Teacher Notes

Determination of Ka of Weak Acids

Classic Chemistry Experiment

Materials Included In Kit

Acetylsalicylic acid, 2-CH3CO2C6H4COOH, 3 g*
Phenolphthalein solution, 0.5%, 30 mL
Potassium dihydrogen phosphate (potassium phosphate, monobasic), KH2PO4, 3 g*
Potassium hydrogen phthalate, KHC8H4O4, 3 g*
Potassium hydrogen sulfate (potassium bisulfate), KHSO4, 3 g*
Potassium hydrogen tartrate (potassium bitartrate), KHC4H4O6, 3 g*
Sodium hydroxide, NaOH, 0.1 M, 250 mL
*Unknown weak acid salts

Additional Materials Required

Buffer solutions, pH 4 and 7, 100 mL each*
Water, distilled or deionized
Balances, 0.01-g precision, 3
Beakers, 150-mL, 12
Erlenmeyer flasks, 125-mL, 12
Graduated cylinders, 100-mL, 12
pH Meters, 12
Pipets, Beral-type, 24
Stirring rods, 12
Wash bottles, 12
Weighing dishes, 24
*Calibrate the pH meters before use.

Safety Precautions

Acids and bases are skin and eye irritants. Avoid contact of all chemicals with eyes and skin. Inform the teacher and clean up all acid and base spills immediately. Phenolphthalein is an alcohol-based solution and is flammable. It is also a possible carcinogen. Wear chemical splash goggles and chemical-resistant gloves and apron. Remind students to wash hands thoroughly with soap and water before leaving the laboratory. Keep sodium bicarbonate and citric acid on hand to clean up acid and base spills, respectively. Please review current Safety Data Sheets for additional safety, handling and disposal information.


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. All solutions, except phenolphthalein, may be disposed of according to Flinn Suggested Disposal Method #26b. Phenolphthalein solution may be disposed of according to Flinn Suggested Disposal Method #18a.

Lab Hints

  • The experiment has been written with the intention of having students identify unknowns from a list of possibilities. Alternatively, the identities of the samples may be revealed to students and students may be asked to calculate Ka values and percent errors. The logarithmic scale of pKa values tends to obscure fairly large percent errors in Ka determinations. Thus, the experimentally determined pKa value of KH2PO4 (7.0) compares favorably with the literature value (7.2). The percent error in the corresponding Ka values—1.0 x 10–7 and 6.2 x 10–8 for the experimental and literature values, respectively—is quite large (38%).
  • For best results, it is necessary that all solids be completely dissolved before performing the “half-neutralization” reactions with sodium hydroxide and that the solutions be thoroughly mixed at every stage. When the sodium hydroxide is added to neutralize the acid, the flask should be constantly swirled to mix the solutions. The addition of sodium hydroxide should stop when the phenolphthalein indicator just remains pink throughout the solution. Remind students to be careful not to overshoot the endpoint. If they overshoot, have them start the determination over with a new sample.
  • The number of drops of sodium hydroxide required for “half-neutralization” of the unknowns varies from about 75 to 150 drops, depending on the the mass of salt used and the molar mass of the unknown. For best results, keep the mass of salt used to less than 0.20 g.
  • Technically, the pKa of an acid may be determined by measuring the pH value in any solution of a weak acid of known concentration if the concentration of its conjugate base is also known. Many factors, especially dissolved carbon dioxide, may interfere with the accuracy of measurements. Under certain conditions the pH of a weak acid may be very sensitive to small changes in the concentrations of the weak acid and conjugate base. In practical terms, therefore, this procedure is most convenient when the solution contains equal molar amounts of the weak acid and its conjugate base—that is, when the solution is an ideal buffer.
  • A buffer is any solution that contains appreciable amounts of both HA and A. By definition, the pH of a buffer is relatively insensitive to the addition of small amounts of strong acids and bases. The properties of a buffer are easily understood by looking at the titration curve (graph of pH versus equivalents of base added) for a weak acid. The titration curve is relatively flat in the “buffer region” (around the midpoint in the titration curve) corresponding to half-neutralization of the weak acid.
  • Although the procedure for this experiment is straightforward, the concepts are challenging. Students must be able to write equilibrium constant expressions for acids and understand the mathematical relationship between pH measurements and H3O+ concentrations. They should also be familiar with the properties of diprotic acids and the nature of an “acid salt” of a diprotic acid.
  • For best results, calibrate the pH meters using standard pH 4 and 7 buffer solutions before use.

Answers to Prelab Questions

Phosphoric acid is a triprotic acid (three ionizable hydrogens). The values of its stepwise ionization constants are Ka1 = 7.5 x 10–3, Ka2 = 6.2 x 10–8 and Ka3 = 4.2 x 0–13.

  1. Write the chemical equation for the first ionization of phosphoric acid with water. 
  2. Write the equilibrium constant expression (Ka1) for this reaction.
  3. What would be the pH of a solution when [H3PO4] = [H2PO4]? Note: pH = –log[H3O+].

    When [H3PO4] = [H2PO4], the hydrogen ion concentration is equal to Ka1.
    [H3O+] =
    Ka1 = 7.5 x 10–3
    pH = –log[H3O+] = –log(7.5 x 10–3) = –(–2.12) = 2.12
    Note: Unless rigorous precautions are taken, pH measurements are only precise to one decimal place. In terms of significant figures, however, both decimal places are allowed in the pH calculation, because the first digit (2.12) corresponds to the exponent.

  4. Phenolphthalein would not be an appropriate indicator to use to determine Ka1 for phosphoric acid. Why not? Choose a suitable indicator from the following color chart.
    The color change for phenolphthalein occurs between 8 and 10. This is significantly higher than the pH value for neutralization of only the first acidic hydrogen in H3PO4. At a pH of 10, in fact, almost 100% of the phosphoric acid initially present would be converted to HPO42–.
    Note: In general, for the conversion of an acid HA to its conjugate base A, 99% of the compound will exist in the form A when the pH is 2 units higher than the p
    Ka value. Thus, in order to obtain complete conversion of H3PO4 to H2PO4, and avoid conversion of H2PO4 to HPO42–, a pH range between 4 and 6 would be optimal. The orange transition color of methyl red would be a suitable indicator.

Sample Data


Answers to Questions

  1. Average the pH readings for each trial (samples 1 and 2) to calculate the average pKa value for your unknown weak acids. Enter the results in the data table.

    All results have been rounded to one decimal place for pKa values.

  2. Comment on the precision (reproducibility) of the pKa determinations. Describe sources of experimental error and their likely effect on the measured pKa (pH) values.

    The reproducibility of pH measurements is excellent. The pH at the “half-neutralization” point is relatively constant because it depends only on the ratio of the acid and its conjugate base. Small errors in the concentrations of the acid and its conjugate base produce only small changes in this ratio. The observed variation in pH measurements appears to be within the limits of precision of the pH meter. Note: The pKa value of an acid depends on temperature. The literature pKa values are given for a temperature of 25 °C.

  3. The following table lists the identities of the possible unknowns in this experiment. Complete the table by calculating the pKa value for each acid. Note: pKa = –logKa.
  4. Compare the experimental pKa value for each unknown with the literature values reported in Question 3. Determine the probable identity of each unknown and enter answers in the data table.
  5. Write separate equations for each unknown potassium salt dissolving in water and for the ionization reaction of the weak acid anion that each of these salts contains.

    Sample equations are shown for potassium dihydrogen phosphate:


    H2PO4(aq) + H2O(l) → HPO42–(aq) + H3O+(aq)

  6. Why was it not necessary to know the exact mass of each acid sample?

    It is not necessary to know the exact mass of each acid as long as we know that half of it was neutralized. Since the [HA] and [A] terms cancel out in the equilibrium constant expression, the exact amounts are not important.

  7. Why was it not necessary to know the exact concentration of the sodium hydroxide solution?

    Again, the only variable that must be controlled is that half of the acid has been neutralized. It does not matter how much sodium hydroxide is added to neutralize the sample.

  8. Why was it necessary to measure the exact volume of distilled water used to dissolve the acid as well as the exact volume of solution transferred from the beaker to the Erlenmeyer flask?

    The important variable is that half of the acid must be neutralized. Exactly half of the initial volume must be transferred to the Erlenmeyer flask for neutralization. Both the initial volume and transfer volume must therefore be accurately and precisely known.

Student Pages

Determination of Ka of Weak Acids


Acids vary greatly in their strength—their ability to ionize or produce ions when dissolved in water. What factors determine the strength of an acid? In this experiment, the strength of acids will be measured by determining the equilibrium constants for their ionization reactions in water.


  • Weak acid
  • Equilibrium constant
  • Conjugate base
  • Neutralization reaction


The modern Brønsted definition of an acid relies on the ability of the compound to donate hydrogen ions to other substances. When an acid dissolves in water, it donates hydrogen ions to water molecules to form H3O+ ions. The general form of this reaction, called an ionization reaction, is shown in Equation 1, where HA is the acid and A its conjugate base after loss of a hydrogen ion. The double arrows represent a reversible reaction.

The equilibrium constant expression (Ka) for the reversible ionization of an acid is given in Equation 2. The square brackets refer to the molar concentrations of the reactants and products.
Not all acids, of course, are created equal. The strength of an acid depends on the value of its equilibrium constant Ka for Equation 1. Strong acids ionize completely in aqueous solution. The value of Ka for a strong acid is extremely large and Equation 1 essentially goes to completion—only H3O+ and A are present in solution. Weak acids, in contrast, ionize only partially in aqueous solution. The value of Ka for a weak acid is much less than one and Equation 1 is reversible—all species (HA, A and H3O+) are present at equilibrium.

Polyprotic acids contain more than one ionizable hydrogen. Ionization of a polyprotic acid occurs in a stepwise manner, where each step is characterized by its own equilibrium constant (e.g., Ka1, Ka2). The second reaction (removal of the second acidic hydrogen) always occurs to a much smaller extent than the first reaction, and so Ka2 is always significantly smaller than Ka1. Sulfuric acid (H2SO4) and phosphoric acid (H3PO4) are examples of polyprotic acids.
The ionization constant of a weak acid can be determined experimentally by measuring the H3O+ concentration in a dilute aqueous solution of the weak acid. This procedure is most accurate when the solution contains equal molar amounts of the weak acid and its conjugate base. Consider acetic acid as an example. Acetic acid (CH3COOH) and the acetate anion (CH3COO) represent a conjugate acid–base pair. The equilibrium constant expression for ionization of acetic acid is shown in Equation 5. If the concentrations of acetic acid and acetate ion are equal, then these two terms cancel out in the equilibrium constant expression, and Equation 6 reduces to Equation 7.
In this experiment, solutions are prepared in which the molar concentrations of an unknown acid and its conjugate base are equal. The pH of these solutions are then equal to the pKa for the acid. The definition of pKa is closely related to that of pH. Thus, pH = –log[H3O+] and pKa = –logKa. The substances tested are salts of polyprotic acids that still contain an ionizable hydrogen. Sodium bisulfate (NaHSO4), for example, is a weak acid salt; it contains Na+ and HSO4 ions. The HSO4 ion is a weak acid—the equilibrium constant for ionization of HSO4 corresponds to Ka2 for sulfuric acid.

Experiment Overview

The purpose of this experiment is to measure the pKa value for ionization of two unknown weak acids. Solutions containing equal molar amounts of the weak acids and their conjugate bases are prepared by “half-neutralization” of the acid. Their pH values are measured and used to calculate the pKa value for the unknowns and thus determine their identities.


Phenolphthalein solution, 0.5%, 1 mL
Sodium hydroxide solution, NaOH, 0.1 M, 15 mL
Unknown weak acids, A–E, about 0.5 g each
Water, distilled or deionized
Balance, 0.01-g precision
Beaker, 150-mL
Erlenmeyer flask, 125-mL
Graduated cylinder, 50- or 100-mL
pH Meter
Pipets, Beral-type, 2
Stirring rod
Wash bottle
Weighing dishes, 2

Prelab Questions

See Student PDF.

Safety Precautions

Acids and bases are skin and eye irritants. Avoid contact of all chemicals with eyes and skin. Inform the teacher and clean up all acid and base spills immediately. Phenolphthalein is an alcohol-based solution and is flammable. It is also a possible carcinogen. Keep the solution away from flames. Wear chemical splash goggles and chemical-resistant gloves and apron. Wash hands thoroughly with soap and water before leaving the laboratory.


  1. Label two weighing dishes 1 and 2.
  2. Obtain an unknown weak acid and record the unknown letter in the data table.
  3. Measure out a small quantity (0.15–0.20 g) of the unknown into each weighing dish. Note: It is not necessary to know the exact mass of each sample.
  4. Using a graduated cylinder, precisely measure 50.0 mL of distilled water into a 150-mL beaker.
  5. Transfer sample 1 to the water in the beaker and stir to dissolve.
  6. Using a graduated cylinder, precisely transfer 25.0 mL of the acid solution prepared in step 5 to an Erlenmeyer flask.
  7. Add 3 drops of phenolphthalein solution to the acid solution in the Erlenmeyer flask.
  8. Using a Beral-type pipet, add sodium hydroxide solution dropwise to the flask. Gently swirl the flask while adding the sodium hydroxide solution.
  9. Continue adding sodium hydroxide dropwise and swirling the solution until a faint pink color persists throughout the solution for at least 5 seconds. This is called the endpoint. Note: A pink color develops immediately when the base is added, but fades quickly once the solution is swirled. When nearing the endpoint, the pink color begins to fade more slowly. Proceed cautiously when nearing the endpoint, so as not to “overshoot” it.
    Note: At this point the solution in the beaker contains exactly one-half of the original amount of acid, essentially all of which is in the acid form, HA. The flask contains an equal amount of the conjugate base A obtained by neutralization.
  10. Pour the contents of the Erlenmeyer flask back into the beaker. Pour the solution back and forth a few times to mix. Note: It is important to transfer the solution as completely as possible from the flask back into the beaker.
  11. Using a pH meter, measure the pH of the resulting solution in the beaker, which contains equal molar amounts of the acid and its conjugate base. Record the pH in the data table.
  12. Dispose of the beaker contents according to the teacher’s instructions and rinse both the beaker and the Erlenmeyer flask with distilled water. Dry the beaker with a paper towel.
  13. Repeat steps 4–12 using sample 2.
  14. Repeat steps 1–13 for one of the remaining unknowns.

Student Worksheet PDF


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