Tabular Difference Between Diamond and Graphite with Structure and Properties
The difference between diamond and graphite is a foundational concept in JEE Main Chemistry, illustrating how atomic structure can drastically alter physical properties. Both are allotropes of carbon, yet they exhibit strikingly different characteristics—diamond being the hardest naturally occurring substance, while graphite is soft and slippery. This contrast arises purely from differences in bonding and structure. JEE aspirants must understand these distinctions for concept-based MCQs and application questions on crystalline solids, hybridisation, and structure-property relationships.
Allotropy and Significance in Chemistry
Allotropy refers to the existence of an element in more than one physical form in the same physical state. Carbon shows allotropy by forming several polymorphs; the most well-known are diamond and graphite. Their comparison is a classic example for questions on chemical bonding, periodic table trends, and real-life material applications.
Structure of Diamond
Diamond consists of each carbon atom covalently bonded to four other carbon atoms, forming a rigid three-dimensional tetrahedral network. This sp3 hybridisation leads to a very strong covalent structure, making diamond extremely hard and transparent. There are no free electrons in this arrangement.
Structure of Graphite
Graphite is composed of layers of carbon atoms arranged in hexagonal rings. Each carbon in a layer is strongly bonded to three neighbors via sp2 hybridisation, while a weak van der Waals force holds the layers together. Delocalised electrons within the layers allow graphite to conduct electricity and make it soft and slippery.
Tabular Difference: Diamond vs Graphite
| Property | Diamond | Graphite |
|---|---|---|
| Nature | Covalent network solid; transparent, crystalline | Layered structure; opaque, crystalline |
| Hybridisation | sp3 hybridisation | sp2 hybridisation |
| Bonding | Each carbon bonded to 4 carbons (tetrahedral) | Each carbon bonded to 3 carbons (hexagonal rings); one free electron per C |
| Hardness | Very hard (hardest natural material) | Soft and slippery |
| Electrical Conductivity | Non-conductor | Good conductor |
| Density (g/cm3) | 3.5 | 2.26 |
| Solubility | Insoluble in all solvents | Insoluble in all solvents |
| Uses | Jewellery, cutting tools, abrasives | Pencils, lubricants, electrodes, batteries |
| Chemical Reactivity | Less reactive | More reactive (burns readily in O2) |
Chemical Bonding and Hybridisation
In diamond, carbon atoms use sp3 hybrid orbitals to form four equivalent covalent bonds. In graphite, carbon uses sp2 hybridisation with one unhybridised p orbital per carbon, leading to the delocalised π-electrons responsible for electrical conductivity within its layers.
Physical Properties and Real-Life Implications
Diamond is colorless, transparent, and extremely hard; it does not conduct electricity due to the absence of free electrons. Its high refractive index and hardness make it suitable for jewellery and industrial machining tools. In contrast, graphite is black, soft, and slippery, thanks to weak forces between layers; its free electrons provide electrical conductivity, so it is used in making electrodes and pencil leads.
Practical and Industrial Uses
- Diamond is valued in cutting instruments and as a gemstone.
- Graphite is essential in pencil manufacture, lubricants, and dry cells.
- Electrodes in batteries are often made from graphite, leveraging its conductivity.
- Diamond is used in high-precision industrial applications requiring extreme hardness.
Similarities and Carbon Polymorphism
Despite their differences, both diamond and graphite are pure forms of carbon. Their distinct structures but identical composition make them polymorphs. These structural variations explain their divergent properties, a key aspect of allotropy in Chemistry.
JEE-Ready Summary: Key Takeaways
- Diamond: sp3 network, hard, non-conducting, high density, tetrahedral bonding.
- Graphite: sp2 layers, soft, conducts electricity, lower density, hexagonal bonding.
- Difference between diamond and graphite is due to their crystal structure and bonding; not atomic mass or composition.
- Common JEE confusion: structure, conductivity, uses, and hybridisation differences.
Explore Related JEE Chemistry Topics
- Hybridisation in carbon: sp3 vs sp2
- Chemical Bonding and Structure
- Diamond: Properties and Applications
- Crystalline and Amorphous Solids
- Covalent Bonding in Carbon
- Difference: Isomers and Allotropes
- Metals vs Non-metals
- Hybridisation of Graphite
- Difference Between Diamond and Graphite
- Allotropes vs Isomers
- Allotropes of Carbon
Mastering the difference between diamond and graphite equips JEE candidates to link atomic structure with material properties, a key Chemistry skill. For more exam-oriented content curated by Vedantu experts, explore the interlinked topics above.
FAQs on Difference Between Diamond and Graphite in Chemistry
1. What is the main difference between diamond and graphite?
Diamond and graphite are both allotropes of carbon but differ in their atomic structures and properties.
- Diamond: Each carbon atom forms four strong covalent bonds in a 3D tetrahedral network (sp3 hybridisation), making it extremely hard and non-conductive.
- Graphite: Each carbon atom bonds to three others in flat hexagonal layers (sp2 hybridisation); layers slide easily, making graphite soft and a good conductor of electricity.
2. Why is diamond hard but graphite is soft?
The hardness of diamond and softness of graphite is due to differences in their atomic bonding and structure.
- In diamond, every carbon is bonded to four other carbons by strong covalent bonds in 3D, creating a rigid, interlocked structure.
- In graphite, carbon atoms are bonded in flat layers, with weak forces (van der Waals) between layers, allowing them to slide over each other easily.
3. Is graphite stronger than diamond?
No, diamond is much stronger and harder than graphite due to its network of strong covalent bonds.
- Diamond is the hardest known natural substance.
- Graphite is soft and can be used as a lubricant or for pencil leads because its layers break apart easily.
4. Why do diamonds turn to graphite under certain conditions?
Under high temperature and specific pressure, diamond can slowly convert to graphite because graphite is thermodynamically more stable at standard conditions.
- This is a slow process at normal temperature and pressure.
- It is an example of polymorphic transformation of carbon.
5. Why is graphite soft but diamond hard?
The softness of graphite is due to its layered structure, where weak forces hold the layers together.
- Layers slide over each other when pressure is applied.
- Diamond is hard due to a rigid 3D network of strong covalent bonds.
6. Is diamond a better conductor than graphite?
No, graphite conducts electricity well, while diamond does not.
- In graphite, each carbon atom has a free electron, allowing electrical conductivity.
- Diamond lacks free electrons due to its full valence shell in the sp3 hybrid structure.
7. Can both diamond and graphite exist naturally and synthetically?
Yes, both diamond and graphite occur in nature and can also be created artificially.
- Natural: Found in Earth's crust (diamond in kimberlite, graphite in metamorphic rocks).
- Synthetic: Manufactured using high pressure and temperature (for diamonds) or chemical vapor deposition (for graphite).
8. What is meant by ‘polymorphs’ with respect to carbon?
‘Polymorphs’ are different crystalline forms of the same element. Diamond and graphite are polymorphs of carbon because they have different atomic arrangements but are both made of carbon atoms.
9. Why does graphite conduct electricity but diamond does not?
Graphite conducts electricity because each carbon atom's fourth electron is free and delocalised within its layers, allowing electric current to flow. Diamond does not conduct because all four valence electrons are used in strong covalent bonding, leaving no free electrons for conduction.
10. How do the structures of diamond and graphite affect their uses in industry?
The distinct structures of diamond and graphite determine their industrial applications:
- Diamond: Used for cutting, grinding, and drilling due to extreme hardness; also valued as gemstone.
- Graphite: Used in pencils, lubricants, batteries, and as electrode material because it is soft, slippery, and conducts electricity.
