was less effective when given subcutaneously.
However, recent studies indicate that a single subcutaneous injection of
N-6-M1 (dose range: 4-64 mg/kg body weight) induced high incidence of
respiratory tract neoplasms in hamsters and mice. In an attempt to
evaluate respiratory infection as a cofactor in the development
of respiratory tract neoplasia, germfree, specific-pathogene-
free and infected (chronic murine pneumonia) rats were given N-6-M1 in drinking
water for 22 weeks. Rats were sacrificed 2 weeks post treatment. The incidence
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of lung neoplasms indicated that chronic respiratory infection enhanced the
neoplastic response of the lungs to a systemic carcinogen • These data
show that cyclic, nonalkylating nitrosocompounds induce neoplasms similar to
those induced by most aliphatic nitrosamines " and alkylating cyclic
119
nitrosamines • Aminopyrene has been shown to yield dimethylnitrosamine
80
In vitro and to induce malignant neoplasms of the liver on continued
81
simultaneous administration with nitrite. The possibility that
nitrosamines are formed from nitrite and secondary amines under the acidic
conditions of the stomach was hypothesised ' and was proven to be a
reality Recently, lung adenomas were induced in Swiss mice chronically
124
fed nitrite plus morpholine, piperazine, or N-methyl aniline • The continued
administration of morpholine or methyl benzyl amine simultaneously with nitrite
to rats in their drinking water induced the tumors expected from the known
125
carcinogenic effects of the corresponding nitrosamines. However, no
neoplasms were induced in rats by feeding the secondary aminoacids orpline,
12fi
hydroxyproline and arginine concurrently with nitrite;" An extension of
the above hypothesis 57,120 js t0 study the formation of nitrosamides from
corresponding alkylureas and nitrite under acid conditions. Sander and
127
Burkel induced a spectrum of neoplasms in rats feeding nitrite together
with N, N -dimethyl urea, methyl-urea, ethylurea, or 2-imidazolidone (a
cyclic urea). Neurogenic neoplasms occur in the offspring of female rats
1 28
fed nitrite and ethylurea during pregnancy which are similar to those
induced in the offspring of rats given nitrosamides during the last half of
129
pregnancy Nitrosamides have been formed in vivo by the concurrent
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CITE
feeding of nitrite and methyl urea and ethyl urea and have produced lung
130
adenomas in Swiss mica • , It is speculated that nitrosamines may be
formed from the absorbed nitrogen dioxide and nitric oxide from the ambient
131
air • -•
The susceptibility of many species of animals (rat, mouse, hamster,
guinea pigs, rabbit, dog and monkey) to the carcinogenic action of nitrosamines
indicates that man is probably equally susceptible.
6.1.4.6 Human Health Hazard—The causal relationship of
nitrosamines to human cancer is suaaested
by the following: a. All mammals subjected to the nitrosamine carcinogenic
study show susceptibility, b. Nitrosamines are highly effective carcinogens
by all routes of administration, c. Some animal neoplasms induced by
nitrosamines are predominantly epithelial and their organotypic and histologic
features resemble those seen in humans, d. Nitrosamines are metabolised rr^
vitro similarly by human and rat tissue, e. Nitrosamines occur in the human
environment - cured meat and fish products; tobacco and tobacco smoke and
others, f. Evidence is overwhelming that nitrosamines or nitrosamides are
formed in vivo, most likely in the mammalian stomach or possibly in the lower
gastrointestinal tract through interaction of secondary and tertiary amines
ingested as food or drugs with nitrite and due to microbial metabolism. The
cause and effect relationship between certain environmental and pharmaceutical
chemicals and the development of certain human neoplasms has been shown beyond
any doubt. Only a few nitrosatable compounds and the respective nitrosamine
derivatives have been shown to induce neoplasms in laboratory animals. The
daily exposure of man to nitrosamines is unknown. However, the incidence of
"cancer" in man is quite large, and the distribution of "cancers" of certain
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types in some segments of the population seems to reflect well defined exposures
to carcinogens. JThe nitrosamine precursors are ubiquitous in the environment.
Continued, chronic exposure to these nitrosamine precursors could conceivably
result in the formation of nitrosocompounds. Nitrosamines are metabolised
similarly in vitro by rat and human tissues. The evidence for the carcinogenic
activity of nitrosamines in laboratory animals including the rat is over-
whelming. Even if one were to assume that nitrosocompounds, were to be
present at a very low level, unremitting insults from these compounds over
several decades may be hazardous to health. Moreover, the distinct possibility
of synergistic effects of other noxious agents on the biological behavior of
nitrosamines should be considered. (The task of detection, identification
and delineation of biological behavior of nitrosamines in the environment is
of an enormous magnitude.
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6.1.4.7 Epidemiological Studies on Humans—The epidemiclogical
evidence for a role of nitrosamines in human cancer is extremely
limited, and the available data are from studies in Africa. In
132
one study of plant products in a small area with excessive
esophageal cancer rates* dimethylnitrosamine was detected in the
juice of the fruit of a solanaceous bush which was used to curdle
milk. The resulting food product is nearly always consumed by
males, although excessive female esophageal cancer mortality existed
133
also. In the second study an analysis of a local alcoholic
spirit was carried out and dimethylnitrosamine detected. A geographic
correlation exists between high consumption of these spirits and
esophageal cancer in Africa. It is well known, however, that
esophageal cancer is associated with the consumption of alcohol in
134
developed countries, but nitrosamtnes were not found in English
gin. There do not appear to be other reports of nitrosamfnes
in distilled alcoholic beverages in developed countries. In addition
to the epidemiologic weakness of these two studies there have been
significant changes in the laboratory methodology for confirming
the presence of dimethyl nitrosamine since these studies were published.
Beyond these two studies there are in the literature only hypo-
theses relating observed phenomena. Geographic studies of gastric
cancer in Chile have been reported with the hypothesis that the
esophageal and gastric cancer rates are related to nitrates in food-
134
stuffs fertilized with the native nitrate fertilizer. The suggestion
has also been made that the high incidence of nasopharyngeal cancer
among the Cantonese is related to the presence of both dimethyl and
diethylnitrosamine in dried fish.
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The British Medical Journal has recently summed up the situation
in an editorial: "At present we have no clear quantitative indication
of the risk that nitrosamines in the environment may present to man.
To some workers in the field the expectation is that the hazards are
low, if present at all, though this is still guesswork. However,
none would question the desirability of promoting research that will
"136
provide more convincing conclusions.
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.KS Research Needs and Priorities ^' Q'E. QR £/JT
Our extremely limited knowledge of the health effects of nitrites, organic
or inorganic, dictates a need for sequential implementation of research.
Some research questions, cannot be properly addressed until certain
information is forthcoming from other studies.
Until the nitrate compounds in ambient air are adequately characterized,
only limited animal and human studies seem appropriate. These are:
(1) animal studies of the pulmonary effects of nitrate
analogs of the sulfate compounds currently under
study,
(2) replication, as currently planned, of the asthma
panel studies wlyich have yielded our strongest
evidence to date of a health effect of nitrates,
and,
(3) studies of methemoglobinemia and its possible
effects in persons exposed to high density urban
traffic.
In the special field of carcinogenesis the first step is to search
for N-m'trosocompounds in the air; because of inadequate analytical
technology' a lack of interest among investigators, or some other, reason, these
have yet to
be detected despite the known presence of the precursor compounds. Animal
studies should follow the isolation and identification of N-nitrosocompounds
from the air.
Special attention should also be directed to the peroxyacetyl
nitrate (PAN) precursor perox.yace.tate*, This powerful oxidant may
react with polycyclic hydrocarbons to form epoxy compounds. Some
carcinogenic compounds have to be in this epoxy form (the ultimate
carcinogen) to exert their biological activity. A study of the interaction
°f peroxvacetate an<* benzo (a) pyrene in a biological model would be
of interest.
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6.2 ECOLOGICAL tFFECTS
Nitrogen is a component of all living matter known to man.
The major source of nitrogen is the earth's atmosphere where, in
molecular form, it is a major constituent. Nitrogen compounds
are soluble, volatile, or readily decomposable, therefore, except in
extremely dry land areas witK little rainfall, no large deposits
exist.
Most organisms are unable to use nitrogen in molecular form, therefore,
it must be converted into other chemical forms. The great
importance of nitrogen as a nutrient element for plants has resulted
in the study and elucidation of the movement of nitrogen through the
biosphere. To become available to most plants and animals nitrogen
must enter the soil or water. Nitrogen enters these media in rain
water as ammonia or nitrates, as particulate nitrate,or as biologically
fixed nitrogen.
The transformations of nitrogen in the Biosphere are regulated
almost entirely by terrestrial and aquatic microorganisms. In
general outline, the nitrogen cycle is identical in terrestrial,
fresh water and oceanic habitats; only the microorganisms which
mediate the various transformations are different.
(1) Biological Nitrogen Fixation - Atmospheric nitrogen gas is
transformed into ammonia, nitrates and other nitrogen-containing com-
pounds. The transformation is carried out by a wide variety of
microorganisms. The microorganisms may be either symbiotic (.living
in the roots of leguminous plants) or nonsymbiotic Giving independently
in the soil) and the process may be accomplished under aerobic or
anaerobic conditions.
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(2) Organic Nitrogen Formation (assimilation) -- Fixed nitrogen
as either nitrates or ammonia is assimilated by plants and converted
to plant protein. Plants are eaten by animals and plant proteins
are converted to animal proteins. In addition, carnivores consume
other animals as a protein source. Nitrogen is bound in plant or
animal protein until the organisms die.or as in the case of animals.
certain products are excreted.
(3) Deamination or Ammonification — A two step process, also
termed mineralization, in which the excretion products of animals and
the proteins in dead plants and animals are broken down through
proteolysis to amino acids. The amino acids in turn are converted
into ammonia. The ammonia may be assimilated by aquatic or
terrestrial plants and microorganisms* may be bound by clay particles
in the soil » or it may be converted by microorganisms to nitrates
in the process termed nitrification. It may also escape into the air.
/4) Nitrification — Formation of nitrates through the microbial
conversion of ammonia first to nitrite and then to nitrate. Nitrates
may be assimilated by plants, washed downward through the soil into ground-
water or through surface runoff into streams, rivers, and oceans, may
be transformed into atmospheric nitrogen or reduced to ammonia.
(5) Nitrate Reduction -- Microorganismal conversion of nitrates
back to ammonia via the nitrite step. These processes are the converse
of the previous transformations.
(6) Denitr1f1cation — Nitrates through bacterial action are converted
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into atmospheric nitrogen. Denitrification is an anaerobic process.
Nitrates (N03~) are converted into nitrites ("Nt^") » to nitrous oxide
(NpO) and finally into nitrogen gas (Np) which goes off into the atmos-
phere. In the soil, nitrites rarely accumulate under acidic conditions,
nitrites decompose spontaneously to nitric oxide (NO), and under
137 138
alpine conditions, they are biologically converted to ^0 and ^ '
It must be emphasized that this process is anaerobic and that conversion
of nitrates to nitrites is extremely sensitive to the presence of
atmospheric oxygen. If atmospheric oxygen is present, the conversion
does not occur.
Figure 6.1 is a diagrammatic presentation of the movement of
nitrogen in the biosphere. The transformation and movement of nitrogen
as explained in the foregoing discussion relate to the biogeochemical cir-
culation of nitrogen. The circulation of nitrogen is a long term
process. Turnover times for the three largest "pools" of nitrogen
o
are: 3 X 10 years for atmospheric nitrogen, 2,500 years for nitrogen
in the seas when nitrates and organic nitrogen compounds are counted together
139
and less than one year for nitrates and nitrites in the soil. A
more detailed account of the distribution and annual transfer rates is
140
shown in Figure 6.2 . As Delwiche points out, the transfer rates
can be estimated only within broad limits. The only two quantities
of nitrogen known with any degree of accuracy are the amount of
nitrogen in the atmosphere and the rate of industrial fixation. The
amount and the length of time nitrogen is in the atmosphere indicates
why the atmosphere is the greatest source of nitrogen, while the
short period of time nitrogen is in the soil emphasizes why nitrogen
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Pool of Nitrogen
in Atmosphere
Volcanic
Action
Electrification
and
Photochemical
Fixation
Nitrogen-fixing
blue-green algae
and bacteria
Excretion
rea, etc
Marine Life
Protein Decay
by
Bacteria and
Fungi
Nitrogen-fixing
Bacteria in
soil and roots
Ammonification
Shallow
marine
sediments
Protein
Synthesis
Nitrate
Bacteria
Nitrite
Bacteria
Nitrate
bacteria
Denitrification
Figure 6.1 Nitrogen Cycle
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Qrgonic N
Sediments
0000
Organic N
Animals
0.00215
. J..
Figure 6.2 The biogeochemical cycle for nitrogen.
Numbers in circles are amounts of nitrogen in pools,
in kilograms per square meter of the earth's surface;
numbers on arrows are transfer rates in kilograms of
nitrogen per s,quare meter per year.
Ref. 139.
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is so often in short supply as a nutrient element. ' W07£ Qft «.-.
6.2.1 Nitrates as Fertilizers
Man's greatest intervention into natural cycles has occurred
because of the shortage of nitrogen as an available nutrient ele-
140
ment in the soil. " It has been estimated that the amount of
ammonia nitrogen (NH^-N) that is converted by microorganisms to
nitrate nitrogen (NOo-N) is equivalent to the net nitrogen as-
similation (0.017 kg-N/m2) by plants each vear.141 To alleviate this shortage
the limiting factor in plant growth, industrial fixation of
nitrogen was developed. At the present time the amount of nitrogen
fixed industrially for the production of fertilizer equals the
amount that was fixed by all terrestrial ecosystems before the
140
advent of modern agriculture. The world's annual output of in-
dustrially fixed nitrogen was 30 million tons in 1968 and it has
been estimated it will reach or exceed 100 million tons by the year
2 '
2000. Consumption of fertilizer nitrogen in the U.S.A. will probably
142
reach 11 million tons in 1980. The impact of this environmental
loading has not, until very recently, been considered.
Inorganic nitrates are soluble in water and the nitrates not
rapidly utilized by plants or bacteria are quickly washed av/ay into
the nearest streams or lakes or may percolate down to the water
table. Not only are these nitrates lost to the farmer, but they can
cause eutrophication problems in the waterways or enter the drinking
water from where they may cause problems in humans.
The use of nitrate fertilizers has been beneficial as they have
.142
increased growth yields tremendously so that greater amounts of food
can be produced on smaller areas. It is the side effect, the fact
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that plants only assimilate approximately 50 percent of the nitrogen
fertilizer placed on the field,142 wnile the
rest is washed off into creeks, rivers and into the ground wa^er,
that creates the problem.
Plants growing in the soil normally obtain their nitrogen
from the soil as NCL-N. The extent to which a plant is exposed to
nitrates depends on the concentration of nitrates within the plant
root surface area. The amount of nitrate in this region depends
on whether the source is the process of microbial nitrification
or whether nitrate fertilizers have been added. Under those soil
conditions in which microbial nitrification is the sole nitrogen
source, high NOo-N levels will probably not be reached. The problem
of high soil nitrate levels has only occurred since man began to
add chemical nitrates to the soil. Therefore, exposure levels of
plants to nitrates is dependent on the amount of fertilization which
takes place through the addition of inorganic nitrates or inorganic
ammonia. If inorganic ammonia fertilizers are added to the soil,
these may be microbially changed to NOo-N or they may be taken up
by plants as NH.-N. The fate of both soluble and slow-release nitrogen
fertilizers is shown in Figure 6.3. The extent to which plants
may be exposed to NOo-N is dependent on losses through(l) leaching
of nitrite (N02~) and nitrate (N03~);(2) biological denitrification
of both N02~ and NO," through microbial completion;^) volatilization
of ammonia (NH3) from improper application of anhydrous or aqueous
ammonia, and surface application of urea and other ammoniacal nitrogen
142
sourcesto alkaline soils; and(4) chemical denitrification. Loss of
fertilizer nitrogen may also result from_ surface runoff and erosion,
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BIO OEfmniflCAllON
SURFACE BUHOFF
AMD
CROSIOM
CHEMO-OENITfiiriCATION
V
FERTILIZER H (SOLUBLE):
VOLATILIZATION
y^~"V
/ SOLUBLE \ OISSOLUTION
son
OHGAIilC H
TOOL:
HUMUS. M!CRO0I».L
CCllS. PLANT ANO
ANII/Al Rf.SiOUES.
nioioc.ic'.iLY rixco
ASO
XONSYMDIOTIC)
l?(Ttf!-LATTICE
flXATION OF NM4
ABSOnPIIOS BY
PLANTS
LEACHING
Figure 6.3. Schematic representation of the fate of
soluble and slow-release N in the soil, and relationship
to the soil organic N pool.
Ref. H2.
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inter-lattice fixation of ammonium (NH4 ) by clay minerals, micro-
bial immobilization,and chemical immobilization involving
142
reactions of fertilizer nitrogen with soil organic components.
Determining the levels of N03-N and the extent to which
the roots of a plant will be exposed to those levels is not possi-
ble at this time with any degree of accuracy. Were it possible
to determine the levels of N03-N and NH^-N in the rhizosphere
and the time of maximum nitrogen assimulation by the plant,
fertilizers could be applied at that time so that the greatest
benefit could occur.
6.2.2 Nitrates in Aquatic Habitats
Runoff waters in agricultural areas frequently carry high
levels of nitrates. Any plant growing in water or on the shores
of a lake or stream which contains high levels of nitrate is likely
to be exposed. Phytoplankton algae and other microorganisms
growing in the surface waters and the algae and vascular plants
growing in benthic areas are the most likely to be exposed to the
nitrates in lakes and ponds.
In addition to the nitrates in runoff water, additional nitrates
may be added to the water through nitrification taking place in the
143 144
sediments • and in the surface waters-
Specific concentrations of nitrates in streams or particularly
in lake waters will depend on the sources of the nitrates. The
concentration of any nitrogen compound in water is the net result
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of the rates of nitrogen immobilization, mineralization, nitrification
144
and denitrification.' These levels vary during the year with the
NOo-N and N02-N levels in lakes being maximal in the spring and
minimal during the middle to late summer. The NH.-N levels, on
the other hand, are usually highest in surface waters in the fall
144
and highest in deep waters during the summer. The levels are
144
influenced by biological utilization and/or denitrification.
The effects of biological activity may also be seen in many lakes
which exhibit a NOg-N distribution pattern in which NOo-N levels are
low in surface and bottom waters with maximum levels in intermediate
depths. In surface waters the NO^-N nitrate is immobilized and in
144
bottom water denitrified. Not all aquatic plants utilize N03-N
Some take up NH4-N preferentially over NOg-N. The "bloom" of
aquatic organisms has been associated chiefly with high NOo-N levels.
6.2.3 Nitrate Accumulation in Plants
High nitrate levels in plants has been found to cause acute
poisoning in cattle, sheep, other livestock and has been suggested
as a possible source of nitrate poisoning in man.
In most plants, nitrate is reduced to ammonia and then used
145
in the synthesis of protein and other organic compounds. Accumu-
lation of nitrate in plants may indicate that rate of assimilation
has not kept up with the rate of uptake. Accumulation of nitrate
146
in such cases is only temporary and diminishes as the plant ages.
The stage of development and a variety of environmental factors
influence the nitrate content. Variation from plant to plant as
.146
well as within species,genus or family may be very great. Plants
which normally have low levels of nitrate may,under certain con-
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ditions accumulate it to very high levels. Perennial forage
grasses have been found to be low in nitrate in many tests, but
have been observed to accumulate nitrate in other cases.
The plant families most often considered to be accumulators
of nitrates are: Amaranthaceae, Chenopodiaceae, Cruci ferae,
Compositae, fcramneae and Solonaceae. No single family has been
146
thoroughly sampled for nitrate accumulation. ,
Nitrate is not uniformly distributed throughout the plant
tissues. Stems usually contain more nitrate than leaves and
leaves more than flower parts. Roots have not been tested to any
degree but appear to contain lower levels than stems. In fodder
sugar beets, the oldest leaves were found to be the highest in
nitrate nitrogen (N03~N).
The site of nitrate reduction in plants is not definitely known.
It has been suggested that in woody plants, nitrate is reduced in
the roots while in herbaceous plants reduction occurs in the leaves. McKee,
however, feels that sufficient information does not exist to
determine the site of reduction with any degree of certainty.
The activity of the enzyme, nitrate reductase, has been said
to influence nitrate levels in plants. Energy is required for nitrate
reduction. Carbohydrate is the energy source. High levels of
nitrogen appear to stimulate the plants to utilize stored carbo-
hydrate for energy to reduce the nitrate through the activity of
nitrate reductase. As a result of rapid nitrate assimilation and
rapid carbohydrate utilization, the level of plant nitrate increases
and the carbohydrate level decreases. Therefore, a plant high in
146
carbohydrate is likely to be low in nitrate.
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Nitrogen uptake by plants is also influenced by the
content of the soil and soil moisture. Nitrate will not accumu-
late in a plant unless the levels in the soil are. sufficient
to permit rapid assimilation by the plants. Potassium nitrate
appears to be more rapidly taken up than either calcium or sodium
146
nitrate.
Molybdenum and manganese deficiencies have been shown to be
associated with accumulations of nitrate in both plants and micro-
organisms. Molybdenum is the metallic component of nitrate reductase
enzyme.
Moisture-dependent processes contribute to the accumulation of
nitrate. Microbial activity which releases nitrate from complex
organic compounds requires moisture. The nitrate released in
this process requires moisture to move through the soil to the
plant roots and move across the cell membranes. Fertilizer nitrogen
also requires moisture if a plant is to utilize it.
Plants which have been under a moisture shortage stress may
accumulate high levels of nitrate in a very few days. The moisture
shortage results in a disturbance of assimilatory processes. As
a result, a drop in nitrate reductase activity may occur while the
plant continues to assimilate nitrates causing a high nitrate accu-
mulation in pi ants.through experimental moisture stress studies have been
146
few and not very successful.
A variety of other factors such as light, herbicides, temperature,
soil type and parasitization by diseases and insects have all been
said to influence nitrate concentrations in plants.
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6.2.4 Exposure to Airborne Nitrates
The movement of nitrogen through the atmosphere is composed
of three separate cycles, the N?0 cycle, the NHL cycle and the NO
147 J x
cycle. The NO cycle has the greatest effect upon plant communities
X
and therefore upon the other organisms which reside in these
communities.
1 £7
The major urban sources of NOX emissions are the varied combustion
processes. NO and N02 may move into the atmosphere
where reactions with 0~ and water result in the formation of HNO-
J 147 J
vapor and eventually nitrate salt aerosol. Rain or dry particulate
deposition eventually brings the nitrate down to the earth's
147
surface. Robinson and Robbins estimate the amount of nitrate brought
down to be 462 X 10- tons per year. The greatest amount of nitrate
fallout occurs over the ocean. Nitrate fallout of the above type has
not been associated with plant damage or with other adverse plant effects.
The nitrogen oxides emitted from combustion processes may, in
urban areas, instead of moving off into the atmosphere enter into
photochemical reactions and form peroxyacyl nitrates. The formation
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of these compounds are discussed in greater detail elsewhere.
The peroxyacyl nitrate family includes peroxyacetyl--.nl trate
(PAN), peroxypropionyl nitrate (PPN), peroxybutyl nitrate (PBN),
and peroxyisobutyryl nitrate (PBN). PAN, the most abundant
member of the family, is responsible for serious plant injury
148
in some polluted areas and because of this has been most studied.
Preliminary studies have shown that PPN is several times as toxic
149
to vegetation as PAN, ' whereas PBN and Pi<;nBN are more toxic than
150
PPN. Since PAN is the only member of the series that has received
much study, and since PPN and PBN usually are not measured in the
ambient air, discussion will be restricted to the effects of PAN.
The presentation of the quantitative effects of ozone and PAN has
to be limited to laboratory and controlled field exposures because,
under ambient conditions, the effects of these compounds cannot
easily be differentiated. The term "oxidant" is usually used when
discussing the toxic materials to which the plants are exposed under
ambient conditions. Research has shown that additional phytotoxi-
151,152
cants may be present in the photochemical complex. Synergistic
effects between the toxicants discussed and other atmospheric con-
149,153,154
taminants may also produce injury of sensitive plant species.
On the basis of available information, ozone is the most important
phytotoxicant of the photochemical complex.
PAN-type injury, characterized by under-surface glazing or
bronzing of the leaves of many plant species, has been observed in
California and in the states along the eastern seaboard of the
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149,155
United States. This type of injury has also been reported
in other parts of the United States and in several metropolitan
156rl58
areas of other countries. Injury occurring in the field
has been identified in spinach, beets, celery, tobacco, endive,
romaine lettuce, Swiss chard, pepper, alfalfa, petunias, snapdragon,
149,159,160,161
primrose, asters, and other plants.
PAN has been accepted as the primary phytotoxicant that
causes the oxidant-type injury initially described by Middleton
162 163
et al., and more completely by Bobrov et al. and Glater, et
164
al. The initial collapse is in the spongy cells surrounding the
air space into which stomata opens. The effect in some cases is
limited for the most part to cells nearest the lower epidermis.
This results in a slight separation of the lower-leaf epidermis,
which produces a characteristic under-surface silvering, glazing,
or bronzing. More acute injury causes the necrosis to extend through
the entire leaf. Injury to the leaves of grasses, petunia, and
tobacco causes a cross-leaf banding associated with the sequential
maturation of cells from the tip to the base of the leaf. A
detailed discussion of PAN symptoms with illustrations has been
165
published.
Acute injury to sensitive petunia and tomato plants Kas occurred
148
from four hours of exposure to approximately 14 ppb of PAN. The
PANs are highly sensitive phytotoxicants and produce visible
symptoms of injury only on leaf tissue of a specific physiological
148
stage of development. ' Very small quantities of PAN are capable
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of producing injury to sensitive plants. It was postulated by
Mudd that the action by PAN must affect some substance in plant
tissue present in small amounts, enzymes for example. Experi-
mental work indicates that PAN in amounts from 80 to 100 ppm can
inactivate enzymes extracts taken from plant cells. However,
the same results have not been obtained using living plants and
PAN in the parts per billion range.
The response of a given species or cultivar of plant to a
specific air pollutant, e.g. PAN, cannot be predetermined on the
basis of the known response of related plants to the same pollu-
tant. Neither can the response be predetermined by the given known
response of the plant to similar doses of a different pollutant.
Genetic susceptibility and environmental influences must, therefore,
be determined for each plant and pollutant.
PAN has been shown to be capable of inhibiting growth in
certain microorganisms, but extensive studies have not been made.
An evaluation of economic loss based on plant damage or
damage to certain sectors of the biosphere is extremely difficult.
An assessment of the economic impact of air pollutants on vegetation
166
has been published. The authors estimate that crop loss due to
oxidant air pollution in 1969 was $77 million and that the loss
in ornamentals was $42 million. The primary losses have been in
Southern California where smogs are most common.
PAN is discussed in some detail in "Air Quality Criteria for
Photochemical Oxidants", AP-63. U. S. Dept. of Health, Education,
l!67
and Welfare.
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6.3 MATERIALS /C ^ C/7£
6.3.1 . Laboratory
Corrosion scientists have long known the important role particles,
which settle out from the air, can play in promoting and accelerating
the corrosion rate of many metals. Only a few studies,
however, have been conducted on the effects of individual chemical
compounds including nitrates. Researchers have found that nitrate parti-
cles, as evidenced by individual studies of sodium nitrate and ammonium
nitrate, accelerate the corrosion rate of mi Id steel when exposed under
168
conditions of high relative humidities (70 and 99 percent). The magnitude
of corrosion-acceleration was, however, considerably less than that
caused by chlorides and moderately less than for sulfates.
Bell Laboratories chemists investigated the cause of stress
corrosion cracking of nickel-brass Ca copper-zincrnickel alloy} wire
spring relays, a problem unique**) California and especially the Los
169,170 •
Angeles area. Field observations and tests Csee 6.3.2) led the
scientists to suspect that some component in particulate matter was
responsible for the failures. Subsequent laboratory studies showed that
of all the inocganic salts evaluated, only nitrates were capable of
causing cracking in test wires at 75 oercent °r lower relative humidity
Furthermore, the rate of crack growth was directly dependent on nitrate
concentration, stress, relative humidity, and temperature. Increases
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in any of these variable increased the growth rate. A positive-
applied electrical potential (independent of magnitude) also increased
crack growth and was necessary at nitrate concentrations (up to about
o 171
15 vg/cm of surface area) found in Los Afigeles,
6.3.2 Field
In 1959 the Pacific Telephone and Telegraph Company noticed
considerable breakage of their nickel-brass wire springs in some relays
169
located in Los Angeles area central offices. Failures often occurred
within 2 years after installation. These failures were totally unexpected
since the wire springs had been used with excellent results for years
throughout the nation. Bell Laboratories investigated the problem,
They noted that breakage occurred on wires that were under moderate
stress and a positive electrical potential, and that atrborne dust
had accumulated on surfaces adjacent to broken wires, They concluded
that the failure mechanism was a form of stress corrosion cracking and
that some component in the dust may have been responsible. The invest!-*
/
gators exposed a number of wire spring assemblies to unftltered
Los Angeles air or filtered air. Cracks began to show up after six
months in the unfiltered afr but not in the filtered atr. An analysis
of Los Angeles dust revealed a nitrate content from 5 to 15 times greater
than in dust from most eastern and midwestern cfties, Subsequent
laboratory studies (see 6.3.1.) showed that high nitrate concentrations
in airborne dust caused the failures,
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The Investigators also surveyed the nitrate accunriiUtlon on equip-
171 -
ment located at 81 southern California locations. They found that
nitrate deposition correlated with relay failure. As a result, the
telephone company was able to develop a rating scheme to estimate the
degree of danger to equipment.
Several steps were taken to correct this cracking problem, On
new parts, wire springs were made from a zinc free copper-nickel
alloy, which is not vulnerable to nitrate caused stress corrosion. In
high nitrate aCeas, existing relays in central offices were protected
by installing high efficiency filters on outside air intakes and by
holding relative humidity below 50 percent.
Bell Laboratories have also reported a totally different type of
corrosion problem that has been observed in such widely scattered
locations as Cincinnati, Cleveland, Detroit, Los Angeles, New York,
171
and Philadelphia. The nickel *brass palladium-capped contacts of
crossbar switches corroded forming bright greenish corrosion products
that gradually crept up over the palladium cap of the contacts, resulting
in electrically open circuits. Investigators concluded that the
"creeping green" corrosion was promoted by the presence of anions,
principally nitrates, 1n accumulated dust.
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00 "Of QUOTE OR
6.4 VISIBILITY
A major effect of nitrate air pollutants is visibility reduction.
Visibility loss was the earliest noted aspect of photochemical
air pollution, and it is still its most prominent feature. Analyses
of fine particulate matter collected in Los Angeles smog show mainly
nitrates and sulfates in solution. Organic materials are the
third major constituent. The fine particles are seen because they
scatter light so much more strongly than the larger particles.
Nitrate concentrations are found to exceed sulfate concentrations
in areas where auto exhaust pollution is predominant, but sulfate
exceeds nitrate where fuel oil or coal burning is predominant. In
all areas where photochemical activity is able to reach an advanced
stage, the nitrate concentration in the fine particles increases as
the smog develops and the visibility goes down.
The detailed chemistry of the nitrates is not yet fully
understood, but here is good evidence that the reaction of NC^
and 0-j to produce NpOg leads to fine particle formation and growth.
The N205 is gaseous, but it hydrolyzes at surfaces to yield
nitric acid solution. The nitric acid may then react'with other
solutes, producing nitrate salts in solution. Development of more
detailed knowledge of the photochemical reactions leading to
nitrate particles is a major object of the program element on
Formation and Decay of Air Pollutants.
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J. Appl. Chem. (London). 6_: 26-44, January 1956.
169. Hermance, H. W. Combatting the Effects of Smog on Wire-Spring
Relays. Bell Lab. Rec. 48-52, February 1966.
170. McKinney, U. and H. W. Hermance. Stress Corrosion Cracking
Rates of a Nickel Brass Alloy Under Applied Potential Stress
Corrosion Testing. In: Stress Corrosion Testing, ASTM STP 425,
Philadelphia, American Society for Testing and Materials, 1967.
p. 274-291.
171. Hermance, H. W., C. A. Russell, E. J. Bauer, T. F. Egan and
H. V. Wadlow. Relation of Airborne Nitrate to Telephone
Equipment Damage. Environmental Science and Technology 5_:
781-785, September 1971.
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7. CONTROL TECHNOLOGY AND REMEDIAL ACTIONS
7.1 STATIONARY SOURCES
To insure that those atmospheric nitrates which may have serious
health or welfare effects can be controlled, an active control tech-
nology program must be pursued which aims at primary fine particulate
and the gaseous precursors of secondary nitrates.
7.1.1 Combustion Processes
7.1.1.1 Source Categories—The major man-made source of nitrogen
oxide (NO ) emissions to the atmosphere is from the combustion of
J\
fossil fuels. In general, this can further be subdivided into
stationary and mobile combustion sources with each class contributing
approxmately one-half of the total NO emissions.
J\
The major categories of stationary combustion processes are:
utility,, industrial, and commercial steam raising boilers; residential
heaters; industrial process furnaces; gas turbines; and internal
combustion engines. For each of these categories there are sub-categories
based on equipment design and these sub-categories must be further
categorized depending on the type of fuel used. Stationary combustion
processes are fired with a variety of fossil fuels (via, coal, heavy
oil, light oil, and natural gas), dependent on user and equipment
design. The control methods which can be applied vary, depending on
each unique categorization.
7.1.1.2 Formation Mechanisms—To place the various control technologies
in the proper context, it is necessary to understand the two mechanisms
by which NO is produced in a combustion process. The first mechanism
X
is fixation of atmospheric nitrogen. The rate of fixation has an ex-
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ponential dependence on temperature and a lesser dependence on oxygen
concentration. Therefore, a reduction in peak temperature offers a
method of control of this "thermal NO". Thermal NO is formed in all
A
combustion processes, but predominates for cleaner fuels such as
natural gas and light oil. The second mechanism is the oxidation of
chemically bound nitrogen in the fuel. The rate of oxidation is
apparently nearly temperature independent, but is strongly dependent
on oxygen availability. Therefore, a limitation of oxygen avail-
ability in the flame zone offers a method of control for this "fuel
NO". There is considerable evidence that for combustion of nitrogen
bearing fuels (i.e. - heavy oil and coal), a minimum of 50% of the
total NO created is from conversion of the bound nitrogen. These
y\
two distinct formation mechanisms obviously have a direct bearing
on the control strategy to be employed.
7.1.1.3 Control Categories—The various approaches to control of
nitrogen oxides can be categorized as either temperature reduction
or oxygen limitation. The temperature reduction techniques include
flue gas recirculation and water injection. They involve introduction
of an inert diluent into the flame zone to reduce the peak combustion
temperature. These techniques can be quite effective for reduction
of thermal NO. The oxygen limitation techniques include staged
combustion, low excess air and burner design changes. These methods
involve creation of local fuel rich regions in the early flame zone
where fuel nitrogen is normally oxidized. This either prevents
formation of NO and/or reduces the NO which has been formed. The
products of the rich combustion are then mixed with excess air to
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complete oxidation and heat release. These two techniques have
the potential for controlling both fuel and thermal NO from all
fuels. Both approaches have been studied extensively in experi-
mental systems, but work is still needed to define the specific
techniques which are most effective.
7.1.1.4 Control Limitations--A1though the general techniques are
well known, application of these techniques to practical combustion
systems is still uncertain. Most applications until now have
been retrofits to existing equipment and they do not approach the
optimum level of NO control which should be achievable. Appli-
cation of the techniques is difficult because they must be applied in a
manner to .prevent increase in emissions of other pollutants (i.e. - CO and
smoke) and losses of system efficiencies. Further development work
is needed, particularly for heavy oil and coal-fired systems. In
addition, the impact of various alternate fuels on the control tech-
niques must be assessed.
7.1.1.5 Control Application Experience—The implementation of com-
bustion modification techniques to date has been almost exclusively
in the area of utility boilers. The Los Angeles APCD has had NO emission
/\
standards for existing gas- and oil-fired utility boilers for some
time (until December 31, 1974, 225 ppm for gas and 325 ppm for oil).
These standards are being met using combinations of staged combustion,
low excess air and flue gas recirculation. In addition, the Los Angeles
APCD has a regulation limiting new units, regardless of size, to
140 IDS. per hour, which means that a 315 Mw boiler can emit about
40 ppm NO (at 3 percent 02). A unit (Scattergood 3) has been installed which
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is projected to meet this limit firing natural gas. This unit
employs the full range of available combustion modification tech-
niques and, if successful, will significantly advance the state
of the art of application to gas-fired utility boilers.
Federal standards have been promulgated for new (ordered after
August 1972) utility boilers greater than 25 Mw for all fuels. The
approximate levels are 168 ppm NO for natural gas, 230 ppm for oil,
and 500 ppm for coal (all concentrations are at 3 percent 02, dry basis),
Boiler manufacturers are currently selling units guaranteed to
meet these standards. Staged combustion, low excess air and burner
design changes are being combined to meet these regulations. The
principle area of uncertainty is for achieving the standards on
coal wall-fired boilers.
Federal standards are being proposed for stationary gas turbine
engines fired on natural gas and oil. The level of control has not
been finalized, but will probably contain an allowance for chemically
bound nitrogen. Currently, compliance is anticipated to be achieved
by water injection to reduce thermal NO.
For the other sources, industrial and commercial boilers, resi-
dential heaters, industrial process furnaces and stationary internal
combustion engines, standards have not been proposed. The R&D for
each of these sources is in various stages of progress and will influ-
ence both standards and implementation schedules.
As previously mentioned, the combustion modification techniques
employed for control of NO may adversely affect emissions of other
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pollutants ( via - CO, HC, and particulates) as well as system
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efficiency. However, since much of the data has been generated on
existing field equipment not designed for application of these
techniques, the significance of these adverse effects must be
carefully assessed. There is good evidence that a system initially
engineered for NO control can also achieve emission standards for
/\
other pollutants at equal or greater efficiency.
7.1.2 Industrial Processes
The primary sources emitting significant quantities of primary
nitrates are probably nitrate fertilizer prilling towers and bulk
blending plants. Analytical data documenting the amount of nitrate
actually emitted are not currently available. The particulate nitrate
emitted by these processes can be controlled by conventional par-
ticulate collection devices, but generally, controls other than
cyclones are not used.
Nitric acid plants may emit up to 10 ppm of HN03 mist, but the
major cause of nitrates from this source is probably from atmospheric
conversion of NO to nitrates.
X
According to the National Academy of Engineering Report,
"Abatement of Nitrogen Oxides Emissions from Stationary Sources",
1972;
Relative to emissions considered nationally,
NO from chemical operations is quite small; but
J\
locally these emissions can be significant. Essentially
all these emissions are associated with the manufacture
or use of nitric acid. About 75 percent of the nitric
acid produced in the United States is consumed in
ammonium nitrate production. The remainder is used
in a variety of processes, with, manufacture of adipic
acid consuming 9 percent.
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Because of the high ratio of N02 to NO in
the stack gases from nitric acid plants, colored
plumes are visible at relatively low NO levels, on the
/\
order of a few hundred ppm or much less, depending
upon the stack diameter. The present method of
control of emissions from nitric acid plants is
primarily by catalytic reduction, using natural
gas. It appears that, in much of the current practice,
N02 is reduced only to NO. This is only decolorization,
not emission control. Complete reduction to N2 using
methane requires complete burnout of the 02 present,
more consumption of natural gas, and closer control
of the equipment. In addition, CO and hydrocarbon
emissions increase with increasing NO reduction.
Selective reduction NO to N2> using ammonia as a fuel,
has been described, but there is insufficient information
to determine whether this process is a practical alternative.
Scrubbing with caustic soda has long been practiced
with NO emissions from nitration reactions, but it presents
* /\
a disposal problem. Recent developments with molecular
sieves indicate that adsorption processes based on their
.use may be capable of reducing emissions to the 10 to
50 ppm level. In a nitric acid plant, the desorbed N02
can then be recycled to the absorption tower. Evaluation
of a molecular-sieve absorption process on a large demon-
stration scale is timely.
The quantities of NO emitted from nitric acid plants
and chemical operations are but a small fraction of the
total man-made emissions on a nationwide basis; but they may
comprise a significant local source of pollution. Technology
for "decolorization" (conversion of N02 to NO) by catalytic
reduction with natural gas is well established. Present
methods of abatement (reduction to N2) are available to reduce
NO from typically 3,000 ppm to 100-500 ppm, but require
careful control. Adsorption by molecular sieves or other
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adsorbents has been shown on a bench scale to result in
even lower emission levels but no commercially proven
process is yet available. Alkaline scrubbing is a proven
process, but would involve liquid waste disposal problems.
7<2 MOBILE SOURCES
7.2.1 Control Strategies
The control strategy for mobile source-derived nitrates is
for the control of emission of the gaseous nitrogen oxides which
become nitrate in the atmosphere. These strategies consist of
modifications to the existing engines of cars currently in operation
and the introduction of new engine types and control systems in
future cars. The National Academy of Sciences Committee on Motor
Vehicle Emissions has reported to EPA that they believe four
engine types can meet the 1975 NO emission standards. These include:
/\
(1) modified, conventional internal-combustion engines with an
oxidation catalyst; (2) the Wankel engine with a thermal reactor
and exhaust gas recirculation; (3) the diesel engine; and (4)
the carbureted stratified-charge engine. The NAS reports the
most favorable one appears to be the stratified-charge engine offered
by Honda. In official tests it easily met the 1975 NO standards
/\
for over 50,000 miles. The NAS says it also can meet the much stricter
NO standards for 1976.
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For older cars there are several methods availaDre^to reduce
their NO emissions. These include: (1) retrofit system for
A
changing the spark timing; (2) exhaust-gas recirculation to lower
combustion temperatures; (3) water injection method to lower
temperature; (4) reductive catalyst of Ni-Cu to reduce the NO
A
gases. There are drawbacks in the use of several of these methods
to reduce NO emissions, in that their use can produce conditions
A
which favor the production of higher CO and hydrocarbon emissions
which have emission standards to be met. For example, the operation
of an engine on the lean side of the air-fuel ratio consumes the
fuel with very little hydrocarbon or CO emission, but favors the
formation of NO gases. Also, the operation of an engine at high
A
temperatures reduces hydrocarbons but promotes NO formation. Thus,
A
there is to be a case of trade-off in the control of all the harmful
emissions formed under opposing conditions. An exception is found
in the very lean-misfire region, where emissions are low but the
combustion is not stable or reliable. Perhaps the best means of
controlling them all will have to be a combination of novel engine
design or modification with exhaust gas treatment of some form.
In a report on nitrogenous compounds in the environment by
2
the Hazardous Materials Advisory Committee, they warn that a
new, potentially serious source of NO emissions exists from the
A
new, more powerful models of jet engines used today, plus the
expanded use of jet-air traffic and the greatly increasing fleets
of jet aircraft today. They point out that NO emissions from jet
A
aircraft engines are already a significant source of NO in all
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large cities with a major airport. Thus, for complete NOX control
in the environment, controls on the automobile engine alone will
not be enough, otner significant sources must also be identified
and corrected.
7.3 REFERENCES
1. Report to EPA, Committee on Motor Vehicle Emission, National
Academy of Sciences, Feb. 1973.
2. "Nitrogenous Compounds in the Environment", by Hazardous
Materials Advisory Committee of EPA, Report No. EPA-SAB-73~001,
December 1973.
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8. SUMMARY OF NERC/RTP CURRENT RESEARCH ACTIVITIES RELATING TO NITRATES
A number of research projects currently are underway at the
National Environmental Research Center, Research Triangle Park,
either in-house or via contract, which will provide further insight
into the problems relating to nitrates in the atmosphere. These
are described briefly.
8.1 MEASUREMENT AND ANALYSIS
A study is underway to improve and optimize electrometric methods
for measurement of nitrate and sulfate in particulates, to evaluate
the effects of substrates and environmental conditions on collection
of nitrates on filters, and to evaluate laboratory analysis methods
for atmospheric nitrates. An ion selective electrode nitrate
monitor, designed to provide a real-time indication of nitrate con-
centration averaged over 1 to 3-hour periods, has been developed and
is now being evaluated.
8.2 HEALTH EFFECTS STUDIES
All ongoing CHESS studies include daily measurements of sus-
pended nitrate levels and consider adverse health effects associated
with this pollutant. Contracts have been awarded for studies on the
relative toxicity of the respirable fraction of total suspended
particulate (TSP), sulfates, and nitrates, and to study the effects
of controlled exposure to sulfates and nitrates on environmental
responses.
8.3 FORMATION AND DECAY OF POLLUTANTS
The research program in Formation and Decay of Pollutants has
included the formation, decay and transport of nitrates as one of its
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major areas of concentration. The majority of the projects 1n
this program can be expected to have some bearing on our under-
standing the nitrate problem. A number of projects are devoted
mainly to nitrates, and are discussed below:
(1) An intramural research program on nitrogen compound photo-
chemistry is being conducted in the EPA Research Triangle Park
Laboratory. At present the principal research facility being
used is a long path infrared photochemical reactor which is coupled
to a Fourier Transform Spectrometer. Information is being acquired
on nitrite photolysis, PAN production, nitric acid formation, and
other topics of interest. These intramural nitrate studies will
be extended to the large 500 cubic foot irradiation chamber which
is presently being put into operation. The experimental studies
are complemented by an intramural modeling program in which the
experimental observations are fit into the overall chemical kinetics
scheme by which pollution photochemistry is quantitatively described.
(2) Field measurements of nitrates are being carried out con-
currently with the intramural laboratory studies. First, there
is an intramural field program which was conducted in 1973 at
Pasadena, California, and will be continued in 1974 at Houston,
Texas, and St. Louis, Missouri. This program includes the measure-
ment of gaseous and particulate nitrates. Secondly, there is a field
program being conducted under contract with the Battelle Institute.
This program is specifically addressed to the problem of accounting
for the fate of all the nitrogen compounds emitted into the air.
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(3) A program specifically addressed to the chemistry of the
PAN compounds is being sponsored at the Illinois Institute of
Technology Research Institute. The program is aimed at develop-
ment of a better understanding of the fate of the PAN's in the
air.
(4) Several of the aerosol characterization projects within
the program are addressing the question of the nitrate component of
the aerosols. Grants at the University of California, the University
of Minnesota, Washington University, St. Louis, and the University
of Washington are all contributing to development of the knowledge
of nitrate aerosols. Under an interagency fund transfer from EPA,
the Brookhaven National Laboratory is studying the formation of
nitrates in power plant plumes.
(5) .An outdoor smog chamber facility is being established at
the University of North Carolina for study of aerosol formation
in polluted air. A special effort will be devoted to the identifi-
cation of the gas-particle interactions in the air. Nitrate
formation will come into this project strongly. A close cooperation
is being initiated between EPA personnel and the University of North
Carolina personnel. The EPA role will be mainly to bring its long
path infrared competence and facilities to bear on the study of the
gas-particle interactions, while the University of North Carolina
role will be the operation of the facility and the characterization
of the particulate reaction products.
(6) A pollution up-take project is being initiated intramurally.
In this study an attempt will be made to identify the rates at which
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nitrogen containing compounds can be lost to surfaces in urban and
rural areas. The gas-surface interaction facility of the Chemistry
and Physics Laboratory Materials Section will be applied in the
study. Initially the study will be addressed to the relative rates
of loss of various pollutants to surfaces under controlled laboratory
conditions.
8.4 CONTROL TECHNOLOGY
8.4.1 Stationary Sources^
The Control Systems Laboratory R&D program is composed of four
major research categories: (i) field testing and surveys, (2) process
research and development, (3) fuels research and development, and
(4) fundamental combustion research. Field testing and surveys includes
studies designed to determine what can be done today to control NO
/\
emissions. This work is conducted on commercial equipment and is generally
performed by research and development organizations familiar with the specific
combustion systems being studied, and often with the financial
and technical assistance of the manufacturers, users, and trade
associations. In addition to developing trends and providing
directional recommendations for industry to minimize emissions with today's
technology, the work also defines where the research and development efforts
should be concentrated by developing emission factors as a fraction
of equipment type and size, and fuel consumed. The field testing
and survey studies are the initial efforts in the development of
control technology and are designed to provide the state of the art
in control of NO emissions from today's commercial combustion systems.
/\
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8.4.1.1 Proces ses—Process research and development work involves studies
with commercial or prototype combustion systems to develop cost
and design information for the application of the optimum NO
J\
control technology to classical combustion systems (existing and
new). The specific objective in this area is to
develop design and operational guidance manuals that can be used
by manufacturers and users to control NO emissions by combustion
A
modification techniques. The results of the studies in this category
provide the basis for the demonstration of combustion control tech-
nology. As in Field Testing and Surveys, these studies involve
industry participation.
8.4.1.2 Fuel_s_--Fuels research development studies are conducted in
versatile experimental combustion systems and are designed to
develop generalized combustion control technology which is applicable
to the optimum control of NO emissions from the combustion of both
A
conventional fuels and fuels of the future (e.g., coal derived fuels). These
engineering research and development studies will provide the future goals for NO
control and will generate the necessary technology to be applied in
the Process Research and Development area.
8.4.1.3 Combustion —Fundamental combustion studies are basic investi-
gations of the chemistry and physics of pollutant formation in com-
bustion systems. The results of the work provide the fundamental
understanding required for the generation of optimum NO control tech-
/\
nology in the Fuels Research and Development studies and for the
general application of the optimum NO control technology in Process
A
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Research and Development. This understanding will be developed
from experimental studies In Idealized combustion systems. The
results will provide the fundamental basis for mathematical simu-
lations of pollutant formation and control during combustion.
8.4.1.4 Combustion Flue-Gas—Combustion flue-gas treatment processes
to remove NO are under development. The control of NO emissions by
X „ A
combustion modifications can be limited by special process requirements J
for very high temperatures, limited space restrictions, or possibly
the inherent fixed nitrogen content of the fuel. Flue-gas treatment
techniques, including selective catalytic reduction, aqueous alkaline
scrubbing, and other selective adsorption and absorption techniques,
are under development.
8.4.1.5 Nitric Acid--Nitric acid manufacturing processes emit small
quantities, but high concentrations of NO . By 1975, control using
A
molecular sieves will be demonstrated in a form which is applicable
to new and existing ammonia oxidation process nitric acid plant industry
sources. Molecular sieves should permit a reduction in NO emissions
A
by at least an order of magnitude below the level attainable with
existing technology.
8.4.1.6' Planned Research and Development—It is imperative that a
better data base for fine particulate, including specific species of
nitrates, from stationary sources be developed. A standard analytical
scheme is under development and will be utilized to develop this
data base during all future sampling of stationary sources - especially
as related to EPA demonstrations of control technology.
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8.4.1.7 Problem Areas—In order for the Agency to make decisions
with respect to the control of nitrates, it is necessary to establish
immediately a better data base on the physical and chemical charac-
terization of fine particulate emissions from stationary sources.
Strategy and methods for the control of gaseous precursors
of secondary fine particulate matter, as well as nitrates, emitted
by stationary sources must be developed. Before the magnitude of
the necessary effort here can be determined, we must know the level
of control required.
8.4.2 Mobile Sources
Research is currently underway to measure nitrate emissions
directly from raw vehicle exhaust. These efforts involve filtration
of particulate matter and wet chemical analysis for nitrate by
colorimetric methods. The sensitivity of this method is less than
one microgram per milliliter of extraction solution. So far, no
nitrate has been positively detected by this approach, due probably
to the sampling technique and lack of program coordination with the
vehicle testing operations. Both catalyst-equipped and conventional
automobiles are being investigated for direct nitrate emission and
the effect of various types of control devices on the nature of the
particulate emissions.
During the course of the above work, efforts were made to adapt
existing chemical analysis procedures to the routine measurement and
possible monitoring of vehicles for nitrate emissions. Some pro-
cedures show good results and have possibilities for the incorpora-
tion into a certification procedure for automotive exhaust.
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