Titration with weak acid and strong base

[stardusto12]

Homework Statement

We did a lab in class about weak acid/strong base titration, but I'm having a bit of trouble. We we only given the molarities of an unknown weak acid and a strong base (NaOH) and the pH of six different mixtures containing different mL of the two, which is below:

Test tube #1: 9mL of weak acid and 1mL of NaOH pH=3.04
Test tube #2: 8mL of weak acid and 2mL of NaOH pH=3.07
Test tube #3: 7.5mL of weak acid and 2.5mL of NaOH pH=3.40
Test tube #4: 7mL of weak acid and 3mL of NaOH pH=3.56
Test tube #5: 6mL of weak acid and 4mL of NaOH pH=3.80
Test tube #6: 5.5mL of weak acid and 4.5mL of NaOH pH=4.21

and the molarity of the solutions were .1176 for the unknown acid and .1059 for NaOH.

Homework Equations

I was able to calculate the intial moles of the acid and base but were I'm having difficulty is figuring out the moles at equilibrium. (of A- and HA in HA + H2O = H3O +A-).

The Attempt at a Solution

I'm very confused and was only able to get the intial moles of the acid and base. I know I'm supposed to use the Henderson-Hasselbaclh equation, but I don't know how to get log of (HA/A-) since I don't have HA or A-.

Thanks!

 

[symbolipoint]

The formalities can be obtained directly from how each solution is prepared. What is the question which you are trying to answer? Are you trying to find a value for the equilibrium constant for the weak acid?

 

[stardusto12]

Yes, I am trying to find the equilibrium constants, but of both. I am trying to find the equilibrium concentrations of the both the acid and base. Afterwards, I can calculate the rest on my own. I just don't know how to derive the equilibrium concentrations from the information given.

 

[symbolipoint]

K = (H)(Fs + H)/(Fa - H), The equilibrium constant expression for a weak acid.

H = the hydronium ion concentration as molarity
Fs = formality of the salt of the weak acid
Fa = formality of the weak acid

Examine that formula carefully to see that the molarity of the salt is Fs + H and the
molarity of the weak acid is Fa - H. The formula should be just fine for you because you
are dealing with only acidic solutions of the weak acid. I would not call your data a set
from an actual titration; the values seem to look wrong for that. You should still be able to
use each one of them to find any unknown value for the K formula.

 

[stardusto12]

thank you! Oh, what do you mean by formality? and is it the ph that's wrong? I believe my partner didn't clean it off well with the distilled water...

 

[symbolipoint]

thank you! Oh, what do you mean by formality? and is it the ph that's wrong? I believe my partner didn't clean it off well with the distilled water...

Formality means the concentration of a substance as formula units per liter based on the way the solution was prepared, with no regard for reactions of the substance - maybe not the best way to describe the meaning of Formality, but to be certain, check the meaning of this in an analytical chemistry textbook.

If you mix 1 mole of acid HA with 0.25 mole NaOH, then you still began with 1 mole of HA. You now have 0.75 moles unneutralized HA and 0.25 moles of NaA.

If you mix 100 ml of 1 M HA with 25 ml of 1 M NaOH, then you have still
0.100/125 Formal HA (you can ignore the reaction, but the MOLARITY will be different).

 

[stardusto12]

Thanks!

 

[symbolipoint]

Formality means the concentration of a substance as formula units per liter based on the way the solution was prepared, with no regard for reactions of the substance - maybe not the best way to describe the meaning of Formality, but to be certain, check the meaning of this in an analytical chemistry textbook.

If you mix 100 ml of 1 M HA with 25 ml of 1 M NaOH, then you have still
0.100/125 Formal HA (you can ignore the reaction, but the MOLARITY will be different).

Actually, I stated the latter part wrongly.
The formality would be (0.100 liter)*(1 mole/liter)/(0.125 liter) for the acid HA.