Oxygen is more electronegative than sulphur, yet \(\mathrm{H}_{2} \mathrm{~S}\) is acidic in nature while \(\mathrm{H}_{2} \mathrm{O}\) is neutral because (a) \(\mathrm{H}-\mathrm{S}\) is weaker than \(\mathrm{O}-\mathrm{H}\) bond (b) \(\mathrm{H}_{2} \mathrm{~S}\) is a gas at ordinary temperature while \(\mathrm{H}_{2} \mathrm{O}\) is a liquid (c) Water has higher boiling point than \(\mathrm{H}_{2} \mathrm{~S}\) (d) Water molecules are associated due to intermolecular hydrogen bonding

In the realm of chemistry, understanding the nature of acids and bases is fundamental. Acidic substances have the unique ability to donate protons (H+ ions). This proton donation process is an essential concept in the Bronsted-Lowry theory of acids and bases.

Considering hydrogen sulfide (H2S), its acidic nature arises from the hydrogen atoms' propensity to split off as protons, thereby classifying it as an acid. In contrast, water (H2O), which is neutral, does not readily release its hydrogen atoms as protons. The dissociation of H2S to release H+ is influenced by several factors, including bond strength and the stability of the resulting conjugate base.

Understanding the role of bond strength in the acidity of H2S can be simplified by examining the breaking of the H-S bond. Compared to the O-H bond in water, the H-S bond is weaker, thus facilitating the release of H+ ions more easily and affirmatively answering why H2S acts as an acid.

Electronegativity is a chemical property that describes the tendency of an atom to attract a shared pair of electrons towards itself in a chemical bond. This property is pivotal in influencing the acidity of molecules.

Typically, the more electronegative an atom, the stronger its pull on the electron pair, leading to a polarization of the bond. If an acid has a bond between hydrogen and a highly electronegative atom, upon ionization, the resultant conjugate base is stabilized by the electronegative atom's ability to spread the negative charge.

In the case of H2S, sulfur is less electronegative than oxygen; however, the H-S bond's polarity is sufficient to allow for proton donation, prompting its acidic character. It is not merely the absolute value of electronegativity that dictates acidity, but also how electronegativity influences bond strength and stability of the conjugate base. Thus, weaker bonds and stable conjugate bases favor acidity.

The study of chemical bonding provides great insight into the behaviors of molecules and compounds. Chemical bonds are the connections between atoms within molecules, and they play a crucial role in determining a substance's properties, including acidity.

Acidic strength can hinge on the type of bonds present in a molecule. For an acid like H2S, the H-S bonds are relatively weaker than the O-H bonds found in water. This difference in bond strength affects the ease with which the compounds can dissociate to release hydrogen ions. Weaker bonds, as seen in H2S, are more conducive to ionization in comparison to the stronger O-H bonds in H2O, resulting in water's neutrality.

The nature of the bonds within a molecule can also dictate other properties, such as melting and boiling points, which are influenced by the intermolecular forces holding the substance together. While such intermolecular forces do not directly determine acidity, they are important factors in the overall behavior of molecular substances.