Graphite is unique among non-metals — it is the only non-metal that conducts electricity. This unusual property makes graphite invaluable in industries ranging from battery manufacturing to electrodes in electrolysis.
Graphite electrodes are used in the electrolytic refining of metals (like copper) and in dry cell batteries as the central carbon rod (the positive electrode/cathode).
Graphite has a distinctive layered (lamellar) structure. Each carbon atom is bonded to three other carbon atoms in a flat, hexagonal pattern, forming sheets (layers). This is called sp² hybridization.
Because each carbon forms only 3 bonds (not 4), one electron per carbon atom remains unhybridized and delocalized across the entire layer as a π (pi) electron cloud. These delocalized electrons can move freely throughout the layer — just like electrons in a metal. It is this free movement of electrons that allows graphite to conduct electricity.
Diamond, the other common allotrope of carbon, has each carbon atom bonded to 4 carbon atoms (sp³ hybridization, tetrahedral structure). There are no free/delocalized electrons at all. With no free electrons to carry charge, diamond is an excellent electrical insulator.
Yes! Graphite is also a good conductor of heat, again due to the delocalized electrons and vibrations within its layers. This is why graphite is sometimes used in high-temperature applications.
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