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Graphene & Graphite - How Do They Compare?

By Mark Walters  Posted by Mark Walters (about the submitter)       (Page 1 of 1 pages)   No comments
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In very basic terms graphene could be described as a single, one atom thick layer of the commonly found mineral graphite; graphite is essentially made up of hundreds of thousands of layers of graphene. In actuality, the structural make-up of graphite and graphene, and the method of how to create one from the other, is slightly different.


Way back when you were at school, it is very likely that you would have come across the term "pencil lead', referring to the central core of a pencil that is able to produce marks on paper and other material. In fact, rather than referring to the chemical element and heavy metal, lead, this central core is most commonly made from graphite mixed with clay. The mistake arose when it was first discovered, at which point because it is a form of carbon and contains a similar molecular make up to other members of the carbon group (though primarily due to the visual similarities), it was thought to be a form of lead.

Graphite is a mineral that naturally occurs in metamorphic rock in different continents of the world, including Asia, South America and some parts of North America. It is formed as a result of the reduction of sedimentary carbon compounds during metamorphism. Contrary to common belief, the chemical bonds in graphite are actually stronger than those that make up diamond. However, what defines the difference in hardness of the two compounds is the lattice structure of the carbon atoms contained within; diamonds containing three dimensional lattice bonds, and graphite containing two dimensional lattice bonds (layers of carbon sheets). While within each layer of graphite the carbon atoms contain very strong bonds, the layers are able to slide across each other, making graphite a softer, more malleable material.

Extensive research over hundreds of years has proved that graphite is an impressive mineral showing a number of outstanding and superlative properties including its ability to conduct electricity and heat well, having the highest natural stiffness and strength even in temperatures exceeding 3600 degrees Celsius, and it is also highly resistant to chemical attack and self-lubricating. However, while it was first identified over a thousand years ago and first named in 1789, it has taken a while for industry to realise the full potential of this amazing material.

Graphite is one of only three naturally occurring allotropes of carbon (the others being amorphous carbon and diamond). The difference between the three naturally occurring allotropes is the structure and bonding of the atoms within the allotropes; diamond enjoying a diamond lattice crystalline structure, graphite having a honeycomb lattice structure, and amorphous carbon (such as coal or soot) does not have a crystalline structure.

While there are many different forms of carbon, graphite is of an extremely high grade and is the most stable under standard conditions. Therefore, it is commonly used in thermochemistry as the standard state for defining the heat formation of compounds made from carbon. It is found naturally in three different forms: crystalline flake, amorphous and lump or vein graphite, and depending on its form, is used for a number of different applications.

As previously touched upon, graphite has a planar, layered structure; each layer being made up of carbon atoms linked together in a hexagonal lattice. These links, or covalent bonds as they are more technically known, are extremely strong, and the carbon atoms are separated by only 0.142 nanometres. The carbon atoms are linked together by very sturdy sp2 hybridised bonds in a single layer of atoms, two dimensionally. Each individual, two dimensional, one atom thick layer of sp2 bonded carbon atoms in graphite is separated by 0.335nm. Essentially, the crystalline flake form of graphite, as mentioned earlier, is simply hundreds of thousands of individual layers of linked carbon atoms stacked together.


So, graphene is fundamentally one single layer of graphite; a layer of sp2 bonded carbon atoms arranged in a honeycomb (hexagonal) lattice. However, graphene offers some impressive propertiesthat exceed those of graphite as it is isolated from its "mother material'. Graphite is naturally a very brittle compound and cannot be used as a structural material on its own due to its sheer planes (although it is often used to reinforce steel). Graphene, on the other hand, is the strongest material ever recorded, more than three hundred times stronger than A36 structural steel, at 130 gigapascals, and more than forty times stronger than diamond.

Due to graphite's planar structure, its thermal, acoustic and electronic properties are highly anisotropic, meaning that phonons travel much more easily along the planes than they do when attempting to travel through the planes. Graphene, on the other hand, being a single layer of atoms and having very high electron mobility, offers fantastic levels of electronic conduction due to the occurrence of a free pi (Ï ) electron for each carbon atom.

However, for this high level of electronic conductivity to be realised, doping (with electrons or holes) must occur to overcome the zero density of states which can be observed at the Dirac points of graphene. The high level of electronic conductivity has been explained to be due to the occurrence of quasiparticles; electrons that act as if they have no mass, much like photons, and can travel relatively long distances without scattering (these electrons are hence known as massless Dirac fermions).


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