Lesson video

Hello, my name is Mr. Gundry and you join me today as we delve into an intriguing aspect of chemistry focusing on the distinctive properties of diamond and graphite.

This lesson, titled Diamond and Graphite, is an integral part of our exploration into the unit of chemistry of carbon.

By the end of our session, you'll be equipped to describe the properties of diamond and graphite.

Understand how these properties relate to their structures and relate these characteristics to their practical applications.

This lesson aligns with our broader quest to understand and answer fundamental big questions such as how do we explain how substances behave and what are things made of? Through this lens, we will see how the versatile bonding of carbon not only creates a myriad of structures, but also gives rise to materials with unique and varied properties.

Let's embark on this learning journey together to uncover the marvels of carbon, enhancing our understanding of the material world.

Before we delve into the complexities of these materials, let's familiarise ourselves with some of the essential terms that will be pivotal in our discussion today.

These are allotrope, giant covalent structure, forces of attraction and delocalised electrons.

Understanding these concepts is crucial as they form the foundation upon which the properties of diamond and graphite are built.

In the next slide, you'll see their definitions, but I won't explain them now, but as we make our way through the lesson.

We're gonna start this lesson by reminding ourselves about what giant covalent structures are and why diamond and graphite have this structure.

In discussing the versatility of carbon, we can consider the vast spectrum of the structures that it can form, ranging from simple molecules to giant covalent structures.

Let's delve into this further.

Simple carbon molecules are characterised by their definitive small size and precise number of atoms. These molecules are held together by covalent bonds, which are strong bonds formed when atoms share pairs of electrons.

In these structures, individual molecules are discreet entities and separate from each other.

On the other hand, we have giant covalent structures.

These are not just large, they are massive lattices of atoms also held together by covalent bonds.

Unlike in simple molecules, the numbers of atoms in these structures are not fixed or defined.

They're incredibly large numbers and the exact number is often undetermined.

These structures are not composed of individual separate molecules, but rather have a continuous network of covalently bonded atoms, the repeat in a regular pattern, creating a solid framework that extends normally in three dimensions.

This distinction is crucial because it leads to a significant difference in properties and behaviours between the substances made of simple molecules and those composed of a giant covalent structure.

Here's a true or false question for you.

All giant structures are held together by covalent bonds.

What are your thoughts on that? Well, I can tell you the answer is false and I'd like you to pause the video to read through the justifications below and press play when you are ready to move on.

While covalent bonds are prevalent in some giant structures such as those found in diamond and graphite, as well as silicon dioxide, which you may have learned about before, this is not a universal binding force.

We also encounter giant ionic structures like sodium chloride, salt, which is held together by the electrostatic force of attraction between ions and giant metallic structures, such as copper, as a metal, held together by metallic bonds where electrons are shared amongst a lattice of positively charged metal ions.

So in essence, the variety of bonds, covalent, ionic, and metallic contribute to the formation of giant structures, each lending different properties, which we'll learn a little bit more about the giant covalent properties later.

So carbon is quite a versatile element, capable of forming a range of substances solely with its own atoms, and the structures that arise from these self combinations are unique in their properties and applications.

These different structural forms of a single element are termed allotropes.

Carbon is not the only element that can form allotropes, but we're only gonna be looking at carbon today.

We're also not going to be looking at all of carbon's allotropes, just two specific ones.

These two notable allotropes that we're gonna look at exemplify the giant covalent structure of carbon in diamond and in graphite.

Both diamond and graphite are composed of vast networks of carbon atoms. The actual number of carbon atoms in each structure depends on the size of the diamond or the sample of the graphite that you are looking at.

Since the number of carbon atoms is not fixed, it's likely in the trillions, quadrillions or even larger numbers.

And since this, we cannot give a precise number or particular molecular formula like you might for a simple molecular compound.

Instead, both diamond and graphite are represented with the simplest formula of just 'C'.

This stands for carbon indicating that both of these substances are pure elements made of entirely carbon atoms. It's more of an empirical formula, if you will, showing that the only substance in this is carbon, albeit it's arranged in different forms for each.

This is why we use the term allotropes for such variations, different structural forms of the same element.

I'd like you to read through the following statements and decide which ones accurately describe diamond and graphite.

Pause the video if you need some time to do this and press play when you are ready to move on.

Now let's examine these statements.

Firstly, it's true that both diamond and graphite have the chemical formula of 'C'.

This formula indicates that they're both made up of entirely carbon atoms, which leads us to the next point that the second statement is also correct 'cause both diamond and graphite only contain carbon atoms. There's no other elements present in these structures.

Obviously, we're assuming we're talking about pure diamond and pure graphite.

In nature, we often find impurities through these structures.

Lastly, they are indeed both allotropes of carbon, meaning that while they are made of the same type of atom, carbon, the arrangement of these atoms is different in each substance, giving them distinct properties.

So all of the statements listed are true about diamond and graphite.

Each one reinforces the fact that these are two substances whilst different in structure and characteristics, they are both pure forms of carbon and represent different structural forms that elemental carbon can take.

The carbon atoms in diamond are covalently bonded to four other carbon atoms, forming a tetrahedral structure that is extremely strong and gives diamond its renowned hardness.

I've built a tetrahedron here to show the four corners representing the four atoms that the central carbon atom has bonded to.

So it's a bit easier to see when it is a 3D structure than on the 2D one that is on the screen in front of you.

So each of these corners here representing one of the four carbon atoms that the central inside this carbon atom would be bonded to.

It's important to note that diamond isn't made up of only four bonds.

Classic response that students give to this description.

But each atom of carbon within the structure has four covalent bonds to other carbon atoms. Graphite structure is quite different.

Each carbon atom is covalently bonded to three others, creating layers of hexagonal rings.

These layers can slide over each other because they are are only weak forces of attraction between them.

And this makes graphite slippery.

We'll cover more on properties later in this lesson.

Both diamond and graphite are considered to be highly ordered, giant covalent structures with a repeating pattern that extends in all three dimensions, diamond and in graphite with weak interlay forces of attraction holding the layers of graphite together.

Coal, however, is not considered a giant covalent structure like diamond and graphite.

It doesn't have a regular structure and is a complex mixture of different molecules containing carbon.

Coal is formed from the remains of ancient plant matter under high pressure and temperature over millions of years, resulting in a substance that does not have the crystalline regularity that defines a giant covalent structure like in diamond and graphite.

Coal is often unlikely to be pure.

It's mostly impure as well, and so it's unlikely you'll ever find a sample that we would class as a truly kind of natural allotrope of carbon.

Here's a moment for reflection on these structures then.

Take a look at the images provided and consider the question, which of these has a giant structure? Pause the video if you need some time to ponder this and press play when you are ready to move on.

Well, the correct answers are diamond and graphite.

Each carbon atom in diamond forms four covert bonds in a very strong tetrahedral arrangement resulting in a three dimensional network.

Graphite, on the other hand, has layers of hexagon arranged carbon atoms bonded in layered sheets.

These sheets are held together by weaker forces that allow them to slide over each other.

Coal, although primarily carbon, does not have a highly ordered, giant structure like diamond and graphite.

The arrangement of these atoms is not regular, and coal is actually a mixture of various different molecules containing carbon.

So it doesn't exhibit a single continuous network like the other two substances.

Coal structure is therefore more amorphous.

That means that it's without a clearly defined shape or form, so it does not fit the criteria for a giant structure.

I'd like you to read through the following statements and consider which you think are correct regarding diamond and graphite.

Pause the video if you need some time to think about this and when you are ready, press play to continue.

Now let's review these statements.

Graphite does not have a regular repeating structure.

Well, this is incorrect 'cause graphite does have a regular repeating structure with layers of carbon atoms arranged in a hexagonal pattern.

Graphite is made of layers of carbon atoms. Well, that's true.

In graphite, each carbon atom is bonded to three others, as we know, in that hexagonal pattern and they are stacked to form multiple layers.

Diamond is made up of hexagonal rings of carbon atoms. Well, that's not true.

Diamond is a tetrahedral structure where each carbon atom is bonded to four others in a three dimensional framework.

Not hexagonal rings.

That's graphite.

So therefore this last statement is correct.

Every carbon atom in diamond is bonded to four others.

Now it's over to you.

For some extended tasks.

You'll need to fill in the blanks with appropriate terms and define some key concepts related to carbon structures.

Pause the video whilst you complete this and press play when you're ready to move on.

So here are the answers.

We're filling in the with the words carbon, diamond and graphite.

So the atoms in diamond are arranged in a rigid, three-dimensional network where each carbon atom is covalently bonded to four other carbon atoms. In contrast, the atoms in graphite are bonded in flat, two-dimensional layers where each carbon atom is covalently bonded to only three others, forming hexagonal rings of carbon.

And the definition of allotrope is that they are different structural forms of elements such as in graphite and diamond.

And giant covalent structures are large regular arrangements of atoms all joined together by covalent bonds.

Well, if you managed to fill in all of those gaps correctly and get those definitions, you're doing great.

We're now going to look at how the structure of these giant covalent substances relates to their properties and uses.

Like all giant covalent structures, graphite has a high melting point, in this case approximately three and a half thousand to 3,600 degrees Celsius.

This high melting point is a direct consequence of the strong covalent bonds that hold the carbon atoms together.

A large amount of energy is needed to break these bonds between the carbon atoms and so a high temperature is required.

Graphite possesses a unique trait that distinguishes it from other giant covalent structures in that the layers are held together by weak forces of attraction.

These weak inter layer forces are relatively easy to break, which allows the layers to slide over each other with ease, making graphite quite soft.

This ability to slide apart with little resistance is precisely what makes graphite such an effective lubricant.

It can be used in applications where two services need to move smoothly past each other, such as in locks or in machinery.

The flake graphite shown demonstrates the usefulness of the layered structure of graphite.

Each flake is a stack of these layers and you can imagine how they might peel off or slide over each other providing lubricating properties.

So let's review the properties of graphite.

Take a moment to consider which statements accurately describe this substance.

Once you've made your selections, we can continue.

So pause here and resume when you're ready for the answers.

So the layers of graphite are indeed held together by weak forces of attraction.

This allows the layers to slide over each other contributing to graphite's lubricating properties.

Graphite does not have a low melting point.

It actually has quite a high melting point due to the strong covalent bonds between the atoms in each layer.

And so it does not require a lot of energy to overcome the forces between the layers of graphite, and that is why the layers can slide over each other very easily.

In graphite, each carbon atom has four electrons in the outermost shell.

Same for any other kind of neutral scenario where we find carbon.

And those electrons are ready to form bonds.

When carbon atoms bond within the graphite structure, they form three strong covalent bonds with adjacent carbon atoms, creating a flat hexagonal lattice.

This bonding uses up three of the four outer shell electrons in a carbon atom.

And the fourth electron, however, not being involved in the bonding, that makes it delocalised through the structure.

These delocalised electrons are free to move through the layers of graphite and their mobility is crucial to graphite's electrical conductivity.

Because each carbon atom contributes just one delocalised electron, the graphite structure as a whole has a large number of free moving electrons.

When graphite is part of an electrical circuit connected to a power supply, these delocalised electrons can flow, moving through the structure carrying charge and current and that's why we can class graphite as a good conductor of electricity.

This property is particularly useful in industrial applications such as in arc furnaces.

Here graphite tipped electrodes are used to pass electricity through metals providing the immense heat that is required to melt them.

The robust structure of graphite can withstand these high temperatures and its electrical conductivity makes it ideal for this.

This property of graphite is highlighted here in this clip where a graphite rod is being connected to a circuit.

Before it is connected, the light is not on.

But as soon as it is connected, the lights switch is on, showing graphite's ability to conduct electricity.

Here's a true or false statement for you.

Graphite conducts electricity as it has free moving charged particles.

What are your thoughts on this? Well, the answer is true.

Pause the video to read through the justifications below and press play when you are ready to move on.

Well, the reason why this statement is true is because the graphite has a structure that allows for one electron from each carbon atom to become delocalised.

Remember, electrons are negatively charged and so as these electrons are free to move across the layers, they can carry a charge.

So the properties of graphite are closely linked to its various uses.

Graphite has a very high melting point 'cause of those strong developed bonds.

This makes it suitable for refractory materials that are used to line kilns and furnaces where high temperatures are common.

The layers in graphite are held together by weak forces, which allows them to slide over each other very easily.

This property makes graphite useful as a lubricant in situations where traditional oil-based lubricants might not be appropriate.

Additionally, this quality is what makes graphite useful in pencils, allowing it to leave marks on paper as the layers slide off onto the surface.

The delocalised electrons in graphite allow it to conduct electricity, and that's why we use it in electrolysis.

So the decomposition of chemical substances using electricity.

So we use it as an electrode so the electricity can be conducted through this.

And we also use it in lithium ion batteries.

Probably use it in a few other types of batteries as well, but most commonly in lithium ion batteries 'cause it's a very good conductor.

Like with graphite, the structure of diamond accounts for its properties and uses.

Diamond's high melting point is a direct consequence of its strong covalent bonds, which form a rigid tetrahedral structure.

Each carbon atom is bonded to four others in a very compact structure, which requires a large amount of energy to break apart.

The same strong covalent bonding and rigid lattice structure gives diamond its high melting point, but also contributing to its extreme hardness.

Diamonds are among the hardest materials known and it's still sort of debated as to whether they are the hardest naturally occurring material.

We know they're not the hardest material known to man 'cause we've made some slightly stronger materials.

But we know diamonds can only be scratched by materials with the same hardness or greater.

This makes diamond some very useful uses in tools which we'll talk about in a little bit.

Since all of diamonds outer shell electrons are involved in bonding, there are no free electrons to carry charge.

This makes diamond an excellent insulator, which means that it does not conduct electricity.

The uses of diamond reflect on these properties.

So in jewellery, the aesthetic appeal of diamonds, especially when cut and polished to create facets that reflect light brilliantly makes them highly prized for jewellery.

The hardness of diamond also makes it very useful for tool tips, particularly for cutting, drilling or grinding hard materials.

Diamond tip tools are used in a wide variety of industries from construction and manufacturing to high tech fields.

I'd like you to read through the statements shown and decide which you think are true about diamond.

Pause the video whilst you do this and press play when you are ready to move on.

So the only correct statement here is the diamond has all atoms bonded strongly together, obviously with covalent bonds, and that's a lot of energy that is going to be needed to break them apart.

So this gives diamond a high melting point.

All the other statements here are false.

True or false, both diamond and graphite are giant covalent structures, so they can both be used for making tools.

What do you think? Well, this is false.

I'd like you to read through the justifications below to decide which best explains why.

Pause the video whilst you do this and press play when you're ready to move on.

So graphite is able to be used as a lubricant because the layers are held together by weak forces of attraction.

This is not the same as in diamond.

Diamond has all carbon atoms connected together in strong covalent bonds, making one structure, whereas graphite is made up of a layered structure.

Therefore graphite is quite soft in comparison to the hard nature of diamond.

So graphite is not able to be used for making tools.

We're now onto our final set of tasks for this lesson.

Here two students, Lucas and Sophia, are discussing diamond and graphite.

Pause the video to read through what they're saying and decide who's correct and also to decide why.

And I'd like you to also make corrections to any incorrect statements.

Pause video whilst you do that and press play when you are ready to move on.

So each student has one correct statement each.

They are that graphite is made up of layers of carbon atoms that can slide over each other and that diamond is a hard substance and it is used to cut other materials.

But we needed to change up the wording for two statements.

So we should have read only graphite rather than both graphite and diamond can conduct electricity as it has free delocalised electrons.

Diamond has no free moving charge particles, so it cannot conduct electricity.

Graphite has a high melting point, not a low melting point, due to all the carbon atoms being bonded together with strong covalent bonds that need large amounts of energy to break apart rather than the weak forces of attraction mentioned.

There are weak forces of attraction, but they're between the layers, not between the atoms. I've tried to give some extra explanation here, so hopefully this helps if you didn't quite manage to spot the errors.

Well done if you did.

The final task for this lesson then involve you completing this table and then explaining why these substances have these properties.

Pause the video as you do this and press play when you're ready to continue.

So diamond has low conductivity compared to graphite's high conductivity.

Graphite is soft and has a very high melting point, and diamond can be used for tool tips and for jewellery.

Diamond and graphite both have high melting points as they have strong covalent bonds holding all the atoms together.

So lots of energy needed to break these apart.

Diamond has a rigid network of carbon atoms making it a hard substance, but graphite is made of layers held with weak forces that slide over other, making it soft.

Graphite has delocalised electrons which can move and carry charge, so it is a good conductor of electricity.

Diamond does not have any free moving charge particles, so cannot conduct electricity.

So in summary then, in diamond, each carbon atom is bonded to four others with strong covalent bonds to form a giant covalent structure.

Diamond is very hard.

It has a very high melting point and does not conduct electricity, which makes it very good for use in cutting tools.

In graphite, each carbon atom is covalently bonded to three others to form layers of hexagonal rings.

There are only weak forces of attraction between the layers in graphite, which can easily be rubbed apart, which makes it a good lubricant.

Graphite conducts electricity because it has delocalised electrons which can move and carry charge and current.

Thank you for learning about diamond and graphite with me today.

I look forward to seeing you in the next lesson.