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Briefly explain how the following mistakes would affect your determined pKa (or pKb) value: a. b. c. You do not transfer all the weighed acid to your flask You add 75 mL of DI water to your flask instead of the 50 mL in the procedure. Your determined equivalence point volume is greater than the true equivalence point.

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Related questions

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a. For the titration of 25.0 mL of 0.1 M HCl (aq) with 0.1 M NaOH (aq), at what volume of NaOH (aq) should the equivalence point be reached and why? If an additional 3.0 mL of 0.1 M NaOH (aq) is then added, what is the expected pH of the final solution?

b. What is the initial pH expected for a 0.1 M solution of acetic acid? For the titration of 25.0 mL of 0.1 M acetic acid with 0.1 M NaOH (aq), at what volume of NaOH (aq) is the equivalence point reached? Is the pH at the equivalence point greater than, less than or the same as in problem #a above? Explain.

c. What is the initial pH expected for a 0.1 M solution of phosphoric acid (H3PO4)? For 25.0 mL of 0.1M H3PO4 (aq), what volume of 0.1 M NaOH (aq) is required to fully titrate all three protons to their end points?

d. In the titration of a weak acid with a strong base, how is the half equivalence point determined and what is its significance? How are the pKa and Ka of the weak acid determined from the half equivalence point?

In step 4 of part A where you add distilled water to the acid KHP, suppose you add 40 mL DI h2o instead of 30mL to the flask to dissolve the KHP. Why would this affect the overall outcome of your data? Why/why not? Hint: try to determine whether or not your calculations are affected by this change.

this was a titration performed between KHP and NaOH solution

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