Diamond and graphite are both allotropes of carbon. However only graphite can be used as an electrode explain why this is so even though they are both forms of carbon?
Diamond is a tightly bound molecule of carbon atoms. Highly covalent bonds. Graphite, however, occurs in layers and electrons can be found in the areas between layers. These de-localized electrons are responsible for the electrical conductivity of graphite. Diamond is not a layered material.
Here is a site about graphite.
http://en.wikipedia.org/wiki/Graphite
Diamond and graphite are both allotropes of carbon, meaning they are different structural forms of the same element. While they are composed of the same carbon atoms, their different arrangements create distinct properties.
Graphite is made up of layers of carbon atoms arranged in a hexagonal lattice, with weak van der Waals forces between the layers. This layered structure allows graphite to have conductivity. The delocalized electrons within the layers can move freely along the planes, making graphite a good conductor of electricity.
On the other hand, diamond is formed by a three-dimensional network of carbon atoms, with each carbon atom covalently bonded to four neighboring carbon atoms. This rigid and dense structure of diamond does not allow for the movement of electrons and, thus, does not conduct electricity effectively.
To understand why graphite can be utilized as an electrode, we need to delve into the concept of electrodes. In simple terms, an electrode is a conductor that carries electric current into or out of a substance during an electrochemical reaction. In various applications, such as batteries or electrolysis, electric current needs to flow through the electrode.
Due to the layered structure of graphite and the presence of delocalized electrons within each layer, graphite can easily conduct electricity through these layers. As a result, graphite is commonly used as an electrode material since it allows the flow of electric current, making it suitable for electrochemical processes.
In summary, although both diamond and graphite are forms of carbon, their different atomic arrangements lead to varied properties. Graphite's layered structure and the presence of delocalized electrons enable it to conduct electricity, making it suitable for electrode applications, while diamond's dense and rigid structure restricts the movement of electrons and limits its electrical conductivity.