Chemistry Lab Demonstrations: Candy Chromatography

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Last lab of the semester today. Next week is the lab final and checkout. This week the students practiced column chromatography. They purified their crude product mixture from last week’s nitration lab. I’ve talked about the theory behind column chromatography before, so I won’t rehash it here in any detail. Suffice it to say that different organic compounds have differing affinities for a stationary phase versus a mobile phase. These differing affinities allow for one compound of interest to be separated from a mixture through the use of column chromatography. Students were aided this week in that their product was bright yellow. They could physically watch it run down the column, then only collect the yellow fractions.

Last lab of the semester means last demo of the semester.  This one’s a do-it-yourself demo, if you’d like.  You can separate the colors contained in M&M shells (or Skittles, or Reese’s Pieces, or Sharpies, etc) through chromatography.  I got my M&M proceedure here.  If you’re interested, other proceedures are available here, and here.   Basic rundown: put drops of water on wax paper, and put a piece of candy on each drop.  Allow for the water to strip the color off the colorful candy shell.  Cut a coffee filter into a rectangle.  Use a toothpick to spot each color onto the coffee filter.  Put the coffee filter into a 1% solution of table salt and allow the water to rise through the coffee filter.  Watch the colors separate like magic!

Couple’a observations I noticed.  Quite interestingly… the stationary phase matters.  A lot.  I started by spotting the colors on my silica gel TLC plates .  I was quite disappointed because the red and yellow both travelled with the solvent front and there was little separation.  I tried several different solvents… no luck.  I also noticed that according to the websites I was looking at, red should have travel the shortest distance.  Then I switched over to filter paper, and all of a sudden I got the results I was expecting.  Who knew?  Also, you should put a crease in the coffee filter before placing it in the solvent.  The paper will start to buckle and it will droop and fall over if it is not creased first.  The more distance you give the colors to separate, the better the results.  I used the largest filter paper we had, and ran the chromatograph twice to get the results shown.

Pop quiz, hot shot: Do you know what the difference between Red 40 and Red 40 Lake are?  I didn’t either.  Turns out… nothing.  At least, not as far as the compound responsible for the hue is concerned.  It’s all in the formulation:

Color additives are available for use in food as either “dyes” or “lakes”.

Dyes dissolve in water, but are not soluble in oil. Dyes are manufactured as powders, granules, liquids or other special purpose forms. They can be used in beverages, dry mixes, baked goods, confections, dairy products, pet foods and a variety of other products. Dyes also have side effects which lakes do not, including the fact that large amounts of dyes ingested can color stools.

Lakes are the combination of dyes and insoluble material. Lakes tint by dispersion. Lakes are not oil soluble, but are oil dispersible. Lakes are more stable than dyes and are ideal for coloring products containing fats and oils or items lacking sufficient moisture to dissolve dyes. Typical uses include coated tablets, cake and donut mixes, hard candies and chewing gums, lipsticks, soaps, shampoos, talc, etc.

There are 5 food coloring agents in M&Ms: Red 40, Yellow 5, Yellow 6, Blue 1, and Blue 2.  As you might expect, green separates into blue and yellow, but surprising the red and yellow of the orange M&M do not separate.  Rather, there is one orange spot with a larger Rf than red.  Brown separates to blue, red and orange.   But it looks like the blue in the blue M&M is a different blue than the blue in the green and brown M&M.

I’ve got lots of pictures from my experience (click for larger).  Note how poorly silica works and how different the Rf’s are between silica and filter paper.  the video is of separating components of felt tip pens, but it’s also neat.

There are no more demos planned, since the lab course is over.  Hope you enjoyed my miniseries.

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Chemistry Lab Demonstrations: Silver Nitrate/Copper Wire

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This week in lab, students performed electrophilic aromatic substitution (and here).  Dissolution of 4-methylacetanilide in 70% nitric acid gives mono nitration.  There are two possible products.  The acetamide is a better ortho/para director than the methyl group, so 2-nitro-4-methylacetanilide is the major product of the reaction.  The reaction was a bit touchy.  In my lab, for the most part the reactions were carried out at room temperature.  This resulted in the reaction not occuring!  Nearly every student got back unreacted starting material, instead of product.  The product is supposed to be a bright, crayon yellow solid.  Other labs ran the reaction on low heat and got excellent results…  But some students left the reaction on the heat too long or at too high of a temperature and the reaction decomposed into this ugly brown oil.  So there is a very small window for success in this reaction.  For the first time this semester (!) students analyzed the reaction mixture by Thin Layer Chromatography.

A brief safety warning.  Nitric acid is a very strong oxidizer.  It reacts explosively with readiliy-oxidizable small organic molecules like alcohols and acetone… acetone being of course what all good lab students rinse their glassware with before they start lab.  There was an explosion in our department last year as a result of improperly mixing nitric acid waste and acetone waste.  Nitric acid MSDS here, incident report on nitric acid explosion (not from our department) here (it’s probably worth glancing through the other incidents on that page as well), and video showing gas evolution from nitric acid oxidation (this time of copper) here – note how much gas is produced in a short amount of time.  Imagine this in a closed container.  Keep watching the video till the end, as it actually makes for a neat demo in itself.

The demo for the week was the silver nitrate/copper wire demo.  Silver nitrate (nitrate being the conjugate base of nitric acid… which is how I’m relating this demo to lab this week…) dissolves readily in water to give a solution of silver nitrate.  Just about everything silver nitrate touches gets stained black.  Not immediately… only after exposure to light.  For this reason, silver nitrate used to be used in early photography.  No stains for me, though it is always a concern.

Dropping copper wire into the silver nitrate solution initiates a redox reaction between the silver ion and copper metal.  The silver is reduced to elemental silver and the copper is oxidized to copper(II):

Cu(0) + 2AgNO₃ = Cu(NO₃)₂ + 2Ag(0)

The silver crystallizes at the surface of the copper and the copper wire quickly becomes coated with a bunch of elemental silver.  At the same time, the copper ions go into solution and the colorless solution turns a characteristic blue as the concentration of copper ions builds.  It’s a pretty dramatic demo of a neat redox reaction.  I was testing the speed of the reaction before I went to lab (silver started to become visible within 30 seconds to a minute), and I ended up leaving the copper wire in the silver nitrate while I went to lab.  I came back several hours later and the crystals had plenty of time to grow and were very nice looking.  I let it go overnight to see how big the crystals would get.  I took a picture when I got to lab the next morning (click for larger).  And then I collected the silver in a scintillation vial to take home with me.  The pictures are from my experience, Noel posted a picture a while back as well, and there is a nice video below.

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Chemisry Lab Demonstrations: Nylon Rope Trick

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This week’s lab involved taking the unknown carboxylic acid from experiments 1 (acid/base extraction) and 2 (purification by recrystallization) and converting it to a primary amide. The acid was refluxed in neat thionyl chloride and added dropwise to 30% ammonium hydroxide. Extract with dichloromethane and evaporate.  The fumes from this lab are pretty nasty.  The reagents are bad enough – the lab smelled like ammonia all day – but sulfur dioxide and HCl gas are liberated during the reaction.  The “fume hoods” in the ugrad labs aren’t so hot, so the students attach a long-stem funnel (where do I sign up to be a glassware photographer?!) to the water aspirator and invert the funnel over the reaction mixture as it refluxes.  This collects the fumes as an extra “mini fume hood” during the reaction.

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Chemistry Lab Demonstrations: Homemade Breathylizers…Sorta

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Undergrads were on spring break last week, so no lab last week.  This week, we performed the Jones oxidation of isoborneol to camphor.  The crude product was put in the bottom of a Hirsch funnel with a cold finger, and the product was purified by sublimation.  One of my favorite chemistry trivia facts is that the opposite of sublimation is deposition.  Now you know.

I must say, I am really, really disappointed with the way this week’s demo turned out.  Plan A was following a patent prep to immobilize Jones reagent on silica gel.  That actually worked really well.  No problems there.  I had a nice orange granular solid.  The next step was to pack it into 5″ pipets between glass wool and plain silica gel to create a tall, narrow column of Jones reagent.  That also worked really well.  I meant to take a picture of the setup before I dumped everything, but I was pretty down after it didn’t work and just poured it all into my chromium waste bottle.  Sorry.

Anyway, the plan was to ask if anyone came to lab drunk that day (no one admitted they did), then have someone volunteer to swish around some Listerine for a while.  Then attach a clean drinking straw to the pipet and blow.  See, Listerine is 21% alcohol, so there’d still be some ethanol vapor in the breath which should flow past the Jones reagent.  The Jones reagent will oxidize the ethanol to acetic acid and will itself be reduced.  The reduced chromium reagent is greenish brown.  With the pipet system, the orange color would slowly change over to green from the start to the finish.  The amount of ethanol in a person’s breath (and therefore the BAC) can be determined by seeing how far up the column the green color extends.  The more solid that is reduced to green, the more wasted the person.

I know there are several potential safety issues with this setup.  Chromium is toxic and shouldn’t be ingested or released into the environment.  The goal was to have only me handle the glass and only the student with clean hands touch the drinking straw – which was not ot be unwrapped until just before use.  When I was testing this, I don’t know what was going on, but I could get no air pressure through the pipet.  With breath not flowing through the pipet, there’s no chance of the Jones reagent being reduced, so no demo.

See, they really did used to use Jones reagent in breathalyzers.  They don’t anymore because chromium is toxic, and well, lots of things can reduce chromium… not just ethanol.  This leads to false positives.   Now they use a fuel cell for better results.  The best part of the Jones reagent story is that drunk people would blow through a solution of Jones reagent, which would reduce some of the chromium.  A UV/Vis detector could measure the absorption of the solution before and after the test.  The absorption is related to the amount of ethanol because the measured absorption (A) is equal to the product of the concentration of chromium (c), the path length of the cell (l), and a constant unique to chromium (ε).  A =εcl.  This relationship, ironically enough, is known as Beer’s law.  That’s right.  Beer’s law can be used to tell how drunk a person is.

Plan B was to have a row of test tubes with a solution of Jones reagent to which was added increasing amounts of ethanol.  This would gradually change the color from yellow to greenish.  A separate test tube with a solution of Jones reagent in lab would have an arbitrary amount of ethanol added to it.  It could be titrated against the standards to show the color change and determine the level of drunkenness.  Problem is the color change is very subtle in dilute solutions.  Too subtle to really see.  No demo again.

I won’t bore you with plan C, but suffice it to say – fail.

So finally, plan D.  Same as plan B, but with a surrogate for Jones reagent… and a surrogate for ethanol.  Basically none of the actual reagents, but would still give a color change.  I went with acid/base chemistry.  A row of test tubes was filled with an acidic solution of bromocresol green.  The acidic solution is yellow.  To each tube was added an increasing amount of sodium hydroxide so the color would gradually change to a deep blue.  The first test tube was yellow, then gradually orange, then muddy brown, then blue.  In fact, the series hit just about every color EXCEPT green (the image looks better.  Mine was not nearly so nice.  I was in a hurry).   Anyway, then an arbitrary amount of NaOH was added to a separate tube with an acidic solution of bromocresol green and titrated against the standard.

It wasn’t a great demo, but it looked like it was supposed to look.

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Chemistry Lab Demonstrations – Silver Mirror

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This week’s reaction was the HWE-modified Wittig reaction (and here) between para-anisaldeyhde and triethyl phosphonoacetate.  Saponification, followed by recrystallization of the acidified product gives nice needles of p-methoxycinnamic acid.  Funny thing, though, p-methoxycinnamic acid is supposed to be a white solid.  Half of the lab had their reactions turn bright yellow.  Good times.

No exicting Wittig demonstrations that I can think of, but since we used an aldehyde as a reagent, I showed off Tollens’ reagent and its utility in qualitative determination of an aldehyde.  Tollens’ reagent is a solution of silver nitrate, ammonia, and potassium hydroxide.  This produces a silver diamine complex [Ag(NH₃)₂]⁺. This is important because simple aqueous silver(I) is much easier to reduce than the silver diamine complex (RSC).  Silver nitrate on its own would be reduced rapidly and colloidal silver would develop.  These more controlled conditions allow for the controlled formation of elemental silver.

The basic solution is mixed with a sample of an aldehyde and shaken.  The silver oxidizes the aldehyde (but not a ketone!) to a carboxylic acid.  Simultaneously, the silver is reduced to elemental silver.  This production of elemental silver slowly coats the flask with a thin layer of silver.  This qualitiatively proves the presence of an aldehyde.  More importantly, it makes a very pretty silver-coated flask.  If rinsed thorougly with water, the flask can be capped and kept indefinitely.  I gave the one I made in the demo away to the student with the highest percent yield for the previous lab.  Here’s the reaction that takes place:

CH₃CH₂CHO(aq) + 2[Ag(NH₃)₂]⁺(aq) + 3OH^-(aq) → 2Ag(s) + CH₃CH₂COO^-(aq) + 4NH₃(aq) + 2H₂O(l)

The demo is often called the sliver mirror because the layer of silver, well, looks like a mirror.  In fact, this is one way mirrors were actually made historically.  A sacrificial adehyde (such as glucose) was poured over glass submerged in a solution of Tollens’ reagent.  The reaction takes place and deposits a nice layer of silver on the glass.  Wash and you have yourself a mirror.

Here’s the prep I used for preparing the demo, but here are some others if you want more references (I like that last one).  And, of course, a video (not me).

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Chemistry Lab Demonstrations – Dehydrating Sugar

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This week, the organic chemistry lab I’m TA’ing dehydrated 2-methylcyclohexanol with phosphoric acid to give a mixture of 1- and 3-methylcyclohexene.  The product ratio was determined by GC analysis… At least it was supposed to be.  The students for the most part did a good job of working through the lab quickly and efficiently, then it was upstairs to the GC room.  about halfway through the GC runs… the power went out.  That of course means the GCs shut down, the voltage sensor stops working, and the data no longer transfers to DataStudio.  Oh well, it’s the thought that counts, right?

While I’ve had some trouble finding demonstrations to fit some of the labs this semester, I had no such problem this week.  I knew exactly what I wanted to do – dehydrate sugar with con’c sulfuric acid (more).  The sulfuric acid easily dehydrates the carbohydrate (i.e. the hydrate of carbon) to give essentially elemental carbon and water.  Oh, and a lot of heat.  The water escapes as steam and the expanding gas (as well as some sulfur dioxide) aerates the residual carbon.  The black mess rises out of the beaker like it’s possessed.  It pretty cool (although kinda gross) looking.

I used 60-mL of ordinary table sugar in a 100-mL beaker.  30 mL of concentrated sulfuric acid (without stirring) gave an approximately 12 inch column of carbon after about 2 minutes.  The video below is a good representation (although, I think it might be in time-lapse mode.  My experience has never formed the column that quickly).  I appreciate that the person in the video is using gloves… but I would not recommend trying this one on your stove at home.  The fumes are not good for you.

There was a J. Chem. Ed. article from Willamette University in Oregon out a few years ago about a variation of the dehydration without using concentrated sulfuric acid.  Instead, they combine the dehydrating sugar demo and the flaming gummy bear demo and use potassium chlorate.  the potassium chlorate/sugar mixture is wetted with ethanol and the mound is ignited.  The same effect occurs, but without the need for concentrated sulfuric acid.  It’s not clear that the fumes are any more safe, but the need for con’c acid is avoided.  Here’s a good description with pictures from the University of Minnesota.

Finally, one of my students asked if this is the same principle behind the Black Snake fireworks.  Eh… essentially yes, but not exactly.  It seems that for the homemade versions at least, a mixture of sugar and sodium bicarbonate are ignited in ethanol.  The combustion produces heat and gas and a carbonaceous black solid.  The solid is again aerated with the expanding gasses and the snake is formed.  So it’s the same, but different!

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Chemistry Lab Demonstrations: Upsidedown Thunderstorm

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S_N2 lab today.  Preparation of (2-methylphenoxy)acetic acid from o-cresol and sodium chloroacetate.  NaOH deprotonates cresol.  Add sodium chloroacetate and reflux.  S_N2 reaction occurs.  Acidify, collect the precipitate, and recrystallize from water.  Unfortunately, today’s lab took a really long time.  I’m not sure why, but it did.

Anyway, I looked around for cool S_N2 demos… and there really aren’t any.  Sure, I can do the Finkelstein at different concentrations and show that one precipitates NaCl faster than another because it is more concentrated.  Woo hoo.   That’s not very exciting.  So I ultimately decided that today’s lab just isn’t going to have anything to do with the S_N2 reaction.

I decided to do the upsidedown thunderstorm.  That’s my name.  I’ve seen ‘thunderstorm in a test tube,’ but that’s not a very flashy name, now is it?

Here’s the deal:  Add a layer of con’c sulfuric acid to a test tube (use all proper precautions for handling con’c acid!).  Slowly and carefully add ethanol down the side to create an ethanol/acid biphase (if you leave it sit too long, the acid and ethanol will mix, so don’t let that happen).  Add a few crystals of finely-ground potassium permanganate.  The reaction occurs, and evolves a gas.  The gas bubbles through the ethanol layer and looks like upsidedown rain.  After a few seconds, “lightning” appears at the phase boundary, and the sulfuric layer becomes cloudy and purple (the storm clouds…).  The ’storm’ continues for several minutes.  Manganese waste should probably not be thrown down the drain, so if you do this, dispose of it according to local regulations.

There appears to be a few different processes leading to ‘lightning’ formation.  Permanganate is converted to the exposive Mn(VII) oxide.  Additionally, ozone is created.  The ozone oxidizes the ethanol resulting in combustion of the organic material.  Wikipedia explains:

Concentrated sulfuric acid reacts with KMnO₄ to give Mn₂O₇, which can be explosive … Potassium permanganate and sulfuric acid react to produce some ozone, which has a high oxidizing power and rapidly oxidizes the alcohol, causing it to combust. As a similar reaction produces explosive Mn₂O₇, this should only be attempted with great care. An approximate equation for the ozone formation is shown below.

6 KMnO₄(aq) + 9 H₂SO₄(aq) → 6 MnSO₄(aq) + 3 K₂SO₄(aq) + 9 H₂O(l) + 5 O₃(g)

When I was practicing the demo, I wanted to see if I could scale it up some to make it more visible than just in a test tube.  I tried it with about 20 mL each of the acid and alcohol in a beaker.  The detonations were sufficiently energetic (and perhaps the ethanol layer was not tall enough) that the ethanol caught fire.  I had a cute little ’sterno’ burner going.  For the actual demo, I switched to an Erlenmeyer flask which helped.  No fire this time.

Storm In A Test Tube !

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Chemistry Lab Demonstrations: Sodium Acetate Crystallization

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Recrystallization lab today.  Standard two-solvent recrystallization of an unknown carboxylic acid from ethanol/water.  My students did awesome – as far as I know everyone obtained crystals on the first try.

The demo today was crystallization of sodium acetate from a supersaturated water solution.  I had to try a few times to get a good supersaturated solution.  First, I heated up some water and tried adding anhydrous sodium acetate one scoop at a time.  That didn’t really work – it was hard to get it all into solution, and when I took it off the hot plate, it ended up crystallizing while cooling.

Then I tried loading up a round bottom with the crystalline trihydrate and spinning it on my rotovap while the hot water bath was turned up real high.  That didn’t work at all.  I was real surprised it didn’t work, but nothing went into solution.  Sodium acetate trihydrate “melts” at around 55 degC (no, contrary to some websites, this is NOT molten sodium acetate.  Anhydrous sodium acetate decomposes above  324 degC).

What ended up working for me was pouring a small layer of the trihydrate into an Erlenmeyer flask and warming that on the hot plate.  The solution was formed, and more trihydrate would be added portionwise.  Then I would add a few small scoops of the anhydrous just to get a bit more sodium acetate into solution.  This formed a nice, clear, colorless solution of sodium acetate.  It could be cooled without auto-crystalization to create the supersaturated solution.

See, at 0 degC, sodium acetate has a water solubility of 760 g/L.  Hot solvents dissolve more than cold solvents, so heating the water allows one to dissolve more than 760 g of sodium acetate in a liter (not that I was working on that large of a scale).  If cooled slowly and without addition of a nucleation site (i.e. dust particle, etc), a supersaturated solution is formed.  That is, if cooled to 0 degC, more than 760 g/L would be dissolved in water.

This supersaturated solution is just begging to crystallize.  You don’t have to coax it very much.  Even the introduction of one tiny seed crystal or dust particle can be enough to get sodium acetate to crystallize.   And once it starts crystallizing… it takes off.  It crystallizes much faster than typical crystals.  So fast, in fact, that it can be used as a demo.  You can either drop a seed crystal at one side of the flask and marvel at how fast the crystals race to the other side, or you can slowly pour it onto a seed crystal and slowly build yourself a Lot’s wife-worthy pillar of salt.

I made two supersaturated solutions, so that I’d have one as a backup.  It’s a good thing, too, because my backup solution auto-crystallized on my walk over to the ugrad lab.

One final note.  The crystallization is exothermic.  This is the exact process used in hand warmers for gloves and such.  there is a disk in a supersaturated solution of sodium acetate.  When the disk is snapped, the solution crystallized, giving off quite a bit of heat.  This warms your hands.  The temperature rises to about 130 degF (around 55 degC).  The second video shows the temperature rise.

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Chemistry Lab Demonstrations: LIQUID CO₂ Extraction!

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It’s the extraction lab this week in the OChem lab I’m TA’ing. It’s a straightforward aqueous base extraction of an acidic unknown from a neutral impurity. Acidify, filter the precipitate, and you’re done. I was trying to come up with a demonstration for the lab. I thought about extracting caffeine from coffee or tea leaves, but that would take a while, and isn’t all that visually appealing.  I’ve only got a few minutes in my pre-lab lecture time.

So I looked around for a while, and finally found this paper by James Hutchison from the University of Oregon (doi:  10.1039/b405810k).  They suggest a new lab for undergraduates involving the extraction of D-limonene from orange peels using liquid carbon dioxide.  That’s right, I said liquid carbon dioxide.

The premise: create a removable filter using copper wire and filter paper to jam into the bottom of a disposable centrifuge tube.  Add grated orange peel.  Add crushed dry ice.  Cap the centrifuge tube tightly (but not TOO tightly! The tube needs to be able to vent so as not to EXPLODE!) and immerse in warm water (T = 40-60 degC).  The pressure rises (naturally) and the temperature increases and you jump into the liquid portion of carbon dioxide’s phase diagram (click for larger)

The liquid carbon dioxide percolates through the orange peels and extracts the limonene.  the oil-in-solvent mixture drains through the filter paper to the bottom of the centrifuge tube.  If you leave the tube in the water long enough, eventually the liquid all evaporates and the pressure decreases.

The goal is that the evaporation of the carbon dioxide leaves the pure oil at the bottom of the tube.  The authors mention that for approximately 2.5 g of freshly-grated orange peel, 0.1 mL of oil should remain after 3 carbon dioxide extractions.  They note this is a yield comparable to typical organic solvent extraction or cold pressing.  I did one extraction on day-old chopped orange peel and did not isolate any oil whatsoever.  Not a drop.  I’m a little disappointed by that, but not really.  It’s still an ok teaching point for the students.  Not all experiments work all the time. I could examine my starting materials and get better quality reagents and it might work.

Now, inside the tube I don’t think we were past the critical point.  I don’t think the temperature inside the centrifuge tube actually makes it up to the temperature of the surrounding water.  I say this because after the examining the tube after the experiment, the orange was cold and there were ice crystals in the tube.  There are two possible explanations for this.  One, the temperature inside didn’t make it past the critical temperature.  Two, when I opened the tube after the experiment, some non-trivial amount of pressure was released.  PV=nRT tells us that a sudden drop in the pressure simultaneously lowers the temperature, and I could have frozen the water out that way.  In fact, the authors note that while exact temperature and pressure readings are impossible with this simple setup, they speculate that the conditions approach the triple point.

In any case, it was a very cool experiment to watch, even if it didn’t do what it was supposed to.  Pictures below.  These pictures are from Monday night when I was practicing the demonstration.  It looked much cooler in person.  The first shows the system when first submerged in the water.  The second is about 15-30 seconds later.  It’s hard to see, but if you look closely, all three phases are apparent in the system.  The third is after the dry ice has completely liquified.  Click for larger.

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