Structure of Graphene

As demonstrated in Figure 1, graphene is a layer of graphite which is an allotrope of carbon. Carbon is the fundamental element that is regarded as the building block of life. It consists of six neutrons, six protons, and six electrons which orbit the nucleus. Out of these six electrons, two electrons lie in the innermost inert orbital (i.e. they are chemically inactive); however, the other four electrons are valence – which means that they can easily interact with external particles.

Figure 1. The structure of monolayer graphene [1]

Electrons in atoms are arranged in various types of orbitals, such as pi and sigma orbitals. Orbitals can also be thought of as being electron probability distributions. As it is impossible to predict the exact location of an electron in an atom at any given time, probability distributions are utilised to make sensible estimates of the electron location. These orbitals have different properties and hence they impact the chemical structure of atoms. For example, pi bonds are much weaker than sigma bonds; however, the combination of these two bonds creates much stronger covalent bonds.

Carbon has the unique 1S2 2S2 2P2 configuration at ground state, as demonstrated by Figure 2. This means that it has two unpaired electrons in the outer shell, which suggests that it can bind with only two additional molecules. However, experimental research suggests that carbon has a binding ability of four electrons [2].

Figure 2. The various states of Carbon [3]

This phenomenon can be explained by the fact that the energy difference between the 2s and the 2p orbital is very small, which makes it easier to excite an electron from the 2s-state into the 2p-state [2]. Such excitations of electrons can be caused by external perturbations (e.g. a Hydrogen atom). When an electron is excited to a higher state, hybridisation occurs – as showcased in Figure 3. Hybridised orbitals are fundamentally known as mixed states. This means that they are formed out of s-orbitals and p-orbitals “adding up”, as explained by the Linear Combination of Atomic Orbitals Theory. These hybridised orbitals give Carbon some of its unique chemical properties.

Figure 3. Hybridisation of Carbon Dioxide [3]

Graphene consists purely of carbon atoms joined together by extremely strong covalent bonds, which arise from the SP2-hybridisation within the carbon atoms. In graphene, only one of the s-orbitals combines with two other p-orbitals (namely p_x and p_y) [2]. These create a bond structure that is confined to a plane with the characteristic angle of 120 degrees between the hybrid orbitals [2]. The last p_z orbital lies perpendicular to this planar arrangement and forms much weaker pi-bonds, which are the weak Van-der-Waals-forces between the graphene layers in graphite [2]. However, the hybridisation of orbitals to form the planar arrangement creates strong covalent bonds which make graphene unique, as demonstrated in Figure 4.

Figure 4. The atomic structure of graphene [3]

The band structure of graphene reveals the zero-band gap between the valence band and the conduction band. Valence band refers to the highest range of electron energies in which an electron is present at absolute zero temperature, whereas the conduction band refers to the lowest range of energies unoccupied by the electrons. This connection, which will be further explored in later articles, highly influences the possibility of an electron excitation. As electrons can be excited independent of the wavelength of the incoming photons, graphene is highly opaque.

Graphene has a very unique atomic structure due to the hybridisation of orbitals. The basic knowledge of this structure can be utilised to further explain the impressive properties that graphene showcases. Moreover, with advancement in nano technology, it might be possible to alter the very fundamental structure of graphene at an atomic level to tailor it for specific uses. For example, various scientists are currently researching methods to engineer a band gap to make graphene suitable for wide ranging electronics applications [4]. Using this article as the basis, we will explore graphene’s properties – which arise due to its structure – and how they are being exploited to revolutionise engineering.

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