2.1: Separation of an Unknown Mixture by Acid/Base Extraction

Experimental Aims: The objective of this exercise is to separate a two-component mixture using extraction techniques and then to identify the isolated components by determining their melting points.

Experimental learning objectives: At the end of this experiment you should be able to:

1. use a separatory/dropping funnel;
2. dry an organic liquid;
3. use a rotary evaporator;
4. identify the organic phase in an immiscible organic/aqueous mixture;
5. use acid/base reactions to impact the solubility of organic compounds and
6. determine melting points.

Each student will be given a mixture of two substances, which belong to two of the three categories listed below.

Possible carboxylic acids	benzoic acid	2-chlorobenzoic acid
Possible phenols	4-tert-butylphenol	2-naphthol
Possible neutrals	1,4-dimethoxybenzene	fluorene

Background Reading: Besides reading the lab manual you will need to do a little bit more. To help you understand the chemical basis of this exercise, you should review Sections 3.5 - 3.7 in Solomons & Fryhl which concerns the properties of acids and their conjugate bases. Pay particular attention to the use of pK_a values. You should also review the appropriate pages in the Mohrig and pay keen attention during your lab talk to acquaint yourself with extraction, washing , drying agents, and recrystallization.

Background Information: Extraction is a particularly useful means of separating organic compounds if one compound in the mixture can be chemically converted to an ionic form. The ionic form is soluble in an aqueous layer and can be extracted into it. Other, non-ionic organic compounds in the mixture will remain dissolved in the organic solvent layer. Separation of the two layers results in the separation of the two compounds.

The extent to which an acid-base reaction proceeds to completion depends upon the relative acidity of the reactants and products. Reactions occur so that stronger acids and bases are converted into weaker conjugate base and conjugate acids, respectively. The pKa value of the acids gives a measure of the acidity of each compound. Stronger acids have smaller pKa values and their conjugate bases are weaker. The position of an acid-base equilibrium can then be predicted from knowledge of the pKa values of the acids involved.

Take a look at the following acid-base reactions in Figure 1, paying particular attention to the position of the equilibrium and its relationship to the pKa values given.

Figure 1:

The reactions of a carboxylic acid and a phenol with bicarbonate ion. Note that the carboxylic acid has a lower pKa than the conjugate acid of bicarbonate ion (carbonic acid). The reaction, therefore, proceeds to products. The reaction of a phenol, however, favors the reactants since the pKa of phenol (10) is larger than that of the carbonic acid (6.4). Acid-base reactions favor the side with the weaker acid (that is, they favor the side with the larger pKa). So, extracting a mixture of these two compounds with bicarbonate results in the ionization and extraction of a carboxylic acid in the presence of phenol thus separating the two compounds from one another.

Now, look at the reaction in Figure 2 where we use a stronger base to do the reaction:

Figure 2:

The reactions of a carboxylic acid and a phenol with hydroxide ion. Note that in both cases, the reactions favor the formation of products. Therefore, extracting with hydroxide ion would result in the ionization and extraction of both compounds at he same time.

A close look at these two figures indicates that separating a mixture of a carboxylic acid and a phenol would best be done using bicarbonate ion since only the carboxylic acid is converted into its conjugate base by bicarbonate. The conjugate base of the carboxylic acid, being an ionic species, is soluble in the aqueous layer while the phenol (left unionized) would remain dissolved in the organic layer. However, if we were to extract with hydroxide ion, both the carboxylic acid and the phenol would be converted into their conjugate bases (see figure2). The conjugate bases, again are both ionic species and therefore soluble in the aqueous layer. This means that both compounds would be extracted at the same time, resulting in no separation.

A neutral compound will not react with either bicarbonate ion or hydroxide ion since a neutral compound does not have protons acidic enough to be removed by these bases. Therefore, such a compound will remain dissolved in the organic layer, no matter which base is added. For example, a mixture of neutral compound and a carboxylic acid can be separated using bicarbonate ion since only carboxylic acid will be ionized by the bicarbonate ion.

Once extracted, the carboxylic acid and phenol can both be recovered by adding \(\ce{HCl}\) to the aqueous solutions. The carboxylate ion and phenoxide will both be protonated by \(\ce{HCl}\), resulting in the formation of the original carboxylic acid and phenol, neither of which is soluble in water so they precipitate from solution. The solid can then be isolated by filtration. Figure 3 shows this chemistry for you.

Figure 3:

The reactions of a carboxylate ion and a phenoxide ion with \(\ce{HCl}\). Since \(\ce{HCl}\) is stronger acid than either of the conjugate acids, the products are favored in both cases. The products, a carboxylic acid and a phenol, are insoluble in aqueous solutions and precipitate from solution. The resulting solids can be isolated and their melting points determined.

The procedure you will use in this exercise exploits the difference in acidity and solubility just described.

1. you will dissolve your unknown in ethyl acetate (an organic solvent). All of the possible compounds are soluble in ethyl acetate.
2. you will extract with sodium bicarbonate to remove any carboxylic acid that is present.
3. you will extract with sodium hydroxide to remove any phenol that is present.
4. you will acidify both of the resulting aqueous solutions to cause any compounds that were extracted to precipitate.

These solids will be isolated by vacuum filtration, dried, and then their melting point ranges determined to identify them. If a neutral compound was present in your unknown, it will remain in the organic layer throughout the extraction procedure. To isolate it, you will simply evaporate the ethyl acetate to leave a solid. The melting point ranges of all solids will be determined. You will also weigh each solid you obtain to determine the percent recovery of your procedure. Remember, though, that you only have two compounds in your unknown mixture so that you should not isolate solids from all of the extracts.

Extraction procedure

General Notes

To measure the small volumes called for in this procedure, it is convenient to measure them in a graduated measuring cylinder. Make sure you label everything so that you can identify which layer you are putting into each flask correctly - label one 125 mL Erlenmeyer flask "bicarbonate", a second one as "hydroxide", and a 50 mL Erlenmeyer flask "ethyl acetate". We are using ethyl acetate in this lab, so avoid excessive exposure.

Be sure you are familiar with the procedure below before starting the lab.

1. Collect an unknown and record the unknown number. Without this number, we cannot grade your report. Label three Erlenmeyer flask as directed above.
2. Dissolve approximately 1.0 g of your unknown mixture in 10 mL of ethyl acetate.
3. Pour the solution into a clean separatory funnel and add 10 mL of 10% aqueous sodium bicarbonate found on your bench.
4. Stopper the funnel and invert it. Slowly open the stopcock to release any built up pressure, then close the stopcock (Figure 4).
5. Gently shake the separatory funnel to allow intimate mixing of the solutions and effect extraction of the compound from the organic mixture. (Caution: When shaken, the mixture may develop pressure; be sure to vent it periodically).
6. Clamp the separatory funnel to a retort stand and allow the mixture to separate into two layers (Figure 5).
7. Remove the stopper and collect the aqueous layer (the lower layer) in the 125 mL Erlenmeyer flask labeled "bicarbonate".
8. Repeat steps 3-7 two more times draining each portion successively into the same flask. At the end of this sequence you will have extracted the organic solution with three 10 mL portions of 10% aqueous sodium bicarbonate.
9. Put the Erlenmeyer flask labeled "bicarbonate" aside in a safe place. Later you will isolate any compound that was extracted by the bicarbonate. Do you remember which functional group that would be?
10. Add 10 mL of 5% aqueous \(\ce{NaOH}\) to the separatory funnel with the remaining ethyl acetate.
11. Stopper the funnel and invert it. Slowly open the stopcock to release any built up pressure, then close the stopcock.
12. Gently shake the separatory funnel to allow intimate mixing of the solutions and effect extraction of the compound from the organic mixture.
13. Clamp the separatory funnel to a retort stand and allow the mixture to separate into two layers.
14. Remove the stopper and collect the aqueous layer in the 125 mL Erlenmeyer flask labeled "hydroxide".
15. Repeat steps 10-14 two more times draining each portion successively into the same flask. At the end of this sequence you will have extracted the organic solution with three 10 mL portions of 5 % aqueous sodium hydroxide.
16. Put the Erlenmeyer flask labeled "hydroxide" aside in a safe place. Later, you will isolate any compound that was extracted by the hydroxide. Do you remember which functional group that would be?

The remaining steps described in this section will allow you to isolate any compound remained in the ethyl acetate layer. Recall, this would be a neutral compound, if you have one.

17. Add 5 mL of saturated aqueous \(\ce{NaCl}\) and 5 mL of distilled \(\ce{H2O}\) to the ethyl acetate layer in the separatory funnel.
18. Separate and set aside the lower, aqueous layer.
19. Pour the organic layer in the 50 mL Erlenmeyer flask and dry with anhydrous \(\ce{Na2SO4}\).
20. Filter the dried organic solution into a dry pre-weighed 50 mL round bottom flask and remove the ethyl acetate on a rotary evaporator. If a solid remains after evaporation of the ethyl acetate, it is a neutral substances and you will determine its weight and melting point.

Instructions follow for isolating the carboxylic acid and / or phenol from aqueous layers you put into the Erlenmeyer flasks labeled "bicarbonate" and "hydroxide", respectively.

1. Take the Erlenmeyer flask labeled "bicarbonate" and carefully acidify the aqueous solution by the dropwise addition of 6M \(\ce{HCl}\).

   Caution

   The bicarbonate solution will vigorously liberate carbon dioxide when neutralized with \(\ce{HCl}\) - that is, it will bubble a lot.

   Check to make sure the solution is acidic with blue litmus paper.
2. If a solid precipitates, add a boiling stone and then gently heat the solution to bring most of the solid back into solution. Cool slowly to room temperature and then use an ice/water bath to complete the precipitation. If no solid precipitates, your unknown did not contain a carboxylic acid. In that case, skip steps 3-4.
3. When the solution is ice cold, isolate the solid precipitate by suction filtration.
4. Filter, rinse the solid with ice-cold water, and determine the weight and melting point range of the carboxylic acid.

Now, we will follow the same procedure to isolate the phenol from the Erlenmeyer flask labeled "hydroxide".

1. Take the Erlenmeyer flask labeled "hydroxide" and carefully acidify the aqueous solution in the centrifuge tube by the dropwise addition of 6M \(\ce{HCl}\). Check to make sure the solution is acidic with blue litmus paper.  

   Caution

   The hydroxide solution will become hot when neutralized with \(\ce{HCl}\).

2. If a solid precipitates, add a boiling stone and then gently heat the solution to bring most of the solid back into solution. Cool slowly to room temperature and then use an ice/water bath to complete the precipitation. If no solid precipitates, your unknown did not contain a phenol. In that case, skip steps 3-4.
3. When the solution is ice cold, isolate the solid precipitate by suction filtration.
4. Filter, rinse the solid with ice cold water, and determine the weight and melting point range of the phenol next week.

Safety Notes

You must wear eye protection at all times. In the event that any reagent used in this investigation comes in contact with your skin or eyes, wash the affected area immediately with lots of water. Notify your instructor.
Avoid excessive exposure to all organic solvents.
Acids and bases can cause severe burns.
No flames should be present in the laboratory during this experiment.

Acid/Base Extraction Flow chart

References

L. M. Harwood and C. 1. Moody, Experimental Organic Chemistry- Principles and Practice, Blackwell Scientific Publications.

C.A. MacKenzie, Experimental Organic Chemistry, Prentice-Hall. 4th Edition

J. A. Moore and D. L. Dalrymple, Experimental Methods in Organic Chemistry, Saunders Golden Sunburst Series, W. B. Saunders Company.

C. F. Wilcox and M. F. Wilcox, Experimental Organic Chemistry- A Small-scale Approach, Prentice-Hall. 2nd Edition.

O. R. Rodig, C. E. Bell Jr. and A. K. Clark, Organic chemistry Laboratory- Standard and Microscale Experiments, Saunders College Publishing.

J.R. Mohrig, C.N. Hammond and P.F. Schatz, Techniques in Organic Chemistry, Freeman Publishers, 2nd Edition.