Difference between Diamond and Graphite

diamond and graphite
Carbon exists in three allotropic forms as diamonds, graphite, and fullerene. Diamond and Graphite are crystalline in nature. In this article, we are going to discuss the difference between diamond and graphite

Difference between Diamond and Graphite


1. In diamond, each carbon atom is sp3 hybridized.

2. It is colorless and transparent in nature.

3. Each carbon atom is joined to four other carbon atoms in a tetrahedral manner. This forms the three-dimensional structure of the diamond.

4. The carbon-carbon bond length is 1.54 Å

5. The bond between two carbon atoms in very strong and stable. That’s why diamond is very hard.

6. The density of diamond is very high because carbon atoms are strongly held together.

7. All electrons are involved in bonding, diamond is a bad conductor of heat and electricity. ( free electrons means electricity).


1. In graphite, each carbon atom is sp2 hybridized.

2. Graphite shows color from black to grey.

3. Each carbon atom is joined to three other carbon atoms to form three covalent bonds. They are in one plane hence form a two-dimensional layer.

4. The carbon-carbon bond length in graphite is 1.42Å and the distance between two layers is 3.35Å

5. The parallel layers are loosely held together by Van der Waals forces. That’s why graphite is very soft.

6. The density of graphite is very low.

7. The graphite is a good conductor of heat and electricity. The reason is each carbon atom has one unused electron in p-orbital which overlaps laterally with p-orbital of other carbon atoms to for delocalized pi bonding.

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