Archive-name: sci/chem-faq/part4
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Last-modified: 2 October 1997
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Subject: 15. Chemical Demonstrations
15.1 Are there any good compilations of demonstrations?
Yes. Good places to start are the four volume "Chemical Demonstrations"
by B.Z.Shakhashiri [1], or the two volume "Chemical Demonstrations - A
Sourcebook for teachers" by Summerlin and Ealy [2]. The Journal of Chemical
Education is also an excellent on-going source of novel demonstrations
and developments of traditional demonstrations.
15.2 What are good outdoor demonstrations for under 12s?
15.3 What are good outdoor demonstrations for over 12s?
15.4 What are good indoor demonstrations for under 12s?
15.5 What are good indoor demonstrations for over 12s?
While waiting for a promised contribution, here is my only contribution,
and some from my sci.chem archives. Unfortunately, enthusiastic editing
by others allows some of the culprits to go uncredited :-).
The ability of water-miscible solvents to mask the hydrophobic nature of
Goretex can be demonstrated. Goretex is just a porous PTFE, the same
material as PTFE filters - such as Millipore HF. You can easily filter
liquid water through porous PTFE, provided the filter is previously wetted
with a water-miscible solvent ( usually ethanol ). If a filter is set up on
a vacuum flask, ensure the filter is completely wetted with ethanol, turn on
the vacuum, and immediately add water - it rapidly filters through. Once it
has stopped, it only takes about 15 seconds for the air to dry the filter,
then ask a student to filter more water from the same flask. No chance.
Pour off the water, surreptitiously add a few mls of ethanol, immediately
followed by the same water - and watch it filter through again :-).
This is the nearest equivalent our laboratory has to the workshop practice
of sending an apprentice out to purchase a spark plug for a diesel engine.
It does relate slightly to the real world - indicating why "breathing"
fabrics like Goretex should not be used with solvents.
From: brom@yoyo.cc.monash.edu.au (David Bromage) Date: Tue, 14 Sep 1993
Subj: Re: Need: A safe chemical display
The so-called "Blue Bottle Reaction" might be useful.
Half fill a 1 litre flask with water and add 10g of NaOH, then add 10g
of glucose and up to 1ml of 1% methylene blue. Stopper the flask and
swirl gently to dissolve the contents. On standing for a few minutes the
solution should turn colourless. When the flask is shaken the solution
will turn blue then decolorise on standing.
Methylene blue exists in solution as a reduced colourless form and an
oxidised blur form. The initially blue dye is reduced by the alkaline
form of glucose and re-oxidised by dissolved oxygen. When the solution
is shaken, atmospheric oxygen enters into solution at a more rapid rate
than when left standing. The dye acts here as a catalyst whose colour
indicates the redox state.
[ This demonstration, and different coloured versions of it, have recently
been discussed in the J.Chem.Ed.[3]. ]
How about a chemical garden?
Make up (or dilute a commercial preparation) of sodium silicate to
1.1g/ml. Place this solution in a large glass container then add 'lumps'
or large crystals of salts to be grown. Lumps should not be more than
0.5cm in diameter. As a salt dissolves it forms an insoluble silicate
which forms a membrane around the lump of salt. The membrane is
permeable to water which enters and dissolves more salt. The resulting
pressure bursts the membrane releasing more salt solution to form more
membrane. As the salt solution is less dense than the silicate solution,
the membrane grows as a convoluted vertical tube.
Salt Colour Growth time
Ferric chloride brown 1 hour
Ferrous sulphate grey-green 3 hours
Cobalt chloride purple 5 hours
Chromium chloride grey-green 6 hours
Nickel sulphate yellow-green ~24 hours
Cupric sulphate blue ~24 hours
Potassium aluminium sulphate white ~1 day
To produce a "garden" which is not completely overgrown with the faster
species it is necessary to take growth rates into account. Distilled
water should be used as Ca and Fe in tap water can cause cloudiness.
If you really want oscillating reactions, I know of two.
A. Iodate reaction.
Make up 3 solutions
1) Dilute 200ml of 100 vol hydrogen peroxide to 500ml
2) dissolve 21g of potassium iodate (KIO3) and 1.5ml of conc sulphuric
acid in 500ml of water.
3) Dissolve 7.8g of malonic acid and 1.4g of manganese sulphate in 400ml
of water and add 1.5g of starch in 100ml of water.
Add equal volumes (50-100ml) of each solution to a flask in any order.
Colourless-blue oscillations should start within 2 minutes. If not, try
10-20% variations in relative volumes. (try increasing 2 first).
Oscillations should last up to 10 minutes but I my experience have lasted
up to 3 hours.
B. The Belusov reaction
Prepare 5 solutions.
1) 58g of malonic acid on 500ml of water
2) 6M sulphuric acid
3) 21g of potassium bromate (KBrO3) in 500ml of water
4) Dilute 5ml of solution 2 to 500ml then add 1.75g of cerous sulphate.
5) 1.6g of 1,10-phenanthroline and 0.7g of ferrous sulphate in 100ml of
water (or commercial ferroin solution to 0.025M)
Mix together 50ml of 1 to 4 and 5ml of solution 5. Blue-pink oscillations
should start within a few minutes.
For either oscillating reaction the choice exists of complete mixing with
uniform oscillations or waves of colour (eg in a measuring cylinder).
Some interchange of reagents is possible. The Bray reaction omits
malonic acid from the Iodate reaction. Malonic acid can be replaced by
citric or succinic acids.
[ There have been several good discussions [4,5], and recipe compilations
[6], for many popular oscillating chemical reactions.]
A particularly dramatic 'trick' is not to burn paper. Make up a solution
containing 57% v/v ethanol and 43% v/v water with 5% w/w sodium
chloride. Soak a filter circle in the solution and hold it near a flame
(with tongs) just long enough to ignite. After the flames die down the
filter circle will still be intact. The ethanol burns but just enough
water remains in the paper to prevent ignition. NaCl is added to provide
a more convincing flame. To add drama, 'burn' a banknote - but ensure
that all of the note, especially the corners, is soaked.
From: lmartin@uclink.berkeley.edu (Lonnie C Martin) Date: 17 Feb 1993
Subj: Re: Growing a Silver Tree in Beaker?
In article <...> xslkkk@oryx.com (kenneth k konvicka) writes:
>Am trying to do a demo for elementary school kids. How do you grow a tree
>of silver using copper wire(?) submerged in a solution of AgNO3? Saw one
>in high school physics class about a thousand years ago at good ol' Reagan
>HS, Austin, Tx. Was really beautiful. The silver formed nice large
>plates. Any demonstration books you could steer me toward?
What you have described is about all there is to it. I do this
demonstration for the chemistry classes here at Berkeley about twice a
year, or so. Just make a "tree" out of copper wire (you might clean it
with sandpaper or steel wool) so that it will fit into a beaker of your
choice (we use 4 litre here), and pour in the silver nitrate solution.
I think we use 0.1 molar, but as long as the concentration is fairly close
to that, it will work just fine.
It is not necessary to make the tree very "bushy". The silver will fill it
out nicely with fuzzy thick hanging globs of crystals. The solution will
change from colourless to blue, as copper nitrate is formed. A very nice
experiment. You can expect this to take on the order of an hour to get
fully developed.
From: flatter@rose-hulman.edu (Neil Flatter) Date: Tue, 14 Sep 1993
Subj: Re: Need: A safe chemical display
We use cobalt (II) chloride in a saturated sodium chloride solution to
demonstrate cooling coils. It changes from red/pink to a blue/purple when
heated and reverses as it is cooled. We cycle it through a condenser from
a distillation to illustrate that portion of a simple set-up.
Subject: Stupid lab tricks -Compiled- VERY LONG Date: 7 Jul 92
From: Nazman <nasadk@rpi.edu>
Ever try taking an empty ditto fluid can, put some water in it, heat it
until steam is coming out, cap it back up and let it cool off?. You would
be surprised what a little air pressure can do. That one amazed me when I
was young.
I was amazed again when I saw a brief description on TV of how science
teachers are trying to make science fun again. Four teachers on stage, set
up a few ring stands and a few bunsen burners, and placed a 55 gallon oil
drum on top. Boiled the water, capped it. Put a hell of a dent in the drum
when it collapsed.
A favourite of mine requires a little preparation, but is great fun. Try
tearing an aluminum can in half. Kinda difficult. Now, if you take an empty
can, gently score around the circumference on the inside, (the inside is
coated to prevent a reaction between the soda and the can) and fill the can
with a solution of warm water and Copper(II)Chloride (CuCl2) so that the
solution is just above the score mark. Let this sit for a few minutes.
You are done with the solution when the outside of the can appears brownish
(blackish) where the score mark is. Gently pour out the solution (keep it)
and let the can sit. When ready, hold the top of the can in one hand and
the bottom in the other, and break like you are breaking a stick in half.
Two bits of advice :
BE CAREFUL!. You will end up with the sharp edge of the can, which can cut
severely!
Try this ahead of time just to make sure you get it right. Wait too long,
and when you pick up the can, it will split due to the weight of the
solution. Don't wait long enough and it won't work. My guess is about
5 minutes
P.S. This is more of a demonstration of structure of an aluminum can, but
if you want to demonstrate the "strength" before you rip it in half, place
the can on the floor, so it is sitting like it normally would, and balance
on one foot off the top of the can. It helps to have something nearby to
hold on to, and the can cannot have any dents. You would be surprised how
strong an empty can is. I weigh about 190 lbs, and have stood on an empty
soda can for 30 seconds, get off the can, and not have it collapse. This
takes some practice, so give it a try.
From: A_ROSATI@GUVAX.GEORGETOWN.EDU <Anthony V. Rosati>
You can followup the bromophenol blue trick by brewing a cup of tea
and, while they watch, add some lemon juice. The color will lighten.
There is an indicator in tea that changes with the acidity of
ascorbic acid.
Another neat trick is to demonstrate the dehydration capacity of
concentrated sulfuric acid. Take a 500 mL beaker about one third full
of white table-sugar. Then add about a half-inch to one inches worth
of concentrated sulfuric acid. (This demonstration _MUST_ be conducted
in a hood) Let it sit for about five minutes. Within that time, the
sulfuric acid will seep in, start turning the color of the sugar
brown, and then black, followed by an intense, hot dehydration. The
sugar will start to form a jet black, smelly, sticky column that rises
out of the beaker. It is really impressive.....
You might want to also look up "oscillating reactions" in your
chemistry library. Many of these are simple to set up and generate
neat color cycles that would impress the kids!
From: mfrancis@ucsd.edu < Lyn Francisco >
1. Take a balloon, blow it up, tie it, then stick it in a vat of
liquid nitrogen. Wait until it shrinks (around 3 s or so), take it
out, and then watch it inflate in your hands. This will very nicely
illustrate the relation between temperature and pressure.
2. We did this during a demonstration to let the world know about
ACS on campus. Take a large container (like one of those 10-gallon
water containers, cut in half or something), fill it up with water,
then put in one can of the original Coke and one can of diet Coke.
Make sure that both cans are unopened. Now, drop a few pieces of
dry ice in the container. The original Coke should drop to the
bottom, and the diet Coke stay up toward the top. It was cool, and
attracted all the frat-types and non-science people to our table.
From: dfield@nike.calpoly.edu < Dan Field >
If you really want to fire them up, my favorite has always been the
hydrogen balloons. Just fill up several balloons, one color with air or
He, another with H2, and another with 2H2 + O2. You can fill them ahead of
time, or better yet if demonstration time allows, use the products of one
of your demonstration reactions to fill the balloons. Light a candle on
a L O N G stick, dim the lights, and pop!, boom!, B O O O M !!.
You'll have instantly created little monsters, young pyromaniacs virtually
guaranteed to associate some excitement with chemistry.
[ Warning - the sound level of such explosions has recently been found to
exceed health and safety guidelines, and people should read the article
before demonstrating this experiment to students or children [7]. ]
From: edremy@d31ha0.Stanford.edu < Eric R.>
There are lots of things you can do with liquid N2. Try freezing a
banana and using it as a hammer. (Follow by using an unfrozen banana:
kids love it!) Simply adding some to a test tube and (lightly!) corking
it is fun, provided you're careful with the cork. Shattering a
superball is also good.
However, my personal favorite for spectacular demos is the HCl fountain.
Ascii graphics follow
---------
\ / Top flask is filled with HCl gas
\ S S=rubber stopper w/ hole for needle
\ /
-|-
/ | \ Run tube from top into bottom
/--|--\ Bottom flask filled with water and
/ | \ acid/base indicator.
-----------
MAKE SURE THAT THESE FLASKS ARE VACUUM SAFE!!!
To start this whole extravaganza, inject 20-30 cc of water into the top
flask. The HCl gas goes into solution, creating a partial vacuum, sucking
the water up from the bottom. As the water spurts out of the tube, it
collects more HCl (And changes color as it becomes acid) and accelerates
the reaction... Quite impressive.
We used to do this for our chemistry magic show every year. The only
problem is that the failure mode is somewhat dangerous: One year the
top flask had a flaw and imploded, sending glass and HCl everywhere.
Best to do behind a shield
From: Bill
I believe that the same thing can be done with ammonia.
The same precautions apply.
From: ?
Bubble H2 through a soap solution and you get bubbles that float up.
Have them float through a bunsen burner flame suspended over the table and
they explode. VERY NEAT effect.
From: joec@morgan.com <Joe>
USUAL WARNINGS: many chemicals are poisonous and some reactions
may be difficult to control. Use your head.
Best done indoors
-----------------
Dissolve silver nitrate in warm water. Get some copper wire
and clean it with steel wool. Insert copper wire (preferably coiled
at one end) in the solution and it will immediately dull. Some time later,
silver crystals will be CLEARLY visibly growing on the copper.
The best effect is to let it sit overnight. The resulting
effect is downright beautiful
Dissolve Cobalt Chloride in warm water. Put some Aluminum foil
in it and watch it tarnish. Clean, polished Aluminum works
best but household aluminum foil also works (just slower).
The Aluminum slowly disappears and Cobalt metal shows up at the bottom.
This is a slow one but it does work.
Light an alcohol lamp, i.e. denatured alcohol and bring a magnet near
the flame but not above it- to the side. Watch the flame get pulled in
the direction of the magnet.
Sprinkle iron filings over the same alcohol lamp and watch sparks fly!
Ignite some Magnesium ribbon and drop into an atmosphere of CO2. It
will continue to burn with lots of noise and sparks. Carbon dust will
rain down as a byproduct.
Mix water and household (3 in 1) oil. Note the phase boundary.
Add soap and shake. Watch the phase boundary disappear.
Heat up a piece of blackboard chalk with a propane torch. Chalk
is CaCO3 - heating it up will drive off CO2, leaving CaO (also known
as lime). Heating up lime will cause the it to emit a whitish light,
which is where the phrase 'limelight' comes from.
[ Note - not all blackboard chalk is CaCO3 - test carefully first ]
Do these outdoors:
------------------
Get some KMnO4 and pour into a small pile. Depress the center of
the pile slightly and add a drop or two of Glycerine and stand back.
Something between 1-5 minutes later, it will burst into flame.
When it dies down, drop some more glycerine on it to have it flare up
again. Be careful disposing of the KMnO4 left over - its a powerful
oxidizer.
We also do THERMITE periodically (Aluminum powder and rust). Details for
those who ask - it burns *BRIGHT* and *HOT*.
Drop some dry ice chunks into a 2 liter PLASTIC soda bottle 1/2 full with
warm water which is then quickly sealed. Get at least 50' ft back rather
quickly. The pressure will build up and detonate with a LOUD *BOOM* after
a brief and unpredictable time. The bottle will break into many hundreds
of parts (don't use GLASS!) and you will get a mist cloud some 20-30'
across. Note: It is quite LOUD and may scare a younger audience.
Make Hydrogen soap bubbles and set them off. Get an erlenmeyer flask and
fit a cork into the top and route a glass tube through the tube and
have it bend down and into a jay of soapy water. Remove the cork and
drop in Zinc metal and pour in somewhat dilute HCl. Put the
cork back in and let the H2 bubble into the soapy water. This will
make H2 soap bubbles. Let them break free and ignite them with a light
match on long pole.
Thermite reaction
First of all....this is a fairly vigorous reaction so take the usual
precautions:
1-Do it outside, preferably on sand or dirt. Since it burns at 4000 degrees
fahrenheit, it will melt most anything. By the way, a nuclear explosion
burns at 8000 and the surface of our Sun burns at 10000. It will readily
melt rock salt, beach sand, etc. You get the idea.
2-It can spray sparks around. Keep it away from combustible materials. The
burning sparks are either molten Aluminum or molten Iron.
3-It is VERY bright so you shouldn't stare at it.
4-It puts out lots of smoke.
Here is how I do it.
Ingredients:
1-Aluminum powder
2-Iron Rust (Red-Fe2O3).
Grind carefully and separately into a powder-like consistency.
Mix in roughly equal proportions, by volume with an excess of rust.
Mix thoroughly to get an even color.
Pour the powder mixture on the ground in a pile.
Get magnesium ribbon and lay it on top of the pile, and press partially
into the pile. Do not smother the Mg ribbon. Ignite the ribbon with a
propane torch and get back quickly.
When done, be careful...it will leave molten, glowing red iron as a
byproduct.
You can make rust by mixing household clorox with steel wool pads and let
sit overnight and then filtering out the rust.
Have fun and be careful.
Usual disclaimers apply
From: gallivan@after.math.uiuc.edu < Justin Gallivan >
This works nicely with soap bubbles in a dish. If you have the H2 and O2
tanks available, Try a few with the H2 only which makes a nice quiet
flame and add the O2 later for a little shock value. You may want to
try this first for safety's sake but it always went off without a hitch in
my general chemistry days.
From: Rob
I hope I'm not too late. An extremely simple trick is done with a chunk
of styrofoam (larger the better) and some acetone, which is an excellent
theta-solvent for styrofoam.
Simply spray the acetone out of a bottle onto the styrofoam, and the
styrofoam rapidly decomposes, losing its structure, and appears to
actually be melting. It is quite a "dramatic" demonstration, and can be
offset against how nicely styrofoam coffee cups hold water/coffee, but not
acetone.
From: ?
I thought this one was neat...
Take a bottle (should be reasonable size, like a ketchup bottle)
fill it to within 2" of the top, color light blue (not opaque!)
with methylene blue. Drop in a NaOH pellet and a few drops of
Karo clear syrup. Other reducing sugars might work; I just know
it works with this syrup. (Or did; the last time I tried it was
almost 20 years ago, and they may have changed the formula since
then.)
Over a period of a few minutes, the blue color will fade. Shake
the bottle, and suddenly it's blue again. Leave it, and it will
slowly fade. It'll last for a couple of days, until random
microbes do in the sugar I suppose.
From: ?
A "Bottle of fire" for lighting bunsen burners and such:
Get a dark, heat-resistant glass bottle, and put just enough pentane
in it to wet the sides. (i.e., rinse it with pentane and dump out
the excess.) Light the top of the bottle. The flame will burn down
into the neck of the bottle a little, but be almost invisible to the
audience. Pick up the bottle, turn it over, and flames will pour
out. Set it down, and the flames seem to go out.
When Dr. Toffel did this, someone said "There's something in the
bottle!" He said "Nope," poured some water from the faucet into the
bottle, dumped fire and water into the sink, then showed that the
bottle would still "pour fire". (This probably takes some practice.)
From: mvp@hsv3.lsil.com < Mike Van Pelt >
Portable bunsen burner:
Bubble air through a test tube of pentane, and run this to
your bunsen burner. You can use a large balloon as your air source,
or have a vict... I mean, volunteer, blow through the tube.
From: Howard Clase.
One experiment that I like was you make a solution of lead nitrate,
which is clear, and a solution of some iodide salt (potassium iodide),
which is also clear. When you mix the two of them together you form
a yellow solid - lead iodide.
This is only half of it! If you don't use too much of the chemicals
to produce your "instant orange juice" - but DON'T let anyone drink it.
You will find that the lead iodide will dissolve if you heat the solution.
On
Cooling it re-precipitates as beautiful golden spangles.
From: mgray1@metz.une.oz.au < Matthew Gray >
Another exciting and easy impress all trick is to get two solutions, one
of Ag(I) and another of Cu(I), usually both hexamine complexes. When
these two are mixed, a redox reaction takes place, producing a silver
mirror effect. Other reducing metals can be used, such as iron, but I
haven't tried these myself.
From: ?
Grind some potassium permanganate to a fine powder (to speed up the
reaction). Put it in a small heap (1 teaspoon) on a tile, make a dent in
the top and pour one drop of glycerine in the hole. After about 10-15
seconds the heap will catch fire.
From: torin.walker@rose.com
Here are some that are rather interesting. All of these tests have been
performed in my workshop and are all safe (with the exception of the
handling of HCl and the irritating effect of experiment #2). Experiment #3
is by far the most fascinating.
1 Copper Sulfate couple grams in a test tube.
Sodium Bicarbonate - same as above.
These two liquids are transparent but when mixed, turn into a soft blue
opaque suspension.
2 Glycerin and HCl
Takes a long time (couple of hours) to complete but when these two clear
liquids are mixed together, it turns from clear to a deep transparent red
and slowly goes brown. Warning - this is extremely irritating to the eyes
if you are exposed to it for a while - usually, an hour is enough to
really get you annoyed.
3 (My favorite) Acetone (you can buy large tins of this stuff (1L) at a
hardware store in the automotive section (usually with the bondo and
other body repair supplies) and styrofoam (a large bag of popcorn type
packaging filler will be needed.)
When styrofoam is placed in acetone, the styrofoam ( large volume of
styrofoam for a small volume of acetone ) dissolves and becomes a wet,
play-dough like substance that feels cold to the touch.
This experiment is harmless unless swallowed :-) and should prove to be
quite interesting to the students.
The coldness is due to the evaporation of the acetone from your skin
(ever use nail polish remover? That's acetone.) The acetone will
eventually all evaporate (a 2 inch sphere of this will take a day or two)
and the result will be a porous (trapped acetone bubbles) material that
can be molded to any shape you wish.
From; David O'Driscoll. University of Central Queensland...
Hope someone hasn't already done this one, I have been studying for
exams so have not been reading all of them.
The one we use at our high school demos are pH clocks....
quite good as they are not static displays.
First, take three or four large (1L) beakers and 3/4 fill them then
take your favourite pH indicators (ones with good colours), and add a
few drops to them, then add some dilute sodium hydroxide or something
to make them slightly basic. Next add a handful of dry ice to each beaker.
This creates a nice bubbling mixture with good visual effects, what happens
is obvious (I hope!!!). Some of the CO2 is dissolved in the water, turning
the mixture acidic and when the end-point of the indicator is reached the
colour changes - sometimes quite dramatically. The kids seem to like it and
the chemistry is not too involved.
From:webbb@mbf.UUCP ( Bryan Webb )
I didn't see the originating message of this thread, but from the
responses that have made it here, I think this is the kind of stuff
you might be looking for. In earlier times, I've done these:
1) Place a small pile (several grams) of powdered magnesium on a surface
you don't care about in an environment provided with plenty of
ventilation. On top of this, place a couple of grams of powdered
iodine (well, as close as you can get to it, though that might not
be crucial). Now, put a couple drops of water on the iodine ...
enough to also contact the magnesium ... and stand back. The heat
of the reaction vaporizes some of the remaining iodine into a purple
vapor.
2) This is pretty dangerous, so be very careful. Take a couple of grams
of red phosphorous and place on top of a couple of grams of potassium
iodate. Rapidly stand back... spontaneous combustion. My experience
was a time delay of a couple of seconds, but I wouldn't want to count
on it... I discovered this accidentally... boy was I surprised. The
speed of the reaction may be related to the humidity.
3) Potassium dichromate is normally bright orange at room temperatures. If
it is cooled to liquid nitrogen temperatures, it becomes yellow. If
heated, it becomes a deeper red color. I'm not aware of any other
inorganics that have this range of color change when the temperature
is varied.
4) Ahhh, my favorite... When I was in high school, I took the 2nd year
chemistry class that was offered. We had the resources of the
school at our disposal, so long as the experiment we wanted to do
was "in a book". The book I had was "Chemistry Magic", and described
an "experiment" where some cotton balls were placed on a fireproof
surface, a few grams of Sodium Peroxide was placed on top, and then
you put a drop or two of water that will wet at least a little bit
of both the peroxide and the cotton.
It's a long story, but I this experiment worked, at least on other
cellulose objects like paper towels. In fact, the fire in the
metal trash basket was hot enough to melt/burn away the bottom,
the linoleum underneath, and some of the concrete in the floor.
The flames formed a "solid" yellow flame and lots of thick white
smoke (containing NaOH dust). You really don't want to breathe
this stuff. We didn't, anyway :-)
5) Oh, another thing we did in that class was take the gas outlet used
for the bunsen burners and direct it into a test tube that was
partially submerged in liquid nitrogen. (The whole system was
sealed.) The gas condenses into a liquid... the only problem
was safe disposal. It helps to plan ahead! :-)
[ Note that nuke@reed.edu subsequently supplied the following warning ]
" If you decide to try this be aware that liquid nitrogen will condense
liquid oxygen in a vessel open to the air immersed in it. Liquid O2
forms explosive mixtures with many organics. IF you still want to try
it, immerse the tube in the nitrogen and then immediately run the gas
in. only do a little bit. How much you get depends on what proportion
of weights of low hydrocarbons the gas contains ( I think methane
condenses at this temp, but not quantitatively like some stuff, unless
there is a large surface area)."
6) One of my classmates made luciferin [sic]. It's a liquid that
glows in the dark for about 12 hours. That was fun too!
Happy researching!
Standard disclaimers apply; I'm not sure my company would have hired
me if they had the foregoing admissions before them.
Non-standard disclaimers too: I don't recommend you do any of these things
either.
From: fred@theory.chem.pitt.edu < fred >
If you would like to condense out methane gas in a relatively safe way,
fill a balloon with the gas and THEN condense the gas with liquid N2.
You can use scissors to cut the balloon, and pour the liquid CH4 into
a beaker with water in it (notice that it floats, forms ice, etc.) and
light it. Only the fumes burn as they mix with atmospheric oxygen.
This makes a fair "olympic torch." Wear goggles etc.
From: Larry (Call me "Lefty") C
One that can be safely performed with a long enough spatula.
Mix Calcium Carbide with any strong oxidizer (KMnO4, NaNO3, even MnO2
works). Proportions aren't real important here.
Using face shield, gloves, lab coat and long spatula, drop a SMALL
amount (say, 1 gram or so) of this into common household bleach.
Acetylene and chlorine are evolved, which immediately, uh... exploded
Delightful chlorinated hydrocarbons result, unfortunately :(
15.6 How do I safely perform the Glowing Pickle experiment?
This experiment consists of electrically heating a vegetable that has been
soaked in a brine solution to conduct electricity. Because this experiment
involves electricity at dangerous voltages, the experiment should be
performed on special apparatus under qualified supervision. I'm not going
to detail the equipment and procedures, as they have been described in
an Journal of Chemical Education article [8]. The experiment has not just
been limited to table salt and pickles, many other vegetables and salts
that produce different colours have been investigated and described in the
same issue of the Journal of Chemical Education [9]. People intending to
perform the experiment should obtain both articles.
15.7 How do I make Slime?.
" Slime " is a trademarked commodity obtained by cross-linking guar gum and
borax, and is marketed by the Mattel Toy Corporation. The slime produced for
demonstrations is usually made by cross-linking a poly vinyl alcohol (PVA)
product using borate. The normal method is to carefully prepare a 4%
mass/volume aqueous solution of a hydrolysed high molecular weight PVA
( >100,000 ) - available from Eastman Kodak. Commercial PVA-based adhesives
( such as Elmer's Glue ) will also produce a reasonable quality slime, as
will polymeric materials that have multiple hydroxyl groups and can form
highly-hydrated gels, such as guar gum - but some experimentation may be
required to ascertain optimum ratios.
High MW polymers are difficult to dissolve in solvents ( including water ),
and the best method is to carefully sprinkle the powder over a beaker of
water that is being gently stirred, and continue gentle stirring until a
uniform solution with no gelatinous lumps is obtained.
Any grade of borax ( Na2B4O7.10H2O ) can be used to prepare a 4% mass/mass
aqueous solution. The slime is made by vigorously mixing the two solutions
in the ratio of 1-2 parts of the borax solution to 10 parts of the PVA
solution using a paddle stirrer. Details of a suitable procedure for use in
classrooms have been published [10]. A firmer, less messy, slime can be
prepared from an 8% PVA solution - using equal ratios of the 4% borax
solution [11].
The properties of slime indicate that the cross-linking mechanism does not
consist of strong covalent bonds. Borax hydrolyses in water to form a boric
acid-borate buffer with an approximate pH of 9.
B(OH)3 + 2H2O <==> (B(OH)4)- + H3O+ pK = 9.2
The borate ion is tetrafunctional when interacting with the alcohol groups
of polyols, and thus builds the cross-linking structure. PVA has about 1-2%
of 1,2 diols amongst the remaining 98-99% of 1,3 diols. To obtain the desired
properties, the bonds between the borate and the PVA must be weak, and it is
believed they are hydrogen bonds ( shown as ... below ).
PVA Borate PVA
H H
| |
O-H...O O-H...O
\ / \ / \ /
H-C \ / C-H
/ B- \
CH2 / \ CH2
\ / \ /
H-C-O...H-O O...H-O-C-H
/ | | \
H H
Although individual hydrogen bonds are weak, the large number of available
OH groups in highly-hydrolysed PVA will result in a hydrated, 3-dimensional,
gel, rather than a borate precipitate. The continual breaking and reforming
of the bonds under low mechanical stress, and the large amount of water
incorporated into the gel, are responsible for the rheological properties of
the hydrated gel. Slime can be broken down by reducing the concentration of
borate by titration with a strong acid, and details of such a procedure have
been recently published [12].
------------------------------
Subject: 16. Laboratory Procedures
16.1 What are the best drying agents for liquids and gases?
The Rubber Handbook lists the traditional information on drying agents
that involve on chemical action. This lists phosphorus pentoxide and
magnesium perchlorate as the most effective desiccants. However, later
work by Burfield [1-9] has demonstrated that much of the traditional
information is misleading. He found that the efficiency of the desiccant
is strongly dependent upon the solvent. He also found that Drierite
( anhydrous calcium sulphate ) is only a moderately efficient desiccant for
organic solvents [9], and that correctly prepared molecular sieves are
often the preferred desiccant [2]. His publications are highly recommended.
16.2 What is the effect of oven drying on volumetric glassware?
Many older laboratory texts insist that volumetric glassware should not
be oven dried because of the danger of irreversible and unpredictable
volume changes. However most modern laboratory glassware is now made of
Pyrex, and work by D.R.Burfield has demonstrated that low temperature
drying does not significantly affect the calibration of volumetric
glassware [10]. He demonstrated that exposing volumetric flasks and
pipettes to 320C, either continuously or thermally cycled, resulted in no
significant detectable change to the calibration. He concluded that
"oven temperatures in the range of 110-150C should provide efficient drying
of glassware with no risk of discernible volume changes, even after
prolonged use, providing that Pyrex glass is the material of construction".
16.3 What does the Karl Fischer titration measure?
In 1935 Karl Fischer used the reaction between iodine, sulfur dioxide, and
water to produce a technique for quantifying water [11]. In aqueous solution,
the reaction can be presented as I2 + SO2 + 2H2O <=> 2HI + H2SO4.
He used anhydrous methanol to dissolve the I2 and SO2, and added pyridine
to move the equilibrium to the right by reacting the acidic products.
Fischer assumed his modifications did not change the reaction and one mole of
iodine was equivalent to two moles of water. Smith et al.[12], demonstrated
that both the methanol and pyridine participate in the reaction and one mole
of iodine is equivalent to one mole of water. They suggested two steps:-
(1) SO2 + I2 + H2O + 3RN -> 2RN.HI + RN(SO2)O
(2) RN(SO2)0 + CH3OH -> RN(SO4CH3)H (where R = base = C5H5 for pyridine)
This was further investigated by E.Scholz [13], who proposed:
(1) CH3OH + SO2 + RN -> (RNH)SO3CH3
(2) H20 + I2 + (RNH)SO3CH3 + 2RN -> (RNH)SO4CH3 + 2(RNH)I (where R = Base)
The advantage of the Karl Fischer titration is that it has few interferences
and can quantify water from < 1ppm to 100% in diverse samples, ranging from
gases to polymers. It will measure all water that is made available to the
reagent. the endpoint is usually ascertained using a dead-stop endpoint,
and for low water levels coulometric techniques are used to quantitatively
produce the iodine by anodic oxidation of iodide. The procedures are
described in detail in ASTM, AOAC etc.
16.4 What does the Dean and Stark distillation measure?
The Dean and Stark procedure can be used to measure the water content of
a diverse range of samples, and has been extensively used in industrial
laboratories to measure water in petroleum oils. The technique can measure
% levels of water, but is not as accurate as the Karl Fischer titration,
and is not applicable to samples where the water is not liberated by the
solvent. The sample is mixed with a solvent ( usually a toluene/xylene mix )
and refluxed under a condenser using a special receiver. There are two common
designs of receivers, one for solvents that are heavier than water, and the
more common one for solvents that are lighter than water - illustrations will
be shown in most laboratory glassware supplier catalogues.
The water and solvent are refluxed, and as they condense the two phases
separate as they run into the receiver. The water remains in the receiver
while the solvent returns to the flask. The Dean and Stark technique is also
useful for removing unwanted water from reactions, eg the synthesis of
dibutyl ether by the elimination of water from two molecules of n-butanol
using acidic conditions. An example of this is provided in the preparation
of dibutyl ether described in Vogel, and detailed procedures for the
determination of water using Dean and Stark are provided in ASTM and AOAC.
16.5 What does Kjeldahl nitrogen measure?
The Kjeldahl procedure is routinely used to measure the protein nitrogen
content of organic compounds, especially natural foodstuffs. Contrary to
popular belief, the procedure does not determine total nitrogen on all
organic compounds, as it is not applicable to materials containing N-O or
N-N linkages without modifications to the method. This discrepancy is
becoming of more significance as automated nitrogen analysers using other
techniques are producing different results because they measure the total
nitrogen present.
The method usually involves high temperature ( 390C ) digestion of the
sample using concentrated sulfuric acid, a catalyst ( Cu, Hg, or Se ),
and a salt to elevate the acid boiling point. In some cases 30% hydrogen
peroxide is also used, making the digestion effectively a high-temperature
piranha solution attack on the organic matter. After digestion, the sample
is made strongly alkaline and the ammonia is steam distilled into a boric
acid solution, and aliquots are titrated against a standard acid using an
indicator solution endpoint.
Some organics compounds require aggressive digestion conditions to make
all the organic nitrogen available, consequently Kjeldahl procedures should
not normally be used on samples that may have N-O or N-N bonds. Details of
procedures for foods are in the AOAC handbooks, and general Kjeldahl
procedures are detailed in the ASTM volumes.
16.6 What does a Soxhlet extractor do?
The soxhlet extractor enables solids to be extracted with fresh warm solvent
that does not contain the extract. This can dramatically increase the
extraction rate, as the sample is contacting fresh warm solvent. The sample
is placed inside a cellulose or ceramic thimble and placed in the extractor.
The extractor is connected to a flask containing the extraction solvent, and
a condenser is connected above the extractor. The solvent is boiled, and the
standard extractor has a bypass arm that the vapour passes through to reach
the condenser, where it condenses and drips onto the sample in the thimble.
Once the solvent reaches the top of the siphon arm, the solvent and extract
are siphoned back into the lower flask. The solvent reboils, and the cycle
is repeated until the sample is completely extracted, and the extract is
in the lower flask.
There is an alternative design where the hot solvent vapour passes around
the thimble, thus boiling the solvent in the thimble - this can be a problem
if low-boiling azeotropes form. Procedures for using soxhlet extractors are
described ( along with illustrations which might make the above description
comprehensible :-) ), in Vogel and many other introductory organic laboratory
texts.
16.7 What is the best method for cleaning glassware?.
As scientific glassware can be used for a variety of purposes, from the
ultra-trace determination of sub-ppq levels of dioxin, to measuring %
concentrations of inorganic elements, there is no single cleaning method that
is "best" for all circumstances. Difficult and intractable deposits often
involve the use of hazardous and corrosive chemicals, and details of the
necessary safety precautions for each cleaning solution should obtained
before attempting to clean glassware. The use of heat and/or ultrasonic
agitation can greatly improve the removal rate of many deposits, especially
inorganic and crystalline deposits.
Whilst the semiconductor industry use piranha solution ( refer Section
12.9 ), and several other reactive and toxic chemicals for cleaning, those
reagents can react dangerously with the residues found in laboratories, and
their use is prohibited in some institutions. Such chemicals should only be
used after extensive prior consultation with laboratory management and safety
staff - to either identify safer alternatives, or to ensure that appropriate
protective and safety systems are in place.
If the probable composition of material deposited on the glassware is known,
then the most appropriate cleaning agent can be readily selected. There are
several safe aqueous cleaning solutions that are routinely used. If possible,
glassware should be washed or soaked immediately with an appropriate solvent
for the residue. This will make subsequent cleaning easier, but all traces
organic solvents must be removed before using any cleaning solution.
The most common aqueous-based soaking solutions are commercial formulations
that usually contain alkalis, chelating agents, and/or surfactants, and can
be used either at ambient temperature, or temperatures up to boiling ( with
ventilation - caustic fumes are noxious ). These are very effective for
general grime, most labels, pyrogens, and many common chemical residues, and
well known examples include RBS-35, Decon, Alconox, and Pyroneg. Their main
advantages are low toxicity and ease of disposal.
The next common strategy involves physical abrasion to remove deposits
inside flasks, usually with a bottle brush and an aqueous cleaning solvent
( like those above ) or a suitable organic solvent. A refinement is to add
sand, pumice, glass spheres, or walnut shell chips, along with some water
or solvent, and shake vigorously. It's important that the sand should not
have sharp edges - as it can scratch the glass. It has been suggested that
table salt in solvent ( eg petroleum spirit, methylene chloride, acetone )
is superior, as it doesn't scratch the glass, can be easily removed by
washing with water, and has minimal disposal problems [14].
The traditional glassware cleaning solution is "chromic acid", and many
analytical chemistry texts detail the preparation [15,16]. Chromium (VI) is
highly toxic ( mutagenic, carcinogenic ), and disposal is expensive, as all
solutions containing more than 5 mg/l of chromium are considered hazardous
waste in the USA. Disposal of chromic acid requires a two-stage process,
involving bisulfite addition to reduce Cr(VI) to Cr(III), followed by
neutralisation of the acid. There have also been several reports of
spontaneous explosions of chromic acid cleaning solutions [17,18,19],
consequently the use of chromic acid for cleaning glassware is declining,
and several alternative glassware cleaners have recently been evaluated [20].
Sodium dichromate dihydrate is usually used to prepare chromic acid, as
potassium dichromate is less soluble in sulfuric acid. One technique is to
dissolve 140g of technical grade sodium dichromate dihydrate in approximately
100 ml of water. Add two litres of technical grade 98% sulfuric acid to a 4-5
litre glass beaker that is sitting in a cold water bath in a fume cupboard.
Carefully stir the acid gently and pour a few mls of the dichromate solution
slowly into the acid. Keep repeating the addition every few seconds - after
the previous dose has been dispersed. As long as the stirring is gentle and
continuous, little or no splattering should occur, but the solution will
become quite warm. Allow to cool before storing in a glass-stoppered reagent
bottle. Always ensure that the stopper is sufficiently loose to release any
gas pressure. Never use a screw-capped or similar types of sealed containers.
If made correctly, the chromic acid solution should have no precipitate, will
be a deep red colour, and will last for years in a glass-stoppered bottle.
Ensure the glassware to be cleaned does not have any residual organic
solvents. Chromic acid is very effective at around 80C, but an overnight soak
at ambient temperature is commonly used. If the solution develops a green
hue, it is exhausted and should be disposed of, or regenerated, using
appropriate procedures. Slowly pouring used acid down a drain with the cold
water tap fully open is no longer considered appropriate. There is a recent
report of a technique to regenerate chromic acid cleaning solution ( by
distillation of water and oleum ) that reduces disposal quantities [21].
The major problems with chromic acid are the multiple rinses, and perhaps
even alkaline EDTA treatment [16], that are necessary to remove all the
chromium from glassware - especially if it is required for cell culture or
trace analysis, and the increasing problems of safe and legal disposal of
spent solutions.
An alternative to chromic acid is "Nochromix", which is commercial solid
formulation that contains 90-95% of ammonium persulfate ( ((NH4)2)S208 )
along with surfactants and other additives. The powder is dissolved in
water and mixed with 98% sulfuric acid. The solution is clear, but turns
orange as the oxidizer is consumed, and further additions of solid are
routinely required. It is available from Godax Laboratories, New York.
A similar bath that is reported to be very effective can be made by the
addition of 19 grams of reagent grade ammonium persulfate to two litres of
reagent grade 98% sulfuric acid [22]. Add more ammonium persulfate and acid
every few weeks, as necessary.
One popular replacement for chromic acid in organic laboratories has been
alcoholic sodium hydroxide or potassium hydroxide solutions. These remove
most deposits, with metals and hydrocarbons greases ( Apiezon ), as notable
exceptions. One advantage they have is that they will remove silicone grease
deposits from joints and stopcocks, especially if warmed to 65C, and the
glassware immersed for up to 10 minutes [23]. Prolonged immersion, even at
ambient temperature, will damage ground-glass joints, dissolve glass sinters,
and will leave glass surfaces translucent or opaque. The solution can be
prepared by either adding two litres of 95% ethanol to 120 mls of water
containing 120 grams of sodium hydroxide [16], or by dissolving 100 grams of
potassium hydroxide in 50 ml of water and, after cooling, make up to one
litre [15].
Solutions based on hydrofluoric acid, usually containing 1-5% of HF, also
rapidly attack glass, and destroy sinters, but are very effective for removal
of carbonaceous and fine silica deposits. They also remove silicone greases,
but alcoholic caustic solutions are preferred [23,24,25]. Hydrofluoric acid
is corrosive and extremely nasty if it comes in contact with humans. It
requires extensive safety precautions before use. For most deposits, only a
few minutes are required, and ultrasonic agitation often assists the removal
of deposits. Cleaned glassware usually remains transparent. Cleaners
containing HF should not be used on volumetric glassware.
Another acidic solution, comprising of a 3:1 mixture of concentrated sulfuric
acid and fuming nitric acid, is also extremely effective for removing grease
and dirt, but also requires extensive safety precautions. The grease and dirt
can often be removed more safely using hot aqueous-based cleaners.
If you have intractable organic-based deposits in flasks without standard
ground glass ( or clear glass ) joints, then some deposits can be carefully
burned off in a glass annealing furnace. The glass needs to carefully
follow a slow heating and cooling schedule to minimise thermal stresses and
distortion. My experience has been that standard joints do tend to freeze
more often after such treatment. Also note that glassblowers may not want
to coat their annealing furnace with your rubbish, so they may prohibit
the use of their furnace for such activity.
------------------------------
Subject: 17. Preparation of chemicals
17.1 Where do I find laboratory-scale procedures for organics?
The best introductory handbooks are practical textbooks, eg "Organic" Vogel
and "EPOC" Vogel. They provide a diverse range of experiments that soon help
develop synthetic skills. If you master the preparations in Vogel you are
at the stage where you can start to obtain papers from organic chemistry
journals and reproduce their syntheses. There are also several texts that
discuss techniques for purifying laboratory chemicals, eg [1] The parameters
of common specialist synthetic procedures usually are fully described in
specialist texts that will only normally be available in chemistry department
libraries ( eg Palladium Reagents in Organic Syntheses [2]). Most educational
institutions will have a structured laboratory programme to develop skills.
17.2 Where do I find laboratory-scale procedures for inorganics?
Most synthetic chemistry of inorganics appears to be devoted more to complex
organometallics, superacids and superconductors than common inorganics, but
it is worth considering that, of the top fifteen industrial chemicals
produced, the only organic compounds are ethylene, propylene, ethylene
dichloride and urea. There are specialist texts available that describe how
to purify inorganic laboratory reagents, eg [1]. I expect some inorganic
chemists to berate me for not knowing the standard inorganic synthesis
textbooks. ;-)
17.3 Where do I find industrial chemical process details?
The standard text for common processes remains Shreve, and I must admit that
I enjoy reading the 1945 first edition to obtain a good overview of an
industry. McKetta provides excellent process design details, along with
comparisons of various processes. Kirk Othmer provides an excellent update
on the various processes and chemicals used extensively today. Kirk Othmer
remains the first port of call, but Ullmann is a close second. Both of these
provide extensive references to more specific texts.
Industry journals, eg Hydrocarbon Processing, offer annual reviews of the
processes used in their industry. Patent literature has to be treated
cautiously, as it is not always immediately obvious which patents detail
actual viable processes. Chemical engineering texts, eg Perry, provide
comprehensive detail of the equipment and operational parameters.
------------------------------
Subject: 18. Sensory properties of chemicals
18.1 How do light sticks work?, and how can I make one?
From: perks@umbc.edu (Mark Perks) Date: 15 Sep 1994
Subject: Re: Chemiluminescence Sticks
Chemical Demonstrations [[1] v.1 p.146- ], by Bassam Shakhashiri, offers a
thorough discussion of Cyalume lightsticks. Professor Shakhashiri is at
the University of Wisconsin, Madison, I believe.
"The Cyalume lightstick contains dilute hydrogen peroxide in a
phthalic ester solvent contained in a thin glass ampoule, which is
surrounded by a solution containing a phenyl oxalate ester and the
fluorescent dye 9,10-bis(phenylethynyl)anthracene...When the ampoule is
broken, the H2O2 and oxalate ester react.."
From: chideste@pt.Cyanamid.COM (Dale Chidester) Date: Mon, 13 Mar 1995
Subject: Re: How to make chemical light ?
The following produce rather spectacular results. Chemicals are
available through Fluka and Aldrich. The dyes are expensive.
Dyes:-
9,10-bis(phenylethynyl)anthracene (BPEA) (yellow) [10075-85-1] Fluka 15146
9,10-diphenylanthracene (DPA) (blue) [1499-10-1] Fluka 42785
5,6,11,12-tetraphenylnaphthacene (rubrene) (red) [517-51-1] Fluka 84027
Other reagents required:-
bis(2-carbopentyloxy-3,5,6-trichlorophenyl)oxalate (CPPO)
[75203-51-9] Aldrich 39,325-8
bis(2-ethylhexyl)phthalate (DOP) (solvent) [117-81-7] Fluka 80032
sodium salicylate (catalyst) [54-21-7] Fluka 71945
35% hydrogen peroxide [7722-84-1] Fluka 95299
Saturate solvent with dye and CPPO. Sonicate to help solvation. Start with
about 50 mg dye (BPEA, DPA or rubrene) in 10 g solvent with 50 mg CPPO and
5 mg sodium salicylate. CPPO is limiting reagent.
Put small quantity (20 drops) in a small vial and add equal volume of
hydrogen peroxide. Mix vigorously. There will be two phases. Avoid skin
contact! Don't cap tightly!
The following explanation of the chemistry was provided:-
From: sbonds@jarthur.claremont.edu (007)
All of the material below is taken from a chemical demonstrations book
[[1], v.1, p.146 ].
The oxidant is hydrogen peroxide contained in a phthalate ester solvent.
The concentration is very low, less than 0.5%. The fluorescing solution
consists of a phenyl oxalate ester and a fluorescent dye. The dye used is
9,10-bis-(phenylethynyl)anthracene (for green) or 9,10-diphenylanthracene
(for blue).
Here is the reaction sequence:
1) (Ph)-O-CO-CO-O-(Ph) + H2O2 --> (Ph)-O-CO-CO-O-OH + (Ph)-OH
2) (Ph)-O-CO-CO-O-OH --> O-O
| | + (Ph)-OH
OC-CO
3) C2O4 + Dye --> Dye* + 2CO2
4) Dye* --> Dye + hv
In 1) The hydrogen peroxide oxidizes the phenyl oxalate ester to a
peroxyacid ester and phenol. The unstable peroxyacid ester decomposes to
the cyclic peroxy compound and more phenol in step 2). The cyclic peroxy
compound is again unstable and gives off energy to the dye as it decomposes
to the very stable carbon dioxide. The dye then radiates this energy as
light.
An alternative chemiluminescence demonstration involves the H2O2 oxidation
of lucigenin ( bis-N-methylacridinium nitrate [2315-97-1] Aldrich B4,920-3 ),
[ [1] v.1 p.180-185 ] which has recently been modified to provide a slow
colour change across the visible spectrum [2]. One of the reagents in that
lucigenin oxidation ( Rhodamine B ) is a mutagen and suspected carcinogen.
18.2 How do hand warmers work?, and how can I make one?
They consist of an aqueous solution of sodium acetate with a small "clicker"
disk to impart a physical shock. The solute is dissolved into solution by
prior warming. when the heat is required, the disk is "clicked" to shock the
solution, and this causes the sodium acetate to crystallise from the now
supersaturated solution. The heat of crystallisation is slowly released.
18.3 What are the chemicals that give fruity aromas?
Most of the desirable food aromas come from low to medium molecular weight
organic compounds - usually alcohols, aldehydes, esters, ketones, and
lactones. These may be " natural " ( extracted from natural sources ),
" nature-identical " ( synthetic, but identical to known natural compounds ),
and " artificial " ( synthetic, not found in nature ). The perceived aroma of
molecules can change dramatically with minor isomeric or structural changes,
and common fruity aromas are usually complex mixtures of several compounds.
Because man-made chemicals are frequently made from chemicals derived from
fossil fuels, the isotopic ratios of the carbon atoms has been used to
discriminate between natural and nature-identical chemicals. Natural
processes usually involve the use of enzymes that selectively produce a
specific isomer, and some man-made aromas are now produced enzymatically.
Chiral chemistry, often utilising chiral chromatography that was initially
developed for pharmaceuticals, is now also being used for the production
and testing of man-made aromas, as enantiomerically-pure aroma chemicals
command premium prices.
Some chemicals are listed below, along with their use in either fragrances
and/or flavours, and boiling point ( 760 mmHg, unless otherwise stated ).
Some of them are also considered toxic, and thus their use may be controlled.
Volume A11 of Ullmann has an excellent monograph on flavours and fragrances,
and more detail can be obtained from the journal Perfumer and Flavorist.
The catalogues of well-known suppliers such as Dragoco GmbH ( Germany ),
L.Givaudin and Cie ( Switzerland ), and Takasago Perfumery Company ( Japan ),
also contain information on chemical composition and health and safety.
Chemical BP CAS RN Application
C (mmHg)
acetoin 148 [513-86-0] butter
acetophenone 202 [98-86-2] orange blossom
benzyl acetate 206 [140-11-4] jasmine
butyl acetate 125 [123-86-4] apple
2,3-butanedione 88 [431-03-8] butter
(+)-carvone 230 [2244-16-8] caraway, dill
(-)-carvone 230 [6485-40-1] spearmint
citral 229 [5392-40-5] lemon
citronellal 207 [2385-77-5] balm mint
decanal 208 [112-31-2] citrus
dihydromyrcenol 78 (1) [18479-58-8] lavender
2,6-dimethyl-2-heptanol 171 [13254-34-7] freesia
ethyl butyrate 120 [105-54-4] pineapple
ethyl 2t-4c-decadienoate 71 (45) [3025-30-7] pear
ethyl hexanoate 168 [123-66-0] pineapple
ethyl isovalerate 132 [108-64-4] blueberry
ethyl 2-methylbutyrate 133 [7452-79-1] apple
geraniol 229 [1066-24-1] roselike
hexyl acetate 169 [142-92-7] pear
hexyl salicylate 168 (12) [6259-76-3] azalea
1-(4-hydroxyphenyl)-3-butanone [5471-51-2] raspberry
isoamyl acetate 143 [123-92-2] banana
(+)-limonene 176 [5989-27-5] lemon
linalool 198 [78-70-6] lily of the valley
linalyl acetate 220 [115-95-7] bergamot
8-mercapto-p-menthan-3-one 57 (8) [38462-22-5] blackcurrant
1-p-methene-8-thiol 40 (1) [71159-90-5] grapefruit
3-methyl-2-cyclopenten-2-ol-1-one [80-71-7] caramel
4-methyl-2(2-methyl-1-propenyl)tetrahydropyran
70 (12) [16490-43-1] rose
myrcenol 78 (50) [543-39-5] lime
2t-6c-nonadien-1-ol 98 (11) [28069-72-9] violet
3-octanol 175 [20296-29-1] mushroom
1-octen-3-ol 84 (25) [3391-86-4] mushroom
phenethyl acetate 238 [103-45-7] rose
phenethyl alcohol 220 [60-12-8] rose
phenethyl isoamyl ether [56011-02-0] chamomile
2-propenyl hexanoate pineapple
alpha-pinene 156 [80-56-8] pine
alpha-terpineol 217 [98-55-5] lilac
alpha-trichloromethylbenzyl acetate [90-17-5] rose
18.4 What is the most obnoxious smelling compound?
Many low molecular weight sulfur-containing compounds tend to induce adverse
reactions in people, even if they have not encountered them before, eg the
glandular emissions of skunk (n-butyl mercaptan, dicrotyl sulfide).
Butyric acid reminds people of vomit, and cadaverine ( 1,5 Pentadiamine )
reminds people of rotten tissue, but without an earlier association, they
may not regard them as unusually obnoxious.
18.5 What is the nicest smelling compound?
Aside from thinking about your stomach, when the smell of cooking foods
is attractive, then most people like the smell of flowers and citrus fruits.
These are volatile, aromatic, oils, whose major components are complex
mixtures of medium volatility compounds, often derived from terpenes, eg
Oil of Rose ( 70 - 75% geraniol = (E)-3,7-dimethyl-2,6-octadiene-1-ol ),
Oil of Bergamot ( 36 - 45% linalyl acetate = 3,7-dimethyl-1,6-octadien-3-yl
acetate ). Many aromatic oils are mixtures of terpene esters ( oil of
geranium = 20 - 35% geraniol esters ) or aldehydes ( oil of lemon grass =
75 - 85% citral = 3,7-dimethyl-2,6-octadienal ). Merck briefly describes
nearly 100 volatile oils, from Oil of Amber to Oil of Yarrow, along with
typical applications. Flower perfumes are complex blends of compounds, and
detailed compositions of your favourite smell are often available in the
journal " Perfumer and Flavorist ".
Expensive flower petal perfumes, such as rose and jasmine, are produced
using extracts obtained by the traditional "enfleurage" process ( refer to
Section 24.4 ).
