ASP Analysis

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CH229 General Chemistry LaboratoryDr. Deborah Exton

VOLUMETRIC ANALYSIS OF ASPIRIN

1. Purpose

In this laboratory experiment, you will assess the purity of a sample of aspirin and becomeacquainted with the concept of titration analysis.

2. Pre-lab Reading

Volumetric Analysis (accompanying handout)

3. Background

The aspirin which you previously synthesized is probably not pure, despite your best efforts.

The most likely impurities are acids - either salicylic acid (unreacted starting material) or acetic

acid. Even commercially prepared aspirin tablets are not 100 percent acetylsalicylic acid. Most

aspirin tablets contain a small amount of binder which helps prevent the tablets from crumbling.

The binder is chemically inert and was intentionally added by the manufacturer, but its presence

means that aspirin tablets do not have 100 percent purity. Moreover, moisture can hydrolyze

acetylsalicylic acid. Thus, aspirin which is not kept dry can decompose. Acetic acid is the

hydrolysis product formed by the reaction of water with acetylsalicylic acid:

You may have noticed the smell of vinegar (acetic acid) when opening an old bottle of aspirin, or a

bottle which has not been properly sealed.

In this experiment you will determine the purity of the aspirin which you previously

synthesized. (If you missed the Aspirin Synthesis laboratory or failed to synthesize enough aspirin

for analysis, you will be given a sample of acetylsalicylic acid.) More specifically, you will

determine the percentage of acetylsalicylic acid in your aspirin sample by means of a titrimetric

analysis. In this procedure, you will incorporate a process known as a back titration. In a normal

titration, an experimenter is able to determine the amount of analyte present in a solution by

carefully adding incremental volumes of a standard solution until reaction between analyte and

titrant is judged to be complete. (If you have not yet read “Volumetric Analysis: Titration” youshould do so now before reading any further.) Occasionally, it is convenient or necessary to add an

excess of the titrant and then titrate the excess with another reagent. This process is called back

titration. In this technique, a measured amount of the reagent, which would normally be the titrant,

is added to the analyte sample so there is a slight excess of reagent present. After this reagent reacts

completely with the analyte, the amount of excess (unreacted) reagent is determined by titration

with another standard solution.



CH229 Volumetric Analysis of Aspirin

The amount of reagent which reacted with the analyte can be found by determining the difference

between the amount of reagent added and the amount in excess:

moles reagent reacted (with analyte) = total moles added - excess moles back-titrated

Once this quantity has been determined, stoichiometric considerations will allow you to find the

moles of analyte initially present.

At room temperature, acetylsalicylic acid can be neutralized with base:

If acetylsalicylic acid were the only acid present in your sample, you could determine the purity

of the aspirin by a simple titration of your sample with sodium hydroxide. However, if acidimpurities are present, titration of the aspirin will neutralize not only the acetylsalicylic acid

(previous equation) but the acid impurities as well. Thus, from such a titration, one could calculate

the total number of moles of acid present in the sample by measureing the volume of standardized

NaOH required to reach the end point. If the stoichiometric ratios between acid and base are 1:1 (as

in this experiment), the total number of moles of acid may be calculated by:

moles acid = moles base = (liters base) x (molarity base)

The titration of an impure sample of aspirin will yield the conjugate bases acetate ion, salicylate

ion, and acetylsalicylate ion. Of these, only the acetylsalicylate ion is an ester. It will react with

additional base reasonably rapidly at elevated temperatures :

This reaction represents what is termed a base-promoted hydrolysis, or saponification, of esters.

The reaction is the reverse of the esterification process which you employed while synthesizing

aspirin. After you have neutralizd all acidic material in the aspirin by titration with base, you will

add a known excess amount of base to cause the saponification to occur. The excess base that isnot consumed in the hydrolysis will be determined by a back-titration with standard HCl. From

your data, you will be able to calculate the grams of acetylsalicylic acid in your aspirin sample.

The following example may help you understand this calculation:

Example:

A 0.5130-g sample of aspirin prepared by a student required 27.98 mL of 0.1000

M NaOH for neutralization. An additional 42.78 mL of 0.1000 M NaOH was




added, and the sample was heated to hydrolyze the acetylsalicylic acid. After

the reaction mixture cooled, the excess base was back-titrated with 14.29 mL of

0.1056 M HCl. How many grams of acetylsalicylic acid are in the sample?

What is the percentage of acetylsalicylic acid (or the purity)?

Solution: First recognize that the 27.98 mL of base was used to neutralize all

acidic material present in the sample. Since we are only interested in the

quantity of acetylsalicylic acid, we must determine the quantity of base requiredfor hydrolysis of the ester. The total number of moles of base added for the

hydrolysis reaction is

moles NaOH = 0.04278 L x 0.1000 M = 4.278 x 10 -3 mol

The number of moles of HCl used in the titration corresponds to the excess

NaOH, or the number of moles not consumed in the hydrolysis reaction:

mol HCl = mol excess NaOH

= 0.1056 M x 0.01429 L

= 1.509 x 10-3 mol

The difference between the number of moles of base added for the hydrolyses

and those which were not consumed equals the number of moles of base that

brought about hydrolysis. This is exactly equal to the number of moles of

acetylsalicylate ion which is equal to the number of moles of acetylsalicylic

acid:

4.278 x 10-3 - 1.509 x 10-3 = 2.769 x 10-3 mol

The number of grams acetylsalicylic acid is found using the molecular weight of

acetylsalicylic acid:

grams = 2.769 x 10-3 mol x 180.2 g/mol

= 0.4989 g acetylsalicylic acid

Thus,

% purity = 0.4989 g/0.5130 g x 100 = 97.25 %

4. Procedure

1. Preparation: Prepare an ice bath. Obtain 75 mL ethyl alcohol (EtOH) in a clean and dry 125-

mL Erlenmeyer flask. Cool the EtOH in the ice bath.

Check out two 25-mL burets from the stockroom window and prepare them for use as describedin “Volumetric Analysis”. Fill one buret with acid (HCl) and one with base (NaOH). Once the

burets have been filled, be sure to record the concentrations of the acid and base in your

laboratory notebook

Verify that you have at least 1.5 g recrystallized aspirin from the aspirin synthesis experiment.

(It is not necessary to weigh anything at this point - just check your final mass from last week’s

synthesis procedure.) If not, obtain a sample of acetylsalicylic acid from the stockroom window.

Prepare a hot water bath by bringing approximately 350 mL water to a boil in a 600 mL beaker

placed on a hot plate. (No open flames in the presence of ethanol!)




2. Sample preparation: Weigh about 0.5 g recrystallized aspirin. (Record mass to the nearest

0.0001 g.) Place in a clean, towel-dried 250-mL Erlenmeyer flask. Add 25 mL chilled EtOH.

Swirl to dissolve. Add 3 drops phenophthalein.

3. Acid Titration: Rapidly titrate with standard (≈0.2 M) NaOH, being sure to record starting and

finishing volumes. Remember that the buret reading can be estimated to + 0.01 mL. The first

10 mL can be added quickly, followed by slow additions of NaOH until a faint pink color

persists. (This indicates the phenolphthalein end point.) The volume of NaOH used for this

titration corresponds to that which is requied to neutralize all acids present in your sample, that

is, impurities as well as the acetylsalicylic acid.

4. Saponification of Aspirin: To saponify (or hydrolyze) the aspirin, you will add additional

NaOH from the buret, keeping careful record of the quanitity used. The quantity of base to be

used is about 8.5 mL more than the volume of base used in the previous titration. It will

probably be necessary to add more NaOH to your buret. Record the initial volume, add the base

to the Erlenmeyer flask, being careful to stop before reaching the lowest graduation mark (25.00

mL) on the buret. Record the final volume. The volume of base added is the difference

between these two readings.

Heat the basic solution in the Erlenmeyer flask in the water bath for 15 min, swirling

occasionally. (Use this waiting period to prep and titrate trials 2 and 3.) If the pink color should

disappear after this time, add 2 more drops of indicator. After the heating period, cool the

solution on an ice bath.

5. Back-titration of excess base: Record the initial volume of HCl in the second buret. Back-

titrate the excess NaOH in the Erlenmeyer flask with the standard (≈0.025 M) HCl solution until

the pink color disappears. Record the final volume.

6. Repeat steps 2 - 5 two more times.

5. Calculations

Calculate the grams of acetylsalicylic acid in each of your aspirin samples and the percentage purity

of the aspirin samples (see example). Calculate the mean percentage purity and the standard

deviation.

6. Discussion / Conclusion

Report the mean percentage purity in your aspirin sample. What are the sources of impurity in

aspirin? Comment on why and how titrimetric analysis is used and why it was necessary to perform

a back titration to determine the percentage purity. What are the major sources of error when

performing a titration?

7. References

John H. Nelson and Kenneth C. Kemp, Chemistry: The Central Science, Laboratory

Experiments, 7/e, Prentice-Hall, 1997.

ASP Analysis

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