Carbon: Structure and Allotropes of Carbon

Carbon in the Periodic Table and Electronic Configuration

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Carbon is a the first member of group IV elements (or group 14 in modern IUPAC numbering) and the fourth element in period 2 of the periodic table. It has the chemical symbol "C", with an atomic number - 6, and atomic (or mass number) - 12.

Carbon has the electronic configuration  →1s² 2s² 2p², which can be simply written simply as "2, 4". The 2, 4 electronic configuration of carbon implies that it has 2 electrons in the first shell (K-shell) and 4 electrons in the outer (valence) shell (L-shell).

Note: The electronic configuration of elements is significant (important) in chemistry because it helps us understand the behavior and properties of each (respective) element.

Please read on families of elements in the periodic table here.

 

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Importance of the Electronic Configuration of Carbon

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• Carbon has 4 valence electrons, making it tetravalent. As a consequence, it forms 4 covalent bonds.

• Being in Group IV, it is equidistant from achieving a noble gas configuration by either losing or gaining (4) electrons, so it almost always forms covalent bonds by sharing electrons (rather than by accepting or giving out 4 electrons through electrovalent (ionic) bonding).

• Carbon atoms can bond to one another repeatedly to form long chains and rings. This property of carbon is called catenation. For this reason, carbon forms millions of compounds (as studied in organic chemistry) and several structural forms of the element itself (allotropes).

You can read a brief introduction to organic chemistry here.

 

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Structure and Allotropes of Carbon

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Allotropy is defined the existence of an element in two or more different physical forms in the same physical state, with the different forms (allotropes) having different physical properties but the same chemical properties.

Carbon exists as allotropes. These are grouped into:

1. Crystalline allptropes: The carbon atoms are arranged in a definite, orderly, repeating pattern (lattice): Examples of crystalline allptropes of carbon include diamond, graphite, fullerenes and graphene.

2. Amorphous allptropes: Carbon atoms are not arranged in a regular lattice: Examples of amorphous carbon allptropes include coal, coke, charcoal, soot (lamp black), and animal charcoal.

Please read on detail the amorphous allotropes of carbon here

 

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Crystalline allptropes of Carbon (Diamond)

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Structure of Diamond

Diamond is a crystalline allotrope of carbon with each carbon atom bonded covalently to 4 other carbon atoms arranged tetrahedrally (bond angle 109.5°). This forms a rigid, three-dimensional (3D) giant covalent (macromolecular) crystal lattice structure (known as the diamond cubic crystal system) extending throughout the crystal.

Each carbon atom making up the diamond molecule is sp³ hybridized and covalently bonded to four other carbon atoms pointing toward the corners of a regular tetrahedron.

The molecular structure of diamond is represented via the image below.

Note: All bonds are strong single covalent bonds of equal length.

You can read on salts in chemistry, their classification, solubility, preparations and colors here.

 

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Properties of Diamond based on its Structure

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• Diamond is extremely hard. In fact, it is the hardest known natural substance on earth. Diamond is very hard because of its rigid 3D network of strong covalent bonds. It's hardness rates a 10 on the Mohs scale (which is the maximum rating on this scale for measuring mineral hardness). 

• It has a high melting point (~3550 °C). In this regard, a lot of energy is needed to break the many strong covalent bonds.

• Diamond does not conduct electricity. This is so because all 4 valence electrons of each carbon atom are used in bonding; and as a result, there are no free, delocalized or mobile electrons.

• Diamond has a high thermal conductivity. In fact, it conducts heat better than any other material, even silver.

• It is insoluble in water and organic solvents.

• It has a high density and high refractive index. This is the reason for its brilliance when cut. It also bends light sharply to create colorful flashes of fire and brilliance. The image below shows this:

• Diamond is made entirely of carbon atoms linked in a rigid 3D covalent crystal structure

You can read on the physical and chemical properties of acids, test and uses of acids here.

 

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Uses of Diamond

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• Diamond is used as cutting, drilling, and grinding tools (e.g glass cutters, drill bits, diamond saws).

• It is used in the production of jewellery (gemstones).

• It is used in polishing hard materials, gems, and carbide tools.

• Diamondsm is used in the manufacture of surgical blades (like scalpel knives) which  provide ultra-fine, precise edges used in delicate procedures like eye surgery (e.g, cataract removal).

• They have the highest thermal conductivity of any natural material, making them ideal for drawing heat away from microprocessors and lasers.

• Nanodiamonds (tiny diamond particles) can help track and transport cancer-fighting drugs inside human cells.

• Windows made of diamond can protect vacuum chambers and high-power laser windows from intense heat and pressure.

You can read on pH indicators, types and importance of pH and pOH here.

 

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Crystalline allptropes of Carbon (Graphite)

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Structure of graphite

Graphite is a crystalline allotrope of carbon with each carbon atom bonded covalently to 3 other carbon atoms in the same plane, forming hexagonal rings. These hexagons join to form flat layers (or sheets).

Note: The 4th electron of each carbon atom is not used in bonding within the layer, making it it delocalized (free and mobile); and moves between the layers. This is shown below:

The layers are held together by weak Van Der Waals forces, allowing them to slide over one another.

You can read on bases and alkalis, their preparation, test and uses here.

 

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Properties of Graphite based on its Structure

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• Graphite has a greyish-black, opaque appearance with a metallic lustre.

• Graphite is a good conductor of heat and electricity due to the delocalized (free) electrons that can move along the layers.

• Graphite has a high melting point (~3600–3700 °C), and it sublimes (turns directly from a solid into a gas without turning into a liquid first). Meanwhile, understand that the covalent bonds within each layer are strong, giving rise to its high melting point.

• The layers of graphite are slippery, and slide over each other easily because of the weak forces between them.

• Graphite is lighter (less dense) than diamond.

• It remains stable at low temperatures, but in the presence of oxygen at high temperatures (above 700 °C), it oxidizes to form carbon dioxide. (This is a chemical property of graphite).

• It resists attack from most acids, alkalis, and organic solvents under normal conditions.

Please read on bases and alkali, their classification, physical and chemical properties here.

 

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Uses of Graphite

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• It is used as lead in pencils (when mixed with clay). Simply put, it is mixed with clay to make the dark marks in pencils.

• Fine graphite powder is used as a dry lubricant for machine parts to stops squeaks and lubricates locks without catching dirt. It is useful where oil would attract dust. 

• It is used as electrodes in electrolysis and dry cells (because it conducts electricity and resists heat). Please read ionic theory and electrolysis here.

• Graphite is a key material used in the anodes (negative side) of lithium-ion batteries.

• It is used in the manufacture of crucibles for melting metals (because it withstands high temperatures).

• It is used in steelmaking where it adds carbon to steel and lines high-heat furnaces since it can withstand extreme temperatures.

• Graphite fibres mixed with plastic helps to build strong, and  light tennis rackets, and golf clubs.

• Since it conducts electricity and heat well, it is used in cooling parts and electrical brushes.

 

The table below shows the differences between diamond and graphite:

Feature	Diamond	Graphite
Structure	3D tetrahedral network	2D hexagonal layers
Bonds per Carbon	4 covalent bonds	3 covalent bonds
Appearance	Clear and transparent	Greyish-black and opaque
Hardness (Mohs)	10 (hardest known natural substance)	1–2 (soft and greasy)
Electrical Conductivity	Insulator (non-conductor)	Conductor (due to free electrons)
Density	~ 3.5 g/cm³	2.2–2.3 g/cm³
Hybridisation	sp³	sp²
Interlayer Forces	None (rigid 3D frame)	Weak van der Waals forces
Common Uses	Jewelry and cutting tools	Pencils and lubricants

 

You can read on the differences between physical and chemical changes in chemistry here.

 

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Crystalline allptropes of Carbon (Fullerenes)

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Fullerene (C₆₀) was discovered more recently (in 1985) by the scientists: Harold Kroto, Richard Smalley, and Robert Curl . They won the Nobel Prize in Chemistry for their discovery in 1996.

The molecules of fullerenes such as buckminsterfullerene or buckyball (C₆₀) consists of carbon atoms arranged in closed hollow cage-like structures (spheres, ellipsoids, or tubes).

A fullerene is an allotrope of carbon whose molecules consist of carbon atoms connected in closed cages, tubes, or ellipsoids. 

 

​​​​​​Structure of Fullerenes

Fullerene is made entirely of carbon atoms joined by single and double bonds. The carbon atoms form fused 5-membered (pentagonal) and 6-membered (hexagonal) rings. The pentagon rings force the structure to curve into a closed cage. Below is an image of fullerene C₆₀.

The shapes of common forms of fullerenes include spherical buckyballs C₆₀ and cylindrical carbon nanotubes.

Note: Each carbon atom in fullerene is sp² hybridized and bonded to three other atoms.

 

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Types of Fullerene

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• Buckminsterfullerene (C₆₀)

• Higher Fullerenes (C₇₀, C₇₆, C₇₈ and C₈₄).

• Carbon Nanotubes (CNTs)

• Heterofullerenes

• Endohedral Fullerenes

• Fullerenols

• Megatubes

You can read on the uses of chemistry here.

 

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Properties of Fullerenes

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• They are dark-colored crystalline solids or powders at room temperature. The buckminsterfullerene C₆₀ is a blackish-brown solid.

• Buckminsterfullerene (C₆₀) is composed of fused pentagonal and hexagonal carbon rings. The carbon atoms arranged in 12 pentagons and 20 hexagons, looking like a football.

• Compared to diamond and graphite, they have a relatively low density (around 1.65g/cm³) for C_60.

• They have hollow interior and can trap other atoms or small molecules inside (endohedral fullerenes).

• They can withstand high pressures and return to their original shape.In this regard, they are highly compressible, have high tensile strength, can withstand high pressures, returning to their original shape. They are resistance to deformation.

• Fullerenes can behave as semiconductors (of heat and electricity) when mixed (doped) with alkaline earth metals like potassium and cesium.

• They are  hydrophobic and completely insoluble in water. However, they are soluble in specific organic solvents like carbon disulfide, toluene and benzene.

• Although, individual molecules of fullerene are hard, a bulk solid fullerene can act as a soft, slippery lubricant (which acts like microscopic ball bearings).

• Pure fullerenes are poor conductor of heat and electricity.

• They have high electron affinity. This implies that they can easily accept electrons.

You can attempt various questions in chemistry here.

 

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Uses of Fullerenes

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• They are used in medicine as targeted drug delivery, research and as antioxidants or antiviral agents.

• They are needed in nanotechnology. In this regard, they are applied in electronics, sensors, and solar cells.

• They are utilized in water purification systems to trap pollutants.

• They can be used  as lubricants.

Please read the continuation of this topic on (amorphous allotropes of carbon) here.