NITROGEN (REVISED)

Note: This article, originally published in 1998, was updated in 2006 for the eBook edition.

Overview

Nitrogen is the first member in Group 15 (VA) of the periodic table. The periodic table is a chart that shows how chemical elements are related to one another. Nitrogen is in a family group named after itself. Other elements in the nitrogen family are phosphorus, arsenic, antimony, and bismuth.

Nitrogen is one of the most interesting of all chemical elements. It is not a very active element. It combines with relatively few other elements at room temperature. Yet, the compounds of nitrogen are enormously important both in living organisms and in industrial applications. Five of the top fifteen chemicals that are produced synthetically by chemical producers are compounds of nitrogen or the element itself. How does such an inactive element end up with so many important compounds?

Nitrogen makes up more than three-quarters of the Earth's atmosphere. It is also found in a number of rocks and minerals in the Earth's surface. It ranks about number 32 among the elements in terms of abundance in the Earth's crust.

SYMBOL
N

ATOMIC NUMBER
7

ATOMIC MASS
14.0067

FAMILY
Group 15 (VA)
Nitrogen

PRONUNCIATION
NYE-tru-jun

Nitrogen was discovered by a number of chemists at about the same time, approximately 1772. But it was not until the early part of the twentieth century, when chemists learned how to make compounds of nitrogen, that the most important uses of the element became known.

By far the most notable use of nitrogen is in the production of ammonia (NH₃). Ammonia is used to make other compounds, such as ammonium sulfate ((NH₄)₂SO₄), ammonium nitrate (NH₄NO₃), urea (CO(NH₂)₂), and nitric acid (HNO₃). These compounds are primarily used to make synthetic fertilizer. Both elemental nitrogen and nitrogen compounds have a number of important industrial uses.

Discovery and naming

Gases were poorly understood by chemists until the late 1700s. What is air "made of?" That question is difficult to answer for a number of reasons. First, air cannot really be "seen." In fact, it took chemists many years to figure out how to capture air so that they could study it. Also, is ordinary "air" an element or a compound? For many centuries, philosophers said that air was an element. They could not imagine how anything as basic as air could be made of other materials.

Also, is ordinary air different from other kinds of "airs" seen in nature? For example, "air" sometimes comes bubbling out of the ground near oil wells. Today, scientists know that kind of "air" as methane gas (CH₄). But early chemists were not sure how "oil air" differed from ordinary air.

Some important breakthroughs in the study of air occurred in the 1770s. The key was a simple experiment that science students still do today. The experiment begins with an empty bottle being turned upside down in a pan of water. The air in the bottle cannot get out.

If a burning candle is placed inside the bottle with the trapped air, the water rises just a bit. Why does this happen? Early chemists thought that a part of the air was used up as the candle burns. Today, they know that part of the air is oxygen gas. Another part of the air is left behind. That part does not disappear when the candle burns.

This simple experiment shows that air is composed of (at least) two different elements: oxygen and something else. One of the first people to discover what the "something else" is was Scottish physician and chemist Daniel Rutherford (1749-1819). Rutherford carried out an experiment like the candle-in-a-bottle research just described.

Some of the greatest chemists of the time were working on this problem at the same time that Rutherford made his discovery. English chemist Henry Cavendish (1731-1810) probably discovered nitrogen before Rutherford did, but did not publish his findings. And in science, the first person to publish the results of an experiment usually gets credit for the work.

It seems likely that English chemist Joseph Priestley (1733-1804) and Swedish chemist Carl Wilhelm Scheele (1742-86) also discovered nitrogen in the early 1770s. (See sidebar on Scheele in the chlorine entry in Volume 1.)

Chemists debated about the name of this new element for some time. Antoine-Laurent Lavoisier (1743-94), a French chemist and the "father of modern chemistry," preferred the name azote meaning "without life." Lavoisier chose this name because nitrogen does not support breathing, the way oxygen does. (See sidebar on Lavoisier in the oxygen entry in Volume 2.)

The modern name of nitrogen was first suggested in 1790 by French chemist Jean Antoine Claude Chaptal (1756-1832). This name made sense to chemists when they realized that the new gas was present in both nitric acid and nitrates. Thus, nitrogen means "nitrate and nitric acid" (nitro-) and "origin of" (-gen).

Physical properties

Nitrogen is a colorless, odorless, tasteless gas with a density of 1.25046 grams per liter. By comparison, the density of air is about 1.29 grams per liter. Nitrogen changes from a gas into a liquid at a temperature of -195.79°C (-320.42°F) . It changes from a liquid to a solid at a temperature of -210.01°C (-346.02°F). When it freezes, it becomes a white solid that looks like snow. Nitrogen is slightly soluble in water. About two liters of nitrogen can be dissolved in 100 liters of water.

Chemical properties

At room temperature, nitrogen is a very inactive gas. It does not combine with oxygen, hydrogen, or most other elements. Nitrogen will combine with oxygen, however, in the presence of lightning or a spark. The electrical energy from either of those sources causes nitrogen and oxygen to form nitric oxide:

Nitric oxide is more active than free nitrogen. For example, nitric oxide combines with oxygen and water in the atmosphere to make nitric acid. When it rains, nitric acid is carried to the earth. There it combines with metals in the Earth's crust. Compounds known as nitrates and nitrites are formed.

Changing nitrogen as an element to nitrogen in compounds is called nitrogen fixation. The reaction between nitrogen and oxygen in the air when lightning strikes is an example of nitrogen fixation.

Certain bacteria have developed methods for fixing nitrogen. These bacteria live on the root hairs of plants. They take nitrogen out of air dissolved in the ground and convert it to compounds, such as nitrates. Those nitrates are used to make protein molecules, compounds vital to the building and growth of cells.

Plants, animals, and humans do not have the ability to fix nitrogen. All living organisms on Earth depend on soil bacteria to carry out this process. Plants can grow because the bacteria fix nitrogen for them. They use the fixed nitrogen to make proteins. Animals and humans can survive because they eat plants. They also depend on the soil bacteria that allow plants to make proteins. So all living creatures rely on soil bacteria to fix their nitrogen for them and, therefore, to survive.

Occurrence in nature

Nitrogen is a fairly common element in the Earth's crust. It occurs primarily as nitrates and nitrites. Nitrogen is by far the most important element in the Earth's atmosphere. It makes up 78.084 percent of the atmosphere.

Nitrogen combines with oxygen in the presence of lightning or a spark. The electrical energy from those sources causes nitrogen and oxygen to form nitric oxide.

Isotopes

Two naturally occurring isotopes of nitrogen exist, nitrogen-14 and nitrogen-15. Isotopes are two or more forms of an element. Isotopes differ from each other according to their mass number. The number written to the right of the element's name is the mass number. The mass number represents the number of protons plus neutrons in the nucleus of an atom of the element The number of protons determines the element, but the number of neutrons in the atom of any one element can vary. Each variation is an isotope.

Five radioactive isotopes of nitrogen are known also. A radioactive isotope is one that breaks apart and gives off some form of radiation. Radioactive isotopes are produced when very small particles are fired at atoms. These particles stick in the atoms and make them radioactive.

None of the radioactive isotopes of nitrogen has any important commercial use. However, nitrogen-15 is used quite often in tracer studies. A tracer is a radioactive isotope whose presence in a system can be detected. Normally, tracer studies use radioactive isotopes. These isotopes give off radiation that can be detected with instruments. Nitrogen-15 is used for a different reason. A compound made with nitrogen-15 will weigh just a little bit more than one made with nitrogen-14. There are simple chemical methods for detecting whether a heavier compound or a lighter one is present in a system. Thus, nitrogen-15 can be used to trace the path of nitrogen through a system.

Extraction

Nitrogen is almost always made from liquid air. Liquid air is made by cooling normal atmospheric air to very low temperatures. As the temperature drops, the gases contained in air turn into liquids. At -182.96°C (-297.33°F) , oxygen changes from a gas into a Liquid. At -195.79°C (-320.42°F), nitrogen changes from a gas into a liquid. And so on. Eventually, all the gases in air can be made to liquefy (change into a liquid).

The reverse process also takes place. Suppose liquid air in a container warms up slowly. When its temperature reaches -195.79°C, liquid nitrogen changes back to a gas. A container can be put into place to catch the nitrogen as it boils off the liquid air. When the temperature reaches -182.96°C, oxygen changes from a liquid back to a gas. Another container can be put into place. The escaping oxygen can be collected. All of the gases in atmospheric air can be produced by this method.

Large amounts of nitrogen gas are produced in this way. In fact, nitrogen is second only to sulfuric acid in terms of production. In 1996, more than a trillion cubic feet of nitrogen gas were produced in the United States alone.

Uses

Nitrogen gas is used where an inert atmosphere is needed. An inert atmosphere is one that does not contain active elements. Ordinary air is not an inert atmosphere. It contains oxygen. Oxygen tends to react with other elements.

Suppose an ordinary light bulb were filled with air. When the bulb is turned on, an electric current runs through the metal filament (wire) inside the bulb. The filament gets very hot, begins to glow, and gives off light.

But a hot metal wire will react quickly with oxygen in ordinary air. The metal combines with oxygen to form a compound of the metal. The metal compound will not conduct an electric current. The bulb will "burn out" very quickly.

An easy solution to that problem is to use nitrogen instead of ordinary air in the light bulb. Nitrogen does not react with other elements very well, even when they get hot. The filament can get very hot, but the metal of which it is made will not combine with nitrogen gas. The nitrogen gas is an inert atmosphere for the bulb.

Another use for inert atmospheres is in protecting historic documents. Suppose the Declaration of Independence were simply left on top of a table for people to see. The paper and ink in the document would soon begin to react with oxygen in the air. They would both begin to decay. Before long, the document would begin to fall apart.

Instead, important documents like the Declaration of Independence are kept in air-tight containers filled with nitrogen gas. The documents are protected from oxygen and other gases in the air with which they might react.

Fairly simple methods are now available for changing nitrogen gas into liquid nitrogen. Liquid nitrogen is used to freeze other materials. The temperature of the nitrogen has to be reduced to -195.79°C (-320.42°F) for this change to occur.

Historic documents like the Declaration of Independence are protected in air-tight containers filled with nitrogen gas. This keeps them from oxygen and other gases that would cause them to decay.

Today, it is possible to buy large containers of liquid nitrogen. The liquid nitrogen can be used, then, to freeze other materials. For example, foods can be frozen simply by dipping them into large vats of liquid nitrogen. The frozen foods in a grocery store are usually produced this way. Liquid nitrogen can also be used to keep foods cold when they are being transported from one place to another.

Compounds

Nitrogen is the starting point for an important group of compounds. First, nitrogen is combined with hydrogen to make ammonia (NH₃). The production of ammonia is sometimes called industrial nitrogen fixation.

The formation of ammonia from nitrogen and hydrogen is very difficult to accomplish. The two elements do not easily combine. Finding a way to make nitrogen and hydrogen combine was one of the great scientific discoveries of the twentieth century.

The nitrogen compound ammonia is used by most farmers in synthetic fertilizers to ensure large crops.

That discovery was made by German chemist Fritz Haber (1868-1934) in 1905. He found that nitrogen and hydrogen would combine if they were heated to a very high temperature with a very high pressure. He also found that a catalyst was needed to make the reaction occur. The catalyst he used was iron metal, though other metals are sometimes used. A catalyst is a substance used to speed up or slow down a chemical reaction. The catalyst does not undergo any change itself during the reaction.

The significance of Haber's discovery was soon apparent. World War I began in 1914. Very soon, Germany was no longer able to get nitrates from Chile. Nitrates were essential to the war effort for making explosives. But ships carrying nitrates from Chile to Germany were usually not able to get across the Atlantic Ocean.

However, nitrates can be made from ammonia using the Haber process. Germany was then able to make all the ammonia it needed. The ammonia was converted to nitrates for explosives. So a lack of nitrates from Chile did not stop the German war machine. Instead, it went on fighting for another four years.

Ammonia usually ranks about number 5 or 6 among the most highly produced chemicals in the United States. The most important use of ammonia is in synthetic fertilizers. A synthetic fertilizer is a mixture of compounds used to make plants grow better. Most farmers use huge amounts of synthetic fertilizer every year to ensure large crops.

In 1996, more than 38 million tons of nitrogen-containing synthetic fertilizer was made in the United States. Nearly half of that was anhydrous ammonia. Anhydrous means "without water." Anhydrous ammonia is simply ammonia gas. It is stored in large tanks. Farmers inject anhydrous ammonia directly into the ground to produce strong and healthy plants.

Ammonia is also found in many household cleaners, especially glass-cleaning and grease-cutting products.

The largest producer of ammonia in the world is China. Other large producers are the United States, India, Russia, Canada, Ukraine, Indonesia, and the Netherlands.

Nitrogen's role in the Oklahoma City bombing

O n April 19, 1995, a bomb exploded at the Alfred P. Murrah Federal Building in Oklahoma City, Oklahoma. The devastating impact of the explosion destroyed the building within eight seconds. Each of the nine floors collapsed on top of one another. The bomb killed 168 people and injured hundreds more. The nation was stunned to learn that the attack had not been committed by international terrorists, but by an American, Timothy McVeigh.

While searching the home of McVeigh's accomplice (a partner in a crime), Terry Nichols, investigators found a receipt for ammonium nitrate. This nitrogen compound is most commonly purchased by farmers. They use it as a fertilizer for their crops. But it can also be used as an explosive. The amount on the receipt was for 2,000 pounds. Since neither Nichols nor McVeigh was farming at the time the fertilizer was purchased, there was no real reason to purchase such a large amount of fertilizer.

Investigators determined that an ammonium nitrate bomb did, indeed, destroy the Murrah Building. The explosion resulted from a mixture of 4,800 pounds of ammonium nitrate and fuel oil from twenty plastic drums.

In June 1997, McVeigh was found guilty of all eleven charges against him, including eight counts of first-degree murder of federal agents. He was sentenced to death later that month.

Ammonia can also be converted into other forms. For example, it can be combined with nitric acid (HNO₃) to form ammonium nitrate (NH₄N0₃). And it can be combined with sulfuric acid (H₂S0₄) to make ammonium sulfate ((NH₄)₂S0₄):

In 1996, about 2.9 million tons of ammonium nitrate and 2.6 million tons of ammonium sulfate were produced as fertilizers. These two compounds usually rank about number 14 and number 30 among chemicals produced in the United States

Ammonium nitrate and ammonium sulfate both have other uses also. For example, ammonium nitrate is used to make explosives, fireworks, insecticides and herbicides (chemicals that kill insects and weeds), and rocket fuel. Ammonium sulfate is also used in water treatment systems, as a food additive, in the tanning of leather, in fireproofing materials, and as a food additive.

Yet another important compound of nitrogen is nitric acid (HNO₃). Nitric acid is made by reacting ammonia with oxygen:

Nitric acid usually ranks about number 13 among chemicals produced in the United States. The major use of nitric acid is to make ammonium nitrate as a synthetic fertilizer. Nitric acid is also used to make explosives, dyes, certain kinds of synthetic rubber and plastics, and in the preparation of metals.

Health effects

Nitrogen is absolutely essential to all Living organisms. It is an important part of all protein molecules. Proteins are the building material in all kinds of cells. They are also used for many other functions. For example, all living organisms use hormones to send chemical messages from one cell to another. Hormones are proteins.

Nitrogen is absolutely essential to all living organisms. It is an important part of all protein molecules.

Nitrogen is also used to make nucleic acids. Nucleic acids have many important functions in living organiams. For one thing, they store the organism's genetic information. The genetic information is the set of instructions that tell every cell what its job in the organism is. It passes on that information from one generation to the next.

Nitrogen

Discovery and naming

General properties

Where it comes from

How nitrogen is obtained

How it is used

Chemistry and compounds

Environmental issues

Resources

Nitrogen—the fifth most abundant element in the universe—is the non-metallic chemical element of atomic number 7, with a symbol N, atomic weight 14.0067, specific gravity 0.96737 (compared to air), melting point -345.74^°F (-209.86^°C), boiling point -320.44^°F (-195.8^°C). Many of its compounds have been discovered in interstellar space. For example, the Far Ultraviolet Spectroscopic Explorer (FUSE), a space-based telescope launched in 1999 and operated by Johns Hopkins University (Baltimore, Maryland), have made such discoveries.

Nitrogen is located in group 15 of the periodic table. The elements that make up this group are sometimes known as the nitrogen family, after nitrogen itself. It has two stable isotopes: nitrogen-14, with an abundance of 99.634%, and nitrogen-15, with an abundance of 0.366%. At least five radioactive isotopes of the element have been prepared, with atomic weights of 12, 13, 16, 17, and 18.

Discovery and naming

Nitrogen was discovered in the early 1770s, probably by a number of different chemists at almost the same time. The reason for this multiple discovery was that chemists were just learning how to capture, store, and study gases during this period. Prior to that time, all gases were thought of as different forms of air. In fact, they were not called gases, but airs. Credit for the discovery of nitrogen is often given to Scottish physician and chemist Daniel Rutherford (1824–1887) because he was the first person to announce his discovery and give a detailed description of nitrogen. Nitrogen was first identified as the product left behind when a substance was burned in a closed sample of air (which, of course, removed the oxygen component of air). It seems likely that English chemist Henry Cavendish (1731–1810) discovered nitrogen at about the same time, if not earlier. However, Cavendish failed to publish his results until somewhat later than Rutherford. Joseph Priestly and Carl Scheele could also claim to have discovered the element at about the same time.

The modern name for the element was suggested in 1790 by the French chemist Jean Antoine Claude Chaptal (1756–1832) based on the fact that it occurs in both nitric acid and nitrates. Thus, nitrogen means nitrate and nitric acid (nitro- ) and origin of (-gen ).

General properties

Nitrogen is a colorless, odorless, tasteless gas composed of diatomic molecules. Its molecules are represented by the formula N₂. The triple bond that holds the two nitrogen atoms together in a nitrogen molecule is very strong, and nitrogen is, therefore, a relatively unreactive element. When a substance burns in air, for example, it reacts with oxygen; however, in most cases, not with the nitrogen that is also present in air. One important exception involves the combustion of magnesium in air, in which case both magnesium oxide and magnesium nitride are formed.

Nitrogen has a number of important industrial and commercial uses, as do many of its compounds. The most common of these compounds are those that contain nitrogen and hydrogen (some form of ammonia or its derivative compounds) or nitrogen, oxygen, and a third element, that is, the nitrates and nitrites.

Where it comes from

Nitrogen is the most abundant element in the atmosphere, making up about 78% by volume of the air that surrounds the Earth. The element is much less common in the Earth’s crust, however, where it ranks 33rd (along with gallium). Scientists estimate that the average concentration of nitrogen in crustal rocks is about 19 parts per million, less than that of elements such as neodymium, lanthanum, yttrium, and scandium, but greater than that of well-known metals such as lithium, uranium, tungsten, silver, mercury, and platinum.

The most important naturally-occurring compounds of nitrogen are potassium nitrate (saltpeter), found primarily in India, and sodium nitrate (Chile saltpeter), found primarily in the desert regions of Chile and other South American nations. Nitrogen is also an essential component of the proteins found in all living organisms.

How nitrogen is obtained

Nitrogen is produced commercially almost exclusively from air, most commonly by the fractional distillation of liquid air. In this process, air is first cooled to a temperature below that of the boiling points of its major components, a temperature somewhat less than-328^°F (-200^°C). The liquid air is then allowed to warm up, allowing the lower-boiling-point nitrogen to evaporate from the mixture first. Nitrogen gas escaping from the liquid air is then captured, cooled, and then liquefied once more.

This process produces a high-quality product that generally contains less than 20 parts per million of oxygen. Both an oxygen-free form of nitrogen (containing less than two parts per million of oxygen) and an ultrapure nitrogen (containing less than 10 parts per million of argon) are also available commercially.

A number of methods are available for preparing nitrogen from its compounds in the laboratory on a small scale. For example, a hot aqueous solution of ammonium nitrite decomposes spontaneously to give elemental nitrogen and water. The heating of barium or sodium azide (NaN₃ or Ba[N₃]₂) also yields free nitrogen. In another approach, passing ammonia gas over a hot metallic oxide will result in the formation of free nitrogen, the free metal, and water. Yet another route is the reaction between ammonia and bromine, resulting in the formation of nitrogen and ammonium bromide.

How it is used

As more and more uses for the element have been found, the demand for nitrogen in the United States increased dramatically over the last several decades of the twentieth century. In 1988, for example, it was the second most widely produced chemical in the United States, with a production of 52.1 billion lbs (23.7 billion kg). However, in 2000, nitrogen production in the United States decreased to 36.4 billion lbs (16.5 billion kg). Since the 1990s, more nitrogen is being imported and less of it is being produced in the United States. In fact, the United States went from the world’s largest exporter of nitrogen fertilizer in the 1980s to the world’s largest importer in the 1990s. Industry experts expect further declines in U.S. production in the 2000s.

The most important applications of nitrogen depend on the element’s inertness. For example, it is used as a blanketing atmosphere in metallurgical processes where the presence of oxygen would be harmful. In the processing of iron and steel, for example, a blanket of nitrogen placed above the metals prevents their reacting with oxygen, forming undesirable oxides in the final products.

The purging of tanks, pipes, and other kinds of containers with nitrogen can also prevent the possibility of fires. In the petroleum industry, for example, the processing of organic compounds in the presence of air creates the possibility of fires, a possibility that can be avoided by covering the reactants with pure nitrogen.

Nitrogen is also used in the production of electronic components. Assembly of computer chips and other electronic devices can take place with all materials submerged in a nitrogen atmosphere, preventing oxidation of any of the materials in use. Nitrogen is often used as a protective agent during the processing of foods so that decay (oxidation) does not occur.

Another critical use of nitrogen is in the production of ammonia by the Haber process, named after its inventor, German chemist Fritz Haber (1868–1934). The Haber process involves the direct synthesis of ammonia from its elements, nitrogen, and hydrogen. The two gases are combined at temperatures of 932 to 1,292^°F (500 to 700^°C) under a pressure of several hundred atmospheres over a catalyst such as finely divided nickel. One of the major uses of the ammonia produced by this method is in the production of synthetic fertilizers.

About one-third of all nitrogen produced is used in its liquid form. For example, liquid nitrogen is used for quick-freezing foods and for preserving foods in transit. Materials can also be processed at the very low temperatures of liquid nitrogen in ways that they cannot be handled at room temperature. For example, most forms of rubber are too soft and pliable for machining at room temperature. They can, however, first be cooled in liquid nitrogen and then handled in a much more rigid form.

Chemistry and compounds

Although molecular nitrogen is relatively inert, it will combine with a number of other elements at high temperatures. When it reacts with metals such as aluminum, magnesium, lithium, calcium, barium, strontium, and titanium, the products are known as nitrides. Lithium nitride (Li₃N), for example, is used to provide nitrogen in a variety of metallurgical operations.

Nitrogen and oxygen combine (again, at high temperatures) directly and indirectly to form a series of compounds that include nitrous oxide (or dinitrogen monoxide; N₂O), nitric oxide (or nitrogen monoxide; NO), dinitrogen trioxide (or nitrous anhydride; N₂O₃), nitrogen dioxide (NO₂), and dinitrogen pentoxide (or nitric anhydride; N₂O₅). Nitrogen and the halogens also react with each other to form a series of very unstable, explosive compounds that include nitrogen trifluoride (NF₃), nitrogen trichloride (NCl₃), and nitrogen triiodide (NI₃).

The most common compounds of nitrogen are those in which the element demonstrates oxidation numbers of 3^-, 3⁺, or 5⁺. Ammonia (NH₃) and its compounds (ammonium compounds) are examples of the first of these, the nitrites (NO₂^-) are examples of the 3⁺oxidation state, and the nitrates (NO₃^-) are examples of the 5⁺ oxidation state.

The process by which nitrogen is cycled through the environment, from plants to animals to the atmosphere and back to plants, is known as the nitrogen cycle. In that cycle, nitrogen gas in the atmosphere is converted (fixed) to a combined form by the action of lightning, in which it is converted to an oxide of nitrogen, or by certain nitrogen-fixing bacteria in the soil, which change it into nitrates and nitrites. The combined nitrogen is then taken up by plants and used to form plant proteins.

Plant proteins are eaten by animals, who convert the proteins into animal proteins. When an animal dies, the proteins are returned to the soil, where denitrifying bacteria break down compounds of nitrogen and return nitrogen to the atmosphere in the form of an element.

Three compounds of nitrogen traditionally rank in the top 25 among those chemicals produced in the largest volume in the United States. They are ammonia, nitric acid, and ammonium nitrate. All three of these compounds are extensively used in agriculture as synthetic fertilizers. More than 80% of the ammonia produced, for example, goes to the production of synthetic fertilizers.

In addition to its agricultural role, nitric acid is also an important raw material in the production of explosives. Trinitrotoluene (TNT), gunpowder, nitroglycerin, dynamite, and smokeless powder are all examples of the kind of explosives made from nitric acid. Slightly more than 5% of the nitric acid produced is also used in the synthesis of adipic acid and related compounds used in the manufacture of nylon.

Environmental issues

Some compounds of nitrogen have been implicated in a variety of environmental questions. For example, sodium nitrate and sodium nitrite have been used as food additives because of their ability to inhibit the growth of disease causing microorganisms. The compounds are most widely used in preserving meats such as bacon, ham, sausage, hot dogs, and bologna, as well as some fish products.

However, questions have been raised about the possible effects of these additives on human health. Nitrites, for example, appear to decrease the ability of a young child’s blood to carry oxygen. In addition, nitrites combine with organic compounds known as amines to form a family of toxic compounds known as the nitrosoamines. These hazards have prompted some scientists and non-scientists alike to call for the ban of nitrates and nitrites as food additives.

Oxides of nitrogen are also involved in problems of air pollution. Although oxygen and nitrogen do not combine with each other at room temperature, they do react at elevated temperatures, such as those produced by an internal combustion engine. As a motor vehicle is operated, nitric oxide is constantly being produced. This oxide, however, readily reacts

KEY TERMS

Fixation —The process by which elemental nitrogen in the atmosphere is converted to a compound, such as a nitrate or an ammonium compound.

Fractional distillation —A process by which two or more substances are separated from each other by allowing them to boil or evaporate at their own distinct boiling points.

Haber process —The chemical process by which nitrogen and hydrogen are combined with each other at high temperature and pressure over a catalyst to produce ammonia.

Isotopes —Two molecules in which the number of atoms and the types of atoms are identical, but their arrangement in space is different, resulting in different chemical and physical properties.

Nitrogen cycle —A series of chemical reactions by which elemental nitrogen in the atmosphere is converted to nitrates and ammonium compounds, those compounds are processed through the plant and animal world, and then are returned to the atmosphere as free nitrogen.

with oxygen in the air to form nitrogen dioxide, a reddish-brown toxic gas. The tan color that is sometimes associated with smog in urban areas is caused by the presence of nitrogen dioxide. Since nitrogen dioxide is harmful to humans and other animals at low concentrations and toxic at higher levels, its presence in polluted air is a serious environmental issue.

See also Element, chemical; Gases, liquefaction of.

Resources

BOOKS

Emsley, John. Nature’s Building Blocks: An A-Z Guide to the Elements. Oxford: Oxford University Press, 2002.

Oxtoby, David W., et al. The Principles of Modern Chemistry. 5th ed. Pacific Grove, CA: Brooks/Cole, 2002.

Siekierski, Slawomir. Concise Chemistry of the Elements. Chichester, UK: Horwood Publishing, 2002.

Trefil, James. Encyclopedia of Science and Technology. The reference Works, Inc., 2001.

Tro, Nivaldo J. Introductory Chemistry. Upper Saddle River, NJ: Pearson Education, 2006.

David E. Newton

Nitrogen

Nitrogen is the non-metallic chemical element of atomic number 7, with a symbol N, atomic weight 14.0067, specific gravity 0.96737 (compared to air), melting point -345.74°F (-209.86°C), boiling point -320.44°F (-195.8°C).

Nitrogen is a non-metallic element located in group 15 of the periodic table . It has two stable isotopes: nitrogen-14, with an abundance of 99.634%, and nitrogen-15, with an abundance of 0.366%. At least five radioactive isotopes of the element have been prepared, with atomic weights of 12, 13, 16, 17, and 18.

Credit for the discovery of nitrogen is usually given to the Scottish physician Daniel Rutherford in 1772, although Henry Cavendish, Joseph Priestly, and Carl Scheele could also claim to have discovered the element at about the same time. Nitrogen was first identified as the product left behind when a substance was burned in a closed sample of air (which, of course, removed the oxygen component of air).

General properties

Nitrogen is a colorless, odorless, tasteless gas composed of diatomic molecules. Its molecules are represented by the formula N2. The triple bond that holds the two nitrogen atoms together in a nitrogen molecule is very strong, and nitrogen is, therefore, a relatively unreactive element. When a substance burns in air, for example, it reacts with oxygen but, in most cases, not with the nitrogen that is also present in air. One important exception involves the combustion of magnesium in air, in which case both magnesium oxide and magnesium nitride are formed.

Nitrogen has a number of important industrial and commercial uses, as do many of its compounds. The most common of these compounds are those that contain nitrogen and hydrogen (some form of ammonia or its derivative compounds) or nitrogen, oxygen, and a third element, that is, the nitrates and nitrites.

Where it comes from

Nitrogen is the most abundant element in the atmosphere, making up about 78% by volume of the air that surrounds the Earth . The element is much less common in the Earth's crust, however, where it ranks 33rd (along with gallium). Scientists estimate that the average concentration of nitrogen in crustal rocks is about 19 parts per million, less than that of elements such as neodymium, lanthanum, yttrium , and scandium, but greater than that of well-known metals such as lithium , uranium , tungsten, silver, mercury, and platinum.

The most important naturally-occurring compounds of nitrogen are potassium nitrate (saltpeter), found primarily in India, and sodium nitrate (Chile saltpeter), found primarily in the desert regions of Chile and other South American nations. Nitrogen is also an essential component of the proteins found in all living organisms.

How nitrogen is obtained

Nitrogen is produced commercially almost exclusively from air, most commonly by the fractional distillation of liquid air. In this process, air is first cooled to a temperature below that of the boiling points of its major components, a temperature somewhat less than - 328°F (-200°C). The liquid air is then allowed to warm up, allowing the lower-boiling-point nitrogen to evaporate from the mixture first. Nitrogen gas escaping from the liquid air is then captured, cooled, and then liquefied once more.

This process produces a high-quality product that generally contains less than 20 parts per million of oxygen. Both an "oxygen-free" form of nitrogen (containing less than two parts per million of oxygen) and an "ultra-pure" nitrogen (containing less than 10 parts per million of argon) are also available commercially.

A number of methods are available for preparing nitrogen from its compounds in the laboratory on a small scale. For example, a hot aqueous solution of ammonium nitrite decomposes spontaneously to give elemental nitrogen and water . The heating of barium or sodium azide (NaN3 or Ba[N3]2) also yields free nitrogen. In another approach, passing ammonia gas over a hot metallic oxide will result in the formation of free nitrogen, the free metal , and water. Yet another route is the reaction between ammonia and bromine, resulting in the formation of nitrogen and ammonium bromide.

How we use it

As more and more uses for the element have been found, the demand for nitrogen has increased dramatically over the past few decades. In 1988, for example, it was the second most widely produced chemical in the United States, with a production of 52.1 billion lbs (23.7 billion kg).

The most important applications of nitrogen depend on the element's inertness. For example, it is used as a blanketing atmosphere in metallurgical processes where the presence of oxygen would be harmful. In the processing of iron and steel , for example, a blanket of nitrogen placed above the metals prevents their reacting with oxygen, forming undesirable oxides in the final products.

The purging of tanks, pipes, and other kinds of containers with nitrogen can also prevent the possibility of fires. In the petroleum industry, for example, the processing of organic compounds in the presence of air creates the possibility of fires, a possibility that can be avoided by covering the reactants with pure nitrogen.

Nitrogen is also used in the production of electronic components. Assembly of computer chips and other electronic devices can take place with all materials submerged in a nitrogen atmosphere, preventing oxidation of any of the materials in use. Nitrogen is often used as a protective agent during the processing of foods so that decay (oxidation) does not occur.

Another critical use of nitrogen is in the production of ammonia by the Haber process, named after its inventor, the German chemist Fritz Haber. The Haber process involves the direct synthesis of ammonia from its elements, nitrogen and hydrogen. The two gases are combined at temperatures of 932–1,292°F (500–700°C) under a pressure of several hundred atmospheres over a catalyst such as finely divided nickel. One of the major uses of the ammonia produced by this method is in the production of synthetic fertilizers .

About a third of all nitrogen produced is used in its liquid form. For example, liquid nitrogen is used for quick-freezing foods and for preserving foods in transit. Materials can also be processed at the very low temperatures of liquid nitrogen in ways that they can not be handled at room temperature. For example, most forms of rubber are too soft and pliable for machining at room temperature. They can, however, first be cooled in liquid nitrogen and then handled in a much more rigid form.

Chemistry and compounds

Although molecular nitrogen is relatively inert, it will combine with a number of other elements at high temperatures. When it reacts with metals such as aluminum , magnesium, lithium, calcium , barium, strontium, and titanium , the products are known as nitrides. Lithium nitride (Li3N), for example, is used to provide nitrogen in a variety of metallurgical operations.

Nitrogen and oxygen combine (again, at high temperatures) directly and indirectly to form a series of compounds that include nitrous oxide (or dinitrogen monoxide; N2O), nitric oxide (or nitrogen monoxide; NO), dinitrogen trioxide (or nitrous anhydride; N2O3), nitrogen dioxide (NO2), and dinitrogen pentoxide (or nitric anhydride; N2O5). Nitrogen and the halogens also react with each other to form a series of very unstable, explosive compounds that include nitrogen trifluoride (NF3), nitrogen trichloride (NCl3), and nitrogen triiodide (NI3).

The most common compounds of nitrogen are those in which the element demonstrates oxidation numbers of 3-, 3+, or 5+. Ammonia (NH3) and its compounds (ammonium compounds) are examples of the first of these, the nitrites (NO-₂ ) are examples of the 3+oxidation state , and the nitrates (NO-₃ ) are examples of the 5+ oxidation state.

The process by which nitrogen is cycled through the environment, from plants to animals to the atmosphere and back to plants, is known as the nitrogen cycle . In that cycle, nitrogen gas in the atmosphere is converted ("fixed") to a combined form by the action of lightning , in which it is converted to an oxide of nitrogen, or by certain nitrogen-fixing bacteria in the soil , which change it into nitrates and nitrites. The combined nitrogen is then taken up by plants and used to form plant proteins.

Plant proteins are eaten by animals, who convert the proteins into animal proteins. When an animal dies, the proteins are returned to the soil, where denitrifying bacteria break down compounds of nitrogen and return nitrogen to the atmosphere in the form of an element.

Three compounds of nitrogen traditionally rank in the top 25 among those chemicals produced in the largest volume in the United States. They are ammonia (number five in 1988), nitric acid (number 12 in 1988), and ammonium nitrate (number 14 in 1988). All three of these compounds are extensively used in agriculture as synthetic fertilizers. More than 80% of the ammonia produced, for example, goes to the production of synthetic fertilizers.

In addition to its agricultural role, nitric acid is also an important raw material in the production of explosives . Trinitrotoluene (TNT), gunpowder, nitroglycerin, dynamite, and smokeless powder are all examples of the kind of explosives made from nitric acid. Slightly more than 5% of the nitric acid produced is also used in the synthesis of adipic acid and related compounds used in the manufacture of nylon.

Environmental issues

Some compounds of nitrogen have been implicated in a variety of environmental questions. For example, sodium nitrate and sodium nitrite have been used as food additives because of their ability to inhibit the growth of disease-causing microorganisms . The compounds are most widely used in preserving meats such as bacon, ham, sausage, hot dogs, and bologna, as well as some fish products.

However, questions have been raised about the possible effects of these additives on human health. Nitrites, for example, appear to decrease the ability of a young child's blood to carry oxygen. In addition, nitrites combine with organic compounds known as amines to form a family of toxic compounds known as the nitrosoamines. These hazards have prompted some scientists and nonscientists alike to call for the ban of nitrates and nitrites as food additives.

Oxides of nitrogen are also involved in problems of air pollution . Although oxygen and nitrogen do not combine with each other at room temperature, they do react at elevated temperatures, such as those produced by an internal combustion engine . As a motor vehicle is operated, nitric oxide is constantly being produced. This oxide, however, readily reacts with oxygen in the air to form nitrogen dioxide, a reddish-brown toxic gas. The tan color that is sometimes associated with smog in urban areas is caused by the presence of nitrogen dioxide. Since nitrogen dioxide is harmful to humans and other animals at low concentrations and toxic at higher levels, its presence in polluted air is a serious environmental issue.

See also Element, chemical; Gases, liquefaction of.

Resources

books

Emsley, John. Nature's Building Blocks: An A-Z Guide to theElements. Oxford: Oxford University Press, 2002.

Greenwood, N.N., and A. Earnshaw. Chemistry of the Elements. 2nd ed. Oxford: Butterworth-Heinneman Press, 1997.

Hawley, Gessner G., ed. The Condensed Chemical Dictionary. 9th ed. New York: Van Nostrand Reinhold, 1977.

Oxtoby, David W., et al. The Principles of Modern Chemistry. 5th ed. Pacific Grove, CA: Brooks/Cole, 2002.

Trefil, James. Encyclopedia of Science and Technology. The Reference Works, Inc., 2001.

periodicals

"Nitrogen." Kirk-Othmer Encyclopedia of Chemical Technology. 4th ed. Suppl. New York: John Wiley & Sons, 1998.

David E. Newton

KEY TERMS

Fixation

—The process by which elemental nitrogen in the atmosphere is converted to a compound, such as a nitrate or an ammonium compound.

Fractional distillation

—A process by which two or more substances are separated from each other by allowing them to boil or evaporate at their own distinct boiling points.

Haber process

—The chemical process by which nitrogen and hydrogen are combined with each other at high temperature and pressure over a catalyst to produce ammonia.

Isotopes

—Two molecules in which the number of atoms and the types of atoms are identical, but their arrangement in space is different, resulting in different chemical and physical properties.

Nitrogen cycle

—A series of chemical reactions by which elemental nitrogen in the atmosphere is converted to nitrates and ammonium compounds, those compounds are processed through the plant and animal world, and then are returned to the atmosphere as free nitrogen.

Nitrogen

melting point: −210°C
boiling point: −196°C
density: 0.0012506 g/cm ³
most common ions: NH ₄⁺, N ³⁻, NO ₂⁻, NO ₃⁻

Nitrogen is a gaseous element that is abundant in the atmosphere as the molecule dinitrogen (N₂). Scottish chemist Daniel Rutherford, Swedish chemist Carl Wilhelm Scheele, and English chemist Henry Cavendish independently discovered the element in 1772. Nitrogen received its name in 1790 from French chemist Jean-Antoine Chaptal, who realized that it was present in nitrate (NO₃⁻) and nitric acid (HNO₃).

Nitrogen is the most abundant terrestrial element in an uncombined state, as it makes up 78 percent of Earth's atmosphere as N₂, but it is a minor component (19 parts per million) of Earth's crust. Nitrogen exists as two isotopes : ¹⁴N (99.63% relative abundance) and ¹⁵N (0.4% abundance). Both isotopes are nuclear magnetic resonance (NMR) active, with the rarer ¹⁵N isotope being utilized more commonly in NMR spectroscopy because of its nuclear spin of one-half.

In its reduced state nitrogen is essential for life because it is a constituent of the nucleotides of deoxyribonucleic acid (DNA ) and ribonucleic acid (RNA ) molecules that encode genetic information) and of the amino acids of proteins. The nitrogen-containing minerals saltpeter (KNO₃) and sodium nitrate (NaNO₃) are found in Chile, India, Bolivia, the former Soviet Union, Spain, and Italy; they were significant as fertilizers and explosives

precursors prior to modern industrial nitrogen fixation. The global nitrogen cycle between the atmosphere and the biosphere is based on continuous exchanges whereby dinitrogen is fixed by the enzyme nitrogenase in symbiotic bacteria associated with some plant roots, by the Haber-Bosch industrial process for the reduction of N₂ with H₂ to ammonia, and by atmospheric oxidation during electrical discharges such as lightning.

Dinitrogen possesses the strongest known chemical bond, with a high bond dissociation energy of 945 kJ mol⁻¹ and a short N-N triple bond length of 109.8 picometers. This colorless, tasteless, odorless gas is relatively unreactive because of its strong N-N triple bond, a stable electronic configuration, and the lack of a dipole moment. Dinitrogen is reduced by lithium metal at room temperature to give the saline (saltlike) lithium nitride (Li₃N).

Dinitrogen is obtained from the atmosphere by either membrane separations or repetitive cycles of compression and cooling (termed liquefaction ), followed by fractional distillation to separate it from other gases. The major uses of dinitrogen are as blanketing atmospheres for chemical processing and metallurgical production, in glove boxes for handling of dioxygen- and water-sensitive compounds, in electronic materials manufacturing, and in food packaging. Liquid dinitrogen is used as a refrigerant in the laboratory and food industry and in the preservation of biological samples.

The major industrial applications of nitrogen-containing compounds are in fertilizers and explosives. The most important nitrogenous compounds are ammonia (NH₃), which is used as a fertilizer, refrigerant, nonaqueous solvent, and precursor for many nitrogen compounds including nylon and plastics; nitric acid (HNO₃); ammonium nitrate (NH₄NO₃), a fertilizer and explosive; fertilizers ammonium phosphate and urea (H₂NC[O]NH₂). Other important oxides include nitrous oxide (N₂O), used as a dental anesthetic and aerosol propellant, and nitric oxide (NO), the simplest stable odd-electron molecule and a short-lived, biologically active neurotransmitter, cytotoxic agent in immunology, vasoconstrictor for blood pressure control, and major component along with NO₂ in acid rain and smog. The strong reducing agent hydrazine (N₂H₄) is used in controlling the attitude of spacecraft and in rocket fuels.

Covalent, intermetallic metal nitrides are among the most stable compounds and are hard, refractory materials that can possess useful properties. For example, titanium nitride (TiN) is used as a gold-colored coating on costume jewelry and as a wear-resistant coating on tool bits; silicon nitride (Si₃N₄) is a strong, thermally stable ceramic material; and gallium nitride (GaN) is a compound semiconductor with optoelectronic applications (e.g., lasers, LEDs).

see also Cavendish, Henry; Gases; Inorganic Chemistry; Scheele, Carl.

Louis Messerle

Bibliography

Greenwood, Norman N., and Earnshaw, A. (1997). Chemistry of the Elements, 2nd edition. Oxford, U.K.: Butterworth-Heinemann.

nitrogen Four-fifths (79%) of the air we breathe consists of nitrogen, nearly all the rest being oxygen. It was known in the late seventeenth century that breathing air with its oxygen removed resulted in death, but only in 1772 did Rutherford isolate nitrogen; soon after, Lavoisier showed that pure nitrogen could not support life, although he misnamed it mephitic or ‘smelly’ air. It is odourless.

The nitrogen we breathe is chemically inert and takes no part in the chemical or metabolic reactions in the body. In this respect it resembles the ‘inert gases’ such as argon and neon which are a small part of the atmosphere. Nitrogen is poorly soluble in water and body liquids, and there is virtually no exchange between the nitrogen we breathe into the lungs and the body itself. However the chemical combinations of nitrogen are crucial for life. It is a definitive component of proteins and their constituents, amino acids. It is present in innumerable other essential chemical components of the body, from vitamins to hormones to enzymes and many other vital molecules; in recent years the ubiquitous importance of nitric oxide (NO) in physiological function has been recognized. It is no exaggeration to say that life only became possible by the creation of nitrogen-containing chemicals. But these chemicals reach their sites in the human body not from inhaled nitrogen, but from ingested plant and animal materials. Only plants (including some bacteria) can convert atmospheric nitrogen to the organic compounds needed for animal life, so plants are the ultimate source of all nitrogenous chemicals in the body.

In proteins nitrogen occurs mostly in amino- (-NH2) groups. During metabolic breakdown of these and other nitrogen-containing substances the nitrogen is not converted to its gaseous form for excretion in the lungs, but forms mainly urea, a small molecular-weight waste product that, as its name implies, is excreted in the urine. Although the metabolism of proteins provides some energy for the body, this is normally far smaller than that due to burning carbohydrates (that contain no nitrogen), and fatty substances (most of which contain no nitrogen). Rather, the amino acids derived from the dietary proteins are taken up by body cells for use in the turnover of their own proteins, which they need to synthesize continually: for their growth and repair, cellular enzymes, secretions and so forth. In health, the necessary daily intake of nitrogen to balance inevitable losses is estimated at about 12 g, which would be contained in about 75 g of protein. In starvation or in the aftermath of serious injury or infection requiring rebuilding of tissues, protein is depleted, mainly from muscle, and is used for the production of glucose by the liver; only adequate nutritional supplements can avoid a state of ‘negative nitrogen balance’, with wasting and weakness.

Although nitrogen is poorly soluble in water, so that little is normally dissolved in body liquids, it is more soluble in fats, which accounts for its role in ‘bends’ or decompression sickness, seen when deep-sea divers breathing air ascend too rapidly to the surface. After a significant time underwater the nitrogen will first have dissolved in the blood, since its pressure is high in the lungs, then passed into the tissues, particularly fat; on rapid ascent (‘decompression’) it comes out of solution to form bubbles in nerves and round joints, causing the characteristic pain of the bends. In practice air is nowadays never used in deep diving, the nitrogen being replaced by helium, which is far less soluble in fat.

Nitrogen under pressure will also cause psychological and neurological disturbances. This condition is called nitrogen narcosis, but long before actual narcosis (sedation and anaesthesia) occurs the nitrogen exerts toxic effects. These include euphoria, fixed and complacent ideas, uncontrollable laughter, and neuromuscular incoordination. Scuba divers may suffer from this, and it has been called the ‘rapture of the depths’. It is not due to any chemical reaction of the nitrogen, since it can also be caused by ‘inert’ gases such as argon, but probably by the solution of the pressurized nitrogen in fatty substances such as the membranes of nerve cells in the brain. Possibly also the nitrogen attracts water to form hydrated forms which disrupt brain cell function. Although nitrogen narcosis may have the same physicochemical basis as decompression sickness, its clinical manifestations are quite different. In many respects it resembles the psychological and neurological effects of acute lack of oxygen, but the mechanisms are probably very dissimiliar.

John Widdicombe