Chapter 22: Problem 166
Why is diamond extremely hard and high melting, whereas graphite is very soft and high melting?
Short Answer
Expert verified
Diamond's hardness comes from its strong 3D covalent bonding, while graphite is soft due to weak forces between its carbon layers.
Step by step solution
01
Understand the Structure of Diamond
Diamond is composed of carbon atoms arranged in a tetrahedral lattice. Each carbon atom is covalently bonded to four other carbon atoms, forming a strong, three-dimensional network.
02
Recognize the Bonding in Diamond
In diamond, the strong covalent bonds between carbon atoms result in each atom being held very rigidly in place, contributing to diamond's notable hardness and high melting point.
03
Understand the Structure of Graphite
Graphite is also made of carbon atoms, but they are arranged in hexagonal sheets. Within each sheet, carbon atoms are bonded strongly; however, the sheets are held together by weaker van der Waals forces.
04
Recognize the Bonding in Graphite
The strong covalent bonds within the layers of graphite contribute to its high melting point, while the weak van der Waals forces between layers account for its softness, as the layers can easily slide over one another.
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Key Concepts
These are the key concepts you need to understand to accurately answer the question.
Covalent Bonding
Covalent bonding is a type of chemical bond where two atoms share one or more pairs of electrons. This type of bond is particularly strong and provides stability to molecules. In the world of chemistry, covalent bonds are crucial for creating complex structures like diamond and graphite. Both these materials are made up of carbon atoms, and it's the covalent bonds that give them distinct properties.
- In diamond, each carbon atom is covalently bonded to four others, forming a strong 3D network. This makes diamond exceptionally hard.
- In graphite, each carbon atom is covalently bonded within layers, creating sheets. These strong intralayer bonds give graphite a high melting point but allow flexibility within the structure.
Carbon Structures
Carbon is a versatile element capable of forming a variety of structures due to its ability to form four covalent bonds. The two most well-known carbon structures are diamond and graphite. Understanding these structures helps explain why their physical properties differ so dramatically.
In diamond, carbon atoms are organized into a tetrahedral lattice, where each carbon forms a covalent bond with four other carbons. This 3D framework results in an incredibly strong and rigid structure, making diamond extremely hard and resistant to melting.
In diamond, carbon atoms are organized into a tetrahedral lattice, where each carbon forms a covalent bond with four other carbons. This 3D framework results in an incredibly strong and rigid structure, making diamond extremely hard and resistant to melting.
- Each carbon in diamond shares electrons equally, maintaining a fixed position in the network.
- The arrangement in graphite allows flexibility and softness, with movements primarily within layers.
Van der Waals Forces
Van der Waals forces are weak attractions between molecules or parts of molecules that result from temporary differences in electron density. Though these forces are much weaker than covalent bonds, they play an essential role in the properties of materials like graphite.
In graphite, the covalent bonds within each layer are very strong, but the layers themselves are held together by van der Waals forces. These weak intermolecular forces allow the layers to slide over one another, giving graphite its characteristic softness and lubricating properties.
In graphite, the covalent bonds within each layer are very strong, but the layers themselves are held together by van der Waals forces. These weak intermolecular forces allow the layers to slide over one another, giving graphite its characteristic softness and lubricating properties.
- Van der Waals forces are responsible for the ease with which graphite layers can move.
- These forces allow graphite to be an effective lubricant and a good conductor of electricity, as electrons can flow more freely between layers.