Diamond vs graphite by a conductivity?

In summary, the difference in the type of hybridization of atomic orbitals and bonding in diamond and graphite is what causes diamond to be electrically insulating and graphite to be a good conductor. The molecular orbital diagrams for graphite show that its p orbitals overlap to form pi-bonds, creating an unfilled conduction band with free electrons. This is why graphite is able to conduct electricity while diamond is not.
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hermtm2
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Diamond vs graphite by a conductivity?

Hello. This is the last question on HW.

Q) Diamond and graphite are both made of carbon atoms. Explain why diamond is elecrically insulationg while graphite is a good conductor in terms of type(s) of hybridization of atomic orbitals and bonds in each material. Draw relevant molecular orbital (MO) diagrams for two canbons forming bonds fro each material and schemetically show these MOs leading to energy bands. (Consider only p-orbitals that form pi bonding, anti-bonding, or non-bonding MOs ---Only graphite)


I knew that no one electron is free to pass an electric current through the Diamond, since each atom's valence electron is taken up in covalent bonding. Other than that I have no clue... Is there MOs for a dimond or graphite? Are they consisted of only carbon, right? Please help me out.


Thanks,

Ryan.
 
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  • #2


In a nutshell, diamond is sp3 hybrid and graphite is sp2 hybrid bonded. p orbitals perpendicular to the plane overlap to form pi-bonds between the layers, effectively an unfilled conduction band
 

What is the difference between diamond and graphite in terms of electrical conductivity?

Diamond and graphite are both forms of carbon, but they have different molecular structures. Diamond has a three-dimensional lattice structure while graphite has a layered structure. This means that diamond is a poor conductor of electricity, while graphite is a good conductor.

Why is diamond a poor conductor of electricity?

As mentioned before, diamond has a three-dimensional lattice structure. This means that all of the valence electrons in diamond are tightly bonded to the carbon atoms and cannot move freely to conduct electricity.

How does the layered structure of graphite make it a good conductor?

In graphite, each layer is made up of hexagonally arranged carbon atoms. The electrons in these layers are only weakly bonded to the carbon atoms, allowing them to move freely and conduct electricity.

Can diamond be made into a good conductor of electricity?

It is possible to make diamond into a good conductor by introducing impurities, a process known as doping. However, this would alter the properties of diamond and it would no longer be a pure form of carbon.

Which form of carbon is better for electrical applications, diamond or graphite?

It depends on the specific application. For most electronic devices, graphite is preferred due to its high electrical conductivity. However, diamond has some unique properties that make it useful for specialized applications, such as its extreme hardness and ability to withstand high temperatures.

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