pH of a solution of acetic acid and ammonium acetate After studying ionic equilibrium, I was just making some wild questions and I was unable to calculate the $\ce{pH}$ of this solution. Assume, 1 mole of each was added to $1L$ solution of water. 

My attempt : 

$\ce{CH3COOH}+\ce{NH4OH&lt;=&gt;CH3COONH4}+\ce{H2O}$

This reaction will happen till all $\ce{CH3COOH}$ is consumed by $\ce{NH4OH}$. More of $\ce{NH4OH} $ will be formed by Le Chatelier's principle. So, the answer must be same as that of a weak acid and weak base salt : $pH=\dfrac{1}{2}(pK_w+pK_a-pK_b)$.
But how will we calculate if there was $2$ mole of acetic acid?
You have some mistaken assumptions.

$\ce{NH4OH}$ and $\ce{CH3COONH4}$ are not molecules that exist.  

The assuption "all $\ce{CH3COOH}$ is consumed" is incorrect.  

Why do you say more $\ce{NH4OH}$ will be formed? How is this possibe? What can it be formed from? 

Review that Henderson–Hasselbalch equation