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"INDUSTRIAL METABOLISM OF NITROGEN"

by

Robert U.AYRES*Vicki NORBERG-BOHM**

92/57/EP

* Professor of Environmental Economics, at INSEAD, Boulevard de Constance, Fontainebleau 77305Cedex, France.

** PhD Student, at Harvard University, Cambridge Massachusetts, U.S.A.

Printed at INSEAD,Fontainebleau, France

Abstract: Industrial Metabolism of Sulfur

This paper quantitatively reviews the sources, production processes and economic uses

of sulfur-based chemicals, especially sulfuric acid and its derivatives. It discusses the

environmental sinks for these chemicals. Finally, it reviews the production of sulfur

oxides (SOx) via combustion of hydrocarbon fuels and the reduction of sulphide metal

ores, notably copper, lead and zinc. Data is provided mainly for the U.S., with some

global estimates.

INDUSTRIAL METABOLISM OF NITROGEN

Robert U. Ayres' and Vicky Norberg-Bohmb

Introduction

Inorganic Nitrogen (N2) is a colorless, odorless, tasteless gas which constitutes almost 80%of the atmosphere. In this form it is chemically rather inert and inaccessible to livingorganisms, as well as other key chemicals in living cells. However, nitrogen is essential toorganic life. Amino acids are nitrogenous compounds that are the basic building blocks of allproteins. Similarly, the basic building blocks of DNA (thymine, cytosine, adenine andguanine) all consist of single or double rings of carbon and nitrogen atoms, with various sidechains. However, the most reduced form of nitrogen, ammonia (NH 3), is an alkali. It is alsoquite toxic to animals, although plants and some bacteria can metabolize it. The carbon-nitrogen combination (cyanide) also has many toxic forms, such as the insecticide methylisocyanate (MIC).

On the other hand, the most oxidized form of nitrogen, the nitrate radical (NO) is stronglyacidic. The natural tendency toward thermodynamic equilibrium, in an oxidizing environment,would be for reduced nitrogen in the -3 valence state (NH3) to be oxidized gradually to the+5 valence state (nitrate, NO3). The result would be an acid ocean, leaving an atmospherewith no free oxygen left in it. No life on such an earth we would be possible. Hence life itselfdepends on the continuing stability of a natural cycle — the nitrogen cycle — whichbiologically reduces oxidized forms of nitrogen. This cycle evolved over billions of years toa balanced and stable state. However, as will be seen, the cycle is now unbalanced as a resultof human activity.

Sources and Production of Fixed Nitrogen for Industry

Human interference in the natural nitrogen cycle began with the cultivation of legumes andthe deliberate recycling of animal manures to soil. However, the magnitude of thisintervention was small and remained so until relatively modern times. It would not beunreasonable to assume that the nitrogen cycle itself remained in balance, through most ofhuman history, although the total amount of biologically available nitrogen l in circulationhas probably increased gradually. In the last century, however, this increase has accelerated.For a century, beginning about 1840, sodium nitrate (saltpeter) was mined in Chile. Nitrogenfertilizer accounted for a minor part of the use of this material (13% in the U.S. in 1910);

INSEAD, Fontainebleau, France

b Harvard University, Cambridge Massachusetts, U.S.A.

1 Biologically available inorganic nitrogen corresponds with "odd-nitrogen", i.e. nitrogen compounds or radicalswith 1 or 3 nitrogen atoms combined with hydrogen, oxygen or carbon. Principal examples include ammo-nia/ammonium (NH3/NH4), nitrite/nitrate (NO2/NO3), and cyanide (CN).

Industrial Metabolism of Nitrogen R.U. Ayres & V. Norberg-Bohm September 24, 1992

most of it was used for explosives (41%), dyestuffs (12%) and for other chemicals (25%) with5% unaccounted for and 4% going to glass manufacturing. However, total Chilean productionpeaked at about 500,000 tons in the 1920's. It is interesting to note that in 1920 there werefive different anthropogenic sources of biologically available ("fixed") nitrogen, producing aworld total of just over 1.55 million metric tons of N-content 2. Of this, Chilean saltpeter (stillthe largest source) accounted for only 30%. The second largest (26.6% in 1920), wasammonium sulfate from by-product coking of coal. This was a process developed in Germanyin the 1880's by Koppers, which is still in use, although nowadays the main by-product iscoke-oven gas. In fact, about 0.86 kg of NH3 can be extracted from a ton of bituminouscoking coal [Russell & Vaughan 1976, Table 3.2], along with 178 kg of high-BTU gas (8800cubic feet). Thus, production of 1 mcf of gas also yielded up to 98 kg of NH 3. Later ammoniawas produced in larger quantities from the coking of coal for the steel industry, but still asa byproduct. Its output could not be expanded above the limited amounts available from coalusage.

The first practical process for direct nitrogen fixation, around 1905, was the cyanamidprocess, yielding calcium cyanamide (CaCN 2) from limestone, coke and atmospheric nitrogen.The calcium cyanamide is dissolved in water to form urea and ammonia. By 1920 the thirdlargest source of synthetic fixed nitrogen (20.9%) was in this form. This process was stillcommercially significant in the 1950's, but it was then obsolescent. The last cyanamide plantin the Western world closed in June, 1971.

The fourth source of fixed nitrogen in 1920 (19.8%) was synthetic ammonia (NH 3), made bythe so-called Haber-Bosch process, developed by the German firm BASF and commercializedin 1913. Karl Bosch adapted this process to industrial production, largely to meet the demandfor explosives during World War I. During World War II many ammonia plants wereconstructed to produce the nitrogenous materials for explosives (almost all of which arenitrogen-based) as well as fertilizers. The availability of substantial capacity for synthesizingammonia gave rise to a growing nitrogen fertilizer industry in the U.S. after WWII.

The fifth, and least important source in 1920 was the so-called Birkeland-Eyde process,developed in Norway and first commercialized in 1904. This process emulated the naturalprocess by which nitrogen is oxidized in the atmosphere in the presence of an electric arc (i.e.lightning). By 1920 this process was already obsolete, because of its dependence on verycheap electricity, and then accounted for only 1.5% of synthetic nitrogen.'

From a world total of 1.55 million metric tons (N) in 1920, the use of synthetic nitrogen hasskyrocketed, especially after the second world war. The world total for 1948-49 was 3.3million metric tons, rising to 4.7 million metric tons just four years later [Shreve 1956,Chapter 20]. 1989 world production of fixed nitrogen, almost all of it in the form of syntheticammonia, was close to 120 million metric tons, i.e. 99 million metric tons N-content. The so-called Haber-Bosch process is the basis for almost all the ammonia manufactured today. Theprocess involves direct combination of nitrogen and hydrogen at high pressure, in the presenceof a catalyst of iron oxide plus small quantities of cerium and chromium at high pressure and

2 For a detailed discussion and breakdown see A.J. Lotka [Lotka 1956, Chap. XVIII].

5 However, in an hypothetical future world in which hydrocarbons are scarce but electricity (from fusion reactorsor the sun) is inexpensive, the Birkeland-Eyde process could conceivably be revived.

2


temperature. There have been many improvements to the basic process utilizing variouscatalysts and increasingly high temperatures and pressures.

Before 1950 the major source of hydrogen for the synthesis of ammonia was the reaction ofcoal or coke and steam via the water-gas process. A small number of plants used waterelectrolysis or coke-oven by-product hydrogen. After 1950 the major source of hydrogen hasbeen natural gas obtained by steam reforming of natural gas (or, less commonly, oil, coal orlignite). The partial oxidation process is also used to produce hydrogen from natural gas andother liquid hydrocarbons. As of 1975, 75 to 80 percent of the world supplies of hydrogenfor the manufacture of ammonia came from steam-reforming of hydrocarbons, of which 65%was from natural gas. The importance of natural gas has continued to grow.

The original source of synthetic nitrogen chemicals has always been air, which is already 70%pure. It is fairly easy to purify the nitrogen further by letting the oxygen react with someother substance, leaving an oxide (H20 is the simplest) that is easy to remove by condensationor solvent extraction. Originally, the nitrogen was obtained either from a liquid-air separationplant or by burning a small amount of hydrogen in the synthesis gas. Modern ammonia plantseliminated the above process steps by the use of secondary reforming, a process in whichmethane is burned in air in the amount required to produce a 3:1 mole ratio of hydrogen tonitrogen synthesis gas.

The Haber-Bosch process as used in the U.S. today is based on the passage of a mixture ofair and natural gas over a catalyst at very high pressure. Under the required conditions themethane in the gas is cracked and the carbon reacts with oxygen to form CO 2 while thehydrogen and nitrogen combine in the presence of a catalyst to form ammonia (NH3). SeeFigure I. There are several variations of the basic process depending on (1) the separationof air to produce nitrogen or (2) the removal of the oxygen by combustion to leave a nitrogenCO2 material that can then be separated. While the wartime plants were mainly in the 200-300tonne/day range, in the 1960's the introduction of efficient single-stage centrifugalcompressors made possible a new industry standard of 1,000 tonnes per day. The newtechnology increased capacity and cut costs sharply and further expanded the market (SeeTable O.

There are several known processes for nitrogen fixation which are not currently exploitedcommercially. These include ionization and chemo-nuclear reactions to obtain oxides ofnitrogen, fixation of nitrogen as metal nitrides or di-nitrogen complexes of transition metals,and reducing nitrogen bound up in certain transition metal complexes to ammonia. Inaddition, biological fixation by nitrogen-fixing microbes through genetic engineering is being

3

Electro-lytic

hydrogen

61.0.del cm.

Hydrogenpurification

Nitrogen

Ammoniasynthesis

gas

Anhy-drous

AirLeas

coke-ovengas

Refineryoff-gas

Naturalgas

Airseparation

Feedstockpurification

I

Purified Purifiedfeedstock Hydrogen

r Coolie' H 9 ---1 Cooling H 9 ....• v v A liy • ,-- A

Cryogenic Partial—1110.

Steampurification oxidation reforming

Carbondioxide

Condeasat}r7N30

rr Hydrogen _I

combustion

Vent

1-y3

Electricity Tir CoMiag H ?

Ammoniasynthesis

fTe Sole•1

- /To SIGs] --IP.

Aquaammonia

Steam

A


Figure 1: Materials-Process Relationships for Ammonia SynthesisSource: IMuehlberg et al 1977; p. 301

4


developed for agricultural applications.

Table1000

I: U.S. Nitrogen Productionmetric tons contained nitrogen

Capacity Production

1989 13136b 12606"1985 1297r 12965'1980 15501d 14736'1975 132341 12346'1970 11916h 10461'1960 4371h 3628g1950 1420h 118981945 1204h 410g1939 340h 244g1929 192h 76g1919 28g1909 13g1899 lg

• Units differ in the original sources and have been converted. Sources as follows:

a. [MINYB 1989 Table 2, p.7421 b. [MINYB 1989 Table 4, p.7441c. (MINYB 1985 Table 4, p.727/ d. [MINYB 1980 Table 4, p.601]e. (MINFP 1985 Table 7, p.5611 f. [MINFP 1975 Table 1, p.7501g. IHISSTAT 1975; Series P249+p250I

h. (HISSTAT 75; Series P315I

Major Industrial Uses of Nitrogen

Virtually all nitrogenous chemicals are now manufactured from ammonia. A list of thesechemicals follows.

Acrylonitrile (ACN), or Vinyl Cyanide CH2CHCN, made by reacting propylene withammonia and air, is the major intermediate for manufacturing acrylic fibers, ABS andSAN plastics and nitrite rubber.

Ammonium Chloride NH4CI, made by reacting ammonia and hydrochloric acid orammonium sulfate and sodium chloride, is used in dry cells, soldering flux, tanning anda number of other uses.

Ammonium Nitrate NH4NO3, made by reacting ammonia and nitric acid, is the mostimportant nitrogen fertilizer and also an important ingredient in explosives.

Ammonium Phosphates are major forms of fertilizer and feed additives; di-ammonium =(14144)2BP°4 and mono-ammonium = (NH4)H2PO4

Ammonium Sulfate (NH4)2SO4 was largely recovered from by-product (aqueous) ammonia,e.g. from coke oven gas, by sulfuric acid scrubbing. It is a major fertilizer.

Trends

5


Aniline C6H5NH2, made from nitrobenzene, is an important intermediate in the manufactureof isocyanates (urethanes), rubber chemicals (nitriles), aniline dyes, hydroquinone, drugsand miscellaneous chemicals.

Ethanolamines known as MEA, DEA, TEA, (mono-CH2OH.CH2NH2; di-(CH2OH.CH2)2NH;tri-(CH2OH.CH2)3N), are made from ethylene oxide and ammonia. They are used indetergents (DEA), gas conditioning (MEA) and petroleum products, emulsion polishes,herbicides, etc.

Ethylenediamine NH2CH2CH2NH2 is the basis of carbamate fungicides.

Hexametbylenetetramine (CH2)6N4 is used as a phenolic thermosetting catalyst and anexplosive (in RDX).

Hydrazine N2H4 is used as a rocket fuel.

Melamine C3N3(NH2)3, made from urea, is an important thermosetting plastic (used withformaldehyde), in laminates, textiles, adhesives, etc.

Nitric Acid HNO3, made by partial oxidation of ammonia, is a strong acid used in manyprocesses, especially in the manufacture of ammonium nitrates, adipic acid, nitroben-zene, isocyanates and numerous other products.

Nitrobenzene C6H5NO2, made from the reaction of nitric acid with benzene, is used inexplosives and to manufacture aniline.

Urea NH2CONH2 NH2CONH2 is a fertilizer made from ammonia and carbon dioxide. It isalso used in manufacturing melamine plastics.

In contrast to the situation in 1910, noted above, 80% of anthropogenic fixed-nitrogen in theU.S. today (90% in the world as a whole) is used in fertilizers. The remaining 20% is usedin 2500 other industrial chemical products, from explosives to plastics (melamine) andsynthetic fibers (nylon). A breakdown of nitrogen consumption by end use for the UnitedStates is given in Table II. A breakdown of nitrogen consumption by major chemicals andthree aggregated end use categories is given in Table III [Considine 1974]. Major nitrogenouschemicals and their key end uses are shown in Figures 2 through 6 which are discussedbelow.

Figure 2 shows the production of ammonia and three of its key derivatives: nitric acid, ureaand ammonium sulfate (more process detail on the production of ammonia was shown inFigure 1). Figure 3 shows the production of fertilizers. Figure 4 shows the production ofresins. Figure 5 shows the production of explosives. Figure 6 shows the production of othernitrogenous compounds and their key end uses. In these charts, products are indicated bycircles and processes are indicated by rectangles. The numbers in the process boxescorrespond to the numbers given for the process information presented in Appendix A. Figure7 shows the major chemical uses in quantitative terms.

6


Table H: U.S. Nitrogen Consumption by End Use1000 metric tons NH3 or equivalents

Use 1960' 1964' 1971' 1975' 1985d 1989d

Fertilizers 3110 5330 8870 1090 8560 9450Explosives 120 250 500 690 490 700Synthetic Fibers & Plastics 210 480 890 1410 420 470Chemicals 160 360 670 750Pulp & Paper 30 40 80 90 690 970Metallurgy 30 40 60 80Animal Feeds' 50 140 400 580 150 110Other" 300 370 510 540Export 130 140 540 270

Total 4140 7150 12520 14590 10300 11520• Units differ in the original sources and have been converted.a Primarily as urea.b Refrigeration, rubber, water treatment, detergents, textiles, dyes, etc.C Source: [Slack & James 1979]d Source: [MINYB 1989]

Table III: Major Areas of Ammonia Consumption"

Percentage of Consumption

Fiber & PlasticConsuming Areas Fertilizer Intermediates Non-fertilizer Total

Ammonia-direct application 26.6 - - 26.6Ammonium nitrate (AN) 17.0 - 4.0 21.0Nitric acid for non-AN uses 1.1 1.6 4.2 6.9Urea 10.1 0.8 1.6 12.5Ammonium phosphate 11.2 - - 11.2Ammonium sulfate 5.9 - - 5.9Nitrogen solution & mixed fertilizers 4.4 - - 4.4Acrylonitrile - 2.2 - 2.2Hexamethylenediamine - 0.7 - 0.7Amides & nitriles - - 0.6 0.6Caprolactum - 0.2 - 0.2Losses (transport, handling, storage) - - - 3.0Miscellaneous other uses - - 4.8 4.8

Percentage of market 76.3 5.5 15.2 100.0

' Based on practice in the United States.Source: [Considine 19741

7

Industrial Metabolism of Nitrogen R.U. Ayres & V. Norberg-Bohm

September 24, 1992

Figure 2: Production of Ammonia & Some Key Derivatives: Nitric Acid, Urea & AmmoniaSulfate

Source: [Ayres, Norberg et al 1989; Figure 5.11

8


Figure 3: Production of Nitrogenous FertilizersSource: [Ayres, Norberg et al 1989; Figure 5.2]

9


Figure 4: Production of Nitrogenous ResinsSource: [Ayres, Norberg et a11989; Figure 5.31

1 0


September 24, 1992

Figure 5: Production of Nitrogenous ExplosivesSource: [Ayres, Norberg et al 1989; Figure 5.41

11

INSECT-ICIDES

OL-AMINES

GASSCRUB-

BING

REACTIM OrETRYEEVECIEZOE •AQUEOUS AMMONIA


Note: Numbers in proms boxes refer to the possess descriptions in Appendix A.Letters is arrows toomet with eonespoodiog lass. in Figures 4 amnesia 8

A

PETROREFINING

WATERCONDI-

TIONING

YLENAMINES

FIBERSTEXTILES

SOLVENTS

AMMON OP la1714NCEREOIONIA

METHYL-AMINES

PHARMACEUTICALS

RUBBER

AMMONIA

NH3 NITRA2701I OPRENEW

NTTRO-BENZENEC6H5NO2 REDUCTION OF

NMIOSENZENE

4\s„.... DEGUSSA PROCESSESANDRVSSOW &

1 UREAI CO 1\.(NH2)2

YDRO-GEN

CYANIDEHCN

HEI4ICINTERME-DIATES

SOLVAY S0114.4StiPROCESS

AMMO-N.m

FAum DOME DECOMPOSITION

SULFATE OF AMMONJUM SUZPA1ESUL & SODIUM' CilLORIDE\(N144)2" 46

Figure 6: Production of Important Nitrogenous Compounds for End Uses OtherFertilizers, Resins & Explosives

Source: [Ayres, Norberg et al 1989; Figure 5.5]

than

12

natio:5,

Synthetics,Adipic Acid(CH2)4

\ (COOH)2 ubber

losive

Use,

Rocket

HydrazineN2H4

0.018

Batteries

matensect-

lcides

AmmoniumChloride,

OtherChemicals Ethanol-

aminesMEA, DEA,

TEA

0.16 / Ethylene-diamine

NH2CH2• CH2NH2

Isocyanates

AcrylonitrileCH2CH3CH

AMMONIANH3

14.0Aniline

C5H5NH2

Melamine

C3N3(NE,

NitrobenzeneC6H5NO2

NitricAcid

HNO3Other

Explosives

Fuels

UreaNH2CONH2

ss• Sulfate(NH4)204

2.0 AmmoniumPhosphates

(NH4)2PO4 &NH4)H2PO4

Direct

Ammonia

Deter-gents,

1°- GasTreatment

Fertilizer

Use

► (76%)


Figure 7: U.S. Ammonia Flows 1980 (million metric tons)Source: [Ayres et al 1988, Figure 10.4]

13

AmmoniaAmmonium nitrateAmmonium sulfateDi-ammonium phosphateMono-ammonium phosphateCalcium nitrateCalcium cyanamidePotassium nitrateSodium nitrateUreaUrea-formaldehyde

NH3N1140NO3(NH4)2SO4NH4H2PO4

% Nitrogen

82.2436.6521.2021.1412.1817.0734.9713.8516.4847.0031.00

(NH4)2HPO4Ca(NO3)2.4H20CaCN2KNO3NaNO3CO(N112.)2CO(NH2)2HCHO


Agricultural Uses and Losses of Nitrogen

As mentioned, nitrogen is an essential element for protein production. It can only be ingested(by plants) as soluble ammonium (NH-4) ions, nitrate (NO;) ions, or soluble organic moleculessuch as amino acids. There are many routes by which soluble nitrogen finds its way into thesoil.

Nitrogen oxides are created in the air either by electric storms or fuel combustion. Theseoxides then react with water vapor and create nitrous or nitric acid (HNO 3) which falls as acomponent of acid rain. Nitric acid subsequently reacts with metallic salts to form nitrates.Ammonia is also emitted by volcanoes.

Rhizobia bacteria in the root nodules of legumes (such as soybeans) are able to fix nitrogenfrom the air. Some other varieties of bacteria can also fix nitrogen. A third source - from apractical point of view - is the recycling of plant and animal matter. Some of this nitrogenfrom decaying organic matter is released in the form of ammonia (NH 3), and, along withammonia from anthropogenic sources, is returned to the soil in rain.

For centuries nitrogen fixation by atmospheric electrical discharges, from legumes and fromdecayed matter (manure) were the only means of restoring lost fertility to soil. Gradually, therecycling process was accelerated through composting and distribution. Organic matter in thesoil may contain about 5-6% nitrogen, but this is not available to plants until it has beenreleased by bacteria or fungi in the form of ammonia or nitrates. Soils with plentiful organicmatter can retain much of this for future use by plants.

With increasing pressure on farmers to increase crop yields the available nitrogen in the soilwas used more rapidly than it could be replaced by the available plant and animal materials.In 1930 the annual nitrogen deficit for the U.S. was estimated at 3.4 million tonnes, despitesome use of nitrogen fertilizers [Shreve 1956, p. 397]. However, fertilizer use has increasedvery rapidly since then.

Table IV: Principal Nitrogenous Fertilizer Materials

Source: (Ayres et al 1988; Table 10.21

Globally, about 90% of all ammonia (76% in the U.S.) is used for fertilizer (see Figure 7).In recent years a growing proportion of nitrogen has been applied by direct injection of

14


anhydrous ammonia into the soil, particularly in the large farms of the Mississippi valley.However, the major portion is still applied either in solid or liquid form that can be handledmore easily and applied in combination with other plant nutrient materials. The principalprimary nitrogenous fertilizer materials are shown in Table IV which also shows thepercentage nitrogen contained in the product.

These products may be applied either as straight materials or in various combinations knownas mixtures which are made up to a standard set of specifications. By custom and law theseare always reported as percentage of nitrogen - phosphate - potash. Because of the growth inthe direct application of high nitrogen materials there has been a tendency for mixtures tostabilize nationally in the 21-24 million ton/year range while total application continues togrow (see Table V).

Table V: U.S. Commercial Fertilizers Consumed & AverageNitrogen Content

Year

Mixed Fertilizers

Gross Amount Percentage1000 metric tons Nitrogen

All Fertilizers with PrimaryNutrients

Gross Amount Percentage1000 metric tons Nitrogen

1968 9,659 8.77 16,840 18.281969 9,672 8.95 17,061 18.501970 9,508 9.25 17,369 19.481971 9,758 9.59 18,099 20.381972 9,757 9.91 17,643 20.111973 10,227 10.21 18,970 19.831974 10,917 10.10 20,395 20.371975 9,365 10.16 18,419 21.181976 10,414 10.53 21,270 22.201977 10,931 10.58 22,271 21.691978 10,029 10.74 20,693 21.821979 10,769 10.84 22,381 21.721980 10,555 10.79 22,902 22.591981 10,671 10.99 23,478 23.041982 9,442 11.23 21,309 23.58

Source: [Ayres et al 1988; Table 10.3)

In well ventilated soils having low organic matter (with a carbon nitrogen ratio of less than15) when the quantity of ammonium ions exceeds the absorption by plants and microbes theexcess ammonia is oxidized to nitrates by a nitrification process. The first step (to nitrite) isbrought about by bacteria such as nitrosomas. The second step (to nitrate) is carried out byother bacteria such as nitrobacter. These nitrates remain in the soil for plant use, up to certainlimits depending on temperature, etc. Most nitrate compounds are soluble and can be leachedfrom the soil - particularly in permeable soils having low organic matter and high rainfallIn clogged soils having insufficient air circulation, fertilizer nitrates may be denitrified(reduced to gaseous nitrogen) by anaerobic microbes which break up the molecules to satisfytheir demand for oxygen. However, in soils with high levels of decaying organic matter, aircirculation, and a high microbial population, the microbes consume the nitrogen in the formof ammonia and effectively store it for future use by plants - resulting in very little loss by

15


leaching. Bacterial action also continuously removes nitrogen from soils altogether byconverting organic nitrogen to volatile NH3, some of which escapes into the atmosphere.

Crutzen has estimated the partitioning of 100 units of agricultural fertilizer into ultimate sinksas follows [Crutzen 1976, cited in USNAS 1978]: (1) accumulation in the soil, groundwateror sediments: 38 units; (2) active recycling through the soil-biosphere system as manure orplant residues: 39 units; (3) denitrification to N2 or N20 which return to the atmosphere: 23units. In this picture about 17 units escape as ammonia, 4 from the point of fertilizerapplication, 12 from manure and 1 from sewage. Another 6 units escape as N2 (5.6 units) orN20 (0.4 units). Crops initially take up 50% (50 units) of the N-content of the fertilizer; ofthis 47 units are consumed by livestock, and 42 units end up in manure and animal urine, ofwhich 15 units go back to the soil, along with 30 units of nitrogen from plant residues and4 units from local deposition of gaseous emissions. But, during each cycle, 10 units of organicN are lost to waterways and ground water, so the amount of organic N "permanently" added(i.e. for the next cycle) is 15 units, while 24 units of inorganic N are also added for the nextcycle. This represents accumulation within the terrestrial biosphere and cultivated lands.Crutzen's allocation neglects NO emissions from soil, which are not well documented, butknown to be significant [Williams et al 1992].

Obviously a significant fraction is taken up by plants and (unless the crop is 'plowed under)much of it is removed from the soil when the crop is harvested. Of course, the specificproportions removed vary from crop to crop and depend on the ratio between what is suppliedand what is required. Analyses have been made as to total removal by crops in the U.S. andit is interesting to note that USDA's figures do not agree with Crutzen's. During the 1970'sthe average N removal in the U.S. was 72.4% of the fertilizer nitrogen applied. This figurevaries somewhat with individual years, from a high of 78.5% in 1973 to a low of 64.0% inthe following year.

The nitrogen removed by harvested crops is partly consumed in food, either directly, as infresh vegetables, or after conversion to meat, eggs, etc. The average per capita foodconsumption comes to about 1000 pounds retail weight per year (exclusive of water contentof fresh milk and fruit juices). While many foods have above average protein (meat, milk andeggs) and other little or none (e.g. sugar), average nitrogen content is at least 2% (equal tothe average of harvested crops), or 20 pounds per capita per year. Nationally, this amountsto 2.5 million tons (2.25 million tonnes), which is only a small fraction of the total nitrogeninput as fertilizer. Most of the N-input must be lost by other routes e.g. from animal feed lotsor food processing plants such as beet sugar production, cheese-making, brewing, flourmilling and meat packing. .

As noted previously, crop harvesting (and erosion) removed nitrogen from the soilconsiderably faster than it was being replaced, probably until some time in the 1960's 4 Thatremoval process resulted in a substantial net flow of nitrogen from rural to urban areas(embodied in food products) from whence much of it was returned to the environment insewage and/or garbage.

When U.S. consumption of synthetic nitrogenous fertilizers presumably reached rough balance with annualremovals by harvesting and erosion. It is interesting to note that the National Research Council estimates a netnitrogen buildup in soil and groundwater of the order of 1.5 million tonnes per year (1972) [USNAS 1978].

16


However, due to the rapid increase in nitrogen fertilizer use, runoff from agricultural areasis now the dominant source of aggregate nitrogen emissions to the water (and has been sincethe 1960's). In addition, large-scale animal feed lots have become major point sources ofnitrogen emissions (both to water and air) by way of manure and urine s. In the extreme caseof a heavily urbanized region, most N-emissions can evidently be attributed to humanconsumption wastes (e.g. sewage).

Anthropogenic Sources and Emissions of Oxides of Nitrogen (N20, NOx)

Oxides of nitrogen are formed naturally by electrical discharges in the atmosphere. They arealso formed when either atmospheric or fuel-bound nitrogen are "burned" in air. This is themajor anthropogenic contribution. The major nitrogen oxides are N20, NO and NO2. In theabsence of fuel-bound nitrogen, typically 95-98% of the NOx in exhaust gas consists of NO.Incompletely oxidized nitrogen is oxidized further to NO2 in the atmosphere.

The reaction chemistry for NOx formation by fossil fuel combustion is not completely known,but the most probable sequence for thermal NOx seems to be the following [Lim et al 1981].

N2 + 0 —• NO + NN2 + 02 —• NO + 0N + OH —• NO + H

where oxygen and hydrogen atom concentrations are in equilibrium.

02 + N2 44. 0 + 0 + N2

H20 + 02 4-i• H ` + OH - + N2

The reaction rate is determined by two factors. The activation energy of the initialendothermic reaction step is quite large (317 KJ/mole) and probably determines the overallreaction rate. The reaction is, therefore, highly temperature dependent: equilibrium NO levelsin heated mixtures of N2 and 02 at atmospheric pressure is an increasing function of

temperature (Figure 8). The mathematical form e k/T, where k is a constant, has beensuggested [Mackinnon 1974]. Recent data for fluidized beds is shown in Figure 9.

Because of the high activation energy for the nitrogen oxidation reaction, it will only occurafter all available carbonaceous fuels are exhausted. Thus, the rate of NO x production alsodepends on the amount of excess oxygen that is present beyond the stoichiometric air/fuelratio (14.6 for the case of gasoline). The equilibrium NO level is also proportional to theproduct of N2 concentration times the square root of 0 2 concentration, times residence timein the flame. Because of these relationships, air/fuel ratio is a dominant factor in determiningNOx emissions in the absence of fuel-bound nitrogen. Figure 10 shows typical NOx (NO,

5 Animal urine accounts for at least 50% of excreted nitrogen, in the form of urea. Under typical feedlot

conditions urea is rapidly hydrolyzed to carbon dioxide and ammonia.

17


0 0.25 0.5 0.75 1

1.25 1.5

1.75NO Yield (%)

Figure 8: NOx Formation vs. Temperature, Based on Measurements at Furnace WallSource: [Ermenc 1956; cited in Engdahl 1968]

Bed Temperature (Degrees F)1400 1600

1800

700

800 900

1000Bed Temperature (Degrees C)

Figure 9: NOx vs. Bed Temperature, Equivalence Ratio 0.847 (18% excess air)Source: [USEPA 1979-e]

NO2) emissions as a function of air/fuel ratio for automobile gasoline engines [Jackson 1968].

For smoldering fires, or "rich" fuel air mixtures, NOx production is negligible. However for"lean" fuel air mixtures, it becomes increasingly significant. Thus, NOx is, for practicalpurposes, not significantly produced in low temperature combustion processes, except to the

18

(Each point = average of tthree determinations)I


September 24, 1992

3500

3000

zV 2500

2000

1500

O1000

500

010 12 14 16

18 20

22Air/Fuel Ratio

Figure 10: Influence of Air-Fuel Ratio on Concentration of NOx in Exhaust GasSource: [Adapted from Jackson 1968, Figure 5AI

extent that fuel-bound (organic) nitrogen is present.

However, organic nitrogen is present in coal and residual oil and in biomass6. Thus, fuel-bound nitrogen in all organic materials, as well as in fossil fuels (in coal and residual oil) alsocontributes to both N20 and NOR emissions. Apparently N20 — one of the most potent"greenhouse gases" — is associated only with fuel-bound nitrogen.

Median nitrogen content of US coals corresponds to approximately 1350 ng/J of NO2,assuming 100% conversion; the 25-75 percentile range is fairly narrow, viz. 1250-1400 ng/J[Lim et al 1981]. Figure 11 shows measured relationships between coal nitrogen content inpercent by weight and measured NOR emissions. The median NO R (as NO2) emission levelis about 270 ng/J. The 25-75 percentile range for bituminous coal combustion correspondsto NOR (as NO2) emissions in the range of 240-300 ng/J [Lim et al 1981]. Based on thisdata, it appears that if there were thermal contribution to NOR, the average conversionfraction for fuel-bound nitrogen to NOR would be 270/1350 = 0.2.

The relative contribution of fuel-bound nitrogen to NOR emissions from combustion is notyet completely understood. However studies have shown that the contribution of fuel boundN to total NOR may be as high as 50% for residual oil, and 80% for coal. Looking at itanother way, 20%-90% fuel-bound nitrogen in oil and 10%-60% in coal may be convertedto NOR, depending on conditions of combustion. No data on wood combustion is available.

6 In coal the range is between 0.6% and 2% by weight (Figure 11). In terrestrial biomass, it has been found thatthe nitrogen/carbon molar ratio is fairly constant (1:80), which means that the C/N weight ratio is 685. This is aboutthe same as the average ratio for coal.

19

0

_______ _ _______ ___ _ ___ ..... ....0 .................. ,c3..

0 .....

0....4 .

.............. A

--0 Bituminous

-A- Sub-bituminous

A AA

E 0.8

0 0.6

O

0.2


September 24, 1992

00.6 0.8 1 1.2 1.4 1.6 1.8

2

2.2Coal Nitrogen Content (% Weight, Dry, Ash-Free)

Figure 11: Nitric Oxide Emission as Measured vs. Coal Nitrogen ContentSource: [USEPA 1979-fl

Emissions vary widely according to the type of combustion equipment, so test data is limitedand generalizations are difficult to make. Taking into account the contribution fromatmospheric nitrogen, EPA has developed estimates of NO R emissions for most fuel uses[USEPA 1978 & undated]. The EPA emissions coefficients for coal burned in electric utilitiesand other industrial boilers, respectively, are about 11.5 kg/ton and 9.5 kg/ton. Natural gasand residual oil combustion emissions by electric utilities are slightly lower than emissionsfrom coal (10.95 and 7.29 kg/ton, respectively). Industrial boilers emit a little less NO R thanutility boilers, on the average, because they are smaller and operate at lower temperatures.Railroad uses of coal can be assumed to be very similar to industrial boilers. Fuel wood isused primarily in residences, or by the paper/pulp industry. High water content tends to keeptemperatures low and minimizes NOR.

Historically, the only important technological change affecting NOR emissions (apart fromautomobile emissions controls) has been the increased size and efficiency of utility boilers,especially since the first uses of pulverized coal in the 1920s. It may be noted that, as of1979, over 99% of U.S. coal-burning utility boilers used either pulverized coal (88%) orcyclones (11%). On the other hand, 60% of industrial boilers use pulverized coal while 40%use stokers. Only 14% of boilers in the commercial residential category (large buildings) usepulverized coal, while the rest use stokers or manual feeds. Western Europe probablyapproximates the U.S. pattern. However in Eastern Europe and Asia it is likely that stokersstill predominate. Use of pulverized coal roughly doubles (11.5/7.1) the NO R emissionsrelative to the use of lump coal in stokers.

The technology of pulverization of coal was initially developed in the nineteenth century.Attempts were made in the late nineteenth century to burn pulverized coal under industrialboilers but they were unsuccessful until approximately 1910. In the twentieth century, agrowing concern over fuel economy stimulated attempts at innovation. Pulverized coal was

20

Industrial Metabolism of Nitrogen

R.U. Ayres & V. Norberg-Bohm September 24, 1992

rapidly adopted in electrical generating and industrial plants after World War 1. Designadvances that increased the efficiency of the furnaces by eliminating slag and improvingcombustion followed, as did improvements in pulverizers. By the early 1920s, pulverized coalwas producing a 3 to 5 percent gain in efficiency compared to firing with stokers as well asa savings in coal costs [de Lorenzi 1948 :9-32-43; Rosin 1947 :17]. Based on mostlyanecdotal historical evidence, we estimate the historical penetration of pulverized coal asshown in Table VI.

Table VI: Penetration of Coal Pulverization

Year Utilities Industrial Comm/Residential

1920 0% 0% 0%1950 50%-60% 25%-30% < 5% (?)1980 99% 60% 14%

Source: [Ayres et al 1987; Table 3.21EPA's national emissions estimates for NOx are given in Table VII. Emission coefficients forNOx are taken from EPA reports [USEPA 1978 & undated] and summarized in Table VIII.

Table VII: EPA Estimates of NO 2 Emissions 1940-1980 (Teragrams/yr)

Source 1940 1950 1960 1970 1980

Highway Vehicles 1.3 2.1 3.6 6.0 7.2Aircraft 0.0 0.0 0.0 0.1 0.1Railroads 0.6 0.9 0.7 0.6 0.8Vessels 0.1 0.1 0.1 0.1 0.1Off-Highway 0.2 0.4 0.5 0.8 1.0

Mobile Source Total 2.2 3.5 4.9 7.6 9.2

Electric Power 0.6 1.2 2.3 4.5 6.4Industrial Boilers *2.3 *2.9 *3.7 *3.9 *3.1Commercial/Institutional Heat 0.2 0.3 0.3 0.3 0.3Residential Heat 0.3 0.3 0.4 0.4 0.4

Stationary Combustion Total 3.4 4.7 6.7 9.1 10.2

Chemicals 0.0 0.0 0.1 0.2 0.2Petroleum Refining 0.1 0.1 0.2 0.2 0.2Iron & Steel Mills 0.0 0.1 0.1 0.1 0.1Pulp Mills - - - - -Mineral Products 0.1 0.1 0.1 0.2 0.2

Industrial Total 0.2 0.3 0.5 0.7 0.7

Incineration 0.0 0.1 0.1 0.1 0Open Waste Burning 0.1 0.1 0.2 0.3 0.1Forest Fires 0.7 0.4 0.2 0.2 0.2Other Burning 0.2 0.2 0.2 0.1 0

Waste & Miscellaneous Total 1.0 0.8 0.7 0.7 0.3

Grand Total 6.8 9.3 12.8 18.1 20.4

• Very questionable (see text)Source: [USEPA 1986]

21


September 24, 1992

Table VIII: 1980 NOx Emissions Coefficients, Uncontrolled (kg as NO2)

Fuel Type

Fuel

le JltonElectricutilities

Stationary

Industrialboilers

Resid./Commer.

Mobile

OffHighway Highway

Anthracite coal 26.6 N 8.25 8.25 N NBituminous coal 26.25 11.5 (1) 9.1 (b) 6.0 (c) N 6.0Lignite 18.0 7. (a) 5.5 (b) 3.4 (`) N NWood (12% H20) 15.5 N 0.75 0.75 N NResidual oil 48.96 7.29 7.02 7.02 N NHeating oil 48.19 2.86 2.86 2.72 N NDiesel oil 50.80 N N N 41.8 54.3Gasoline (d) 51.54 N N N 15.7 urban 17.8

17.8 ruralJet fuel (kerosine) 51.38 N N N N NNatural gas 53.40 10.95 3.13 2.23 N N

"N" = Not applicable.(a) For 1950 assume 8.85 for coal and 5.30 for lignite;

For 1920 assume 5.60 for coal and 3.20 for lignite.(b) For 1950 assume 7.10 for coal and 4.15 for lignite;

For 1920 and earlier assume 5.6 for coal and 3.2 for lignite.(c) For 1950 and earlier assume 5.6 for bituminous coal and 3.2 for lignite.(d) For 1950 assume 19.6 urban (70%) and 22.5 rural (30%), or 20.5 average;

For 1920 assume a composite coefficient of 10, because of lower compression ratios.Source: [Ayres et al 1987; Table 3.5]

References

[Ayres, Ayres & Tarr 1987] Ayres, Robert U., Leslie W. Ayres & Joel A. Tarr, An Historical Reconstructionof Major Atmospheric Gaseous Emissions in the U.S., 1880-1980, Technical Report, Variflex Corporation,Pittsburgh PA, 1987.

[Ayres et al 1988] Ayres, Robert U., Leslie W. Ayres et al, An Historical Reconstruction of Major PollutantLevels in the Hudson-Raritan Basin 1800-1980, Technical Memorandum (NOS OMA 43), NationalOceanic & Atmospheric Administration, Rockville MD, October 1988. [3 Vols]

[Ayres, Norberg et al 1989] Ayres, Robert U., Vicky Norberg-Bohm, Jackie Prince, William M. Stigliani &Janet Yanowitz, Industrial Metabolism, the Environment, & Application of Materials-Balance Principlesfor Selected Chemicals, Research Report (RR-89-11), International Institute for Applied Systems Analysis,Laxenburg, Austria, October 1989.

[Considine 1974] Considine, D.M., Chemical & Process Technology Encyclopedia, McGraw-Hill BookCompany, New York, 1974.

[Crutzen 1976] Crutzen, Paul J., The Nitrogen Cycle & Stratospheric Ozone, Nitrogen Research ReviewConference, United States National Academy of Sciences, Fort Collins CO, October 12-13, 1976.

[de Lorenzi 1948] de Lorenzi, Otto, Combustion Engineering, Technical Report, Combustion Engineering, Inc.,New York, 1948.

22


[Engdahl 1968] Engdahl, Richard B. "Stationary Combustion Sources", in: Stem(ed), Air Pollution: Sourcesof Air Pollution & Their Control, Chapter 32 :4-54 [Series: Environmental Sciences] DI, Academic Press,New York, 1968. 2nd edition.

[Jackson 1968] Jackson. "Re. NOR", in: Stem(ed), Air Pollution: Sources of Air Pollution & Their Control[Series: Environmental Sciences] III, Academic Press, New York, 1968. 2nd edition.

[Lim et al 1981] Lim K.J. et al. "NOx Combustion Modification", in: Martin(ed), Emission Control forIndustrial Boilers, Chapter 3 :222-279, Noyes Publications, Park Ridge NJ, 1981. [Reprinted from"Technology Assessment Report for Industrial Boiler Applications: NO R Combustion Mod." for EPA1979]

[Lotka 1956] Lotka, Alfred J., Elements of Mathematical Biology, Dover Publications, New York, 1956. 2ndReprint edition. [Original title: Elements of Physical Biology, 1924]

[Lowenheim & Moran 1975] Lowenheim, F.A. & M.K. Moran. "Nitrogen", in: Faith, Keyes & Clark (eds),Industrial Chemicals, :462ff, Wiley Interscience, New York, 1975, 4th edition.

[MacKinnon 1974] MacKinnon, DJ., "Nitric Oxide Formation at High Temperature", Journal of the AirPollution Control Association 24(3), March 1974.

[Muehlberg et al 1977] Muehlberg, et al. "Phosphate Rock & Basic Fertilizer Industry", in: Industrial ProcessProfiles for Environmental Use, Chapter 22(EPA-6002-77 023v), Government Printing Office,Washington DC, Dow Chemical Company & Radian Corporation, February 1977.

[Rosin 1947] Rosin, P.O. "The History of Pulverized Fuel as Reflected by its Economic & TechnicalProblems", in: Conference on Pulverized Fuel :13-44, The Institute of Fuel, London, 1947.

[Russell & Vaughan 1976] Russell, Clifford S. & William J. Vaughan, Steel Production: Processes, Products& Residuals, Johns Hopkins University Press, Baltimore MD, 1976.

[Shreve 1956] Shreve, R. Norris, The Chemical Process Industries, McGraw-Hill Book Company, New York,1956.

[Slack & James 1979] Slack, A.V. & G.R. James, Ammonia, Marcel Dekker, New York, 1979.

[Williams et al 1992] Williams, E.J., A. Guenther, & F.C. Fehsenfeld, "An Inventory of Nitric Oxide Emissionsfrom Soils in the United States", Journal of Geophysical Research 97(D7), May 20, 1992 :7511-7519.

[HISSTAT 1975] United States Bureau of the Census, Historical Statistics of the United States, Colonial Timesto 1970, Bicentennial Edition, Part 2, United States Government Printing Office, Washington DC, 1975.

[MINYB 1989] Bureau of Mines, United States Department of the Interior, Minerals Yearbook 1989, UnitedStates Government Printing Office, Washington DC, 1991.

[MINYB 1985] Bureau of Mines, United States Department of the Interior, Minerals Yearbook 1985, UnitedStates Government Printing Office, Washington DC, 1986.

[MINFP 1985] Bureau of Mines, United States Department of the Interior, Mineral Facts & Problems, 1985Edition, Bureau of Mines Bulletin 675, United States Government Printing Office, Washington DC, 1985.

[MINFP 1975] Bureau of Mines, United States Department of the Interior, Mineral Facts & Problems, 1975Edition, United States Government Printing Office, Washington DC, 1975.

[USEPA 1978] United States Environmental Protection Agency, Mobil Source Emission Factors, TechnicalReport (EPA-400/9-78-005), United States Environmental Protection Agency, Washington DC, September1978.

23


[USEPA 1979] United States Environmental Protection Agency, Methods of Controlling NOx in Combustion,Technical Report (EPA 600/7-79-178 e,f,g), United States Environmental Protection Agency, CincinnatiOH, 1979.

[USEPA 1986] United States Environmental Protection Agency Office of Air Quality Planning & Standards,National Air Pollution Emission Estimates, 1940-1984, Technical Report (EPA-450/4-85-014), UnitedStates Environmental Protection Agency Office of Air Quality Planning & Standards, Research TrianglePark NC, January 1986

[USNAS 1978] National Academy of Sciences (NAS/NRC), Nitrates: An Environmental Assessment, WashingtonDC, 1978.

APPENDIX A: NITROGEN PROCESSES

#1 Ammonium and Ammonium Sulfate from Coal Gas:

Ammonia can be recovered from coke-oven operation, principally as aqua ammonia. Onlya small amount of commercial ammonia, perhaps 1% is from this source. The ammoniain the coal gas reacts with sulfuric acid. In the high temperature coking process, 15 to 20%of the nitrogen leaves the oven as ammonia. There are three general methods formanufacturing ammonium sulfate. In each of them the hot gases from the ovens are pre-cooled and the gases are eventually passed through an acid bath saturator. About 25% ofammonium sulfate production is currently from this method. Until the middle of WWII,this was the major source of ammonium sulfate. An increase in demand for ammoniumsulfate fertilizers during the war caused the development of the synthetic ammoniumsulfate industry. [Lowenheim & Moran 1975; Slack & James 1979]

#2 Ammonia by Catalytic Synthesis from Nitrogen and Hydrogen (Haber Bosch Process):

N2(g) + 3H2(g) --0 2NH3(g) + heat

Haber Process: Nitrogen and hydrogen react in a 1:3 ratio in the presence of a catalyst athigh temperatures (200 to 700 degrees C) and pressures (100 to 1000 atmospheres). Thenitrogen is derived from the air by liquefaction, the producer gas reaction or by burningout the oxygen in the air with hydrogen. The hydrogen is obtained from water gas, coke-oven gas, natural gas, fuel oil, catalytic reformer gases, or the electrolysis of water orbrine. Since WWII the major source of hydrogen has been natural gas. [Lowenheim &Moran 1975]

24



#3 Ammonium Sulfate from Gypsum (Merseburg Reaction):

2NH3 + CO2 + Hp -4. (NH4)2CO3

(NH4)2C0 3 + CaSO4 • 21120 (NH4)2504 + CaCO3 + 21120

This is an important process outside the U.S. The calcium sulfate may be in the form ofgypsum, anhydrite or phosphor gypsum from wet-process phosphoric acid production.[Lowenheim & Moran 1975, Considine 1974; Slack & James 1979]

#4 Ammonium Sulfate from Synthetic Ammonia and Sulfuric Acid:

N11'3 + H2SO4 (NH4)2SO4

Ammonia is directly neutralized with sulfuric acid and water is removed by evaporation.The ammonium sulfate product is recovered by crystallization. About 75% of ammoniumsulfate production is from this method. [Lowenheim & Moran 1975]

#5 Nitric Acid from Ammonia:

4NH3 + 502 —• 4NO + 6H202NO + 02 --• 2NO2

3NO2 + Hp —. 2HNO3 + NO

The only current commercially important method for the manufacture of nitric acid is thecatalytic oxidation of ammonia with air. The manufacture of nitric acid from sodiumnitrate (Chile saltpeter) and sulfuric acid, historically the first commercial method, is nowobsolete. The high temperature oxidation of atmospheric nitrogen in an electric arc(Berkeland-Eyle Process) is also not of commercial significance. Efforts to developprocesses for making nitric acid directly from atmospheric nitrogen have not beensuccessful. [Lowenheim & Moran 1975; Slack & James 1979]

#6 Urea from Ammonia and Carbon Dioxide:

2NH3 + CO2 -1. N1-12COONH4 NI-12CONH 2 +

Anhydrous liquid ammonia and gaseous carbon dioxide are reacted at elevated temperatureand pressure to form ammonium carbamate (NH 2COONH4) which is then dehydrated toform urea. Conversion of carbamate is 50 to 75%. The remainder is decomposed toammonia and carbon dioxide. There are many variations of the process, depending largelyon the economical disposal of the waste gases. These waste gases can be partially recycled,totally recycled, or used to produce ammonium salts. The majority of current ureamanufacturers use total-recycle processes. [Lowenheim & Moran 1975, Considine 1974;Slack & James 1979]

25



#7 Ammonium Phosphates:

Anhydrous NH3 is reacted with H3PO4. Reaction ratios of NH3/H3PO4 are between 1 and2, depending on the desired grade of product.

There are 3 types of ammonium phosphates: tri-ammonium phosphate, (NH4) 3PO4?Di-ammonium phosphate, (NH4)2HPO, and mon-ammonium phosphate, NH 4H2PO4. Onlymono- and di- ammonium phosphates are used for fertilizers, alone or in combination withother salts.

The reaction for producing mon-ammonium phosphate is:

NH3 + H3PO4 —• NH4H2PO4

The reaction for producing di-ammonium phosphate is:

2NH3 + H3PO4 (NH4)2FIPO4

[Considine 1974, Slack & James 1979, Encyclopedia Britannica]

#8 Nitro-phosphates (Nitric Phosphates) and Calcium Nitrate from Phosphate Rock:

Phosphate rock is dissolved in nitric acid to produce mixtures of calcium nitrate,phosphoric acid and mono-calcium phosphate, according to the amount of acid used. Oneexample is:

3Ca3(PO4)2 • CaF2 + 18HNO3 4H3PO4 + CaH2(HPO4)2 + 9Ca(NO3)2 + 2HF

[Considine 1974; Slack & James 1979]

#9 Calcium Nitrate from Limestone:

This involves the direct reaction of limestone and nitric acid. This method is in limited useat present.

#10 Ammonium Nitrate from By-Product Calcium Nitrate:

Ca(NO3)2 + 2NH3 + CO2 + H20 -4, 21VH4NO3 + CaCO3

By-product calcium nitrate from nitro-phosphate plants is reacted with ammonia andcarbon dioxide.

26


#11 Ammonium Nitrate from Ammonia and Nitric Acid:

N1-13 + HNO3 —ftN1-1414,T03

Ammonia and nitric acid are reacted to form ammonium nitrate in either solution or in amolten form. From here it is processed to a crystal or granular form. The concentration ofthe nitric acid is typically 57% to 60% but can range from 40% to 65%. Processes differmainly in the method used to remove the solid phase from solution. Five processes in usetoday are the prilling process, the continuous vacuum crystallization process, the Stengelprocess, open pan graining, and pan granulation. [Lowenheim & Moran 1975, Considine1974; Slack & James 1979]

#12 Melamine from Urea:

Urea is heated in a fluidized bed and decomposes to ammonia and isocyanic acid whichin the presence of a catalyst form melamine.

#13 Acrylonitrile from Propylene (SOHIO Process):

C3H6 + N113 + 402 CH2 = CHCN + 3H20

Acrylonitrile is produced by the catalytic reaction of propylene, ammonia and air. Avariation of this process uses a mixture of propylene and nitric oxide in stoichiometricamounts highly diluted with nitrogen. Two former processes, from ethylene cyanohydrinand from acetylene and hydrogen cyanide, have not been used commercially for manyyears. [Lowenheim & Moran 1975]

#14 Hexamethylenediamine (HMDA, 1,6-Diaminohexane, (CH 2)6(NH2)2) from Adiponitrileby Hydrogenation:

(1VH3)

CN(CH2)CN + 4H2 --P (CH2)6(N2)2

Hexamethylenediamine is produced from adiponitrile by hydrogenation. Adiponitrile isproduced from acrylonitrile, butadiene, or adipic acid. [Lowenheim & Moran 1975, Slack& James 1979]

#15 Hexamethylenediamine (HMDA, 1,6-Diaminohexane, (CH 2)6(NH2)2) from Cyclohexaneand Nitric Acid:

Cyclohexane is oxidized with air and nitric acid in separate successive steps to yield adipicacid. Part of the adipic acid is used later for the nylon 66 condensation, the rest isammoniated to ammonium adipate which is successively dehydrated and hydrogenated toyield HMDA. [Slack & James 1979]

27



#16 Caprolactam (NH(CH2)5CO, Aminocaproic Lactam) from Cyclohexanone (BeckmannRearrangement):

(H2SO4,1VH3)C611100 + Z (NH2OH)2 - H2SO4 • C611 ioNOH + (1VH4 )2S)4 + H20

(H2SO4 , SO3) (4NH3, H20)C611 ioNOH • C61111NO .HzSO4 • C6H11N0 + 2(NH4)2504

Cyclohexanone is a key intermediate in the majority of caprolactam processes. Cyclo-hexanone is derived from phenol or cyclohexane. Caprolactam is produced via Beckmannrearrangement by the addition of hydroxylamine sulfate to cyclohexanone. Ammonia isused to neutralize the acid solution producing ammonium sulfate as a by-product. As of1975, there were several new processes that were not used in the U.S. The newer processesproduce less or no by-product ammonium sulfate. For details on these see Lowenheim &Moran (1975). Adipic Acid is also a by-product of caprolactam production. [Lowenheim& Moran 1975]

#17 Hexamethylenetetramine (Hexamine, HMTA) from Formaldehyde and Ammonia:

6CH20 + 4NH3 –• (CH))N4 + 6H20

[Lowenheim & Moran 1975; Slack & James 1979]

#18 Melamine-Formaldehyde Resins:

This process is similar to that of urea-formaldehyde resins. See process (3.8).

C3N3(NH2)3 + 3HCHO TrimethylolmelamineMelamine + Formaldehyde --• Trimethylolmelamine

[McGraw-Hill]

#19 Urea-Formaldehyde Resins:

Formaldehyde and urea are reacted to form a water soluble methylolurea. The addition ofheat or a catalyst causes the methylolurea to form a hard, insoluble, infusible resin. Theurea-formaldehyde ratio is between 1:1.5 and 1:2.5, depending on the end use.

CO(NH2)2 + HCHO NH2CONHCH2OHUrea + Formaldehyde —• Mono-methylolurea

CO(NH2)2 + 2HCHO HOCH2NHCONHCH2OHUrea + Formaldehyde —0 Di-methylolurea

[Encyclopedia Britannica, McGraw Hill]

28


#20 Polyurethane Resins:

Polyurethane resins are produced by reacting di- or polyhydroxyl compounds with di- orpoly-isocyanates, preferably aromatics. The isocyanate is generally 2,4- or 2,6-toluene di-isocyanate (TDI) or a mixture. Nitric acid is used to nitrate the isocyanate. [Slack & James1979]

#21 Polyester Resins:

Polyester resins are formed by the condensation reaction of polyfunctional acids andalcohols. Nitric acid is used in the oxidation of the cyclic poly-acids used in themanufacture of polyester resins. [Slack & James 1979]

#22 Acrylic Polymers from Oxidation by Nitric Acid:

Isobutylene is oxidized with nitric acid to hydro-oxyisobutyric acid which reacts withmethyl alcohol to yield the methacrylate. [Slack & James 1979]

#23 Acrylic Polymers from Acrylonitrile:

Acrylonitrile is dehydrated to the acrylamide using sulfuric acid. The acrylamide is reactedwith alcohol to yield the acrylate. Ammonium sulfate is produced as a by-product. [Slack& James 1979]

#24 Acrylic polymers by reaction with Hydrogen Cyanide:

There are two different processes:

1. Ethylene oxide is reacted with basic hydrogen cyanide to form cyanohydrin which isreacted with the desired the alcohol to yield the acrylate.

2. Acetone is reacted with hydrogen cyanide to form acetone cyanohydrin which isdehydrated with sulfuric acid to methyl methacrylamide. The methyl methacrylamideis reacted with the alcohol yielding the methyl methacrylate. This is the most widelyused process. [Slack & James 1979]

#25 Nylon 66:

NH2(CH2 ) 6NH2 + (CH2) 4 (COOH) 2 —1° [NH(CH2

) 6NHCO(CH-2) 2C0]n + H20

Hexa-methyenediamene + Adipic Acid -+ Poly -hexamethylene-adipamide + Water

Nylon 66 is produced by the condensation of adipic acid and HMDA in an aqueoussolution adjusted to pH 7.8 with acetic acid. [Considine 1974]

29


#26 Nylon 6 from the Polymerization of Caprolactam:

NH(CH2)5C0 —• INH(CH2)5C01"E-Caprolactam Polycaprolactam

[Considine 1974]

#27Phenolic Resins from Hexa-methylene-tetramine:

Hexamethylenetetramine is used in the second stage polymerization of the phenolicmonomer. [Slack & James 1979]

#28 Cellulose Nitrate:

C611/605 + xHNO3 —• C6H(o-x)0(s.x)(0NO2)x + xH20

[Slack & James 1979]

#29 Trinitrotoluene (TNT):

CH3C6FI3(NC0)2 + HNO3 + H2SO4 -, CH3C112(NO2)3Toluene + Nitric Acid + Sulfuric Acid ---• TNT


#30 Glycerol Tri-nitrate (Nitroglycerine, NG), CH2NO3CHNO3CH2NO3:

Produced from glycerol and a mixture of nitric and sulfuric acid. [Slack & James 1979]

#31 PETN, Penta-erythritoletranitrate:

Produced by batch nitration of pentaerythritol with nitric acid. [Slack & James 1979]

#32 Hydrazine (N2H4) from Ammonia and Sodium Hypochlorite (Raschig process):

NH3 + Na0C1—• NH2CL + NaOHNH2C1 + NH3 —• N2H4 + Hcl

[Lowenheim & Moran 1975; Slack & James 1979]

30


#33 Hydrazine from Urea:

(HN2)2C0 + NaOCI + 2NaOH -1. N2H4 + NaCI + Na2CO3 + H20

[Lowenheim & Moran 1975]

#34 RDX (cyclonite, Trimethylenetrinitramine):

Hexamine is reacted with 50% nitric acid yielding a di-nitrate. Ammonium nitrate,dissolved in 100% nitric acid is added. The crude RDX is precipitated, separated, andrecrystallized from acetone. The product contains 10% tetra-methylenetetramine (HMX),also an explosive. [Slack & James 1979]

#35 Ethylene-amines:

Ethylene-amines are produced by the reaction of ethylene di-chloride with ammonia vaporor aqueous ammonia under varying temperatures and pressures. [Slack & James 1979]

#36 Fatty Nitriles and Derivatives:

These are produced by the reaction of fatty acids (tallow, coco, cottonseed, soya, tall oil,etc.) with ammonia at elevated temperatures and pressures in the presence of a metallicoxide catalyst. Their basic formula is RCN, where R is an odd-numbered, straight-chainalkyl group in the C7 to C21 range. [Slack & James 1979]

#37 Ethanol-amines:

These are produced by the pressurized reaction of ethylene oxide and aqueous ammoniaat moderate temperatures. Their chemical formulas are:

Mono- NH2CH2CH 20HDi- NH(CH2CH2OH)2Tri- N(CH2CH2OH)3


#38 Methyl-amines:

These are produced by the reaction of NH 3 and methanol in the presence of metallic oxidecatalysts (AL203, Si02) or phosphoric acid or phosphate salt dehydration catalysts. [Slack& James 1979]

#39 Aniline from Chloro-benzene by Ammonolysis:

31


C6H5C1 + 2NH3 (aqua) —. C6H5tVH 2 + NI-14C1 (solution)


#40 Aniline from Nitrobenzene by Catalytic Vapor-Phase Hydrogenation:

C6H5NO2 + 3H2 C6Ff5NH2 + 2H2O


#41 Aniline from Nitrobenzene by Reduction:

(HC1)4C6H5N)2 + 9Fe + 4H20 4 C6H5NH2 +3 Fe30 4


#42 Hydrogen Cyanide from the Andrussow and DEGUSSA Processes:

Andrussow Process: 2NH3 + 302 + 2CH4 2HCN + 6H20

This reaction takes place in the presence of a platinum-rhodium catalyst.

DEGUSSA Process: NH3 + CH4 HCN + 3H2

This reaction takes place in the presence of a platinum catalyst. [Lowenheim & Moran1975]

#43 Biuret:

Formed by heating aqueous urea solutions at atmospheric pressure. [Considine 1974]

#44 Ammonium Chloride (Sal-Ammoniac) as a By-Product from the Manufacture ofSodium Sulfite:

2NH3 H20 + SO2 —0 (NH4)2S03(NH4)2S03 + 2NaCI --• Na2S03 + 2NH4C1

This is a minor source of ammonium chloride. Ammonium chloride can also be manufac-tured from ammonia and hydrochloric acid, but the process is not usually economic.[Lowenheim & Moran 1975]

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#45 Ammonium Chloride (Sal-Ammoniac) as a By-Product from the Ammonia-SodaProcess:

CaC12 + 2NH3 + CO + Hp —D. 2NH4C1 + CaCO3

This is the major source of ammonium chloride. [Lowenheim & Moran 1975; Slack &James 1979]

#46 Ammonium Chloride (Sal-Ammoniac) by reaction of Ammonium Sulfate and SodiumChloride Solutions:

(NH4)2SO4 + 2NaC1 —,. Na2SO4 + 2NH4C1

This is a major source of ammonium chloride. [Lowenheim & Moran 1975; Slack & James1979]

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industrial metabolism of nitrogen - inseadflora.insead.edu/fichiersti_wp/inseadwp1992/92-57.pdf ·...

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