Final answer:
The statement is false because a stronger H-X bond does not increase acidity; rather, weaker bonds make a compound more acidic, as they donate protons more readily. Acid strength is also influenced by inductive effects and charge delocalization.
Step-by-step explanation:
The statement 'The stronger the H-X bond, the more acidic the compound.' is false. When we discuss the strength of an acid, we're often considering how easily a compound can donate a proton (H+). A key principle in understanding this concept is that the weaker the bond between hydrogen and the rest of the molecule, the more readily it will donate a proton, increasing the acidity of the compound. For example, the hydrogen halides demonstrate this principle, with HCl (hydrochloric acid) being a stronger acid than HF (hydrofluoric acid), even though the H-F bond is stronger.
However, other factors also affect acidity. Inductive effects, where electron density is shifted through a molecule, can influence acidity. If these effects withdraw electron density from an O-H bond, for instance, they increase the compound's acidity. Similarly, charge delocalization that stabilizes the conjugate base will also enhance acidity.
Considering oxyacids and charged species, it is observed that acids with a higher electronegativity or those on positively charged species tend to be more acidic because the bonds are more polar and the protons are easier to remove.