CR Scientific: Chemistry Experiments: Salicylic Acid and Salicylates

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CAUTION: The following experiment involves strong acids and/or strong bases and may also involve flammable solvents. If you choose to proceed, you do so entirely at your own risk. To use this website you must read and agree to the Terms of Use.

Aspirin & Salicylic Acid
Ester hydrolysis and preparation of an organic salt

Synopsis: Acetylsalicylic acid is hydrolyzed to salicylic acid (SA). The SA is reacted with NaOH to form sodium salicylate. The relevant chemical principles are discussed.

Introduction:

Acetylsalicylic acid (C₆H₄OCOCH₃COOH) is the ubiquitous pain reliever known as aspirin. Chemically, aspirin is both an aromatic acid and an ester; analgesic properties notwithstanding, the compound is an excellent showcase for at least two or three principles of organic chemistry.

Salicylic acid (C₆H₄OHCOOH; aspirin's parent compound) and sodium salicylate (C₆H₄OHCOONa) are useful for detecting iron compounds; they form colored complexes with the Fe³⁺ ion. They can also be used in several organic chemistry experiments which illustrate important principles.

For convenience we may abbreviate acetic acid as HOAc or AcOH. This refers to an acetate group ("OAc"; Figure 1) with a hydrogen ("H"). Acetic acid loses an H⁺ when it ionizes to acetate, which is then abbreviated as OAc^-.

Because the shorthand system becomes ambiguous with more complex molecules, we'll refer to salicylic acid as "SA" and acetylsalicylic acid as "ASA". (Consider that HOSal could be salicylic acid, if we assume the dissociable H+ comes from the carboxylic acid group.).

Figure 1. The difference between an acetyl group and an acetate group. In chemical shorthand, these are written as Ac and AcO (or OAc), respectively. Ethyl acetate would be EtOAc, while acetyl bromide would be AcBr. Though organic shorthand uses Ac, which is also the chemical symbol for Actinium, the latter is not usually encountered in practical chemistry. "Ac" in an organic chem textbook is always "acetyl", not "actinium".

Salicylic acid is a natural analgesic present in the leaves and bark of certain plants. It is generally unsuitable for internal use, since it is a strong gastric irritant and can cause internal bleeding. In fact, aspirin was invented for this very reason; the acetylated molecule isn't as rough on the digestive tract, although it does hydrolyze to some degree in the stomach.

Safety:

Safety goggles MUST be worn at all times when doing the procedure or handling the compounds, even when just opening their containers. Gloves are also a good idea.

Salicylic acid is more corrosive than acetylsalicylic acid. It can cause burns. Do not get salicylic acid on the skin or in the eyes. Do not ingest salicylic acid or its salts (or, for that matter, any compounds in the chem lab).

Alcohol, acetone, and ether are highly flammable (in increasing order of fire hazard). There mustn't be any ignition sources nearby. Heating must be carried out with a sealed-element hot plate designed for lab use. The laboratory must be well-ventilated while the solvent is evaporating; it must at the same time, however, be inaccessible to children and unauthorized people.

Hydrochloric acid, especially when it's strong or concentrated, should be handled in a fume cupboard or outdoors (the latter while wearing an approved respirator for HCl vapors). Be careful storing concentrated or strong HCl; vapor always seems to find its way through the container walls or seals and into the surrounding air. A plastic bottle of ammonia stored a few feet away is a good indicator for leaking HCl vapors: within a few weeks, there will eventually form a thin layer of NH₄Cl on nearby objects.

Solutions containing HCl must be kept in the fume hood during any heating that takes place.

Materials:

Acetylsalicylic acid ("ASA") (aspirin tablets)
Hydrochloric acid, 1M
Acetone OR isopropyl alcohol
Ethanol, 95%
Safety goggles
Filtering flask
Büchner funnel
Filter paper
Glass rod, spatula, or rubber policeman
Water aspirator or vacuum pump
Petri dish or Crystallizing dish
Distilled water
Beaker made of heat-resistant glass
Round-bottom flask
Support stand and clamps
Crushed ice
Diethyl ether (optional)
Separatory funnel (optional)

Methods & Observations:

I. Crystallizing Acetylsalicylic Acid

Fifteen or twenty aspirin tablets are crushed and covered with about 250 mL of acetone or isopropyl alcohol. The mixture is stirred and the insoluble matter allowed to settle. The clear liquid is decanted and filtered through a Buchner funnel on a filtering flask, to which is applied a mild vacuum such as that obtained from a hand pump or water aspirator.

This solution is allowed to evaporate slowly at room temperature in a shallow container such as a petri dish. It may take several days. Do not heat it at this point; in fact, keep it far away from any potential ignition sources such as space heaters. Tabular crystals of acetylsalicylic acid (ASA) should form on prolonged standing (Figure 2.)

Acetylsalicylic acid crystallizing from 2 propanol

Figure 2. Acetylsalicylic acid (ASA) crystals.

Look carefully at the crystals under magnification. Is there only one kind of crystal? Why or why not? How about the odor: is there a slight hint of vinegar? Minor hydrolysis is normal, but the bulk of the crystals should be ASA.

II. ASA to SA: Acid Hydrolysis

When all the liquid has evaporated, put the crystals into a 250-mL beaker and add 150 mL of water. Add enough HCl to adjust the pH to around 2. ( Instead of HCl and acidity, one can go the other way entirely; a solution of NaOH can be used to make the pH around 12. The hydrolysis reaction is catalyzed by either alkaline or acidic conditions. To keep it simple, use HCl.)

Put some crushed ice and water into a 250-mL boiling flask and, using a clamp and support stand, place this on top of the beaker (Fig. 3):

Primitive reflux setup using an ice flask

Figure 3. Hydrolysis of acetylsalicylic acid. The ice flask condenses most of the vapors and drips them back into the beaker. Don't forget to hold the ice flask in place with clamps and a ringstand.

Maintaining constant volume in the beaker keeps the pH stable (not that it's really that critical for this experiment). The ice flask's main benefit is re-capture of most of the HCl that evaporates.

Notice the magnetic stir bar in the bottom of the beaker. The speed doesn't need to be high. The stirrer is optional in this experiment but helps even the heat.

The beaker is shown resting directly on the hot plate, but a water bath would be better.

Gradually heat the beaker on a hot plate; 60-70 °C is sufficient, though in this experiment the liquid was heated to about 85 °C. Do not allow the liquid to boil. At first the crystals will not dissolve, but when the solution gets hot enough they will disappear. For best results, the hot plate should have a built-in magnetic stirrer to maintain even heating.

Keep the temperature between 60 and 85 °C for at least an hour. The ice in the round-bottom flask will melt, so periodically replace it with fresh ice. The ice flask re-condenses most of the vapors, maintaining a nearly constant volume in the beaker throughout the experiment.

III. Separating and Crystallizing the Salicylic Acid

The liquid may thicken somewhat. When the heat is turned off, the salicylic acid (SA) will come out of solution as a thick, white precipitate (Fig. 4). The HOAc freed from the hydrolysis will impart the characteristic odor to the solution, but don't inhale the vapors. Remember there's also hydrochloric acid in there.

Salicylic acid precipitates upon cooling

Figure 4. The precipitate can be chilled, filtered, and washed on a Büchner funnel with ice-cold distilled water. SA dissolves only sparingly in cold water.

A crust of salicylic acid may cling to the stirrer bar; this can be scraped off, or it can be heated to near boiling in 100 mL of water to dissolve the caking. The solution is then cooled to recrystallize the compound..

At this stage one can allow more of the liquid to evaporate at room temperature over the course of a couple days. The salicylic acid can then be washed with ice-cold distilled water on a Buchner funnel. Save the washed SA and allow it to dry completely.

An alternate method might be to extract the SA from the water-HOAc solution using a separatory funnel and some diethyl ether. The SA goes into the ether layer, which separates on top. The only problem is that HOAc is about as soluble in ether as it is in water. An ether/water separation thus takes care of most of the HCl, but only about half the HOAc.

As the ether solution is evaporated (make sure your ether has peroxide inhibitors in it!), acicular crystals of SA will form (fig. 5). The amateur-scientist reader might try engine starter if pure diethyl ether is not available, though the starter may contain petroleum-based lubricants that can impact crystal morphology. Regardless, be very careful with ether; take extra care to avoid any possible sources of ignition.

Figure 5. Salicylic acid crystals from diethyl ether. About 3 mL of solution was allowed to evaporate in a micro beaker.

The rest of the salicylic acid was saved for suction filtration and washing with chilled water.

We can test our product to see whether it contains salicylic acid, as well as to determine whether it's pure. A solution of a ferric compound (e.g. FeCl₃) will cause a reddish-purple color if there is salicylic acid present.

Salicylic acid melts at about 158°C; ASA melts around 135°C. We can best perform the melting point determination by placing some crystals into a capillary tube which has been melted shut at one end, then fastening this tube to the side of a thermometer which is immersed in mineral oil. The oil is heated slowly and the temperature noted. Does the melting take place at one temperature (or at least a very narrow range), or does it continue through a wide range of temperature? A poorly-demarcated melting point tends to indicate an impure substance.

IV. Sodium Salicylate

HOSal + NaOH ----> NaOSal + H₂O

One mole of salicylic acid reacts with one mole of sodium hydroxide in aqueous medium to form a mole of sodium salicylate and a mole of water.

Salicylic acid: molecular weight 138.12
Sodium hydroxide: molecular weight 40.01

The molar mass of NaOH is approximately 0.290 times that of salicylic acid; thus, 10.00 grams of salicylic acid would be reacted with 2.90 grams of NaOH. It is important to use a 1:1 mole ratio so there is no excess NaOH or salicylic acid in the final product.

Set a few grams of the SA aside for future experiments (such as preparation of sulfosalicylic acid, used in determination of proteins); put it in a clean, sealed container and label it. Weigh the remainder of the SA to be used in the present experiment. Whatever this weight is, multiply this by 0.290 to give the molar equivalent weight of NaOH that will be needed to react with the SA. This is because NaOH has a formula weight of 29.0% that of salicylic acid.

Dissolve the NaOH in as little distilled water as required to solubilize it completely. Dissolve the SA in another beaker, adding a little ethanol if necessary to dissolve it completely.

Mix the two solutions and let stand for at least an hour. Heat carefully at just below the boiling point to evaporate the liquid down until it thickens and becomes pale yellow. Continue heating at 65 to 80 °C, stirring periodically, until crystals precipitate out and eventually become completely dry.

The crystals are sodium salicylate. If the reactants were weighed accurately, there should be no excess NaOH or salicylic acid.

Compare these crystals under the microscope with crystals of SA. Are there any noticeable differences, or would these two compounds be virtually indistinguishable? How about the solubilities?

Save the sodium salicylate in a labeled container. It will be useful in future experiments.

Figure 6. Sodium salicylate after evaporation of all the liquid at 65°C.
Notice the color: it's a hard-to-place shade of gray with a hint of tan, and it has a vaguely pearly or soapy luster. Contrast this with the nondescript white of salicylic acid.

Discussion:

Acetylsalicylic acid is an ester formed by reaction of the -OH group of salicylic acid with the -COOH group of acetic acid. The more "acid" part of salicylic acid (in other words, its own -COOH group) is not part of the ester reaction, even though its presence may cause confusion to students. In other words, only the alcohol group of salicylic acid participates. (Recall that organic acid + alcohol = ester + water. HOAc + EtOH, for example, gives EtOAc + H₂O). This alcohol group is also called a "phenolic" group, since it is attached to an aromatic (benzene) ring. (technically speaking, phenol is both an acid and an alcohol).

The reaction can go either way depending on conditions. Hydrolysis of acetylsalicylic acid is an equilibrium reaction catalyzed by acid (as in this case) or by base. Normally the hydrolysis is driven to the right by low pH, as well as by losses of HOAc to the air. Thus, the top reaction in Figure 6 shows only a one-way arrow. One must remember that equilibrium means equal rates, not equal concentrations; at equilibrium, this reaction favors a high concentration of SA and a very small one of ASA.

Organic reactions in general are sensitive to a number of conditions. The bottom reaction (Fig. 7) shows what would happen if strong or concentrated sulfuric acid were used instead of dilute HCl. The H₂SO₄ would break the ester bond as a matter of course, but it wouldn't stop there. The net product would be mostly phenolsulfonic acid, since H₂SO₄ is a powerful sulfonating agent. Moreover, it can decarboxylate aromatic rings via electrophilic attack; the leaving group is CO₂ (March, 1992). When concentrated sulfuric acid gets hot enough, it actually turns organic compounds into CO₂ and H₂O (this is why boiling, concentrated H₂SO₄ is so dangerous-- organic matter is charred instantly).

Figure 7. The choice of acid makes a remarkable difference. Dilute HCl produces salicylic acid and acetic acid; concentrated sulfuric acid yields mostly phenolsulfonic acid and CO₂.

In an aqueous environment made pH 2.0 with HCl, de-esterification should be the only reaction of concern. As noted above, concentrated sulfuric would lead to a different set of reactions. A dry or almost-dry reaction, such as heating the solid ASA with alkali, would give still a different outcome.

The reaction of salicylic acid with sodium hydroxide illustrates another important concept in chemistry: an organic molecule having two, acidic functional groups will react with alkali in a manner dependent on the pK_a values of those functional groups. Specifically, how can one be sure the alkali will attack only the carboxylic acid group of SA and not the phenolic group? This is actually quite simple: the pK_a of the carboxylic acid group is only about 3, but the pK_a of the phenolic group is 13.4. Unless an excess of NaOH is used at pH 14, the phenolic group of salicylic acid will not lose its proton. In other words, an NaOH solution of e.g. pH 10 will not remove H⁺ from the phenolic OH group of salicylic acid. On the other hand, the H+ will be dissociated from the COOH group. This H⁺ will react with the OH^- from NaOH to give H₂O; the Na⁺ will react with the COO^-, forming sodium salicylate. One might predict a very strong solution of NaOH (e.g., pH 14) to deprotonate both the carboxylic acid and the phenolic groups, giving disodium salicylate. For example, an early article on "Acidum Salicylicum" in King's American Dispensatory (1898) mentions the compound. However, it so happens that strong enough NaOH will also decarboxylate the aromatic ring, leaving the sodium salt of phenol.

Alkali and heat will decarboxylate both salicylic acid and sodium salicylate, the latter more readily (even without alkali). In other words, the carboxylic acid or carboxylate group is removed completely from the aromatic ring; this is similar to what H₂SO₄ does, except that no sulfonate group attaches to the ring. This base-mediated decarboxylation would therefore leave phenol [C₆H₅OH, sometimes written shorthand as PhOH] and probably a tarry mess of polymerization products; this is potentially the subject of another experiment. Depending on how much alkali were present, we could instead get sodium phenoxide [C₆H₅ONa, sometimes written as NaOPh or PhONa].

This has been a fairly simple investigation into the chemistry of acetylsalicylic acid. This experiment has shown only a couple of the potential reactions in which this interesting and useful molecule can participate.

References:

Felter, H. and Lloyd, J. King's American Dispensatory. 1898. On-line version at Henriette's Herbal.

March, J. Advanced Organic Chemistry. New York: Wiley, 1992.

Merck Index, 10th Edition. Rahway, New Jersey: Merck and Company, Inc., 1983.

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