Hydrogen bonding is a phenomenon characterized by the formation of hydrogen bonds, a distinct category of attractive intermolecular forces. These bonds result from the dipole-dipole interaction between a hydrogen atom, bonded to a highly electronegative atom, and another nearby highly electronegative atom. In the case of water molecules ([math]H_2O[/math]), hydrogen forms a covalent bond with the more electronegative oxygen atom. Consequently, hydrogen bonding arises in water molecules due to the dipole-dipole interactions between the hydrogen atom of one water molecule and the oxygen atom of another [math]H_2O[/math] molecule. In this article, we will delve into the concept of hydrogen bond, how this bond is formed as well as the examples that explain hydrogen bonding.

A hydrogen bond is characterized by the electrostatic attraction between a hydrogen atom and an atom possessing a lone pair of electrons within a substance. In order for a hydrogen bond to form, the hydrogen must be bonded to an electronegative atom.

It's essential to note that, despite the covalent nature of the molecules within the substance, the hydrogen bond itself is non-covalent.

When a hydrogen atom forms a covalent bond with an electronegative atom, the shared pair of electrons tends to be attracted to the more electronegative partner. Consequently, the hydrogen atom acquires a partial positive charge, while the electronegative atom assumes a partial negative charge. This development is recognized as polarity, and molecules exhibiting this property are termed polar molecules. In substances where two or more such polar molecules coexist, the positive end of one molecule is drawn towards the opposing end of another. This specific type of interaction is identified as a hydrogen bond.

The water molecule comprises two hydrogen (H) atoms covalently bonded to an oxygen (O) atom.

In the O-H bond, the bond pair of electrons is situated very close to the oxygen nucleus, due to the significant electronegativity difference between oxygen and hydrogen. Consequently, the oxygen atom acquires a partial negative charge (-δ), while the hydrogen atom develops a partial positive charge (+δ).

Hydrogen bonding arises as a result of the electrostatic attraction between the positively charged hydrogen atom of one water molecule (with +δ charge) and the negatively charged oxygen atom of another water molecule (with -δ charge). This phenomenon characterizes hydrogen bonds as a unique category of intermolecular attractive forces, specifically observed in compounds where hydrogen is bonded to a highly electronegative atom.

Comparatively, hydrogen bonds exhibit greater strength than typical dipole-dipole and dispersion forces. However, their strength is notably weaker when contrasted with true covalent or ionic bonds.

Ammonia is composed of three hydrogen (H) atoms bonded to a nitrogen (N) atom. Due to nitrogen's higher electronegativity compared to hydrogen, it attracts the shared pair of electrons, acquiring a slight negative charge. Conversely, the hydrogen atoms exhibit a slight positive charge, leading to the formation of hydrogen bonds within the molecule.

The capacity for hydrogen bonding in ammonia is limited by the fact that each nitrogen atom possesses only one lone pair in its outermost shell. Consequently, there are insufficient lone pairs of electrons in a cluster of ammonia molecules to fulfill all the hydrogen bonding possibilities. Each ammonia molecule can engage in one hydrogen bond using its lone pair and another involving one of its positively charged hydrogens. The remaining hydrogens do not participate in hydrogen bonding. This limited bonding capacity results in ammonia having a lower boiling point compared to water.

Fluorine has the highest value of electronegativity and forms the strongest hydrogen bond.