Understanding the intricacies of covalent structures is pivotal for grasping the differences in properties of some elemental forms. This becomes especially clear when looking at the allotropes of carbon and comparing them with silicon structures.
Silicon and Silicon Dioxide
Silicon
- Atomic Number: 14
- Electron Configuration: [Ne] 3s² 3p²
Silicon exists primarily in a giant covalent structure where each silicon atom is bonded tetrahedrally to four other silicon atoms. This formation leads to a stable, yet hard and brittle crystalline structure. Due to its strong covalent bonds, silicon has a high melting and boiling point.
Image courtesy of Greg Robson
Silicon Dioxide (SiO₂)
- Common Name: Quartz
- Structure: Each silicon atom is bonded to two oxygen atoms forming a continuous three-dimensional network of SiO₄ tetrahedra.
Silicon dioxide is exceptionally stable and hard. It also possesses a high melting point due to the strong covalent bonds between the silicon and oxygen atoms. SiO₂ is predominantly found in nature as sand and is used extensively in the glass industry.
A tetrahedral structural building unit.
Image courtesy of Babbage
Carbon Allotropes
Carbon's ability to form an array of allotropes stems from its electron configuration, which allows it to form strong covalent bonds with itself. Here are some of its notable allotropes:
Diamond
- Each carbon atom in diamond is tetrahedrally bonded to four other carbon atoms, creating a three-dimensional giant covalent structure.
- Properties include:
- High Hardness: Diamond is one of the hardest known natural materials.
- Transparency: Diamond is transparent as it lacks free electrons or impurities.
- Non-conductivity: Diamond does not conduct electricity.
Graphite
- Carbon atoms in graphite are bonded to form flat planes of hexagonal rings.
- Properties include:
- Conductivity: Graphite can conduct electricity due to delocalised electrons between the layers.
- Slippery: The layers in graphite can slide over each other, making it a good lubricant.
Fullerene (C₆₀)
- Structured in a series of hexagonal and pentagonal rings.
- Can be viewed as a hollow sphere, ellipsoid, or tube.
- Used in a variety of applications including drug delivery systems in the medical field.
Graphene
- A single layer of carbon atoms arranged in a two-dimensional honeycomb lattice.
- Properties include:
- High electrical conductivity: Making it suitable for use in electronics.
- Strong: It's incredibly strong due to the nature of its covalent bonds.
(a) Diamond, (b) amorphous (polycrystal of graphite and diamond), (c) fullerene, (d) single-walled carbon nanotube and (e) graphite.
Image courtesy of Jozef Sivek
Comparison of Properties
The unique properties of each allotrope and silicon structure are primarily due to their distinct bonding and structural patterns.
- Melting and Boiling Points: Diamond, being a giant covalent structure, has a significantly high melting and boiling point. Similarly, silicon and silicon dioxide also exhibit high melting and boiling points.
- Electrical Conductivity: While diamond and silicon are insulators, graphite and graphene are excellent conductors. This difference arises due to the delocalised electrons present in graphite and graphene, which are absent in diamond and silicon.
- Hardness: Diamond's tetrahedral arrangement of carbon atoms makes it extremely hard. In contrast, graphite's layered structure provides it with its soft and slippery nature.
Diamond and graphite with their respective structures.
Image courtesy of Itub
Strength Disparity Between Si-Si and C-C Bonds
The strength of a covalent bond depends on several factors, including bond length and the effective nuclear charge experienced by the bonded atoms. Carbon atoms are smaller than silicon atoms. As a result:
- Carbon-carbon bonds (as in diamond or graphite) are shorter and stronger than silicon-silicon bonds.
- The bond energy of a C-C single bond is about 348 kJ/mol, while that of a Si-Si single bond is around 226 kJ/mol.
- The shorter bond length in carbon allotropes, combined with the greater effective nuclear charge of carbon compared to silicon, contributes to the strength disparity.
In conclusion, while carbon and silicon both belong to Group 14 of the periodic table and exhibit similar bonding patterns, their physical properties differ significantly due to the differences in their atomic structures, bond lengths, and bond strengths. Understanding these differences and the reasons behind them is essential for both academic and practical applications in chemistry.
FAQ
Though silicon and carbon belong to the same group in the periodic table, their behaviours are different due to their positions in the table. Silicon is a metalloid, lying just below carbon. In nature, silicon tends to react with oxygen to form silicon dioxide (SiO2). This is because silicon's outer electrons are farther from the nucleus compared to carbon's, making them more easily accessible for reactions. As a result, silicon readily reacts with oxygen in Earth's atmosphere, forming compounds like SiO2 and preventing it from existing freely. Hence, most silicon is found combined with oxygen in minerals like quartz.
The tetrahedral arrangement in diamond is a result of each carbon atom forming four sigma (σ) covalent bonds with four adjacent carbon atoms. This spatial arrangement optimally reduces repulsion between the bonding pairs of electrons, providing a stable electron configuration. The covalent bonds in diamond are exceptionally strong due to the small size of carbon atoms, allowing the electrons to be held tightly between the nuclei. This combination of the tetrahedral arrangement and strong covalent bonds gives diamond its incredible rigidity and hardness, making it one of the hardest naturally occurring materials on Earth.
The versatility of silicon dioxide (SiO2) comes from its unique chemical and physical properties. Structurally, SiO2 forms a three-dimensional network of silicon atoms bonded covalently to oxygen atoms. This gives it a high melting and boiling point, and makes it hard and chemically inert. In its crystalline form, it is used as quartz in watches and in computer chips due to its piezoelectric properties. Its amorphous form, commonly found as sand, is a key ingredient in making glass and ceramics. Additionally, its insulating properties make it valuable in the semiconductor industry, where it is used to isolate components in integrated circuits.
Graphite's lubricating property arises from its unique crystal structure. Unlike diamond which has a three-dimensional tetrahedral arrangement, graphite is composed of flat layers of hexagonal carbon rings stacked on top of one another. These layers are held together by weak van der Waals forces. This means that when a force is applied, these layers can easily slide over each other, allowing graphite to act as a lubricant. These weak interlayer attractions, contrasting with the strong covalent bonds within each layer, grant graphite its slippery nature, making it useful in applications like pencil leads and industrial lubricants.
Fullerenes, often referred to as "buckyballs", have a distinct molecular structure. They are made up of carbon atoms arranged in a spherical shape, resembling a football. The most common fullerene, C60, consists of 60 carbon atoms forming a combination of hexagonal and pentagonal rings. This contrasts with the tetrahedral arrangement in diamond and the flat hexagonal planes in graphite. Fullerenes are molecular in nature, meaning they exist as discrete molecules, unlike diamond and graphite which are giant covalent structures. The unique structure of fullerenes leads to properties distinct from other carbon allotropes, including potential applications in medicine and electronics.
Practice Questions
Diamond consists of a three-dimensional giant covalent structure where each carbon atom is tetrahedrally bonded to four other carbon atoms. This gives diamond its extreme hardness. There are no free electrons, which makes diamond an electrical insulator. On the other hand, graphite has flat planes of hexagonal rings of carbon atoms, with weak forces between these layers, allowing them to slide over one another, accounting for graphite's slippery nature. Within each layer, there are delocalised electrons, enabling graphite to conduct electricity, unlike diamond.
The strength of a covalent bond is influenced by bond length and the effective nuclear charge experienced by the bonded atoms. Carbon atoms are smaller than silicon atoms, leading to shorter carbon-carbon bond lengths compared to silicon-silicon bond lengths. Shorter bond lengths result in stronger covalent bonds. Additionally, carbon, being higher up on the periodic table, has a greater effective nuclear charge than silicon, which means its nucleus exerts a stronger pull on bonding electrons. Consequently, the C-C bond is stronger than the Si-Si bond due to its shorter bond length and the greater effective nuclear charge of carbon compared to silicon.
