Do you know how many individuals unknowingly use products that contain graphite? For instance, did you know that graphite makes most automotive lubricants, electric motor brushes, and the “lead” in pencils? Speaking of which, does graphite conduct electricity?
Knowing this information will be interesting. By reading the article, learn whether graphite has a high melting point, conducts heat and electricity, or whether it is poisonous. You will also learn about graphite’s many properties, uses, and whether it’s stronger than diamond.
Can Graphite Conduct Electricity?
Graphite conducts electricity. It has free electrons that conduct electricity by carrying charge from one location to another. Each carbon atom in graphite has a covalent link with three other atoms, freeing one atom. The free electron moves across the graphite material and does not bond to any particular atom. Free atoms can leave that layer or move to one neighboring.
The weak forces between the layers of graphite enable free electrons to flow without facing any resistance. Because each atom is a charge carrier and can move around freely, it can transfer electric current.
Here is a brief demonstration of the electrical conductivity of graphite: Graphite Conducts Electricity
Why Do Graphite Conduct Electricity?
Carbon atoms with four electrons in their outer shell make up graphite. The fourth electron is “delocalized,” or free to move about, while the other three form a solid bond. Single and double bonds alternate among the layers of graphite. With this kind of structure, electrons can move freely along the layers.
The weak forces separating the graphite layers allow the layers to slide over one another. An electric current can flow through graphite thanks to the weak intermolecular forces of attraction or van der Waals forces. The van der Waals forces weaken graphite, making it useful in manufacturing lubricants and pencil lead.
Does Graphite Conduct Heat?
Graphite can conduct heat because it has free electrons; only three of the four valence bonds in graphite form a cohesive bond. The fourth electron does not form a bond and will move from one carbon atom to the next under the control of applied potential. The free electron is left free to transfer heat.
The vibration of atoms is the leading cause of heat transmission. The electrons in a medium with mobile electrons will vibrate and transmit heat. Graphene sheets form when the carbon atoms in graphite arrange themselves into a hexagonal shape. The carbon atoms in each graphene sheet arrange themselves in a honeycomb lattice.
Interlayer bonds between the structure’s layers are weak. The atoms can quickly vibrate due to the thin, weak layers.
Does Graphite Have a High Melting Point?
The melting point of graphite is a high 6512°F (3730°C). The high temperatures that are needed ensure that graphite does not follow the normal process of igniting. Graphite does not become liquid at this temperature; instead, it transforms from a solid to a gas. Because of the need to break the strong covalent bonds in graphite, high temperatures are necessary.
All the atoms will share the delocalized atom in every layer of graphite. The strength of the bond between the atoms increases by sharing the delocalized free electron, boosting the stability of the graphite. These two characteristics make bond breakdown extremely complex and energy-intensive, which accounts for the high melting temperature.
Is Graphite Poisonous to Humans?
In most cases, graphite is not poisonous to humans. Some claim you can become poisoned if you chew the lead in your pencil. The “lead” in your pencil is graphite, not actual lead. You might be surprised to learn that graphite and your body structure contain carbon atoms.
In reality, carbon makes up 12% of the human body. Hence, there is no real danger if you breathe or consume graphite. However, this doesn’t mean you should excessively expose yourself to graphite. If you inhale graphite particles for a long time, you could develop graphitosis.
When inhaling graphite dust for an extended period, it can lead to a deadly form of pneumoconiosis called graphitosis. For graphite poisoning to occur, you should have consumed around 428 micrograms of graphite for every milliliter of blood. It would be nearly impossible for somebody to consume this much graphite.
Additionally hazardous when burned graphite. Carbon dioxide and carbon monoxide are two byproducts of the burning of graphite. Since carbon dioxide is the same gas you exhale, it is not dangerous. However, you will suffocate if carbon dioxide replaces the air you are inhaling.
Carbon monoxide, the other product, is poisonous. Your lungs’ oxygen receptors bind to the gas, and it won’t release that connection again. As a result, red blood cells will never carry oxygen in your body if carbon monoxide clings to them.
Which is Stronger Diamond or Graphite?
Due to how the carbon atoms in diamonds bond together, diamond is stronger than graphite. Diamonds have a stronger covalent bond, which makes them denser. The diamond’s carbon atoms arrange themselves in a tetrahedral structure as they bond, forming four covalent connections. Diamonds don’t contain any free electrons floating around.
In contrast, the carbon atoms in graphite make three covalent connections to form a hexagonal structure. The intermolecular interactions between the carbon atoms in graphite appear minimal, allowing the layers to slide over one another. It is this movement that makes graphite soft and slippery.
You might have observed that a bit of graphite remains on the paper after using a pencil. The graphite layers slide off one another as a result. The sheets stick to the paper as they glide and scrape against one another.
Because there is a considerable amount of space between its layers, graphite has a lower density than diamond.
Read: Is Duct Tape Flammable?
Properties of Graphite
Graphite is one of the most beneficial natural elements. The method of refining graphite is a complex chemical process that involves the heat decomposition of polymeric film.
Because of that, graphite has unique properties, as highlighted below:
- Conducts electricity and heat: Graphite has free electrons that carry electric current freely without facing any resistance.
- High melting point: If you want to melt graphite, you’ll need a lot of heat to break the strong covalent bonds.
- A soft, slippery feel: Due to its layers that slide over each other, it can rub off surfaces such as paper. That’s why most lubricants contain graphite, and the inner material of pencils consists of graphite.
- Insoluble in water and organic solvents: When carbon atoms connect, non-polar covalent bonds form.
- Flexible: Graphite is soft and cracks under mild pressure. However, though flexible, it is not elastic. It is the only non-metal that has some properties of metal.
- A greyish-black, opaque substance: Its greasiness and softness will leave a black mark when you touch it with your hand.
Various Uses of Graphite
Graphite has a wide range of applications in manufacturing and engineering due to its unique properties. Graphite consists of different forms, and highlighted are three of its most famous form and uses:
1. Amorphous graphite
Tiny crystals of amorphous graphite can contain up to 80% carbon. It is also the cheapest and most common among the three. Here are some examples of its uses:
- Graphite is a primary component in most lubricants, including grease and forging lubricants.
- It is a primary ingredient in electrolytes for printing.
- The black track that graphite leaves on the paper make it suitable for making the “lead” in pencils.
2. Lump/vein graphite
Lump graphite is the most expensive and has a carbon content of up to 90%. It is not easily found and resembles solid lumps. Here are some examples of its applications:
- Manufacturers use lump graphite in the production of paints. Graphite is naturally water-repellent and best for providing a protective coating on wood.
- Due to its high thermal and electrical conductivity, lump graphite makes high-friction applications such as car brakes and clutches.
3. Crystalline flake graphite
Crystalline graphite is rare and contains between 85 and 90% carbon. It appears as giant crystals or as flakes that range in size from coarse to fine. The cost of graphite is a little higher than that of amorphous graphite. Here are a few examples of its applications:
- It helps produce plates, brushes, and carbon electrodes needed in dry cell batteries and the electrical sector since it has a low resistance to electrical conductivity.
- Manufacturers use it in the manufacturing of crucibles.
- Because of its superior crystalline shape, manufacturers use it to manufacture lubricants. It has low friction and the capacity to create an oil suspension, hence its preference in making lubricants.
- It serves as the primary component in repellant solutions. Metal protectors are among the most popular crystalline graphite repellents.
Graphite is a good conductor of electricity and heat. Its ability to do both is due to how its carbon atoms bond. The electrons in the outer shell leave one electron free, which carries an electric charge. The free electron can vibrate hence transmitting heat.
Graphite is weaker than diamond because of its weak intermolecular force. You don’t have to fear using graphite as it’s only harmful when you inhale its dust for an extended period.