Last Updated 02 Aug 2020

Aspirin Experiment

Category Aspirin, Chemistry, Water
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Synthesis of Aspirin and Oil of Wintergreen INTRODUCTION: Synthesis and use of organic compounds is an extremely important area of modern chemistry. Approximately half of all chemists work with organic chemicals. In everyday life, many if not most of the chemicals you come in contact with are organic chemicals. Examples include drugs, synthetic fabrics, paints, plastics, etc. Synthesis of Aspirin and Methyl Salicylate. The two compounds we will be preparing, aspirin (acetylsalicylic acid) and oil of wintergreen (methyl salicylate), are both organic esters.

An ester is a compound that is formed when an acid (containing the COOH group) reacts with an alcohol (a compound containing an -OH group). O C R1 O H O + H O C R2 R1 O R2 + H O H acid alcohol ester water Here R1 and R 2 represent groups such as CH3 - or CH3 CH2 -. The reaction type shown above may be called a condensation reaction because the small molecule H 2 O is eliminated from the reactants while the remaining bits of the reactant condense together to give the main product. This reaction may also be called an esterification, since the product of the reaction is an ster, a compound containing the CO2 R group (see chapter 11 for definitions of acids, esters, and alcohols). Esters usually have pleasant, fruit-like odors and are the chemicals responsible for the odors and flavors of many fruits (oranges, bananas, pineapples) and flowers. In most cases, such natural products get their properties from a mixture of organic compounds. In this experiment you will prepare two esters of o-hydroxybenzoic acid, more commonly known as salicylic acid. One of the esters, acetylsalicylic acid, is aspirin, the common analgesic. We will synthesize aspirin by mixing salicylic acid with acetic anhydride.

The second ester product is oil of wintergreen, or methyl salicylate, which we prepare by allowing salicylic acid to react with methyl alcohol. This compound, which has a familiar odor is used as a flavoring agent and in rubbing ointments. Both of these reactions are shown below. Preparation of acetylsalicylic acid H H C C H C H C C C C O O H O H + O H3C H3C C O C O H H+ H H C C C H C C C C O O O C O CH3 + O C H3C O H H salicylic acid acetic anhydride acetylsalicylic acid (aspirin) acetic acid Preparation of methyl salicylate H H C C H C H C C C C O O H O H + H O H H CH3 H+ H C C C C C C H

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O H + O H H O C O CH3 salicylic acid methanol methyl salicylate (oil of wintergreen) water This experiment illustrates several properties of organic synthesis. While both product compounds in the experiment are esters of the same compound (salicylic acid), they are quite different in structure. Aspirin involves a reaction of the -OH group of salicylic acid, while methyl salicylate involves a reaction of the COOH group of salicylic acid. Organic chemistry is the broad field of studying the tremendous variety of such reactions of organic functional groups. Purification by Recrystallization.

After preparing the aspirin, we will purify it. Most solids can be purified by recrystallization, at a cost of lower percent yield. Recrystallization is usually done by dissolving the substance in a suitable solvent, which is hot. If insoluble particles are present, the hot solution is filtered to remove them (we will skip this hot filtration step). The solution is allowed to cool slowly, and is eventually cooled in ice. The crystals that form slowly are more pure than the original solid. Characterization by Melting Point. A simple characterization technique that can be very useful in determining purity is melting point.

It does not, however, tell much about the identities of the impurities. Pure materials usually have characteristic temperatures at which they melt, or a narrow temperature range (less than one degree) over which they melt. Impure compounds usually melt at a lower temperature, over a wider range. HAZARDS: Both acetic anhydride and sulfuric acid are reactive chemicals that can produce a serious burn on contact with the skin, and are irritating to the eyes. In case of contact with these chemicals, wash the skin thoroughly with soap and water. Do not dispose of any chemicals down the sink. Instead use the waste containers provided.

NOTE: The aspirin you will make is impure and must not be taken internally! LABORATORY OBSERVATIONS AND DATA: Be sure you make plenty of good qualitative observations, noting initial colors, odors, etc. , and any changes that occur during the experiment. Clearly label all numerical data. We will spend a little over one class day on this lab. In the first day, you need to prepare and recrystallize your aspirin. If time permits, you can also prepare the methyl salicylate. If not, this can wait until the second day when you will also take melting points of your crude and purified aspirin samples.

PROCEDURE: Synthesis of aspirin. Weigh out approximately 2. 1 g of salicylic acid (record exact mass), and transfer it to a clean, dry 6 inch test tube. Use the dispenser to carefully add 3 mL of acetic anhydride (density = 1. 08 g/mL) to the salicylic acid. Then add 3 drops of concentrated sulfuric acid, H2 SO4 , to the reaction mixture (it acts as a catalyst and speeds up the reaction). Put the test tube in a beaker of boiling water in a hood and heat for five minutes after most of the solid has dissolved. Stir the mixture with a glass rod to break up any lumps.

Pour the contents of the test tube into a 50 mL Erlenmeyer flask containing 25 mL of water. Swirl the flask for a few minutes to mix the solutions and get rid of any unreacted acetic anhydride. (The acetic anhydride reacts with water to produce acetic acid. ) Place the flask in an ice bath and watch for a white solid to crystallize out. Occasionally a reaction will yield an oily product that resists crystallization. If that happens, scratch the bottom and sides of the flask with a glass stir rod to help start crystal formation, or warm the mixture just until the oil dissolves, and then re-cool.

Allow 10 minutes for crystallization to occur. Meanwhile put a wash bottle of distilled water in some ice and prepare a Buchner funnel. Filter the solid, being sure to use a trap flask between the Buchner funnel flask and the aspirator. Wash the solid with a small amount of cold distilled water. Discard the liquid filtrate in the designated waste container. Pre-weigh an empty watch glass, then scrape your aspirin product off the filter paper onto the watch glass. Record this yield of crude aspirin. Use a bit of the solid product to pack a melting point capillary tube to use the second week to find the melting point of your crude product.

Recrystallization of the crude aspirin to form pure aspirin Put 15-20 mL of distilled water on a hot plate to begin warming. Dissolve your crude aspirin product in about 5 mL of 95% ethanol in a 50 mL Erlenmeyer flask. If some of your aspirin fails to dissolve, do the following: Prepare a warm water bath by using a beaker of water (about 50 mL in a 250 mL beaker), using a hot plate to heat the water bath. DO NOT USE A FLAME OR BUNSEN BURNER THE ETHANOL IS FLAMMABLE. Warm the Erlenmeyer flask containing the aspirin and ethanol in the warm water bath.

When the aspirin has dissolved, add 15 mL of warm distilled water (50 o C approximately). If any crystals form at this point, reheat the mixture in the water bath to re-dissolve them. Let the solution cool slowly, with the mouth of the flask covered by a watch glass. When it is at room temperature, place it into the ice bath and leave it there a full ten minutes. After crystallization is complete, filter the crystals in a Buchner funnel, wash them with a little ice cold distilled water (put your squeeze bottle in the ice), and suction for several minutes. Discard the liquid filtrate in the designated waste container.

Scrape the solid into a pre-weighed beaker and put in your drawer, lightly covered with a tissue. Do not cover it tightly because we want your product to finish drying until the next class period. You will need to get a final mass of this purified aspirin after allowing it to dry for a day. You will also take a melting point of the purified aspirin. When finished with the experiment, place your aspirin product in a test tube and stopper it. Label the test tube with your name, the name of the compound, and the date. Your instructor will collect this product. Synthesis of methyl salicylate.

Place 0. 5 g of salicylic acid and 3 mL of methyl alcohol in a large test tube. Add 2 drops of concentrated sulfuric acid and then place the test tube in the hood in a water bath at 70 o C for 15 minutes. The boiling point of methyl alcohol is 64. 6 o C, so point the mouth of the tube away from others and avoid overheating, to minimize "bumping". Note the odor before and after heating. Allow your methyl salicylate to cool to room temperature, then stopper the test tube. Add a label with your name, the compound name, and the date. Your instructor will collect this product.

Determination of the Melting Point of your aspirin. You should already have a capillary tube packed with your impure aspirin. Now pack a tube with your pure aspirin. Put your two tubes in the melting point apparatus and slowly heat your samples. Record the temperature at which each starts to melt and the temperature at which it has all melted. (Your instructor will give you more instruction on these procedures. ) RESULTS: Calculate the theoretical yield of aspirin from the balanced equation given in the introduction. You will need to add up molar masses by counting the atoms shown in the structures.

Be sure to determine the limiting reagent, either salicylic acid or acetic anhydride. You will need to use acetic anhydride’s density, 1. 08 g/mL. Show this work clearly in your report. Watch significant figures and units. Also report the masses of your crude and purified aspirin samples, and the percent yield of your final, purified aspirin. DISCUSSION: Describe what your melting points say about the purity of your initial crude product and your recrystallized product. Pure aspirin has a melting point of 135o , while salicylic acid has a melting point of 157-159o .

Impure compounds normally have lower melting points and broader melting ranges than if pure, even if the impurity would have a higher melting point itself. List all the compounds that could be mixed in with your aspirin product as impurities, i. e. all reactants, solvents, and other products. You should have six compounds besides aspirin. Briefly describe the source of each compound. Considering factors such as limiting reagent (which you just calculated), and procedural steps which may have removed some of these compounds, which compound(s) do you think are most likely to contaminate your aspirin product? Explain

QUESTIONS: 1. Infrared spectra are often used to get a quick look at the purity of a product. IR spectra of aspirin, as well as the salicylic acid and acetic anhydride used to prepare it Selected Infrared Frequencies are shown below. Consult the structures of these three Absorption Range, cm-1 Bond Type compounds from the introduction. Recall from our experiment 3600-3200 (usually broad) on IR that different bonds show up at different frequencies in O - H the IR spectrum. Prepare a listing for each compound, C - H 3300-2800 showing bonds and their IR frequencies taken from the B - H 2650-2300 spectra below.

You need consider only those bonds listed in C ? N 2260-2220 the table at right. (It is hard to read the frequencies on the attached spectra accurately, so just approximate. The frequency values listed on the x-axis in the figures are 4000, 3000, 2000, 1500, and 1000 cm-1. ) 2. Look at the IR spectra again. Assume that you ran an IR spectrum of your aspirin, but that it was contaminated with unreacted salicylic acid. At what frequency in the spectrum would you look for evidence of this contamination? Explain your reasoning. IR Spectrum of Aspirin (KBr pellet)

IR Spectrum of Salicyclic Acid (KBr pellet) IR Spectrum of Acetic Anhydride (liquid thin film) 3. 1 H NMR spectra of salicyclic acid and aspirin are shown below. Note that both have small, broad peaks near 11 ppm that were artificially enhanced to make them obvious on these spectra. The peaks around 7-8 ppm are all doublets, triplets, or messy multiplets; the other peaks are singlets. How many types of H should be expected for each compound based upon their structures? The two structures share many common features, and thus their spectra should be similar for these common features.

Likewise, each structure has some unique type of hydrogen not found in the other. These should result in differences between the spectra. By simply comparing the structures and the spectra, decide which hydrogens on the structures give rise to which peaks. You won’t be able to assign all the peaks this way, but do as much as you can and explain your reasoning. 4. The Merck Index is an extremely valuable reference that can be found in the Reference Section of the Fintel Library as well as several reading rooms on 5th floor Trexler.

It is an especially good place to find basic information on organic compounds. Look up aspirin in The Merck Index. Summarize the information it gives about the solubility and decomposition of aspirin. Record the edition number and where you found this book. 5. The CRC Handbook of Chemistry and Physics, often simply called “The CRC,” provides less information on each compound than The Merck Index does, but covers more inorganic compounds and includes hundreds of pages of other facts useful to chemists.

In the CRC find the multipage table entitled Physical Constants of Organic Compounds. Within this table, find salicylic acid (some old editions may list it as 2-hydroxybenzoic acid). Record the information given about salicylic acid’s melting point, boiling point, and solubility. You will probably need to consult the listing of Symbols and Abbreviations given one or two pages in front of this huge table. Record the edition number of the book and where you found it.

Aspirin Experiment essay

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