Question

Choose the stronger acid in each of the following pairs: (a) \(\mathrm{H}_{2} \mathrm{Se}\) or \(\mathrm{H}_{3} \mathrm{As}\) (b) \(\mathrm{B}(\mathrm{OH})_{3}\) or \(\mathrm{Al}(\mathrm{OH})_{3}\) (c) \(\mathrm{HBrO}_{2}\) or HBrO

Short answer

The stronger acids are \( \text{H}_{2} \text{Se} \, \text{B}(\text{OH})_{3} \, and \(\text{HBrO}_2 \)

Step-by-step solution

Choose the stronger acid between \( \text{H}_{2} \text{Se} \) and \( \text{H}_{3} \text{As} \)

Determine which atom selenium (Se) or arsenic (As) has higher electronegativity and follows the periodic trend. Moving down a group in the periodic table, acid strength generally increases due to larger atomic size and weaker bond strength between hydrogen and the non-metal. Since Se is below As in the periodic table, \( \text{H}_{2} \text{Se} \) is the stronger acid.

Choose the stronger acid between \( \text{B}(\text{OH})_{3} \) and \( \text{Al}(\text{OH})_{3} \)

Consider the elements boron (B) and aluminum (Al). Aluminum being a metal forms a typical weak base with hydroxide ions, whereas boron, a metalloid, forms \( \text{B}(\text{OH})_{3} \) that acts as a Lewis acid. \( \text{B}(\text{OH})_{3} \) is the stronger acid due to its ability to accept hydroxide ions.

Choose the stronger acid between \( \text{HBrO}_{2} \) and \(\text{HBrO} \)

Compare the acids based on the oxidation state of bromine (Br). \(\text{HBrO} \) has a lower oxidation state (+1) compared to \( \text{HBrO}_2 \) which has an oxidation state of +3. Acids with higher oxidation states are generally stronger. Therefore, \( \text{HBrO}_2 \) is the stronger acid.

Key concepts

electronegativity

Electronegativity is a measure of an atom's ability to attract and hold onto electrons. The higher the electronegativity, the stronger the atom pulls electrons towards itself. When comparing acids like \(\text{H}_{2} \text{Se} \) and \(\text{H}_{3} \text{As} \), consider that a higher electronegativity often leads to stronger acids. However, other factors such as atomic size play a role too. Selenium (Se) is more electronegative than arsenic (As), but it lies below arsenic in the periodic table, meaning \(\text{H}_{2} \text{Se} \) is a stronger acid due to a larger atomic size and weaker bond with hydrogen.

periodic trends

Periodic trends help us understand how properties of elements change across the periodic table. For acids, some trends like atomic size and electronegativity are key. As you move down a group, atoms become larger, making hydrogen bonds weaker, thus creating stronger acids. For example, in the pair \(\text{H}_{2} \text{Se} \) and \(\text{H}_{3} \text{As} \), selenium (Se) is below arsenic (As), thus \(\text{H}_{2} \text{Se} \) is the stronger acid due to a weaker bond with hydrogen, following the periodic trend of increasing acid strength down a group.

oxidation states

The oxidation state of an atom in an acid can influence its strength. Higher oxidation states generally mean stronger acids because they pull more electron density away from the hydrogen, making it easier to lose a proton. For example, \(\text{HBrO}_{2} \) has a higher oxidation state (+3) compared to \(\text{HBrO} \) (which is +1), making \(\text{HBrO}_{2} \) the stronger acid. This is because the Br in \(\text{HBrO}_{2} \) is more electron-withdrawing due to its higher oxidation state, enhancing its acid strength.

Lewis acids

A Lewis acid is a substance that accepts a pair of electrons. In our comparison, \(\text{B}(\text{OH})_{3} \) acts as a Lewis acid because boron can accept electrons due to its empty p-orbital. Aluminum in \(\text{Al}(\text{OH})_{3} \), being a metal, typically forms a weak base. Therefore, \(\text{B}(\text{OH})_{3} \) is the stronger acid as it is capable of accepting electron pairs, whereas \(\text{Al}(\text{OH})_{3} \) does not accept electrons as effectively, making it less acidic.