CH229 General Chemistry Laboratory
Dr. Deborah Exton
VOLUMETRIC ANALYSIS OF ASPIRIN
1. Purpose
In this laboratory experiment, you will assess the purity of a sample of aspirin and become
acquainted 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” you
should 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 acid
impurities 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 is
not 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
CH229 Volumetric Analysis of Aspirin
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 required
for 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 described
in “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!)
CH229 Volumetric Analysis of Aspirin
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. Backtitrate
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.
موضوع: بحث كامل عن الكيمياء التحليلية للاسبرين بالغة الانكليزية 
الجمعة أغسطس 07, 2009 4:10 am
