Question

Consider following statements: 1\. In diamond, each carbon atom is linked tetrahe drally to four other carbon atoms by \(\mathrm{sp}^{3}\) bonds. 2\. Graphite has planar hexagonal layers of carbon at oms held together by weak Van der Waal's forces. 3\. Silicon exists only in diamond structure due to its tendency to form \(\mathrm{p} \pi-\mathrm{p} \pi\) bond to itself. In this: (a) Only 1 is correct (b) Only 1 and 2 are correct (c) Only 2 and 3 are correct. (d) All are correct statement here.

Short answer

The correct answer is (b) Only 1 and 2 are correct.

Step-by-step solution

Analyze Statement 1

Statement 1 says that in diamond, each carbon atom is linked tetrahedrally to four other carbon atoms by \( \mathrm{sp}^3 \) bonds. This is true, as diamond's structure consists of a tetrahedral arrangement where each carbon is bonded to four other carbons through \( \mathrm{sp}^3 \) hybridization. This forms a three-dimensional network.

Analyze Statement 2

Statement 2 claims graphite has planar hexagonal layers of carbon atoms held together by weak Van der Waals forces. This is correct. Graphite is composed of layers of carbon atoms arranged in a hexagonal pattern, with weak Van der Waals forces between these layers allowing them to slide over each other.

Analyze Statement 3

Statement 3 suggests that silicon exists only in diamond structure due to its tendency to form \( \mathrm{p}\pi-\mathrm{p}\pi \) bonds. While silicon does form diamond-like structures due to its covalent bonding, it does not typically form \( \mathrm{p}\pi-\mathrm{p}\pi \) bonds. The statement is therefore incorrect regarding the bond type's influence.

Conclusion

Based on our analysis: Statement 1 is correct, Statement 2 is correct, but Statement 3 is incorrect because silicon does not favor \( \mathrm{p}\pi-\mathrm{p}\pi \) bonds as carbon does.

Key concepts

Diamond Structure

Diamond is a fascinating form of carbon known for its striking beauty and remarkable physical properties. In diamond, every carbon atom makes connections with four other carbon atoms. This forms a strong and stable tetrahedral structure. Each of these bonds is known as an sp³ bond, which means that the carbon's s orbital and three p orbitals mix to form four equivalent hybrid orbitals. These orbitals are arranged so that each bond points towards the corners of a tetrahedron, creating a very dense and tough crystal lattice.
This three-dimensional network results in diamond's well-known characteristics such as extreme hardness and high thermal conductivity. Because these bonds are exceptionally strong and evenly distributed, diamonds are not only durable but also have unique optical properties, making them brilliant pieces in jewelry.
Interestingly, the strong structure of diamond also makes it a very poor conductor of electricity. This is because there are no free electrons to move around, as they are all involved in maintaining the tight bonding between the atoms.

Graphite Structure

Graphite is another allotrope of carbon, distinctively different from diamond despite being made from the same element. Unlike the three-dimensional structure of diamonds, graphite consists of two-dimensional planes of carbon atoms arranged in a hexagonal lattice. Within each plane, the carbon atoms form strong covalent bonds with three neighboring atoms using sp² hybridization. This sp² bonding involves mixing one s orbital and two p orbitals to form three planar orbitals that make connections with neighboring carbon atoms.
The layers in graphite are not tightly bonded to each other. Instead, they are held together by very weak van der Waals forces, which allow these layers to slide over one another easily. This sliding attribute gives graphite its characteristic properties, such as a slippery feel and excellence as a lubricant. Additionally, this structure allows graphite to conduct electricity, as electrons can move freely along the planes.

• This layered structure makes graphite useful in pencils, where the layers peel off to leave a mark on paper.
• Graphite is also valuable in industries as a conductor of electricity and high-heat resistance material.

Silicon Crystal Structure

Silicon shares some similarities with diamond, especially with its crystal structure. Like diamonds, silicon atoms are arranged in a lattice where each silicon atom is covalently bonded to four other silicon atoms, creating a repeating tetrahedral structure. This arrangement is often called the diamond cubic crystal structure due to its resemblance to a diamond crystal.
However, a distinctive difference lies in the bonding properties. While carbon atoms in diamonds engage in very strong covalent sp³ bonds, silicon atoms form slightly weaker covalent bonds. This slight difference is due to silicon's larger atomic size compared to carbon, which affects the overlap of orbitals when forming bonds.

• The diamond-like structure of silicon is crucial in the semiconductor industry, as it allows for a controllable flow of electricity.
• Silicon's ability to withstand high temperatures while maintaining structure makes it ideal for electronic components and photovoltaic cells.

Silicon doesn't naturally form π-bonds like carbon does in certain conditions, making the initial statement about its (\( ext{p} ext{π}- ext{p} ext{π} \)) bond tendency incorrect. Instead, its utility and function largely lie in applications where stability and conductivity are needed.