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Ammonia, Ammonium Ion NH3, NH4+

Ammonia (NH3) is very soluble in water. The Henry's law constant for NH3 is 57.5 mol/(L-atm) at 25 C. NH3 can combine with a proton (H+) to form an ammonium ion (NH4+).

The concentrations of NH3 and NH4+ at equilibrium are related by the following equation,

where Ka is a constant (5.5x10-10 at 25ºC) and square brackets indicate concentrations (mol/L).The value of Ka at other temperatures is given by the following equation,

where T is temperature in K (Wright et al. 1961). Figure 1 shows how NH3 and NH4+ concentrations vary with pH. In the pH range of most natural waters, NH4+ is the predominant form of inorganic N in the -3 oxidation state. NH3 is the species that is toxic to fish, not NH4+.


Figure 1. Percentage of NH4+ and NH3 as functions of pH.


NH4+ is a major component of atmospheric aerosols. (see atmosphere)

NH4+ participates in cation exchange reactions in which dissolved ions replace ions bound by negatively charged sites on mineral particles, such as clays and natural organic matter.

Ion exchange may be an important reaction for NH4+ in soils, sediments, and aquifer systems. For example, the retardation coefficient of NH4+ (rNH4) in sandstone was estimated from ion exchange measurements to be 16<rNH4<80 (Drever 1982). A pulse of NH4+ would move at a rate between 1/80th and 1/16th that of water in the sandstone aquifer.

Calculations of ion exchange equilibria can sometimes be used to qualitatively describe the ion exchange behavior of cations. Two useful empirical ion exchange parameters are the cation exchange capacity (CEC) and the selectivity coefficient. The CEC of a material (clay, soil, ...) is measured by saturating the material with one cation (e.g. NH4+), displacing the first cation with another cation (e.g., Mg2+), and measuring the concentration of displaced NH4+. Table 5 presents the CEC values of some clay minerals.


Table 5. Cation exchange capacities of
some clay minerals (meq/100g) (Drever 1982)

Smectites 80-150
Vermiculites 120-200
Illites 10-40
Kaolinite 1-10
Chlorite <10



The selectivity coefficient can sometimes be used in the following equation to describe ion-exchange equilibria involving NH4+ and Na+.

XNH4 and XNa indicate the mole fractions of NH4+ and Na+ in the exchanger phase, KNH4Na is the selectivity coefficient, square brackets indicate aqueous concentrations, and n is an empirical coefficient with value close to 1 for singly charged ions and 0.7<n<0.9 for doubly charged ions. Similar equations apply for exchange of other ions having the same charge, including Na+/K+, NH4+/K+, and Ca2+/Mg2+. For ions of different charge the equation is more complicated. For example,

The following equation describes NH4+ sorption in some aquifer systems (Drever and McKee 1980).

{NH4+}clay is the concentration of sorbed NH4+ (meq/kg), M is the total cation concentration (meq/kg), and the other symbols are as defined above.

Both the CEC and selectivity coefficient depend on the pH and solution composition. Thus, even for the same sample of the same material, neither parameter is a constant. Nevertheless, the ion exchange parameters allow the rationalization of cation behavior in many systems.

Nitric Acid/Nitrate Ion HNO3, NO3-

Nitric acid and nitrate salts are all very soluble in water. HNO3 is a strong acid and completely dissociates in water. HNO3 is volatile. NO3- absorbs solar UV radiation. It is a major source of hydroxyl radical in surface water and atmospheric water droplets.

Nitrous Acid/Nitrite Ion HNO2, NO2-

Nitrite (NO2-) is an intermediate product in many N transformations. It is produced by the oxidation of NH4+ in the first stage of nitrification (Nitrosomonas) and by reduction of NO3- in the first stage of denitrification. It is oxidized to nitrate by bacteria of the genus Nitrobacter in the second stage of nitrification and reduced to N2O, NO, or N2 in the second stage of denitrification. HNO2 is a weak monoprotic acid with pKa 5.2. The following figure shows how the relative concentrations of HNO2 and NO2- vary with pH. At the pH values of most natural water systems, the predominant form is NO2-.

Figure 1. Percentage of HNO2 and NO2- as functions of pH.

Nitric Oxide NO

NO is produced by combustion and by denitrification. Combustion produces both NO and NO2, collectively denoted NOx, but most of the N oxides produced by combustion is NO. Combustion produces NO both by combination of N2 and O2 and by oxidation of any organic N in the fuel. In the United States, transportation accounts for approximately 40% of anthropogenic NOx emissions. Combustion from stationary sources produces approximately 56%. (Wark and Warner 1981)

NO is slightly soluble in water. The Henry's law constant for NO is 1.9x10-3 mol/atm. The typical atmospheric partial pressure of NO is 2x10-10 atm. NO reacts rapidly with O2 to form NO2. NO plays an important part in some biological processes. NO biochemistry is the subject of much current research.

Nitrogen Dioxide NO2

NO2 is produced by combustion, degradation of organic matter, and oxidation of NO. It is one of the few colored gases and gives a brownish tint to polluted air. It is visible at concentrations as low as 1 ppmv. NO2 is oxidized to HNO3 in the atmosphere. It also hydrolyzes and disproportionates to give HNO3 and HNO2.

NO2 is slightly soluble in water. The Henry's law constant for NO2 is 1.0x10-2 mol/atm. The typical partial pressure of NO2 in the atmosphere is 2x10-9 atm.

Nitrous Oxide N2O

N2O is the second most abundant N species in the atmosphere. Its partial pressure is

3x10-7 atm. It is relatively unreactive in the troposphere. The main sink for N2O is photochemical reactions in the stratosphere (see atmosphere). It is produced by denitrification. N2O is fairly soluble in water. Its Henry's law coefficient is 2.6x10-2 mol/atm.

Organic N

Organic N includes all substances in which N is bonded to C. It occurs in both soluble and particulate forms. The largest fraction is made up of amino acids and peptides and is often called amino N. Particulate organic N includes small organisms (algae, bacteria, ...), both living and dead, and fragments of organisms. Soluble organic N is from wastes excreted by organisms or from the degradation of particulate organic N.

Organic-N concentrations in natural waters, soils, and sediments are operationally defined. In the Kjeldahl method, a sample of water, soil, or sediment is heated with H2SO4 and a catalyst and N from amino acids is converted to NH4+. Total Kjeldahl N (TKN) includes N from amino acids and any NH4+. Organic-N is calculated by subtracting NH4+ (determined separately) from TKN. In another method, a water sample is oxidized using various combinations of potassium persulfate, heat, and ultraviolet light and all N is converted to NO3-. Organic N is calculated by subtracting NO3- (determined separately) from the NO3- in the oxidized sample. Concentrations of individual amino acids can be determined chromatographically.

Typically, most N in soils and surficial sediments occurs in organic form. The amount of organic N is soils and sediments is influenced by climate - all else being equal, increasing with moisture and decreasing with temperature in the United States. It is also influenced by vegetation. In Illinois equivalent soils developed under prairie had twice the organic N of soils developed under forest. It is also influenced by topography. Soils such as the more upland prairie of the Morrow plots are estimated to originally have had 0.3 percent by weight N whereas soils developed under prairie wetlands originally had 2.2 percent N. The amount of N is also influenced by the particle size of the soil and sediment (more accumulating in fine-grained material) and the amount of mineral nutrients (especially phosphorus) in the soil and sediment. It is also influenced by the age of the land surface - older surfaces generally being lower in C, N, and mineral nutrients. Agricultural practices have reduced the N content in the plow layer of cornbelt soils by about 40 percent, on average, from their estimated virgin condition.

Organic N sometimes makes up a significant fraction of soluble and particulate N in natural waters. See Examples of Organic N Measurements.

Urea CO(NH2)2

Urea is an organic N compound that is manufactured in large quantities. It is used as a fertilizer, including controlled-release fertilizers, and in many industrial processes. The U.S. production of urea was 9,330,000 tons in 1999 (Anon. 2000). Urea is the soluble form of N that is excreted by mammals. Therefore, large amounts of urea are added to the surface soils of feedlots and dairies. However, urea is rapidly hydrolyzed to NH3 and CO2, so it is not expected to make up a significant part of organic N in water bodies, soils, or sediments.

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