|Group||14||Melting point||3825 oC, 6917 oF, 4098.15 K|
|Block||p||Density (kg m-3)||3513 (diamond); 2267 (graphite)|
|Atomic number||6||Relative atomic mass||12.011|
|State at room temperature||Solid||Key isotopes||12C, 13C, 14C|
|Electron configuration||[He] 2s22p2||CAS number||7440-44-0:1333-86-4|
|ChemSpider ID||4575370||ChemSpider is a free chemical structure database|
The three crowns represent the three states of Carbon in nature, plus the suggestion of a heraldic motif, emphasising Carbon’s regal status, the "King of the Elements", in the Periodic Table.
There are a number of pure forms of this element including graphite, diamond, fullerenes and graphene.
Diamond is a colourless transparent crystalline solid, the hardest known material. Graphite is black and shiny but soft, and the nano-forms, fullerenes and graphene appear, in bulk, as black or dark brown soot-like powders
Carbon is unique among the elements in its ability to form strongly bonded chains, sealed off by hydrogen atoms. These hydrocarbons, extracted naturally as fossil fuels (coal, oil and natural gas), are mostly used, when combusted with oxygen, as a source of energy for transport, electrical energy generation and industry. A small but important fraction are used as feedstock for the petrochemical industries producing polymers, fibres, paints solvents, plastics etc. Impure carbon in the form of charcoal (from wood) and coke (from coal) is used in metal smelting, especially for the iron and steel industry. Industrial diamonds are used for cutting rocks and drilling. More recently the discovery of carbon nano-tubes, other fullerenes and atom-thin sheets of graphene are revolutionising hardware developments in the electronics industry and in nano-technology generally.
150 years ago the natural concentration of carbon dioxide in the earth’s atmosphere was 280 ppm. In 2013, as a result of combusting fossil fuels with oxygen, there were 390 ppm. Atmospheric carbon dioxide allows visible light in but prevents some infra-red escaping (the natural greenhouse effect), keeping the Earth warm enough to sustain life. However, an enhanced greenhouse effect due to human-induced rise in atmospheric carbon dioxide is being felt by living things as our climate changes.
The carbon-based molecules of life were once thought to be obtainable only from living things – they were thought to contain a ‘spark of life’. The synthesis of urea in 1828 from inorganic reagents united the two branches of chemistry and this theory was forgotten. Living things get almost all their carbon from carbon dioxide, either from the atmosphere or dissolved in water, through photosynthesis in green plants and photosynthetic plankton. The sun’s energy splits water into oxygen (released to the atmosphere, into fresh water and the seas) and hydrogen (which joins with carbon dioxide to produce carbohydrates).
With the addition of other elements, especially nitrogen and phosphorus, some of the carbohydrates are used to form the other monomer molecules of life, including bases and sugars for RNA and DNA, and amino acids for proteins. Life forms which do not photosynthesise have to rely on consuming other living things for their source of carbon molecules. Their digestive systems break the polymers in their diet into the simple monomers to build their own cellular structures. All these reactions need a source of energy, obtained by respiration, most commonly by allowing the oxygen previously vented to the environment by photosynthesis to rejoin the carbohydrate, thus reforming carbon dioxide and water, and making energy available for the cells. Anaerobic respiration (without oxygen) provides a small amount of energy by re-arranging the bonds in, usually, sugar molecules.
Carbon is found in the sun and other stars formed from the debris of a previous supernova, and it is built up by nuclear fusion in bigger stars. It is present in the atmospheres of many planets usually as carbon dioxide (currently 390 ppm and rising for our own atmosphere). Graphite is found naturally in many locations. Diamond is found in the form of microscopic crystals in some meteorites. Natural diamonds are found in the mineral kimberlite, sources of which are in
Molar heat capacity
(J mol-1 K-1)
|6.155 (diamond); 8.517 (graphite)||Young's modulus (GPa)||1050 (diamond)|
|Shear modulus (GPa)||478 (diamond)||Bulk modulus (GPa)||542 (diamond);33 (graphite)|
Carbon occurs naturally as anthracite (a type of coal), graphite, and diamond. More readily available historically was soot or charcoal. Ultimately these various materials were recognised as forms of the same element. Not surprisingly, diamond posed the greatest difficulty of identification. Naturalist Giuseppe Averani and medic Cipriano Targioni of Florence were the first to discover that diamonds could be destroyed by heating. In 1694 they focussed sunlight on to a diamond using a large magnifying glass and the gem eventually disappeared. Pierre-Joseph Macquer and Godefroy de Villetaneuse repeated the experiment in 1771. Then, in 1796, the English chemist Smithson Tennant finally proved that diamond was just a form of carbon by showing that as it burned it formed only CO2.
|Listen to Carbon Podcast|
Chemistry in Its Element - Carbon
You're listening to Chemistry in its element brought to you by Chemistry World, the magazine of the Royal Society of Chemistry
Hello, this week to the element that unites weddings, wars, conflicts and cremations and to explain how, here's Katherine Holt.
Any chemist could talk for days about carbon. It is after all an everyday, run-of
-the-mill, found-in-pretty-much-everything, ubiquitous element for us carbon-based life forms. An entire branch of chemistry is devoted to its reactions.
In its elemental form it throws up some surprises in the contrasting and fascinating forms of its allotropes. It seems that every few years a new form of carbon comes into fashion - A few years ago carbon nanotubes were the new black (or should I say 'the new bucky-ball') - but graphene is oh-so-now!
But today I'm going to talk about the most glamorous form that carbon can take - diamond. For millennia diamond has been associated with wealth and riches, as it can be cut to form gemstones of high clarity, brilliance and permanence. Diamonds truly are forever! Unfortunately, diamond also has a dark side - the greed that diamond induces leads to the trade of so-called 'conflict diamonds' that support and fund civil wars.
Mans desire for diamond has led alchemists and chemists over many centuries to attempt to synthesise the material. After many fraudulous early claims diamond was finally synthesised artificially in the 1950s. Scientists took their inspiration from nature by noting the conditions under which diamond is formed naturally, deep under the earth's crust. They therefore used high temperatures (over 3000oC) and high pressures (>130 atms) to turn graphite into carbon. This was an impressive feat, but the extreme conditions required made it prohibitively expensive as a commercial process. Since then the process has been refined and the use of metal catalysts means that lower temperatures and pressures are required. Crystals of a few micron diameter can be formed in a few minutes, but a 2-carat gem quality crystal may takes several weeks.
These techniques mean its now possible to artificially synthesise gemstone quality diamonds which, without the help of specialist equipment, cannot be distinguished from natural diamond. It goes without saying that this could cause headaches among the companies that trade in natural diamond! It is possible to turn any carbon based material into a diamond - including hair and even cremating remains! Yes - you can turn your dearly departed pet into a diamond to keep forever if you want to! Artificial diamonds are chemically and physical identical to the natural stones and come without the ethical baggage. However, psychologically their remains a barrier - if he really loves you he'd buy you real diamond - wouldn't he?
From the perspective of a chemist, materials scientist or engineer we soon run out of superlatives while describing the amazing physical, electronic and chemical properties of diamond. It is the hardest material known to man and more or less inert - able to withstand the strongest and most corrosive of acids. It has the highest thermal conductivity of any material, so is excellent at dissipating heat. That is why diamonds are always cold to the touch. Having a wide band gap, it is the text book example of an insulating material and for the same reason has amazing transparency and optical properties over the widest range of wavelengths of any solid material.
You can see then why diamond is exciting to scientists. Its hardness and inert nature suggest applications as protective coatings against abrasion, chemical corrosion and radiation damage. Its high thermal conductivity and electrical insulation cry out for uses in high powered electronics. Its optical properties are ideal for windows and lenses and its biocompatibility could be exploited in coatings for implants.
These properties have been known for centuries - so why then is the use of diamond not more widespread? The reason is that natural diamond and diamonds formed by high pressure high temperature synthesis are of limited size - usually a few millimeters at most, and can only be cut and shaped along specific crystal faces. This prevents the use of diamond in most of the suggested applications.
However, about 20 years ago scientists discovered a new way to synthesise diamond this time under low pressure, high temperature conditions, using chemical vapour deposition. If one were to consider the thermodynamic stability of carbon, we would find that at room temperature and pressure the most stable form of carbon is actually graphite, not diamond. Strictly speaking, from a purely energetic or thermodynamic point of view, diamond should spontaneously turn into graphite under ambient conditions! Clearly this doesn't happen and that is because the energy required to break the strong bonds in diamond and rearrange them to form graphite requires a large input of energy and so the whole process is so slow that on the scale of millennia the reaction does not take place.
It is this metastability of diamond that is exploited in chemical vapour deposition. A gas mixture of 99 % hydrogen and 1 % of methane is used and some activation source like a hot filament employed to produce highly reactive hydrogen and methyl radicals. The carbon-based molecules then deposit on a surface to form a coating or thin film of diamond. Actually both graphite and diamond are initially formed, but under these highly reactive conditions, the graphitic deposits are etched off the surface, leaving only the diamond. The films are polycrystalline, consisting of crystallites in the micron size range so lack the clarity and brilliance of gemstone diamond. While they may not be as pretty, these diamond films can be deposited on a range of surfaces of different size and shapes and so hugely increase the potential applications of diamond. Challenges still remain to understand the complex chemistry of the intercrystalline boundaries and surface chemistry of the films and to learn how best to exploit them. This material will be keeping chemists, materials scientists, physicists and engineers busy for many years to come. However, at present we can all agree that there is more to diamond than just a pretty face!
Katherine Holt extolling the virtues of the jewel in carbon's crown. Next week we're heading to the top of group one to hear the story of the metal that revolutionised the treatment of manic depression.
Its calming effect on the brain was first noted in 1949, by an Australian doctor, John Cade, of the Victoria Department of Mental Hygiene. He had injected guinea pigs with a 0.5% solution of lithium carbonate, and to his surprise these normally highly-strung animals became docile. Cade then gave his most mentally disturbed patient an injection of the same solution. The man responded so well that within days he was transferred to a normal hospital ward and was soon back at work.
And it's still used today although despite 50 years of medical progress we still don't know how it works. That was Matt Wilkinson who will be here with the story of Lithium on next week's Chemistry in its Element, I do hope you can join us. I'm Chris Smith, thank you for listening and goodbye.
Chemistry in its elementis brought to you by the Royal Society of Chemistry and produced by thenakedscientists dot com. There's more information and other episodes of Chemistry in its element on our website at chemistryworld dot org forward slash elements.
Mining and Sourcing data: British Geological Survey – natural environment research council.
Text: John Emsley Nature’s Building Blocks: An A-Z Guide to the Elements, Oxford University Press, 2nd Edition, 2011.
Additional information for platinum, gold, neodymium and dysprosium obtained from Material Value Consultancy Ltd www.matvalue.com
Data: CRC Handbook of Chemistry and Physics, CRC Press, 92nd Edition, 2011.
G. W. C. Kaye and T. H. Laby Tables of Physical and Chemical Constants, Longman, 16th Edition, 1995.
Members of the RSC can access these books through our library.