Problem 29 - 43 Years JEE ADVANCED (1978-2020) + JEE MAIN Chapterwise & Topicwise Solved Papers Chemistry

Question

With respect to graphite and diamond, which of the statement(s) given below is (are) correct? \([2012]\) (a) Graphite is harder than diamond. (b) Graphite has higher electrical conductivity than diamond. (c) Graphite has higher thermal conductivity than diamond. (d) Graphite has higher C-C bond order than diamond.

Step-by-Step Solution

Answer

Only statement (b) is correct: Graphite has higher electrical conductivity than diamond.

Step 1: Understanding Graphite and Diamond

Graphite and diamond are two allotropes of carbon that differ in their atomic structure and properties. Graphite is composed of layers of carbon atoms with weak Van der Waals forces between layers, making it soft and slippery. Diamond, on the other hand, consists of a three-dimensional network of carbon atoms linked by strong covalent bonds, making it extremely hard.

Step 2: Analyzing Statement (a)

Statement (a) suggests that graphite is harder than diamond. We know that diamond is one of the hardest natural substances due to its robust covalent bonding structure. In contrast, graphite is soft and used as a lubricant. Therefore, statement (a) is incorrect.

Step 3: Evaluating Electrical Conductivity

Graphite has a higher electrical conductivity than diamond due to its structure. Each carbon atom in graphite forms three sigma bonds and has one delocalized electron, which can move freely between layers, conducting electricity. Diamond lacks free electrons, which makes it a poor electrical conductor. Therefore, statement (b) is correct.

Step 4: Evaluating Thermal Conductivity

Statement (c) addresses thermal conductivity. Diamond is known for having remarkable thermal conductivity due to its strong covalent bonds and lattice structure. Graphite, although conducting electricity better, does not have as high thermal conductivity as diamond. Therefore, statement (c) is incorrect.

Step 5: Analyzing C-C Bond Order

In diamond, each carbon atom forms four covalent bonds in a tetrahedral structure, whereas in graphite, each carbon atom forms three bonds with a resonance structure contributing delocalized electrons. The bond order in graphite remains essentially single, while in diamond, each bond is a strong single bond. Thus, statement (d) is incorrect.

Key Concepts

Graphite vs Diamond

Graphite and diamond are fascinating examples of carbon allotropes, which are different forms of carbon distinguished by their atomic arrangements and properties. Graphite has a structure where carbon atoms are arranged in stacked sheets with weak Van der Waals forces between them. These layers can slide over each other, making graphite feel slippery and soft. In contrast, diamond is famous for its hardness. Its structure is a rigid three-dimensional network of carbon atoms, each bonded strongly to four other atoms. This makes diamond incredibly hard, suitable for cutting tools, while graphite is useful for lubricants and pencils.

• Graphite: soft, slippery, sheets of carbon atoms
• Diamond: hard, rigid, 3D network of carbon atoms

Despite being composed entirely of carbon, these two substances exhibit vastly different characteristics due to their structures.

Electrical Conductivity

The difference in electrical conductivity between graphite and diamond is another fascinating comparison. Graphite is an excellent conductor of electricity, whereas diamond is not.
This stark contrast stems from the atomic structure of each material. In graphite, each carbon atom forms three covalent bonds with other carbon atoms, leaving one electron free. These free electrons, also known as delocalized electrons, can move freely along the layers, allowing graphite to conduct electricity efficiently throughout the material.
On the other hand, diamond's structure consists entirely of strong covalent bonds where each carbon atom is bonded to four others. No free electrons are available to move and carry charge. This lack of mobile charge carriers makes diamond an electrical insulator.

• Graphite: conducts electricity well due to free electrons
• Diamond: non-conductive due to no free electrons

Thermal Conductivity

Although diamond and graphite differ significantly in electrical conductivity, their thermal conductivity tells a different story. Diamond is one of the best thermal conductors known to man, while graphite lags behind in this aspect.
The thermal conductivity of diamond can be attributed to its strong network of covalent bonds, which allow phonons (vibrations that carry heat) to travel quickly through the material. This efficient transfer of heat makes diamond very effective at conducting thermal energy, which is why it's used in thermal management applications.
In comparison, graphite does not share this trait, despite its good electrical conductivity. The weak forces between the sheets in graphite do not support the same level of phonon movement, which results in lower thermal conductivity.

• Diamond: exceptional thermal conductor due to robust covalent bonds
• Graphite: lower thermal conductivity, performed by weaker interlayer forces

C-C Bond Order

C-C bond order is a term that describes the nature and strength of carbon-carbon bonds in a molecular structure. In the context of graphite and diamond, this concept plays a crucial role in understanding their different properties.
In diamond, each carbon atom is bonded to four others in a three-dimensional tetrahedral geometry, creating a bond structure considered as single bonds. These strong single bonds contribute to diamond's remarkable hardness and stability.
Graphite's structure, however, involves each carbon atom being bonded to three others, with the presence of delocalized π-electrons due to resonance. These contribute to the strength of its bonds but overall result in a bond order that is less than that of diamond. This is because the resonance and layering in graphite introduce a level of variability in bond strength.

• Diamond: has strong single C-C bonds in tetrahedral arrangement
• Graphite: includes resonance and π-bonding, affecting bond order