Carbon

Description and uses

Carbon is the 6th element of the periodic table, and one of the two elements (the other one being arsenic) with a sublimation point at 1 atm pressure, instead of a melting point. This is because the triple point (the point at which the melting and boiling point are the same) is very high, at 5.2 atm pressure.

Carbon forms several allotropes at S.T.P., the most stable one being graphite (meaning that it will not change form). At S.T.P, diamond is not completely stable, meaning that it will eventually turn into graphite (although it will take a huge amount of time). Graphite is formed of hexagonally bonded sheets of carbon stacked on top of each other. It has delocalised electrons, allowing it to conduct electricity. It is soft, as layers can easily slide over each other due to weak intermolecular forces.

Carbon can easily form fuels in the form of hydrocarbons (molecules exclusively containing carbon and hydrogen), and other hydrogen-carbon compounds such as alcohols. This is because carbon covalent bonds can store large amounts of energy, which can be broken and released during combustion. Examples of carbon fuels include coal, crude oil, natural gases and fats. Carbon easily form strongly bonded chains surrounded by hydrogen, sometimes with other elements such as oxygen and nitrogen. These can act as large stores of energy.

Carbon has 3 naturally occurring isotopes being ¹²C, ¹³C and ¹⁴C. ¹²C is the most abundant, taking up around 98.9% of all carbon atoms. The definition of 1 amu (atomic mass unit) is exactly 1/12 the mass of a carbon-12 atom at S.T.P. ¹³C is the next most common isotope, with an abundance of 1.1%. ¹⁴C (also known as radiocarbon) is radioactive with a half life of around 5700 years, which forms when ¹⁴N in our atmosphere absorbs a neutron from a star, and it releases a proton, becoming ¹⁴C.