| Discovery date | 1824 |
| Discovered by | Jöns Jacob Berzelius |
| Origin of the name | The name is derived from the Latin 'silex' or 'silicis', meaning flint. |
| Allotropes | amorphous Si, crystalline Si |
Si
Silicon
14
28.085
| Group | 14 | Melting point | 1414°C, 2577°F, 1687 K |
| Period | 3 | Boiling point | 3265°C, 5909°F, 3538 K |
| Block | p | Density (g cm−3) | 2.3296 |
| Atomic number | 14 | Relative atomic mass | 28.085 |
| State at 20°C | Solid | Key isotopes | 28Si, 30Si |
| Electron configuration | [Ne] 3s23p2 | CAS number | 7440-21-3 |
| ChemSpider ID | 4574465 | ChemSpider is a free chemical structure database | |
Image explanation
The image is based on a diatom. Diatoms are photosynthesising algae. They are unique in that their cell walls are made of silica (hydrated silicon dioxide).
Appearance
The element, when ultrapure, is a solid with a blue-grey metallic sheen.
Uses
Silicon is one of the most useful elements to mankind. Most is used to make alloys including aluminium-silicon and ferro-silicon (iron-silicon). These are used to make dynamo and transformer plates, engine blocks, cylinder heads and machine tools and to deoxidise steel.
Silicon is also used to make silicones. These are silicon-oxygen polymers with methyl groups attached. Silicone oil is a lubricant and is added to some cosmetics and hair conditioners. Silicone rubber is used as a waterproof sealant in bathrooms and around windows, pipes and roofs.
The element silicon is used extensively as a semiconductor in solid-state devices in the computer and microelectronics industries. For this, hyperpure silicon is needed. The silicon is selectively doped with tiny amounts of boron, gallium, phosphorus or arsenic to control its electrical properties.
Granite and most other rocks are complex silicates, and these are used for civil engineering projects. Sand (silicon dioxide or silica) and clay (aluminium silicate) are used to make concrete and cement. Sand is also the principal ingredient of glass, which has thousands of uses. Silicon, as silicate, is present in pottery, enamels and high-temperature ceramics.
Silicon carbides are important abrasives and are also used in lasers.
Biological role
Silicon is essential to plant life but its use in animal cells is uncertain. Phytoliths are tiny particles of silica that form within some plants. Since these particles do not rot they remain in fossils and provide us with useful evolutionary evidence.
Silicon is non-toxic but some silicates, such as asbestos, are carcinogenic. Workers, such as miners and stonecutters, who are exposed to siliceous dust can develop a serious lung disease called silicosis.
Natural abundance
Silicon makes up 27.7% of the Earth’s crust by mass and is the second most abundant element (oxygen is the first). It does not occur uncombined in nature but occurs chiefly as the oxide (silica) and as silicates. The oxide includes sand, quartz, rock crystal, amethyst, agate, flint and opal. The silicate form includes asbestos, granite, hornblende, feldspar, clay and mica.
Elemental silicon is produced commercially by reducing sand with carbon in an electric furnace. High-purity silicon, for the electronics industry, is prepared by the thermal decomposition of ultra-pure trichlorosilane, followed by recrystallisation.
Silica (SiO2) in the form of sharp flints were among the first tools made by humans. The ancient civilizations used other forms of silica such as rock crystal, and knew how to turn sand into glass. Considering silicon’s abundance, it is somewhat surprising that it aroused little curiosity among early chemists.
Attempts to reduce silica to its components by electrolysis had failed. In 1811, Joseph Gay Lussac and Louis Jacques Thénard reacted silicon tetrachloride with potassium metal and produced some very impure form of silicon. The credit for discovering silicon really goes to the Swedish chemist Jöns Jacob Berzelius of Stockholm who, in 1824, obtained silicon by heating potassium fluorosilicate with potassium. The product was contaminated with potassium silicide, but he removed this by stirring it with water, with which it reacts, and thereby obtained relatively pure silicon powder.
| Atomic radius, non-bonded (Å) | 2.10 | Covalent radius (Å) | 1.14 |
| Electron affinity (kJ mol−1) | 134.068 |
Electronegativity (Pauling scale) |
1.90 |
|
Ionisation energies (kJ mol−1) |
1st
786.518
2nd
1577.134
3rd
3231.585
4th
4355.523
5th
16090.571
6th
19805.55
7th
23783.6
8th
29287.16
|
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|
| Common oxidation states | 4, -4 | ||||
| Isotopes | Isotope | Atomic mass | Natural abundance (%) | Half life | Mode of decay |
| 28Si | 27.977 | 92.223 | - | - | |
| 29Si | 28.976 | 4.685 | - | - | |
| 30Si | 29.974 | 3.092 | - | - | |
|
|
|
Specific heat capacity (J kg−1 K−1) |
712 | Young's modulus (GPa) | Unknown | |||||||||||
| Shear modulus (GPa) | Unknown | Bulk modulus (GPa) | 100 | |||||||||||
| Vapour pressure | ||||||||||||||
| Temperature (K) |
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| Pressure (Pa) |
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| Listen to Silicon Podcast |
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Transcript :
Chemistry in its element: silicon(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) Meera Senthilingam For this week's element we enter the world of science fiction to explore life in outer space. Here's Andrea Sella. Andrea Sella When I was about 12, my friends and I went through a phase of reading science fiction. There the were the fantastic worlds of Isaac Asimov, Larry Niven and Robert Heinlein, involving impossible adventures on mysterious planets - the successes of the Apollo space programme at the time only helped us suspend our disbelief. One of the themes I remember from these stories was the idea that alien life forms, often based around the element silicon, abounded elsewhere in the universe. Why silicon? Well, it is often said that elements close to each other in the periodic table share similar properties and so, seduced by the age-old red herring that "carbon is the element of life", the writers selected the element below it, silicon. I was reminded of these readings a couple of weeks ago when I went to see an exhibition of work by a couple of friends of mine. Called "Stone Hole" it consisted of stunning panoramic photographs taken at extremely high resolution inside sea caves in Cornwall. As we wandered through the gallery a thought occurred to me. "Could one imagine a world without silicon?" Every single photograph was, not surprisingly, dominated by rocks based on silicon and it was a powerful reminder of the fact that silicon is the second most abundant element in the earth's crust, beaten to first place by oxygen, the element with which it invariable entangled. Silicate rocks - those in which silicon is surrounded tetrahedrally by four oxygen atoms - exist in an astonishing variety, the differences being determined by how the tetrahedra building blocks link together, and what other elements are present to complete the picture. When the tetrahedra link one to the next, one gets a mad tangle of chains looking like an enormous pot of spaghetti - the sorts of structures one gets in ordinary glass. The purest of these chain-like materials is silicon dioxide - silica - found quite commonly in nature as the colourless mineral quartz or rock crystal. In good, crystalline quartz, the chains are arranged in beautiful helices and these can all spiral to the left. Or to the right. When this happens the crystals that result are exact mirror images of each other. But not superimposable - like left and right shoes. To a chemist, these crystals are chiral, a property once thought to be the exclusive property of the element carbon, and chirality, in turn, was imagined to be a fundamental feature of life itself. Yet here it is, in the cold, inorganic world of silicon. Most grandiose of all, one can make porous 3D structures - a bit like molecular honeycombs - particularly in the presence of other tetrahedral linkers based on aluminium. These spectacular materials are called the zeolites, or molecular sieves. By carefully tailoring the synthetic conditions, one can build material in which the pores and cavities have well defined sizes - now you have a material that can be used like a lobster traps, to catch molecules or ions of appropriate size. But what of the element itself? Freeing it from oxygen is tough, it hangs on like grim death and requires brutal conditions. It was Humphrey Davy, the Cornish chemist and showman, who first began to suspect that silica must be a compound, not an element. He applied electric currents to molten alkalis and salts and to his astonishment and delight, isolated some spectacularly reactive metals, including potassium. He now moved on to see what potassium could do. Passing potassium vapour over some silica he obtained a dark material that he could then burn and convert back to pure silica. Where he pushed, others followed. In France, Thénard and Gay-Lussac carried out similar experiments using silicon fluoride. Within a couple of years, the great Swedish analyst Jöns Jakob Berzelius had isolated a more substantial amount of the material and declared it an element. Silicon's properties are neither fish nor fowl. Dark gray in colour and with a very glossy glass-like sheen, it looks like a metal but is in fact quite a poor conductor of electricity, and there in many ways, lies the secret of its ultimate success. The problem is that electrons are trapped, a bit like pieces on a draughts board in which no spaces are free. What makes silicon, and other semiconductors, special is that it is possible to promote one of the electrons to an empty board - the conduction band - where they can move freely. It's a bit like the 3-dimensional chess played by the point-eared Dr Spock in Star Trek. Temperature is crucial. Warming a semiconductor, allow some electrons to leap, like salmon, up to the empty conduction band. And at the same time, the space left behind - known as a hole - can move too. But there is another way to make silicon conduct electricity: it seems perverse, but by deliberately introducing impurities like boron or phosphorus one can subtly change the electrical behaviour of silicon. Such tricks lie at the heart of the functioning of the silicon chips that allow you to listen to this podcast. In less than 50 years silicon has gone from being an intriguing curiosity to being one of the fundamental elements in our lives. But the question remains, is silicon's importance simply restricted to the mineral world? The prospects do not seem good - silicate fibres, like those in blue asbestos are just the right size to penetrate deep inside the lungs where they pierce and slash the inner lining of the lungs. And yet, because of its extraordinary structural variability, silicon chemistry has been harnessed by biological systems. Silicate shards lurk in the spines of nettles waiting to score the soft skin of the unwary hiker and inject minuscule amounts of irritant. And in almost unimaginable numbers delicate silicate structures are grown by the many tiny life-forms that lie at the base of marine food chains, the diatoms. Could one therefore find silicon-based aliens somewhere in space? My hunch would probably be not. Certainly not as the element. It is far too reactive and one will always find it associated with oxygen. But even linked with oxygen, it seems unlikely, or at least not under the kinds of mild conditions that we find on earth. But then again, there is nothing like a surprise to make one think. As the geneticist J B S Haldane put it, "The universe is not queerer than we suppose. It is queerer than we can suppose". I live in hope. Meera Senthilingam So although unlikely there could be some silicon based surprises lurking out in space. That was the ever hopeful Andrea Sella from University College London with the life forming chemistry of silicon. Now next week we hear about Roentgenium the element that we need to get just right. Simon Cotton The idea was to make the nickel ions penetrate the bismuth nucleus, so that the two nuclei would fuse together, making a bigger atom. The energy of the collision had to be carefully controlled, because if the nickel ions were not going fast enough, they could not overcome the repulsion between the two positive nuclei and would just fly off the bismuth on contact. However, if the nickel ions had too much energy, the resulting "compound nucleus" would have so much excess energy that it could just undergo fission and fall apart. The trick was, like Goldilocks' porridge, to be "just right", so that the fusion of the nuclei would occur, just. Meera Senthilingam And join Simon Cotton to find out how successful collisions were created by the founders of the element roentgenium in next week's Chemistry in its Element. Until then I'm Meera Senthilingam and thank you for listening. (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.
