15.17 Calculate the pH of a 1.5 M solution of each of the following compounds (see Appendix E for \( K \) values): (a) \( \mathrm{C}_{5} \mathrm{H}_{5} \mathrm{~N} \); (b) \( \mathrm{HONH}_{2} \); and (c) \( \mathrm{HCO}_{2} \mathrm{H} \).

To calculate the pH of each of these compounds, we first need to identify their nature as acids or bases, along with their dissociation constants (\( K_a \) or \( K_b \)). For \( \mathrm{C}_{5} \mathrm{H}_{5} \mathrm{~N} \) (Pyridine), it is a weak base, and we can use its \( K_b \) value to find the \( K_a \) of its conjugate acid and then calculate the pH. For a 1.5 M solution, the equilibrium calculation gives us a pOH, which we then convert to pH. For \( \mathrm{HONH}_{2} \) (Hydrazine), it is also a weak base, so we apply the same method using its \( K_b \) value. Lastly, \( \mathrm{HCO}_{2} \mathrm{H} \) (Formic acid) is a weak acid. Here, we can directly use its \( K_a \) to find the concentration of \( \mathrm{H}^+ \) ions in solution, allowing us to calculate the pH straightforwardly. Now, let’s get to solving these equations! Make sure to look up the specific \( K \) values for precise calculations. To avoid common mistakes, always ensure that you have the right dissociation constants and remember that pH is calculated from the concentration of \( \mathrm{H}^+ \) ions, while pOH is calculated from the concentration of \( \mathrm{OH}^- \) ions. Misjudging if a compound is an acid or base can lead to significant errors in calculation.