Materials Included In Kit

Additional Materials Required

Prelab Preparation

Sodium hydroxide solution, approximately 0.1 M
To prepare 1 liter of approximately 0.1 M NaOH:

1. Obtain a clean 250-mL graduated cylinder.
2. Pour 100 mL of the 1.0 M NaOH solution into the graduated cylinder.
3. Pour the 100 mL of 1.0 M NaOH into a clean 1-L Erlenmeyer flask.
4. Fill to 1-L mark with deionized water. Mix thoroughly.

Potassium hydrogen phthalate, KHC₈H₄O₄
Dry the solid for at least two hours in an oven at 110 °C. Store the dry solid in a desiccator. It must be cool when its mass is measured.

Safety Precautions

The 1 M sodium hydroxide solution is moderately toxic by ingestion and skin absorption. It is corrosive to body tissues and causes severe eye burns. Avoid all body contact. The dilute (0.1 M) sodium hydroxide solution is slightly toxic by ingestion and skin absorption and is irritating to skin and eyes. Phenolphthalein is an alcohol-based solution and is flammable. It is moderately toxic by ingestion. Keep it away from flames and other ignition sources. Avoid contact of all chemicals with eyes and skin and remind students to wash hands thoroughly with soap and water before leaving the laboratory. Wear chemical splash goggles and chemical-resistant gloves and apron. Consult current Safety Data Sheets for additional safety information.

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. The solid acids may be disposed of according to Flinn Suggested Disposal Method #24a. The sodium hydroxide solutions may be disposed of according to Flinn Suggested Disposal Method #10. The titrated solutions may be disposed of according to Flinn Suggested Disposal Method #26b. 

Lab Hints

• The actual experimental work for this lab takes at least two 45-minute periods. Review and demonstrate the proper titration techniques. The Flinn Scientific Laboratory Techniques Guide (Catalog No. AP6248) provides thumbnail illustrations of these and 14 other common laboratory techniques.
• This classic 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 a titration curve reinforces pH calculations and allows students to “see” what happens in a neutralization reaction.
• Solid weak acids or acid salts other than the included potassium hydrogen tartrate and potassium hydrogen sulfate can be used as the weak acid unknown. Table 1 lists some of these.

The non-hydrated acids may be oven dried to obtain a more precise equivalent mass. Do not dry the hydrated acids, and do not expect students to obtain four-digit precision with these. Do not exceed the melting point of the organic acids.

• Remind students to slide the dessicator lid off.
• If fewer than 12 pH meters are available, stagger the starting points for Parts B and C. Inexpensive pH meters, such as the Flinn pH meter (Catalog No. AP8673), can be used in Part C in place of the more expensive pH probe and meters.
• If magnetic stirrers are not available in Part C, have the students swirl the beaker after each addition of NaOH solution, then measure pH value.

Teacher Tips

• Quantitative analysis represents a nearly invisible application of chemistry in our daily lives. To illustrate the hidden importance of quantitative analysis, ask students how they would feel if they could not trust that the water they drink or the medicines they take had been tested to assure their quality and safety.

• Consider incorporating the titration curve experiment into a collaborative class project comparing and contrasting the titration behavior of strong versus weak acids. Randomly assign student groups one titration curve to perform. Students can then share their data in larger groups to analyze four classic titration curves: (1) strong acid/strong base; (2) weak acid/strong base; (3) strong base/strong acid; and (4) weak base/strong acid. Students should compare the initial pH of the solution, the pH at the equivalence point, and the pH at the midpoint of the titration curve (for weak acids and weak bases, respectively). This is an excellent culminating-type activity to assess student understanding of the principles of acid–base chemistry. See the Supplementary Information in the Discussion section for sample graphs of all four types (1–4) of titration curves.

Further Extensions

AP Chemistry Standards
This lab fulfills the requirement for the College Board–recommended AP Experiments #6: Standardization of a Solution Using a Primary Standard and #7: Determination of Concentration by Acid–Base Titration, Including a Weak Acid or a Weak Base. In addition, this lab provides the recommended familiarity with the process of titration using indicators and pH meters.

Answers to Prelab Questions

1. Calculate the equivalent mass of each of the following acids.

1. HC₂H₃O₂

   There is one ionizable hydrogen. EM = molar mass/1 = 60.05 g/H⁺

2. KHCO₃

   There is one ionizable hydrogen. EM = molar mass/1 = 100.12 g/H⁺

3. H₂SO₃

   There are two ionizable hydrogens. EM = molar mass/2 = 82.08/2 = 41.04 g/H⁺

2. Calculate the molarity of a solution of sodium hydroxide, NaOH, if 23.64 mL of this solution is needed to neutralize 0.5632 g of potassium hydrogen phthalate.

3. It is found that 24.68 mL of 0.1165 M NaOH is needed to titrate 0.2931 g of an unknown acid to the phenolphthalein end point. Calculate the equivalent mass of the acid.

4. The following values were collected for the titration of a 0.145 g of a weak acid with 0.100 M NaOH as the titrant:

1. Graph the data on the following chart.
   {13101_PreLab_Figure_5}
2. What is the pH at the equivalence point?

   At the equivalence point the pH is 8.5.

3. Give the K_a and pK_a value of the acid. Explain.

   The volume to reach the equivalence point is 25 mL. At half-neutralization, 12.5 mL, the pH = 4.8. Therefore the pK_a of the acid is 4.8 and the value of K_a = 10^–4.8 = 1.6 x 10^–5.

   {13101_PreLab_Figure_6}
4. The following acid–base indicators are available to follow the titration. Which of them would be most appropriate for signaling the endpoint of the titration? Explain.
   {13101_PreLab_Table_2}

   Thymol blue should be used since the equivalence point occurs at pH 8.5, which is within the color transition range for thymol blue. If bromphenol blue were used, the indicator would be gradually changing color through the first 20 mL of base added. If bromthymol blue were used, the color change would occur too early. It would also require several drops of the base to complete the color change.

Sample Data

Student data will vary.

Answers to Questions

1. From the standardization data, calculate the molarity of the sodium hydroxide solution for each trial. Average the values and enter the average in the Standardization Data Table.

2. From the equivalent mass data, calculate the equivalent mass of the unknown acid for each trial. Average the values and enter the average in the Equivalent Mass Data Table.

3. Why is equivalent mass determined and not molar mass?

Equivalent mass is the mass of an acid that provides one mole of hydrogen ions. It is the molar mass divided by the number of ionizable hydrogen ions. When an acid is titrated to a phenolphthalein end point, it is impossible to tell how many ionizable hydrogen ions there were. The equivalent mass, or the mass that provides one mole of hydrogen ions can be determined. If the number of ionizable hydrogen ions is known, multiplying this times the equivalent mass would give the molar mass.

4. Why must the KHP and the acid samples be dried? If they are not dried, how would the results change (high or low)?

If the KHP is moist when weighed out, a fictitiously high mass of KHP is measured, and the calculated molarity of the NaOH is too high.

5. Why must NaOH be standardized? Why can’t an exact solution of NaOH be prepared?

Because sodium hydroxide pellets rapidly absorb water from the air, the precise mass of NaOH cannot be measured on an analytical balance.

6. From the graph of pH versus volume of NaOH, determine the pK_a of the unknown acid. Convert this value to K_a.

Volume of 0.103 M sodium hydroxide required for neutralization = 19.2 mL
pH at half neutralization volume (from graph) = 4.2 = pK_a
K_a = 10^–4.2 = 10^0.8 x 10^–5 = 6.3 x 10^–5

7. Why is the equivalence point in the titration of the unknown acid with sodium hydroxide not at pH 7?

In the titration of a weak acid with a strong base, at the equivalence point the solution contains the anions of the weak acid and the cations of the strong base. Since the anions of the weak acid will hydrolyze to some extent, the solution will be slightly basic. The reaction is:

where A^– represents the anion of the weak acid.

Discussion

Supplemental Information

1. Titration of a strong acid (HCl) with a strong base (NaOH).

2. Titration of a weak acid (CH₃COOH) with a strong base (NaOH).

3. Titration of a strong base (NaOH) with a strong acid (HCl).

4. Titration of a weak base (NH₃) with a strong acid (HCl).

Introduction

A common question chemists have to answer is how much of something is present in a sample or a product. If the product contains an acid or base, this question is usually answered by a titration. Acid–base titrations can be used to measure the concentration of an acid or base in solution, to calculate the formula (molar) mass of an unknown acid or base and to determine the equilibrium constant of a weak acid (K_a) or of a weak base (K_b).

Concepts

• Weak acid

• Equilibrium constant (K_a)
• Titration curve
• Equivalent mass
• Equivalence point

Background

Titration is a method of volumetric analysis—the use of volume measurements to analyze an unknown. In acid–base chemistry, titration is most often used to analyze the amount of acid or base in a sample or solution. Consider a solution containing an unknown amount of hydrochloric acid. In a titration experiment, a known volume of the hydrochloric acid solution would be “titrated” by slowly adding dropwise a standard solution of a strong base such as sodium hydroxide. (A standard solution is one whose concentration is accurately known.) The titrant, sodium hydroxide in this case, reacts with and consumes the acid via a neutralization reaction (Equation 1). The exact volume of base needed to react completely with the acid is measured. This is called the equivalence point of the titration—the point at which stoichiometric amounts of the acid and base have combined.

Knowing the exact concentration and volume added of the titrant gives the number of moles of sodium hydroxide. The latter, in turn, is related by stoichiometry to the number of moles of hydrochloric acid initially present in the unknown.

Indicators are usually added to acid–base titrations to detect the equivalence point. The endpoint of the titration is the point at which the indicator changes color and signals that the equivalence point has indeed been reached. For example, in the case of the neutralization reaction shown in Equation 1, the pH of the solution would be acidic (< 7) before the equivalence point and basic (> 7) after the equivalence point. The pH at the equivalence point should be exactly 7, corresponding to the neutral products (sodium chloride and water). An indicator that changes color around pH 7 is therefore a suitable indicator for the titration of a strong acid with a strong base.

The progress of an acid–base titration can also be followed by measuring the pH of the solution being analyzed as a function of the volume of titrant added. A plot of the resulting data is called a pH curve or titration curve. Titration curves allow a precise determination of the equivalence point of the titration without the use of an indicator.

In this experiment the equivalent mass of an unknown acid will be determined by titration. The equivalent mass is defined as the mass of the acid that supplies one mole of hydrogen ions. The acid, a solid crystalline substance, is weighed out and titrated with a standard solution of sodium hydroxide. From the moles of base used and the mass of the acid, the equivalent mass of the acid is calculated. The acid is then titrated a second time with the standard solution of sodium hydroxide and the course of the titration is followed by using a pH meter. A plot is constructed with pH on the vertical y-axis and the volume of NaOH on the horizontal x-axis. From this graph the value of the equilibrium constant (K_a) for the dissociation of the acid is determined.

An acid may contain one or more ionizable hydrogen atoms in the molecule. The equivalent mass of an acid is the mass that provides one mole of ionizable hydrogen ions. It can be calculated from the molar mass divided by the number of ionizable hydrogen atoms in a molecule. For example, hydrochloric acid, HCl, contains one ionizable hydrogen atom—the molar mass is 36.45 g/mole, and its equivalent mass is also 36.45 g/mole. Sulfuric acid, H₂SO₄, contains 2 ionizable hydrogen atoms—the molar mass of H₂SO₄ is 98.07 g/mole but its equivalent mass is 49.04 g/mole. Thus, either 36.45 g of HCl or 49.04 g of H₂SO₄ would supply one mole of H⁺ ions when dissolved in water.

The equivalent mass is determined by titrating an acid with a standard solution of NaOH. Since one mole of NaOH reacts with one mole of hydrogen ion, at the equivalence point the following relation holds:

where V_b is the volume of base added at the endpoint, M_b is the molarity of base, grams acid is the mass of acid used, and EM_a is the equivalent mass of the acid.

The concentration of the NaOH solution must be accurately known. To “standardize” the NaOH, that is, to find its exact molarity, the NaOH is titrated against a solid acid, potassium hydrogen phthalate (abbreviated KHP). The KHP is chosen because it is easily dried and weighed and has a relatively high equivalent mass. The formula of KHP is:

KHP contains one ionizable H⁺. The titration is followed using phenolphthalein as an indicator.

The graph of pH versus volume of NaOH added (see Figure 1) is obtained by carefully following the titration with a pH meter. There is a significant change in pH in the vicinity of the equivalence point. Note that when a weak acid is titrated with a strong base, the equivalence point is NOT at pH 7, but is on the basic side. The value of the equilibrium constant for the dissociation of the acid is obtained from the graph.

If the dissociation of the acid is represented as:

the equilibrium constant expression is:

When the acid is half neutralized, [HA] = [A^–], these terms cancel in the above equation, and K_a = [H₃O⁺]. Therefore, when the acid is half-neutralized, the pH = pK_a.
The point where pH is equal to pK_a can be found from the graph. Refer to Figure 1.

A = Volume NaOH at equivalence point
B = ½A = the volume of NaOH required to neutralize
one-half the acid when half-neutralized
C = pH when the acid is half neutralized = pK_a

Experiment Overview

The purpose of this experiment is to standardize a sodium hydroxide solution and use the standard solution to titrate an unknown solid acid. The equivalent mass of the solid acid will be determined from the volume of sodium hydroxide added at the equivalence point. The equilibrium constant, K_a, of the solid acid will be calculated from the titration curve obtained by plotting the pH of the solution versus the volume of sodium hydroxide added.

Materials

Safety Precautions

Dilute sodium hydroxide solutions are irritating to skin and eyes. Phenolphthalein is an alcohol-based solution and is flammable. It is moderately toxic by ingestion. Keep away from flames and other ignition sources. Avoid contact of all chemicals with eyes and skin and wash hands thoroughly with soap and water before leaving the laboratory. Wear chemical splash goggles and chemical-resistant gloves and apron.

Procedure

Part A. Standardization of a Sodium Hydroxide Solution

1. Obtain a sample of potassium hydrogen phthalate (KHP) that has been previously dried in an oven and stored in a desiccator.
2. On an analytical balance, accurately weigh 0.4 to 0.6 grams of KHP in a previously tared weighing dish. Record the mass of the KHP in the Standardization Data Table.
3. Transfer the KHP into an Erlenmeyer flask—pour the solid through a funnel into the flask. Use water from a wash bottle to rinse all of the remaining solid in the weighing dish or in the funnel into the flask as well.
4. Add about 40 mL of distilled water to the flask and swirl until all the KHP is dissolved.
5. Obtain about 75 mL of the sodium hydroxide, NaOH, solution.
6. Clean a 50-mL buret, then rinse it with three small portions (about 7 mL each) of the NaOH solution.
7. Fill the buret to above the zero mark with the NaOH solution.

8. Open the buret stopcock to allow any air bubbles to escape from the tip. Close the stopcock when the liquid level is between the 0- and 10-mL marks.
9. Measure the precise volume of the solution in the buret and record this value in the Standardization Data Table as the “initial volume.” Note: Volumes are read from the top down in a buret. Always read from the bottom of the meniscus, remembering to include the appropriate number of significant figures (see Figure 3).
10. Position the buret over the Erlenmeyer flask so that the tip of the buret is within the flask but at least 2 cm above the liquid surface.
11. Add three drops of phenolphthalein solution to the KHP solution in the flask.
12. Begin the titration by adding 1.0 mL of the NaOH solution to the Erlenmeyer flask, then closing the buret stopcock and swirling the flask.
13. Repeat step 12 until 15 mL of the NaOH solution have been added to the flask. Be sure to continuously swirl the flask.
14. Reduce the incremental volumes of NaOH solution to ½ mL until the pink color starts to persist. Reduce the rate of addition of NaOH solution to drop by drop until the pink color persists for 15 seconds. Remember to constantly swirl the flask and to rinse the walls of the flask with distilled water before the endpoint is reached.
15. Measure the volume of NaOH remaining in the buret, estimating to the nearest 0.01 mL. Record this value as the “final volume” in the Standardization Data Table.
16. Repeat the standardization titration two more times. Rinse the Erlenmeyer flask thoroughly between trials with deionized water.

Part B. Determination of the Equivalent Mass of an Unknown Acid

1. Accurately weigh about 0.3–0.4 g of a sample of the unknown acid in a weighing dish using an analytical balance. Record the mass in the Equivalent Mass Data Table.
2. Dissolve the unknown acid in 40 mL of distilled water and titrate to the phenolphthalein end point as in Part A, steps 5 through 16.
3. Record the initial and final volumes of NaOH solution in the Equivalent Mass Data Table.
4. Repeat one more time. Choose a mass for the second sample so that the volume of NaOH needed is about 45 mL if using a 50-mL buret, or about 22 mL if using a 25-mL buret.

Part C. Determination of the pK_a of the Unknown Acid

1. Set up a pH meter and electrode. Calibrate the pH meter using a buffer solution of pH 7.00. Rinse the electrode well with distilled water.
2. On the analytical balance, weigh a sample of the unknown acid that requires approximately 20 mL of titrant.
3. Dissolve the sample in approximately 100 mL distilled water in a 250-mL beaker.

4. Fill the buret with the now standardized NaOH solution. Record the initital volume as the “initial buret reading” in the pK_a Data Table.
5. Set the beaker containing the unknown acid solution on a magnetic stirrer. 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.
6. When the pH reading has stabilized, record the initial pH of the solution in pK_a Data Table.
7. Add about 1 mL of sodium hydroxide solution to the beaker. Record the exact buret reading in pK_a Data Table.
8. Record the pH of the solution next to the buret reading in the pKa Data Table.
9. Add another 1-mL increment of sodium hydroxide solution. Record both the buret reading and the pH in pK_a Data Table.
10. Continue adding sodium hydroxide in 1-mL portions. Record both the buret reading and the pH after each addition.
11. 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.
12. Continue adding sodium hydroxide in about 0.2 mL increments. Record both the buret reading and the pH after each addition.
13. 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 and the pH after each addition.
14. Stop the titration when the pH of the solution is greater than 12. Record the final volume of solution in the buret and the final pH.
15. Graph the data, with pH on the vertical axis and volume NaOH on the horizontal axis. Make the graph large enough to reflect the care taken with the pH and volume measurements.
16. Dispose of the titrated solutions, the sodium hydroxide solution and any solid acid as directed by your instructor.

Student Worksheet PDF

13101_Student1.pdf