Nitrogen exists in nature in several oxidation states: N(-III) as in NH3, NH+, and various organic compounds; N(III) in nitrites, NO- , and N(V) in nitrates, NO- as well as N(0) in N2 in addition to other formal oxidation states in oxides of which nitrous oxide (N2O), nitric oxide (NO), and nitrogen dioxide (NO2), are most important environmentally.

The elemental form N2 contains a triple bond with a large bond energy (946kJ/mol). Consequently, reactions that require the N—N bond to be broken are likely to take place with difficulty, even if the overall energy change of the reaction is favorable. As a result, N2 is relatively inert. Some of its most important environmental reactions are produced by microorganisms, which can provide a reaction mechanism of low activation energy to convert N2 to ammonia and amines.

The aqueous redox chemistry of nitrogen involves primarily NO-, NO-, and NH+; these take part in oxidation-reduction processes, which are expressed in the following equations:

These reactions depend on pH, and the equilibrium composition of a nitrogenous system depends on this as well as on the redox potential pE of the system (Section 9.4). Nitrite has a comparatively narrow range of pE over which it can exist at significant concentrations (Figure 10-5); most commonly the stable forms of nitrogen are NH+ or ammonia in reducing environments, and NO-in oxidizing ones. At pH ranges near neutrality, the NO2- stability region lies near pE 6.5.

The essential features of the nitrogen cycle are shown in Figure 10-6. There are several processes of importance.

1. Nitrogen fixation refers to the conversion of atmospheric N2 to another chemical form, most frequently N(-III). In nature, formation of N(-III), amine nitrogen, is a microbial process, most importantly involving bacteria that have a symbiotic relationship with the roots of certain plants. Legumes such as clover, peas, and alfalfa that are associated with Rhizobium bacteria are best

FIGURE 10-5 The relative concentrations of the nitrogen species in solution as a function of pE at pH 7.

known in this respect, but some aquatic bacteria and blue-green algae are important in marine systems, and some trees such as alder (with bacteria of the Frankia genus) contribute in forest regions. This complex biochemical process depends on a large metal-containing enzyme, nitrogenase, in which

FIGURE 10-6 The nitrogen cycle. Terms are defined later in the text.

both iron and molybdenum take part. The process leads to amine nitrogen, which is incorporated into amino acids in the plant. Plants generally, however, absorb nitrogen from the soil in the form of the nitrate ion. Other natural fixation processes—for example, formation of nitrogen oxides in the atmosphere by lightning discharges—also contribute to a lesser degree. Much nitrogen is fixed industrially, as discussed later.

2. Nitrification is the conversion of amine nitrogen to nitrate. Decay of protein material produces NH3, which is oxidized through nitrite to nitrate. This is also a bacterial process. Nitrification is reversible, under anaerobic conditions, since reduction of nitrate can provide an energy source for bacteria through the net change

(CH2O represents the approximate composition of the organic material that is used up in this process.)

An intermediate in the reduction of nitrate is nitrite. The nitrite ion NO2 is relatively toxic because of its interaction with hemoglobin. Nitrite in the blood results in the oxidation of the Fe(II) of hemoglobin to Fe(III), forming methemoglobin, which has no oxygen-carrying ability; this disease is called methemoglobinemia. While cases of direct nitrite poisoning are rare, the nitrate ion can be reduced to nitrite in the stomach, and for this reason food and water with high nitrate contents are dangerous. This reduction is especially possible in the stomachs of infants, where the low acidity allows growth of nitrate-reducing microorganisms. Some evidence exists that nitrites in the body can react with organic amines to form possibly carcinogenic nitrosa-mines,

Sodium nitrate and nitrite (the latter often produced from nitrate in situ) are common additives in cured meats, where they act to prevent the growth of bacteria, but nitrite is also added to produce flavor and an attractive color through the formation of nitrosylmyoglobin. Concerns about the risks from nitrate and nitrite have led to reductions in their use.

3. Denitrification is the formation of gaseous N2 from nitrate to return nitrogen to the atmosphere. Primarily it is carried out by soil bacteria and involves reduction by organic carbon compounds with the production of CO2. This process can be used in an anaerobic process to remove nitrates from wastewater by addition of a reducing agent such as methanol as food for the bacteria:

TABLE 10-3

Global Nitrogen Reservoirs

TABLE 10-3

Global Nitrogen Reservoirs

Nitrogen in

Amount (g x 1015)



< 109 (99.99% N2; 99% of the rest is N20)



< 104 (land); 8 x 102 (ocean)



< 107 (dissolved N2); 6 x 105 (dissolved inorganic); 2-5 x105 (organic)


6 x

< 104 (organic); 1 x 104 (inorganic)

Rocks; sediments

6 x

< 108

Source: Data from P. M.Vitousek, J. D. Aber, R. W. Howarth, G. E. Likens, P. A. Matson, D. W. Schindler, W. H. Schlesinger, and D. G. Tilman, Ecol. Appl, 7, 737 (1997).

Source: Data from P. M.Vitousek, J. D. Aber, R. W. Howarth, G. E. Likens, P. A. Matson, D. W. Schindler, W. H. Schlesinger, and D. G. Tilman, Ecol. Appl, 7, 737 (1997).

Some nitrogen is returned to the atmosphere as N20 generated bacterially in soils. Nitrous oxide is comparatively inert, but it is rapidly destroyed by processes that are not entirely clear. It can be photolyzed to N2 at high altitudes, where radiation exists at wavelengths short enough to be absorbed by N20 (see Chapter 5), but it does not seem likely that this reaction can be responsible for the low N20 levels usual at low altitudes. Other processes for N20 removal must exist at or near the earth's surface. It is a greenhouse gas, and its concentration in the atmosphere has been increasing recently by a few tenths of a percent per year from the effects of fertilizers, land clearing, and industrial processes, although the details are not well understood.

Some ammonia is also released naturally from decay of organic materials and is present in air either as NH3 gas or as an ammonium salt aerosol. These are removed from the atmosphere in rain. Ammonia is the main basic material in the atmosphere.

Table 10-3 gives estimates of the amount of nitrogen contained in the various environmental reservoirs. The atmosphere contains most of the earth's nitrogen as N2. Inputs of "fixed" nitrogen, oxidized or reduced forms, are given in Table 10-4. There are widely varying estimates for some of these quantities; for example, estimates of nitrogen oxides produced by lightning discharges vary from 2 to 20 x 1015 g of N per year, and fossil fuel combustion from 14 to 28 x 1015 g of N per year.9 As can be seen from these data, anthropogenic inputs exceed natural ones. The cycle is unbalanced as the amount of fixed nitrogen is continually increasing.

9See, for example, G. Brasseur, J. J. 0rlando, and G. S. Tyndall, eds., Atmospheric Chemistry and Global Change, 0xford University Press, 0xford, U.K., 1999, Chapter 5, for some of these estimates.

TABLE 10-4

Input of Fixed Nitrogen into the Environment

Nitrogen input Amount (g x 1012 / yr)


Nitrogen input Amount (g x 1012 / yr)











Crops (legumes, etc.)



Fossil fuel combustion


Biomass combustion


Land use changes


Total anthropogenic


Source: Data From D. A. Jaffe, in Global Biogeochemical Cycles, S. S. Butcher, R. J. Charlson, G. H. Orians, and G. V. Wolfe, eds., Academic Press, San Diego, CA, 1992.

Source: Data From D. A. Jaffe, in Global Biogeochemical Cycles, S. S. Butcher, R. J. Charlson, G. H. Orians, and G. V. Wolfe, eds., Academic Press, San Diego, CA, 1992.

Fixed nitrogen includes gaseous compounds such as NO*, N2O and NH3 as well as nitrates and so on. NO* released to the atmosphere annually amounts to about 64 x 1012g of N, about half from natural sources; NH3 is about 53 x 1012g of N, over half of which has an anthropogenic origin, and N2O, about 15 x 1012 g of N. There is some interconversion of these species as one is oxidized or reduced to another.

Fertilizers make up the largest source of anthropogenic input of nitrogen to the environment (see also Section 10.5). Because nitrogen is an essential nutrient for plants, and often a limiting one, large-scale use of nitrogen-containing fertilizers has become commonplace. These may take the form of nitrates, which are immediately available to the plant, or ammonia or organic nitrogen-based fertilizers. These must be converted to inorganic nitrate before use by plants, and since this conversion takes place over a period of time, they can provide a longer lasting source of nitrogen than nitrate fertilizers. Nitrates tend to be quite soluble and weakly held by ion exchange forces, and so may easily be leached from the soil and wasted. Organic nitrogen fertilizers may or may not be soluble, but soluble materials can be formulated to resist dissolution (e.g., pellets may be coated with sulfur and wax). However, it should be emphasized that movement of materials through the soil is highly complicated by absorption and ion exchange processes; even readily soluble materials can be retained for long periods in some types of soil (Section 12.3). Thus ammonia itself, although a very water-soluble gas, is an effective fertilizer, as are its aqueous solutions. It is held in the soil as the ammonium cation by ion exchange. Other widely used compounds are NH4NO3, (NH4)2SO4, ammonium phosphates, and urea. Note that since nitrogen is generally assimi-

lated by plants as the nitrate ion, the original source of the nitrogen does not affect the plant. Rather, choice of the form a fertilizer should take is based on the rate of release of NO", extent of leaching (which results in waste and ultimately water pollution), and other economic factors. Estimates are that half of the nitrogen applied as fertilizer is never taken up by plants but enters into runoff water.

Industrial nitrogen fixation to produce fertilizers (and other nitrogen compounds) is based primarily on the Haber process:


The source of hydrogen is normally natural gas or petroleum, either of which react with steam under the action of a catalyst in a process such as

Consequently, fertilizer production and cost are closely linked to the supply of these fossil fuels. Ammonia is the starting material for most other industrial nitrogen compounds; for example, nitrates are produced through reactions that include oxidation of ammonia by air in the presence of a platinum catalyst.

Oxides of nitrogen in the atmosphere are a cause of concern with respect to air pollution problems, as discussed in Chapter 5. The main compounds of concern are NO and NO2, which are by-products of combustion processes. They are formed from the reaction of N2 in the air with O2. The primary product is NO, but some NO2 is produced as well. Production of nitrogen oxides (often called NO*) is directly related to the temperature in the combustion zone of a furnace or engine and can be reduced by operating at lower temperatures. This generally reduces the efficiency of the device, however. Reduction of the amount of excess air in the combustion chamber also will reduce NO* emission, but at the expense of an increase in the amount of incompletely burned fuel. A two-stage combustion, the first at high temperature with an air deficiency, followed by completion of the combustion at a lower temperature, can be quite effective in reducing NO* emission. Figure 10-7 shows NO* equilibrium concentrations as a function of temperature at one particular fuel/air ratio, equivalent to a 20% excess of oxygen. The reaction between oxygen and nitrogen to form the oxides is kinetically complex, requiring first the reaction of oxygen with another molecule to generate oxygen atoms, and reaches equilibrium in a finite time only at high temperatures. Once formed, the NO and NO2 concentrations will by "frozen in" when the temperature drops to a value at which the rate becomes insignificant.

FIGURE 10-7 Equilibrium pressures of NO and NO2 in air with a 20% excess of oxygen over fuel. Redrawn from S. S. Butcher and R. J. Charleson, An Introduction to Air Chemistry, Copyright © 1972. Used by permission of Academic Press.

FIGURE 10-7 Equilibrium pressures of NO and NO2 in air with a 20% excess of oxygen over fuel. Redrawn from S. S. Butcher and R. J. Charleson, An Introduction to Air Chemistry, Copyright © 1972. Used by permission of Academic Press.

If production of NO* is not reduced, absorption of NO* from the combustion gases is an alternative process to reduce pollution levels. Such a practice is used for SO2 removal, as will be discussed in Section 10.4, but is more difficult for NO*. Various methods have been proposed for both acid and alkaline scrubbing of stack gases. Alkaline scrubbers can remove SO2 and NO* simultaneously, but neither NO nor NO2 is efficiently absorbed by a simple basic solution. However, a mixture of NO and NO2 can be absorbed through an equilibrium reaction,

Although this equilibrium lies far to the left under normal conditions, it is shifted to the right if the N2O3 is absorbed into a basic solution; this occurs readily, since N2O3 is the anhydride of nitrous acid,

The nitrite salts formed in a basic solution can be oxidized to nitrate. This process, however, requires equimolar amounts of both NO and NO2 in the gas stream, or recycling of the excess.

Catalytic reduction of NO* is possible; reaction with CO or CH4, for example, can produce N2:

This reaction is favored thermodynamically but is very slow under normal circumstances. Various metals (e.g., Ag, Cu, Ni, Pd) are effective catalysts but are subject to poisoning in practical use and are sometimes expensive. Some oxides (e.g., copper chromite [Cu2Cr2O4] and Fe2O3) are also effective catalysts for this process. Commercial systems for catalytic NO* removal from flue gases use ammonia as the reducing agent in the following reactions:

Only a few such systems are in use , but up to 80% removal of NO* is claimed for a 1000-MW coal-fired plant in Japan.

The nitrogen cycle, like the carbon cycle, is not balanced because of human activities (Table 10-4). Although less well known than the carbon dioxide problem, the increase in fixed nitrogen compounds can have comparably significant environmental effects. Industrial nitrogen fixation plays a major role in this imbalance, but also involved is release of fixed nitrogen from soils and biomass from land clearing and wetland draining, and burning of biomass and fossil fuels. "Natural'' fixation has increased through planting of more leguminous and other plant crops. Increased fixed nitrogen in soils increases plant growth, which in fact is helpful regarding the carbon dioxide problem, as more carbon is sequestered in biomass. However, this is coupled with more release of nitrous oxide to the atmosphere in denitrification processes as a consequence of more nitrogen compounds available to soil bacteria. Atmospheric nitrogen compounds contribute to acid rain and to soil acidity; this, along with the acidity increase that results from the bacterial activity, encourages the solubilization of other elements.

Excessive nitrogen fertilization has several direct consequences. One is a change in the number of plant species, which occurs because those that thrive on the nitrogen drive out those that prefer low nitrate conditions. Major ecosystem changes can result. A second is disruption of the soil chemistry. As mentioned earlier, soil acidity changes can change the availability of other nutrients, and with excessive nitrogen, one or another of these can be growth limiting. Conversion of these other nutrients to forms that can be leached away can ultimately decrease soil fertility. Another serious concern is the runoff of the excess nitrogen from soils. Nitrate salts are soluble, and not strongly held by ion exchange processes. When they enter freshwater systems, they can increase algal growth and eutrophication, although in most freshwaters, phosphate is limiting. The major effect is felt in estuaries and coastal areas of the

4N0 + 4NH3 + 02 ^ 4N2 + 6H20 6N02 + 8NH3 ^ 7N2 + 12H20

ocean, where nitrogen is normally limiting and additional input can lead to harmful algal blooms and damage to important marine life. Consequences of the nitrogen overload are most obvious in northern Europe and to a lesser extent in the United States, where fertilizer use has been highest and nitrogen inputs to forests and freshwater systems have increased to 10-20 times the natural level.

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