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### Topic: HCl added to a methylamine solution (buffer)  (Read 315 times)

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#### Baksu888

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##### HCl added to a methylamine solution (buffer)
« on: March 11, 2020, 06:26:54 PM »
How many mol of HCl would you have to add to 0.25 liter of 0.200 M methylamine solution to make a buffer with a pH of 11.00?

I tried plugging it into the Henderson Hasselbalch Equation:

CH3NH2 + H+    CH3NH3

$$K_a = \frac {1.0*10^{-14}} {4.4*10^{-4}} = 2.3*10^{-11}$$
$$0.25L * 0.200M = 0.050 \ mol$$
$$mol \ of\ CH_3NH_2=0.050-x$$
$$mol\ of\ CH_3NH_3=x$$
$$11.00 = -log(2.3*10^{-11})+log \frac {x} {0.050-x}$$

When I solved the equation, I got 0.035 mol of HCl. However, the answer for this question is supposed to be 0.015 mol. Also, when I tried to enter the equation into an online smart calculator (wolfram), I ended up with no real answers.

Was my method of calculation incorrect? Did I make a mistake in my calculations?

EDIT: I realized my stupid mistake. The Henderson Hasselbalch Equation calls for base over acid, and I accidentally did acid over base. So it should be the following:
$$11.00 = -log(2.3*10^{-11})+log \frac {0.050-x} {x}$$
After solving the equation, I got 0.015 mol of HCl, which was the right answer.