Materials Included In Kit

Additional Materials Required

Safety Precautions

All the acids and bases used in this lab are irritating to eyes, skin and other body tissues. The phenolphthalein solution is a flammable liquid, a fire risk and is moderately toxic. Methyl orange indicator solution is toxic by ingestion. The thymolphthalein solution is a flammable liquid and a fire risk. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Have students wash hands thoroughly with soap and water before leaving the lab. Please review current Safety Data Sheets for additional safety, handling and disposal information.

Disposal

Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulation that may apply, before proceeding. The sodium hydroxide solution and ammonia solution may both be disposed of according to Flinn Suggested Disposal Method #10. The hydrochloric acid solution may be disposed of according to Flinn Suggested Disposal Method #24b. The acetic acid solution may be disposed of according to Flinn Suggested Disposal Method #24a. The titrated solutions may be disposed of according to Flinn Suggested Disposal Method #26b.

Lab Hints

• The actual experimental work for this lab takes about two 50-minute periods. Review and demonstrate the proper pipetting and titration techniques. The Flinn Scientific Laboratory Techniques Guide (Catalog No. AP6248) provides thumbnail illustrations and 14 other common techniques.
• This titration lab teaches students how to use volumetric glassware and encourages them to develop good laboratory technique. Having students collect pH data and analyze the shape of the titration curve reinforces pH calculations and allows students to “see” what happens in a neutralization reaction.

Teacher Tips

• The student calculations of equivalence point pH and indicator selection are done in Parts 1 and 2 of the Prelab Calculations and Questions section. Question 3 provides an excellent assessment of the students’ understanding of the principles of acid–base chemistry.
• This lab is ideal for use with technology. Allows students to quickly gather data and graph.

Further Extensions

AP^® Standards
This lab fulfills the requirement for the College Board recommended AP Experiment #11: Determination of the Appropriate Indicators for Various Acid–Base Titrations. In addition, the lab provides the recommended familiarity with the process of titration using indicators and pH meters.

Answers to Prelab Questions

1. In Part 1 of this laboratory, 25.0 mL of a 0.10 M solution of the weak acid acetic acid, CH₃COOH, is titrated with a 0.10 M solution of the strong base sodium hydroxide, NaOH. K_a of acetic acid is 1.8 x 10^–5.
   1. Calculate the pH of the equivalence point. Enter this value in the Part 1 Data Table.

      At the equivalence point,
      moles of acetate ion (A^–) = initial moles of acetic acid

      {13904_Answers_Equation_16}
   2. Five indicators, along with their pK_a values and their color change ranges, are listed in the following table.
      {13904_PreLab_Table_1}

      Based on the pH calculated in 1a, select the appropriate indicator for the weak acid–strong base titration. Enter the selection in the Part 1 Data Table.

      The indicator with a pK_a about one unit above the equivalence point pH of 8.7 is phenolphthalein (9.4).

2. In Part 2, 25.0 mL of a 0.10 M solution of the weak base ammonia, NH₃, is titrated with a 0.10 M solution of the strong acid hydrochloric acid, HCl. K_b of ammonia is 1.8 x 10^–5.
   1. Calculate the pH of the equivalence point. Enter this value in the Part 1 Data Table.

      At the equivalence point,
      moles of ammonium ion (NH₄⁺) = initial moles of ammonia (NH₃)

      {13904_Answers_Equation_17}
   2. Select the appropriate indicator for the weak base–strong acid titration. Enter the selection in the Part 2 Data Table.

      The indicator with a pK_a value about one unit below the equivalence point pH of 5.3 is methyl red (4.95).

3. 25 mL of a 0.10 M solution of the weak acid hydrofluoric acid, HF, is titrated with a 0.10 M solution of the weak base ammonia, NH₃. Will the pH at the equivalence point be less than 7, equal to 7, or greater than 7? Explain. K_a for HF is 7.2 x 10^–4 and K_b for NH₃ is 1.8 x 10^–5.

   When a weak acid is titrated by a weak base, the equivalence point pH is dependent on the strengths of the conjugate base (A^–) of the weak acid (HA) and the conjugate acid (BH⁺) of the weak acid (B).
   
   If K_a of HA > K_b of B; then K_a of BH⁺ > K_b of A^–. At the equivalence point, BH⁺ produces more hydronium ions (H₃O⁺) in solution than A^– produces hydroxide ions (OH^–). The net result is a pH < 7 at the equivalence point.
   
   If K_a of HA < K_b of B; then the pH at the equivalenct point is > 7.
   
   K_a = 7.2 x 10^–4 for HF
   
   K_b = 1.8 x 10^–5 for NH₃
   
   Since K_a > K_b, the solution pH at the equivalence point will be < 7.

Sample Data

Part 1. Titration 0.10 M CH3COOH with 0.10 M NaOH

pH Equivalence Point (calc) ___8.7___ Selected Indicator ___Phenolphthalien___

Part 2. Titration 0.10 M NH3 with 0.10 M HCl

pH Equivalence Point (calc) ___5.3___ Selected Indicator ___Methyl Red___

Answers to Questions

Graph the pH versus the mL of titrant for each of the titrations. Make sure the graph is large enough to reflect the care taken with measuring the pH and volume. Draw the best fitting smooth curve for the data. Label the equivalence point. Indicate the indicator color for each data point. Were the indicators selected appropriate for the two titrations? If not, why?

Introduction

Acids and bases vary in strengths and are normally classified as strong or weak. In any acid–base titration the neutralization, or equivalence point, occurs when the moles of acid in solution are equal to the moles of base. However, the pH of the solution at this point can vary widely and depends on the strengths of both the acid and the base. How is an indicator selected that fits a particular acid–base titration?

Concepts

• Weak acid
• Equivalence point
• pH indicator
• Weak base
• Acid–base titration
• Conjugate acid–base pairs

Background

In acid–base titrations, the plot of pH versus volume of titrant results in an S-shaped curve (see Figure 1).

The steepness of the curve and the pH value at the equivalence point depend on the strength of both the acid and the base. If both the acid and base are strong, the curve is very steep and the equivalence point pH value is 7.

If a weak acid is titrated by a strong base, the titration curve is less steep and the equivalence point pH value is >7. At the equivalence point, The overall neutralization reaction is, At this point, the initial moles of the weak acid (HA) have been completely converted to its conjugate base (A^–). This conjugate base is a weak base and equilibrates with water to form a basic solution. K_b, for this reaction, is where K_w is 1 x 10^–14 and K_a is the dissociation constant of the weak acid (HA).

The pH at the equivalence point is found by first calculating the pOH, or –log[OH], of this solution of the weak base A^– and water.

The initial concentration of A^–, before its reaction with water, is equal to the initial moles of weak acid, HA, present in the solution, divided by the volume of solution at the equivalence point.

When the weak base A^– reacts with water, at equilibrium,

[HA] = [OH] = x If [A^–]>>[HA], substituting these values into Equation 3 yields

Since pOH = –log[OH–] and pH + pOH = 14.00
pH = 14.00 – pOH

Once this pH value is determined, an appropriate indicator can be selected for the titration. Indicators are mostly complex organic molecules that are themselves weak acids. If the indicator is represented by HIn, then in solution, The HIn form has one color in solution and the In^– form has another. If Equation 3 is rearranged, then As base is added in the titration, H₃O⁺ ions are removed and the equilibrium shifts right, forming more In^– ions. A color change starts to occur when [In^–] is about one-tenth [HIn].

At this point, In terms of pH For a specific titration of an acid by a base, an indicator is selected that has a pK_a one unit above the pH value of the equivalence point. The color transition of the indicator is complete when Thus the transition range for most indicators is two pH units, or pK_a ±1.

When a weak base is titrated with a strong acid, all the weak base (B) is converted to its conjugate acid (BH⁺) at the equivalence point. The BH⁺ produced equilibrates with water to form an acidic solution. At equilibrium of the weak acid BH⁺ and water, [B] = x = [H₃O⁺] if, at equilibrium, [BH⁺]>>[B]; then As K_a increases, the pH at the equivalence point decreases. Relating this to the K_b values, the weaker the base, the lower the pH at the equivalence point.

In basic solution, the indicators initially exists in the In^– form. Now the color transition occurs when the HIn concentration is one-tenth the In^– concentration, or

Experiment Overview

The appropriate indicators are selected for two titrations—a weak acid solution titrated with a strong base solution and a weak base solution titrated with a strong acid solution. The indicators are added to the solutions and the solutions are titrated. Titration curves of pH versus volume of titrant are generated and used to verify the appropriateness of the selected indicators.

Materials

Prelab Questions

1. In Part 1 of this laboratory, 25.0 mL of a 0.100 M solution of the weak acid acetic acid, CH₃COOH, is titrated with a 0.100 M solution of the strong base sodium hydroxide, NaOH. K_a of acetic acid is 1.8 x 10^–5.
   1. Calculate the pH of the equivalence point. Enter this value in the Part 1 Data Table.
   2. Five indicators, along with their pK_a values and their color change ranges, are listed in the following table.
      {13904_PreLab_Table_1}
      Based on the pH calculated in 1a, select the appropriate indicator for the weak acid–strong base titration. Enter the selection in the Part 1 Data Table.
2. In Part 2, 25.0 mL of a 0.10 M solution of the weak base ammonia, NH₃, is titrated with a 0.10-M solution of the strong acid hydrochloric acid, HCl. K_b of ammonia is 1.8 x 10^–5.
   1. Calculate the pH of the equivalence point. Enter this value in the Part 1 Data Table.
   2. Select the appropriate indicator for the weak base–strong acid titration. Enter the selection in the Part 2 Data Table.
3. 25 mL of a 0.10 M solution of the weak acid hydrofluoric acid, HF, is titrated with a 0.10 M solution of the weak base ammonia, NH₃. Will the pH at the equivalence point be less than 7, equal to 7, or greater than 7? Explain. K_a for HF is 7.2 x 10^–4 and K_b for NH₃ is 1.8 x 10^–5.

Safety Precautions

All the acids and bases used in this lab are irritating to eyes, skin and other body tissues. The phenolphthalein solution is a flammable liquid, a fire risk and is moderately toxic. Methyl orange indicator solution is toxic by ingestion. The thymolphthalein solution is a flammable liquid and a fire risk. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the lab. Clean up all spills immediately. Please review current Safety Data Sheets for additional safety, handling and disposal information.

Procedure

Tips

1. Label all solutions.
2. Practice pipetting technique.
3. Always rinse glassware with distilled or deionized water.

Part 1. Titration of a Weak Acid with a Strong Base

1. Place about 75 mL of 0.10 M acetic acid solution in a clean 250-mL beaker.
2. Using a clean 25-mL volumetric pipet, quantitatively transfer 25.0 mL of 0.10 M acetic acid solution to a clean 150-mL beaker (see Figure 2).
   {13904_Procedure_Figure_2_Using a volumetric pipet}
3. Obtain about 100 mL of the 0.10 M NaOH solution in a clean, 250-mL beaker.
4. Clean a 50-mL buret thoroughly with tap water, then rinse it with several small portions of the standard NaOH solution, being sure to run some solution through the tip.
5. Attach the buret clamp to the ring stand. Place the buret in the buret clamp.
6. Fill the buret above the 0-mL mark, then lower the meniscus back to the 0-mL mark.
7. Set up a pH meter and electrode on the ring stand. Calibrate the pH meter using a buffer solution of pH 7.00. Rinse the electrode well with distilled water.
8. Set the beaker containing the acetic acid solution on a magnetic stirrer so that the buret tip is just below the lip of the beaker. Clamp the pH electrode so it is submerged in the acid solution (see Figure 3). Be sure the stir bar does not hit the electrode. Set the stir bar gently spinning.
   {13904_Procedure_Figure_3_Setup}
9. When the pH reading has stabilized, record the initial pH of the solution in Part 1 Data Table.
10. Add 3 drops of the indicator selected for this titration in the Prelab Calculations section. Record the solution color in the Part 1 Data Table.
11. Add about 1 mL of sodium hydroxide solution to the beaker. Record the exact buret reading in Part 1 Data Table.
12. Record the pH and the color of the solution next to the buret reading in the Part 1 Data Table.
13. Add another 1-mL increment of sodium hydroxide solution. Record the buret reading, the pH, and the solution color in Part 1 Data Table.
14. Continue adding sodium hydroxide in 1-mL portions. Record the buret reading, the pH, and the solution color after each addition.
15. When the pH begins to increase by more than 0.3 pH units after an addition, decrease the portions of sodium hydroxide added to about 0.2 mL.
16. Continue adding sodium hydroxide in about 0.2 mL increments. Record the buret reading, the pH, and the solution color after each addition.
17. When the pH change is again about 0.3 pH units, resume adding the sodium hydroxide in 1-mL increments. Continue to record both the buret reading, the pH, and the solution color after each addition.
18. Stop the titration when the pH of the solution is greater than 12. Record the final volume of solution in the buret, final pH, and solution color.
19. Repeat the titration. If the indicator selected did not change color at the equivalence point in the first titration, review the Prelab Calculations and select another indicator.

Part 2. Titration of a Weak Base with a Strong Acid

1. Place about 75 mL of 0.10 M ammonia (NH₃) solution in a clean 250-mL beaker.
2. Using a clean 25-mL volumetric pipet, quantitatively transfer 25.0 mL of 0.10 M NH₃ solution to a clean 150-mL beaker.
3. Obtain about 100 mL of 0.10 M hydrochloric acid (HCl) solution in a clean, 250-mL beaker.
4. Rinse the 50-mL buret with three small portions of deionized water, then rinse it with several small portions of 0.10 M HCl solution. Place the buret back in the buret clamp.
5. Fill the buret with 0.10 M HCl solution above the 0-mL mark, then lower the meniscus back to zero.
6. Set the beaker containing the 0.10 M NH₃ solution on a magnetic stirrer. Clamp the pH electrode so it is submerged in the basic solution. Be sure the stir bar does not hit the electrode. Set the stir bar gently spinning.
7. When the pH reading has stabilized, record the initial pH of the solution in Part 2 Data Table.
8. Add three drops of the indicator selected for this titration in the Prelab Calculations section. Record the solution color in the Part 2 Data Table.
9. Add about 1 mL of 0.1 M HCl solution to the beaker. Record the exact buret reading in Part 2 Data Table.
10. Record the pH of the solution and the solution color next to the buret reading in the Part 2 Data Table.
11. Add another 1-mL increment of hydrochloric acid solution. Record the buret reading, the pH, and the solution color in Part 2 Data Table.
12. Continue adding hydrochloric acid in 1-mL portions. Record both the buret reading, the pH, and the solution color after each addition.
13. When the pH begins to decrease by more than 0.3 pH units after an addition, decrease the portions of 0.10 M HCl added to about 0.2 mL.
14. Continue adding 0.10 M HCl in about 0.2 mL increments. Record the buret reading, the pH, and the solution color after each addition.
15. When the pH change is again about 0.3 pH units, resume adding the 0.10 M HCl in 1-mL increments. Continue to record the buret reading, the pH, and the solution color after each addition.
16. Stop the titration when the pH of the solution is less than 2. Record the final volume of solution in the buret, the final pH, and the final solution color.
17. Repeat the titration. If the indicator selected did not change color at the equivalence point in the first titration, review the Prelab Calculations and select another indicator.
18. Dispose of all solutions as directed by your instructor.

Student Worksheet PDF

13904_Student1.pdf