Problem 191 - A Complete Resource Book in Chemistry for JEE Main

Question

Graphite is a soft, solid, lubricant, extremely difficult to melt. The reason for this anomalous behaviour is that graphite [2003] (a) has molecules of variable molecular masses like polymers(b) has carbon atoms arranged in large plates of rings of strongly bound carbon atoms with weak interplate bonds (c) is an allotropic form of diamond (d) is a non crystalline substance

Step-by-Step Solution

Answer

The correct answer is (b) because graphite has strong in-plane bonds and weak inter-layer bonds.

Step 1: Identifying the Properties of Graphite

Graphite is known for its unique structure. It is composed of layers, where each layer has carbon atoms arranged in a hexagonal pattern. Between these layers, the bonds are very weak, which allows the layers to slide over each other easily, making it an excellent lubricant.

Step 2: Analyzing Each Option

Let's look at the options: (a) talks about variable molecular masses like polymers, which is not typical for graphite. (b) correctly describes graphite's structure with strong intra-layer carbon bonds and weak inter-layer bonds. (c) mentions graphite as a form of diamond, which is true, but doesn't explain the behavior. (d) says graphite is non-crystalline, which is incorrect since it has a crystalline structure within each layer.

Step 3: Choosing the Correct Option

Based on the analysis, option (b) provides the accurate explanation for graphite's properties. The arrangement in large plates with strong in-plane bonds and weak out-of-plane bonds explains why graphite is soft and a good lubricant but hard to melt.

Key Concepts

Allotropy

Graphite is an interesting example of allotropy in carbon. Allotropy refers to the property of some chemical elements to exist in two or more different forms, known as allotropes, in the same physical state. For carbon, its most well-known allotropes are diamond and graphite.

Despite both being composed entirely of carbon atoms, diamond and graphite have very different physical properties. In diamond, each carbon atom forms four strong covalent bonds in a tetrahedral structure, making it extremely hard and not electrically conductive. On the other hand, graphite forms layers of carbon atoms in a hexagonal arrangement where the atoms within each layer are strongly bonded, while the layers themselves are held together by weaker forces. This difference in structure between the allotropes leads to their vastly different characteristics, such as graphite's softness and lubricating ability, contrasting with diamond's hardness.

Hexagonal Structure

In graphite, each carbon atom is bonded to three other carbon atoms in a plane to form a flat, hexagonal lattice. This is a geometric arrangement where each hexagon is made up of six carbon atoms. These flat, interconnected hexagons stack on top of one another to form layers. The remarkable thing about this hexagonal structure is how it contributes to the properties of graphite.

The strong covalent bonds within each hexagonal layer make the layers themselves quite robust. However, it's the arrangement into a hexagonal pattern that allows the layers to overlap and stack. This type of bonding and layering gives graphite its characteristic slipperiness, as the layers can easily slide over each other. This property is crucial for the use of graphite in applications like pencil lead and lubricants. Though the individual layers are strong due to their covalent bonds, the overall structure is soft, once again highlighting the unique nature of graphite's hexagonal arrangement.

Interlayer Bonds

The interlayer bonds in graphite are crucial in defining its mechanical and chemical properties. These bonds are significantly weaker than the covalent bonds within the layers. As a result, while the hexagonal carbon layers are strong individually, the layers themselves can easily shift over each other. This weakness between the layers is due to van der Waals forces, which are much weaker than covalent bonds.

Van der Waals forces in graphite make it possible for the layers to glide, undoubtedly related to graphite’s effectiveness as a lubricant. Since the layers are not tightly bound, they can move with minimal resistance, allowing for the smooth glide over surfaces. This is one reason why graphite is used in places where reducing friction is necessary. Additionally, the weak bonds between layers attribute to graphite's very high melting point - considerable energy is needed to overcome even the weak interlayer attractions when transitioning graphite from a solid to a liquid state. This characteristic makes graphite a very stable form of carbon.