THE NITROGEN CYCLE

CONCEPT

Contrary to popular belief, the air we breathe is not primarily oxygen; by far the greatest portion of air is composed of nitrogen. A colorless, odorless gas noted for its lack of chemical reactivity— that is, its tendency not to bond with other elements—nitrogen plays a highly significant role within the earth system. Both through the action of lightning in the sky and of bacteria in the soil, nitrogen is converted to nitrites and nitrates, compounds of nitrogen and oxygen that are then absorbed by plants to form plant proteins. The latter convert to animal proteins in the bodies of animals who eat the plants, and when an animal dies, the proteins are returned to the soil. Denitrifying bacteria break down these compounds, returning elemental nitrogen to the atmosphere.

HOW IT WORKS

Chemistry and Elements

The concepts we discuss in this essay fall under the larger heading of geochemistry. A branch of the earth sciences that combines aspects of geology and chemistry, geochemistry is concerned with the chemical properties and processes of Earth. Among particular areas of interest in geochemistry are biogeochemical cycles, or the changes that particular elements undergo as they pass back and forth through the various earth systems and particularly between living and nonliving matter. The elements involved in biogeochemical cycles are hydrogen, oxygen, carbon, nitrogen, phosphorus, and sulfur (see Biogeochemical Cycles and Carbon Cycle).
An element is a substance composed of a single type of atom, which cannot be broken down chemically into a simpler substance. Each element is distinguished by its atomic number, or the number of protons (positively charged subatomic particles) in the nucleus, or center, of the atom. On the periodic table of elements, these fundamental substances of the universe are listed in order of atomic number, from hydrogen to uranium—which has the highest atomic number (92) of any element that occurs in nature—and beyond. The elements with an atomic number higher than that of uranium, all of which have been created artificially, play virtually no role in the chemical environment of Earth and are primarily of interest only to specialists in certain fields of chemistry and physics.
Everything that exists in the universe is an element, a compound formed by the chemical bonding of elements, or a mixture of compounds. In order to bond and form a compound, elements experience chemical reactions, which are the result of attractions on the part of electrons (negatively charged subatomic particles) that occupy the highest energy levels in the atom. These electrons are known as valence electrons.


Chemical changes

A chemical change is a phenomenon quite different from a physical change. If liquid water boils or freezes (both of which are examples of a physical change resulting from physical processes), it is still water. Physical changes do not affect the internal composition of an item or items; a chemical change, on the other hand, occurs when the actual composition changes—that is, when one substance is transformed into another. Chemical change requires a chemical reaction, a process whereby the chemical properties of a substance are altered by a rearrangement of atoms.
Liquid nitrogen.
Liquid nitrogen.
There are several clues that tell us when a chemical reaction has taken place. In many chemical reactions, for instance, the substance may experience a change of state or phase—as, for instance, when liquid water is subjected to an electric current through a process known as electrolysis, which separates it into oxygen and hydrogen, both of which are gases. Another clue that a chemical reaction has occurred is a change of temperature. Unlike the physical change of liquid water to ice or steam, however, this temperature change involves an alteration of the chemical properties of the substances themselves. Chemical reactions also may encompass changes in color, taste, or smell.

Nitrogen’s Place Among the Elements

With an atomic number of 7, nitrogen (chemical symbol N) is one of just 19 elements that are nonmetals. Unlike metals, nonmetals are poor conductors of heat and electricity and are not ductile—in other words, they cannot be reshaped easily. The vast majority of elements are metallic,however, the only exceptions being the non-metals as well as six “metalloids,” or elements that display characteristics of both metals and non-metals.
Nitrogen is also one of eight “orphan” non-metals—those nonmetals that do not belong to any family of elements, such as the halogens or noble gases. All six of the elements involved in biogeochemical cycles, in fact, are “orphan” non-metals, with boron and selenium rounding out the list of eight orphans. Sometimes nitrogen is considered the head of a “family” of elements, all of which occupy a column or group on the periodic table.
These five elements—nitrogen, phosphorus, arsenic, antimony, and bismuth—share a common pattern of valence electrons, but otherwise they share little in terms of physical properties or chemical behavior. By contrast, chemicals that truly are related all have a common “family resemblance”: all halogens are highly reactive, for instance, while all noble gases are extremely unreactive.

Abundance

The seventeenth most abundant element on Earth, nitrogen accounts for 0.03% of the planet’s known elemental mass. This may seem very small, but at least nitrogen is among the 18 elements considered relatively abundant. These 18 elements account for all but 0.49% of the planet’s known elemental mass, the remainder being composed of numerous other elements in small quantities. The term known elemental mass takes account of the fact that scientists do not know with certainty the elemental composition of Earth’s interior, though it likely contains large proportions of iron and nickel. The known mass, therefore, is that which exists from the bottom of the crust to the top layers of the atmosphere.
Elemental proportions too small to be measured in percentage points are rendered in parts per million (ppm) or even parts per billion (ppb). Within the crust itself, nitrogen’s share is certainly modest: a concentration of 19 ppm, which ties it with gallium, a metal whose name is hardly a household word, for a rank of thirty-third. On the other hand, this still makes it more abundant in the crust than many quite familiar metals, including lithium, uranium, tungsten, silver, mercury, and platinum.
In Earth’s atmosphere, on the other hand, the proportion of nitrogen is much, much higher. The atmosphere is 78% nitrogen and 21% oxygen, while the noble gas argon accounts for 0.93%. The remaining 0.07% is taken up by various trace gases, including water vapor, carbon dioxide, and ozone, or O3.
In the human body, nitrogen’s share is much more modest than it is in the atmosphere but still 10 times greater than it is in relation to the planet’s total mass. The element accounts for 3% of the body’s mass, making it the fourth most abundant element in the human organism.

Properties and Applications of Nitrogen

The Scottish chemist Daniel Rutherford (1749-1819) usually is given credit for discovering nitrogen in 1772, when he identified it as the element that remained when oxygen was removed from air. Several other scientists at about the same time made a similar discovery.
Because of its heavy presence in air, nitrogen is obtained primarily by cooling air to temperatures below the boiling points of its major components. Nitrogen boils (that is, turns into a gas) at a lower temperature than oxygen: -320.44°F (-195.8°C), as opposed to -297.4°F (-183°C). If air is cooled to -328°F (-200°C), thus solidifying it, and then allowed to warm slowly, the nitrogen boils first and therefore evaporates first. The nitrogen gas is captured, cooled, and liquefied once more.
Nitrogen also can be obtained from such compounds as potassium nitrate or saltpeter, found primarily in India, or from sodium nitrate (Chilean saltpeter), which comes from the desert regions of Chile. To isolate nitrogen chemically, various processes are undertaken in a laboratory—for instance, heating barium azide or sodium azide, both of which contain nitrogen.


Reactions with other elements

Rather than appearing as single atoms, nitrogen is diatomic, meaning that two nitrogen atoms typically bond with each other to form dinitrogen, or N2. Nor do these atoms form single chemical bonds, as is characteristic of most elements; theirs is a triple bond, which effectively ties up the atoms’ valence electrons, making nitrogen an unreactive element at relatively low temperatures.
Even at the temperature of combustion, a burning substance reacts with the oxygen in the air but not with the nitrogen. At very high temperatures, on the other hand, nitrogen combines with other elements, reacting with metals to form nitrides, with hydrogen to form ammonia, with O2 (oxygen as it usually appears in nature, two atoms bonded in a molecule) to form nitrites, and with O3 (ozone) to form nitrates. With the exception of the first-named group, all of these elements are important to our discussion of nitrogen.
Nitrogen combines with hydrogen to form ammonia. Ammonium nitrate, a fertilizer, is also a dangerous explosive. It was used in April of 1995 to blow up the Alfred P. Murrah Federal building in Oklahoma City, killing 168 people.
Nitrogen combines with hydrogen to form ammonia. Ammonium nitrate, a fertilizer, is also a dangerous explosive. It was used in April of 1995 to blow up the Alfred P. Murrah Federal building in Oklahoma City, killing 168 people.

Some uses for nitrogen

In processing iron or steel, which forms undesirable oxides if exposed to oxygen, a blanket of nitrogen is applied to prevent this reaction. The same principle is applied in making computer chips and even in processing foods, since these items, too, are affected detrimentally by oxidation. Because it is far less combustible than air (magnesium is one of the few elements that burns nitrogen in combustion), nitrogen also is used to clean tanks that have carried petroleum or other combustible materials.
As noted, nitrogen combines with hydrogen to form ammonia, used in fertilizers and cleaning materials. Ammonium nitrate, applied primarily as a fertilizer, is also a dangerous explosive, as shown with horrifying effect in the bombing of the Alfred P. Murrah Federal Building in Oklahoma City on April 19, 1995—a tragedy that took 168 lives. Nor is ammonium nitrate the only nitrogen-based explosive. Nitric acid is used in making trinitrotoluene (TNT), nitroglycerin, and dynamite as well as gunpowder and smokeless powder.

Introduction to the Nitrogen Cycle

The nitrogen cycle is the process whereby nitrogen passes from the atmosphere into living things and ultimately back into the atmosphere. In the process, it is converted to nitrates and nitrites, compounds of nitrogen and oxygen that are absorbed by plants in the process of forming plant proteins. These plant proteins, in turn, are converted to animal proteins in the bodies of animals who eat the plants, and when the animal dies, the proteins are returned to the soil. Denitrifying bacteria break down these organic compounds, returning elemental nitrogen to the atmosphere.
Note what happens in the nitrogen cycle and, indeed, in all biogeochemical cycles: organic material is converted to inorganic material through various processes, and inorganic material absorbed by living organisms eventually is turned into organic material. In effect, the element passes back and forth between the realms ofthe living and the nonliving. This may sound a bit mystical, but it is not. To be organic, a substance must be built around carbon in certain characteristic chemical structures, and by inducing the proper chemical reaction, it is possible to break down or build up these structures, thus turning an organic substance into an inorganic one, or vice versa. (For more on this subject, see Carbon Cycle.)

Steps in the cycle

Plants depend on biologically useful forms of nitrogen, the availability of which greatly affects their health, abundance, and productivity. This is particularly the case where plants in a saltwater ecosystem (a community of interdependent organisms) are concerned. Regardless of the specific ecosystem, however, fertilization of the soil with nitrogen has an enormous impact on the growth yield of plant life, which can be critical in the case of crops. Therefore, nitrogen is by far the most commonly applied nutrient in an agricultural setting.
There are several means by which plants receive nitrogen. They may absorb it as nitrate or ammonium, dissolved in saltwater and taken up through the roots, or as various nitrogen oxide gases. In certain situations, plants have a symbiotic, or mutually beneficial, relationship with microorganisms capable of “fixing” atmospheric dinitrogen into ammonia. In any case, plants receive nitrogen and later, when they are eaten by animals, pass these nutrients along the food chain—or rather, to use a term more favored in the earth and biological sciences, the food web.
When herbivorous or omnivorous animals consume nitrogen-containing plants, their bodies take in the nitrogen and metabolize it, breaking it down to generate biochemicals, or chemicals essential to life processes. At some point, the animal dies, and its body experiences decomposition through the activity of bacteria and other decomposers. These microorganisms, along with detritivores such as earthworms, convert nitrates and nitrites from organic sources into elemental nitrogen, which ultimately reenters the atmosphere.

REAL-LIFE APPLICATIONS

Important Forms of Nitrogen

As noted earlier, dinitrogen, or N2, is the form in which nitrogen typically appears when uncombined with other elements. This is also the form of nitrogen in the atmosphere, but it is so chemically unreactive that unlike oxygen, it plays little actual part in sustaining life. Indeed, because nitrogen in the air is essentially “filler” as far as humans are concerned, it can be substituted with helium, as is done in air tanks for divers. This prevents them from experiencing decompression sickness, or “the bends,” which occurs when the diver returns too quickly to the surface, causing nitrogen in the blood to boil.
The dinitrogen in the air is a holdover from long ago in Earth’s development, when volcanoes expelled elements from deep in the planet’s interior to its atmosphere. Owing to its lack of reactivity, dinitrogen never went anywhere. For it to play a role in the functioning of Earth cycles, it must be “fixed,” as discussed later in this essay. In addition to dinitrogen, nitrogen appears in a number of other important inorganic compounds, including nitrite and nitrate; ammonia and ammonium; and nitric oxide, nitrogen dioxide, and nitrous oxide.
Smog blankets Los Angeles in a haze. Nitric oxide reacts with oxygen in the air to form nitrogen oxide, a reddish brown gas that colors smog.
Smog blankets Los Angeles in a haze. Nitric oxide reacts with oxygen in the air to form nitrogen oxide, a reddish brown gas that colors smog.
Nitrite and nitrate are two ionic forms of nitrogen. An ion is an atom or group of atoms that has lost or gained electrons, thus acquiring a net electric charge. Both nitrite and nitrate are anions, or negatively charged ions, designated by the use of superscript minus signs that indicate that each has a net charge of negative 1. Thus, nitrite, in which nitrogen is chemically bonded with two atoms of oxygen, is rendered as NO2-, while the formula for nitrate (nitrogen with three oxygen atoms), is designated as NO3-.

Ammonia and ammonium

Nitrification is a process in which nitrite is produced, whereupon it undergoes a chemical reaction to form nitrate, the principal form of nitrogen nutrition for most plant species. The chemical from which the nitrite is created in the nitrification reaction is ammonium (NH4+), which is formed by the addition of a hydrogen cation, or a positively charged ion (H+), to ammonia, or
NH3. The latter, which is probably familiar to most people in the form of a household cleaner, is actually an extremely abundant compound, both in natural and artificial forms.
Ammonium is soluble, or capable of being dissolved, in water and often is used as a fertilizer. It is attracted to negatively charged surfaces of clays and organic matter in soil and therefore tends to become stuck in one place rather than moving around, as nitrate does. In acidic soils, typically plants receive their nitrogen from ammonium, but most nonacidic soils can use only nitrate. As noted earlier, ammonium may be combined with nitrate to form ammonium nitrate—both a powerful fertilizer and a powerful explosive.

Oxides

Nitrogen reenters the atmosphere in the form of the gas nitric oxide (NO), emitted primarily as the result of combustion reactions. This may occur in one of two ways. Organic nitrogen in bioenergy sources, such as biomass (organisms, their waste products, and their incompletely decomposed remains) or fossil fuels (e.g., coal or oil), may be oxidized. The latter term means that a substance undergoes a chemical reaction with oxygen: combustion itself, which requires the presence of oxygen, is an example of oxidation.

KEY TERMS

Atom: The smallest particle of an element, consisting of protons, neutrons, and electrons. An atom can exist either alone or in combination with other atoms in a molecule.
Atomic number: The number of protons in the nucleus of an atom. Since this number is different for each element, elements are listed on the periodic table in order of atomic number.
Bioenergy: Energy derived from biological sources that are used directly as fuel (as opposed to food, which becomes fuel).
Biogeochemical cycles: The changes that particular elements undergo as they pass back and forth through the various earth systems and specifically between living and nonliving matter. The elements involved in biogeochemical cycles are hydrogen, oxygen, carbon, nitrogen, phosphorus, and sulfur.
Chemical bonding: The joining, through electromagnetic force, of atoms that sometimes, but not always, represent more than one chemical element. The result is usually the formation of a molecule.
Compound: A substance made up of atoms of more than one element chemically bonded to one another.
Decomposers: Organisms that obtain their energy from the chemical breakdown of dead organisms as well as from animal and plant waste products. The principal forms of decomposer are bacteria and fungi.
Decomposition reaction: A chemical reaction in which a compound is broken down into simpler compounds or into its constituent elements. In the earth system, this often is achieved through the help of detritivores and decomposers.
Detritivores: Organisms that feed on waste matter, breaking organic material down into inorganic substances that then become available to the biosphere in the form of nutrients for plants. Their function is similar to that of decomposers; however, unlike decomposers—which tend to be bacteria or fungi—detritivores are relatively complex organisms, such as earthworms or maggots.
Diatomic: A term describing a chemical element that typically exists as molecules composed of two atoms. Nitrogen and oxygen are both diatomic.
Ecosystem: A term referring to a community of interdependent organisms along with the inorganic components of their environment.
Electron: A negatively charged particle in an atom, which spins around the nucleus.
Element: A substance made up of only one kind of atom. Unlike compounds, elements cannot be chemically broken into other substances.
Eutrophication: A state of heightened biological productivity in a body of water, which is typically detrimental to the ecosystem in which it takes place. Eutrophication can be caused by an excess of nitrogen or phosphorus in the form of nitrates and phosphates, respectively.
Food web: A term describing the interaction of plants, herbivores, carnivores, omnivores, decomposers, and detritivores, each of which consumes nutrients and passes it along to other organisms.
Geochemistry: A branch of the earth sciences combining aspects of geology and chemistry, that is, concerned with the chemical properties and processes of Earth—in particular, the abundance and interaction of chemical elements and their isotopes.
Ion: An atom or group of atoms that has lost or gained one or more electrons and thus has a net electric charge. Positively charged ions are called cations, and negatively charged ones are called anions.
Leaching: The removal of soil materials that are in solution, or dissolved in water.
Molecule: A group of atoms, usually but not always representing more than one element, joined in a structure. Compounds are typically made up of molecules.
Periodic table of elements:A chart that shows the elements arranged in order of atomic number along with their chemical symbols and the average atomic mass for each particular element.
Proton: A positively charged particle in an atom.
Reactivity: A term referring to the ability of one element to bond with others. The higher the reactivity, the greater the tendency to bond.
Soluble: Capable of being dissolved.
Valence electrons: Electrons that occupy the highest principal energy level in an atom. These are the electrons involved in chemical bonding.
On the other hand, nitric oxide may enter the atmosphere when atmospheric dinitrogen is combined with oxygen under conditions of high temperature and pressure, as, for instance, in an internal-combustion engine. In the atmosphere, nitric oxide reacts readily with oxygen in the air to form nitrogen dioxide (NO2), a reddish-brown gas that adds to the tan color of smog over major cities.
Yet nitric oxide and nitrogen dioxide, usually designated together as NOx, are also part ofthe life-preserving nitrogen cycle. Gaseous NOx is taken in by plants, or oxidized to make nitrate, and circulated through the biosphere or else cycled directly to the atmosphere. In addition, denitrification, discussed later in this essay, transports nitrous oxide (N2O) into the atmosphere from nitrate-rich soils.

Nitrogen Processes

In order for most organisms to make use of atmospheric dinitrogen, it must be “fixed” into inorganic forms that a plant can take in through its roots and leaves. Nonbiological processes, such as a lightning strike, can bring about dinitrogen fixation. The high temperatures and pressures associated with lightning lead to the chemical bonding of atmospheric nitrogen and oxygen (both of which appear in diatomic form) to create two molecules of nitric oxide.
More often than not, however, dinitrogen fixation comes about through biological processes. Microorganisms are able to synthesize an enzyme that breaks the triple bonds in dinitro-gen, resulting in the formation of two molecules of ammonia for every dinitrogen molecule thus reacted. This effect is achieved most commonly by bacteria or algae in wet or moist environments that offer nutrients other than nitrate or ammonium. In some instances, plants enjoy a symbiotic, or mutually beneficial, relationship with microorganisms capable of fixing dinitrogen.

AMMONIFICATION, NITRIFICATION, AND DENITRIFICATION

Dinitrogen fixation is just one example of a process whereby nitrogen is processed through one or more earth systems. Another is ammonification, or the process whereby nitrogen in organisms is recycled after their death. Enabled by microorganisms that act as decomposers, ammonification results in the production of either ammonia or ammonium. Thus, the soil is fertilized by the decayed matter of formerly living things.
Ammonium, as we noted earlier, also plays a part in nitrification, a process in which it first is oxidized to produce nitrite. Then the nitrite is oxidized to become nitrate, which fertilizes the soil. As previously mentioned, nitrate is useful as a fertilizer only in non-acidic soils; acidic ones, by contrast, require ammonium fertilizer.
In contrast to nitrification is denitrification, in which nitrate is reduced to the form of either nitrous oxide or dinitrogen. This takes place under anaerobic conditions—that is, in the absence of oxygen—and on the largest scale when concentrations of nitrate are highest. Flooded fields, for example, may experience high rates of denitrification.

The Role of Humans

Humans are involved in the nitrogen cycle in several ways, not all of them beneficial. One of the most significant roles people play in the nitrogen cycle is by the introduction of nitrogen-containing fertilizers to the soil. Because nitrogen has a powerful impact on plant growth, farmers are tempted to add more and more nitrate or ammonium or both to their crops, to the point that the soil becomes saturated with it and therefore unable to absorb more.
When the soil has taken in all the nitrogen it can hold, a process of leaching—the removal of soil materials dissolved in water—eventually takes place. Nitrate, in particular, leaches from agricultural sites into groundwater as well as streams and other forms of surface water. This can lead to eutrophication, a state of heightened biological productivity that is ultimately detrimental to the ecosystem surrounding a lake or other body of water. (See Biogeochemical Cycles for more about eutrophication.)
Yet another problem associated with overly nitrate-rich soils is an excessive rate of denitrification. This happens when soils that have been loaded down with nitrates become wet for long periods of time, leading to a dramatic increase in the denitrification rate. As a result, fixed nitrogen is lost, and nitrous oxide is emitted to the air. In the atmosphere nitrous oxide may contribute to the greenhouse effect, possibly helping increase the overall temperature of the planet (see Carbon Cycle and Energy and Earth).

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