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AP Chemistry Lab 17 Determination of the Dissociation Constant of Weak Acids PRELIMINARY LAB ASSIGNMENT For oxalic acid, H2C2O4 the values for the acid dissociation constants are: Kal = 6.5 X 10-2 Ka2 = 1.5 X 10-13 1. Write the chemical equation for the first dissociation of oxalic acid in water.
2. Write the Ka1 expression for the reaction in question 1.
3. What would be the pH of a solution in question 1 be when [H2C2O4] = [HC2O4-]?
4. Write the chemical equation for the second dissociation of oxalic acid in water.
5. Write the Ka2 expression for the reaction in question 4.
6. What would be the pH of a solution in question 4 be when [HC2O4-] = [C2O4-2]?
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AP Chemistry Lab 17 Determination of the Dissociation Constant of Weak Acids
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INTRODUCTION When a weak acid is dissolved in water, it breaks apart or dissociates to a slight extent. A proton from the acid is donated to a water molecule. The equations for the equilibrium and the equilibrium constant expression are as follows: HA + H2O H3O+ + A+ Ka = [H3O ] [A ]
[HA]
where A- represents the anion of the weak acid and the square brackets indicate molar concentrations of the species. For most weak acids the percent of acid that dissociates is less than 5%. The value of the equilibrium constant, Ka, indicates to what extent the reaction occurs. The greater the value of K a, the stronger the acid, and the greater the amount of dissociation. Polyprotic acids contain more than one ionizable hydrogen. The dissociation process occurs stepwise, and there is an equilibrium constant for each of the steps. The second reaction always occurs to a much smaller extent than the first, so K a2 is always a smaller value than Kal. H2A + H2O H3O+ + HAHA- + H2O H3O+ + A2+ Ka1 = [H3O ] [HA ]
[H2A]
+ 2Ka2 = [H3O ]- [A ]
[HA ]
Some values for Ka and pKa (pKa = -logKa) that cover a wide range of acid strengths are listed in Table 1. Table 1. Acid Dissociation Constants Acid Formula Iodic HIO3 Benzoic HC7H5O2 Acetic HC2H3O2 Potassium hydrogen phthalate KHC8H4O4 Carbonic H2CO3 Sulfurous H2SO3 Sodium hydrogen sulfate NaHSO4 Sodium hydrogen sulfite NaHSO3 Potassium dihydrogen phosphate KH2PO4 Sodium bitartrate NaHC4H4O6H2O Hypochlorous HClO Hydrocyanic HCN
Ka1 0.17 6.5 x 10-5 1.8 x 10-5 3.9 x 10-6 4.3 x 10-7 1.7 x 10-2 1.2 x 10-2 6.4 x 10-8 6.2 x 10-8 4.6 x 10-5 3.0 x 10-8 4.9 x 10-10
Ka2
5.6 x 10-11 6.4 x 10-8
pKa1 0.77 4.19 4.74 5.41 6.37 1.77 1.92 7.19 7.21 4.34 7.52 9.31
pKa2
10.25 7.19
This experiment is designed to determine the Ka and pKa values of a number of weak acids. Acetic acid, HC2H3O2, will be used as an example for the experimental procedure. When acetic acid is dissolved in water, an equilibrium exists in which a mixture of acetic acid, hydronium ions, and acetate ions will all be present: HC2H3O2 + H2O H3O+ + C2H3O2-
AP Chemistry Lab 17 Determination of the Dissociation Constant of Weak Acids
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Acetic acid and acetate ions are conjugate acid-base pairs. A conjugate acid is a substance that has one more proton in its structure than its corresponding conjugate base. This combination also results from a mixture of a weak acid, acetic acid, and its salt, sodium acetate. The equilibrium constant expression is: + Ka = [H3O ] [C2H3O2 ] = 1.8 x 10-5 [HC2H3O2]
If a solution contains equal concentrations of HC2H3O2 and C2H3O2-, these concentration terms cancel out in the above equation so that Ka = [H3O+] = 1.8 x 10-5, and pH = pKa = 4.74. You will prepare solutions in which the concentrations of acid and its anion are equal. The value of the pH of the solution will then equal the pKa for the acid. Some of the substances tested will be salts of diprotic acids that still contain one ionizable hydrogen. For example, NaHSO4 ionizes in solution forming Na+ and HSO4-. The HSO4- then reacts with water in the equilibrium: HSO4- + H2O H3O+ + SO42The value of Ka that is found when equal concentrations of HSO4- and SO42- are in solution is Ka2 for sulfuric acid.
DEFINITIONS: pH, acid strength, 1st acid dissociation constant, 2nd acid dissociation constant, polyprotic acid.
MATERIALS several weak acids 0.1 M NaOH 0.1 M phenolphthalein solution pH indicator paper and/or pH meter balance
SAFETY 1. Acids and bases are harmful to skin and eyes. Wash spills off skin with lots of water. Neutralize acid spills on the table with baking soda. 2. Phenolphthalein is dissolved in alcohol, so it is flammable. Keep the solution away from flames. 3. Wear goggles.
AP Chemistry Lab 17 Determination of the Dissociation Constant of Weak Acids
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PROCEDURE 1. Measure out about 0.2 g of the weak acid to be tested. For liquid acids, use about four drops. It is not necessary to know the exact amount. 2. Measure precisely 50.0 mL of distilled water into a beaker, add the weak acid, stir to dissolve and mix well. If the acid does not readily dissolve, warm gently to dissolve, then cool to room temperature. 3. Pour precisely 25.0 mL of the weak acid solution into an Erlenmeyer flask. Add 2 drops of phenolphthalein solution to the acid solution in the flask, and then add NaOH solution dropwise while swirling the flask. Stop adding the NaOH when the first pink color persists throughout the solution. This process converts all of the weak acid, HA, in the flask into its conjugate base, A-, according to the neutralization reaction OH- + HA H2O + A-
(one way arrow)
At this point the beaker contains exactly one-half of the original acid, essentially all of which is in the undissociated form, HA, and the flask contains an equal amount of the anion of the weak acid. 4. Pour the contents of the flask into the beaker and mix the solution. 5. Measure the pH of this solution using both a pH meter and pH indicator paper. The pH is the pKa of the acid. Calculate the value of Ka of the acid. 6. The solutions may be washed down the drain with an excess of water. ANALYSIS 1. Prepare a table of results, showing the pH determined with a pH meter and with pH paper for each acid. Show the calculated Ka for both determinations for each acid. 2. Assume the acid dissociation constant for the acid salt NaHSO4 is to be determined. a. Write the chemical equation to show this salt ionizing in water. b. Write the chemical equation showing the anion acting as an acid in water. c. Write the equilibrium expression for the acid dissociation of the anion. d. Explain how to determine the acid dissociation constant using the expression from c and the experimental procedure used in this lab. 3. Why is it not necessary to know the exact mass of the acid whose Ka is to be determined? 4. Why is it not necessary to know the exact concentration of the NaOH solution used? 5. Why is it necessary to precisely measure the volume of distilled water used to dissolve the acid? EXPERIMENTAL ERROR Calculate the percent error of your Ka (compare to those in Table 1). Which method, pH meter or pH paper, gave better results? What is the effect on the value of Ka if an extra drop of NaOH solution is added past the end point in step 3?