Which of the following statement is not correct? (a) Decreasing order of bond angle of hydrides of chalcogens is \(\mathrm{H}_{2} \mathrm{O}>\mathrm{H}_{2} \mathrm{~S}>\mathrm{H}_{2} \mathrm{Se}>\mathrm{H}_{2} \mathrm{Te}\) (b) Decreasing order of stability of hexafluorides of \(S\), Se and Te is \(\mathrm{SF}_{6}>\mathrm{SeF}_{6}>\mathrm{TeF}_{6}\) (c) Decreasing order of tendency to undergo hydrolysis is \(\mathrm{TeF}_{6}>\mathrm{SeF}_{6}>\mathrm{SF}_{6}\) (d) All are correct

Understanding the bond angles in hydrides of chalcogens is foundational to predicting molecular shapes and properties. These bond angles are crucial as they influence the physical and chemical characteristics of water and other similar molecules.

When discussing hydrides such as water (H2O), hydrogen sulfide (H2S), hydrogen selenide (H2Se), and hydrogen telluride (H2Te), we observe a general trend where the bond angle decreases as we move from oxygen to tellurium (down the group in the periodic table). The reason behind this is attributed to the increase in atomic size leading to less effective repulsion between the lone pairs of electrons. Moreover, oxygen, being the smallest, has a higher electronegativity which results in a greater repulsion between the two lone pairs on the oxygen atom, thereby increasing the H-O-H bond angle compared to its counterparts.

This concept helps to explain many of the unique properties of water, such as its relatively high boiling point and its solvency powers. When discussing or solving problems related to these angles, one should remember that the bond angle is also affected by the hybridization of the central atom and the presence of lone pairs which can alter the expected linear geometry.

Hexafluorides, particularly of group 16 elements like sulfur, selenium, and tellurium, are an important class of compounds with fascinating chemical properties. Their stability is of particular interest in both inorganic chemistry and practical applications.

The stability of hexafluorides, such as SF6, SeF6, and TeF6, tends to decrease as we move down the group in the periodic table. Sulfur hexafluoride (SF6) is known for its exceptional stability which arises from the sulfur atom being small enough to form six strong bonds to the fluorine atoms, resulting in an octahedral geometry that is highly symmetrical and nonpolar.

This stability decreases with selenium and tellurium hexafluorides, as these larger atoms introduce longer bonds that are also more flexible. Consequently, they are less able to hold onto the fluorine atoms as tightly as sulfur does. This reduction in bond strength translates to a decrease in overall stability, making SF6 the most stable compared to SeF6 and TeF6. This stability is a significant factor in why SF6 is widely used as an insulating gas in high-voltage equipment.

The hydrolysis of hexafluorides involves reactions with water that can lead to decomposition and the formation of new products. This chemical behavior is especially relevant when considering the environmental impact and safety of these compounds.

Hexafluorides of sulfur, selenium, and tellurium demonstrate an increasing tendency to undergo hydrolysis as their atomic size increases from sulfur to tellurium. TeF6 is more susceptible to hydrolysis than SeF6, which in turn is more susceptible than SF6. The relatively stronger S-F bonds in SF6 provide remarkable resistance to hydrolysis, resulting in its inertness and why SF6 is popular for applications where chemical stability is essential, such as in electrical insulators.

In contrast, the weaker Te-F bonds in TeF6 result from the larger size of the tellurium atom, which makes its bonds more polarizable and less tightly held. This weaker bond strength underlies the greater reactivity of TeF6 with water. Understanding the hydrolysis of hexafluorides is vital for chemists, as it influences how these compounds are handled, stored, and disposed of to avoid degradation and release into the environment.