Among \(\mathrm{NH}_{3}, \mathrm{H}_{2} \mathrm{O}, \mathrm{HF}\) and \(\mathrm{H}_{2} \mathrm{~S}\) which would have highest magnitude of hydrogen bonding? (a) HF due to maximum pobrity. (b) \(\mathrm{H}_{2} \mathrm{O}\) due to lone pairs of electrons. (c) \(\mathrm{NH}_{3}\) due to small size of nitrogen. (d) \(\mathrm{H}_{2} \mathrm{~S}\) due to higher electron affinity of sulphur

Electronegativity is a chemical property that describes how strongly atoms attract electrons in a chemical bond. It's crucial for understanding the behavior of atoms in different molecules. The higher the electronegativity, the greater the atom's ability to attract electrons to itself. In the context of hydrogen bonding, it plays a key role because hydrogen bonds usually form between hydrogen and an atom with high electronegativity, such as nitrogen (N), oxygen (O), or fluorine (F).

Why is this important for hydrogen bonding? When an atom like oxygen, with a high electronegativity, is bonded to hydrogen, the uneven distribution of electrons creates a partial negative charge (delta-) on the oxygen and a partial positive charge (delta+) on the hydrogen. This polarity is a prerequisite for hydrogen bonding and affects how molecules will interact with each other. For example, in water (H₂O), oxygen's high electronegativity allows it to pull the bonding electrons closer, creating a significant dipole that leads to strong hydrogen bonds with neighboring water molecules.

Dipole-dipole interactions occur between polar molecules, which possess regions of partial positive and negative charge, known as dipoles. These interactions are another type of intermolecular force that plays a critical part in the behavior of molecules. Molecules with greater dipoles will have stronger dipole-dipole interactions, leading to higher boiling points and melting points compared to similar, nonpolar molecules.

In regards to hydrogen bonding, which is a specific form of dipole-dipole attraction, the partially positive hydrogen atom of one molecule is attracted to the partial negative charge of an electronegative atom in a nearby molecule. This attraction contributes to the unique properties of substances like water, where each molecule can form multiple hydrogen bonds, resulting in a high boiling point and surface tension. Remembering that the polarity of a molecule enhances its ability to participate in dipole-dipole interactions and hydrogen bonding will help you anticipate the physical properties of substances.

Lone pairs of electrons refer to the pair of valence electrons that are not shared with another atom in a chemical bond. These lone pairs are particularly important in the formation of hydrogen bonds. Molecules with lone pairs on highly electronegative atoms like oxygen or nitrogen can act as hydrogen bond acceptors. This means they can interact with the partially positive hydrogen on a polar molecule to form a hydrogen bond.

For instance, in water (H₂O), the oxygen atom has two lone pairs of electrons which can attract the hydrogen atoms of adjacent water molecules, leading to the formation of a strong network of hydrogen bonds. This explains why water has a relatively high boiling point and unique properties such as its solid form (ice) being less dense than its liquid form. Oxygen's ability to form two hydrogen bonds per molecule because of its lone pairs results in a three-dimensional lattice when frozen, making ice less dense than the liquid. As illustrated in the exercise, the presence of lone pairs on highly electronegative atoms is a key factor in determining the strength of hydrogen bonding in a molecule.