STRUCTURE OF DIAOND
In diamond, each carbon atom is sp3 hybridised and is linked tetrahedrally to four other carbon atoms. C-C bond lengths are equal to 1.54 A (154 pm) and angle is 109° 28′. As result, in diamond there is a three dimensional network of strong covalent bonds (Fig. 14.2). This makes diamond extremely hard with very high melting point (3843 K) and high density 3.51 g/cm3. Diamond is a poor conductor of electricity because all the valence electrons of carbon are involved in carbon-carbon sigma covalent bonds and hence are localized and not free to conduct electricity.
USES OF DIAMOND
Diamond is the hardest naturally occurring substance. Due to its extreme hardness, it is used for making tools for cutting and grinding other hard materials and also used in oil-well drills for drilling holes through the earth’s rocky layers.
It is used in dies for the manufacture of tungsten filaments for electric light bulbs.
Diamond is transparent and has very high refractive index. It possesses extraordinary brilliance. Because of these properties it is used for making jewellery.
Sharp-edged diamonds are used by eye surgeons to remove cataract from eyes with high precision.
5. Because of its extraordinary sensitivity to heat rays, it is used in high precision thermometers.
6. It is used in protective windows for space probes as it can keep out harmful radiations.
The value of diamond depends upon its weight and freedom from impurities. Weight of diamond is expressed in terms of carats. One carat is equal to 0.2 g or 200 mg.
STRUCTURE OF GRAPHITE
In graphite, each carbon atoms is sp2 hybridised and is linked to three other carbon atoms directly in the same plane to form hexagonal rings. These rings constitute huge sheets or layers of atoms as shown in Fig. 14.3. The unhybridized p-orbitals overlap sidewise with each other to form 1t clouds. The C-C bond length in rings is 142 pm which indicates the presence of 1t bond character. The different sheets of carbon atoms are held by weak van der Waals’ forces and are separated by a distance of 340 pm. These layers can easily slide over one another. This explains the slippery nature of graphite and its use as lubricant. The electrical conductivity of graphite is attributed to the mobile (delocalized) 1t electrons. The low density (2.26 g/cm3) of graphite is due to large distance between different layers of carbon atoms.
USES OF GRAPHITE
Graphite is used as lubricant either as a powder or as a dispersion in oil or water. The dispersion of graphite in oil is known as oil dag and in water is known as aqua dag.
Mixed with clay it is used in ‘Lead’ pencils.
Since it is good conductor of electricity and is inert, it is used for making carbon electrodes in electrolytic cells and in dry cells.
Because of its high melting point it is used for making graphite crucibles. Crucibles made of graphite are not attacked by dilute acids or fused alkalies.
It is a component of printers’ ink.
It is used as moderator in nuclear reactor.
Example 14.1. Identify the type of chemical bonds in Fe, NaCl, SiO2‘ I2, diamond, graphite.
Substance Type of solid Interparticle bonds
Fe Metal Metallic bonds
NaCI Ionic compound Ionic bonds
Si02 Network solid Covalent bonds
I2 Molecular solid Dispersion. forces
Diamond Network solid Covalent bonds
Graphite Network solid Covalent bonds and
Example 14.2. Why solid NaCl does not conduct electricity while solid iron does?
Solution. Sodium Chloride is ionic solid in which the position of ions is fixed. Due to immobility of ions in solid NaCI it is not able to conduct electricity.
Iron, on the other hand is metal in which positive kernels occupying fixed positions are surrounded by moving electrons. The free movement of electrons in metal crystals is responsible for the conduction of electricity.