Question

Choose the stronger acid: (a) \(\mathrm{H}_{2} \mathrm{~S}\) or \(\mathrm{H}_{2} \mathrm{Se},\) (b) \(\mathrm{H}_{2} \mathrm{Te}\) or \(\mathrm{HI},\) (c) \(\mathrm{PH}_{3}\) or \(\mathrm{NH}_{3}\). Give your reasons.

Short answer

\(\mathrm{H}_{2}Se\) is a stronger acid than \(\mathrm{H}_{2}S\) due to weaker bond strength. \(\mathrm{HI}\) is a stronger acid than \(\mathrm{H}_{2}Te\) due to bond polarity. \(\mathrm{NH}_{3}\) is a slightly stronger acid than \(\mathrm{PH}_{3}\), though both are very weak acids.

Step-by-step solution

Understanding Acid Strength

The strength of a binary acid (one composed of hydrogen and one other element) in the same group of the periodic table increases as the size of the central atom increases. This is due to the bond strength decreasing as the atoms get larger, making it easier for the hydrogen to ionize. Likewise, for non-metal hydrides, bond polarity also affects acidity, with more polar bonds typically lending to stronger acids.

Comparing Acids in Group 16 (Chalcogens)

Compare \(\mathrm{H}_{2}S\) and \(\mathrm{H}_{2}Se\) which are both in Group 16. Since selenium (Se) is larger than sulfur (S), the bond in \(\mathrm{H}_{2}Se\) is weaker than in \(\mathrm{H}_{2}S\). Hence, \(\mathrm{H}_{2}Se\) is a stronger acid than \(\mathrm{H}_{2}S\). Similarly, compare \(\mathrm{H}_{2}Te\) with \(\mathrm{HI}\). Tellurium (Te) is larger than iodine (I), so based on size, \(\mathrm{H}_{2}Te\) could be expected to be more acidic. However, because \(\mathrm{HI}\) is a binary acid of a halogen and due to its significant bond polarity, it is a stronger acid than \(\mathrm{H}_{2}Te\).

Comparing Non-Metal Hydrides

When comparing \(\mathrm{PH}_{3}\) and \(\mathrm{NH}_{3}\), neither is a strong acid. However, due to the increase in electronegativity from phosphorus to nitrogen and the corresponding increase in bond polarity, \(\mathrm{NH}_{3}\) would be expected to be a slightly stronger acid. In practice, both of these compounds exhibit very weak acidic behavior, and their basic behavior tends to be more relevant.

Key concepts

Understanding Binary Acids

In chemistry, binary acids are simple compounds consisting of hydrogen and one other element. These acids are most commonly found across the nonmetals of the periodic table. One fundamental factor in determining the strength of binary acids is the size of the central atom. Specifically, as you move down a group in the periodic table, atoms increase in size. Larger atoms bond to hydrogen with less force, which makes the bond easier to break and the hydrogen ionize. The result is an increase in acid strength. For example, hydrogen sulfide (ewline{H}_{2}S) is a binary acid with sulfur as its central atom. When comparing it to hydrogen selenide (ewline{H}_{2}Se), which has the larger selenium atom, one can deduce that ewline{H}_{2}Se is the stronger acid because selenium forms weaker bonds with hydrogen due to its larger size.

Another point to consider when studying binary acids is their tendency to donate protons to water molecules, releasing hydronium ions (ewline{H}_{3}O^{+}) and thus exhibiting acidic characteristics. The ease of this proton donation is a clear indicator of acid strength. By applying this knowledge along with periodic trends, students can compare and predict the strengths of various binary acids.

Periodic Trends in Acidity

As we further explore acidity throughout the periodic table, it's important to grasp the correlation between periodic trends and acid strength. Acidity typically increases from left to right across a period because of increasing electronegativity, which allows bonds to hydrogen to become more polar and prone to ionization. However, for binary acids specifically, the trend that predominates is the one down a group. As the atomic radius of the central non-metal atom increases, the bond to hydrogen becomes weaker, causing these acids to readily release hydrogen ions and act as stronger acids.

ewline{H}_{2}Te, a compound larger in atomic size than ewline{HI}, would be expected to be a stronger acid purely based on atomic size, but electronegativity and bond polarity play a defining role. Iodine's high electronegativity generates a strong bond polarity with hydrogen, making ewline{HI} more acidic than one would predict if considering size alone. This illustrates how periodic trends can sometimes interplay in a nuanced fashion, and why understanding the full context of element properties is critical when assessing acid strength.

Bond Polarity and Acidity

Bond polarity is a concept that arises due to differences in electronegativity between two atoms bonded together. When assessing the acidity of a molecule, paying attention to the polarity of the bond between hydrogen and the other atom is crucial. A greater difference in electronegativity leads to a more polar bond, and polar bonds typically mean a hydrogen ion can be more easily removed, yielding a strong acid.

For instance, in non-metal hydrides like ewline{PH}_{3} and ewline{NH}_{3}, although both show limited acidic behavior, the bond polarity gives us insight into their relative strengths. Nitrogen has a higher electronegativity compared to phosphorus, meaning the ewline{N-H} bonds in ewline{NH}_{3} are more polar than the ewline{P-H} bonds in ewline{PH}_{3}. Therefore, ewline{NH}_{3} would theoretically be a stronger acid. However, it's important to remember that acidity is also influenced by other factors such as molecular structure and the presence of additional atoms or groups within the molecule, which can affect the overall distribution of electron density and consequently, acid strength.