ALLOTROPES OF CARBON

CONTENT
⦁ Carbon Allotropes and their Structures
⦁ The Atomic Structure of Carbon
⦁ Allotropes of Carbon
⦁ Diamond
⦁ Graphite
⦁ Differences between Diamond and Graphite
⦁ Chemical Properties of Carbon

Carbon Allotropes and their Structures
Carbon is a non-metal found in group 4 of the periodic table. It occurs naturally as diamond and graphite. These physically different forms of the same chemical element are known as allotropes. Other elements that exist in different forms in their free states include tin, sulphur and phosphorus. Also, it occurs in an impure form as coal and in the combined state as petroleum, wood and natural gases.
The Atomic Structure of Carbon
Carbon is the sixth element in the periodic table. the nucleus its atom it composed of six proton and six neutron and surrounded by six electron;2 in the first energy level (k-shell) and four in the 2nd energy level ( L-shell). Its orbital configuration is 1S22S22P2. Carbonatom has four valency electrons. Carbon atoms are able to catenate, i.e. join to one another by covalent bonds in a continuous fashion. The phenomenon is called catenation.
Allotropes of Carbon
Allotropy is the existence of an element in the same physical state but indifferent forms. The various forms are called allotropes. Carbon exhibits allotropy. Other elements that also exhibit allotropy are sulphur, Tin, Oxygen and Phosphorus. The two crystalline allotropes of carbon are: Diamond and graphite. Charcoal is amorphous allotrope of carbon.
Diamond
Structure and Bonding
In diamond, each carbon atom is at the center of a regular tetrahedron and covalently bonded to four other atoms in a strong compact fashion. The covalent linking between the atoms is continuous to produce a single infinite tightly locked 3-dimensional giant (macro) molecule, which has a network octahedral structure see the diagram below:

The rigidity of the structure is responsible for the hardness of diamond; it is the hardest known substance. Since the covalent bonds are strong and the molecule is compact, the melting point of diamond is very high.
Physical Properties of Diamond
⦁ Pure diamond is hard, colourless and transparent.
⦁ Forms octahedral crystals with high refractive index.
⦁ It is very hard; it has a density of 3.5gcm-3, and melting point of 36000C.
⦁ It does not conduct electricity, because all the four valence electrons per carbon atom are used in bonding i.e.no mobile electrons.
⦁ It is insoluble in any solvent.
Uses of Diamond
⦁ Because of its high refractive index and amazing metallic luster when cut, it is used in jewelries.
⦁ Because of its hardness, diamond is used in cutting glasses, in drilling rocks; in boring of holes; in making bearings in engines; and as an abrasive, i.e. to smoothen rough surfaces.

Graphite
Structure and Bonding
In graphite, each carbon atom is covalently bonded to three other atoms to produce an infinite two-dimensional flat hexagonal layer structure, which is strong and hard. See the diagram above. The flat hexagonal layers in graphite are held together by the weak van der Waals attractive forces, which allow movement of the planes parallel to each other, and make the graphite to be soft and slippery. The fourth electron in the valence shell of each carbon atom in graphite is mobile, because it is not used in bonding, and account for its electrical conductivity.
EVALUATION
⦁ Define the term allotropy.
⦁ What is the structure of? (a) diamond (b) graphite.
⦁ Give the reason why diamond is hard, while graphite is soft.
Physical Properties of Graphite
⦁ Graphite form soft, black and opaque hexagonal crystals, which are greasy to feel. The softness is due to the ability of the adjacent layers to slide over one another.
⦁ It is hard. Its density is 2.3gcm-3, and melts at about 35000
⦁ It is good conductor of heat and electricity due to the presence of a mobile electron per carbon atom.
⦁ It is soluble in any solvent. Graphite is an example of a non-metallic conductor. It is a metalloid.
Uses of Graphite
⦁ Graphite is used as a lubricant, because of its flat hexagonal layer which can slide over one another.
⦁ Used as inert electrodes during electrolysis and for brushes of electric motors been a good conductor of electric current.
⦁ When mixed with clay, graphite forms lead, which is used in making lead pencils. The hardness of a pencil depends on the amount of clay in the mixture. Soft pencils contain more of graphite, while hard pencils contain more of clay.
⦁ Used in making crucible, because of its high melting point.
⦁ Used in nuclear reactors; being soft and with a high melting point.
Differences between Diamond and Graphite

S/N DIAMOND PROPERTIES GRAPHITE PROPERTIES
1 Diamond is a transparent
solid that sparkles when
cut and polished Graphite is an opaque solid,
with a metallic luster
2 It is octahedral in shape It is hexagonal in shape
3 Its density is 3.5 Its density is 2.3
4 It is a poor conductor of
electricity It is a good conductor of
electricity
5 It is an inert substance but at
9000c, it burns in air to form
carbon(iv) oxide and
combines with fluorine It is a more reactive substance.
It burns in air to form
carbon (iv) oxide at 7000c.
It also reacts with oxidizing agents
to form oxides, it also reacts
with fluorine and
tetraoxosulphate(vi) acid
6 It is the hardest substance
known It is one of the softest minerals
known.

EVALUATION
⦁ List 3 uses of graphite with reasons.
⦁ Give four differences between diamond and graphite.
Chemical Properties of Carbon
1. Combustion reaction
Carbon burns in limited supply of air to form carbon (II) oxide equation for the reaction is given as
2C(s) + O2(g)    →      2CO(g)
In excess air, complete combustion takes place and carbon(Iv) oxide is formed. Equation for the reaction is given as.
C(s) + O2    →     CO2(g)
When charcoal is used as fuel, it burns released is used for cooking (exothermic reaction)
NOTE: Carbon occurs most abundantly both naturally as diamond and graphite and in numerous compounds including petroleum coal, natural gas among others.
2. Combination reaction:
Carbon combines directly with certain elements such as Sulphur, hydrogen, calcium and aluminium at a very high temperature.
C  +  2S    →    CS2
2C  +  Ca   →   CaC2
3. As a reducing agent:
Carbon is a strong reducing agent. It reduces the oxides of a less active metals.
Fe2O3  +  3C   →   2Fe   +  3CO
4. Reaction with strong oxidizing agent. Trioxonitrate (v) acid oxidizes carbon to carbon (iv) oxide
4HNO3   +  C    →    CO2  +   4NO2   + 2H2O