Predict which acid in each of the following pairs is the stronger and explain your reasoning for each. (a) \(\mathrm{H}_{2} \mathrm{O}\) or \(\mathrm{HF}\) (b) \(\mathrm{B}(\mathrm{OH})_{3}\) or \(\mathrm{Al}(\mathrm{OH})_{3}\) (c) \(\mathrm{HSO}_{3}^{-}\) or \(\mathrm{HSO}_{4}^{-}\) (d) \(\mathrm{NH}_{3}\) or \(\mathrm{H}_{2} \mathrm{S}\) (e) \(\mathrm{H}_{2} \mathrm{O}\) or \(\mathrm{H}_{2} \mathrm{Te}\)

Understanding acid-base reactions is foundational in chemistry. An acid is a substance that can donate a proton (H⁺), while a base is a substance that can accept a proton. The strength of an acid or base is indicative of how readily it donates or accepts protons. In water (H₂O) versus hydrofluoric acid (HF), HF is the stronger acid because it donates protons more easily, as seen in its reaction with water or other bases.

These reactions are also governed by equilibrium constant expressions called acid dissociation constants (K_a), which provide a quantitative measure of an acid's strength. The larger the K_a, the stronger the acid. For instance, HF has a larger K_a than H₂O, confirming its stronger acidic properties.

Electronegativity is a key factor in determining acid strength. It describes how strongly an atom attracts electrons in a chemical bond. Stronger acids often have a more electronegative element involved, as it can stabilize the negative charge on the conjugate base after the acid donates its proton. For instance, fluorine in HF is highly electronegative, making it a stronger acid than H₂O where the central atom is the less electronegative oxygen.

High electronegativity in atoms like fluorine helps to distribute the negative charge more evenly when it becomes the fluoride ion (F^-), the conjugate base of HF. This stability is a driving force for the acid to release its proton, thus increasing acid strength.

Oxyacids contain hydrogen, oxygen, and another element—often a non-metal. The acidity trend for oxyacids can be explained by the number of oxygen atoms present. Oxyacids with a greater number of oxygen atoms are generally stronger acids because these additional oxygens can more effectively delocalize the negative charge through resonance in the conjugate base.

For example, in comparing bisulfite (HSO₃^-) and bisulfate (HSO₄^-), the bisulfate with four oxygen atoms is the stronger acid due to the better charge distribution across multiple oxygen atoms. This trend holds true in many similar comparisons of oxyacid strength.

Binary acids, composed of hydrogen and one other non-oxygen element, display a trend in strength based on the polarizability of the anion. Polarizability refers to how easily an electron cloud can be distorted. Larger anions are more polarizable, which leads to a higher propensity to hold onto the negative charge when the hydrogen ion is released.

Take, for example, ammonia (NH₃) and hydrogen sulfide (H₂S). Since sulfur is larger than nitrogen, H₂S is the stronger acid. This principle can be applied across the periodic table to predict the strength of binary acids, with larger central atoms correlating to stronger acids.

The stability of a conjugate base plays a crucial role in the strength of its corresponding acid. A stable conjugate base results from an acid releasing a hydrogen ion (H⁺), and the more stable this base, the stronger the original acid. This stability is influenced by several factors, including electronegativity, resonance, and atomic size.

For instance, the fluoride ion (F^-), being the conjugate base of HF, gains stability from fluorine's high electronegativity, while the bisulfate ion (HSO₄^-), the conjugate base of sulfuric acid, is stabilized by resonance structures. In general, if the resulting conjugate base can distribute its negative charge evenly or achieve lower energy through stability, its corresponding acid will be stronger.