Question

What kinds of attractive forces exist between particles (atoms, molecules, or ions) in (a) molecular crystals, (b) covalentnetwork crystals, (c) ionic crystals, (d) and metallic crystals?

Short answer

In (a) molecular crystals, attractive forces are intermolecular forces like hydrogen bonding, dipole-dipole interactions, and London dispersion forces. In (b) covalent network crystals, the forces are covalent bonds within continuous networks of atoms. In (c) ionic crystals, attractive forces are ionic bonds between oppositely charged ions. In (d) metallic crystals, metallic bonding occurs between metal ions and a "sea" of delocalized electrons.

Step-by-step solution

Molecular Crystals

\ In molecular crystals, the attractive forces are intermolecular forces. These include forces like hydrogen bonding, dipole-dipole interactions, and London dispersion forces (or van der Waals forces). These forces result from the electrostatic attraction between molecules, which arise due to the differences in electronegativity between the constituent atoms.

Covalent Network Crystals

\ Covalent network crystals, also known as network solids, are characterized by a continuous network of covalently bonded atoms throughout the crystal. The attractive forces within these crystals are covalent bonds. These bonds arise due to the sharing of electrons between the atoms. Some common examples of covalent network crystals include diamond, silicon dioxide (quartz) and silicon carbide.

Ionic Crystals

\ In ionic crystals, the attractive forces are ionic bonds between positively and negatively charged ions. These bonds are formed due to the electrostatic attraction between the oppositely charged ions. Ionic crystals often have high melting and boiling points due to the strong attractive forces holding the ions together. Examples of ionic crystals include sodium chloride (table salt), potassium iodide, and magnesium oxide.

Metallic Crystals

\ Metallic crystals (or metals) have unique bonding arrangements, known as metallic bonding. This involves a lattice of positively charged metal ions surrounded by a "sea" of free-moving or delocalized electrons. The attractive forces between the metal ions and the delocalized electrons are what hold the metallic crystal together. These bonds give metals their characteristic properties, such as electrical conductivity and malleability. Examples of metallic crystals include gold, copper, and iron.

Key concepts

Intermolecular Forces

Intermolecular forces are essential in holding particles together in molecular crystals. These forces include a variety of interactions like hydrogen bonding, dipole-dipole interactions, and London dispersion forces, often referred to as van der Waals forces. Although these forces are generally weaker compared to covalent or ionic bonds, they play a crucial role in determining the physical properties of substances.

• Hydrogen bonding: This type of intermolecular force occurs when hydrogen is bonded to highly electronegative atoms like oxygen, nitrogen, or fluorine. The result is a strong attraction between the hydrogen atom of one molecule and the electronegative atom of a neighboring molecule.
• Dipole-Dipole Interactions: Molecules with permanent dipoles experience dipole-dipole interactions, with the positive end of one molecule attracted to the negative end of another, creating an electrostatic attraction.
• London Dispersion Forces: These are the weakest intermolecular forces and arise due to temporary dipoles caused by fluctuations in electron distribution. They are present in all molecular crystals, even those with only nonpolar molecules.

Understanding intermolecular forces helps explain why some substances exist as gases, liquids, or solids under different conditions.

Covalent Bonds

Covalent bonds are the strong forces holding atoms together in covalent network crystals, also known as network solids. These bonds are formed when atoms share electrons to achieve stability, creating a continuous three-dimensional network. Unlike intermolecular forces, covalent bonds are strong and directional, which means they give these crystals unique properties.
Diamonds, silicon dioxide, and silicon carbide are classic examples of covalent network crystals. In diamonds, each carbon atom shares electrons with four other carbon atoms, forming a rigid lattice that gives diamond its famous hardness. In silicon dioxide, each silicon atom bonds with oxygen atoms in a tetrahedral structure that makes quartz durable.

• Properties of Covalent Network Crystals: These materials typically have high melting and boiling points because the strong covalent bonds need significant energy to break.
• Structure: The repeated, endless pattern of covalent bonds ensures the crystal's rigid and durable form.

This bonding explains why covalent network crystals are often used in abrasive applications and as semiconductors in electronics.

Ionic Bonds

Ionic bonds are the key interactions in ionic crystals. These bonds occur when electrons are transferred from one atom to another, resulting in a pair of oppositely charged ions. The electrostatic forces between these positive and negative ions generate the strong attraction needed to form ionic structures.
Common examples of ionic crystals include sodium chloride (table salt), potassium iodide, and magnesium oxide. Their crystal lattice structures are characterized by the regular arrangement of ions, where each positive ion is surrounded by negative ions and vice versa.

• Properties of Ionic Crystals: They generally have high melting and boiling points due to the strong ionic bonds. These substances often dissolve in water, and the resulting solution conducts electricity.
• Structure: Ionic crystals form repeating patterns in which the electrostatic attractions balance the repulsions, creating a stable crystal lattice.

The stability and strength of ionic bonds make these crystals essential in a range of applications, from culinary uses to industrial ones.

Metallic Bonding

Metallic bonding is unique to metals and gives rise to the properties observed in metallic crystals. In this type of bonding, positively charged metal ions are positioned in a lattice, while delocalized electrons move freely around them. This 'sea of electrons' creates a strong attractive force that holds the lattice together and allows metals to conduct electricity efficiently.
Examples of metals that form metallic crystals include gold, copper, and iron, each showcasing metallic bonding characteristics like malleability and ductility. These properties arise from the ability of atoms to slide over one another without breaking the bond, facilitated by the delocalized nature of the electrons.

• Electrical Conductivity: The free-moving electrons provide metals with excellent electrical conductivity, as they can easily flow in response to an electric field.
• Malleability and Ductility: Metals can be reshaped and drawn into wires due to the layers of atoms sliding over one another without disrupting the metallic bond.

This bonding model provides insights into why metals are crucial in construction, wiring, and various technological applications.