Exp 6 Determination of the Ionization Constant of a Weak Acid-2 PDF

Preview

Summary

Inorganic Chem Lab 102 assignments from Dr. Patricia Smith...

Description

CHEM-102

Experiment #6

Determination of the Ionization Constant of a Weak Acid

Pre-Lab Activities: 

After reading the lab, complete items a, b, c, and d (title, purpose, chemicals and equipment, and summary of procedure) as described on page10 of Exp. 1 on an 8 1/2 x 11 sheet of paper.



Bring a USB drive to store your lab data



Answer the following questions on 8½ x 11 sheet of paper or in your laboratory notebook if one is required by your instructor: 1. What are the largest sources of error in this experiment? 2. If 30.15 mL of 0.0995 M NaOH is required to neutralize 0.180 g of an unknown acid, HA, what is the molecular weight of the unknown acid?

Lab Activities: 

Go over the prelab questions with your lab instructor



Complete lab and fill in data sheet for the Standardization of NaOH (part A). Parts B-D will be completed the following lab period.

Objective: 

To experimentally determine the ionization constant (Ka) and molecular weight (MW) of an unknown weak acid



To identify the unknown acid

Introduction: NaOH stock solution is standardized by titration against primary standard potassium hydrogen phthalate, KHC8H4O4. The concentration and Ka of the unknown acid is then determined by titrating the acid against the standardized NaOH using a pH probe and a computer program. The unknown acid may be identified from a list of possibilities using both the Ka and molecular weight found experimentally. 1|Page

According to the Bronsted-Lowry acid-base theory, the strength of an acid is related to its ability to donate protons. All acid-base reactions are then competitions between bases of various strengths for these protons. For example, the strong acid HCl reacts with water according to Equation (1): HCl + H2O  H3O+ + Cl-

(1)

This acid is completely dissociated in dilute aqueous solution, which means that the [H3O+] of 0.1 M HCl is 0.1 M. Thus, HCl is a stronger acid than water and completely donates a proton to water to form H3O+. By experimentation, acetic acid, CH3COOH (abbreviated HOAc), is a weak acid and is not fully dissociated, as shown by Equation (2): H 2O +

HOAc

H3O+ +

OAc-

[2]

Its dissociation constant, as shown by Equation (3), is therefore small:

[3]

Acetic acid only partially dissociates in aqueous solution, and an appreciable quantity of undissociated acetic acid remains in solution. For the general weak acid HA, the dissociation reaction and dissociation constant expression are: HA + H2O

H3O+ + A-

[4]

[5] Recall that pH is defined as: [6]

We, therefore, can rearrange Equation (5) and combine it with Equation (6) in the following way: [7] 2|Page

[8]

[9]

During a titration of a weak acid, HA, with a strong base, there will be a point in the titration where 50 percent of the acid has been titrated to produce A- and 50 percent remains as HA. At this point [HA] = [A-], the ratio [A-]/[HA] = 1, and log [A-]/[HA] = 0. At one-half the equivalence point, Equation (9) becomes: pH = pKa

[10]

By performing a titration of a weak acid with a strong base and recording the pH versus the volume of base added, we can determine the ionization constant of the weak acid. From the resultant titration curve, we obtain the ionization constant as explained in the following paragraph. From the titration curve (Figure 1), we see that at the point denoted as 1/2 eq. point [HA] = [A-], the pH is 4.3. Thus, from Equation [10] at this point pH = pKa (where pKa is the -log Ka). pKa = 4.3, so -log Ka = 4.3, and log Ka = -4.3 therefore, Ka = antilog (-4.3) = 5.01 x 10-5 A graphical method for locating the equivalence point on the titration curve is to extrapolate from the vertical portion of the curve straight down to a point on the x-axis.

Figure 1: Titration curve for the titration of a weak acid with a strong base. 3|Page

In contrast, a titration curve for the titration of a strong acid with a strong base would appear as shown below.

Figure 2: Titration curve of a solution of a strong acid with a strong base

Compare and contrast these two curves. The equivalence point pH of Figure 1 is higher than that of Figure 2. Why does this occur? (Hint: Think about the products being formed.)

4|Page

Procedure: LAB SESSION I WEAR YOUR SAFETY GLASSES WHILE PERFORMING THIS EXPERIMENT A - Preparation of Approximately 0.100 M Sodium Hydroxide Heat 500 ml of distilled water to boiling in a 600 ml beaker, and after cooling under the water tap, transfer 200 mL of this carbonate free water to a 500 mL glass bottle fitted with a rubber stopper*. Add 1 ml of stock solution of sodium hydroxide (approximately 19 M) to the 500 mL bottle and shake vigorously for at least 1 minute. *

A rubber stopper should be used for a bottle containing sodium hydroxide solution. A strongly alkaline solution tends to cement a glass stopper so firmly that it is difficult to remove.

Clean a 25 ml burette using alconox solution, rinse and then fill with the prepared NaOH solution (recall all techniques for preparing a burette for titration). Weigh triplicate samples between 0.25 - 0.40 g each of pure potassium hydrogen phthalate (KHP) and put into a 125 mL or 250 mL Erlenmeyer flasks. Record masses on Data Section A. Add to each sample about 50 ml of distilled water that has been freed from carbon dioxide by boiling and warm gently with swirling until the salt is completely dissolved. Cool solution to room temperature by running cold tap water over the side of the flask. Add to each flask 2 drops of phenolphthalein indicator solution. Titrate each sample with the standard NaOH solution, recording all data. Calculate the molarity of your standard NaOH solution. LAB SESSION II: B: Determination of the Mass of Unknown Acid Per 22.00 mL NaOH ** This step is necessary to obtain a nice titration curve ** Obtain your unknown acid from the laboratory instructor. Record the initial mass of a clean dry 125 ml Erlenmeyer flask, add about 0.1 g of unknown acid, and reweigh. Record the exact amount of unknown acid. Add enough CO2 free water to dissolve your acid (~25ml). Add 2 drops of phenolphthalein solution to the acid and swirl. Fill the 25 ml burette with the standard sodium hydroxide solution. Titrate the unknown acid solution with the standard sodium hydroxide solution until the acid solution has a light pink color. Record the volume of NaOH solution used. Calculate the mass of unknown acid needed per 22 ml of NaOH solution based on the above data. *

The water is boiled to remove carbon dioxide, which would react with the sodium hydroxide and change its molarity. (Recall the results: H2O + CO2  ?).

5|Page

C: Molecular weight of An Unknown Acid Titrate two more times using the new calculated mass of the same unknown acid in part B for 22.00 ml of the NaOH solution, and find the unknown acid’s molar mass. Add ~45 ml of carbonate-free water to each Erlenmeyer flask containing your sample. Note: You should use 22.00 ml of the NaOH solution for each titration here IF the titration you did in step B is accurate AND the new calculated mass is measured out accurately. D: Determination of the pKa of an Unknown Acid Using either LoggerPro or Lab Quest (See instruction on pages 17) Weigh out another new calculated mass of the same unknown acid in part B that would use 22.00 ml of the NaOH solution and transfer it to a 250 ml beaker. Use ~45.00 mL of carbonate-free water to dissolve the unknown acid in the 250 ml beaker. Set the beaker underneath the burette tip and place the pH probe inside the beaker. NOTE: The pH probe should never be used to stir the solution. You must use a stir rod or swirl the beaker to mix the contents in the beaker. Swirl the beaker and record the initial pH of the unknown acid solution. Then add 1.0 ml of NaOH solution from the burette to the beaker, mix the solution in the beaker thoroughly and then record the pH. Repeat this process until ~25 ml of NaOH solution is added to the beaker. If you need to add more NaOH, you will need to refill the burette. Record the data under section D. Plot the pH of the solution on the ordinate and the volume of standard sodium hydroxide solution added on the abscissa. Calculate the pKa of the unknown acid using Equation (10) (Eq. 10)

pKa = pH

where the pH is the pH at one half the equivalence point in the titration of the acid. Calculate the acid dissociation equilibrium constant (Ka) of the unknown acid using pKa = - logKa

6|Page

CHEM-102

Name: ____________________

Experiment #6

Date: ____________________ Unk # ____________________

Data A.

Standardization of NaOH (grams of KHP) M(NaOH) = (MW KHP)(Volume of NaOH in Liters) Trial 1

Trial 2

Trial 3

Mass of KHP used

_______

_______

_______

Final burette reading

_______

_______

_______

Initial burette reading

_______

_______

_______

Volume of NaOH used in Liters

_______

_______

_______

Molarity of NaOH

_______

_______

_______

Average Molarity NaOH (for 3 trials)

_______

Percent error for each trial (use the average Molarity as true value) _______

_______

B.

_______

Determination of the Mass of the Unknown Acid that would use 22.00 mL NaOH Exact mass of unknown acid: ____________________ Vol. NaOH used to titrate the exact amount above (mL) ____________________ Calculate the mass of the unknown acid required to react with 22.00 ml of the standard NaOH solution using: (Exact mass of the unknown acid)

(X g unknown acid) =

(volume of the NaOH used)

(22.00 mL NaOH)

Mass of unknown acid (X) that would require 22.00 ml NaOH = _______ g. Use this mass for Part C and D. 7|Page

8|Page

C.

Molar Mass (MW) of an Unknown Acid (grams of unknown acid, X) MW (acid) = (M(Base))(V(Base) in Liters) Trial 1

Trial 2

Mass unknown acid (X)

_______

_______

Molarity NaOH (from A of the data sheet)

_______

_______

Final burette reading

_______

_______

Initial burette reading

_______

_______

Volume NaOH used (L)

_______

_______

MW of Unknown Acid

_______

_______

Average MW

_______

Show All calculations below:

9|Page

10 | P a g e

D.

Determination of pKa of Unknown Acid pH Meter Program Data: ml NaOH 0 1 ____2____ ____3____ ____4____ ____5____ ____6____ ____7____ ____8____ ____9____ ___ 10___ ___ 11___ ___ 12___ ___ 13___ ___ 14___ ___ 15___ ___ 16___ ___ 17___ ___ 18___ ___ 19___ ___ 20___ ___ 21___ ___ 22___ ___ 23___ ___ 24___

pH _________ _________ _________ _________ _________ _________ _________ _________ _________ _________ _________ _________ _________ _________ _________ _________ _________ _________ _________ _________ _________ _________ _________ _________ _________

Plot pH v.s. Volume of NaOH, and include the graph in the lab report.

11 | P a g e

12 | P a g e

Volume of NaOH at eq. point

__________ (from graph)

Volume of NaOH at ½ eq. point

__________ (from graph)

Experimental Value

Table Value

pKa

Ka

MW avg

Identify the unknown acid by comparing its experimental pKa and experimental molar mass to those of the known acids given in the table on page 15.

Name of Unknown Acid: __________________________________

13 | P a g e

14 | P a g e

REFERENCE TABLE

Acid Name

Molar Mass

pKa

Ka

Ascorbic Acid

176.14

4.10

7.94 x 10-5

Benzoic Acid

122.13

4.19

6.46 x 10-5

Chloroacetic Acid

94.47

2.85

1.40 x 10-3

B-Chloropropionic Acid

108.47

3.98

1.04 x 10-4

Citric Acid Monohydrate

70.05

4.47

1.68 x 10-4

Crotonic Acid

86.09

4.69

2.03 x 10-5

Glutaric Acid

66.07

5.41

3.89 x 10-6

Glycolic Acid

76.05

3.83

1.48 x 10-4

Maleic Acid

58.04

1.83

1.48 x 10-2

Malonic Acid

52.03

2.83

1.48 x 10-3

Oxalic Acid Dihydrate

63.04

4.19

6.40 x 10-5

Phenol

94.11

9.89

1.28 x 10-10

Potassium Hydrogen Phthalate

204.23

5.41

3.90 x 10-6

Propionic Acid

74.08

4.87

1.35 x 10-5

Sulfanilic Monohydrate

191.21

3.23

5.90 x 10-4

2-Tartaric Acid 75.04 2.98 1.04 x 10-3 ───────────────────────────────────────────────────────

15 | P a g e

16 | P a g e

PROCEDURE FOR PART D ON PAGE 6 USING LOGGER PRO ON LAB COMPUTERS 1. Obtain and wear goggles. 2. Double click the Vernier pH Titration icon to open the program. 3. Locate the Vernier LoggerPro interface, plug the pH sensor into CH1. 4. Take off the buffer solution bottle from the sensor and rinse the sensor with distilled water. 5. Prepare the unknown weak acid solution as described on page 6, part D. 6. Place the beaker on a magnetic stirrer and add a stirring bar. If no magnetic stirrer is available, you need to stir with a stirring rod during the titration.

Figure 1 7. Use a utility clamp to suspend a pH Sensor on a ring stand as shown in Figure 1. Position the pH Sensor in the acid solution and adjust its position so that it is not struck by the stirring bar. 8. Obtain a 25 mL buret and rinse the buret with a few mL of the NaOH solution. Use a utility clamp to attach the buret to the ring stand as shown in Figure 1. Fill the buret a little above the 0.00 mL level of the buret with the NaOH solution. Drain a small amount of NaOH solution so it fills the buret tip and leaves the NaOH at the 0.00 mL level of the buret. Record the precise concentration of the NaOH solution in your data table. Dispose of the waste solution from this step as directed by your 17 | P a g e

instructor. CAUTION: Sodium hydroxide solution is caustic. Avoid spilling it on your skin or clothing. 9. Before adding NaOH titrant, click Collect and monitor pH for 5-10 seconds. Once the displayed pH reading has stabilized, click Keep . In the edit box, type “0” (for 0 mL added). Press the ENTER key to store the first data pair for this experiment. 10. You are now ready to begin the titration. This process goes faster if one person manipulates and reads the buret while another person operates the computer and enters volumes. a. Add the NaOH titrant in 1.0 ml increment. When the pH stabilizes, click Keep . In the edit box, type the current buret reading, to the nearest 0.01 mL. Press ENTER. You have now saved the second data pair for the experiment. b. Continue adding NaOH solution in 1.0 ml increments until 30.0 ml of NaOH solution is added to the beaker. 11. When you have finished collecting data, click as directed by your instructor.

Stop

. Dispose of the beaker contents

12. Print a copy of the Table and Graph window.

PROCEDURE FOR PART D ON PAGE 6 USING LAB QUEST 1. Obtain and wear goggles. 2. Connect the pH Sensor to LabQuest and choose New from the File menu. 3. Take off the buffer solution bottle from the sensor and rinse the sensor with distilled water. 4. Prepare the unknown weak acid solution as described on page 6, Part D.

Figure 1 18 | P a g e

5. Use a utility clamp to suspend a pH Sensor on a ring stand as shown in Figure 1. Position the pH Sensor in the acid solution. 6. Obtain a 25 mL buret and rinse the buret with a few mL of the NaOH solution. Use a utility clamp to attach the buret to the ring stand as shown in Figure 1. Fill the buret a little above the 0.00 mL level of the buret with the NaOH solution. Drain a small amount of NaOH solution so it fills the buret tip and leaves the NaOH at the 0.00 mL level of the buret. Record the precise concentration of the NaOH solution in your data table. Dispose of the waste solution from this step as directed by your instructor. CAUTION: Sodium hydroxide solution is caustic. Avoid spilling it on your skin or clothing. 7. On the Meter screen (top left), tap Mode (top right). Change the data-collection mode to Events with Entry. Enter the Name (Volume) and Unit (mL) and select OK.

19 | P a g e

8. Conduct the titration carefully, as described below. This process goes faster if one person manipulates and reads the burette while another person operates the Lab Quest and enters volumes. a. Start data collection (green arrow). b. Before you have added any NaOH solution, tap Keep and enter 0 as the burette volume in mL. Select OK to store the first data pair.

c. Add the next increment of NaOH titrant (enough to raise the pH about 0.15 units). Make sure you are stirring the contents of the beaker as you are adding the titrant. When the pH stabilizes, tap Keep, and enter the current burette reading as precisely as possible. Select OK to save the second data pair. d. Continue adding NaOH solution in increments that raise the pH by about 0.15 units and enter the burette reading after each increment. Remember that as you are adding titrant, you should be stirring the contents of the beaker with the stir rod.

e. After a pH value of approximately 10 is reached, again add larger increments that raise the pH by about 0.15 pH units, and enter the burette reading after each increment. 20 | P a g e

f. Continue adding NaOH solution until the pH value remains constant or the graph levels out.

9. Stop data collection (red square) to view a graph of pH vs. volume. 10. Save your data on a USB drive and print out a copy of the Table and Graph windows on the lab computer using Logger Pro software

21 | P a g e...