| Discovery date | 1669 |
| Discovered by | Hennig Brandt |
| Origin of the name | The name is derived from the Greek 'phosphoros', meaning bringer of light. |
| Allotropes | White P, Red P, Black P, P2 |
P
Phosphorus
15
30.974
| Group | 15 | Melting point | 44.15°C, 111.47°F, 317.3 K |
| Period | 3 | Boiling point | 280.5°C, 536.9°F, 553.7 K |
| Block | p | Density (g cm−3) | 1.823 (white) |
| Atomic number | 15 | Relative atomic mass | 30.974 |
| State at 20°C | Solid | Key isotopes | 31P |
| Electron configuration | [Ne] 3s23p3 | CAS number | 7723-14-0 |
| ChemSpider ID | 4575369 | ChemSpider is a free chemical structure database | |
Image explanation
The image is of a ball-and-stick model of white phosphorus. It has a tetrahedral shape and has the formula P4.
Appearance
The two main forms of phosphorus are white phosphorus and red phosphorus. White phosphorus is a poisonous waxy solid and contact with skin can cause severe burns. It glows in the dark and is spontaneously flammable when exposed to air. Red phosphorus is an amorphous non-toxic solid.
Uses
White phosphorus is used in flares and incendiary devices. Red phosphorus is in the material stuck on the side of matchboxes, used to strike safety matches against to light them.
By far the largest use of phosphorus compounds is for fertilisers. Ammonium phosphate is made from phosphate ores. The ores are first converted into phosphoric acids before being made into ammonium phosphate.
Phosphorus is also important in the production of steel. Phosphates are ingredients in some detergents, but are beginning to be phased out in some countries. This is because they can lead to high phosphate levels in natural water supplies causing unwanted algae to grow. Phosphates are also used in the production of special glasses and fine chinaware.
Biological role
Phosphorus is essential to all living things. It forms the sugar-phosphate backbone of DNA and RNA. It is important for energy transfer in cells as part of ATP (adenosine triphosphate), and is found in many other biologically important molecules. We take in about 1 gram of phosphate a day, and store about 750 grams in our bodies, since our bones and teeth are mainly calcium phosphate. Over-use of phosphates from fertilisers and detergents can cause them to pollute rivers and lakes causing algae to grow rapidly. The algae block out light stopping further photosynthesis. Oxygen dissolved in the water soon gets used up and the lake dies.
Natural abundance
Phosphorus is not found uncombined in nature, but is widely found in compounds in minerals. An important source is phosphate rock, which contains the apatite minerals and is found in large quantities in the USA and elsewhere. There are fears that ‘peak phosphorus’ will occur around 2050, after which our sources will dwindle.
White phosphorus is manufactured industrially by heating phosphate rock in the presence of carbon and silica in a furnace. This produces phosphorus as a vapour, which is then collected under water. Red phosphorus is made by gently heating white phosphorus to about 250°C in the absence of air.
Phosphorus was first made by Hennig Brandt at Hamburg in 1669 when he evaporated urine and heated the residue until it was red hot, whereupon phosphorus vapour distilled which he collected by condensing it in water. Brandt kept his discovery secret, thinking he had discovered the Philosopher’s Stone that could turn base metals into gold. When he ran out of money, he sold phosphorus to Daniel Kraft who exhibited it around Europe including London where Robert Boyle was fascinated by it. He discovered how it was produced and investigated it systematically. (His assistant Ambrose Godfrey set up his own business making and selling phosphorus and became rich.)
When it was realised that bone was calcium phosphate, and could be used to make phosphorus, and it became more widely available. Demand from match manufacturers in the 1800s ensured a ready market.
| Atomic radius, non-bonded (Å) | 1.80 | Covalent radius (Å) | 1.09 |
| Electron affinity (kJ mol−1) | 72.037 |
Electronegativity (Pauling scale) |
2.19 |
|
Ionisation energies (kJ mol−1) |
1st
1011.812
2nd
1907.467
3rd
2914.118
4th
4963.582
5th
6273.969
6th
21267.395
7th
25430.64
8th
29871.9
|
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|
| Common oxidation states | 5, 3, -3 | ||||
| Isotopes | Isotope | Atomic mass | Natural abundance (%) | Half life | Mode of decay |
| 31P | 30.974 | 100 | - | - | |
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Specific heat capacity (J kg−1 K−1) |
769 | Young's modulus (GPa) | Unknown | |||||||||||
| Shear modulus (GPa) | Unknown | Bulk modulus (GPa) | 10.9 (red); 4.9 (white) | |||||||||||
| Vapour pressure | ||||||||||||||
| Temperature (K) |
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| Pressure (Pa) |
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| Listen to Phosphorus Podcast |
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Transcript :
Chemistry in its element: phosphorus (Promo) You're listening to Chemistry in its element brought to you by Chemistry World, the magazine of the Royal Society of Chemistry. (End promo) Chris Smith Hello - this week fertilisers, fire bombs, phossy jaw and food additives. What's the connection? Here's Nina Notman. Nina Notman Phosphorus is a non-metal that sits just below nitrogen in group 15 of the periodic table. This element exists in several forms, of which white and red are the best known. White phosphorus is definitely the more exciting of the two. As it glows in the dark, is dangerously flammable in the air above 30 degrees, and is a deadly poison. Red phosphorus however has none of these fascinating properties. So where did it all begin? Phosphorus was first made by Hennig Brandt in Hamburg in Germany in 1669. When he evaporated urine and heated the residue until it was red hot. Glowing phosphorus vapour came off and he condensed it under water. And for more than 100 years most phosphorus was made this way. This was until people realised that bone was a great source of phosphorus. Bone can be dissolved in sulfuric acid to form phosphoric acid, which is then heated with charcoal to form white phosphorus. White phosphorus has found a range of rather nasty applications in warfare. It was used in the 20th century in tracer bullets, fire bombs, and smoke grenades. The scattering of phosphorus fire bombs over cities in World War II caused widespread death and destruction. In July 1943, Hamburg was subject to several air raids in which 25,000 phosphorus bombs were dropped over vast areas of the city. This is rather ironically considering where phosphorus was first made. Another group of warfare agents based on phosphorus are nerve gases such as sarin. Sarin is a fluorinated phosphonate that was used by Iraq against Iran in the early to mid-1980s. And was also released in a Tokyo subway in 1995, killing 12 people and harming nearly a thousand others. White phosphorus has also found a wide range of other uses. One of these was in phosphorus matches that were first sold in Stockton-on-Tees in the UK in 1827. This created a whole new industry of cheap lights - but at a terrible cost. Breathing in phosphorus vapour led to the industrial disease phossy jaw, which slowly ate away the jaw bone. This condition particularly afflicted the girls who made phosphorus matches. So these were eventually banned in the early 1900s and were replaced by modern matches which use either phosphorus sulfide or red phosphorus. As well as in matches, today phosphorus has found other uses in lighting. Magnesium phosphide is the basis of self-igniting warning flares used at sea. When it reacts with water it forms the spontaneously flammable gas, diphosphine which triggers the lighting of the flare. Super pure phosphorus is also used to make light emitting diodes. These LEDs contain metal phosphides such as those of gallium and indium. In the natural world the elemental form of phosphorus is never encountered. It is only seen as phosphate, and phosphate is essential to life for numerous reasons. It is part of DNA, and also constitutes a huge proportion of teeth enamel and bones in the form of calcium phosphate. Organophosphates are also important, such as the energy molecule ATP and the phospholipids of cell membranes. A normal diet provides our bodies with the phosphate it needs. With tuna, chicken, eggs and cheese having lots. And even cola provide us with some, in the form of phosphoric acid. Today most of our phosphorus comes from phosphate rock that is mined around the world, and then converted to phosphoric acid. Fifty million tonnes are made every year and it has multiple uses. It is used to make fertilisers, animal feeds, rust removers, corrosion preventers, and even dishwasher tablets. Some phosphate rock is also heated with coke and sand in an electric furnace to form white phosphorus which is then converted to phosphorus trichloride and phosphorous acid. And it is from these that flame retardants, insecticides, and weed-killers are made. A little is also turned into phosphorus sulfides which are used as oil additives to reduce engine wear. Phosphate is also environmentally important. It naturally moves from soil, to rivers, to oceans, to bottom sediment. Here it accumulates until it is moved by geological uplift to dry land so the circle can start again. During its journey, phosphate passes through many plants, microbes, and animals of various eco-systems. Too much phosphate however can be damaging to natural waters because it encourages unwanted species like algae to flourish. These then crowd out other forms of desired life. There is now a legal requirement to remove phosphate from wastewaters in many parts of the world, and in the future this could be recycled as a sustainable resource so that one day the phosphate we flush down sinks and toilets might reappear in our homes in other guises such as in dishwasher tablets and maybe even in our food and colas. Chris Smith Nina Notman with the tale of Phosphorus, the element extracted from the golden stream, otherwise known as urine. Next time Andrea Sella will be joining us with the explosive story of element number 53. Andrea Sella In 1811 a young French chemist, Bernard Courtois, working in Paris stumbled across a new element. His family's firm produced the saltpetre needed to make gunpowder for Napoleon's wars. They used wood ash in their process and wartime shortages of wood forced them instead to burn seaweed. Adding concentrated sulphuric acid to the ash, Courtois, obtained an astonishing purple vapour that crystallized onto the sides of the container. Astonished by this discovery he bottled up the greyish crystals and sent them to one of the foremost chemists of his day Joseph Guy-Lussac who confirmed that this was a new element and named it iode - iodine - after the Greek word for purple. Chris Smith And you can hear more about how Iodine exploded onto the world's stage on next week's Chemistry in its Element, I hope you can join us. I'm Chris Smith, thank you for listening and goodbye. (Promo) Chemistry in its element is brought to you by the Royal Society of Chemistry and produced by thenakedscientists.com. There's more information and other episodes of Chemistry in its element on our website at chemistryworld.org/elements. (End promo)
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Visual Elements images and videos
© Murray Robertson 1998-2017.
W. M. Haynes, ed., CRC Handbook of Chemistry and Physics, CRC Press/Taylor and Francis, Boca Raton, FL, 95th Edition, Internet Version 2015, accessed December 2014.
Tables of Physical & Chemical Constants, Kaye & Laby Online, 16th edition, 1995. Version 1.0 (2005), accessed December 2014.
J. S. Coursey, D. J. Schwab, J. J. Tsai, and R. A. Dragoset, Atomic Weights and Isotopic Compositions (version 4.1), 2015, National Institute of Standards and Technology, Gaithersburg, MD, accessed November 2016.
T. L. Cottrell, The Strengths of Chemical Bonds, Butterworth, London, 1954.
John Emsley, Nature’s Building Blocks: An A-Z Guide to the Elements, Oxford University Press, New York, 2nd Edition, 2011.
Thomas Jefferson National Accelerator Facility - Office of Science Education, It’s Elemental - The Periodic Table of Elements, accessed December 2014.
Periodic Table of Videos, accessed December 2014.
Derived in part from material provided by the British Geological Survey © NERC.
Elements 1-112, 114, 116 and 117 © John Emsley 2012. Elements 113, 115, 117 and 118 © Royal Society of Chemistry 2017.
Produced by The Naked Scientists.
Created by video journalist Brady Haran working with chemists at The University of Nottingham.
© Murray Robertson 1998-2017.
Data
W. M. Haynes, ed., CRC Handbook of Chemistry and Physics, CRC Press/Taylor and Francis, Boca Raton, FL, 95th Edition, Internet Version 2015, accessed December 2014.
Tables of Physical & Chemical Constants, Kaye & Laby Online, 16th edition, 1995. Version 1.0 (2005), accessed December 2014.
J. S. Coursey, D. J. Schwab, J. J. Tsai, and R. A. Dragoset, Atomic Weights and Isotopic Compositions (version 4.1), 2015, National Institute of Standards and Technology, Gaithersburg, MD, accessed November 2016.
T. L. Cottrell, The Strengths of Chemical Bonds, Butterworth, London, 1954.
Uses and properties
John Emsley, Nature’s Building Blocks: An A-Z Guide to the Elements, Oxford University Press, New York, 2nd Edition, 2011.
Thomas Jefferson National Accelerator Facility - Office of Science Education, It’s Elemental - The Periodic Table of Elements, accessed December 2014.
Periodic Table of Videos, accessed December 2014.
Supply risk data
Derived in part from material provided by the British Geological Survey © NERC.
History text
Elements 1-112, 114, 116 and 117 © John Emsley 2012. Elements 113, 115, 117 and 118 © Royal Society of Chemistry 2017.
Podcasts
Produced by The Naked Scientists.
Periodic Table of Videos
Created by video journalist Brady Haran working with chemists at The University of Nottingham.
