an Identifying Unknown Weak Acid
Hunter Morgan, Katelyn Heck, and Jamie Elise King Chemistry 1212 Lab Matthew Morgan
November 26, 2007
Introduction The acidity of a solution is a very important aspect of life when it comes to food, water, and agriculture among other things. The strength and quality of acid is a very important aspect of our daily lives. The strength of an acid must be measured in order to find the quality of these solutions with respect to their purpose in our lives. One way to measure the strength of an acid is through the acid ionization constant, Ka. This is the quantitative measure of the strength of the acid through its ability to donate protons to a base. The following equation of the ionization of a generic acid, HA, can be used to find the corresponding acid ionization constant expression: [H3O+][A-] Ka= ---------------[HA] The corresponding acid ionization constant expression, Ka, is a characteristic of an acid and is often used to identify an unknown acid. The Ka constant represents the strength of the acid; thus, a stronger the acid has a larger the Ka value, and a weaker the acid has a lower Ka value. In this experiment, we will determine an unknown, monoprotic acid solution by using three experimental techniques. The first technique we will use to determine the unknown acid is titration of the acid with 0.100 M NaOH. We will titrate 80 mL of the 0.100 m unknown acid solution with sodium hydroxide, which will produce a titration curve plotting the pH of the acidic solution versus the volume of sodium hydroxide added. Once all of the unknown acid has completely reacted with the sodium hydroxide, the equivalence point has been reached. The equivalence point is the center region of the titration curve in which the pH sharply increases. The charted equivalence point is used
to find the half-equivalence point, which is the point in the titration that exactly one half of the neutralization of the acid has occurred, which also shows half of the needed base has been used. At the half-equivalence point, the concentration of the unknown acid in the solution, [HA] is equal to the neutralizing base [A -]. Which is written as: [HA]=[A-] This allows the first equation to be simplified into: Ka=[H3O+] The negative logarithm of each side will allow us to derive an equation to determine the uknown acid's pH: -log(Ka)= -log[H3O+] pKa= pH The pKa for the acid is equal to the pH of the acid at the half-equivalency point. The following equation is used to determine the Ka of the acid: Ka=10-pKa The second experimental technique we will use is determining the percent ionization of the acid from the freezing point of depression. A curve will be produced plotting the freezing point of the acid which will allow us to determine the unknown solution based on its freezing point relative to water. Once the solution freezes, the freezing point of depression of the unknown acid will allow us to identify the acid. The third experimental technique we will use to determine the unknown acid is by first determining the pH of the acid and the initial concentration of the acid in the solution. Once the pH of the acid is determined, we can determine the ion concentrations at equilibrium. The ion concentration is related to the pH of the solution by:
[H3O+]=10-pH At equilibrium a fraction of the HA molecules are ionizing, which allow us to determine these concentrations by using the first given equation with the ion change in concentration. Once the Ka value is determined by using the three experimental techniques, we will find the average Ka value and compare it to common Ka values in order to determine the actual identity of the unknown monoprotic acid.
Procedure: Part A Determination of a weak acid by titration with NaOH
1. Obtain approximately 40 mL of an unknown weak acid solution. Record the unknown number on the Data Sheet. 2. Add 10.00 mL of the unknown acid solution to a clean, dry 250-mL beaker. 3. Rinse a 50-mL buret with distilled water and make sure to rinse the tip. Close the buret stopcock. Add approximately 3-4 mL of NaOH solution to the buret and rinse the buret by tilting it on its side and twirling it. Drain the NaOH through the tip of the buret. Discard the NaOH in the sink and flush the sink with water. 4. Close the buret stopcock and fill it with 50.0 mL of the NaOH solution. 5. Set up the MeasureNet workstation: a. Press the On/Off button. b. Press Main Menu, then F3 pH/mV, and then F2 pH vs. Volume. c. Press Calibrate. The pH probe will be sitting in a beaker of pH 7.00 buffer solution. Enter the temperature of the buffer solution (measured with a thermometer) and then press Enter. d. When asked for the pH of the buffer, enter 7.00 and then press Enter. e. Press F1 for a 1 point calibration. f. Press Display to accept the values. 6. Remove the pH electrode from the buffer solution, rinse it with distilled water, and then dry the tip with a Kimwipe. 7. Support the pH probe with the drop counter, and then place the probe in the beaker containing the acid. If the tip of the probe is not submerged, add enough water to cover the tip of the probe. 8. Put a stir bar into the acid solution, and then place the beaker on a magnetic stirrer. Turn the magnetic stirrer onto a low setting. 9. Position the buret filled with NaOH over the beaker of the acid solution. The tip of the buret must be centered in front of the "eye" of the drop counter. 10. Press Start. Read the amount of NaOH in the buret and enter the initial volume at the workstation, and then press Enter. (If your buret is completely filled, enter 0.00 mL as the initial volume.) 11. Press Start. Open up the buret valve and vegin adding the NaOH by drops until the titration is complete. 12. When the titration is complete, close the buret stopcock and pres Stop. 13. Read the volume of the remaining NaOH in the buret and enter the final volume of the NaOH at the workstation. Press Enter. 14. Press File Options. Press F3 to save the scan. Enter a 3 digit code to name the file, then press Enter. 15. Rinse the pH probe with distilled water and then place it back in the buffer solution. 16. Press Display to clear the previous scan. 17. Empty the NaOH from the buret and rinse the buret with distilled water. 18. Turn off the MeasureNet workstation. 19. Repeat steps 3 through 18. 20. From the scans that you saved, prepare the titration curves using Excel. 21. From the curves, determine the volume of NaOH required to neutralize the weak acid solution (equivalence point) in each titration. Record the equivalence points of each titration on the Data Sheet. 22. Calculate the Ka value for titrations 1 and 2. 23. Calculate the average Ka value for the weak acid. 24. Using the average Ka value, identify the unknown weak acid.
25. Calculate the molarity of the unknown acid for each titration, and then average the two values.
Part B: Identifying a weak acid by determination if the pH of the acid
1. Add 20.0 mL of the unknown weak acid solution (the same one used in Part A) into a clean, dry 50-mL beaker. 2. Take the pH probe from the pH 7 buffer solution. Rinse the probe with distilled water and then dry the tip with a Kimwipe. 3. Insert the probe into the beaker containing the unknown acid. After the pH reading stabilizes, record the pH on the Data Sheet. 4. Put the acid solution into the waste container. 5. Repeat steps 1 through 4 for a second measurement. 6. Rinse the pH probe with distilled water and place it back in the buffer solution. 7. Using the molarity calculated in step 25 of Part A, and the pH of the unknown acid solution, calculate the Ka value for each trial. Calculate the average Ka value. 8. Using the average Ka value, indentify the unknown weak acid from Table 1.
Part C: Determination of the Ka value by determination of the percent ionization of the acid from freezing point depression
1. Determining the freezing point of the pure solvent a. Press On/Off on the MeasureNet workstation. b. Press Main Menu, then F2 Temperature, and then F1 Temperature vs. Time. c. Fill a 150-mL beaker with half ice and half water. In order to calibrate the temperature probe, press the Calibrate. Enter 0.00C as the actual temperature of the ice bath. Press Enter. d. Press SetUp and then F1 to set the limits for a new scan. Use the left and right arrow keys to switch between min and max. The Y axis is the temperature axis. Set the minimum to -100C and press Enter. Set the maximum to 250C and press Enter. The X axis is the time axis. The minimum should be set to 0 seconds and the maximum to 1500 seconds. Press Display to accept all values. e. up an ice bath by putting 200 mL of water into a 600-mL beaker. Put 15 to 20 grams of NaCl or CaCl2 into the water and stir until the salt is completely dissolved. Fill the rest of the beaker with ice. f. Put together the stopper/wire/temperature probe by spreading the slit hole in the 2-hole rubber stopper and insert the temperature probe. Then put the wire stirrer (one end is a loop) into the other hole in the stopper. Make sure that the temperature probe passes through the loop. Put the stopper/wire/temperature probe onto a ring stand using a utility clamp. g. In a 25 X 200 mm test tube, add 15-20 mL of distilled water. h. Put the stopper/wire/temperature probe into the test tube. Fit the stopper into the test tube. Put the entire assembly into the ice bath beaker. i. Press Start to begin recording a thermogram on the MeasureNet workstation. Move the wire stirrer up and down until the water in the test tube is frozen. The water is frozen when the temperature remains constant for 20-30 seconds. j. Once the water freezes, push Stop. Then press File Options and then F3. Enter a 3 digit number to save the thermogram. Press Enter and then Display. 2. Determining the percent ionization of the unknown acid by freezing point depression
a. Obtain 15-20 mL of a 0.100 m unknown acid solution. Record the unknown number on the data sheet. b. Pour the acid solution into a 25 X 200 mm test tube, and put the stopper/wire/temperature probe into the test tube. Make sure that the bottom of the temperature probe is covered with the acid. c. Put the entire assembly into the beaker with the ice bath and secure it to a ring stand with a utility clamp. Press Start to begin recording the thermogram. Move the wire stirrer up and down to freeze the solution inside the test tube. d. Once the solution freezes, press Stop. Then push File Options, and then F3. Put in a 3 digit number to save the file name for the thermogram and press Enter. Then press Display. e. Repeat steps a through d for a second trial. f. Prepare a thermogram for each trial and for pure water on Excel. i. Start by opening an Excel worksheet. Click on File Open, and open the saved MeasureNet tab delimited file with the temperature versus time data for the solvent and then on click Finish. Copy and paste the first two columns in the tab delimited file onto column A (time) and column B (temperature) in the worksheet. Close the tab delimited file for the solvent. ii. Click on File Open, and open the saved MeasureNet tab delimited file with the temperature versus time data for the solution and then click on Finish. Copy and paste the first two columns in the tab delimited file onto column C (time) and column D (temperature) in the worksheet. Close the tab delimited file for the solution. iii. Find the Chart Wizard icon on the toolbar and click on it. Then click on XY Scatter Plot and then Smooth Line Type. Click on Next, then Series, then Add. iv. When the Name box appears, type in the name of the curve (water or solvent). Then click inside the x-values box and then highlight the cells in column A (time for solvent). Then click inside the y-values box and then highlight the cells in column B (temperature for solvent). Then click on the Add button. v. For columns C and D, repeat step 23 to show the time and temperature data for the solution. vi. Click on Next, then Titles. In this box enter the title for your graph, for example Freezing Point Depression. Label the x-axis in the x-value box (Time, s). Label the y-axis in the y-value box (Temperature, oC). vii. Repeat steps i through vi for the second trial. g. Record the freezing points for water and the acid solution on the Data Sheet. h. Evaluate freezing point depression for both trials and record them on the data sheet. i. Calculate the van't Hoff factor for both trials and then evaluate the average van't Hoff factor for the unknown solution. Record all of this on the data sheet. j. Using the average van't Hoff factor, calculate m effective and record it onto the Data Sheet. k. Calculate m ionized and record it onto the Data Sheet. l. Calculate the percent ionization and find the concentrations of H+, A-, and HA and record on the Data Sheet. m. Calculate the Ka value by using the concentrations and record onto the Data Sheet. Then identify the unknown by looking at table of Ka values.
Data Sheets Experiment 28 Table 1- Ka values for different acids Acid Formula Acetic acid CH3COOH Benzoic acid C6 H5COOH Carbonic acid H2CO3 Formic acid HCOOH Hypochlorous acid HOCl Dihydrogen phosphate ion H2PO4Hydrogen phosphate ion HP O32Hydrogen carbonate ion HC O3Nitrous acid HNO2 Phenol C6H6O Potassium hydrogen phthalate KC8H5O4
Ka 1.8 6.3 4.2 1.8 3.5 6.2 3.6 4.8 4.8 4.5
10-5 10-5 10-7 10-4 10-8 10-8 10-13 10-11 10-11 10-4
5.3 10-6
Part A: Determination of the Ka Values of a Weak Acid by Titration with NaOH
Titration of Unknown Acid with NaOH (trial 1)
14 12 10 8 6 4 2 0
pH
pH
0
5
10
Volume, mL
15
20
Titration of Unknown Acid with NaOH (trial 2)
14 12 10 8 6 4 2 0 0 5 10 Volume, mL 15 20
pH
pH
1. Unknown Acid #: 2. Equivalence Point: 3. Half Equivalence point pH at the half-equivalence point File name of graph 4. Calculate the Ka Values for titration 1 and 2: HA (aq) + H20 (l) H30+ (aq) + A- (aq) Ka = [H30+] [A-] [HA] Ka = [H30+] -log(Ka) = -log[H30+] pKa = pH Ka = 10-pKa [H30+] = 10-pH Trial 1: Ka = 10(-4.71)= 1.95 X 10(-5) Trial 2:
Trial 1 Unknown Acid 11.846 5.923 4.71 4123 ml ml
Trial 2
10.946 5.473 4.74 4456
ml ml
.
Ka = 10(-4.74)= 1.82 X 10(-5) 5. Average Ka values for the unknown weak acid 1.89 X 10(-5)
((1.95 X 10(-5)) + (1.82 X 10(-5)) ) / 2 = 1.89 X 10(-5)
6. Using the Table of Ionization Constants (Ka ) for weak acids at 25 C determine the identity of the unknown acid you are working with. Acetic Acid .
7. Calculate the molarity of the unknown acid solution for each titration. Trial 1: .100 mol NaOH 1 mol acid 1.00L 1 11.846mLof NaOH 1000mL 1 L NaOH 1 mol NaOH 10 mL acid Trial 2:
1000 mL L
0.118M
10.946 mLof NaOH
1.00L 1000mL
.100 mol NaOH 1 L NaOH
1 mol acid 1 mol NaOH
1 10 mL acid
1000 mL L
0.109M
8. Calculate the average molarity of the unknown acid solution. (0.118M + 0.109M) / 2 = 0.114M
Part B: Identifying a weak acid by determination if the pH of the acid
Trial 1 9. Determine the pH of the unknown acid solution 10. Calculate the Ka for each Trial HA (aq) Initial 0.10 M + + H20 (l) H30+ (aq) 0 + A- (aq) 0 2.89
Trial 2 2.87 .
Change - [Value of 10-pH]
[Value of 10-pH] + [Value of 10 -pH]
.
Equilibrium (0.10 M (Value of 10-pH) M [Value of 10-pH] + [Value of 10 -pH] Ka = [H30+] [A-]
[HA] Trial 1 Initial 0.10 M 0 + 1.29 X 10-3
-2
0 + 1.29 X 10-3
.
Change -1.29 X 10-3 Equilibrium 9.87 X 10 M Ka = [H30+] [A-] [HA]
1.29 X 10 M = [1.29 X 10-3] [1.29 X 10-3] [9.87 X 10-2]
-3
1.29 X 10 M = 1.69 X 10-5
-3
Trial 2 Initial 0.10 M 0 + 1.35 X 10-3
-2
0 + 1.35 X 10-3
. -3
Change - 1.35 X 10-3 Equilibrium 9.87 X 10
1.35 X 10 M = [1.35 X 10-3] [1.35 X 10-3] [9.87 X 10-2]
1.35 X 10 M = 1.85 X 10-5
-3
Ka = [H30+] [A-] [HA]
11. Average Ka for the weak acid
1.77 X 10-5
.
12. Using the Table of Ionization Constants (Ka ) for weak acids at 25 C determine the identity of the unknown acid you are working with. Acetic Acid .
Part C: Determination of the Ka value by determination of the percent
ionization of the acid from freezing point depression
Freezing Point of Water (trial 1)
15 Temp, deg C
10
5 0 -5 0 100 200 Time, s 300 Temperature/de gC
Freezing Point of Unknown Acid (trial 1)
20 Temp, deg C 10 0 -10 0 500 Time, s 1000 Temperature/de gC
Freezing Point of Water (trial 2)
Temperature, deg C 20 15 10 5 0 Temperature/de gC 200 400 Time, s 600
-5 0
Freezing Point of Unknown Acid (trial 2)
30 Temp., deg C 20
10
0 -10 0
Temperature/de gC
200 Time, s
400
13. Unknown Acid #: 0.100M Unknown Trial 2 14. File name of graph 5100 / 5200 4741 / 4852 15. Freezing Point of Water o o -0.683 C -0.367 C
o
. Trial 1
.
.
-0.921 C
Freezing Point of Acid o -0.667 C
. 0.238 C
o
16. Tfobserved of acid solution: o 0.300 C . solution Trial 1: -0.683 C (-0.921 C) = 0.238 C
o o o
Tf = freezing pt of pure solvent freezing pt of
Trial 2: -0.367 C (-0.667 C) = 0.300 C
o
o
o
o
Tfpredicted of acid solution:
o
0.186 oC _____
o
0.186
C____
Trial 1: Tf = Kf m = (1.86 C/ m) 0.100 m = 0.186 C Trial 2: Tf = Kf m = (1.86 oC/ m) 0.100 m = 0.186 oC
17. Calculate the van't Hoff factor, I i= Tf (observed)__________ Tf (predicted if non electrolyte)
1.28
1.61
.
Trial 1: i = 0.238 oC / 0.186 oC = 1.28 Trial 2: i = 0.300 oC / 0.186 oC = 1.61
18. Calculate the average van't Hoff factor (1.28 + 1.61) / 2 = 1.45
1.45
.
19. Calculate the m effective m effective = i (m stated ) = 1.45 (0.100m) = 0.145m 20. Calculate m ionized
0.145m
.
0.0450m
.
m ionized = ( m effective) (m stated ) = 0.145m 0.100m = 0.0450m 21. Calculate the % ionized % ionized = [0.0450] ionized x 100 % [0.100 ] initial 22. Record the Concentrations: [HA] = 0.100m 0.0450m = 0.0550m [H+] = 0.0450m [A-] = 0.0450m 23. Calculate Ka : Ka = [0.0450] [0.0450] / [0.0550] = 3.68 X 10-2 24. Using the Table of Ionization Constants (Ka ) for weak acids at 25 C determine the identity of the unknown acid you are working with. Formic Acid . 45% .
Discussion Our goal in this lab was to determine the acid ionization constant, Ka, using three different methods and comparing our results in the end. We were able to find Ka, by titrating the weak acid with sodium hydroxide, measuring the pH of the solution, and freezing point depression. Each of these processes allowed us to determine the ionization constant in a different way. Similar results would infer that our experimentation procedures were precise and accurate; however slight differentiation in our Ka values would be due to human error during the course of the experiment. Variation and precision errors can occur in each of the three procedures. When titrating the weak acid with sodium hydroxide the elements could have been set up incorrectly so that it did not measure the amount of drops correctly. While dealing with the pH probe, we could have forgotten to wipe it and rinse with water distilled water. The freezing point depression could have miscalculated due to the fact that there was not enough ice to give an accurate freezing point. These sorts of mistakes may seem minimal individually but collectively lead to the variation within the results that we got in our experiment.
Conclusion In this experiment, three different methods were used to determine the Ka value of an unknown weak acid. The first method was the titration of the acid with 0.100 M NaOH. The second method used was to identify the weak acid by determining the pH of the acid. The third method used was determining the percent ionization of the acid through the freezing point of depression. Unfortunately, after performing each of these methods, the data and the calculations led to three different average Ka values. The average Ka value from the titration of the weak acid with 0.100 M of NaOH was 1.89e-5. Whereas, the average Ka value determined from the pH of the weak acid was 1.77e-5. Finally, the average Ka value determined by the percent ionization of the weak acid through the freezing point of depression was 3.68e-2. The average Ka from all three methods is 1.23e-2; however, we found our unknown weak acid to have a Ka value closest to Acetic Acid. Since the same unknown weak acid was used in all three methods, the data collected should have allowed us to calculate Ka values that are much closer in numerical value. After checking over each of the calculations several times, the calculations appear to be correct. Therefore, it is likely experimental sources of error in some, if not all three, of the preformed experimental procedures. Common sources of error that could have occurred during the titration of the weak acid with 0.100 M NaOH include not properly cleaning the burette used for the NaOH with distilled water, not properly cleaning the lab equipment before use, or problems with the MeasureNet technology or the drop counter. Any of these could have led to an error in the data collected. In addition, expiremental errors could have occurred while using the method to determine the pH of the weak acid,
including: not cleaning the pH MeasureNet probe properly, with an insufficient amount of distilled water, or not properly cleaning the lab equipment before use. Finally, experimental errors that are common while determining the percent ionization of the weak acid through the freezing point of depression include: not having a cold enough ice bath, not stirring the solution in the ice bath sufficiently, using insufficiently cleaned lab equipment, and recording the data before the solution is completely frozen.
MeasureNet File # 4123 4456 5100 5200 4741 4852
Description Titration of Unknown Weak Acid with NaOH (trial 1) Titration of Unknown Weak Acid with NaOH (trial 2) Freezing Pt of Water (trial 1) Freezing Pt of Unknown Acid (trial 1) Freezing Pt of Water (trial 2) Freezing Pt of Unknown Acid (trial 2)