Categories of Solids

Categories of Solids Based on the Solid Pack

Solids can be divided into three categories on the basis of how the particles that form the solid pack.

Crystalline solids are three-dimensional analogs of a brick wall. They have a regular structure, in which the particles pack in a repeating pattern from one edge of the solid to the other.

Amorphous solids (literally, "solids without form") have a random structure, with little if any long-range order.

Polycrystalline solids are an aggregate of a large number of small crystals or grains in which the structure is regular, but the crystals or grains are arranged in a random fashion.
The extent to which a solid is crystalline has important effects on its physical properties.

Examples: The polyethylene used to make sandwich bags and garbage packs is an amorphous solid that consists of more or less randomly oriented chains of (-CH2-CH2-) linkages. Milk bottles are made from a more crystalline form of polyethylene, and they have a much more rigid structure.

Categories of Solids Based on Bonds that Hold the Solid Together


Solids can be classified on the basis of the bonds that hold the atoms or molecules together. This approach categorizes solids as either molecular, covalent, ionic, or metallic.

Iodine (I2), sugar (C12H22O11), and polyethylene are examples of compounds that are molecular solids at room temperature. Water and bromine are liquids that form molecular solids when cooled slightly; H2O freezes at 0oC and Br2 freezes at -7oC.

Molecular solids are characterized by relatively strong intramolecular bonds between the atoms that form the molecules and much weaker intermolecular bonds between these molecules. Because the intermolecular bonds are relatively weak, molecular solids are often soft substances with low melting points.

Dry ice, or solid carbon dioxide, is a perfect example of a molecular solid. The van der Waals forces holding the CO2 molecules together are weak enough that dry ice sublimes--it passes directly from the solid to the gas phase--at -78oC.



Covalent solids, such as diamond, form crystals that can be viewed as a single giant molecule made up of an almost endless number of covalent bonds. Each carbon atom in diamond is covalently bound to four other carbon atoms oriented toward the corners of a tetrahedron, as shown in the figure below. Because all of the bonds in this structure are equally strong, covalent solids are often very hard and they are notoriously difficult to melt. Diamond is the hardest natural substance and it melts at 3550C.

Ionic solids are salts, such as NaCl, that are held together by the strong force of attraction between ions of opposite charge.


Because this force of attraction depends on the square of the distance between the positive and negative charges, the strength of an ionic bond depends on the radii of the ions that form the solid. As these ions become larger, the bond becomes weaker. But the ionic bond is still strong enough to ensure that salts have relatively high melting points and boiling points.

To understand metallic solids we have to clear up a common misconception about chemical bonds. Ionic and covalent bonds are often imagined as if they were opposite ends of a two-dimensional model of bonding in which compounds that contain polar bonds fall somewhere between these extremes.

ionic ........ polar ........ covalent

In reality, there are three kinds of bonds between adjacent atoms: ionic, covalent, and metallic, as shown in the figure below. The force of attraction between atoms in metals, such as copper and aluminum, or alloys, such as brass and bronze, are metallic bonds.

Molecular, ionic, and covalent solids all have one thing in common. With only rare exceptions, the electrons in these solids are localized. They either reside on one of the atoms or ions or they are shared by a pair of atoms or a small group of atoms.

Metal atoms don't have enough electrons to fill their valence shells by sharing electrons with their immediate neighbors. Electrons in the valence shell are therefore shared by many atoms, instead of just two. In effect, the valence electrons are delocalized over many metal atoms. Because these electrons aren't tightly bound to individual atoms, they are free to migrate through the metal. As a result, metals are good conductors of electricity. Electrons that enter the metal at one edge can displace other electrons to give rise to a net flow of electrons through the metal.


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