Classwork Series and Exercises {Chemistry – SS1}: Carbon

Week 4

Topic: Carbon

Introduction

Carbon forms the largest number of compounds, next only to hydrogen. It ranks seventeenth in the order of abundance in the earth’s crust. Carbon occurs in the free native state as well as in the combined state. Carbon and its compounds are widely distributed in nature.

In its elemental form, carbon occurs in nature as diamond and graphite. Coal, charcoal and coke are impure forms of carbon. The latter two are obtained by heating wood and coal in the absence of air, respectively. In the combined state, carbon is present as carbonate in many minerals, such as hydrocarbons in natural gas, petroleum etc. In air, carbon dioxide is present in small quantities, (0.03%).

Our food also contains carbon in the combined form. All living systems contain carbon compounds. Indeed, life as we know today, would be impossible without such carbon compounds.

Carbon is a non-metallic element and the first member of group 4 of the periodic table

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.

Structure of diamond

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

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.

Properties of diamond

• Hardest substance known to man
• Brittle (not malleable)
• Insulator (non-conductor)
• Insoluble in water
• Very high melting point

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.

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.

 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

 Properties of Graphite

• It conducts electricity
• It writes well on paper
• It has a density of 2.3g/cm³    
• It is a soft black material which feels greasy to the touch

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.

 Comparison of the properties of diamond and graphite

Diamonds and graphite are two crystalline allotropes of carbon. Diamond and graphite both are covalent crystals. But, they differ considerably in their properties.These differences in the properties of diamond and graphite are due to the differences in their structures.