Phosphorus

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Date: 2007
Publisher: Gale, a Cengage Company
Document Type: Topic overview
Length: 1,015 words
Content Level: (Level 4)
Lexile Measure: 1250L

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Phosphorus is a waxy, nonmetal element represented by the atomic symbol P. It has an atomic number of 15 and an atomic mass of 30.9738. Phosphorus is highly reactive, so it does not occur freely in nature. It is typically found in combination with minerals such as apatite, and it is found in living cells.

Phosphorous exists in three main allotropic forms (forms with different crystalline structure), including white, red, and black. White phosphorus melts at 111.47°F (44.15°C); it is also phosphorescent, which means it glows in the dark when exposed to air, as a result of phosphorus vapor slowly combining with oxygen in the air. White phosphorus is one of the most dangerous substances known. It is so reactive that it must be stored underwater; otherwise it will ignite spontaneously. Skin contact with white phosphorus can cause serious burns. When white phosphorus is heated in its own vapor to 482°F (250°C), it becomes red phosphorus. This form is much less reactive than the white allotrope. Black phosphorus, which is not especially reactive, has a texture similar to graphite.

Phosphorus was the first element to truly be discovered; it was genuinely new--neither scientists nor ancient cultures recognized any form of it. In the mid-seventeenth century, German alchemist Hennig Brand was searching for the philosopher's stone, a hypothetical substance able to change common metals into gold. Brand became convinced that the human body contained such an agent, and like many other alchemists he experimented with urine. Around 1669, he evaporated water from urine and burned the concentrated residue, along with some sand. Instead of the philosopher's stone, Brand produced a white, waxy substance that glowed mysteriously in the dark and ignited spontaneously when exposed to air. He named the substance phosphorus, after the Greek word for "light-bearer."

Brand was initially very protective of his secret recipe for making phosphorus. In 1680, Irish chemist Robert Boyle (1627-1691) independently prepared phosphorus from urine. Unlike Brand, Boyle believed that scientific work should be reported so that other investigators could duplicate an experiment and confirm the results. Boyle published his results, and when other chemists succeeded in isolating phosphorus, a fierce controversy erupted over who had done it first. Boyle's approach and his results signified a step away from superstitions and toward modern chemistry.

Boyle went on to define the chemical and physical properties of phosphorus and phosphoric acid, and other researchers soon publicized improved methods for preparing phosphorus. In 1740, German chemist Andreas Marggraf (1709-1782) burned the element to create white "flowers"--actually oxides of phosphorus--and noted that the product showed an increase in weight. This observation later proved critical to the research on combustion and oxygen that Antoine Laurent Lavoisier (1743-1794) carried out.

In the late eighteenth century, Swedish chemist Karl Wilhelm Scheele (1742-1786) developed an economical method of obtaining phosphorus from animal bones. People knew that bones were useful for fertilizing crops, but in the early nineteenth century scientists learned that the phosphorus in bones was a far more effective fertilizer when treated with sulfuric acid to make it soluble in water. The product, called superphosphate, was first manufactured commercially by Irish physician James Murray (1788-1871), who observed that it could also be produced from rocks that contained phosphate. In 1842, British agricultural scientist John Lawes (1814-1900) patented a process for manufacturing superphosphate from mineral phosphate. Toward the end of the nineteenth century, Lawes's factory was making approximately 40,000 tons of the fertilizer per year. Superphosphate is still used as a fertilizer, along with a more potent version that supplies more than twice as much phosphorus.

Plants use phosphorus in photosynthesis, but they also need it to establish roots, mature, and ripen. Phosphate's importance made it a popular fertilizer, but when large amounts of phosphorus fertilizers drain into rivers and lakes, the nutrient stimulates the growth of algae that consume the available oxygen. This can kill fish and has negative effects on other plants and animals in the environment. For this reason, various agencies are trying to reduce the impact of environmental phosphates.

All animal cells need phosphorus as well. It is a crucial part of many biological compounds, such as ATP (adenosine triphosphate), which is involved in practically every metabolic reaction. In the human body, most phosphorus is found in bones and teeth. Even though it is very important to life, phosphorus is much less common than other biologically important elements like oxygen and nitrogen. Some scientists believe that iron meteorites may have provided a boost of phosphorus long ago that set the development of life in motion.

One of the most practical uses for phosphorus compounds is in the production of strike-anywhere safety matches. Phosphorus-tipped matches first came into widespread use in the nineteenth century, but their manufacture involved white phosphorus. Match-factory workers who were regularly exposed to this toxic element developed a condition called "phossy jaw" from breathing phosphorus fumes. People with this disease have high concentrations of phosphorus deposited in the jaw bone, causing their gums to swell, their teeth to die, and the jaw itself to disintegrate. It also causes brain damage, anemia, and loss of appetite. Patients' lives could sometimes be saved by removing the affected bone, but this entailed the risk of being disfigured and unable to eat solid food.

Alice Hamilton (1869-1970), an American physician and pioneer in occupational disease, publicized the effects of phossy jaw and helped eliminate the use of white phosphorus in matches. European chemists solved this problem by developing an alternative, nonpoisonous form of phosphorus. Red phosphorus, which is less reactive, was substituted for the harmful white version.

Despite its dangers, white phosphorus is used to make common products such as detergents, water softeners, animal foods, insecticides, steel, and plastics. It is also part of military nerve gases and explosives such as grenades and mortar shells. In addition to its role in the manufacture of safety matches, red phosphorus is used to make various phosphorus compounds such as phosphoric acid and phosphorus trichloride. These compounds go on to play other important roles; for example, phosphoric acid is used in making baking powder and carbonated beverages.

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Gale Document Number: GALE|CV1648500457