A Brief Note on Polar Molecular Solids

The majority of molecular solids are insulators, are generally soft, and have low densities, among other characteristics. Sugar, solid halogens, sulphur, and ice are examples of molecular solids, as are other organic compounds (made of water). Molecular solids are further subdivided into a number of various categories.

Non-polar Molecular Solids 

Because such solids contain a symmetrical distribution of electrons, there is no excess charge on any side of the solid. Charges that are diametrically opposed cancel each other out. Methane, chlorine, hydrogen, and oxygen are just a few examples. They are either liquids or gases at standard room temperature and pressure. Vander Waals forces are the forces that hold the molecules in these solids together by holding them together (weak dispersion or London forces). In comparison to ionic or covalent bonds, these forces are weaker.

Polar Molecular Solids

The geometry of these solids is such that one side has a negative charge and the other side has a positive charge on the same side. The dipole — dipole force of attraction – that holds them together is the driving force. Their melting and boiling temperatures are higher than those of non-polar molecular solids, but they are still comparatively low compared to other solids. Polar compounds include ethanol and ammonia, to name a couple of examples.

Hydrogen-Bonded Molecular Solids

The intermolecular forces in these types of materials are dominated by strong hydrogen bonding. When compared to polar and non-polar molecular solids, their boiling and melting temperatures are significantly higher. At room temperature and pressure, they exist as volatile liquids or soft solids, respectively. Water is an example of a molecular liquid that is hydrogen-bonded.

1.Covalent solids

What are covalent solids, and how do they work? It is covalent bonds that keep the individual atoms of molecules connected in these substances. The construction of a large molecule is facilitated by the formation of a network of interconnecting covalent bonds throughout the crystal. They have high melting and boiling points and, with the exception of graphite, are generally poor conductors, unless they are melted together (due to the fourth electron being free).

2. Metallic Solids 

What are metallic solids and how do they differ from one another? The constituent particles in metallic solids are metal atoms, which have valence electrons that can be given or lost, resulting in the metal atoms becoming positively charged as a result of this process. It is possible to move around in the sea of electrons available since they are dispersed throughout the crystal. The creation of metallic bonds is caused by an attractive attraction between positively charged ions and a sea of electrons, which attracts them together. This is the gravitational force that holds the metal ions together. Metallic solids have a regular structure and have high melting and boiling temperatures, which distinguish them from other solids. Because of the large number of electrons in the system, they have excellent thermal and metallic conductivity. All metals and alloys are metallic solids, which means they have no liquid state.

3. Ionic Solids

Anions (negatively charged particles) and cations are the constituent particles in ionic solids of the Crystalline Solids family (positively charged). An ion is surrounded by a usual amount of charges that are in opposition to it. For example, the Na+ ion is surrounded by 6 Cl- ions in the compound NaCl. Strong electrostatic forces hold the ions in these substances together, preventing them from separating. It is important to note that they have high melting and boiling temperatures, and that they are soluble in polar but not non-polar solvents.

CONCLUSION

A polar molecule is defined in chemistry as a molecule that contains a charge on one side of the molecule that is not cancelled off by the other side of the molecule. It has a partial charge in one of its regions. There are two ends that are somewhat positive and one end that is slightly negative. They are often asymmetrical, with an unequal distribution of electrons across the structure.