Allotropes of Carbon

Carbon is a non-metal that is necessary for life. The Latin word ‘carbo’ which means ‘coal,’ is the source of the term. This is because it is the most important component of coal. Carbon is the most prevalent chemical in most organic stuff, from fossil fuels to complex molecules. The Earth’s crust contains just 0.02 % carbon in the form of minerals, and carbon dioxide makes only 0.03% of the atmosphere. Despite the fact that there is only a little quantity of carbon in nature, the relevance of carbon appears to be enormous: carbon compounds are found in fossil fuels, polymers, soaps, detergents, and the majority of the medications we use.

Bonding in Carbon: Carbon has an atomic number of 6, which indicates that the electronic configuration of the carbon atom is 2,4. Because a carbon atom’s outermost shell has four electrons, it shares those electrons and reaches the inert gas state. As a result, the carbon atom creates covalent connections with other atoms.

What are Allotropes?

Allotropy is the occurrence of an element existing in two or more forms with distinct physical qualities but similar chemical properties, and the different forms are known as allotropes.

Allotropes of Carbon

Allotropy is a phenomenon that occurs in carbon, and it appears in two sorts of allotropic forms:

• Carbon in crystalline allotropic forms: Carbon has four allotropes with well-defined crystal structures are,
  • Diamond
  • Graphite
  • Fullerenes
  • Carbon Nanotubes
• Carbon in amorphous allotropic forms: Amorphous allotropic forms of Carbon include:
  • Coal
  • Coke
  • Wood charcoal
  • Animal charcoal
  • Sugar charcoal
  • Lampblack
  • Gas Carbon

Diamond

Diamond is a kind of carbon that has its atoms organised in a diamond cubic crystal structure. Another solid form of carbon known as graphite is the chemically stable form of carbon at normal temperature and pressure, although diamond nearly never transforms to it.

Diamond has the highest hardness and thermal conductivity of any natural substance, characteristics that make it ideal for cutting and polishing equipment in the industry.

Structure of Diamond

Jewel has a three-dimensional organization of carbon molecules combined through solid covalent bonds. Every Carbon iota is in the condition of sp³ hybridization and connected tetrahedrally to four adjoining Carbon molecules. This organization stretches out into three measurements. All Carbon-Carbon (C–C) bonds are equivalent and equivalent to 154 pm, and every C–C–C bond point is 109^∘28′.

Structure of Diamond.

Properties of Diamond

• It is the most perfect and densest assortment of Carbon. Its thickness is 3.51 gcm^–3.
• It is the hardest regular known substance and has an extremely high softening point (3843K) It is insoluble in all solvents.
• It is straightforward and measures a high refractive record.
• It is a terrible conveyor of power. This is on the grounds that all the valence electrons of every Carbon are engaged with the Carbon-Carbon Sigma (σ) bonds, and no unpaired electrons are left in the precious stone.
• Synthetically, jewel is impervious to practically all acids, antacids, and salts. In any case, it is followed up by melded Sodium Carbonate. When warmed with a combination of potassium dichromate and sulphuric corrosive to 475 K, it gradually gets oxidized to Carbon dioxide.
• The worth of a jewel relies on its size and shading. Pale blue-white jewels are more valuable than those having a low tone. Dark precious stones are the least expensive and not utilized in adornments.

Uses of Diamond

1. Diamonds are used to cut glass cutters, marble saws, and rock drilling tools, among other things.
2. Because of their exceptional brightness, diamonds are utilised in jewellery.
3. Sharp-edged diamonds are used by eye surgeons to remove cataracts from the eyes with remarkable precision.
4. Diamond dies are used to draw very thin metal wires such as tungsten.

Graphite: A Non-metallic Conductor

Each Carbon atom in graphite is in a state of sp2 hybridization, which means it is covalently linked to three other carbon atoms in the same plane. Planar hexagonal rings are produced as a result. The length of the Carbon-Carbon bond in this ring is 142pm.

Layers are formed by hexagonal rings. Van der Waal’s force holds the layers together, while 142pm separates them. Because these layers are able to slide over one another, graphite is soft and lubricious.

Properties of Graphite

1. It has a metallic sheen and is dark grey in colour.
2. Touching it feels extremely soft and oily.
3. The fourth valence electron of each Carbon is free to travel since only three electrons of each Carbon are required to make hexagonal rings in graphite. As a result, graphite is a strong heat and electrical conductor.
4. Dilute acids, alkalis, and chlorine do not harm it. It is slowly oxidised to carbon dioxide using a combination of potassium dichromate and sulphuric acid.

Uses of Graphite

• It’s used to make carbon arcs and electrodes.
• It’s a lubricant for equipment that operate at high temperatures.
• Lead pencils are made with this material. Graphite powder is combined with clay and formed into sticks. Pencils are made from these sticks.
• It is employed in atomic reactors as a moderator.
• It’s utilised in steel production as a reducing agent.
• It’s a component of high-strength composites.
• It’s utilised to make Crucibles, which can resist extremely high temperatures.

Buckminsterfullerenes: A Synthetic Allotrope of Carbon

Fullerenes are the main unadulterated type of Carbon since they don’t have astonishing edges or surface securities that draw in different iotas, as on account of graphite or jewel. 

Fullerene is an enormous circular particle of arrangement C_2n where n≥30. Fullerene is for all intents and purposes delivered by warming graphite in an electric bend in an inactive gas, for example, helium or argon when a dirty material is framed by the buildup of C_n little atoms.  

The dingy material so shaped predominantly comprises of C₆₀ with a more modest amount of C₇₀ and hints of different fullerenes. The C₆₀ and C₇₀ Fullerenes can be promptly isolated from the fullerenes residue by extraction with benzene or toluene followed by chromatography over alumina.

Structure of C₆₀ Fullerenes

Structure of fullerenes

In honour of American architect Robert Buckminster Fuller, the C60 fullerenes are sometimes known as Buckminsterfullerene or simply fullerene. Fullerene is a 60-vertices saucer-ball-shaped molecule having a Carbon atom at each vertex. It has 20 rings with six members and 12 rings with five members.

Six-membered rings can be fused to other six-membered rings as well as five-membered rings, whereas five-membered rings can only be fused to other six-membered rings. Single and double bonds with Carbon-Carbon lengths of 142pm and 138.3pm, respectively, are present.

Properties of Fullerene

1. When the temperature is changed, the behaviour and structure of fullerene changes. The fullerene is transformed to the C₇₀ form at a greater temperature.
2. Under changing pressures, the structure of fullerene changes.
3. Fullerene has an ionisation enthalpy of 7.61 electron volts.
4. Fullerene has an electron affinity of 2.6 to 2.8 electron volts.
5. In chemical processes, fullerene (C₆₀) mimics an electrophile.
6. Fullerene has the ability to behave as an electron acceptor. It can readily receive three or more electrons. As a result, it has the potential to act as an oxidising agent.
7. To achieve superconductivity, fullerenes are doped with alkali or alkaline earth metals.
8. Fullerene has the property of ferromagnetism.
9. Fullerene is abundant in carbon compounds. It’s extremely soluble in organic solvents as a result of this.

Use of Fullerene

• Conductors made of fullerene are utilised.
• It has the ability to absorb gases.
• Lubricants made of fullerene are utilised.
• Fullerenes are utilised in the manufacture of cosmetics-related products in several forms.
• Graphene sheets make up carbon nanotubes.
• Fullerenes are utilised in biological applications in some ways.

Sample Questions

Question 1: What is the purest form of Carbon?

Answer:

Because it lacks the glittering edges and surface bonds that attract other atoms seen in graphite and diamond, fullerene is the cleanest form of carbon.

Question 2: How is fullerene obtained?

Answer:

When graphite is heated in an electric arc in an inert atmosphere such as helium or argon, a sooty substance is produced by the condensation of C_n tiny molecules, resulting in fullerene. Extraction with benzene or toluene followed by chromatography over alumina separates the C₆₀ and C₇₀ fullerenes found in sooty material from the fullerenes soot.

Question 3: What are the uses of Graphite?

Answer:

1) It’s used to make carbon arcs and electrodes.

2) It’s a lubricant for equipment that operate at high temperatures.

3) Lead pencils are made with this material. Graphite powder is combined with clay and formed into sticks. Pencils are made from these sticks.

4) It is employed in atomic reactors as a moderator.

5) It’s utilised in steel production as a reducing agent.

6) It’s a component of high-strength composites.

7) It’s utilised to make Crucibles, which can resist extremely high temperatures.

Question 4: How does carbon form a bond with other atoms?

Answer:

Carbon has an atomic number of 6, which indicates that the electronic configuration of the carbon atom is 2,4. Because a carbon atom’s outermost shell has four electrons, it shares those electrons and reaches the inert gas state. As a result, the carbon atom creates covalent connections with other atoms.

Question 5: What are the uses of Diamond?

Answer:

• Diamonds are used to cut glass cutters, marble saws, and rock drilling tools, among other things.
• Because of their exceptional brightness, diamonds are utilised in jewellery.
• Sharp-edged diamonds are used by eye surgeons to remove cataracts from the eyes with remarkable precision.
• Diamond dies are used to draw very thin metal wires such as tungsten.