Carbon Structures

Carbon is a pretty cool element! It’s in lots of things you might not expect, like your favourite cotton t-shirt, your diamond wedding ring, and even your energy drink. In fact, it makes up almost 20% of your body! You can find it in places like your bloodstream and muscle cells. It bonds with other elements like hydrogen to form special compounds we call organic molecules, which you’ll learn more about in Organic Chemistry. But did you know that carbon can also form structures all on its own? These structures are called carbon allotropes. Allotropes are different forms of the same element that have unique structures and properties. Carbon allotropes include diamond and graphite, which are super interesting to study. But first, let’s talk about how carbon bonds with other elements.

How does carbon bond?

Carbon is a non-metal with an atomic number of 6, meaning it has six protons and six electrons. It has the electron configuration . If you aren’t sure what this means, check out Electron Configuration and Electron Shells for further information.

Carbon has atomic number 6 and mass number 12, to one decimal place

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Ignoring sub-shells, we can see in the image below that carbon has four electrons in its outer shell, also known as its valence shell.

Carbon's electron shells. It contains four valence electrons

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Carbon is pretty picky when it comes to bonding with other atoms. It prefers to form four covalent bonds, which means it shares four electrons with other atoms. This gives it a stable, full outer shell of electrons. Carbon can bond with lots of different elements like hydrogen, nitrogen, and oxygen. But for this article, we’re focusing on when carbon bonds with other carbon atoms to make different structures, which we call carbon structures or allotropes. Two of the most well-known carbon structures are diamond and graphite. Let’s learn more about them!

What is diamond?

Diamond is a unique carbon structure because it’s made up entirely of carbon atoms bonded together in a huge lattice structure. We call diamond a macromolecule because it’s a very large molecule made up of hundreds of carbon atoms all covalently bonded together. In diamond, each carbon atom forms four single covalent bonds with the other carbon atoms around it, creating a giant lattice that stretches in all directions. A lattice is just a regular, repeating arrangement of atoms or molecules. In the case of diamond, the lattice contains a large but unspecified number of carbon atoms.

The properties of diamond

That's right! Diamond has some unique properties due to its strong covalent bonds. Firstly, it has a high melting and boiling point because a lot of energy is required to overcome these strong bonds, which means that diamond is solid at room temperature. Additionally, diamond is extremely hard and strong, again due to the strength of its covalent bonds. It is also insoluble in both water and organic solvents. Finally, diamond doesn't conduct electricity because there are no charged particles free to move within the structure.

What is graphite?

Graphite is indeed another allotrope of carbon, just like diamond. However, each carbon atom in graphite only forms three covalent bonds with other carbon atoms. This creates a trigonal planar arrangement with an angle between each bond of approximately 120 degrees, as predicted by electron pair repulsion theory. The carbon atoms in graphite form a 2D hexagonal layer that resembles a sheet of paper. When stacked up, there are no covalent bonds between the layers, only weak intermolecular forces.

Although each carbon atom in graphite has one remaining electron, it moves into a region above and below the carbon atom, merging with the electrons from other carbon atoms in the same layer. All these electrons can move anywhere within this region, and we refer to them as delocalized. This delocalization of electrons is very similar to the sea of delocalization in a metal, which is a key feature of metallic bonding.

*The angle between each of the bonds in graphite is 120°*

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The properties of graphite

Graphite's unique structure gives it different physical characteristics from diamond. Some of these properties include being soft and flaky due to the weak intermolecular forces between its layers, making it easy for the layers to slide past each other and rub off (which is why it's used as the lead in pencils). Graphite also has high melting and boiling points because each carbon atom is still bonded to three other carbon atoms with strong covalent bonds, much like in diamond. It is insoluble in water, much like diamond, and is also a good conductor of electricity due to the delocalized electrons that are free to move between the layers of the structure and carry a charge.

Graphene

A single sheet of graphite is called graphene. It is the thinnest material ever isolated - it is just one atom thick. Graphene has similar properties to graphite. For example, it is a great conductor of electricity. However, it is also low density, flexible and extremely strong for its mass. In the future you might find wearable electronics made from graphene embedded in your clothing. We currently use it for drug delivery and solar panels.

Comparing diamond and graphite

Although diamond and graphite have many similarities, they also have their differences. The following table summarises this information.

A table summarising the similarities and differences between diamond and graphite

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You've summarized the key takeaways from our discussion of carbon structures very well! Carbon atoms are extremely versatile and can form multiple different structures, each with unique properties. Allotropes are different forms of the same element, and carbon has several allotropes, including diamond and graphite. Diamond consists of a giant lattice of carbon atoms, each joined together by four covalent bonds, which makes it hard, strong, and have a high melting point. Graphite, on the other hand, contains sheets of carbon atoms each joined by three covalent bonds, and its spare electrons are delocalized above and below each carbon sheet, making graphite soft, flaky, and a good conductor of electricity.