Diamond & Graphite (with exam predictions) - GCSE & IGCSE Chemistry Revision 2024

Summary

This video explores the properties and uses of diamond and graphite, two allotropes of carbon. Diamond is a hard material with a high melting point, featuring strong covalent bonds where each carbon atom is bonded to four others. It does not conduct electricity due to the lack of free electrons. In contrast, graphite has a layered structure with strong covalent bonds within the layers and weak forces between them, allowing delocalized electrons to move freely and conduct electricity. The video also highlights their practical uses in jewelry, drilling, and as electrodes in electrolysis.

Takeaways

• 😀 Diamond and graphite are both allotropes of carbon, meaning they are different structural forms of the same element.
• 😀 Diamond has a giant covalent structure with strong covalent bonds between carbon atoms, making it hard and unable to conduct electricity.
• 😀 In diamond, each carbon atom is bonded to four other carbon atoms, creating a rigid, non-conductive structure.
• 😀 Graphite, another allotrope of carbon, has a giant covalent structure but consists of layers of carbon atoms bonded to three others.
• 😀 Graphite's layers are held together by weak intermolecular forces, which make graphite soft and slippery.
• 😀 Graphite has delocalized electrons between its layers, which allow it to conduct electricity.
• 😀 Diamond is used in applications like jewelry, drilling, cutting, and engraving due to its hardness.
• 😀 Graphite is used for electrodes in electrolysis and as the material in pencils because of its slippery properties.
• 😀 To explain the hardness of diamond, you need to mention that diamond’s structure consists of strong covalent bonds that require a lot of energy to break.
• 😀 To explain why graphite conducts electricity, it’s important to mention delocalized electrons between the layers of carbon atoms that can move and carry a charge.
• 😀 It's essential to remember that the number of covalent bonds in each carbon atom and the presence of free electrons are key factors in distinguishing diamond from graphite.

Q & A

What is an allotrope?

An allotrope refers to one or more different forms of a chemical element. These forms have different bonding arrangements, which lead to distinct physical properties.

How is diamond structured and why is it so hard?

Diamond has a giant covalent structure where each carbon atom is bonded to four other carbon atoms. The strong covalent bonds throughout the structure require a lot of energy to break, which makes diamond very hard.

Why doesn't diamond conduct electricity?

Diamond doesn't conduct electricity because there are no free electrons in its structure. Each carbon atom is bonded to four others, leaving no delocalized electrons to carry a charge.

What are some common uses of diamond, and why are they suitable?

Diamond is used in jewelry, drilling, cutting, engraving, and as scalpel edges. Its hardness makes it ideal for cutting and engraving, while its lack of electrical conductivity is suitable for non-electrical applications.

How does the structure of graphite differ from that of diamond?

In graphite, each carbon atom is bonded to three other carbon atoms, forming layers of hexagons. Unlike diamond, graphite has weak intermolecular forces between layers, making it soft and slippery.

Why is graphite able to conduct electricity?

Graphite can conduct electricity because each carbon atom in its structure has one free electron that is delocalized. These delocalized electrons can move between the layers and carry an electrical charge.

What is the significance of delocalized electrons in graphite's structure?

Delocalized electrons in graphite play a key role in its ability to conduct electricity. These electrons are free to move between the layers, enabling the flow of current through the structure.

Why is graphite slippery, and what does this mean for its uses?

Graphite is slippery because of the weak intermolecular forces between its layers. This allows the layers to slide over each other easily, which makes graphite useful as a lubricant and for pencil leads.

What is the melting point of diamond and graphite, and how do they compare?

Both diamond and graphite have high melting points due to their strong covalent bonds. However, diamond’s structure makes it harder and more rigid, while graphite's layered structure makes it less rigid but still maintains a high melting point.

What are some key properties of diamond and graphite that make them useful for different purposes?

Diamond is hard, has a high melting point, and does not conduct electricity, making it suitable for cutting tools and jewelry. Graphite, on the other hand, is soft, slippery, has a high melting point, and conducts electricity, making it ideal for use in pencils and electrodes.