Nitrogen (N), phosphorus (P), arsenic (As), antimony (Sb), and bismuth (Bi) form Group 5A of the periodic table. Nitrogen, a gas, is an essential component of all living matter, phosphorus is a highly reactive nonmetal, arsenic and antimony are toxic metalloids, and bismuth is a true metal. Nitrogen and phosphorus are similar in some respects—both have nonmetal-lic characteristics and are essential ingredients of living tissue. Otherwise, phosphorus is very different from nitrogen. Phosphorus is solid at room temperatures and exists in three different forms, or allotropes—one of which is poisonous, as are several of its compounds.
The other elements of Group 5A are unlike nitrogen but similar to phosphorus. Arsenic and antimony exist as different allotropes, some compounds of which are toxic. Arsenic, antimony, and bismuth are increasingly metallic. Because of the way their outer (valence; electrons are arranged, each can form up to five chemical bonds with other elements or groups of elements.
This colorless, odorless gas makes up about 78 per cent (by volume) of the earth’s atmosphere, as well as about 16 per cent of animal and vegetable proteins—the other 84 per cent consists of carbon, hydrogen, oxygen, and sulfur. Most naturally occurring nitrogen compounds dissolve in water. The only large deposits, consisting of sodium nitrate (Chile saltpeter—NaN03), are found in Chile and in other arid climates. These deposits have been mined for fertilizer and are also used in making explosives.
Nitrogen was first recognized by the French chemist Antoine Lavoisier (1734-1794). Fie named it azote, meaning “without life,” because of its inability to support life. Flowever, the element’s discovery in 1772 is credited to the Scottish physician Daniel Rutherford (1749-1819). Its present name, coined in 1790, means “niter forming,” because of nitrogen’s presence in niter (potassium nitrate). Its atomic number is 7, and its atomic mass is 14.0067. Its melting point is -209.9° C, and its boiling point is -195.8° C.
The Haber-Bosch process, developed in the early twentieth century, is still an important method of making ammonia. Nitric acid and, in turn, explosives and fertilizers, are manufactured from ammonia. The key reaction is the combination of one molecule of nitrogen (from air) and three molecules of hydrogen (now obtained from natural gas). The combination gives two molecules of ammonia. Significant amounts of ammonia can be obtained only by using high pressures (between 200 and 250 atmospheres), high temperatures (about 500° C), and a catalyst to help speed the process.
Nitrogen itself can be isolated by cooling air until it liquefies and then cooling it still further until the boiling point of nitrogen is reached, at which point the gas can be collected.
Most nitrogen is used in making fertilizer from ammonia (NFH3) or from nitric acid (F1N03). The gas itself is used in the chemical, electrical, and metals industries. It is easy to prepare in liquid form for use as a refrigerant.
Although nitrogen is an essential ingredient of living tissue, only a few organisms can take it directly from the air to make the proteins needed for growth and tissue maintenance. Most plants take nitrogen from the soil in the form of nitrates, nitrites, and ammonium salts. Animals and humans get most of their nitrogen from eating plants or other animals. The need for atmospheric nitrogen thus results in an interdependence of soil, plants, and animals—a relationship known as the nitrogen cycle.
Complex processes cycle nitrogen out of the air and into soil and water. From there it is absorbed by living tissue in plants, animals, and humans. When tissue dies, bacteria release the nitrogen back into the soil and air. Nitrogen enters the soil in rain water as dilute nitric and nitrous acid (different combinations of N, FI, and O). These acids form in the atmosphere when lightning causes nitrogen and oxygen to react and combine with rain. When they fall to earth, they are neutralized by reacting with chemicals in the soil bases to form nitrates and nitrites, which plants use to make proteins. Also, certain bacteria and algae take nitrogen directly from the air and convert it to ammonia (NF13), which plants then use. As already mentioned, bacteria also play a role in releasing nitrogen, as ammonia, from decaying organic matter. This builds up in the soil as ammonium salts, which other bacteria convert into the nitrates and nitrites that plants need.
Bones and teeth contain phosphorus—in fact, every living cell, both plant and animal, contains phosphorus compounds of one kind or another. In pure form, this soft, nonmetallic element exists as three allotropes: white, red, and black.
Phosphorus was first isolated from urine by the German alchemist Flennig Brand in 1669. In 1680, it was discovered independently by the English chemist Robert Boyle. The name comes from the Greek phosphoros, meaning light-bearer. The white form glows in the dark. Its atomic number is 15, and its atomic mass is 30.9738. In its white form, the melting point is 44.1° C, and its boiling point is 280° C. In its red form, its melting point is 600° C (under pressure).
The white allotrope is a soft, waxy solid that is unstable in air, giving off a faint green glow called phosphorescence as it turns first yellow and then red. It burns easily, emitting toxic fumes, so it is kept underwater, with which it does not react. The red allotrope, which is much more stable, is used to make matches.
It is prepared by heating white phosphorus in the absence of air for a period of several hours. Black phosphorus is prepared by heating white phosphorus for eight days with mercury as a catalyst
Plants need phosphorus compounds called phosphates in order to grow. They absorb these compounds from the soil. Unlike nitrogen, phosphorus is not returned to the soil in a natural cycle, so phosphate fertilizers must be continuously added to the soil. Phosphorus is also used in animal feeds, steel, china, pesticides, safety matches, and—as phosphoric acid or its salts—in a variety of drugs, soft drinks, flavoring agents, water softeners, and detergents. The use of phosphate detergents is discouraged because it causes environmental problems by triggering growth of algae in lakes and bays into which wastewater is discharged.
Arsenic, antimony, and bismuth
These three members of Group 5A exist in bright metallic forms that are stable in air.
They occur in pure form in nature or combined with other elements, usually sulfur. In industrial use, they improve the strength and hardness of alloys.
Arsenic in its most stable form is a shiny gray metalloid. Less common are the two other allotropes—yellow and black arsenic. All three are highly toxic, and the compounds are used in making fungicide, weedkiller, rat poison, insecticide, and drugs to treat certain infections.
Arsenic was identified by Albert Magnus about 1250. The Latin word arsenicum means yellow orpiment (a pigment containing arsenic and sulfur). Its atomic number is 33, and its atomic mass is 74.9216. It sublimes [passes directly into a vapor without melting) at 613° C
Antimony, a bluish-white brittle metal, is alloyed with lead to make batteries, electric cables, and metal for printer’s type.
Antimony occurs mainly as its sulfide mineral stibnite (Latin stibnum), from which it derives its chemical symbol Sb. It was known to the Greeks and Romans. Its atomic number is 51, and its atomic mass is 121.75. Its melting point is 630.74° C, and its boiling point is 1586.85° C
Bismuth, a brittle, silvery-pink metal, goes into alloys that melt at low temperatures (as low as 47° C). It is used in electrical fuses, safety plugs in steam boilers, and automatic sprinkler systems. Because it is nontoxic, its compounds are found in drugs, medicines, and cosmetics.
Bismuth was known in the ancient world. It was isolated by Caspar Newmann (1683-1737). Its name may derive from the Old German vissmuth (white matter). Its atomic number is 83, and its atomic mass is 208.98. Its melting point is 271.3° C, and its boilinq point is about 1560+5° C.