Analysis of Aspirin1

Frederick R. WightManatee Community College

Introduction

Acetylsalicyclic acid, aspirin, was synthesized from salicylic acid in CHM2045L by the reaction in Figure 1:Figure 1 Aspirin synthesis:

Aspirin is not stable indefinitely and can, over time, react with moisture to revert back to salicylic acid by a process known as hydrolysis, Figure 2.

Figure 2: Aspirin Hydrolysis

In this laboratory we will analyze lab synthesized and commercial aspirin for purity by both qualitative and quantitative methods.

Qualitatively, the purity of an aspirin sample can be determined from its melting point. Qualitative chemical analysis is concerned with identifying the elements and compounds present in a sample of matter. The melting point of a substance is essentially independent of atmospheric pressure, but it is always lowered by the presence of impurities (recall colligative properties: freezing point depression and boiling point elevation). The degree of lowering of the melting point depends on the nature and concentration of the impurities. The term melting point is somewhat of a misnomer since substances rarely melt at a single temperature. Rather they melt over a range of temperatures and the smaller the temperature range the purer the substance. Note also that even though

1 Partly adapted from a procedure by Kerber, Akhtar & Schneider, SUNY Stony Brook. Parts of this lab have been taken from “Laboratory Manual for Principles of General Chemistry”, J.A. Beran

Aspirin Synthesis 1

salicylic acid has a higher melting point than aspirin, if the aspirin has a trace amount of salicylic acid the melting point of the mixture will be lower than pure aspirin.Qualitatively, it is also possible to detect the presence of salicylic acid using ferric chloride. Salicylic acid forms a colored complex with acidified iron (III) ion which is readily visible at low concentrations. Figure 3.

Figure 3: The addition of acidified iron (III) ion produces the violet tetraaquosalicylatroiron (III) complex.

Aspirin does not form this complex. This forms the basis for a simple qualitative test for the presence of unreacted salicylic acid in your aspirin.

Quantitatively, the purity of an aspirin sample can be determined by a simple acid-base titration. Both aspirin and salicylic acid are weak organic acids but with different properties, Table 1.

Table 1Compound: Salicylic Acid Aspirin

Formula: C7H6O3 C9H8O4

Molar Mass: 138.12 180.15Melting point: 158-160°C 140-142°C

Ka 1.08 x 10-3 2.72 x 10-5

pKa 2.99 4.57Solubility (g/100mL) 0.18 0.25

A standardized NaOH solution titrates the acid to the phenolphthalein endpoint, where

(1) Volume NaOH, L x Molarity of NaOH, mol/L = mol NaOH = moles acid

Note that salicylic acid has a lower molar mass. This means if we titrate the same weight of each substance it will take more NaOH to reach the endpoint with salicylic acid. Thus if our aspirin is contaminated with salicylic acid it will require more NaOH to reach the endpoint than if our sample were 100% pure aspirin. In fact, by titrating a known weight of our sample with a standard NaOH solution we can determine the “effective molar mass” (EMM).

EMM = grams of sample/moles acid from titration

Aspirin Synthesis 2

If the sample consists of ONLY salicylic acid and aspirin, then the percent composition can be derived from the EMM according to the relationship:

% Aspirin = 100(EMM-138.12)/(180.15-138.12) = 2.379(EMM-138.12)

If the EMM is greater than 180.15 or less than 138.12 then the formula does not apply.

Procedure

Part I1. Place a few crystals of pure aspirin, salicylic acid and the lab synthesized aspirin

in 3 separate test tubes.2. Add approximately 3-5 drops of 95% ethanol to dissolve the aspirin. Then add 1-2

ml of water to each sample. 3. Add 2 drops of Iron (III) Chloride solution to each and observe the color of each

sample. Record the observations in your notebook.

Part IIa1. Calculate the mass of pure aspirin that requires 20 mL of a 0.1 M NaOH to

reach the stoichiometric equivalence point. Show your instructor to verify2. Weigh out this mass of pure aspirin and place it into a 250 mL Erlenmeyer

flask.3. Add 10 mL of 95% ethanol (Make sure you add this before the water!!)4. Add 50 mL of DI water5. Add 2 drops of phenolphthalein indicator and swirl to dissolve the aspirin6. Fill a clean 50 mL buret with standardized 0.1 M NaOH7. Read the initial volume8. Slowly add the NaOH solution to the flask with the aspirin in it until the

endpoint is reached (very light pink color)9. Record the final buret reading along with the exact molarity of the base10. Repeat the titration.

Part IIb

1. Repeat the procedure above using the aspirin synthesized in CHM2045L.

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Data Sheet

Part I Observations on addition of Iron(III)chloride

Salicylic acid:____________________________________________________________

Pure aspirin:_____________________________________________________________

Lab synthesized aspirin:____________________________________________________

Part IIa (Pure Aspirin)Calculate the mass of aspirin needed to reach the stoichiometric equivalence point with 20 mL of a 0.1 M NaOH solution. (Show calculation below)

.

Trial 1 Trial 2

Mass of aspirin sample (g)

Molarity of NaOH (mol/L)

Initial buret reading (mL)

Final buret reading (mL)

Volume of NaOH used (mL)

Moles of NaOH used

Effective Molar Mass (g/mol) ____________ ____________

Percent purity of aspirin (%) (Show calculations below for 1st run)

Average % purity of aspirin (%)

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Part IIb (Lab Synthesized Aspirin)Calculate the mass of aspirin needed to reach the stoichiometric equivalence point with 20 mL of a 0.1 M NaOH solution. (Show calculation below)

Trial 1 Trial 2

Mass of aspirin sample (g)

Molarity of NaOH (mol/L)

Initial buret reading (mL)

Final buret reading (mL)

Volume of NaOH used (mL)

Moles of NaOH used

Effective Molar Mass (g/mol) ____________ ____________

Percent purity of aspirin (%) (Show calculations below for 1st run)

Average % purity of aspirin (%)

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Pre-Lab Questions

1. A 0.536 g sample of aspirin prepared in the laboratory is dissolved in 95% ethanol, diluted with water, and titrated to the phenolphthalein endpoint with 18.5 mL of a 0.135 M NaOH

a) How many moles of acetylsalicyclic acid are present in the sample?

b) Calculate the % purity of acetylsalicyclic acid in the aspirin sample

c) Determine the number of grams of acetylsalicyclic acid that will react with 25.00 mL of a 0.100 M NaOH solution

d) In the experiment 3.00 g of salicyclic acid reacts with an excess amount of acetic anhydride. Calculate the theoretical yield of acetylsalicyclic acid for this synthesis

2. Looking at the reaction used to prepare aspirin, can you suggest a reason why it might be possible to obtain an EMM of less than 138.12?

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