# How to tell the pH of HC2H3O2?

Find the $\mathrm{pH}$ of a $\pu{0.500 M}$ solution of $\ce{KC2H3O2}$.

I am preparing for an exam on equilibrium and I have two questions:

1. How would I go about calculating this?
2. Is the $\ce{H3O}$ in the middle of $\ce{KC2H3O2}$ one hydroxide ion?

Here is what I have attempted to do: find the conjugate acid/bases, write out a balanced equation, then take the $-\log{[\ce{H+}]}$ of the conjugate acid. But I am a little stuck on the best course of action to find the conjugate acid/base.

• Welcome to Chemistry Stack Exchange! Please add what you have attempted towards solving the problem into the body of your question. For more information, see the site's homework policy for how to ask homework questions. Thanks! – jonsca May 15 '14 at 22:28
• Thanks for the reply. I have added more info and context to the question. – Jake Chasan May 15 '14 at 22:30

First of all let us introduce the given compound in a more bonding descriptive way: $\ce{H3C-COOK}$. This compound is commonly referred to as potassium acetate: It reacts with water in the following fashion: $$\ce{H3C-COOK_{(aq)} + H2O <=> H3C-COOH_{(aq)} + K+_{(aq)} + {}^{-}OH_{(aq)}}$$

Then the mass action law can be applied in the following way: $$K_c=\frac{\ce{[H3C-COOH]}\ce{[K+]}\ce{[{}^{-}OH]}}{\ce{[H3C-COOK]}\ce{[H2O]}}$$

And you can then rearrange to form the base dissociation constant: $$K_b=\frac{\ce{[H3C-COOH]}\ce{[{}^{-}OH]}}{\ce{[H3C-COOK]}}$$

The $\mathrm{p}K_b$ for potassium should be around $4.7$.

As dissenter pointed out, you do not need the Henderson-Hasselbalch equation to calculate $\mathrm{p}\ce{H}$ values for a pure solution of potassium acetate. You can have a look at wikipedia for an example calculation.

• Hydroxide Ion: $\ce{{}^{-}OH}$ or $\ce{OH-}$ or $\ce{HO-}$ ...
• Hydronium Ion: $\ce{H3+O}$ or $\ce{H3O+}$ or $\ce{H+OH2}$ ...
You need the $K_\mathrm{a}$ value of the solute of interest. That would be $\ce{[C2H3O2]-}$ as $\ce{K+}$ has highly limited Brønsted-Lowry acid/base properties in water.