Network Solids and Carbon: Crash Course Chemistry #34

Summary

This Crash Course Chemistry episode explores the contrasting properties of diamond and graphite, both composed of carbon atoms. Despite being the same element, their atomic arrangements result in vastly different characteristics. Diamonds, with a three-dimensional network of covalent bonds, are incredibly hard and thermally conductive but brittle. Graphite, with a two-dimensional sheet structure, is soft, slippery, and a good lubricant, but only conducts electricity due to its pi bonds. The video also touches on the transformation of graphite into diamonds under extreme heat and pressure, a process not feasible at home, highlighting the influence of chemical bonding on material properties.

Takeaways

• 💎 Diamonds are the hardest natural material on Earth due to their unique three-dimensional network structure of carbon atoms bonded together with covalent bonds.
• ✏️ Graphite, used in pencil leads, is soft because of its two-dimensional sheet-like structure where carbon atoms are bonded in a hexagonal pattern with weak van der Waals forces between layers.
• 🔗 The difference between diamond and graphite lies in the arrangement of carbon atoms: diamonds have a 3D tetrahedral structure, while graphite has a 2D hexagonal structure.
• 🔗 The strength and stability of network solids come from the way atoms are linked in a network, which can be a chain, sheet, or three-dimensional structure.
• 🔋 Graphite can conduct electricity because of the free movement of electrons in its pi bonds, unlike diamond, which is an electrical insulator due to its rigid sigma bonds.
• 🔥 Both diamond and graphite are good conductors of heat because of the strong covalent bonds that facilitate the rapid spread of thermal energy.
• 💥 Graphite can be transformed into diamond under extreme heat (around 3000 degrees Celsius) and pressure (about 15 million kPa), which rearranges its atomic structure into a 3D network.
• 🚫 Turning diamonds into graphite is extremely difficult and practically impossible under natural conditions on Earth due to the high activation energy required.
• 🌐 The properties of materials are significantly influenced by the hybridization of atomic orbitals, which dictates how electrons are arranged and bonded within the structure.
• 🎓 This Crash Course Chemistry episode highlights the importance of understanding chemical bonding and network structures in determining the physical, chemical, and electrical characteristics of materials.

Q & A

What makes diamond the hardest natural material on Earth?

Diamond is the hardest natural material on Earth because its carbon atoms are arranged in a three-dimensional network structure, forming covalent bonds in every direction, which resists forces very effectively.

Why is graphite soft despite being made of the same element as diamond?

Graphite is soft because its carbon atoms are arranged in a two-dimensional sheet network with weak van der Waals forces holding the sheets together, allowing them to slide over each other easily.

How does the atomic arrangement in network solids affect their properties?

The atomic arrangement in network solids significantly affects their properties. For instance, the three-dimensional network of diamond makes it extremely hard, while the two-dimensional network of graphite makes it soft and slippery.

What is the role of hybridized orbitals in determining the properties of network solids like diamond and graphite?

Hybridized orbitals play a crucial role in determining the properties of network solids. In diamond, SP3 hybridization creates a rigid three-dimensional structure, while in graphite, SP2 hybridization forms a two-dimensional sheet with pi bonds for added strength within the sheets.

Why does graphite conduct electricity while diamond does not?

Graphite conducts electricity because the pi bonds in its two-dimensional sheets allow electrons to move freely between carbon atoms. In contrast, diamond's carbon atoms share sigma bonds, which restrict electron movement, making it an electrical insulator.

How can graphite be transformed into diamond under extreme conditions?

Graphite can be transformed into diamond under extreme conditions of high temperature (around 3000 degrees Celsius) and pressure (about 15 million kPa), which provide the energy for breaking and reforming bonds, allowing the atoms to reorganize into a three-dimensional network.

What is the significance of covalent bonds in the thermal conductivity of diamond and graphite?

Covalent bonds are significant for the thermal conductivity of diamond and graphite because they require that when one atom vibrates, all surrounding atoms vibrate in the same way, allowing thermal energy to spread quickly. Diamond conducts heat better than graphite due to its three-dimensional network structure.

Why is diamond brittle despite its hardness?

Diamond is brittle because its rigid three-dimensional network structure does not allow for much flexibility. When stress is applied, it tends to break along cleavage planes rather than bending.

How do the properties of diamond and graphite demonstrate the power of chemistry?

The properties of diamond and graphite demonstrate the power of chemistry by showing how the same material, carbon, can have vastly different characteristics based on the nature and arrangement of its chemical bonds.

What is the activation energy required to break the covalent bonds in diamond, and why is it significant?

The activation energy required to break the covalent bonds in diamond is so high that it makes the breakdown of diamond essentially impossible under normal Earth conditions. This is significant because it contributes to diamond's stability and resistance to change.

Outlines

00:00

💎 The Chemistry of Diamond and Graphite

This paragraph explores the contrasting properties of diamond and graphite, both composed of carbon but with vastly different atomic structures. Diamonds, being the hardest natural material, have a three-dimensional network structure where each carbon atom forms covalent bonds with four others, creating a rigid and strong lattice. This structure makes diamonds resistant to pressure and suitable for cutting and drilling applications. On the other hand, graphite's carbon atoms are arranged in a two-dimensional sheet structure with covalent bonds within the sheets and weak van der Waals forces between them, making it soft, slippery, and useful as a lubricant and in pencil leads. The paragraph also touches on the concept of network solids, explaining how the bonding and arrangement of atoms in these solids contribute to their unique properties.

05:00

🔥 Transforming Graphite into Diamonds

The second paragraph delves into the transformation of graphite into diamonds, a process that requires extreme heat and pressure. At around 3000 degrees Celsius and under approximately 15 million kPa of pressure, graphite's two-dimensional sheet network can rearrange into a three-dimensional network, similar to that of a diamond. This transformation is not feasible under normal conditions but is possible in industrial settings. The paragraph also contrasts the electrical conductivity of graphite, which allows electrons to move freely due to its pi bonds, with the insulating properties of diamonds, where sigma bonds restrict electron movement. Both materials are good thermal conductors due to the strong covalent bonds that facilitate the rapid spread of thermal energy. The discussion concludes with a philosophical reflection on the differing values and uses of diamonds and graphite, suggesting that while diamonds are more precious, graphite's versatility and utility might have had a broader positive impact on the world.

Keywords

💡Covalent Bonds

Covalent bonds are a type of chemical bond formed when two atoms share one or more pairs of electrons. In the context of the video, covalent bonds are crucial in determining the properties of materials like diamond and graphite, both of which are composed of carbon atoms bonded together by covalent bonds. The difference in their hardness and structure is due to the way these bonds are arranged: in diamond, a three-dimensional network makes it extremely hard, while in graphite, a two-dimensional sheet structure makes it soft and slippery.

💡Network Solids

Network solids are a category of atomic solids where atoms are bonded together in a network structure, creating a continuous framework. The video explains that network solids are characterized by their strength and stability due to the interconnectedness of their atoms. Diamond and graphite are both network solids made of carbon, but they exhibit vastly different properties because of the dimensionality of their atomic networks: diamond has a 3D network, while graphite has a 2D network.

💡SP2 Hybridization

SP2 hybridization refers to a type of atomic orbital hybridization where one s orbital and two p orbitals combine to form three sp2 orbitals. In the video, this concept is used to describe the structure of graphite, where carbon atoms are arranged in a flat, hexagonal structure due to sp2 hybridization, leading to the formation of strong, two-dimensional sheets that can slide over each other.

💡SP3 Hybridization

SP3 hybridization is a type of atomic orbital hybridization where one s orbital and three p orbitals combine to form four sp3 orbitals. The video uses this term to explain the structure of diamond, where each carbon atom forms four covalent bonds with four other carbon atoms, creating a rigid, three-dimensional network that accounts for diamond's extreme hardness.

💡Graphite

Graphite is a form of carbon where the atoms are arranged in a two-dimensional sheet structure held together by covalent bonds. The video highlights graphite's softness and its use in pencils, as well as its ability to conduct electricity due to the delocalized electrons that can move freely within the sheets. Graphite's lubricating properties are also mentioned, which are due to its layered structure allowing the sheets to slide over each other.

💡Diamond

Diamond is another allotrope of carbon, where the atoms are arranged in a three-dimensional network structure. The video emphasizes diamond's hardness, which is the result of its strong covalent bonds in all directions. It also mentions that diamonds are electrical insulators because the electrons are not free to move as they are in graphite, which has a different bonding arrangement.

💡Van der Waals Forces

Van der Waals forces are weak intermolecular forces that occur between molecules. In the context of the video, these forces are what hold the layers of graphite together. They are much weaker than the covalent bonds within the layers, which is why graphite can be easily cleaved into thin sheets, a property that makes it useful in pencils and as a lubricant.

💡Pi Bonds

Pi bonds are a type of covalent bond where the electron density is distributed above and below the plane of the atoms involved. The video explains that in graphite, pi bonds form between layers, contributing to the strength within each layer. However, these bonds also allow electrons to move freely, which is why graphite can conduct electricity.

💡Sigma Bonds

Sigma bonds are the strongest type of covalent bond where the electron density is concentrated between the nuclei of the bonding atoms. The video points out that in diamond, carbon atoms form sigma bonds with each other, creating a very rigid and strong three-dimensional network. This structure does not allow for electron movement, which is why diamond does not conduct electricity.

💡Hybridization

Hybridization in chemistry refers to the concept where atomic orbitals combine to form new hybrid orbitals, which can then participate in bonding. The video uses the term to describe how the arrangement of electrons in carbon atoms leads to different structures in graphite and diamond, with sp2 and sp3 hybridization respectively, resulting in their distinct properties.

💡Thermal Conductivity

Thermal conductivity is the ability of a material to conduct heat. The video explains that both diamond and graphite are good thermal conductors due to the strength of their covalent bonds, which allow vibrations to be transmitted efficiently. Diamond is noted to be a better conductor of heat than graphite because it can transfer heat in three dimensions, unlike graphite which only does so in two.

Highlights

Diamond's hardness comes from its unique carbon atom arrangement in a three-dimensional network.

Graphite's softness is due to its two-dimensional sheet-like structure of carbon atoms.

Different atomic networks result in different material properties, despite being composed of the same element.

Chemical bonding and the resulting network structures are crucial for material applications.

Glycogen is an example of a chain network solid, but it lacks durability.

Proteins in silk form a two-dimensional network that is stronger within the plane but weak between planes.

Three-dimensional networks provide strength and stability through multiple directional covalent bonds.

The way atoms bond in network solids significantly influences their properties.

Diamond's hardness is due to its uniform, three-dimensional carbon network resisting stress.

Graphite's layers slide over each other due to weak van der Waals forces, making it useful for writing and lubrication.

Graphite conducts electricity because of the free movement of electrons in its pi bonds.

Diamond's rigid structure makes it brittle and prone to breaking along cleavage planes.

Diamonds are electrical insulators due to the lack of free electron movement in sigma bonds.

Both diamond and graphite are good thermal conductors because of their strong covalent bonds.

Graphite can be transformed into diamond under extreme heat and pressure.

The transformation of diamond into graphite is extremely difficult and not naturally occurring on Earth.

The differences between diamond and graphite demonstrate the power of chemistry in determining material properties.

Despite being made of the same material, the arrangement of atoms in diamond and graphite leads to vastly different characteristics.

Graphite's versatility, including its use in pencils, might have had a more significant impact on the world than diamonds.

Transcripts

The diamond in this ring will cut into any substance on Earth.

It's the hardest, natural material on the planet.

But it's just carbon, the same element that makes up the graphite in this pencil lead,

which is so soft, that it's intended purpose is to rub off on stuff.

The atoms are all of the same element and they're all bonded together with covalent bonds,

so what's the difference?

Well the atoms are arranged differently, they form different atomic networks and therefore, different network solids.

In chemistry we place networks in various categories

because each type of network has characteristics that make them useful for all sorts of handy applications.

You can have a chain of identical molecules, each molecule is connected on two sides.

So glycogen which is just a chain of glucose molecules is bigger than glucose

but it's not super durable as substances go.

A sheet of particles like the proteins that make up silk is a more sophisticated kind of network

because the particles are connected on all sides,

but its connections are still two-dimensional so it's only stronger in those directions.

Now, you can stack sheets on top of each other in order to build a network into the third dimension,

but then the sheets just slide around on each other

and they're not much more useful or durable than they were before.

A truly three-dimensional network branches out in multiple directions,

forming covalent bonds in a structure that resist forces much better than a chain or a sheet of the same material.

So in the most basic sense, the way atoms and molecules bond to form network is what make many materials what they are.

It explains why diamond is so hard while graphite is so soft

and how the main ingredient in sand can also be made into the silicon wafers that make our electronics.

It's also the key to how we can turn this, into this.

Though before you send the kiss off email to your student loan officer,

I should warn you that you won't be able to do that at home.

[Theme Music]

Network solids are one of the three types of atomic solid we've talked about,

materials that are made of individual atoms rather than molecules or ions.

The cool and special thing about network solid is that the atoms exist in, you guessed it, a network structure,

that is each atoms is linked to several others in various directions.

Like all networks, this makes the arrangement both strong and stable.

It also provide some other properties that may surprise you.

If you've seen our episode on atomic orbitals you know the way the electron in a chemical bond orient themselves,

that is the way the orbitals hybridize, has a big effect on the properties of a substance.

If water's orbitals hybridize differently for example, it might not be polar and life as we know it wouldn't be possible.

It's the same with network solids. Orbitals and their hybridizations make a huge difference.

Thankfully though, orbitals can't be different. They are the way they are because of the laws of physics.

So first, let's look at carbon. Pure carbon can bond with itself in two different ways.

In the first scenario, the outer electrons of each atom are arranged into SP2 hybridized orbitals

with two other atoms for a total of three lobes in a flat trigonal arrangement.

This creates a sheet with a hexagonal structure and with one unhybridized P orbital left on each atom.

These orbitals form pi bonds that merge into an extensive network on their own

and this gives the structure its real strength.

You know this carbon network by its street name, graphite.

The pi bonds within each sheet is really strong which allows graphite to withstand a lot of pressure.

But as with any sheet-like formation, only in two dimensions.

The sheets are only held to other sheets with really weak van der Waals forces.

They allow the sheets to slide on top of each other so they can easily be removed layer by layer.

That's why graphite is so fantastic for writing with.

When you use a pencil, layer after layer of graphite is being transferred to the paper.

That slipping sheet structure also makes graphite an excellent lubricant.

It can break down into an extremely fine powder of slippery platelets that almost any substance can slide easily over.

Which is why it's great for things like sticky locks.

And because the electrons that form the pi bonds in the graphite are able to move from one atom to another,

graphite also conducts electricity.

Most of the properties of graphite, you'll notice, don't apply to the other type of network that carbon can form, diamond.

Even though the graphite in the pencil is made of the exact same thing as diamond, pure carbon atoms,

they're different from each other in almost every possible way.

And all of these differences are simply because of the way their atoms are bonded.

In diamond, the bonding electrons in each atom are arranged in SP3 orbitals.

Four lobes that are as far away from each other as possible.

Again, information you should have in you from the orbital episode.

Each lobe in that tetrahedral structure overlaps and bonds with one on an adjacent carbon atom

creating a totally uniform, three dimensional network of carbon in every direction.

So any stress on this network will be resisted by multiple atoms and multiple bonds.

This stability is what makes diamonds so famously, incredibly hard.

But the downside is that this rigid structure also makes them quite brittle.

In other words, they break before they bend.

And like any crystal and solid they tend to break along the seam between atoms called a cleavage plane.

And these are the places that the diamond cutters attack first.

And here's another way that the diamond's network structure makes it totally different from graphite.

Diamonds are electrical insulators, not conductors.

This is because the carbon atoms in diamonds share sigma bonds, not pi bonds,

which doesn't allow the electrons much freedom of movement.

And if an electron can't move around, it can't be used to transfer energy in the form of electricity or anything else.

But even though diamond is terrible at conducting electricity,

both graphite and diamond are good conductors of heat.

That's because of the strength of their covalent bonds.

These type bonds require that when one atom vibrates, all the atoms around it vibrate

in exactly the same way.

So when thermal energy increases in one region, it spreads quickly throughout the diamond or graphite.

Even though there is some space between the sheets of graphite,

the structure of the sheets is so rigid that the vibrations transfer through the whole mass.

And as you may have guessed, diamond conducts heat even better than graphite

because each atom spreads that energy in three directions rather than two.

Graphite has another interesting response to heat.

When it's heated to very high temperatures, around 3000 degrees Celsius, and placed under extreme pressure,

about 15 million kPa, the atoms rearrange themselves from a two dimensional sheet network into a 3D network.

Now that's roughly the amount of pressure of 15 metric tons on one square centimeter.

Basically five elephants sitting on the heel of a high-heeled shoe.

That tremendous heat provides energy for the breaking and reforming of bonds,

which allows the atoms to reorganize themselves to relieve all of that pressure.

And that my friends is how you can turn common graphite into diamonds.

So no, you cannot do that at home.

But just like with you and me sometimes what seems like an excessive amount of pressure

can end up forcing changes that stabilize things in the long run.

Now it seems kinda odd but the opposite reaction, diamonds turning into graphite, is even more difficult.

So much so that it's functionally impossible for it to happen naturally. At least on this planet.

Even though it is possible under those incredibly hot elephants on heel circumstances to break

the covalent bonds in graphite,

the activation energy required to break those bonds in diamond is so high

that the rate of its breakdown is essentially zero for all earthly purposes.

The similarities and especially the differences between diamonds and graphite

are an amazing display of the power of chemistry.

Even though they're both made of the exact same material,

the mere difference in the nature and arrangement of their chemical bonds completely changes their physical

and chemical and electrical characteristics.

Also, one is pretty and one is pencil lead.

They explain why diamond can be used to cut through glass and drill through rock,

while graphite is soft enough to write and lubricate with.

And yes, diamonds are more valuable,

but you could argue that graphite with its easily erasable mass producible pencils

might have done more good for the world.

The pencil is maybe mightier than the diamond.

Thank you for watching this episode of Crash Course Chemistry, by the way.

If you paid attention, you learned that networks lend strength and widely distributed stability to network solids.

You learned that both diamonds and graphite are network solids made up of pure carbon atoms

but that the arrangement of those atoms in 2 and 3 dimensions respectively give them completely different properties.

This episode of Crash Course was written by Edi Gonzales, edited by Blake de Pastino.

Our chemistry consultant is Dr. Heiko Langner. It is filmed, edited, and directed by Nicholas Jenkins.

The script supervisor was Michael Aranda, who is also our sound designer,

and our graphics team, of course, is Thought Café.