Is Carbon a Conductor? (+ 3 Fascinating Facts to Know)

Carbon can exist in different forms, including both conductive and non-conductive forms. Graphite, a form of carbon, is a good conductor of electricity due to its layered structure that allows free movement of electrons. ¹ On the other hand, diamond, another form of carbon, is an insulator and does not conduct electricity due to its tightly bonded carbon atoms. ²

Well, this was just a simple answer. But there are few more things to know about this topic which will make your concept super clear.

So let’s dive right into it.

Why is graphite a good conductor?

Graphite is a good conductor because it possesses a unique structure that allows it to conduct electricity efficiently. Its layered arrangement of carbon atoms creates a network of delocalized electrons that are free to move, enabling the flow of electric current. ³

In more detail, graphite consists of layers of carbon atoms arranged in a hexagonal lattice. Within each layer, the carbon atoms are bonded tightly together in a covalent structure. However, the bonds between the layers are relatively weak, allowing the layers to slide past each other. ⁴

This structure results in the formation of “pi” bonds, where electrons are shared above and below the plane of the carbon layers.

The presence of these “pi” bonds creates a network of delocalized electrons that are not associated with any particular atom. These electrons are free to move throughout the graphite structure, making it an excellent conductor of electricity.

When a voltage is applied, the delocalized electrons can easily flow through the layers, carrying the electric current.

Furthermore, graphite also exhibits good thermal conductivity due to its layered structure, where heat can be conducted along the planes of carbon atoms. ⁵ This property makes graphite useful in applications such as electrical contacts, electrodes, and as a lubricant in high-temperature environments.

Why is diamond an insulator?

Diamond is an insulator because its structure lacks free electrons or mobile charged particles that can conduct electricity. The strong covalent bonds between carbon atoms in diamond make it a poor conductor of electricity. ⁶

Diamond is composed of carbon atoms arranged in a tightly bonded lattice structure. Each carbon atom forms four covalent bonds with its neighboring carbon atoms, resulting in a three-dimensional network of interconnected carbon atoms. ^{7 8}

Unlike graphite, where the carbon atoms form layers with delocalized electrons, diamond’s structure does not have any free or loosely bound electrons. The covalent bonds in diamond are very strong, requiring a significant amount of energy to break them and release electrons for conduction.

As a result, diamond does not conduct electricity under normal conditions. It lacks the presence of free electrons or charged particles that can carry an electric current.

This property makes diamond an excellent insulator and a valuable material in applications where electrical conductivity is undesirable or where electrical insulation is required, such as in high-voltage applications or electronic devices that require excellent heat dissipation.

Further reading

Why is Graphite Conductive?
Is Graphite a Metal?
Is Graphite an Element?
Is Diamond an Element or a Compound?
Is Diamond a Mineral or a Rock? 

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6. Study.com https://homework.study.com/explanation/diamond-and-graphite-are-both-made-of-carbon-diamond-is-an-electrical-insulator-is-the-hardest-material-in-the-world-and-is-often-used-as-a-cutting-tool-in-drastic-contrast-graphite-is-an-excellent-electrical-conductor-is-easy-to-be-cut-and-is-ofte.html
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8. Giant covalent structures. (n.d.). Giant Covalent Structures. https://www.chemguide.co.uk/atoms/structures/giantcov.html