Articles by: azmanam

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!

By February 25, 2009 5 comments fun

Chemistry Lab Demonstrations: Upsidedown Thunderstorm

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SN2 lab today.  Preparation of (2-methylphenoxy)acetic acid from o-cresol and sodium chloroacetate.  NaOH deprotonates cresol.  Add sodium chloroacetate and reflux.  SN2 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 SN2 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 SN2 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 !

By February 17, 2009 12 comments fun

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.

By February 10, 2009 2 comments fun

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.


By February 4, 2009 31 comments fun