What Are The Main Properties of Graphite?

Graphite is characterized by a giant covalent structure where each carbon atom bonds covalently to three neighboring carbon atoms. This bonding pattern results in the formation of layers, with the carbon atoms arranged in a hexagonal pattern within each layer.

What’s interesting about these layers is that, while the atoms within a layer are tightly bound, the forces holding one layer to the next are relatively weak. This unique arrangement accounts for many of graphite’s distinctive properties.

Physical Properties of Graphite

Graphite, much like diamond, possesses a remarkably high melting point. This characteristic stems from its structure; simply loosening the layers from one another is not sufficient to melt it. Instead, the covalent bonds that run throughout the entire structure must be disrupted—a process requiring significant energy.

Despite its impressive thermal resilience, graphite stands out for being soft and slippery to the touch. This quality explains its widespread use in everyday items such as pencils, where layers are easily transferred onto paper, and as a dry lubricant for mechanisms like locks.

The structure of graphite can be compared to a stack of cards: while each individual layer is robust, the layers themselves slide over one another with ease, sometimes even separating completely.

Another notable property of graphite is its relatively low density when compared to diamond. Additionally, graphite is insoluble in both water and organic solvents, mirroring the insolubility seen in diamond, due to the strength of the covalent bonds that hold the structure together.

Graphite stands out for having a lower density compared to diamond. This difference comes down to its structure—there is simply more empty space between the individual layers of atoms, which makes the material less compact overall.

When it comes to solubility, graphite doesn’t dissolve in water or organic solvents. The explanation here mirrors what we see with diamond: the covalent bonds holding the carbon atoms together are just too strong. No matter how much you mix graphite with a solvent, the attraction between the solvent molecules and the carbon atoms is never enough to pull the structure apart.

One of graphite’s most important features is its ability to conduct electricity. This is because the electrons in graphite aren’t all locked into place; some are free to move across the layers. As a result, when you connect graphite into an electrical circuit, electrons can enter at one end, travel along the layers, and exit at the other, keeping the current flowing.

If you look at graphite, you’ll usually see a grayish-black, opaque material with a subtle, metallic sheen. It’s surprisingly soft and can be broken with very little pressure. Its specific gravity is low, which means it feels lighter than you might expect for a material made of carbon. Interestingly, graphite straddles the line between metals and nonmetals, showing characteristics of both.

Despite being flexible, graphite doesn’t snap back if you bend it; it isn’t elastic. Still, it stands out for its high electrical and thermal conductivity, which opens the door to a wide range of uses, especially in metallurgical processes and various areas of manufacturing.

Chemical Properties of Graphite

Color	Steel gray to black
Chemical Classification	Native element
Streak	Black
Luster	Metallic, sometimes earthy
Diaphaneity	Opaque
Cleavage	Perfect in one direction
Mohs Hardness	1 to 2
Specific Gravity	2.1 to 2.3
Diagnostic Properties	Color, streak, slippery feel, specific gravity
Chemical Composition	C
Crystal System	Hexagonal

What are the main properties of graphite?

Properties of Graphite:

• A greyish-black, opaque substance.
• Lighter than diamond, smooth and slippery to the touch.
• A good conductor of electricity (Due to the presence of free electrons) and a good conductor of heat.
• A crystalline solid
• Very soapy to the touch.
• non-inflammable.
• Soft due to weak Vander wall forces.
• The conductor of electricity.

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