Classification of Solids – Crystalline and Amorphous Solids

Solids may be split into two categories: crystalline and amorphous. The first and most frequent kind, known as crystalline, contains regular crystal lattices, or long-range order, whereas the second and less common variety, known as amorphous. It is the regular arrangement of their atoms that allows these solids to maintain their stability. Their distinctive melting and boiling temperatures, regular geometric forms, and flat faces when split or sheared are only a few of their distinguishing characteristics. Sodium chloride, ice, metals, and diamonds are examples of substances that fall within this category.

As opposed to hard solids, amorphous solids are characterised by the absence of repeating periodicity or long-range order in their structure. Neither the solid nor liquid phases of these compounds are clearly distinguishable from one another. Amorphous solids lack a recognisable geometry, exhibit equal characteristics along all axes, melt across a broad range of temperatures, and shatter to produce curved or irregular forms when they are broken apart.

It is vital to remember that these two phrases represent two extremes on a spectrum of meaning. The majority of amorphous solids exhibit some degree of short-range order. Even with the use of modern analytical techniques such as x-ray diffraction and transmission electron microscopy, it is difficult to discern between the two distinct kinds on an atomic size.

Molecular Solids

Molecular crystalline solids, such as ice, sucrose (table sugar), and iodine, are solids that are made up of neutral molecules as their component units and have a crystalline structure. London dispersion forces, dipole-dipole interactions, and hydrogen bonds are among the weak intermolecular forces that hold these molecules together. These forces determine the characteristics of the molecules.

As a result of these differences, the intensity of the attraction forces between the units present in various crystals varies greatly, as seen by the differences in melting temperatures of different crystals.

H₂, N₂, O₂, and F₂ are small symmetrical nonpolar molecules that have negligible dispersion forces and produce molecular solids with extremely low melting temperatures (below 200 degrees Celsius). Substances made up of bigger, nonpolar molecules have stronger attractive forces and melt at greater temperatures than other substances.

Molecular solids, which are made up of polar molecules with permanent dipole moments, melt at even greater temperatures than liquid crystals. Solid SO₂ and table sugar are examples of such substances. As observed in frozen water or ice, intermolecular hydrogen bonding is primarily responsible for the preservation of the three-dimensional lattice structure of such molecular solids as they form.

Ionic Solids 

Sodium chloride, for example, is an ionic crystalline solid that is kept together by strong electrostatic interaction between its positive and negative ions. Because of the strong ionic interaction, ionic solids have extremely high melting points. An approximate measure of the strength of ionic interaction between cations and anions in an ionic solid is the electrostatic force, which may be approximated by Coulomb’s equation as follows:

This equation contains the constant of proportionality, the distance between charges (represented by r), and charge symbols for anions (represented by qa) and cations (represented by qc). The greater the charge on the cations and anions, the greater the strength of the ionic attraction force between them. Additionally, dense packing of anions and cations in the crystal lattice lowers the distance between charges, resulting in higher forces of ionic attraction between the ions and the charges.

Ionic solids are hard, but they are also brittle, and they shatter rather than bend when they are broken. The existence of both attractive (cation-anion) and repulsive (cation–cation and anion–anion) interactions in the crystal lattice is linked to the brittleness of the crystals in question. Ionic solids cannot conduct electricity because the ions are unable to move freely due to the intense coulombic forces acting on them. The ions, on the other hand, become free to move around and conduct electricity when the metal is molten or when it is dissolved in water.

Metallic Solids

Atoms of a metal can coalesce into a variety of metallic solids, including crystals of copper, aluminium, and iron. A “sea” of delocalized electrons is used to characterise the structure of metallic crystals, which is commonly defined as having a uniform distribution of atomic nuclei inside the structure. The atoms that make up a metallic solid are kept in place by a special force known as metallic bonding, which is responsible for the formation of a wide variety of beneficial and interesting bulk characteristics.

The capacity to be shaped easily, a metallic sheen, strong thermal and electrical conductivity, and malleability are characteristics shared by all solid metallic substances. Many are extremely tough and have impressive levels of strength. They do not shatter and are consequently suitable as construction materials because of their malleability, which is the ability to flex under pressure or hammering. The melting points of the different metals are very different from one another. The alkali metals all melt at temperatures lower than 200 degrees Celsius, whereas mercury is a liquid at ambient temperature. In contrast to the transition metals, which do not melt until temperatures higher than 1000 degrees Celsius are reached, certain post-transition metals have a low melting point. These variances are due to the fact that different metals have varying degrees of strength in their metallic bonding.

Covalent Solids

Covalent solids, also known as network solids, are solids that are bound together by covalent connections. Covalent solids may be found in a variety of forms. So they have localised electrons (which are shared among the atoms) and their atoms are organised in defined geometries. Only the breakage of covalent sigma bonds can result in a distortion away from this geometrical configuration. Consequently, the melting point of covalent solids is extraordinarily high. They are often highly hard material that will break into fragments rather than changing shape in a smooth manner. They are rigid and brittle, as the expression goes.

Conclusion

Metallic crystals are composed of metal cations that are surrounded by a “sea” of freely moving valence electrons . They are also known as delocalized electrons since they do not belong to any one atom and are capable of travelling throughout the entire crystal. Therefore, metals are excellent conductors of electrical current. Molecular crystals are generally composed of molecules arranged at the crystal’s lattice points and kept together by weak intermolecular forces, as opposed to solid crystals (see figure below). In the case of nonpolar crystals, the intermolecular interactions may be dispersion forces, whereas in the case of polar crystals, the intermolecular forces may be dipole-dipole forces. Some molecular crystals, such as ice, are kept together by hydrogen bonds, which are formed between molecules. A covalent network crystal is made up of atoms that are located at the crystal’s lattice points, with each atom being covalently connected to the atoms that are located next to it (see figure below). The three-dimensional covalently connected network has a huge number of atoms and is composed of covalently bound atoms. Diamond, quartz, and a variety of metalloids, as well as transition metal and metalloid oxides, are examples of covalent solids.