Network solids are a special type of material where atoms are connected in a continuous network. This structure gives them amazing strength and stability. Two famous examples of network solids are diamond and graphite. Both are made entirely of carbon, but they have very different properties because of how their atoms are bonded together.
Diamonds are known for being incredibly hard, which is why they can cut through almost anything. This hardness comes from the way carbon atoms are bonded in a three-dimensional network with strong covalent bonds. In diamonds, each carbon atom forms a tetrahedral shape using SP3 hybridized orbitals. This means the bonds are evenly distributed in all directions, making diamonds very stable and resistant to stress. However, this also makes them brittle, so they can break if hit hard enough.
Graphite, on the other hand, has a layered structure. In graphite, carbon atoms are bonded in flat sheets using SP2 hybridized orbitals, forming a trigonal arrangement. These sheets have strong pi bonds, which help graphite handle pressure in two dimensions. But the sheets are only loosely held together by weak van der Waals forces, allowing them to slide over each other easily. This makes graphite a great lubricant and perfect for pencil lead, as the layers can easily transfer onto paper.
The differences between diamond and graphite show how important atomic bonding and arrangement are in determining a material’s properties. Even though both are pure carbon, their bonds give them very different characteristics:
Under extreme conditions, graphite can be turned into diamonds. This happens when graphite is exposed to high temperatures (around $3000$ degrees Celsius) and immense pressure (about $15$ million kPa). These conditions rearrange graphite’s two-dimensional sheets into a three-dimensional network. This transformation requires a lot of energy to break and reform bonds, showing how pressure can lead to stabilization and change.
Interestingly, the reverse process—turning diamonds into graphite—is nearly impossible under natural conditions on Earth. The energy needed to break the covalent bonds in diamonds is so high that this process almost never happens, making diamonds very stable and long-lasting.
The different properties of diamonds and graphite highlight the power of chemistry. Even though they are made of the same element, their atomic arrangements and bonding create unique physical, chemical, and electrical properties. Diamonds are valued for their hardness and beauty, while graphite is useful for writing and lubrication. This exploration of network solids shows the fascinating relationship between atomic structure and material properties, revealing the wonders of chemistry.
Using molecular model kits, build the structures of diamond and graphite. Pay attention to the differences in bonding and geometry. Discuss with your classmates how these structural differences affect the properties of each material.
Conduct an experiment to test the electrical conductivity of graphite. Use a simple circuit with a battery, wires, and a light bulb. Replace one of the wires with a piece of graphite and observe what happens. Discuss why diamond would not work in this setup.
Watch a demonstration or video showing the thermal conductivity of diamond and graphite. Discuss why diamond is a better thermal conductor than graphite, considering their atomic structures.
Use a computer simulation to explore how graphite can be transformed into diamond under high pressure and temperature. Adjust the parameters to see how they affect the transformation process. Discuss the energy changes involved in this transformation.
Write a short story from the perspective of a carbon atom transitioning from graphite to diamond. Describe the conditions and changes in bonding that the atom experiences. Share your story with the class and discuss the scientific concepts illustrated in your narrative.
Network – A chemical structure where atoms are bonded in a continuous, repeating pattern throughout the material. – In a network solid, such as silicon dioxide, the atoms are bonded in a three-dimensional lattice.
Solids – A state of matter characterized by particles arranged in a fixed structure, giving it a definite shape and volume. – Metals are examples of solids that have a crystalline structure, allowing them to conduct electricity.
Carbon – A nonmetallic element found in various forms, known for its ability to form a vast number of compounds. – Carbon is essential in organic chemistry, forming the backbone of many molecules, including hydrocarbons.
Diamond – A form of carbon where each atom is bonded to four others in a tetrahedral structure, making it extremely hard. – Diamond is used in cutting tools due to its hardness, which results from its strong covalent bonding.
Graphite – A form of carbon where atoms are arranged in layers, allowing them to slide over each other, making it a good lubricant. – Graphite’s layered structure gives it excellent conductivity and is used in electrodes.
Bonding – The process by which atoms or molecules are held together in a compound through attractive forces. – Covalent bonding involves the sharing of electrons between atoms, as seen in a water molecule.
Properties – Characteristics of a substance that define its physical and chemical behavior. – The properties of metals, such as malleability and ductility, are due to the metallic bonding between atoms.
Conductivity – The ability of a material to allow the flow of electric current or heat through it. – Copper is widely used in electrical wiring because of its high electrical conductivity.
Structure – The arrangement of atoms or molecules within a substance. – The crystalline structure of salt is responsible for its cubic shape and brittleness.
Atoms – The basic units of matter, consisting of a nucleus surrounded by electrons. – Atoms combine to form molecules, which are the building blocks of all substances.
