Joseph K. Myers
2-3-04
(separation of benzoic acid, 2-napthol, napthalene)
What is phase distribution? What is selective dissolution?
The general concept is that under a mixture of two non-soluble organic solvents, a solute with is soluble in both of them will reach an equilibrium of different levels of concentration in each separate layer of the organic solvents. Because they do not mix, the organic solvents themselves can then be separated. Thus, the solute can be purified from other solutes which have different equilibrium levels inside of the same organic solvents. These ideas are what phase distribution and selective dissolution refer to. (The concept can easily be understood without knowing the names that refer to it.)
In this experiment, we are going to apply this theory by creating a first solution of three solutes. After they are dissolved, they cannot be separated from each other until selective dissolution is used. Certain additional solvents will be added to the solution which are non-compatible with each other, and at each step one of the solutes will separate into one or the other of the unmixed layers of the main solutions.
(These are included in text as well as on attached paper.)
(1) C6H5-COOH + NaHCO3 -> C6H5-COONa + (H2CO3 -> H2O + CO2)
(2) C10H7OH + NaOH -> C10H7ONa + H2O
See [1] p. 145.
For (1):
NaHCO3 dissociates into ions Na+ and -O-C{=O,-O-H}. An electron is abstracted from the six available valence electrons upon -O. This is transferred to the H atom in -COOH, and in turn transferred to the adjoining O atom.
This allows the H+ ion to separate, and the intermediate C6H5-COO- joins with the sodium ion Na+ to form a final product C6H5-COONa.
Temporarily, the H+ ion forms H2CO3 with HCO-O2, which reforms into final products H2O + CO2.
For (2):
It is similar to that for (1).
NaOH dissociates into Na+ and O-H.
An electron moves off of O-H onto H within R-OH, breaking the O-H bond, and then moves onto the oxygen atom, leaving the hydrogen proton ion.
The Na+ ion is now free to join with R-O-, and so it does, forming the final product C10H7ONa.
Afterwards, the free hydrogen ion combines with O-H to form the final product H2O.
| Compound | M.W. | m.p. (C) | b.p. | d (g/cm3) |
|---|---|---|---|---|
| Benzoic acid | 122.122 | 122.35 | 249.2 | 1.2659 |
| 2-napthol | 144.170 | 121.5 | 285 | 1.28 |
| Napthalene | 128.171 | 80.26 | 217.5 | 1.0253 |
This particular process is best followed by aid of an understanding of the events which proceed from one another. In a flowchart, the ideas are not plain as they are when straight-forwardly explained. Hence, a student should not look off a flowchart except one that is the student's own, that the student has self-prepared, after being taught to understand what is to be done.
Doing so guides that student, even though the flowchart plays no lasting purpose (as explained).
There are 11 steps to be outlined in a flowchart, which make little sense in and of themselves. (In addition, the final instructions given differed in small part from the main text.)
Specific details such as "Xg of Cx," are extremely confusing in the text of a flowchart; such as the abstraction are the only conveyable images.
1. Prepare the four-part solution.
2. Put solution into separatory funnel.
3. Extract first: with bicarbonate.
(Here, the described result will be repeated: the solution separates into layers; the lower layer is drained from the funnel into a flask.)
4. Extract second: with hydroxide.
5. Remove the remainder of solution from the separatory funnel to a container. (Prior to doing this, the text specified that two spatula tips of anhydrous sodium sulfate should be inserted in this container.)
7. Cool the two extracts.
8. Use addition of HCl to form precipitates from the two extracts.
9. Collect the precipitates by vacuum filtration.
10. Wash solid with cold water.
11. Use simple distillation / evaporation to remove the ether from the so-called "neutral solution."
Details of the extraction solutions are given in Observations and entries.
The fuller flowchart, prepared by pen, is attached with this document.
In preparations:
The beginning of the record for the experiment is the amount of reagents used.
This included .52 g of 2-napthol, .52 of napthalene, and .52 of benzoic acid (amounts are given to 100th of a gram, and it is not hand-waving that they ended up being all the same amounts).
Of diethyl ether, 30 mL was used to initially dissolve the substances.
The bicarbonate (NaHCO3) was dissolved in the amount of 3.00 g, within 30 mL of water.
This bicarbonate solution had a bit of a precipitate (incomplete dissolution).
The sodium hydroxide (NaOH) preparation was produced by 3.00 g of that ingredient, dissolved also in 30 mL of water.
In procedure:
Upon addition of 10 mL of aqueous NaOH or sodium bicarbonate solution, the partitioning within the separatory funnel was complete after one minute.
In obtaining precipitates and products:
No precipitate (benzoic acid) was forthcoming from addition of HCl to the bicarbonate extract.
On the other hand, the precipitate from hydroxide extract (2-napthol) was rather good, and with good yield. The separated layer from the separatory funnel was a brownish color, which turned to foam and precipitate after addition of HCl.
The last extracted item (by evaporation of ether) was napthalene, but it was rather impure, and had a large melting point span.
| Compound | yield | % yield | m.p. exp. (C) | m.p. lit. |
|---|---|---|---|---|
| Benzoic acid | none | n/a | n/a | 122.35 |
| 2-napthol | .50 | 96.1 | (110) 115-125 (+) | 123 |
| Napthalene | .47 g | 90.3 | 82-83 | 80.26 |
Note: it is assumed in this literature that, following temperatures with a temperature unit given, temperatures are specified in the same unit (C). Also, only where it is not clear that a number given refers to a temperature (or other type of value) will such an indication be repeated. For example, within a series of temperatures, degrees C will not be repeated; but in another paragraph, some indication may be given. It is generally harmful to repeat discussed information, except where it is productive, such as in a table of reagents, even though the same information may be present in various isolated places.
Note two: the melting point of the 2-napthol was very broad; this is judged to be for an impure product of 2-napthol. Specifically, a very small top fraction of the product melted at 110; the main product remained intact until temperature 115 to 125, where it slowly and completely melted; a very small remainder never melted all the way to 150.
It helps a student in organic chemistry if the student knows something of the ways of obtaining organic compounds. Without this knowledge, teaching organic chemistry is rather like, "Voila!" and showing the student the magical compounds in the laboratory, where the student knows nothing of the source. Of course, this leads to confusion and frustration.
Although it is difficult to teach the derivation of many organic substances, at least some experience, from labs such as this, leads the student to a feeling of control of the concepts of organic chemistry--which feeling would be absent were organic chemistry only the experience of learning by book and having hands-on experience only within a cordoned-off university.
Seeing the compounds of chemistry as "magical," or obtainable only by the special permission of higher powers is harmful to the science. Seeing the compounds as obtainable only by the knowledge of the science itself is more beneficial to the student and to the furtherance of organic chemistry.
This is the main significance of this experiment, because there was nothing significant in and of itself within the experiment, other than learning these things.
The sources of error for the failure of a bicarbonate extract precipitate are unknown. It may be that the separated layers were not well-enough formed before the lower layer was drained from the separatory funnel.
1. Gilbert JC, Martin SF. Experimental Organic Chemistry. 2002. 3rd Edition. p141-158.