Name: Relation Between Edge Length and Radius Ratios
Uploaded: 2019-07-19
Duration: 14 min 49 s
Description: The ratio of the radius of cation to anion is called radius ratio. Watch the video to learn the tricks to remember the relation between edge length and radiu...

Question

What are atomic solids? Discuss the structures of diamond and graphite.

Accepted Answer

Atomic solids: They are also referred to as network or covalent solids. The constituent particles of these solids are held by strong covalent bonds. The bonds in these solids are extended throughout the crystal and result in the formation of giant $3$ −D structures. The presence of covalent bonds makes them hard and thus have high melting points. They are usually bad conductors of electricity due to the absence of free electrons except for graphite. Some delocalised electrons allow graphite to conduct electricity and heat. Other examples of covalent solids are diamond, carborundum, and quartz. Atomic solids are also referred to as network or covalent solids.

Structure of diamond and graphite:

Diamond: The structure of diamond is based on fcc but with extra carbon atoms inside the cube. Each carbon atom is ${sp}^{3}$ hybridized that is bonded tetrahedrally to four other carbon atoms by single covalent bonds. The coordination number of each carbon atom is $4$ . Thus, it is giant $3 - D$ molecule in which $C - C$ bond length is $154 pm$ and each bond angle is $109^{o} - 28'$ or $109 . 5^{o}$ . This symmetry of diamond offers the number of characteristics properties that are following:

The strong covalent bonds between the carbon atoms promise the hardness of the diamond.

Melting diamond requires a very high temperature.

Since all the valence electrons are involved in covalent bonding, leaves no free electron to conduct electric current and makes them a bad conductor of electricity.

Graphite: Unlike diamond, graphite has a two-dimensional sheet-like structure. Each carbon atom is ${sp}^{2}$ hybridized and is bonded to the other three carbon atoms through covalent bonds forming hexagonal planer rings. The $C - C$ covalent bond distance is $142 pm$ . Any two successive sheets of graphite are about $340 pm$ apart. Thus, this large distance between the layers do not permit them to make covalent bonds and offers the number of characteristic properties that are mentioned below:

The weak Van der Waal's forces of attraction between the layers make them sliding one over another. Consequently, gives it a soft and slippery texture that can be used as lubricants.

The delocalised electrons in graphite solid serve as thermal or electrical conductors.

Within each layer, carbon atoms are firmly attached by strong covalent bonds. Thus, they have a high melting point.

It has a relatively lower density than diamond due to the large distance between the successive layers.

What are atomic solids? Discuss the structures of diamond and graphite.

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