krit.club logo

Chemical Bonding - Giant Covalent Structures (Diamond, Graphite, SiO2)

Grade 9IB

Review the key concepts, formulae, and examples before starting your quiz.

🔑Concepts

•

A giant covalent structure (or macromolecule) consists of a vast number of atoms held together by a continuous network of strong covalent bonds in a regular lattice.

•

In Diamond, each Carbon (CC) atom is covalently bonded to four other carbon atoms in a rigid tetrahedral arrangement. This makes diamond extremely hard and gives it a very high melting point.

•

In Graphite, each Carbon (CC) atom is bonded to three others in flat hexagonal layers. The fourth valence electron is delocalized and free to move between layers, allowing graphite to conduct electricity.

•

The layers in Graphite are held together by weak intermolecular forces (Van der Waals forces), which allows them to slide over each other, making graphite soft and slippery.

•

Silicon Dioxide (SiO2SiO_{2}), also known as silica or quartz, has a structure similar to diamond. Each Silicon (SiSi) atom is bonded to four Oxygen (OO) atoms, and each Oxygen atom is bonded to two Silicon atoms.

•

Giant covalent structures have very high melting and boiling points because a massive amount of energy is required to break the strong covalent bonds throughout the entire lattice.

•

These structures are generally insoluble in water and organic solvents because the attraction between the solvent molecules and the atoms is not strong enough to break the covalent bonds.

📐Formulae

C(s) (Diamond or Graphite)C_{(s)} \text{ (Diamond or Graphite)}

SiO2 (Silicon Dioxide)SiO_{2} \text{ (Silicon Dioxide)}

Bond Angle in Diamond=109.5∘\text{Bond Angle in Diamond} = 109.5^{\circ}

💡Examples

Problem 1:

Explain why Graphite is used to make electrodes for electrolysis, while Diamond is used in cutting tools.

Solution:

Graphite conducts electricity because it has delocalized electrons that can move through the structure. Diamond is used in cutting tools because its rigid tetrahedral lattice of strong covalent bonds makes it the hardest known natural substance.

Explanation:

In Graphite, only 33 out of 44 valence electrons per CC atom are used in bonding, leaving one free. In Diamond, all 44 electrons are locked in bonds, providing maximum structural integrity.

Problem 2:

Compare the melting points of SiO2SiO_{2} and CO2CO_{2}.

Solution:

SiO2SiO_{2} has a very high melting point (approx. 1710∘C1710^{\circ}C) because it is a giant covalent structure. CO2CO_{2} has a very low melting point (approx. −78∘C-78^{\circ}C) because it is a simple molecular substance.

Explanation:

Melting SiO2SiO_{2} requires breaking strong covalent bonds between SiSi and OO atoms. Melting CO2CO_{2} only requires overcoming weak intermolecular forces (London dispersion forces) between discrete O=C=OO=C=O molecules.

Problem 3:

Calculate the number of covalent bonds formed by 11 mole of Carbon atoms in a Diamond lattice.

Solution:

Since each Carbon atom forms 44 bonds, and each bond is shared between 22 atoms, the total number of bonds is: 4×NA2=2×NA\frac{4 \times N_{A}}{2} = 2 \times N_{A} where NA≈6.022×1023N_{A} \approx 6.022 \times 10^{23}.

Explanation:

In a giant lattice, each bond connects two atoms. Therefore, the total number of bonds is half the total number of 'bond-ends' created by the atoms.