🤓 Based on our data, we think this question is relevant for Professor Marks' class at ASU.

*Whenever we have a conjugate acid and a weak base we have a buffer.*

**When we are dealing with buffers, we can use the ****Henderson-Hasselbalch equation:**

$\overline{){\mathbf{pOH}}{\mathbf{}}{\mathbf{=}}{\mathbf{}}{{\mathbf{pK}}}_{{\mathbf{b}}}{\mathbf{}}{\mathbf{+}}{\mathbf{}}{\mathbf{log}}\mathbf{\left(}\frac{\mathbf{Conjugate}\mathbf{}\mathbf{acid}}{\mathbf{weak}\mathbf{}\mathbf{base}}\mathbf{\right)}}$

pK_{b} of (CH_{3})_{3}N is 4.194

Calculate pOH:

$\mathbf{pOH}\mathbf{}\mathbf{=}\mathbf{}\mathbf{4}\mathbf{.}\mathbf{194}\mathbf{}\mathbf{+}\mathbf{}\mathbf{log}\mathbf{\left(}\frac{\mathbf{0}\mathbf{.}\mathbf{12}}{\mathbf{0}\mathbf{.}\mathbf{08}}\mathbf{\right)}\phantom{\rule{0ex}{0ex}}\mathbf{pOH}\mathbf{}\mathbf{=}\mathbf{}\mathbf{4}\mathbf{.}\mathbf{194}\mathbf{}\mathbf{+}\mathbf{}\mathbf{0}\mathbf{.}\mathbf{176}$

**pOH = 4.37**

Use information from Appendix D in the textbook to calculate the pH of the following solutions.

Calculate the pH of a solution that is 0.080 *M* in trimethylamine, (CH_{3})_{3}N, and 0.12 *M* in trimethylammonium chloride, ((CH_{3})_{3}NHCl).