Determining Ka by the
Half-Titration of a Weak Acid
A common analysis of a weak acid or a weak base is to conduct a titration with a base or acid of known molar concentration to help determine the equilibrium constant, Ka, for the weak acid or weak base. If this titration is conducted very carefully and very precisely, the results can lead to a valid approximation of an equilibrium constant. In this experiment, however, you will use a different technique to determine the Ka for a weak acid, acetic acid.
Your primary goal in this experiment is to calculate the Ka of acetic acid. The data that you will use to complete your calculations will come from the reaction of acetic acid with a solution of NaOH. Recall from your work with weak acid-strong base titrations that the point at which a reaction is half-titrated can be used to determine the pKa of the weak acid. In this experiment, the half-titration point will exist when you have added half as many moles of HC2H3O2 as moles of NaOH . Thus, OH– will have reacted with half of the HC2H3O2, leaving the solution with equal moles of HC2H3O2 and C2H3O2 –. At this point, according to the Henderson-Hasselbalch equation,
Indigestion and Titration: An Acid-Base Titration
Imagine yourself as the Lead Analytical Chemist at Kaplan Industries. Your first big assignment is to investigate the strength of several commercial antacids for
Indicators are used to determine the end point of the titration. An indicator is used in acid-base and oxidation-reduction titrations. The color change of the indicator should be near the equivalence point of the reaction. The following charts show commonly used indicators and their color changes.
In acid-base titration experiment, a solution of accurately KHP
concentration was added gradually to another solution of NaOH concentration
until the chemical ... H+ ion) and
strong base ( contained OH ) were 100% ionized in water and they were all
The equilibrium that determines the pH of the solution isNH3(aq) + H2O(l) → NH4+(aq) + OH-(aq)As the strong acid is added to the solution, there is a quick dip in the pH as the concentration of weak base is depleted by the titration.
tablets counteract more of the original HCl, leaving the solution closer to neutral before the NaOH is added.)
Here are your results:
Mass of one dose
mL NaOH used in back-titration
1. Which is the strongest antacid, on a single-dose basis? Which is the weakest? Explain and show your calculations.
2. Which are the strongest and weakest, on a by-weight (mass) basis?
3. When people do back titrations, they usually watch the solution for a color change when the solution becomes neutral. What might you have used in the above
Fourth, slightly open the cork of the buret and add the standard reagentinto the unknown solution. Around the expected equivalence pointof the titration, you need to drop the solution very slowly and mix thesolutions very well because, around the equivalence point, just one dropof solution from the buret can make a radical pH change in the mixed solution. If the color of the solution in the erlenmeyer flask changes, record thevolume of the solution in the buret and add a few drops of the solutionto make sure the the equivalence point you found is correct.Finally, using the data from your acid-base titration, you can calculatethe concentration of the unknown solution.
How is it that when the strong acid is added to the solution of weak base its conjugate acid, in roughly equal concentrations, the pH changes very little?
When the strong acid is added to the solution of the weak base and its conjugate acid, the acid is immediately converted to acid through the neutralization reactionNH3(aq) + H3O+(aq) → NH4+(aq) + H2O(l)The addition of an extra quantity of acid will not effect the pH of the solution in a substantial way, and the pH changes very little.
In this experiment, you will
• Conduct a reaction between solutions of a weak acid and sodium hydroxide.
• Determine the half-titration point of an acid-base reaction.
• Calculate the Ka and pKa for the weak acid.
LabQuest 1.00 M sodium hydroxide, NaOH, solution
LabQuest App 1.00 M acetic acid, HC2H3O2, solution
Vernier pH Sensor phenolphthalein indicator solution
50 mL buret distilled water
buret clamp magnetic stirrer and stirring bar
250 mL beaker plastic Beral pipets
two ring stands utility clamp
1. Obtain and wear goggles.
2. Use a buret clamp to connect a 50 mL buret to a ring stand. Rinse and fill the buret with
1.00 M acetic acid solution. Handle the acetic acid with care. It can cause painful burns if it comes into contact with the skin.
3. Transfer precisely 25.0 mL of the acetic acid solution to a 250 mL beaker.
4. Use a plastic Beral pipet to remove a small volume of the acetic acid from the 250 mL beaker. Draw enough acetic acid into the pipet so that the bulb is about 1/4 full. Carefully set aside the pipet of acid, to be used later.
5. Add 1–2 drops of phenolphthalein indicator solution to the beaker of acetic acid.
6. Connect the pH Sensor to LabQuest and choose New from the File menu. If you have an older sensor that does not auto-ID, manually set up the sensor.
7. Obtain about 50 mL of 1.00 M NaOH solution. CAUTION: Sodium hydroxide solution is caustic. Avoid spilling it on your skin or clothing.
8. Begin the half-titration.
a. Place the beaker of acetic acid on a magnetic stirrer and add a stirring bar.
b. Set up a ring stand and clamp to hold the pH Sensor in place (see Figure 1). Position the pH Sensor in the beaker so that the tip of the probe is completely immersed.
c. Gently stir the acetic acid solution.
d. Do not start data collection. Monitor the pH of the reaction mixture on LabQuest.
e. Use a new plastic Beral pipet to slowly add the 1.00 M NaOH solution, in ~1 mL increments, to the beaker of acetic acid solution (see Figure 1).
9. Conduct the titration carefully. As the reaction approaches the equivalence point, at about pH 6, add the NaOH solution drop by drop. When you reach the equivalence point, the pH will increase rapidly and the indicator will change color. If necessary, add another drop of NaOH, so that the reaction is slightly past the equivalence point. Remember that the pH will not increase rapidly beyond the equivalence point (pH ~10).
10. Add all of the acetic acid from the Beral pipet, which you removed in Step 4, to the beaker of reaction mixture. Check the pH readings and observe the indicator color. The mixture should be slightly acidic once again.
11. Carefully add NaOH, drop by drop, to the beaker of reaction mixture, until you reach the equivalence point as precisely as possible. A very slight pink color of the phenolphthalein indicator is visible. This is your half-titrated solution, because you have neutralized precisely 25.0 mL of the original 50.0 mL of acetic acid that you measured out into the buret.
12. Transfer the remaining 25.0 mL of acetic acid from the buret to the 250 mL beaker of reaction mixture. Stir the solution in the beaker thoroughly. Read and record the pH of the solution in the beaker.
13. When you have finished the testing, dispose of the reaction mixture as directed. Rinse the pH Sensor with distilled water in preparation for a second trial. Repeat the necessary steps to test a new sample of the acetic acid solution.
To titrate an unknown acid/base solution, take a certain amount of theunknown solution and add a standard reagent of the known concentrationcarefully until the neutralization reaction is completed.