AP Chemistry

Name: ___________________________

Three Complimentary Titration Labs Purpose: A. To standardize a base and use it to determine the concentration of acetylsalicylic acid (ASA) in an Aspirin tablet. B. To accurately determine the mass content of acetylsalicylic acid in an Aspirin tablet – a common analgesic (painkiller) medicine by titrating it with a strong base. C. To sketch the titration curve of acetylsalicylic acid by titrating it with a strong base and determine its Ka. Introduction: A. A chemical analysis that is performed with the aid of volumetric glassware (e.g. pipettes, burets, volumetric flasks) is called volumetric analysis. For a volumetric analysis procedure, a known quantity or a carefully measured amount of one substance reacts with a to-be determined amount of another substance with the reaction occurring in aqueous solution. The volumes of all solutions are carefully measured with volumetric glassware. The known amount of the substance for analysis is generally measured and available in two ways. 1. As a primary standard: a precise mass (thus moles) of a solid substance is measured on a balance, dissolved in water and then reacted with the substance to be analyzed. 2. As a standard solution: a measured number of moles of substance is present in a measured volume of solution – a solution of known concentration, generally expressed as molarity of the solution. A measured volume of the standard solution then reacts with the substance being analyzed. The reaction of the known substance with the substance to be analyzed, occurring in aqueous solution, is conducted by a titration procedure. A titration requires a buret to dispense a liquid, called the titrant, into a flask containing the analyte. The titrant may be a solution of known or unknown concentration. In this experiment, the titrant is sodium hydroxide and the analyte is an acid. A reaction is complete when stoichiometric amounts of the reacting solution substances are combined. In a titration this is the stoichiometric point. In the experiment, the stoichiometric point for the acid-base titration is detected by using phenolphthalein indicator. The point at which the phenolphthalein indicator changes color is called the endpoint of the indicator. Indicators are selected so that the stoichiometric point in the titration coincides with the endpoint of the indicator. Too much indicator in a solution may cause the pH to change thus required more titrant than necessary. For Part A, the reaction of sodium hydroxide with KHP is KHC8H4O4(aq) + NaOH(aq)  H2O(l) + NaHC8H4O4(aq)

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AP Chemistry

Name: ___________________________

History of Aspirin The name "aspirin" is composed of a- (from the acetyl group) -spir- (from the plant genus Spiraea) and -in (a common ending for drugs at the time). It has also been stated that the name originated by another means. "As-" referring to AcetylSalicylic and "-pir-" in reference to one of the scientists who was able to isolate it in crystalline form, Raffaele Piria. Finally, "-in", because it was a common ending for drugs at the time. Aspirin as a Weak Acid Aspirin is a weak acid with aa Ka= 3.27 x10-4. The chemical structures of the acid and its conjugate base are as follow. In this lab, we will use the short form HASA for acetylsalicylic acid and ASA− for the acetylsalicylate.

Indicators Phenolphthalein and the Universal Indictor Solution will be used in this lab. The Universal Indicator Solution has the following colours at various pH ranges. pH 3 4 5 6

Colour red-orange orange yellow greenish-yellow

pH 7 8 9 10

Colour green blue-green blue-gray violet

Materials Ring Stand pH meter Stirring Rod Buret Graduated Cylinder

Buret Funnel and Buret Clamp

NaOH (aq) (0.100 M)

3 Small / Medium Erlenmeyer Flasks Mortar and Pestle Electronic Balance 1 Small Beaker and 1 Medium Beaker

Aspirin Tablets (325 mg each) Phenolphthalein Universal Indicator Solution Ethanol Potassium hydrogen phthalate

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AP Chemistry

Name: ___________________________

Part A: Standardizing the Base OVERVIEW: An accurate molar concentration of NaOH (titrant) will be determined by using dry potassium hydrogen phthalate as a primary standard. The NaOH solution, now a secondary solution, is then used to determine the molar concentration of vinegar (and acid). Procedure 1. Prepare the primary standard acid. Measure approximately 0.50 g of potassium hydrogen phthalate (KHP) on weighing paper and transfer it to an Erlenmeyer flask. Record the exact mass. Similarly, prepare one more sample while you are occupying the balance. Dissolve the KHP in about 50.0 mL of deionized water and add 2 drops of phenolphthalein indicator. 2. Prepare the buret. Wash the 50 mL buret and funnel with water and soap. Flush the buret with tap water and rinse several times with deionized water. Rinse the buret with three 5-mL portions of the NaOH solution making certain that the inside of the buret is wet with NaOH solution. Any wastes may go down the drain. 3. Fill the buret. Using a clean funnel, fill the buret with the NaOH solution. After 10-15 seconds, read the volume by viewing the bottom of the meniscus. Bring the buret down to read the volume. Record the initial volume on your observations sheet. Place a white sheet under the Erlenmeyer flask containing the acid 4. Titrate the Primary Standard acid. Slowly add the NaOH solution from the buret to the first acid sample prepared in step 1 above. Initially add the NaOH in 1-2 mL increments. As the endpoint nears, the color fade of the indicator occurs more slowly. Occasionally rinse the walls of the Erlenmeyer flask with water from your wash bottle. The endpoint should be 1⁄2 drop of a slight pink and should persist for 30 seconds. Record the final volume of the NaOH in the buret. 5. Repeat the Analysis of the Remaining Standard acid samples. Refill the buret with NaOH and repeat the titration TWO more times with accurate masses of KHP. Part B: Titration of Aspirin with the standardized Base OVERVIEW: 1. Write the balanced molecular equation for the neutralization of HASA (aq) with NaOH (aq).

2. Given that each Aspirin tablet is 325 mg, calculate the volume of 0.100 M of NaOH (aq) required to completely neutralize the tablet.

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AP Chemistry

Name: ___________________________

3. Determine the both endpoint colours of phenolphthalein and universal indicator appropriate for this experiment.

Procedure 1 . Set up the titration apparatus with the ring stand, buret clamp, buret and buret funnel and perform a cleaning of the buret with the standardized base. 2 . Fill the buret with NaOH (aq) using the buret funnel. Be sure not to pass the 0 mL mark. 3 . Record the starting volume of the NaOH (aq). 4 . Measure and record the mass of an aspirin tablet. 5 . Label a medium beaker as ethanol. Obtain approximately 100 mL of ethanol from your instructor. 6 . Grind the aspirin tablet into powder using mortar and pestle. Pour about 10 mL of ethanol into the mortar to dissolve the powder (not all the powder will dissolve). Using a stirring rod as a guide, transfer the content into an Erlenmeyer flask. Repeat until all the powder from the mortar is transferred into the flask. (Alternatively, place the tablet directly into an Erlenmeyer flask. Add about 20 mL of ethanol into the flask. Using the stirring rod, break the tablet apart and dissolve it as much as possible into the ethanol solvent. Wash the stirring rod with a small amount of ethanol when finished.) 7 . Add a few drops of phenolphthalein in the flask. 8 . Begin titration of the aspirin. Swirl the Erlenmeyer flask when adding the NaOH (aq). Depending on the amount of ethanol used as the initial solvent for the aspirin tablet, the equivalence point is around pH 8.0 – 8.4. Stop the titration when the endpoint is reached. Record the final volume of the NaOH (aq). Calculate the net volume of base added. 9 . Repeat the titration with the other two Erlenmeyer flasks. Be sure to record the initial and final volume of the buret each time. Try to adjust the buret valve in such a way so NaOH (aq) is added one drop at a time around the endpoint.

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AP Chemistry

Name: ___________________________

Part C: Titration of Aspirin with NaOH using the Universal Indicator Solution and a pH meter. Procedure 1. Set up the titration as outlined in Part B, except we will use a small beaker instead of an Erlenmeyer flask and universal indictor solution instead of phenolphthalein. 2. Use a pH meter to measure and record the pH of the aspirin-alcohol solution. Record the initial volume of the buret. Add about 1.0 mL of NaOH into the beaker. After mixing thoroughly, measure and record the pH again. Record the final volume of the buret. 3. Repeat the previous step until you are 0.5 mL before the volume equivalence point. Add 0.1 mL at a time and record the pH and volumes. (For example: if you find that the equivalence volume is at 22.5 mL from the other trials, you should start measuring the pH at every 0.1 mL interval between 22.0 to 23.0 mL of NaOH added.) 4. Continue to add at 1.0 mL increments five more times for five more pH. Observations

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AP Chemistry

Name: ___________________________

Part B: Titration of Aspirin with the standardized Base

Part C: Volumes of 0.100 mol/L of NaOH Added to Aspirin Tablet and measured pH Mass of Aspirin Tablet used: _________________ Total Initial Volume

Final Volume

Volume Added

NaOH Volume Added

pH

Colour of Universal Indicator

Total Initial Volume

Final Volume

Volume Added

NaOH Volume

pH

Colour of Universal Indicator

Added

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AP Chemistry

Name: ___________________________

Analysis 1. Average the mass of the three Aspirin tablets used.

2. For each trial in Part B, determine the experimental mass of HASA used from the endpoint volumes of Table 1. Average these calculated experimental HASA masses.

3. Find the average mass percentage of acetylsalicylic acid in an average tablet.

4. Using the data in Table 2, graph the pH Curve between the Aspirin tablet and 0.100 M of NaOH (aq). Indicate the equivalence pH, volume of base added at stoichiometric point, the buffer region, and the pKa. 5. Using the pH recorded when there is no NaOH added, calculate the theoretical concentration of HASA (aq) −4 in the small beaker (Ka of HASA is 3.27 × 10 ).

6. Calculate the experimental Ka of acetylsalicylic acid.

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AP Chemistry

Name: ___________________________

Evaluation 1. What was the purpose of completing part A of this lab?

2. Compare the equivalence volume form the pH curve with the ones obtained in Part B of this experiment.

3. Given that the stated amount of acetylsalicylic acid in each tablet is 325 mg, calculate the % error for the mass of HASA.

4.

Contrast any difference in the experimental Ka obtained from the pH curve and the theoretical Ka of −4 3.27 × 10 by determining the % error of Ka. How is this % error compared to the mass % error?

5. Predict and explain what would happen to the calculated mass of HASA calculated when there is/are: a. water or ethanol left in the Erlenmeyer flask when the tablet is added to the flask.

b. water left in the buret when NaOH (aq) is added.

c. air bubbles in the buret when NaOH (aq) is added that is dislodged during the titration.

d. air bubbles in the buret when NaOH (aq) is added that is not dislodged during the titration.

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AP Chemistry

Name: ___________________________

5.

Besides human errors, what other sources of error could contribute to the % errors calculated above? (Hint: think about the nature of the solvent.) How might this error affect your average mass of HASA calculated?

6.

Why is alcohol used as the solvent instead of water?

7. Why is it not a good idea to take aspirin with alcoholic beverages?

8. How is aspirin different from Tylenol? (Research on the Internet.)

9. What amounts of aspirin do regular, extra strength, and children aspirin have in them?

10. Why is phenolphthalein a good choice of an indicator for this experiment?

11. In most cases, each tablet of aspirin has more mass than the indicated 325 mg of acetylsalicylic acid. What are the other substances that make up this difference? What is the main role of these “filler” substances? (Research on the Internet.)

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AP Chemistry

Name: ___________________________

12. Why are aspirin substitutes used by many people?

Conclusion 1. Summarize what you have learned from this lab.

2. State three sources of error

3. Explain three ways to improve the lab

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