This experiment is adapted from a method used by Dr. Dennis Barnum, an instructor, at Portland State University. An unknown sample of an organic acid is dissolved in boiled deionized water and titrated with a standard solution of sodium hydroxide. The titration reaction is followed potentiometrically by taking pH measurements with a pH meter as the standard base is added to the unknown acid solution. The data is used to construct a potentiometric titration curve from which the pK_a of the acid and the equivalence point are determined. Based on the equivalence point, the concentration of the unknown acid is determined.

dried primary standard potassium hydrogen phthalate (KHP)

phenothalein indicator

pH 7.0 buffer

pH 4.0 buffer

0.10 N NaOH (prepared in boiled distilled water)

dried organic acid unknown

50 mL buret

pH meter

capped liter bottle

Procedure 1 : Preparation of Standardized Sodium Hydroxide

1. Boil approximately 800 mL of deionized water, cover with a watch glass, and cool to room temperature. A portion of this water will be used for the next procedure. After the water has cooled pour it into a capped bottle.
2. While the deionized water is boiling, calculate the weight of KHP required so that 40 mL of 0.10 N NaOH will be used in the titration.
3. Weigh out four samples of KHP into 250-mL Erlenmyer flasks and dissolve in about 50 mL of boiled water. Dissolution will be faster while the water is still warm, but the indicator should not be added until room temperature has been reached.
4. Add five drops of phenothalein indicator.
5. Fill a 50 mL buret with 0.10 N NaOH and titrate the KHP solution to the first permanent pink color.
6. Repeat the titration for each of the KHP solutions.
7. Calculate the average normality of the sodium hydroxide solution.

Procedure 2: Titration of an Unknown Organic Acid

1. Measure out three 50.0 mL aliquots of unknown into 250 mL Erlenmyer flasks. Record the volumes in your laboratory notebook.
2. Calibrate the pH meter using the pertinent manual that accompanies the instrument. A two point calibration is required using the pH 4.00 and 7.00 buffers.
3. Fill the buret with the standardized 0.10 N NaOH. Adjust the miniscus to the 0.00 mL mark. Record this reading in your laboratory notebook.
4. Place the pH probe into the solution of unknown. Be sure to stir the solution as the titrant is added. Watching the pH meter, added enough titrant to cause a pH change of 0.2 to 0.3 pH units, then stop and record the pH and buret reading in your laboratory notebook. In the flat portion of the titration curve the pH changes rather slowly, so relatively large portions of titrant must be added in order to cause a change of 0.2 pH units.
5. Continue to add the titrant, taking pH readings every 0.2 to 0.3 pH units and buret readings. As the equivalence point is approached it will take smaller increments of volume to produce the desired change in pH. Record both the pH and buret reading in your laboratory notebooks.
7. Beyond the equivalence point larger volumes will again be required to change the pH by 0.2 to 0.3 pH units. Continue to add titrant until a total of 50 mL of NaOH has been added.
8. Check the standardization of the pH meter again to make sure it has not drifted.
9. Repeat the titration for each of the other unknown samples.
10. When you have finished the titrations, rinse off the pH electrode and immerse the probe in a solution of pH 7.00 buffer.

Using the first set of data, plot the entire titration curve, with mL of NaOH along the x-axis and corresponding pH readings along the y-axis. (See the sample titration curve that is attached.) This curve is to be turned in with your report. From the titration curve, determine the region where the equivalence point lies, then plot a second expanded graph that extends from 2 mLs before the equivalence point to 2 mL beyond the equivalence (See the sample expanded curve that is attached.) Using this expanded curve, determine where the inflection point is, and the corresponding mL of NaOH. Mark this point on your curve as is done in the attached sample. This is the equivalence point. Plot only the extended graphs for each of your other three unknowns. Determine the equivalence points for each of them. Calculate an average equivalence volume of NaOH. Using this volume, determine the number of equivalents of present in the unknown sample and the equivalent weight of the unknown sample. From the equivalence point volume of NaOH, determine the NaOH volume where pH = pK_a (Hint: Remember that pK_a = pH at the half-way point for the reaction.). Calculate pK_a for all four samples and determine the average pK_a for this acid.

Write up a laboratory report using the standard format. Include your calculations to standardized the NaOH. Include your calculations to determine the number of equivalence in the unknown sample, the equivalent weight for the unknown sample, and the pK_a for the unknown sample.. Include in the report your one complete titration curve showing where the pK_a point marked and the four expanded curves with the equivalence points marked.