13.11

Moderators: Chem_Mod, Chem_Admin

Josh Ku 3H
Posts: 25
Joined: Wed Sep 21, 2016 2:59 pm

13.11

Postby Josh Ku 3H » Wed Nov 30, 2016 5:01 pm

In this question it presents you with a buffered solution with acetic acid in an equal concentration with sodium acetate. The question asks for the pH change when different amounts of NaOH are introduced.

By introducing hydroxide ions it causes the concentration of acetic acid to decrease and the concentration of acetate ion to increase. What I'm confused on is why the concentration of the acetate ion increases by the exact amount of the concentration of the hydroxide ions. The solution manual says it's because NaOH is a strong base but shouldn't the amount of acetate ion formed be based on the strength of acetic acid and its willingness to give up its protons?

Chem_Mod
Posts: 18139
Joined: Thu Aug 04, 2011 1:53 pm
Has upvoted: 421 times

Re: 13.11

Postby Chem_Mod » Wed Nov 30, 2016 6:04 pm

By introducing NaOH, OH- ions react with acetic acid, specifically protons (H+). These two ions neutralize each other. Since the stoichiometry is all at a 1:1 ratio, the mols of H+ depleted will equal the number of moles of acetic acid that need to dissociate to retain equilibrium in the buffer, which is equal to the number of moles of acetate ion formed as a result.


Return to “*Making Buffers & Calculating Buffer pH (Henderson-Hasselbalch Equation)”

Who is online

Users browsing this forum: No registered users and 1 guest