Wednesday
October 22, 2014

Homework Help: Chemistry

Posted by Donna on Tuesday, May 13, 2008 at 5:23am.

You are carrying out the titration of 100.0 mL of 1.000M acetic acid with 1.000 M sodium hydroxide. Ka = 1.76x10^-5 of acetic acid.

(a) Calculate the initial pH of your acetic acid sample.
(b) Calculate the pH of the solution after the addition of 25.0mL of NaOH.

For (a) should I just take the -log of the given Ka value, which is 4.75?

or should I do this:
CH3COOH <--> CH3COO- + H+
1.0 0 , 0
-x +x , +x
1.0-x x , x

1.0-x can be approximated to 1.0
so, x^2/1.0=1.76x10^-5
x = 4.20x10^-3, and x = [H+]
-log [H+] = pH = 2.38

I'm not sure if should just be 2.38 or 4.75


then part B, I set it up like this:
so convert both CH3COOH and NaOH to moles
.1 CH3COOH and .025 NaOH

CH3COOH + NaOH <--> CH3COONa + H2O



I subtracted moles of NaOH from CH3COOH on the reactants side, and added moles of NaOH to CH3COONa (which starts at zero) on the products side

.1-0.025 = 0.075 of CH3COOH, and 0+0.025 = 0.025 of CH3COO-

because these are conjugates, I can use Henderson-Hasselbalch:

pH = pKa + log 0.025/0.075
pH = 4.27

Is this correct? Can the pH go down when I'm adding a strong base (even if it's a smaller amount) to a weak acid?

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