How does the Lewis model for covalent bonding account for the relatively low melting and boiling points of molecular compounds (compared to ionic compounds)?

Intermolecular forces play a vital role in determining the physical properties of a molecule, including its phase at room temperature and its melting and boiling points. Unlike ionic bonds that hold together ions in a crystal lattice, intermolecular forces are forces of attraction between separate molecules.

There are several types of intermolecular forces, with varying strengths.

• London dispersion forces, which are often the weakest, occur due to temporary dipoles in molecules.
• Dipole-dipole interactions occur when two polar molecules are attracted to each other due to the positive end of one attracting the negative end of the other.
• Hydrogen bonds, the strongest of the three, occur when hydrogen is bonded to a highly electronegative element such as oxygen, nitrogen, or fluorine, allowing for a strong, directional interaction with another electron-rich region.

These forces are crucial as they dictate how molecules interact with each other, leading to varying physical states and thermal properties of molecular compounds.

Understanding the difference between ionic and covalent bonds is fundamental in chemistry. Ionic bonds form when electrons are transferred from a metal to a nonmetal, resulting in positive and negative ions that attract each other in a crystal lattice. Covalent bonds, on the other hand, occur when nonmetal atoms share electron pairs to achieve a more stable electron arrangement.

Ionic compounds, due to the strong electrostatic forces between ions, typically exhibit high melting and boiling points. Conversely, molecules with covalent bonds are held together by shared electrons resulting in a lower energy bond. This is key when considering their thermal properties; covalent compounds generally have lower melting and boiling points than ionic compounds as less energy is required to break the intermolecular forces compared to the ionic bonds.

The melting and boiling points of molecular compounds are significantly influenced by the type and strength of intermolecular forces present. Since these forces are weaker compared to the ionic bonds in salts, molecular compounds typically have lower melting and boiling points.

Simple molecular compounds with only London dispersion forces, for example, have very low melting and boiling points. Substances with stronger intermolecular forces, such as hydrogen bonding, will have higher melting and boiling points but still lower than most ionic compounds. The energy required to overcome these forces during phase changes is relatively small, explaining why many molecular compounds are gases or liquids at room temperature.

Electronegativity is a measure of an atom's ability to attract and hold onto electrons within a bond. When there's a significant difference in electronegativity between two atoms sharing electrons, a polar covalent bond is formed. The more electronegative atom pulls the electron cloud closer, creating a partial negative charge, while the other atom takes on a partial positive charge.

Polarity is, therefore, a result of these uneven electron distributions in a molecule, influencing the molecule's physical properties, such as solubility and melting and boiling points. Even though polar molecules may have slightly higher melting and boiling points than nonpolar ones due to dipole-dipole interactions, these effects are minimal when compared to ionic compounds, which are held together by the strong attraction between oppositely charged ions.