Full version Separation Of Acid, Base And Neutral Compounds From A Solid Mixture By Liquid-Liquid Extraction

Separation Of Acid, Base And Neutral Compounds From A Solid Mixture By Liquid-Liquid Extraction

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Experiment #1: Separation of Acid, Base, and Neutral Compounds from a Solid Mixture by Liquid-Liquid Extraction


Liquid-liquid extraction is a technique used to separate chemical substances in order to purify or identify the various components of a mixture. Flavors, spices, perfumes, and medicines are just some of the everyday things that are extracted from plants and other natural sources [1]. The basic principle used to carry out this separation is the mixing of two liquids that are immiscible with each other. This creates layers of liquid, which can be separated and then isolated to help to identify compounds. Another important concept to remember in this experiment is that ionic salts are polar and therefore water soluble, and neutral molecules are non-polar and will usually not dissolve in water.

The student is given an unknown containing either an acid and a neutral compound or a base and a neutral compound, but since the components are unknown, the student conducts the experiment as if all 3 compounds are present. In this particular experiment, the compounds are benzocaine, salicylic acid, and naphthalene. An appropriate organic solvent is used to create the ether phase. Diethyl ether is used because it has low solubility in water but will dissolve the components of the mixture (creating liquid layers), is not acidic or basic, is less dense than water, and can be easily removed by evaporation. To begin the separation, the student needs to produce a salt that is water-soluble, so that the ether phase can separate from an aqueous phase. To accomplish this, dilute HCl is added to the solution. The HCl protonates the base (benzocaine) in the solution to create an organic polar salt, and it therefore dissolves in water. This aqueous phase is drained from the ether phase and then deprotonated again with NaOH to reproduce the original neutral benzocaine.

The same concept is used to isolate the salicylic acid. Sodium hydroxide is added to deprotonate the salicylic acid to leave the conjugate base. This salt of the acid dissolves into the aqueous phase and is then drained from the funnel where it is then protonated again with HCl, leaving the naphthalene alone in the ether phase.

The student finishes with solid naphthalene (by rotary evaporation) and either solid salicylic acid or benzocaine. Solids are filtered off and weighed to find the percent recovery. Finally, melting points of each product are taken to confirm the identities and purities of the solids.

Experimental Procedure

As outlined in the CHEM 265L manual [2]. No deviations were made.

Results and Observations

Upon addition of HCl to diethyl ether and unknown, gas was produced and separatory funnel needed to be vented about 5 times. Funnel was shaken for 2 or 3 seconds and stopcock opened to release pressure each time. A clear line was observed between the ether and aqueous phases. The ether phase was a little cloudy, while the aqueous phase was clear. After draining the aqueous phase containing the salt of the base, some residue was observed floating on the top of the liquid in the beaker. This may have been contamination, or some diethyl ether that was accidentally released with the aqueous phase.

The addition of NaOH was similar to HCl where lots of gas was produced and the funnel needed to be vented 5-6 times before no gas was released from the solution. This time, the aqueous phase started out very cloudy and the ether phase was clear, but after letting the solutions sit for a minute or so and venting a few more times, the aqueous phase looked a lot clearer and could be extracted.

When adding the drying agent to the naphthalene solution, 1.05g was measured out and added. The sodium sulfate immediately clumped into the middle and stuck to the glass, while the solution was a little cloudy. An additional 0.8g sodium sulfate was added. After swirling the beaker, some sodium sulfate was floating around freely. The now clear solution was decanted into round bottom flask and rotary evaporated. White solid crystals were observed after 1 minute of rotation. The solution was completely evaporated after 3 minutes.

15 NaOH pellets were added to the salt of the base and pH paper was red (acidic), with no precipitate. 5 more pellets added and the pH paper turned blue (basic), but there was still no precipitate. With the salt of the acid, 2 full disposable pipette amounts of HCl were added and the pH paper turned blue (basic), with no precipitate. Another 1.5 pipettes added and an immediate white powdery precipitate observed. The pH then tested as acidic.

Unknown 306 – 1.02g Salicylic Acid Naphthalene

Obtained Melting Point 154В°C - 158В°C 76В°C - 79В°C

Amount Recovered .52g .45g

% Yield 101.96% 88.24%

Physical White solid, fine granules White solid, crystals

Figure 1.0

Sample Calculations:

Amount of salicylic acid recovered:

Final Round Bottom Flask Weight: 47.50g

- Initial Round Bottom Flask Weight: 47.05g

.45g salicylic acid

Amount of naphthalene recovered:

Final Watchglass Weight: 21.95g

-Initial Watchglass Weight: 21.43g

.52g naphthalene

The starting unknown weight was 1.02g; therefore, the maximum yield is .51g of each component, if the unknown was a 50/50 mixture.

% Yield of salicylic acid:

Actual = .52g x 100 = 101.96% yield salicylic acid

Theoretical .51g

% Yield of naphthalene:

Actual = .45g x 100 = 88.24% yield naphthalene

Theoretical .51g


The procedure of this experiment was accurately followed, and ultimately, the components of unknown 306 were found to be salicylic acid and naphthalene. Firstly, when NaOH pellets were added to the salt of the base, no precipitate formed in the solution even with the solution was basified. This was the first identification that the unknown contained salicylic acid. When the salt of the acid was acidified, a large amount of white powdery precipitate formed in the solution, thus confirming the presence of salicylic acid. The powder was filtered off and found to have a weight of .52g. This was one gram over the expected yield, and may have been caused by not drying for a long enough time in the suction filter. Although the precipitate looked dry, there still may have been traces of solution on the powder, giving a 101.96% yield for the salicylic acid compound. Ideally, the precipitate should be left for ~15 hours to dry and be filtered fully, but in this lab, time constraints do not permit.

The weight of naphthalene found was .45g, which is .06g lower than expected, and an 88.24% yield. Some naphthalene may have been accidentally filtered out with the aqueous solution, near the beginning of the experiment, reducing the % yield. Also, the process used to separate the ether solution from the sodium sulfate drying agent was by decanting, which may have not been the most effective way, as some liquid could still be caught in the drying agent. A different method that may have given better results would be to filter the solution. A long stem funnel with a little amount of glass wool pushed into the neck is also a good way to filter off the drying agent. The ether solution is poured into the funnel and the glass wool should filter off the sodium sulfate and allow all of the solution to run into the flask.

Once all products were isolated and weighed, melting points were taken, as seen in figure 2.0.

Observed Melting Point (В°C) Literature melting Point (В°C)

Salicylic Acid 154 – 158 159.0

Naphthalene 76 - 79 80.26

Figure 2.0[3]

In both compounds, the observed melting points were depressed to about 1В°C lower than expected from the purest form. Some impurities in the compounds could explain this difference in melting point. In an organic solid, impurities cause defects in the crystal lattice. These defects interrupt and lower the amount of attractive interactions in the structure of the solid. Therefore, less heat energy is needed to break the bonds and melt the solid.

These impurities are relatively small and could thus be filtered out using recrystallization techniques. This involves dissolving the compound in a hot solvent, then allowing the solvent to cool [4]. If the solution cools at a slow enough rate, the crystal formation is also slow and impurities can be “rejected” into the solution and the pure molecules can be selected. This would be an excellent way to further purify the sample as long as an appropriate solvent is used (i.e. it should dissolve the desired compound at a high temperature, but not at a low temperature), and time was available to perform the recrystallization.

The percent yield of the overall product was 95.09%, which is fairly good, and melting points were accurate, but depressed by about 1В°C for each product. With more time to perform this experiment, the products could be purified further using different filtering techniques, different drying agents (magnesium sulfate perhaps), and recrystallization procedures.



[1] Chemical Engineering Tools and Information. Optimize Liquid-Liquid Extraction. <http://www.cheresources.com/extraction.shtml>. (2008). Accessed May 19, 2008.

[3] Lide, David R. “Physical Constanst of Chemical Compounds”. CRC Handbook of Chemistry and Physics 88th Edition (Internet Version 2008. CRC Press/Taylor and Francis. Boca Raton, FL. (2008).

[4] Milligan, G. St.Martin’s University: Separation by Recrystallization. <http://homepages.stmartin.edu/fac_staff/gmilligan/>. (2007). Accessed May 19, 2008.