You prepare 1.0 L of a 0.25 M acetic acid solution with a final pH of 6.0. What are the molar concentrations of all relevant acetic acid species ([HA] and [AX−]) given that the Ka for acetic acid is 1.74⋅10−5 M?
I am getting confused with this problem.
Since the pH is given, I know what the [HX+] is. So now when I try to do the ICE table
HAHX+AX−Initial0.2500Change−x+x+xEnd0.25−x+x+x
And from here I begin to assume [HX+]=[AX−], which I am not sure about. Then I set x=10−6.0=1⋅10−6 so I get [AX−]=1⋅10−6 M and [HA]=0.24999 M which I think is incorrect, and to even further ensure my that it's incorrect, when I attempt to check the Ka value with this, it does not match.
My second approach: [HA]=0.25 M
In this I determine the pKa from the Ka which turns out to be 4.759, which indicates that there should be more [AX−] than [HA].
I now use the Henderson–Hasselbalch equation: 6.0=4.759+log([AX−][HA])17.40=[AX−][HA]17.40=[AX−]0.25 M[AX−]=4.35 M
I feel more confident about my second answer.
Can someone please help me out with this particular problem and perhaps tell me procedure I should use as well as what the correct answer should come out to be and why?
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