Phosphorus
Phosphorus is a waxy, non-metal element that is represented by the atomic symbol, P. It has an atomic number of 15 and an atomic weight of 30.9738. It does not occur freely in nature, but is typically found in combination with minerals such as apatite.
Three main allotropic forms of phosphorus are found including white, red and black. White phosphorus has melting point of 111.47°F (44.15°C). White phosphorus exhibits the phenomenon of phosphorescence, which means it glows in the dark when exposed to air. This results from the slow combination of phosphorus vapor 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, or it will ignite spontaneously; it can also cause serious burns on contact with flesh. When white phosphorus is heated in its own vapor up to 482°F (250°C), it changes to red phosphorus. It is much less reactive than the white allotrope. Black phosphorus has a texture similar to graphite. It is also less reactive.
Phosphorus was the first element to be truly discovered in the sense that it had never been known in any form before. In the mid-1600s, a German chemist named Hennig Brand (died c. 1692) began searching for the philosopher's stone, a hypothetical substance that alchemists thought would transform common metals into pure gold. Brand became convinced that the human body contained such an agent, so 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 mysteriously glowed in the dark and ignited spontaneously when exposed to air. He named the substance phosphorus after the Greek word for light-bearer.
At first Brand guarded his secret recipe for making phosphorus. But in 1680, the Irish chemist Robert Boyle independently prepared phosphorus from urine. Unlike Brand, Boyle believed that scientific experiments should be reported so that others could repeat and confirm them. When other chemists succeeded in isolating phosphorus, a fierce controversy erupted over who had done it first. At any rate, the discovery signified a step away from ancient superstitions and toward modern chemistry.
Boyle went on to define the chemical and physical properties of phosphorus and phosphoric acid. Methods of preparing phosphorus were soon publicized and improved. In 1740, Andreas Marggraf (1709-1782), a German chemist, combusted the element to produce white flowers (oxides of phosphorus) and noted that the product showed an increase in weight. This observation later proved critical in Antoine-Laurent Lavoisier's definition of oxygen.
In the 1770s, Swedish chemist Carl Wilhelm Scheele developed a more economical method of obtaining phosphorus from animal bones. Although bones were recognized to be of value in fertilizing crops, scientists learned in the early 1800s that the phosphorus found in bones was far more effective as a 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 mineral phosphate rock. In 1842, Sir John Lawes (1814-1900), a British agricultural scientist, patented a process for manufacturing superphosphate from the mineral. By the 1870s, Lawes's factory was making some 40,000 tons per year of the fertilizer. Superphosphate is still in use today, along with a more potent version that supplies more than twice as much phosphorus nutrient.
Besides using phosphorus in photosynthesis, plants require it to establish roots, to mature and ripen. But when phosphorus fertilizers drain into rivers and lakes, the nutrient stimulates the growth of algae, which consume the available oxygen. This can have a detrimental effect on fish and other animals in the environment. For this reason, efforts have been made to reduce the impact of environmental phosphates.
Despite its toxic nature, phosphorus is required by all plant and animal cells. The element makes up an important part of many biological compounds such as ATP ( adenosine triphosphate) which is involved in practically every reaction in metabolism and photosynthesis. In the human body, most phosphorus exists in the bones and teeth.
One of the most practical uses for phosphorus compounds is in the production ofstrike-anywhere safety match. Phosphorus-tipped matches first came into widespread use during the 1800s, but the element's highly poisonous nature gravely affected the health of match factory workers, who developed a condition called phossy jaw from breathing phosphorus fumes. The disease kills the roots of the teeth and deteriorates the jawbone, creating intense pain; it also causes anemia and loss of appetite. Some patients were horribly disfigured by jaw surgery and had to live on liquid food for the rest of their lives. To avoid these conditions, European chemists developed an alternative, nonpoisonous form of phosphorus. This red phosphorus, which is less reactive, was substituted for the harmful white phosphorus in match production and was eventually incorporated on the matchbox striker instead of the match tip. In the United States, Alice Hamilton (1869-1970), a physician and pioneer in occupational disease, publicized the effects of phossy jaw disease and helped eliminate the use of white phosphorus in matches.
White phosphorus is used to make such common products as detergents, water softeners, animal food, insecticides, steel, and plastics. It is also a component of military nerve gases and explosives such as grenades and mortar shells. In addition to producing safety matches, red phosphorus is used in the manufacture of various phosphorus compounds such as phosphoric acid and phosphorus trichloride.
This is the complete article, containing 902 words
(approx. 3 pages at 300 words per page).
