Allotropes of carbon including fullerenes

Allotropes of carbon

The existence of one element in different forms, having different physical properties, but similar chemical properties is known as allotropy. Different forms of an element are called 'allotropes' or allotropic forms. Carbon shows allotropy. The various allotropic forms of carbon can be broadly classified into two classes.

Crystalline form

Diamond and graphite are crystalline forms.

Amorphous form

Coal, coke, charcoal (or wood charcoal), animal charcoal (or bone black), lamp black, carbon black, gas carbon and petroleum coke are amorphous forms.

Diamond

Diamonds are chiefly found in the Union of South Africa, the Belgian Congo, Brazil, British Guiana, India etc.

Diamond was discovered for the first time in India. The famous 'Kohinoor diamond' (186 - carat) and the 'Regent or Pitt' (studded in Napoleon's state sword, 136.2 carat) were found near Kistna river in South India.

The 'Cullinan diamond', the largest ever found weighed 3025.75 carat (about 600 g) was mined in South Africa in 1905.

Diamonds occur in the form of transparent octahedral crystals usually having curved surfaces and do not shine much in their natural form. To give them their usual brilliant shine they are cut at a proper angle so as to give rise to large total internal reflections.

Moissan (1893) prepared the first artificial diamond by heating pure sugar charcoal and iron in a graphite crucible to a temperature of about 3000°C in an electric arc furnace.

Graphite

Graphite is found widely distributed in nature, viz., in Siberia, Sri Lanka, USA, Canada.

Large quantities of graphite are also manufactured from coke or anthracite in electric furnaces.

Diamonds and graphite are two crystalline allotropes of carbon. Diamond and graphite both are covalent crystals. But, they differ considerably in their properties.

Comparison of the properties of diamond and graphite

These differences in the properties of diamond and graphite are due to the differences in their structures.

Structure of diamond

In diamond, the carbon atoms are arranged tetrahedrally (sp³ hybridisation of C): each C atom is linked to its neighbors by four single covalent bonds. This leads to a three-dimensional network of covalent bonds.

�� Fig:13.1 - Structure of diamond

It is due to this, that diamond is very hard, and has high melting and

boiling points. In diamond, each carbon atom is bonded to the other through regular covalent bonds. The electrons thus are held tightly between the nuclei, and there are no mobile electron to conduct electricity i.e. all the valence electrons of carbon are used up in forming the covalent bonds. Hence diamond does not conduct electricity. Diamond is also denser than graphite (density: Diamond = 3.52 g cm^-3 Graphite =2.25 g cm^-3 ) as the Diamond structure is a closely packed structure, while the layer-to-layer large distance makes graphite to have an open structure.

Structure of graphite

In graphite, the carbon atoms are arranged in flat parallel layers as regular hexagons. Each layer is bonded to adjacent layers by weak Van der Waals forces. This allows each layer to slide over the other easily. Due to this type of structure graphite is soft and slippery, and can act as a lubricant. Graphite is also a good conductor of electricity. In graphite, carbon atoms in each layer are bonded to three other carbon atoms by special covalent bonds. This gives some double-bond character to the C-C bonds. This gives it the presence of delocalized p-electron system. These mobile electrons explain the electrical conductivity of graphite.

Uses of diamond

• The unusual brilliant shine of diamond makes it an invaluable precious stone in jewellery.

• Making high precision cutting tools for use in medical field.

• Because of it's hardness it is used in manufacturing tools/cutting drills for cutting glass and rock.

• In making dyes for drawing very thin wires of harder metals. Tungsten wires of thickness 1/6th that of human hair, can be drawn using diamond dyes.

Uses of graphite

• As a lubricant at higher temperatures.

• As a refractory material of making crucibles and electrodes for high temperature work.

• In electrotyping and in the manufacture of gramophone records: Graphite is used for making the non-conducting (generally wax) surface, so that electroplating can be done.

• For manufacturing lead pencils and stove paints.

Amorphous Forms of Carbon

Some amorphous forms of carbon are described below:

Coal

Coal is formed in nature by the carbonization of wood. Conversion of wood to coal under the influence of high temperature, high pressure and in the absence of air is termed carbonization.

Amongst coal varities, anthracite is the purest form. It contains about 94 - 95% of carbon. The common variety is bituminous coal; it is black, hard and burns with smoky flame.

Uses

Coal is mainly used,

• As an industrial fuel in steel, power generation plants etc. It is also a domestic fuel to a limited extent.

• For manufacture of producer gas and water gas, which are used as fuel gases.

• For manufacturing coal tar, coke and coal gas.

• Anthracite coal is used for preparing graphite.

• For the manufacture of synthetic petrol by catalytic hydrogenation of coal.

Wood charcoal

When wood is heated strongly in a very limited supply of air, wood charcoal is obtained. This is called destructive distillation of wood. The volatile products are allowed to escape.

Charcoal is a black, porous and brittle solid. It is a good adsorbent. Charcoal powder adsorbs coloring matter from solutions and poisonous gases from the air. Charcoal is also a good reducing agent.

Uses

• As a fuel.

• As a deodorant and in gas masks to filter pollution.

• As a discoloring agent for decolorizing oils, etc.

• In making gun powder.

Animal charcoal

Animal charcoal (or Bone charcoal) is obtained by destructive distillation of bones. It contains about 10-12% of amorphous carbon.

Sugar charcoal

It is obtained by heating sugar in the absence of air. Sugar charcoal is the purest form of amorphous carbon.

Sugar charcoal becomes activated charcoal when it is powdered to particle size of about 5 micron and heated at about 1000 K in vacuum. Activated charcoal has an increased adsorption capacity.

Lamp black

Lamp black is manufactured when tar and vegetable oils (rich in carbon) are burnt in an insufficient supply of air and the resulting soot is deposited on wet blankets hung in a room.

Lamp black is a velvety black powder. It is used in the manufacture of Indian ink, printer's ink, carbon papres, black paint and varnishes.

Carbon black

When natural gas is burnt in limited supply of air, the resulting soot is deposited on the underside of a revolving disc. This is carbon black and it is then scraped off and filled in bags.

It differs from lamp black in being not so greasy. Carbon black is added to the rubber mix used for making automobile tyres, and has replaced the use of lampblack for a number of purposes.

Gas carbon and petroleum coke

Carbon scraped from the walls of the retort used for the destructive distillation of coal is called gas carbon. During refining of crude petroleum, petroleum coke is deposited on the walls of the distillation tower.

Both, gas carbon and petroleum coke are used for making electrodes in dry cells and are good conductors of electricity.

Fullerenes

Fullerenes are recently discovered (1985) allotropes of carbon. They have been found to exist in the interstellar dust as well as in geological formations on earth. They are large cage like spherical molecules with formulae C₃₂, C_50, C₆₀, C₇₀, C₇₆, C₈₄ etc. The most commonly known fullerene is C₆₀which is named as 'buckminster fullerene' after the designer of the geodesic dome, American architect Buckminster.

Structure of C

₆₀ molecule

C₆₀ molecule has marvelously symmetrical structure. It is a fused-ring of aromatic system containing 20 hexagons and 12 pentagons of sp² hybridized C atoms. The structure bends around and closes to form a soccer ball shaped molecule. C₆₀ is therefore, called buckyball also. The diameter of ball cage is about 70 pm. It is about 6-10 times as large as H atom is. The ball cages are highly stable and do not break up till 1375 K. It is a highly symmetrical structure in which all the carbon atoms occupy identical position.

Fig: 13.2 - Buckminster fullerene (C₆₀)

Appearance and applications

C₆₀ fullerene looks different from diamond and graphite. It is a yellow powdery substance, which turns pink on dissolution in solvents like toluene. It polymerizes on exposure to U.V. radiations.

Fullerenes are fascinating because they show unusual characteristics and applications like:

• They are wonderful lubricants because the balls can roll between the surfaces.

• Alkali compounds of C₆₀ (A₃C₆₀) are super conducting materials even at high temperatures of the order of 10-40 K.

Problem

4. Diamond is extremely unreactive at room temperature, but graphite reacts more readily, why?

Solution

In diamond, all the four valencies of each carbon atom are satisfied by bonding to other carbon atoms through regular covalent bonds. The electrons are held tightly between the nuclei, and there are no mobile electrons. Thus diamond structure is a closely packed structure and is extremely unreactive at room temperature.

In graphite, carbon atoms in each layer are bonded to three carbon atoms. This gives some double-bond character to C-C bonds giving it the presence of delocalized p-electron system. The layer-to-layer large distance makes graphite to have an open structure and so it reacts more readily.