Graphite and diamond are some of the most popular allotropes of carbon.
In today’s industrial applications, these two carbon allotropes play an important role – from manufacturing, parts production to final products.
But, a question always a rises: how does graphite compare to carbon? That’s exactly what we shall explore in this article.
What is Graphite?
Graphite is an allotrope of carbon. The carbon atoms in graphite are organized in a hexagonal manner forming stacked layers. The layers in graphite have a weak bond, allowing them to slide over each other easily. This loose bond provides for graphite’s slippery nature making it a good lubricant for your application.
Graphite’s carbon atoms are connected to three more carbons, resulting in a two-dimensional sheet. This property provides the excellent capacity to conduct heat and electricity perfectly.
What is a Diamond?
Diamond is an allotrope of carbon. Diamond’s carbon atoms are joined in a three-dimensional structure. This structure forms the diamond’s strong and rigid shape making it one of the hardest materials.
The carbon atoms in diamonds are strongly bonded providing for graphite hardness. Diamond can conduct heat perfectly but it cannot conduct electricity.
You can leverage diamond in applications that require cutting, drilling, and polishing because of its hardness.
Where Diamond and Graphite Are Found
· Naturally Occurring Diamond and Graphite
- Naturally Occurring Diamond:
Diamonds are formed deep within the Earth’s mantle. Naturally occurring volcanic eruptions bring diamonds to the Earth’s crust. The volcanic eruptions form special rocks where diamonds are found. These rocks known as kimberlite and lamporite pipes undergo mining for the diamonds to be extracted.
Alluvial deposits are deposits resulting from river activities and diamonds can also be extracted from them.
- Naturally Occurring Graphite:
Graphite is created in metamorphic rocks that result from conditions with high temperature and pressure. They are commonly found in schists, marbles, and gneisses.
You can extract natural graphite by using open-pit mining or Underground mining. In open-pit mining, the layers of the surface are removed while in underground mining tunnels are dug by miners to get to the deposits.
· Synthetic Graphite and Diamond
- Synthetic Diamond:
High-pressure high temperature(HPTP)and Chemical Vapor Deposition(CVD) are the methods you can utilize in making synthetic diamonds. HPTP grows diamond crystals by leveraging excess heat and temperature deep within the earth’s mantle.
CVD on the other hand, diamond is gradually formed by a gas combination being used on a substrate where atoms of carbon are placed one layer at a time.
- Synthetic Graphite:
You can make your synthetic graphite by heating petroleum coke or coal tar pitch under a very intense temperature of about 3000 degrees Celsius. This can result in manufacturing graphite similar to natural forms but offers the advantage of fulfilling your specific requirements.
Analyzing Diamond Vs Graphite Properties
· Thermal Expansion
Diamond has an expansion has a low expansion rate because of its strong covalent bonds and tightly packed structure. Its tetrahedral structure enables it to withstand the rise in temperature preventing significant changes in its size. You can use diamonds in your applications that require intense temperatures and precise dimensions need to be maintained.
Graphite’s thermal expansion is dependent on the direction. Its size reduces slightly when heated within the layers portraying a negative thermal expansion coefficient. It however expands more across layers because they are held with weak forces that allow them to move apart when heated.
· Thermal Conductivity
Diamonds are among the materials perfect for conducting heat. It has good thermal conductivity this is because of its tetrahedral structure and strongly bonded carbon atoms. Its rigid structure prevents the scattering of phonons making it defect-free hence heat can spread very fast.
Graphite conducts heat too but not exceptionally as diamond. Although carbon atoms in graphite are bonded tightly in a hexagonal arrangement, the layers are bonded with weak forces limiting the transfer of heat from one layer to the other. This results in graphite conducting heat better along the layers unlike across the layers.
· Electrical Conductivity
Diamond’s conductivity is significantly lower compared to Graphite. It is a good insulator. Its structure limits electrons from moving freely thus it does not conduct electricity under normal conditions. Under certain conditions such as when it is doped with components like boron, it can act as a semi-conductor.
Graphite is an excellent conductor of electricity because its electrons move freely within its structure. Delocalized electrons enable efficient electrical conduction as they can smoothly move within layers. Graphite has better conductivity along the layers compared to across the layers because of the weak forces holding them together.
· Coefficient of Friction
Diamond has a low coefficient of friction that is between 0.05 and 0.1 when it is used against other materials. It can withstand abrasion and wear because of its incredible hardness and how smooth its surface is and this minimizes friction. You can use diamond coatings in your applications to minimize friction and enhance robustness.
Graphite has a higher coefficient of friction than diamond which ranges between 0.1 and 0.3. Its layers are held together by weak forces thus they can easily slide over the other making it an excellent material for dry lubricants.
· Density
Diamond is considered dense because it has a density of around 3.51g/cm3. Its high density is a result of its tetrahedral arrangement. They are packed tightly and have a rigid crystal lattice. As a result, diamonds have not only an incredible hardness but they are also very dense.
Graphite, compared to diamond has a lesser density of 2.2 g/cm3. Layers in graphite are weakly bonded by weak van der Waals forces and hence can easily slide over each other. There are spaces between layers that are caused by the loose stacking and this minimizes the density of graphite.
· Hardness
Diamond’s carbon atoms are joined in a three-dimensional structure. This structure forms the diamond’s strong and rigid shape. It is one of the hardest materials known and it has a Mohs hardness of 10.
Graphite is a soft material with a Mohs hardness of between 1 and 2. Its softness makes it a perfect lubricant for your applications.
· Stability
Diamond is metastable under room temperature and normal atmospheric pressure. It is formed deep within the earth’s crust at very intense temperatures, making its structure stable.
Graphite is a more stable allotrope of carbon at room temperature and pressure. In this environment its hexagonal layer arrangement is energetically favorable thus it is not tampered with for long it will ultimately transform into graphite. Its stability results from its resilient and flexible properties.
· Color
Diamonds with impurities have several colors such as yellow and blue. Pure diamonds, however, are colorless. This is because the arrangement of carbon atoms prevents the absorption of visible light.
Graphite has a black or dark gray color. This results from delocalized electrons taking in and scattering light within its structure.
· Entropy for Diamond Vs Graphite
Diamond has a highly organized and rigid structure that limits the freedom of atoms to move. This equates to fewer microstates that result in relatively low entropy value.
Graphite has a higher entropy than diamond because its structure allows for movement between layers. Delocalized electrons in graphite surge microstates present and the freedom and flexibility enhance graphite’s entropy value.
· Melting Point
Diamond is held by strong covalent bonds in its structure making it require a lot of energy to break. Diamonds have a melting point above 35500C.
Graphite has a lower melting point than diamond under intense pressure. Its carbons are covalently bonded but its layers are held together by weak forces hence they can withstand temperature but not as well as diamond.
· Magnetic Properties
All electrons in a diamond are covalently bonded hence there are no free magnetic movements hence diamond is non-magnetic. However, it can display weak magnetic properties if it is doped with other elements like nitrogen.
Graphite also lacks unpaired electrons making it diamagnetic but they can however display paramagnetic properties in specific environments.
· Toughness
Diamond has low toughness despite having incredible hardness. This is because its rigid crystal lattice makes it brittle and a huge impact can cause fractures on it.
Graphite can take in stress without breaking because its layered structure offers flexibility and prevents fractures. You can utilize graphite in applications that require flexibility.
· Tensile Strength
Diamond has excellent tensile strength along certain directions because its carbon atoms are covalently bonded. Yet, its tensile strength is limited by its brittleness.
Graphite has low tensile strength because of the weak forces that hold its layers. These forces do not provide much resistance to pressure resulting in layers breaking.
· Dielectric Constant
Diamond has a low dielectric constant because its electrons are closely bound and have a non-metallic property.
Graphite has anisotropic dialectic properties. Along the basal plane, graphite has high electrical conductivity and low dielectric constant while perpendicular to the layers, graphite has a higher dielectric constant.
Bonding in Diamond Vs Graphite
Diamond has a covalent network structure. Each carbon atom is covalently bonded to four other carbon atoms. It has a rigid three-dimensional lattice.
Graphite has a layered structure. Each atom is covalently bonded to three other carbon atoms and layers are held together by weak forces.
Intermolecular Forces in Graphite and Diamond
Diamond has a covalent network structure and thus does not display intermolecular forces. Graphite however has its layers held together by weak van der Waals forces.
Exploring Uses of Diamond and Graphite
You can use diamonds to make jewelry. Also diamond is a perfect material for making cutting and drill tools because of its incredible hardness.
Furthermore, you can use diamonds with nitrogen-vacancy in quantum computing and sensing.
On the other hand, graphite can be used in electrodes for batteries, crucible and furnace parts. Additionally, graphite is a perfect lubricant besides being neutron moderator in nuclear reactors.
Lastly, graphite can be used as a material for pencils and other writing materials.
FAQs
1. Which is Stronger, Diamond or Graphite?
Diamond is stronger than graphite. It has a Mohs hardness of 10 while graphite has a Mohs hardness between 1 and 2.
2. Which one is more Expensive, Diamond or Graphite?
Graphite is more expensive than graphite because of factors like demand, rarity, and craftsmanship needed in its processing.
3. Can Graphite Cut Diamonds?
No, graphite cannot cut diamonds because diamond is harder than graphite.
4. Can we Convert Graphite into Diamond?
Yes, under high pressure and temperature conditions, we can convert graphite to diamonds although it is an energy-intensive process.
5. What is the difference between Diamond and Graphite?
Diamond ‘s carbon atoms are covalently bonded to four carbon atoms making it rigid. On the other hand, graphite’s carbon atoms are covalently bonded to three other carbon atoms with weak forces between layers.
Conclusion
Both graphite and diamonds are important allotropes of carbon and can be tailored for your specific applications.
Once you understand your application requirements, you can easily choose graphite or diamond for your specific applications.
