Making Connections with Graphene, an Emerging 2D Nanomaterial
Graphene is composed of a one atom thick layer of sp2 hybridized carbons arranged like a honeycomb lattice (Fuchs, 2008) (See Figure. 1). For a long time, scientists knew of its existence but didn’t know how to isolate it from graphite. In 2004, Andre Geim and Kostya Novoselov succeeded in isolating graphene by studying the material removed from graphite with adhesive tape- the material was graphene. They studied it under a microscope and discovered its wide range of properties, such as its light weight, high strength, flexibility, and conductivity of electricity and heat (University, 2014).
Over time, a method called chemical vapour deposition (CVD) was developed to obtain graphene. CVD extracts carbon atoms from a carbon-rich source. Its main challenge is extracting single graphene layers so as not to damage the graphene structure (De La Fuente, n.d.).
In IUPAC, graphene is known as a graphene layer. Graphene does not have any common names (Steinbeck, 2010). The name “graphene” was derived from “graphite”, which is an allotrope of carbon (Larousserie, 2013). An allotrope is “any of several crystalline forms of a chemical element” (allotrope, n.d.). For example, there are allotropes of oxygen: diatomic oxygen (O2) and ozone (O3). Allotropes are formed from network solids, which are solids formed from a large network of covalent bonds (Atkins, 2014).
Allotropes, while they have the same chemical formula, have different structures and properties. Diamond and graphite are well-known allotropes of carbon with contrasting properties (See Table. 1). Fullerene has properties of both (See Figure.2). Diamond is made up of sp3 hybridized carbons while graphite is made up of sp2 hybridized carbons. Diamond has a 3D structure because each carbon is bonded to four other carbons at 109.5° resulting in each carbon having a tetrahedral shape (See Figure. 3). Graphite has a 2D structure because of the sp2 hybridization of the carbons in each layer resulting in flat trigonal planar structures with carbons bonding at 120° (See Figure. 4). Diamond is hard due to its continuous sigma bonds and graphite is soft due to weak Van der Waal forces holding its layers together. Diamond is lustrous while graphite is opaque. Diamond does not conduct electricity because there are no free electrons as they are all used in bonding, while graphite is a good conductor of electricity because the pi bonds formed result in electrons orbiting beyond the plane which are used to conduct electric current. Pi bonds are known to be delocalized because the electrons extend over more than the bonded pair of atoms, as they have a high amount of energy. They only form after a sigma bond has been formed (See Figure. 5). Sigma bonds are known to be localized because electrons do not extend past the bonded pair of atoms because of their lower energy (See Figure. 6) (Segal, 1985). The chemical formula of all allotropes of carbon formula is C (The Editors…, 2014).
Graphene and graphite have different properties even though graphene layers make up graphite (See Table. 2). Graphene is a single layer of sp2 bonded carbons while graphite is made up of layers of graphene bonded by weak Van der Waal forces. The forces are easy to break, so layers are easy to rub off (graphite in pencils rubs off onto paper as graphene layers) (Segal, 1985). Singular graphene layers are stronger than graphite, and are 40 times stronger than diamond (for a long time known as the hardest substance to mankind) (Berger, 2014). Graphene also makes up 1D carbon nanotubes, which are rolled up sheets of graphene used in various fields such as medicine (Katsnelson, 2014).
Both graphene and graphite are electrical conductors due to pi bond electrons in the py orbitals, but graphene is a better conductor because its py orbital electrons can travel very freely throughout the