Post Tagged with: "Chemistry Demonstrations"

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.

jones

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.

bromocresolgreen

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.

By March 19, 2009 2 comments fun, synthetic chemistry

Chemistry Lab Demonstrations – Silver Mirror

hwe

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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).

By March 5, 2009 6 comments fun, Uncategorized

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