Chem 534 Name_______________________

Titration and Ka Acidic Partner_________________

Purpose: To neutralize a weak acid of unknown concentration with NaOH and to calculate the Ka of the weak acid.

Procedure:

1. Make sure that the plastic handle on the burette makes a 90^o angle with the glass column--in other words, it has to be closed.

2. Place a funnel in the burette(or if no funnel is available, use the dropper bottle), and slowly pour out 0.10 M solution of NaOH until the volume is a bit above 0.00 mL.

3. Let out some NaOH into a waste beaker until it reads 0.00 mL. The purpose of this step is to fill any air bubbles at the tip of the burette.

4. a. Into an erlenmeyer flask, carefully measure 20.0 mL of HCH₃COO using a burettete.

b. Measure the pH of the acid with pH paper.

pH of acid =

5. Add 2 drops of phenolphthalein indicator to the HCH₃COO in the erlenmeyer flask. Swirl to mix.

6. To the same erlenmeyer, begin your titration by adding NaOH from the burette, slowly, drop by drop, until a permanent light pink colour appears. Swirl the flask continuously.

7. Record the volume of NaOH needed to neutralize the acid.

V =Volume of NaOH (mL)

Analysis

1. Write an equation to show what happens when NaOH neutralizes HCH₃COO.

2. Use moles = CV (C = concentration of NaOH) to find out the number of moles of NaOH that were needed to neutralize the HCH₃COO.

3. Using the molar ratio from equation 1, calculate the number of moles of HCH₃COO present in the original acidic solution..

4. Divide that number of moles by the volume used (see step 4a) to get the concentration of the acid.

5. The calculated concentration of the acid is the equilibrium concentration of HCH₃COO. Use this information and the pH from step 4b to calculate K_a. pH will correspond to the equilibrium concentration of H⁺¹.

HCH₃COO = H⁺¹ + CH₃COO^-1.

	HCH3COO	H^+1	CH3COO^-1.
moles/L at equilibrium

K_a =

Conclusion: