Animated Solution for Chemistry - s and p-Block Elements: The correct order of the thermal stability of hydrogen halides (H—X) is

When we talk about the thermal stability of a molecule, we are essentially asking a very simple question: How much heat can this molecule withstand before it breaks apart?

For diatomic molecules like the hydrogen halides ($H-X$), thermal stability is directly proportional to the bond dissociation energy. If the bond holding the hydrogen and halogen atoms together is incredibly strong, it will take a massive amount of thermal energy to break it. Conversely, a weak bond will snap at much lower temperatures.

To understand the strength of the $H-X$ bond, we must look at the periodic table. As we descend Group 17 (the halogens) from Fluorine ($F$) to Iodine ($I$), the principal quantum number increases. This means new electron shells are being added, and the atomic radius of the halogen atom expands significantly.

Bond Strength

Imagine the hydrogen atom as a tiny sphere trying to hold hands with the halogen atom. When the halogen is small, like Fluorine, the two nuclei can get very close to each other. This results in a short bond length and highly effective orbital overlap, creating a very strong bond.

However, as the halogen atom gets larger (like Iodine), its outermost electrons are far away from its nucleus. The hydrogen atom cannot get as close, resulting in a much longer bond length.

There is a fundamental rule in chemistry: Longer bonds are weaker bonds. Because the shared electron pair is further from the nuclei, the electrostatic attraction holding the atoms together is diminished. Therefore, the bond dissociation energy decreases down the group.

We can also verify this trend using thermodynamics. The standard free energy of formation, $\Delta G_f^{\circ }$, tells us about the intrinsic stability of a compound relative to its constituent elements.

For Hydrogen Fluoride ($HF$), $\Delta G_f^{\circ }=-273.2{\text{ kJ mol}}^{-1}$. This highly negative value indicates that $HF$ is exceptionally stable.

On the other extreme, for Hydrogen Iodide ($HI$), $\Delta G_f^{\circ }=+1.72{\text{ kJ mol}}^{-1}$. A positive free energy of formation means that $HI$ is thermodynamically unstable at standard conditions and readily decomposes into $H_2$ and $I_2$ when heated.

Combining all these concepts, we can confidently conclude that thermal stability decreases as we move down the halogen group. The correct order of thermal stability is: