Difference Between Diamond and Graphite for JEE Main 2024

Diamond and graphite are two allotropes of carbon that exhibit vastly different physical and chemical properties. These two forms of carbon are well-known for their diverse applications and fascinating structures. In this article, we will delve into the dissimilarities between diamond and graphite, shedding light on their crystal structures, properties, uses, definitions ‘what is diamond and graphite’ and characteristics of diamond and graphite.  This knowledge is especially valuable for JEE (Joint Entrance Examination) Mains aspirants, as it highlights the significance of these allotropes and their relevance in various scientific disciplines.

Crystal Structure:

Diamond and graphite possess distinct crystal structures, which contribute to their dissimilar properties. Diamond is a three-dimensional covalent network crystal, wherein each carbon atom is bonded to four other carbon atoms in a tetrahedral arrangement. This network of strong covalent bonds gives diamond its exceptional hardness and makes it an excellent thermal conductor. On the other hand, graphite has a layered structure, consisting of stacked layers of carbon atoms arranged in a hexagonal lattice. Each carbon atom in graphite is bonded to three other carbon atoms, forming strong covalent bonds within the layers. However, weak van der Waals forces between the layers allow them to slide over each other, resulting in the characteristic slippery nature of graphite.

Physical Properties:

1. Hardness:

Diamond is renowned for its extreme hardness, ranking 10 on the Mohs scale. It is the hardest known natural substance, making it highly resistant to scratching or abrasion. Conversely, graphite is relatively soft and has a Mohs hardness of only 1-2. This difference arises due to the distinct bonding and crystal structures of the two allotropes.

2. Electrical Conductivity:

While diamond is an excellent thermal conductor, it is an insulator when it comes to electricity. The strong covalent bonds in diamond leave no free electrons to conduct electricity. In contrast, graphite is an excellent electrical conductor due to its layered structure. The delocalized π-electrons in graphite's hexagonal layers can move freely along the planes, facilitating electrical conductivity.

3. Optical Properties:

Diamond possesses exceptional optical properties, such as high refractive index and strong dispersion of light, which give rise to its brilliant luster and fire. It exhibits strong light reflection and refraction, contributing to its exceptional sparkle. Graphite, on the other hand, is opaque and exhibits a dull appearance, primarily due to its layered structure and the presence of numerous carbon layers.

Chemical Properties:

1. Reactivity:

Diamond is chemically inert and highly resistant to most acids and alkalis. Its strong covalent bonds make it highly stable, preventing chemical reactions under normal conditions. Graphite, however, is relatively reactive compared to diamond. It readily undergoes oxidation when exposed to oxygen at high temperatures, forming carbon dioxide.

2. Combustibility:

Diamond does not burn and is unaffected by heat. It requires extreme conditions, such as temperatures above 800°C in the presence of oxygen, to oxidize. On the other hand, graphite is combustible and readily burns in the presence of oxygen, releasing carbon dioxide.

Uses and Applications:

Diamond and graphite find various applications owing to their contrasting properties.

Diamond:

• Jewelry: Diamond's exceptional hardness, brilliance, and fire make it highly desirable for use in jewelry.

• Cutting and Polishing Tools: Diamond's hardness makes it ideal for cutting and polishing applications, such as in the manufacturing of diamond blades and grinding wheels.

• Industrial Applications: Diamond's exceptional thermal conductivity and hardness make it suitable for use in various industrial applications, including heat sinks, high-power lasers, and cutting tools for machining.

• Electronics: Diamond exhibits unique electrical properties, making it promising for applications in electronics, such as high-frequency transistors and diodes.

Graphite:

• Lubricants: Graphite's slippery nature and low friction properties make it an excellent dry lubricant, particularly in high-temperature environments.

• Batteries: Graphite is widely used as an anode material in lithium-ion batteries due to its ability to intercalate lithium ions.

• Pencils: The softness and ability to leave marks on paper make graphite a common choice for pencil leads.

• Nuclear Reactors: Graphite's ability to slow down neutrons makes it suitable for use as a moderator in nuclear reactors.

Diamond and Graphite Difference

This table provides a simplified overview of the differences between diamond and graphite. There may be additional nuances and properties that are not included here.

Characteristics of Diamond and Graphite:-

Diamond:

• Hardness: Diamond is the hardest known natural substance, ranking 10 on the Mohs scale of mineral hardness. It owes its exceptional hardness to its three-dimensional covalent network structure, where each carbon atom is bonded to four others in a tetrahedral arrangement.

• Optical Properties: Diamond exhibits brilliant luster and dispersion of light, resulting in its sparkling appearance. This is attributed to its high refractive index and strong light reflection and refraction properties.

• Thermal Conductivity: Diamond is an excellent thermal conductor due to its rigid structure and strong covalent bonds. It efficiently transfers heat, making it suitable for applications requiring thermal management, such as heat sinks.

• Electrical Insulator: Despite its thermal conductivity, diamond is an electrical insulator. The absence of free electrons in its covalent network structure restricts the flow of electricity.

• Chemical Inertness: Diamond is highly resistant to chemical reactions, making it chemically inert. It withstands exposure to most acids and alkalis, maintaining its integrity.

Graphite:

• Layered Structure: Graphite has a layered structure, consisting of stacked carbon layers held together by weak van der Waals forces. Each layer is made up of hexagonally arranged carbon atoms bonded to three neighboring carbon atoms.

• Lubricating Properties: Due to its layered structure, graphite exhibits a slippery texture and low friction. These characteristics make it an excellent dry lubricant, particularly in high-temperature environments.

• Electrical Conductor: Graphite is an excellent electrical conductor. The presence of delocalized π-electrons within the layers allows for the movement of electrons, enabling the conduction of electricity.

• Softness: Unlike diamond, graphite is relatively soft and has a low Mohs hardness of 1-2. The layers can slide over each other easily due to the weak interlayer forces, imparting a characteristic "pencil-lead" texture.

• Combustibility: Graphite is combustible and readily burns in the presence of oxygen, producing carbon dioxide. This property is utilized in various applications, including in the manufacturing of electrodes and as a fuel in some industries.

Summary:-

Diamond and graphite, two carbon allotropes, showcase distinct properties arising from their contrasting crystal structures and bonding arrangements. Diamond is renowned for its exceptional hardness, optical properties, and thermal conductivity, while graphite exhibits electrical conductivity and lubricating properties. Understanding the dissimilarities between these two allotropes is essential for JEE Mains aspirants, as it underscores the significance of carbon in various scientific applications and industries.