Lesson Video: Allotropes of Carbon Chemistry

In this video, we will learn how to describe the allotropes of carbon: graphite, graphene, diamond, fullerenes, and nanotubes and compare their physical properties.


Video Transcript

In this video, we will learn about the allotropes of carbon. Specifically, we will learn what an allotrope is and have a look at the different allotropes the element carbon can exist as. We will investigate their structures and how this influences their physical properties. Lastly, we will discuss the uses of these allotropes.

Let’s begin by discussing what an allotrope is. Did you know that the gray tips of pencils and shiny, colorless diamonds that we saw in the cartoon at the beginning of this video are both composed of the same element? This element is carbon. The tip of a pencil and diamond are two different forms of the element carbon. We call these substances allotropes of carbon. An allotrope can be defined as one of the different forms of the same element. In other words, some elements can exist in various forms.

Not all elements can form allotropes, but only some of the elements in groups 13 to 16 of the periodic table, namely carbon, phosphorus, boron, tin, arsenic, oxygen, sulfur, and selenium. Here we will only discuss carbon allotropes. It is important to remember that an allotrope is not the same as a compound. A compound is a substance formed when different elements chemically bond together. So an allotrope refers to the same element existing in different forms, while a compound refers to different elements bonded together.

Now that we know what an allotrope is, let’s have a look at the various allotropes that carbon can exist as. The different forms of carbon can be categorized into two main groups. The crystalline carbon allotropes and amorphous carbon. The term crystalline tells us that the carbon atoms in these allotropes are arranged in a highly regular and ordered pattern in the lattice. Amorphous refers to a noncrystalline structure, in other words, an atom arrangement that does not have a clearly defined structural shape at the microscopic level.

Sometimes coal, soot, and charcoal are included in the category of amorphous carbon as allotropes. But strictly speaking, these three materials are not 100 percent amorphous, nor are they 100 percent made of carbon. So to be 100 percent accurate, we should not include these in the category of amorphous carbon allotropes.

The crystalline allotropes are what we will be focusing on in this lesson. We will have a look at diamond, graphene structures, and fullerenes. The graphene structures include graphene itself and graphite, whose top views are shown in the little diagram but whose side views differ, and nanotubes. Carbon nanotubes can also be classified as fullerenes although sometimes nanotubes are put into a category all of their own on the same level as diamond, graphene, and fullerenes.

There are a growing number of crystalline carbon allotropes with interesting shapes and properties. This is because research is advancing as scientists do much investigation to synthesize new forms of carbon. These new allotropes are referred to as exotic allotropes, but that’s a discussion for another day. Let’s have a look in depth at the commonly known and researched allotropes. Let’s start with diamond. The diamonds that we know from jewelry are cut in the way that they are to maximize the amount of light reflected from them to make them look all sparkly and pretty. But there’s much more to diamonds than use in jewelry.

Imagine if we could zoom in on this structure. Of course, a magnifying glass wouldn’t work; we’d need an electron microscope. We’d see the following internal arrangement of atoms. Diamond is a network covalent structure, otherwise known as a giant covalent structure. Each carbon atom is covalently bonded to four other carbon atoms. So if we take this carbon atom, we can see that it is bonded to one, two, three, four other carbon atoms. There is a tetrahedral arrangement of the four carbon atoms around the central carbon atom.

And because almost every carbon atom in the network covalent structure could be considered a central carbon atom except for those at the edges each with their own tetrahedral array around them, we end up with a three-dimensional lattice composed of these tetrahedral structures all joined together, radiating out in every direction. This structure gives diamond specific properties and related uses. Because of this chemical structure, diamond is rigid and very hard, which makes diamond useful in applications such as in polishing, drilling, and cutting tools. Diamond coatings can be used to strengthen the surface of materials and be used to reflect light from surfaces.

And because of diamond’s highly reflective nature, when cut in a certain way, it is perfect for use in jewelry. Diamond is not very useful as an electrical conductor. This is because the four valence electrons in each carbon atom of diamond are used in covalent bonding and are not free to flow in an electrical circuit. What is interesting, though, is that diamond does conduct heat fairly well. Let’s move on and learn a bit about graphene structures.

Let’s start with graphene itself. Unlike diamond, in graphene, each carbon atom is covalently bonded not to four but three other carbon atoms. If we take this carbon atom here, we can see that it is bonded to one, two, three other carbon atoms in graphene. All the atoms are bonded together in one flat layer or sheet one atom thick. The arrangement of atoms in graphene give it a honeycomb appearance because of the repeating hexagons bonded together.

Now, the three bonds that each carbon atom makes use up three of the valence electrons in each carbon atom. And one of the valence electrons, the fourth valence electron in each carbon atom, we say is delocalized. But what does this mean? If we were to rotate the graphene sheet by 90 degrees in the direction shown, it would just look like a flat sheet. The fourth valence electron in each carbon atom is free to move. If a potential is applied across the graphene sheet, these electrons will move. And that is why they are called delocalized electrons because they are not located in one specific place at all times.

Because of these delocalized electrons, graphene is a good thermal and electrical conductor. Much research is still needed, but scientists and engineers hope that in the future graphene sheets will be useful in microelectronics, batteries, and even fuel cells for the storage of energy. Graphene is very strong, and because of this, it is already being used in some composite materials to improve strength. Because of their flat nature, graphene sheets have a very high surface area and have a potential application as supports for catalysts. Individual graphene sheets were only synthesized for the first time in 2004, so this allotrope of carbon, synthetically made, is relatively new.

But what about its naturally occurring and long-produced cousin, graphite? Graphite is much like graphene. Each carbon atom is also covalently bonded to three other carbon atoms. There is a honeycomb structure of repeating hexagons, and each carbon atom has one delocalized valence electron. So what’s the difference? If we rotate graphite by 90 degrees, we won’t see one layer but many layers stacked on top of each other. Graphite consists of several to many individual graphene sheets stacked on top of each other.

The delocalized electrons between the layers as well as weak interlayer forces of attraction give this allotrope certain properties and thus certain applications. Because of the weak interlayer forces, the layers can slide on top of each other. We say that graphite is soft and slippery, and this makes it a great lubricant to reduce friction between certain surfaces and machinery in industry. An example of this is that natural graphite is often used in the brake linings of heavy vehicles.

What about electrical conductivity? Because of these delocalized electrons, if a potential difference is applied across a piece of graphite, for example, a graphite rod, electrons will flow creating an electrical current. And so graphite is a very good electrical and thermal conductor. The electrodes used in the scrap metal industry are usually made of graphite. Polymers of graphite are lightweight and strong and find use in carbon fiber materials. Carbon fiber is often used to reinforce plastics. Many tennis record frames and bicycle frames, these days, are reinforced with carbon fiber.

Now we know about graphene and its cousin graphite. Let’s move on to the third type of carbon allotrope, the fullerenes. Fullerenes are an interesting family of carbon allotropes. The first fullerene, called fullerene and shown in the diagram, was named after Richard Buckminster Fuller. He did not synthesize the first fullerene, but was an author, inventor, and architect. He designed geodesic domes. These are dome-shaped buildings, which look much like half of a soccer ball. For this reason, when fullerene was first synthesized and its structure was identified to look pretty much like a soccer ball, this architect’s name was chosen for this new class of carbon allotropes.

The first fullerene consisted of 60 carbon atoms joined together as shown in the picture. Other fullerenes in this group might not have 60 carbon atoms, but they all have each carbon atom bonded covalently to three other carbon atoms. Let’s look at this carbon atom. It is bonded to one, two, three other carbon atoms. There is a mixture of single and double bonds in these large molecules, as well as five-, six-, or seven-membered carbon rings, which means that these molecules contain ring structures consisting of five, six, or seven carbon atoms.

In the carbon 60 fullerene, there are only five- and six-membered rings. Here is a five-membered ring, in other words, a pentagon, of carbon atoms. And here is a six-membered ring, in other words, a hexagon, of carbon atoms. Some fullerenes, for example, the carbon 60 fullerene, are spheroidal which means they are 3D geometric structures whose overall shape resembles a sphere, but they have flat faces or flat surfaces. Other fullerenes, not shown here, are ellipsoidal, they look like ellipsis, or are tubelike. And some are referred to as being nested structures because they are structures within structures within structures.

Now the common name of carbon 60 is a buckyball after Buckminster Fuller and because it looks like a ball. The simple 2D picture of a buckyball within a buckyball within a buckyball, in other words, a nested structure, has the fun common name a bucky onion because of the many layers. Let’s investigate the properties and uses of this class of carbon allotropes. There are weak intermolecular forces between individual fullerene molecules, making this material slippery. And so in special applications, they are used as lubricants. Because of their high surface area, they can potentially be used as catalysts and catalyst supports.

Although this is not strictly a physical property of fullerenes, one of their abilities is the ability to contain within their cage other atoms. Drug companies are able to insert certain atoms into the cages of fullerenes or to attach compounds to the outside of fullerenes. And these altered fullerenes are then used as carriers to deliver drugs into the body in medicine. These altered fullerenes are used in MRI imaging and in X-rays.

One special group of fullerenes which has gained much research over the last few decades are the nanotubes. Carbon nanotubes are very much like other fullerenes. Each carbon atom is covalently bonded to three other carbon atoms. Bonds between carbon atoms can be single or double, and there is a hexagonal arrangement of rings, much like in graphene and graphite. In fact, carbon nanotubes can be considered to be rolled-up sheets of graphene. These tubes can be single or nested, which we call single-walled carbon nanotubes or multiwalled carbon nanotubes.

Because of their structure, nanotubes are excellent electrical and thermal conductors, much like graphene and graphite. Much research is being done on the use of carbon nanotubes in electronics. They can be used in composite materials where high strength and flexibility are both needed. And, like graphene, graphite, and fullerenes, they have a very high surface area and so are well suited as supports for metal catalysts. In this video, we have discussed only the main properties and uses of carbon allotropes. Much research is being done and is still needed for the potential applications of these allotropes. And the field of research is widening rapidly. Let’s summarize what we have learned.

We have learned that an allotrope is one of the different forms of the same element. We learned what the crystalline carbon allotropes are and a classification system for them, although that classification system is not shown here. We saw that in diamond, each carbon atom is bonded to four other carbon atoms by single bonds and that, in the remaining allotropes, each carbon atom is bonded to three other carbon atoms and there is a mixture of single and double bonds.

We learned that diamond forms a network 3D array of carbon atoms, that graphene is composed of a layer of carbon atoms, while graphite stacked layers of carbon atoms, that fullerene molecules are large cage-like molecules, and that nanotubes are essentially rolled-up sheets of graphene. The main properties we investigated were hardness for diamond, slipperiness for graphene, graphite, and fullerenes, good to excellent conducting abilities for graphene, graphite, and nanotubes, and high strength for carbon nanotubes.

So the next time you use a pencil, remember that you are sliding off onto the paper a thin layer of graphite. And remind yourself of the carbon allotropes and their many interesting properties and uses.

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