NITROGEN SYMBOL:N PERIOD:2 GROUP:15(VA) ATOMICNO:7
NITROGEN SYMBOL:N PERIOD:2 GROUP:15(VA) ATOMICNO:7
ATOMICMASS:14.0067amu VALENCE:1,2,3,4,and5 OXIDATIONSTATES:–1,
+1,–2,+2,+3,+4,+5 NATURALSTATE:Gas
Guide to the Elements |
ORIGINOFNAME:FromthetwoGreekwordsnitronandgenes,whichtogetherstandfor “sodaorsaltpeterforming.” ISOTOPES:Thereare19isotopesofnitrogen,twoofwhicharestable.Thestableonesand theirproportiontothenaturalabundanceofnitrogenonEarthfollow:N-14=99.634% andN-15=0.366%.Theother17isotopesareradioactiveandman-madeinnuclear reactorsandhavehalf-livesrangingfromafewnanosecondsto9.965minutes.
ELECTRONCONFIGURATION EnergyLevels/Shells/Electrons Orbitals/Electrons
s2,p3
Properties In its natural gaseous state, nitrogen is a relatively inert diatomic molecule (N 2 ) that is
colorless, odorless, and tasteless, yet it is responsible for hundreds of active compounds. It makes up about 78% of the air we breathe. We are constantly taking it into our lungs with no stimulation or sensation; therefore, we really do not detect its presence. When liquefied, it is still colorless and odorless and resembles water in density. The melting point of nitrogen is
–209.86°C, its boiling point is –195.8°C, and its density as a gas is 0.0012506 g/cm 3 . Characteristics
There are approximately 4,000 trillion tons of gas in the atmosphere, and nitrogen makes up about 78% of these gases. It is slightly soluble in water and alcohol. It is noncombustible and is considered an asphyxiant gas (i.e., breathing pure nitrogen will deprive the body of oxygen).
Although nitrogen is considered an inert element, it forms some compounds that are very active. Of the diatomic molecules, such as CO 2 , it is difficult to separate the two atoms in nitrogen’s molecules because of their strong binding energy. This is the reason that, along with carbon dioxide, nitrogen gas is stable. However, once separated, the individual atoms of nitro- gen (N) become very reactive and do combine with hundreds of other elements.
Nitrogen can be liquefied easily, making it useful in many applications wherein sustained cooling is needed. At high temperatures, nitrogen reacts with many metals to form nitrides.
AbundanceandSource Nitrogen is the 30th most abundant element on Earth. There is an almost unlimited source
of nitrogen available to us considering that our atmosphere constitutes 4/5, or over 78%, of the nitrogen by volume. Over 33 million tons of nitrogen is produced each year by liquefy- ing air and then using fractional distillation to produce nitrogen as well as other gases in
210 | The History and Use of Our Earth’s Chemical Elements the atmosphere. During this process the air is cooled and then slowly warmed to fractional
temperature points at which each specific gas in the air will “boil” off. (Note: Oxygen, argon, carbon dioxide, and nitrogen all have specific boiling points and these gases can be used to collect the specific gas during the fractionation process.) When the temperature –reaches –195.8°C, the nitrogen is boiled off and collected.
There is a balance of nitrogen with other gases in the atmosphere that is maintained by what is called the nitrogen cycle. This cycle includes several processes, including nitrogen fixa- tion of bacteria in the soil by legumes (bean and pea plants). Lightning produces nitrogen, as do industrial waste gases and the decomposition products of organic material (i.e., organic proteins and amino acids in plants and animals contain nitrogen). In time, these sources replace the nitrogen in the atmosphere to complete the cycle.
Ammonia (NH 3 ) is the first binary molecule discovered in outer space of our galaxy, the Milky Way. It may also be the main compound that forms the rings of the planet Saturn.
History In 1772 a student, Daniel Rutherford (1749–1819), at the suggestion of his mentor, Joseph
Black (1728–1799), conducted an experiment in which he burned a candle in a closed container of air. (Joseph Black was famous for his concept of “fixed air,” which was an important step in understanding gases in chemistry.) Chemists of the time already knew that air contained at least two gases, one that supported life and one that did not. Rutherford started his experiment by placing a mouse in a sealed glass jar until it suffocated and had reduced the volume of air by 1/16. He repeated this with a candle and noticed that there was still a large quantity of gas in the container after the burning candle had consumed the oxygen. There was no carbon dioxide because he had chemically removed that gas. Rutherford experimented further and found that this leftover gas could not support combustion or life, so he called it noxious air. He never did identify nitrogen but came up with some several suggestions. That led to the identification of a new element, which was later named “nitrogen” for the Greek word meaning “niter producer,”
after the compound, potassium nitrate (saltpeter KNO 3 ), which contained nitrogen. At the same time Rutherford conducted his experiments, three other chemists—Priestley, Cavendish, and Scheele—were also investigating “fixed air” gases, including nitrogen. However, Rutherford was given credit for discovering nitrogen.
CommonUses Nitrogen has many uses. It is the second most commonly produced chemical in the United
States. Its chemical and physical properties, along with the five electrons in its outer shell, make it a versatile element that can react as a metal or nonmetal to produce numerous com-
pounds. Some of its uses are based on its inertness as a gas (N 2 ) and its ability to be liquefied to provide very low temperatures. When recovered as a gas in the atmosphere, it is used to produce anhydrous ammonia (NH 3 ), which is the fifth most commonly produced chemical in the United States. It is also used as the basis for making many nitrogen compounds. At one time it was believed to be impossible to combine hydrogen with nitrogen to form ammonia, a natural product of animal waste that was used as a fertilizer and textile bleach, among other things. In 1905 the German chemist Fritz Haber (1868–1934) demonstrated that it was possible to combine hydrogen with nitrogen in a process that directly produced ammonia. The Haber process requires
211 high temperatures (500°C) and very high pressure. It is the main source of ammonia today.
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Millions of tons of ammonia are produced every year worldwide. Haber received the 1918 Nobel Prize in Chemistry for his work.
Most of us are familiar with the liquid form of ammonia known as ammonium hydroxide (NH 4 OH), a colorless liquid that, with its strong odor, is irritating to the eyes and potentially harmful to the moist mouth and nose, throat, and lungs if its vapors are breathed. Weak solu- tions of NH 4 OH are ingredients in household cleaning ammonia. Concentrated ammonium hydroxide has many industrial uses, including the manufacture of rayon, fertilizers, refriger- ants, rubber, pharmaceuticals, soaps lubricants, inks, explosives, and household cleaners.
Many fertilizers are based on ammonia compounds. Modern agriculture requires more nitrogen in soils than is normally replaced by the nitrogen cycle, lightning, decaying plants and animals, and other natural means
Nitric acid (HNO 3 ) is an important commercial chemical and was manufactured com- mercially to produce fertilizers and explosives as well as plastics and many other products. In 1902 a German chemist, Wilhelm Ostwald (1853–1932), developed a process wherein at high temperatures he used platinum catalysts to convert ammonia into nitric acid. When nitric acid is reacted with glycerol, the result is nitroglycerine—an unstable explosive unless dissolved in inert material, such as clay. It can then be stabilized as dynamite.
Nitrates are formed when nitric acid is neutralized by a base such as sodium hydroxide (NaOH). These nitrates change form to become nitrites that are used as preservatives, particu- larly in canned goods and to keep meat looking fresh.
Sodium azide (NaN 3 ) is an explosive salt of nitrogen that produces large quantities of gas upon its explosion. This quality has made it ideal as the chemical contained in automobile air safety bags. When triggered it explodes immediately, producing the expanding gases that fill the bag.
The radioisotope nitrogen-13 has a relatively short half-life of about 10 minutes that pro- duces a positron as it decays. This makes N-13 useful in PET (Positron Emission Tomography) scan technology, in which it is injected into the patient. The positive electrons (positrons) interact with the patient’s negative electrons to produce an image similar to an X-ray.
Oil companies force nitrogen under great pressure into depleted oil wells to force residual crude oil to the surface. By far, nitrogen compounds are of the utmost importance to the diets and welfare of both plants and animals. Nitrogen is essential to living things. Plants require nitrogen, and so do animals, which get their nitrogen from eating plants and other animals.
ExamplesofCompounds Nitrogen has numerous positive and negative oxidation states. For example, it can form six
different compounds with oxygen using the oxidation states of +1 through +6. Ammonium nitrate (NH 4 NO 3 , also known as “Norway saltpeter”) is mainly used as a fer- tilizer. It is also known as the chemical that was mixed with diesel fuel to create the explosion that demolished the Murrah Federal Building in Oklahoma City in 1995.
Nitrogen oxides (NO x ; the “x” represents the proportion of nitrogen to oxygen atoms in the various oxidation states of nitrogen) have many uses, including the production of nitric acid. Most of the oxides of nitrogen, especially NO 2 , are toxic if inhaled. On the other hand, nitrous oxide (N 2 O), although explosive in air, is known as “laughing gas” and is used as an anesthetic in dentistry and surgery.
212 | The History and Use of Our Earth’s Chemical Elements Nitroglycerine, as mentioned, is an unstable explosive. It is also used as a vasodilator to
reduce high blood pressure and angina pectoris by dilating the blood vessels of heart patients whose hearts are not receiving an adequate blood supply.
Recently a team of chemists discovered a new allotrope of nitrogen that consisted of five nitrogen atoms in the form of a V. This surprised them because it was thought that any struc- ture with three or more nitrogen atoms would be unstable—and indeed, the V form turned out to be so unstable that it created an extreme explosion that destroyed the equipment being used to analyze it.
Hazards Nitrogen is nontoxic, but it is an asphyxiate gas that cannot, by itself, support oxidation
(combustion) or support life. If you breathe pure nitrogen for any period of time, you will die— not because the nitrogen gas is a poison, but because your body will be deprived of oxygen.
Nitrogen oxides are formed under certain conditions when nitrogen combines with oxy- gen, thus contributing to pollution. One source is from the internal combustion engine that produces NO similar to lightning. Once released, it combines with more oxygen to form NO 2 , which is a very reactive polluting gas. Nitrogen dioxide NO 2 is the main cause of “brown” smog over some cities and is harmful to plants, animals, and humans. To make matter worse, if there is adequate sunlight at the time of the smog, the ultraviolet light of the sun will break
down the N and O of the NO 2 to form free radicals of oxygen that are reactive, forming ozone (O 3 ), which is itself a strong oxidizing agent that adds to pollution. Several of the oxygen, hydrogen, and halogen compounds of nitrogen are toxic when inhaled. A common error made in using household cleaners is to mix or use together ammonia cleaning fluids (containing nitrogen) and Clorox-type cleaning fluids (containing chlorine). The combined fumes can be deadly in any confined area. NEVER mix Clorox with ammonia- type cleaning fluids.