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There has been a lot of news about Graphene since its discovery in 2004. And as we are all told it is a revolutionary material which is very strong, conductive and transparent.

But what is it about the structure of Graphene which makes it so strong?

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In a perfect graphene sheet, all carbon atoms are sp$^2$-hybridized, with three in-plane $\sigma$-orbitals and two out-of-plan $\pi$-orbitals. This means that each carbon atom can form equivalent s-bonds to each of its three neighboring atoms. The bonding energy of one C-C bond in graphene amounts to 4.93 eV [doi:10.1088/0953-8984/14/4/312]. These strong (covalent) $\sigma$-bonds are responsible for the extraordinary mechanical properties of graphene.

The out-of-plane orbitals are responsible for the $\pi$ bands (Dirac cones) and lead to the nice electric properties of graphene, but they do not significantly contribute to the mechanical "strongness" of graphene.

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Graphene's strength mainly comes from the strong covalent bonds of the carbon atoms.

Graphite is made of layers of graphene but it is weaker because the layers making up graphite are bonded to each other through London forces hence why the layers can slide past each other and the material is soft. These weak inter-layer London forces provide a weak point in graphite's structure which doesn't exist in graphene.

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What makes Graphene so strong is its electrostatic forces resulting from delocalized electrons flowing through positively charged carbon atoms. This diffrence in charge creates a strong electrostatic attraction that holds Graphene together. This phenomenon also explains why it is such a strong conductor.

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