Titration Using a pH Meter

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Titration Using a pH Meter Source file content last modified: 4/20/15 12:31:55 PM Display Content Selector See Previous: Standardization of Acids and Bases See Next: Conductivity of Acids and Bases Your browser does not support the video tag. Description Solutions of sodium hydroxide are standardized by titration of measured masses of standard reagents in water. These standard sodium hydroxide solutions are used to titrate unknown solutions of strong acids (HCl) and weaks acids (CH 3COOH). The endpoint is detected with a pH meter. Hazards Sodium hydroxide damages tissue and causes blindness. Phenolphthalein is toxic. pH electrodes are fragile. Precautions Wear eye protection. Handle electrodes with care. Cushion electrodes from shocks or jarring. Do not permit electrodes to dry out. Procedure Calibrating pH meter 1. Consult the manual for the pH meter to determine a calibration procedure. Use two buffers for the calibration. 2. Most electrodes are stored under water solutions (usually dilute buffer standards which also limit growth of organisms). Remove the electrode and rinse with distilled water. 3. Immerse the electrode in the first buffer solution (pH 7). Turn on the meter and make the first calibration adjustment (zero adjustment). 4. Turn the meter to standby. Remove, rinse, and wipe the electrode. 5. Immerse in the second buffer solution. Choose pH 4 if the critical measurements are slightly acidic or neutral; choose a pH 10 buffer if critical measurements are very basic. Turn on the meter and make the second adjustment (usually the range or temperature knob). 6. Turn the meter to standby and remove the buffer. Rinse the electrode and store it in dilute buffer or distilled water. 7. Adjust the temperature control on the meter, if necessary. (Use only if there is a large temperature difference. Room temperature is compensated with the standards.) Starting the Titration 1. Select an acid to be titrated. Choose a weak monoprotic acid such as acetic acid for the first titration. The sample can be obtained by any number of suitable methods. An aliquot or portion of vinegar is obtained using a pipet. 2. Quick disconnect bulbs are available for ordinary pipet bulbs. 3. Pipet a sample of acid using a volumetric pipet. Draw the acid in above the pipet calibration line. 4. Remove the pipet from the acid. Wipe the tip. Drain to the calibration line. Transfer the sample to a 250-mL beaker. 5. Place a stirring bar in the beaker. Add some distilled water to provide enough volume to allow clearance between the magnetic stirring bar and the electrode. Add 2-3 drops of 1 % phenolphthalein indicator. Adding the indicator permits the conventional endpoint of the titration to be detected. Therefore, it permits a comparison between the pH meter technique and the indicator technique. 6. Rinse the electrode with distilled water. Place the beaker on the magnetic stirrer and immerse the electrode. 7. Record the buret reading. 8. Turn on the pH meter. 9. Start and adjust the magnetic stirrer. Record the pH. 10. Set the buret stopcock to deliver a steady stream. 11. NOTE the volume when pH begins to change rapidly. Use this as an approximate end point. 12. If you are using a pH meter, start the buret flow. Add an increment of about 2 mL. (The added increments do not need to be the same size. They must be recorded accurately, however.) Turn off the buret. 13. Read and record the buret volume and the pH. 14. Continue adding titrant, reading the pH meter, and reading the buret volume. 15. Call attention to the relationship between the pH change at the equivalence point and the indicator color change. 16. Repeat the titration. Use the data from the first titration to determine the volume required to reach within 1 mL of the end point for the second titration. Add titrant 1 or 2 drops at a time once the point is reached. Read and record the volume and pH throughout this critical portion of the experiment. 17. Plot the volume of NaOH on the x-axis vs the pH on the y-axis. Determine the volume that corresponds to the midpoint of the sharply rising vertical portion of the curve. Using this volume, the volume of vinegar in the aliquot, and the concentration of the standard base, determine the molarity of the unknown acid. 18. EquivA = EquivB; MAVA = MBVB 19. Repeat the procedure described above, but use an aliquot of 1 M HCl. Record the data in the Data Table; calculate the molarity of hydrochloric acid. Handout Name ___________________________ Class ________ Teacher__________________________ DoChem 103 Titration Using A pH Meter Weak acid Vol mL pH

pH

Strong acid Vol mL pH

(trial 1) (trial 2)

pH

(trial 1) (trial 2)

1. The pH of a neutral solution is 7.0. Account for the difference between the pH of a neutral solution and the pH's of the various equivalence points or stoichiometric end points. (The point at which the number of moles of acid neutralize the number of moles of base). Identify indicators that could be used instead of a pH meter for these titrations. Is indicator selection equally critical for the two titrations, or does the selection make more difference in one of the two? 2. How does buffering affect the appearance of the titration curve in a weak acid vs. a strong base? 3. Use the data from the titration of CH 3COOH to estimate the acid dissociation constant for CH 3COOH. Handout Makeup Name ___________________________ Class ________ Teacher__________________________ DoChem 103 Titration Using A pH Meter Predict the outcome of titrating an acid, such as oxalic acid, that releases two moles of protons per mole of acid. Use this data to answer the questions below. Vol mL pH HCl pH CH 3COOH 0.0

1.00

2.87

1.0

0.95

3.21

2.0

0.98

3.52

4.0

1.05

3.85

6.0

1.11

4.06

8.0

1.18

4.21

10.0

1.24

4.35

12.0

1.31

4.46

14.0

1.39

4.57

16.0

1.45

4.66

18.0

1.55

4.77

20.0

1.65

4.87

22.0

1.77

4.97

24.0

1.91

5.08

26.0

2.10

5.20

28.0

2.42

5.34

29.0

2.73



29.1

2.78



29.2

2.83



29.3

2.89



29.4

2.96



29.5

3.04



29.6

3.13



29.7

3.26



29.8

3.44



29.9

3.73



30.0

7.00

5.51

30.1

10.20



30.5

11.00



31.0

11.20



32.0

11.50

5.75

34.0



6.20

34.1



6.24

34.2



6.28

34.3



6.33

34.4



6.38

34.5



6.44

34.6



6.51

34.7



6.59

34.8



6.69

34.9



6.82

35.0

11.90

8.94

36.0

11.20



37.0

11.50



40.0

12.20

11.80

1. The pH of a neutral solution is 7.0. Account for the difference between the pH of a neutral solution and the pH's of the various equivalence points or stoichiometric end points. (The point at which the number of moles of acid neutralize the number of moles of base). Identify indicators that could be used instead of a pH meter for these titrations. Is indicator selection equally critical for the two titrations, or does the selection make more difference in one of the two? 2. How does buffering affect the appearance of the titration curve in a weak acid vs. a strong base? 3. Use the data from the titration of CH 3COOH to estimate the acid dissociation constant for CH 3COOH. Teachers Guide Purpose To use a pH meter to determine the concentration of an unknown acid solution. To plot and interpret titration curves. Materials (per demonstration) 1 pH meter standard buffers, such as pH = 7.00 and pH = 4.00. These may be bought as premixed powders, or as liquids. 150 mL 1 M HCl (8.6 mL of 11.6 M acid (36%) per 100 mL of solution) 250 mL of a standard NaOH solution (~ 1.0 M; 10 g NaOH per 250 mL of solution) 150 mL 5% acetic acid (vinegar) magnetic stirrer 4 250-mL beaker 1 25- or 50-mL buret 1% phenolphthalein (in aqueous ethanol) 1 support stand with single buret clamp distilled water in a 250- or 500-mL squeeze bottle 1 25.00- or 50.00-mL volumetric pipet and pipet bulb 1 thermometer paper towels Lab Hints Use this experiment for all students if you have enough pH meters. Increase materials list by factor of 5. A standard solution of NaOH can be made by titrating a 1 M solution of NaOH (10 g/250 mL) against a solid acid like KHP (potassium hydrogen phthalate). The solution need not be exactly 1.00 M but should be known to three significant figures. Be sure to inform the students of the exact molarity. If students save their NaOH solutions from the classical titration experiment it will save a lot of time. Read the manual that comes with the pH meter. Electrodes that have gone dry may need a soaking period of several days. Older electrodes may need filling; check for a filling hole under a flexible band near the top of the electrode. (See instructions with the electrodes.) Modern electrodes are much more rugged than older all glass electrodes. The new electrodes will require less soaking and last much longer. These electrodes are sealed and do not require filling.

Time Teacher preparation: 10-15 minutes if solutions are already prepared; otherwise add 10 min per solution for preparation Presentation: 20-35 minutes Precautions Wear eye protection. Provide a working eye wash. An emergency shower should be available. Handle electrodes with care. Cushion electrodes from shocks or jarring. Do not permit electrodes to dry out. Disposal The solutions used in this experiment may be disposed of safely at the sink after neutralization. Discard paper towels with ordinary trash. Save any unused stock solutions in bottles. Background In this experiment, titrations will be performed using a pH meter and an indicator. The calculations of the concentration of an unknown solution are the same as in previous experiments. By plotting the data in a graph of "pH of the Solution" vs "Volume of NaOH Added", the volume that is needed to neutralize all of the acid can be found on the graph. This volume of titrated acid and the volume and molarity of the NaOH titrant can be used to find the molarity of the acid. It is instructive to titrate a strong acid, a weak acid, and a polyprotic acid when strong base is available. Presentation? Presentation Question: Predict the outcome of titrating an acid, such as oxalic acid, that releases two moles of protons per mole of acid. This is often a source of confusion for many students. They usually guess that twice as much titrant will be used which is correct. However, they usually don't guess that the second mole of protons is more difficult to remove than the first, so that the titration actually takes place in two steps. Sample Data Vol mL pH HCl pH CH 3COOH 0.0

1.00

2.87

1.0

0.95

3.21

2.0

0.98

3.52

4.0

1.05

3.85

6.0

1.11

4.06

8.0

1.18

4.21

10.0

1.24

4.35

12.0

1.31

4.46

14.0

1.39

4.57

16.0

1.45

4.66

18.0

1.55

4.77

20.0

1.65

4.87

22.0

1.77

4.97

24.0

1.91

5.08

26.0

2.10

5.20

28.0

2.42

5.34

29.0

2.73



29.1

2.78



29.2

2.83



29.3

2.89



29.4

2.96



29.5

3.04



29.6

3.13



29.7

3.26



29.8

3.44



29.9

3.73



30.0

7.00

5.51

30.1

10.20



30.5

11.00



31.0

11.20



32.0

11.50

5.75

34.0



6.20

34.1



6.24

34.2



6.28

34.3



6.33

34.4



6.38

34.5



6.44

34.6



6.51

34.7



6.59

34.8



6.69

34.9



6.82

35.0

11.90

8.94

36.0

11.20



37.0

11.50



40.0

12.20

11.80

Closure? Closure Questions: 1. The pH of a neutral solution is 7.0. Account for the difference between the pH of a neutral solution and the pH's of the various equivalence points or stoichiometric end points. (The point at which the number of moles of acid neutralize the number of moles of base). Identify indicators that could be used instead of a pH meter for these titrations. Is indicator selection equally critical for the two titrations, or does the selection make more difference in one of the two? 2. How does buffering affect the appearance of the titration curve in a weak acid vs. a strong base? 3. Use the data from the titration of CH 3COOH to estimate the acid dissociation constant for CH 3COOH. Answers to Closure Questions: 1. For HCl, the titration has a large steep region, and the pH at the equivalence point is 7. For CH 3COOH, the equivalence region is less steep, and occurs at a pH well above 7. Phenolphthalein (pH 8.2-10) will serve as an indicator for either titration. Bromthymol blue (pH 6-8) or thymol blue (pH 8-9.6) will work for the HCl titration. 2. Buffering causes the weak acid titration curve to change direction three times: first it bends down toward a buffer region; then it bends up toward the equivalence point; and, after the equivalence point, it bends down as more base is added.

3. Since: Ka = [H +] [CH 3COO-] / [CH 3COOH] When the concentration of the acid and its corresponding base are equal, as in the midpoint of the buffer region, the ratio is equal to one (unity). At this point, the pH is equal to the negative logarithm of the acid dissociation constant. Since this region is fairly flat, it is rather easy to estimate the pKas of acids during weak acid titration. Key Words pH pH meter concentration acid base endpoint buffer titration This static was created at 12:31:55 PM on Monday, April 20, 2015

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Titration Using a pH Meter

Titration Using a pH Meter Source file content last modified: 4/20/15 12:31:55 PM Display Content Selector See Previous: Standardization of Acids and ...

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