Acid–vBase Titration with Conductivity

Introduction

Light up your students’ interest with this eye-catching demonstration! While demonstrating an acid-base titration, the conductivity of a solution will be tested. In addition to viewing the acid–base reaction, students will also observe a precipitation reaction.

Concepts

• Acids and bases
• Conductivity
• Precipitation

Materials

Safety Precautions

Barium hydroxide is acutely toxic by ingestion. Barium hydroxide is severely corrosive to eyes, skin and other tissue. Sulfuric acid solution is severely corrosive to eyes, skin and other tissue. It is also toxic by ingestion. Avoid all body tissue contact with all chemicals. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the laboratory. Follow all laboratory safety guidelines. Review current Safety Data Sheets for additional safety, handling and disposal information before performing the demonstration.

Disposal

Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulations that may apply, before proceeding. When the solution is colorless, filter the solids from the liquid and dispose of the solid according to Flinn Suggested Disposal Method #26c. Neutralize any leftover filtered liquid according to Flinn Suggested Disposal Method #24b. Leftover barium hydroxide solution can be precipitated by using Flinn Suggested Disposal Method #27h. Excess sulfuric acid can be neutralized according to Flinn Suggested Disposal Method #24b.

Procedure

1. Measure 250 mL of the 0.01 M barium hydroxide solution and place it in the 400-mL beaker.
2. Add a magnetic stir bar to the beaker and place the beaker on the magnetic stirrer.
3. Add 2–3 drops of phenolphthalein.
4. Set up the buret on a ring stand.
5. Rinse the buret twice with approximately 7 mL of 0.1 M sulfuric acid.
6. Fill the buret with 0.1 M sulfuric acid between the 0- and 5-mL mark.
7. Attach the conductivity tester to the other ring stand with a clamp. Make sure the conductivity tester can be placed in the beaker. See Figure 1 to view the set up.

Student Worksheet PDF

14098_Student1.pdf

Teacher Tips

• This kit contains enough chemicals to perform the demonstration seven times: 2 L of 0.01 M barium hydroxide, 500 mL of 0.1 M sulfuric acid solution and 30 mL of phenolphthalein solution.
• A conductivity tester is needed and sold separately. Flinn’s conductivity tester and light bulb work for this demonstration (AP5355 and AP5886).
• To graph data, a conductivity sensor, such as AP8113 and TC1507 can also be used.
• Turning the lights off in the classroom will help students to view the lightbulb.
• Remember to safely fill the buret. Before filling the buret, lower the buret. Never add chemicals above your head. Using a funnel, slowly add small amounts of the titrant.
• This demonstration can also be used in parallel with an AP Chemistry Sample Free-Response question.
• The demonstration kit, Audio Conductivity Tester—Multi-Demonstration Kit (Flinn Catalog No. AP7279) can be used to further explore conductivity of solutions.
• The student laboratory kit, Acidity of Beverages—Inquiry Lab Kit for AP^® Chemistry (Flinn Catalog No. AP7645) can also be used to further explore acid–base titrations.

Sample Data

Observations
The solution starts clear and the conductivity tester has a slight glow. When phenolphthalein is added, the solution turns pink. As sulfuric acid is added, the light bulb begins to dim. When the endpoint is reached, the bulb is completely out and the solution is a white/milky suspension. When the teacher added more sulfuric acid, the bulb began to glow again.

Answers to Questions

1. Given the sulfuric acid was 0.1 M, what was the initial concentration of the barium hydroxide solution?

   Moles of sulfuric acid = (0.1 M) x (0.02675 L) = 0.002675 moles sulfuric acid.
   At the equivalence point, moles of sulfuric acid equals moles of barium hydroxide.
   0.002675 moles of Ba(OH)₂/0.250 L = 0.0107 M Ba(OH)₂(aq) = 0.01 M Ba(OH)₂(aq)

2. When sulfuric acid is added to the initial barium hydroxide solution, what are the two balanced net-ionic equations that occur?
   {14098_Answers_Equation_6}
3. Why does the conductivity of the solution decrease when sulfuric acid is added?

   When sulfuric acid is added to the solution, two reactions occur. The sulfate reacts with the barium ions, forming a solid barium sulfate, which does not conduct electricity. And the hydronium and hydroxide ions react to form more water. Therefore, as sulfuric acid is added, the ions in solution decrease until all the aqueous barium hydroxide is reacted.

4. Given a K_sp of 1.0 x 10^–10 for BaSO₄, what is the concentration of Ba²⁺ at the equivalence point?

   [Ba²⁺(aq)[SO₄ ^2–(aq)] = Ksp = 1.0 x 10^–10
   [Ba²⁺(aq)] = [SO₄ ^2–(aq)]

   {14098_Answers_Equation_7}
5. When another 0.5 mL of 0.1M sulfuric acid is added past the equivalence point, the concentration of Ba²⁺(aq) in the solution decreases. Why?

   Adding more sulfate ions to a reaction in equilibrium causes a shift in the reaction and more barium sulfate is formed. Therefore, the Ba²⁺ concentration decreases.

   {14098_Answers_Equation_8}

Discussion

This demonstration combines acid–base chemistry, precipitation, and conductivity. With this titration, the equivalence point and endpoint are observed. The equivalence point of the titration is the point where stoichiometric amounts of the acid and base have combined. The endpoint of the titration is the point at which the indicator changes color and signals that the equivalence point has indeed been reached. Ideally, the equivalence point and endpoint should be close to each other. In this lab, phenolphthalein is used as the indicator.

Initially, in the demonstration, the solution contains only aqueous barium hydroxide which conducts electricity.

When the conductivity meter is placed in the solution of barium hydroxide, it lights up. Also, since the solution is basic, the solution turns pink when phenolphthalein is added.

As sulfuric acid is added, a white precipitate begins to form and the hydroxide ion is neutralized. Gradually, the amount of ions in solution decreases. As the solution gets closer to the equivalence point of the titration, the conductivity of the solution decreases and the bulb goes out. At the equivalence point, only water and barium sulfate remain and the solution is no longer pink. Barium sulfate has a very small K_sp and only very slightly dissolves. The amount of barium ions in solution at the equivalence point can be found with a simple calculation. In this case, the barium ion concentration is found to have a concentration of 1.0 x 10^–5 M at the equivalence point. If sulfuric acid is added after the equivalence point, the conductivity of the solution begins to increase again. When more sulfuric acid is added, the ion concentration increases and the solutions conductivity also increases.

References

AP^® Chemistry Course and Exam Description, page 142–144, accessed September 26, 2016. http://media.collegeboard.com/digitalServices/pdf/ap/ap-chemistry-course-and-exam-description.pdf

Shakhashiri, B. Z. Chemical Demonstrations: A Handbook for Teachers in Chemistry; University of Wisconsin: Madison, WI; 1983; Vol. 3, p. 152.