# Writing Common Ion Concentration in Buffer Eq. Expression

1. May 24, 2013

### JeweliaHeart

1. The problem statement, all variables and given/known data
What is the pH of a buffer solution created by combining 100 mL of 0.2 M acetic acid and 400 mL of 0.10 M sodium acetate? Ka= 1.8 x 10^(-5)

2. Relevant equations

1.8 x 10-5=([C2H3O2-][H+])/([HC2H3O2])

3. The attempt at a solution

I know what I'm supposed to do to solve this problem, but I'm not sure how to set up the concentrations in the expression.

I think it should go like this:

1.8 x 10-5= (0.08+x)(x)/(0.04-x)

The explanation in my book has basically the same thing with just one tiny exception. In the equilibrium expression, it used only 0.08 as a concentration for C2H3O2-, without the x.

I thought the x would be necessary b/c 0.08 is only the initial concentration of the acetate ion and the acetic acid will dissociate more before reaching the equilibrium.

Any help would be nice. Thanks in advanced.

Last edited: May 24, 2013
2. May 24, 2013

### AGNuke

There's a simple formula to it. Search for it. I believe it's Henderson-Hasselbalch Equation.

3. May 31, 2013

### JeweliaHeart

Oh, thanks. I just figured it out.

4. May 31, 2013

### Staff: Mentor

While you are right about the fact acid can dissociate a little bit, in most typical situations changes induced by the dissociation are so small, we can safely ignore them. Then calculating pH of a buffer is just a matter of calculating concentrations of acid and conjugate base, and plugging them into Henderson-Hasselbalch equation.

Actually you don't even need to calculate concentrations, it is enough to calculate numbers of moles of acid and conjugate base, as volume cancels out.