CR Scientific Minerals & Experimental Science newsletter
Issue #11
In this issue:
I. Miscellany
II. A Solar Retort Heater
III. Phenol from Salicylic Acid
IV. Mineral Analysis of Uranophane - Ammonia Precipitation
V. Site News
While the information in this newsletter is thought to be accurate to the best of the authors' current knowledge, it is not guaranteed to be free of errors or to be suitable for any particular use. The procedures and experiments outlined within can be dangerous or even fatal if carried out improperly. If you choose to attempt any of the procedures described in this newsletter or on this web site, you do so entirely at your own risk.
I. Miscellany
If there's any further proof needed that one shouldn't
work when overtired, distracted, or caught in the grip of a high fever, it's
the simple numerical errors that found their way into Issue 10 of the newsletter.
They took the form of temperature conversion mistakes.
Simply writing down the numbers on paper and checking the work would have
prevented the errors in the first place. Getting sidetracked can lead
to trouble, as the reader has probably noticed from any number of other endeavors
in life. For those readers who may someday end up sending people
to Mars, here's a humble word of advice: don't try to save time or
impress people by doing math in your head. Do every little step on
a piece of paper. Double check your work with a calculator. Most
of all, don't talk on the phone while doing even the most basic calculations.
Thanks to our readers who sent in comments, corrections, and questions.
Readers of previous issues have noticed there's been quite a delay in getting
issue 11 finished. The phenol / salicylic acid experiment didn't work
out as planned, so an alternate method was sought. One thing led to
another, the idea was put on the back burner, and pretty soon a great deal
of time had gone by with no alternate method. Finally the writer got
around to trying sodium salicylate instead of salicylic acid; it produced
better results.
Finally, a few readers have suggested various topics they'd like to see in
the newsletter; some of these have materialized in our articles section.
There are still a couple of articles in the works, and they'll be uploaded
as soon as they're ready.
II. A Solar Retort Heater
The premise of this mini-experiment was to make, using plaster of Paris,
a mantle that conformed to the bottom of a retort or distilling flask.
The mantle was then painted flat-black to absorb solar heat.
To begin, the writer obtained a cardboard container somewhat larger around
than the maximum diameter of the retort or flask. An iced-tea mix container
was chosen.
The carton was cut off about 3 or 4 inches above the base. It was filled
with wet plaster. The bottom of the retort was coated with petrolatum
(petroleum jelly) and pressed into the wet plaster. It was clamped
in place until the plaster set completely, around 36 hours. | Figure 1.
The plaster set, and the retort was taken away. Excess
petrolatum was removed with xylene and then acetone (making sure there were
no ignition sources nearby)
When the plaster was thoroughly dry, it was spray-painted
flat black.
Note the cut-off cardboard canister remaining around the
outside of the plaster. This can be left in place. |
|
After drying, petrolatum residue was removed with xylene, followed by acetone. Once the plaster was degreased and dried, it was spray painted with flat black paint and allowed to dry.
Tap water was put in the retort, which in turn was set into the "warmer" and
clamped in place with a ringstand. The retort's long spout went into
an 18x150 mm test tube that was immersed in crushed ice. On a hot day
the ice wasn't expected to last long; a protective cover was made of
white posterboard. This was folded like the peak of a tent and placed
over the receiving test tube and the portion of the retort spout that went
into it. The retort and warmer remained exposed to the sun. 
|
| Figure 2. The solar retort still, shown without the "tent" that would cover the receiving tube (E.). A. Plaster mantle, painted black. B. Retort 1/2 full of tap water. C. Clamps and ringstand. D. Polyolefin lab film wrapped around the joint between the test tube / retort spout; for this experiment, plastic wrap would work just as well. E. Test tube where distilled water condenses (normally covered by a "tent" made from a piece of posterboard or a white cloth). |
Results and Discussion:
The procedure took advantage of three factors. First, water has an
appreciable vapor pressure even when it's not near the boiling point.
Second, the water in the receiving vessel [ideally] has a smaller surface
area than the water in the retort, so fewer water molecules will leave this
area and go back to the tap water reservoir (question: suppose a large puddle
of water had an exposed surface area of 1 square kilometer. Then suppose
the same puddle were put into a very tall and narrow glass column so that
it now had an exposed surface area of only 1 square centimeter. Which
would lose more water to evaporation?). A test tube or other tall,
thin vessel is therefore a good choice for the receiving container.
The experiment will work even without this second factor, but it's very important
to shade the receiving vessel from direct sunlight, which is the third factor.
A piece of white poster board or white sheet covering the test tube or receiving
flask allows for a temperature differential between the tap water reservoir
(the retort) and the condensation vessel (the test tube).
The author actually did this experiment in early spring, but summer would
provide better heating. Maximum daily temperatures were typically 68-70
Fahrenheit (20-21°C), often below this. The absolute ambient temperature
peak during the experiment was 81 F (27.2 C). The measured temperatures
were as follows:Air temp 59°F (15.0°C).....Water temp 98.6°F (37.0°C)
Air temp 72°F (22.2°C).....Water temp 105.8°F (41.0°C)
Air temp 81°F (27.2°C).....Water temp 115.7°F (46.5°C)
Intermittent rainy and cloudy days made it difficult to do real data gathering. Obviously, three points aren't enough to graph a reliable trend, but one gets a good idea from the data (and from common sense) that the retort warmer should work best on a hot, sunny day. A stretch of cloudless days at 30-35°C would be particularly good for further experiments.
Aluminum foil was tried instead of the black coating on the plaster mantle.
Predictably, the foil didn't work nearly as well as the black.
It is not a trivial task to build a decent solar oven from aluminum foil
and get the sun's rays to focus on the water in the retort... especially
when the sun keeps moving. A light-absorbing black background is the
better choice.
The retort warmer has the advantage of no flames and no electricity, meaning
that it's safer to leave unattended. However, don't use the retort
warmer to distill acids or other hazardous materials if there is any possibility
that humans or animals could disturb it. If it's left unattended, chances
are someone or something will get
into it. The experimentation area should be cordoned off and marked conspicuously.
The solar retort warmer, of course, is at the mercy of moving sun and clouds.
A constant temperature is therefore difficult to maintain. Still, the
warmer might be good for separating solvents from dissolved solids, or perhaps
for solvent mixtures whose component boiling points and vapor pressures are
significantly different from each other.
This writer has lately been experimenting with sub-boiling distillation of
HCl to recycle metal-contaminated acid. It would be interesting to
try it in this solar retort still and titrate the product to determine the
final concentration. We may do this in a future article. In the
meantime, we did try another mini-apparatus
for purifying concentrated HCl that was full of contaminating metal ions.
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Figure 3. Don't store phenol in plastic containers. Shown above: phenol crystals gradually melted the plastic where they contacted the bottle walls. There may be some polymers that can resist phenol, but the plastic shown here evidently isn't one of them. This writer's first guess is that this plastic (polystyrene, judging from its appearance) is rich in aromatic rings. The phenol may have taken part in a cycle of gradual opening, closing, and re-opening of bonds. It would be interesting if we could know just how many phenol molecules started life as the contents but ended up as part of the container.
The phenol also became contaminated from the iron compounds that were used to give the plastic its amber color. Hence, the once-colorless crystals had acquired a very faint purplish cast.
III. Experiment: Phenol from Salicylic Acid
Introduction
Practically every chemical laboratory in existence around the dawn of the
20th Century had a bottle of "carbolic acid" sitting on the shelf somewhere.
Now more commonly known as phenol
(fig. 4), "carbolic acid" certainly ranks among the most important industrial
chemicals of modern history. 
|
Figure 4. The phenol molecule, drawn
with the carbons and ring hydrogens assumed (as part of the benzene ring
structure). In other words, there is a carbon at each vertex of the
hexagon, and each one-- except where the OH is attached-- is in turn bonded
to a hydrogen. The formula is usually written C6H5OH,
although sometimes it appears as C6H6O (the latter
being less informative with regard to structure). |
The study of phenol is a good introduction to aromatic organic chemistry.
Since a survey that attempted to be anywhere near comprehensive would take
up an entire book, we'll touch here only upon a few interesting facts, then
get on to an experimental preparation of phenol.
Phenol is the building block for many, many other organic compounds (such
as bisphenol A, phenolphthalein, Bakelite, and many others). While
phenol isn't something one would want floating around inside the body in
appreciable amounts, the phenol structure is incorporated into many biological
compounds (such as the anthocyanins from
red cabbage). Any compound where one or more -OH groups is attached
to a benzene ring can be called a "phenolic" compound. Pyrocatechol,
resorcinol, thymol, salicylic acid, and many others can therefore be thought
of as "phenolics". Compounds with multiple phenols, whether the
rings are fused together or are joined by a single carbon-carbon bond, are
called "polyphenols". Plant polyphenols are the subject of much study.
Phenol is very interesting for several reasons. First of all, it is
technically both an alcohol and an acid. The -OH group of course makes
it an alcohol (hence the "ol" in "phenol"), but its attachment to the aromatic
ring system also means the -OH can lose a hydrogen ion to become a negatively
charged oxide group. This forms the phenolate anion, also known
as phenoxide. For example, it may join up with a sodium ion to form
sodium phenoxide.
A second noteworthy property of phenol is that it's deadly to microbes and
is therefore a powerful disinfectant. Its only drawback is that it
is also deadly to human tissues, at least in appreciable concentrations.
The compound absorbs rapidly through the skin; exposure to large areas
can be fatal. Those who live through such severe carbolic acid burns
may suffer areas of dead tissue (necrosis). Phenol has a mild anaesthetic
effect, so there's the danger that skin contact can happen without the victim's
noticing. In low concentrations, however, the local anaesthetic
effect of phenol is useful for certain topical preparations.
Phenol is one of those compounds that has the interesting 1 tendency
to polymerize with itself under some conditions. Phenol exposed to heat,
catalysts, and / or air can undergo oxidative polymerization; the phenol
molecules form radicals, and before long there's a brown tar in the bottom
of the test tube. Many different substances can catalyze phenol
polymerization. Even catalase from horseradish can do it, though catalase
isn't what one would think of as a phenol enzyme (its normal substrate is
hydrogen peroxide). There are references on the Internet and in the
biochemical literature about catalase's action on phenol; we may explore
some of these in a future article.
Industrially, phenol is made either from cumene hydroperoxide, sodium benzenesulfonate,
or chlorobenzene. It was at one time also produced from aniline (q.v. Vogel's Practical Organic Chemistry). The
sodium benzenesulfonate route is probably the most commonly-used one for laboratory
preparation. In this experiment, however, we've attempted to prepare
some phenol by decarboxylating salicylic acid and / or sodium salicylate.
This is probably one of the oldest methods for preparing phenol, even though
it fell out of wide use in industry long ago.Procedure
WARNING: Phenol is corrosive to living tissues. It is also toxic.
Phenol can kill you via inhalation, ingestion, or even skin contact. This procedure also uses the highly caustic compounds Calcium Oxide and Sodium Hydroxide. It is not intended for beginners. If you choose to attempt it, you do so entirely at your own risk; you must wear safety goggles and gloves at all times and work in an area having abundant ventilation (in other words, a fume cupboard). Do not work with phenol unless you are familiar with its hazards and also with its safe handling.
There are two main routes to phenol from salicylic acid. The first
is to heat salicylic acid in the presence of a catalyst. The second
is to convert salicylic
acid to sodium salicylate and then heat the salicylate.
Salicylic acid can be decarboxylated by heat in the presence of soda-lime
(a mixture of CaO and NaOH), but the salicylic acid tends to sublimate before
reaction is complete. It will therefore heavily contaminate the final
product; since phenol and salicylic acid are fairly difficult to separate
from one another, it is better to avoid this route. One study (Toland,
1961) suggests decarboxylation of salicylic acid in the presence of CuO,
with a high-boiling solvent such as xylene added to hinder the escape of
the salicylic acid from the reaction mixture. Steam may then be fed
through the reaction mixture to carry away the phenol (1961).
Sodium salicylate, on the other hand, can be heated without sublimation.
It decarboxylates with relative ease when exposed to heat, even without the
aid of soda-lime. Trial I. Salicylic acid ---> Phenol
As a sort of "proof of concept" study, this writer first tried the direct
approach, decarboxylation of salicylic acid with soda-lime. All reagents
were used dry, having been stored at low humidity. The soda-lime started
to get hot after brief exposure to the moisture in the air, but by this time
it was in a borosilicate test tube getting ready for the torch flame.
Safety goggles and gloves were of course used in case of a mishap.
A test tube was fitted with a one-hole stopper and a double-90°-bend
gas delivery tube which led into a flask chilled on an ice bath. The
contents of the test tube were heated with an alcohol burner flame.
A dense, white vapor issued from the mixture upon strong heating; this
distilled over into a receiving flask that was chilled on an ice bath.
After the whole setup cooled, there was the distinct odor of phenol (don't inhale this;
these vapors are toxic).
This indicated the reaction had worked at least to some extent. The
next step was to use a common test to try detecting the product by some means
other than odor.
Phenol
reacts with ferric ions (Fe3+) to give a purple color.
Salicylic acid,
on the other hand, forms more of a purplish-red color
in the presence of ferric ions. In a clean spot plate well, a solution
of Fe3+ ions was prepared by covering a few crystals of ferric
chloride hexahydrate with distilled water and letting them dissolve.
Some of this was added to three separate spot plate wells containing (A)
salicylic acid, (B) the suspected phenol, and (C) U.S.P. grade phenol.
As the test showed, this first trial yielded a product that was heavily contaminated
with salicylic acid (fig. 5).
Figure 5. The Iron test for phenol and salicylic acid, conducted on the product from Trial I. Each test substance was dissolved in distilled water.
"A" contains salicylic acid, originally prepared from hydrolysis of acetylsalicylic acid. This salicylic acid had no odor whatsoever and presumably contained no appreciable phenol.
"B" contains the "phenol" condensate collected from the receiving flask. This condensate had a very definite phenol odor but still gave more of a burgundy hue with Fe+++.
"C" contains U.S.P.-grade phenol. Pure phenol crystals dissolve quite rapidly in water.
Even though the one marked "B" almost certainly contains phenol, the color suggests masking by salicylic acid.
1.17 grams of sodium salicylate and 0.41 grams calcium oxide (corresponding
to about 7.3 millimoles of each) were mixed together thoroughly and placed
in a 10 mL micro flask.
The flask was fitted with a 1-hole stopper having a gas delivery tube. This
was led into another flask fitted with a 1-hole stopper. The stopper
in flask 2 was kept loose to prevent dangerous pressure buildup; ventilation
was used to remove the phenol vapors that would inevitably escape.
Strong heating caused dense, white vapors to fill the apparatus. Liquid
phenol condensed in the delivery tube, some of it making its way into the
receiving flask. However, the heating was evidently too strong.
The phenol soon turned red-brown when too much heat was applied.Some of the condensed liquid was placed on a microscope slide and allowed to crystallize (fig. 6)
|
Figure 6. Crystals of the putative
phenol, allowed to solidify at room temperature on a microscope slide. Magnification: None. These crystals emitted a definite phenol odor. Melting point determination was not tried, but the crystals had uneven areas of color that suggested impurity. |
The material was tested with ferric
ions and compared to a known phenol sample. While it did give a positive
test for phenol, the color reaction was not as intense as the known phenol
and seemed to have more of a grayish cast to it.
DiscussionTrial I produced phenol, but it was heavily contaminated with salicylic acid that had also vaporized during heating. In fact, salicylic acid probably far outweighed phenol in the condensate.
Trial II produced phenol that was presumed to be free from salicylic acid, but too much heat was applied to the reaction mixture. Undetermined condensation / oxidation products gave the liquid a red-orange to brown color. Future attempts would require temperature control, or at least more gradual heating.
Despite the apparent similarity of the two routes, March (1992) suggests
the decarboxylations of aromatic acids and their salts follow entirely different
mechanisms. They both involve loss of CO2 , but the latter
type involves formation of an aromatic carbanion (a negatively-charged carbon
species) by the SE1 mechanism (1992). Carboxylate departs
from the ring as CO2, leaving behind the aromatic ring with an
unshared electron pair. Something nearby has to donate an electron-poor
species (typically, H+) to make the ring stable again. Without
an extensive journey into the literature, we cannot pretend to know exactly
what transient species form in that reaction vessel.
It is easy to draw structures and try to guess what mechanisms might happen.
It is also easy to be wrong. When high temperatures are involved and
there could be any number of undetermined side reactions, we cannot be sure
just what is happening. So we make a "swag" (actually, a S.W.A.G.).
This is an acronym for "Scientific", "Wild", "Guess", and the traditional
name for a donkey. In some circles, the term is used interchangeably
with "hypothesis".
We do know that unassociated ions cannot distill over (except, perhaps, at
temperatures that would vaporize the distilling apparatus and the laboratory
with it), so there must be overall charge balance. If a carbanion forms
in the reaction vessel, that ionized species is going to have to get a hydrogen
ion from somewhere or react with something to alleviate the charge imbalance.
A "minus" needs a "plus"; if there's a phenolic carbanion, it cannot
distill out of the reaction vessel as neutral phenol until that carbanion
abstracts something positive (generally, H+) from somewhere.
So, here's our SWAG. We started with sodium salicylate. We are
reasonably sure the carboxylate group departed the reaction vessel as carbon
dioxide, leaving behind a carbanion and a temporarily unattached sodium cation
("temporarily" must be emphasized, as this situation may last for mere millionths
of a second, perhaps less). We could guess the carbanion to be attacked
instantly by that sodium, perhaps giving us "2-hydroxyphenylsodium".
In any case, we might expect the thing to rearrange pretty quickly to sodium
phenoxide, a much more stable compound (again, this is conjecture).On the other hand, we know that in our collection vessel we are getting phenol, not sodium phenoxide. If we presume, as part of our guess, that atmospheric water vapor played some role in the reaction, we could easily satisfy the problem by assuming water to give an H+ to the aromatic carbanion (or phenoxide ion) and an OH- to the Na+, leaving behind NaOH in the reaction vessel as the phenol distilled away. However, we might also wildly guess that some of the reactive carbanions in the flask might attack other rings, causing unwanted polymers and other products that would decrease our yield of phenol.
It happens that the above may be more of a genuine "hypothesis" than a swag.
We know that forcing the reaction the other way (introducing excess CO2
into the heated system under pressure) will give us the Kolbe-Schmitt reaction,
turning sodium phenoxide back into sodium salicylate. Thus, our wild
little guess about the mechanism takes on some dignity, even if it should
turn out to be wrong.
The actual intermediates that form during pyrolysis reactions (i.e., dry distillations) have until recent
decades been mostly a mystery. Only with some very sophisticated equipment
has it been possible to analyze transient species and confirm or disprove
some of the guessing that has taken place over the years.
Lacking such equipment, however, an experimenter has to make optimum use
of what's available. As stated before, it's quite certain that one
can't distill free ions, electrons, or protons under normal circumstances.2 Whatever happens
in that reaction vessel has to include one or more steps that satiate the
carbanion, probably with a hydrogen taken from somewhere else. There
are certainly molecules that could in theory give up a hydrogen to this carbanion
and yet remain neutral species; consider, for example, how an alkane
(electrically neutral) can be dehydrogenated to yield an alkene (also neutral).
As our experimental procedure stands at the moment, however, there's
nothing in the reaction vessel that can do so without ultimately taking from
the reservoir of future phenol molecules.
As mentioned before, alkali metal-carbon bonded organic species are interesting
possibilities, but they'd decompose rapidly into e.g. sodium phenoxide, putting us back
at species that can't distill over into the receiving flask without first
taking H+ from some source. Thus, ring fusions, cross-linking,
ring openings, self-condensation products, and myriad other, side reactions
are almost certain unless we introduce some water vapor into the mix.
Although transient species eventually must settle down to something stable,
the chemical possibilities during those moments of heating are intriguing.
At temperatures high enough to make borosilicate glass turn bright orange,
nearly anything is possible.
It's become a recurring theme in our newsletter and in our other articles:
the study of one thing has led us to the possible study of numerous others.
If one were to engineer the decarboxylation reaction for maximal yield, one
might start by choosing something that could be added to the reaction mixture
and which:1.) does not form permanent adducts with salicylate or phenol
2.) does not boil away easily
3.) has the capability to give up hydrogens to satisfy the unstable carbanion, forming phenol molecules which could then leave the reaction vessel.
In the meantime, we've seen that it's possible to form small amounts of phenol
by heating sodium salicylate. We may not know exactly how it works
at the moment, and we may not know offhand what modifications will improve
it, but we know that it works.
While phenol for experiments is better obtained ready-made from a chemical
supplier, the lab preparation of this compound illustrates some very interesting
science.
As always, safety is an important consideration. Keep the goggles on
your face and the ventilation system running.
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Notes - "Phenol from Salicylic Acid"
1 Sometimes the tendency to polymerize isn't interesting; it's just annoying. Phenol is not nearly as obnoxious in that respect as styrene (which can self-polymerize even from agitation or contact with air), but it can go from white crystals to brown tar easily enough. Back to article
2 A person once telephoned wanting to know if it was possible to distill electrons. The answer is no, unless the apparatus can be made as hot as the Sun. Of course, a cathode ray tube from a television can be thought of as distilling off electrons, because the electrons vaporize off a very hot filament. The electrons can't exactly be collected and stored in a beaker, however.
Back to article
References - "Phenol from Salicylic Acid"
March, J. Advanced Organic Chemistry. New York: Wiley, 1992.
Merck Index, 10th edition. Rahway, New Jersey: Merck & Co., Inc., 1983.
Toland, William. "Oxidative Decarboxylation of Aromatic Acids to Isomeric Aryloxy Derivatives". J. Am. Chem. Soc. 83(11): 2507-2512 (1961).
Vogel, Arthur. A Textbook of Practical Organic Chemistry. London: Longman Group Ltd., 1957
IV. Analysis of Uranophane using Ammonia Precipitation
Because of its length, the uranophane article is in its own section. Please click the link to go to Part 2 of Newsletter 11.
V. Site News
Belomo loupes
are currently in stock; however, we received news the factory is going
to be raising the price by 45%, effective immediately. We are therefore
not sure if we are going to be getting new stock. At the moment, we
have about 30 pcs. left, and they're going fast.
There are quite a few new articles on the site,
including one about growing ferrous sulfate
crystals from H2SO4 and steel. We also have a
new article about salicylic acid that
relates to the phenol experiment.The articles on electrolysis are improved and expanded. For those readers interested in microscopy, there's a new article about fixatives and one about a stain formulation.
We've added various electrodes to the
on-line catalog.
There are also vacuum desiccators,
a couple new types and sizes of test tube racks,
and nickel crucibles.
In fact, we've got quite a few new items - have a look!
That concludes this issue of the CR Scientific Newsletter. Until next time, stay safe, learn some science, and have fun with it!
While the information in this newsletter is thought to be accurate to the best of the authors' current knowledge, it is not guaranteed to be free of errors or to be suitable for any particular use. The procedures and experiments outlined within can be dangerous or even fatal if carried out improperly. If you choose to attempt any of them, you proceed entirely at your own risk. To use this website or any of the information contained herein, you must read and agree to our Terms of Use.
You may print out or make photocopies of this newsletter for educational and personal use (Click here for Copyright). However, this newsletter may not be copied or distributed, in whole or in part, for commercial or for-profit use.
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