Nitrogen and ammonia

Where Found

Nitrogen gas (N₂) constitutes 78 percent of the Earth’s atmosphere. There are deposits of potassium nitrate in India and of sodium nitrate in Chile, but nitrogen ranks only thirty-third among the elements in crustal abundance, with an average concentration of 0.03 percent by weight. Natural gas, petroleum, and coal contain nitrogen, and all plants and animals contain nitrogen in the form of proteins. The human body is about 3 percent nitrogen by weight.

Ammonia occurs as ammonium chloride salt in volcanic ejecta, but industrially produced ammonia is the predominant form used. Ammonia from sewage, agricultural runoff, and industrial activities can be a water pollutant.

Primary Uses

The largest use of nitrogen compounds—urea, ammonium nitrate, ammonium phosphates, nitric acid, and ammonium sulfate—is in fertilizer for crops such as wheat, corn, and soybeans. Nitrogen gas is used as a protective gas in the food, electronics, and metals industries.

Although fertilizer uses are predominant, ammonia is also used as a refrigerant and as a chemical intermediate in the manufacture of nitric acid, nitrogen-containing plastics and fibers (polyamides, polyacrylonitrile, and polyurethane), and explosives.

Technical Definition

Nitrogen (symbol N), atomic number 7, belongs to Group 15 of the periodic table of the elements. The has two naturally occurring isotopes and an atomic weight of 14.007. Nitrogen occurs as a colorless, odorless gas weighing 1.25 grams per liter (0.0° Celsius and 1 pressure). Liquid nitrogen is a colorless liquid boiling at -195.8° Celsius and freezing to a colorless solid at -210° Celsius. Gaseous, liquid, and solid nitrogen all consist of diatomic (N₂) “dinitrogen” molecules.

Ammonia (NH₃) is a pungent, toxic gas, weighing 0.76 gram per liter (0.0° Celsius and 1 atmosphere pressure). Liquid ammonia boils at -33.4° Celsius and freezes at -78° Celsius. Ammonia is soluble in water to the extent of 28 percent by weight, and it forms explosive mixtures with air.

Description, Distribution, and Forms

Nitrogen resources in the atmosphere amount to 4 × 10²¹ grams of N₂, a virtually inexhaustible supply. Nitrogen generally ranks second among all chemicals produced in the United States. Ammonia ranks sixth. Nitrogen from the atmosphere must undergo “fixation” (conversion to ammonia or oxy compounds) before it is available for plant nutrition. Fixation of nitrogen occurs in the atmosphere during fires or thunderstorms, when temperatures rise enough to make nitrogen and oxygen react. In the soil, occurs with the mediation of an enzyme, nitrogenase, present in Rhizobiumbacteria that live in the root nodules of peas, clover, and alfalfa. Industrial production of ammonia by the high-pressure, catalyzed reaction of hydrogen and nitrogen (the Haber-Bosch process) probably accounts for less than half of all nitrogen fixation. Fixed nitrogen ultimately returns to the atmosphere through decay of plants and animals and the action of denitrifying bacteria. This process is called the nitrogen cycle.

The world’s growing population of both humans and cattle creates increasing demand for the fixed nitrogen that goes into fertilizer and animal feed. Ammonia plants have been built worldwide to meet this need. Worldwide, approximately 150 million metric tons of ammonia were produced in 2023.

Nitrogen oxides (“NOx” compounds) produced when fuels burn in air are toxic and contribute to acid rain, since they react with water to produce nitric acid. Nitrates and nitrites from agricultural runoff fertilize the growth of algae in lakes and streams (causing eutrophication); the increase in nitrogen runoff from fertilizers contributes, for example, to the summer hypoxic (oxygen-deficient) zone in the Gulf of Mexico. Drinking water containing nitrate interferes with red blood cells, causing methemoglobinemia in infants. Federal drinking water standards require that there be less than 10 milligrams of nitrate per liter of water.

Nitrogen oxides in the atmosphere contribute to photochemical smog and catalyze the destruction of stratospheric ozone. Photochemical smog is associated with automobile exhaust: the unburned and nitrogen oxides are acted on by sunlight. A particularly irritating substance called peroxyacetyl nitrate (PAN) can form. The federal clean air amendments contain a limit on the allowable concentration of nitrogen oxides of only 0.05 part per million, annual arithmetic mean.

Nitrogen oxides that reach the stratosphere can cause the destruction of ozone by a catalytic process. Supersonic aircraft flying in the stratosphere would emit nitrogen oxides and reduce the concentration of ozone. This in turn might result in a harmful increase in the of ultraviolet radiation at the surface of the Earth, since the that absorbs ultraviolet light exerts a protective effect. Atmospheric chemistry is very complex, and there are many other gases that interact with ozone.

Another nitrogen oxide, nitrous oxide, or dinitrogen (N₂O), is emitted by certain industrial processes. Although it is less toxic than other oxides of nitrogen and is odorless, it does absorb infrared radiation and may contribute to (the “greenhouse effect”). Nitrous oxide also enters the atmosphere as the result of microbial action in the soil.

Ammonia is toxic, with a maximum allowable concentration in the workplace of 50 parts per million in air. Air pollution by ammonia is rare except in cases of accidental release. Because ammonia is extensively transported by truck and pipeline, there are occasional releases, necessitating evacuation of the surrounding area. Since it is less dense than air and is soluble in water, ammonia tends to dissipate rapidly after a spill.

Humans obtain most of their nitrogen from the proteins in meat, milk, or legumes. The recommended daily allowance of protein for an adult male is 50 to 70 grams, and protein deficiency results in a debilitating condition called kwashiorkor, suffered mainly by children in underdeveloped countries in Africa.

History

Nitrogen was isolated about 1770 by Daniel Rutherford, Carl Wilhelm Scheele, and Henry Cavendish. Ammonia was isolated by Joseph Priestley in 1774. He prepared the gas by heating ammonium chloride with lime, and he used a pneumatic trough filled with mercury to collect the gas. In 1862, Justus von Liebig suggested that nitrogen is essential for plant nutrition and theorized that plants obtain it from the atmosphere; however, the details of microbial nitrogen fixation were not clear until much later.

Obtaining Nitrogen and Ammonia

The commercial production of nitrogen became possible after the development of the Lindé process for liquefaction of air in 1895. Nitrogen is separated from the other elements in liquid air by fractional distillation, selective adsorption on zeolites, or membrane technology.

Large-scale production of ammonia from hydrogen and nitrogen (the Haber-Bosch process) began in Germany in 1913. Fritz Haber had developed the ammonia synthesis on a small scale, and he demonstrated it to management at I. G. Farben, the giant German chemical company. Carl Bosch led the team that designed the first ammonia plant, inventing the technology of high-pressure hydrogen reactions in the process. The Nobel Prize in Chemistry was awarded to both Haber (1918) and Bosch (1931) for their achievements.

Nitrogen and hydrogen react at 400° to 550° Celsius and 100 to 1,000 atmospheres of pressure in the presence of an iron catalyst. The hydrogen is obtained by reacting steam with over a nickel catalyst (“steam reforming”), and energy requirements are about 25 million British thermal units (Btus) per ton of ammonia. The gaseous reactants used must be purified to free them of substances that might interfere with the action of the catalyst.

The cost of feedstock accounts for more than half the cost of producing ammonia. Ammonia prices are sensitive to the cost of energy and of natural gas. There is a futures market in liquid ammonia that helps users hedge against possible price increases. Minor amounts of ammonia are recovered from coke-oven gases, usually directly converted to ammonium sulfate by reaction with sulfuric acid.

Uses of Nitrogen

As previously mentioned, nitrogen gas, because of its low chemical reactivity, is used to protect foods, pharmaceuticals, electronic parts, and hot metal surfaces from damage by oxygen or other reactive gases. Liquid nitrogen is used to freeze biological samples (such as blood and semen) and foods, and to solidify rubbery materials that need to be pulverized or ground up. Nitrogen gas also inflates the airbags in automobiles, being evolved from sodium azide in a rapid exothermic reaction triggered by a collision.

Although most ammonia is used in fertilizers, some is used as a chemical intermediate to make other nitrogen compounds. Mixtures of ammonia and air, passed over a platinum/rhodium catalyst at around 500° Celsius, produce oxides of nitrogen that combine with water to produce nitric acid. Mixtures of ammonia and methane can be catalytically oxidized to make hydrogen cyanide, and with propene in place of methane, acrylonitrile can be produced. Explosives such as gunpowder, nitroglycerin, and dynamite (trinitrotoluene, TNT) require nitric acid for their manufacture. Hydrogen cyanide is used in making sodium cyanide for the mining industry and methyl methacrylate, the precursor of Plexiglas, while acrylonitrile is used to make polyacrylonitrile synthetic fibers for clothing and carpets.

A host of other synthetic compounds, including many dyes and pharmaceuticals, derive their nitrogen content from nitric acid or ammonia. Two examples are synthetic indigo, used to dye blue jeans, and acetaminophen, an over-the-counter headache remedy.

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