Question

Graphite is an example of : (a) Ionic solid (b) Covalent solid (c) vander Waals' crystal (d) Metallic crystal

Short answer

Graphite is an example of a covalent solid.

Step-by-step solution

Understanding Types of Crystalline Solids

Crystalline solids are categorized based on the types of chemical bonds that hold their constituent particles together. Ionic solids are held together by ionic bonds, covalent solids by covalent bonds, van der Waals' crystals by van der Waals forces, and metallic crystals by metallic bonds.

Recognizing the Chemical Structure of Graphite

Graphite is composed of layers of carbon atoms with each layer resembling a hexagonal lattice. The carbon atoms within each layer are bonded covalently, forming a planar network.

Determining the Appropriate Category for Graphite

Since graphite has a structure where carbon atoms are covalently bonded to one another, it does not fit the criteria for ionic, van der Waals', or metallic crystals. It is, therefore, an example of a covalent solid.

Key concepts

Ionic Solid

Ionic solids are a class of crystalline materials characterized by a lattice structure formed by ions held together by ionic bonds, which are electrostatic attractions between positively charged cations and negatively charged anions. The classic table salt, or sodium chloride (NaCl), is a prime example of an ionic solid.

Key features of ionic solids include high melting and boiling points, due to the strong coulombic forces between ions. These solids are typically hard and brittle, and when dissolved in water, they often form electrolyte solutions capable of conducting electricity, which is significant for various applications ranging from batteries to biological systems.

Understanding ionic solids involves exploring lattice energy, the energy released when the ions come together to form the solid structure, and this energy is key to the stability of the crystal. It's also interesting to note that, despite their rigid structure, ionic solids can exhibit some plasticity when external pressure is applied, leading to ionic conduction.

Covalent Solid

Covalent solids, also known as network solids, are formed by networks or chains of atoms linked together by covalent bonds. These bonds happen when atoms share electrons, leading to a highly stable and strong bond that gives covalent solids their unique properties.

One of the well-known covalent solids is diamond, where each carbon atom is tetrahedrally bonded to four other carbon atoms, creating an incredibly hard material. Other than their hardness, covalent solids are noted for their high melting points and poor electrical conductivity, as the shared electrons are typically localized and not free to move about. However, some covalent solids like graphite can conduct electricity due to their layered structure, where one layer of atoms is weakly bonded to the next, allowing electrons to flow.

As a three-dimensional network of covalently bonded atoms, these materials do not dissolve in water and are generally unreactive due to the strength of covalent bonds. While students often identify covalent bonds with molecules, it's important to differentiate covalent solids, where the bonding extends throughout the material.

Van der Waals' Crystal

Van der Waals' crystals are named after the weak intermolecular forces that hold the structural components of the material together. These forces include London dispersion forces, dipole-dipole interactions, and hydrogen bonds. Due to the delicate nature of these interactions, van der Waals' crystals such as molecular solids or noble gases in their solid form exhibit low melting and boiling points.

Examples of van der Waals' crystals include solid forms of CO₂ or dry ice, iodine crystals, and fullerenes. These materials are generally soft and easily deformed, as the van der Waals' forces provide minimal resistance to external stresses. Additionally, they're often good insulators, since the electrons are situated within discrete molecules and not free to move about as in metals or semiconductors.

In explaining van der Waals' crystals, it is crucial to emphasize the variations in strength among the different van der Waals' forces and how they influence the properties of materials. For instance, hydrogen bonds can be relatively strong compared to other van der Waals' forces and lend higher melting points to materials like ice.

Metallic Crystal

Metallic crystals encompass an array of solid materials with atoms arranged in a high-symmetry crystal lattice, bound together by metallic bonds. This particular bond is characterized by a sea of delocalized electrons, which are not bound to any specific atom and can move freely throughout the crystal. This 'electron sea' model helps explain many of the hallmark properties of metals, such as their luster, good electrical and thermal conductivity, and malleability.

When we consider common metals like copper, iron, or gold, it's important to recognize that these metallic crystals can be malleable and ductile, meaning they can be hammered into shapes or stretched into wires without breaking. This plastic deformation is due to the ability of atoms to roll over each other while maintaining the metallic bond.

Metallic crystals also exhibit a range of melting points and can have magnetic properties, as seen in iron or cobalt. Through the study of metallic bonding and crystal structures, students gain insights into the design of alloys, which are mixtures of metals that combine the advantageous properties of their constituent elements.