Tuesday, December 13, 2011


Thanks to my good friend and fellow science teacher, Josh, who makes the high speed videos.
Who doesn’t love the thermite reaction?  This reaction is exciting, dramatic, and a bit scary.  Shakhashiri describes it like this, “Flame, flying sparks, smoke, and dust are produced.  Molten iron runs through the hole in the pot into the sand bath.”  I pulled out this demo today as part of the “5-types of chemical reactions” unit and a great example of an exothermic reaction.  Thermite is a single displacement reaction with a high activation energy.  The thermite mixture, aluminum powder and iron oxide, can sit on the shelf quietly without much concern because it takes another reaction to get it started. 

Inserting the ignition stick into the thermite mixture.
I tried to use the potassium permanganate reaction with glycerin (as described by Shakhashiri) to get the reaction going.  I did the potassium permanganate reaction yesterday in class, so it seemed like a nice follow up today.  (By the way, it filled up the room with smoke and we all poured out of the classroom coughing as the bell rang at the end the period.  I’ll do it outside next time!) The permanganate reacted vigorously with the glycerin, but it did not produce enough heat to spark the thermite.  (You can watch this “dud” reaction at the end of the “Thermite Balls” video.)

Take a look at the chunk of iron that we collected from the water bath.
A close-up of the iron drops that formed from the reaction.
So I went back to the stand-by:  thermite ignition sticks that we bought from the chemical supply company.  These sticks are essentially really fat sparklers.  I used a Bunsen burner to light the stick, and then I ran outside with my burning sparkler.  I held it in the thermite mixture until the reaction started.  Once it got going, all that was left to do was stand back and enjoy the show.  Watch for the molten iron dropping out of the bottom of the clay pot, making the water in the tank boil.  I fished out the iron pellet from the tank to show the kids the iron drops that formed.  

I followed up the thermite reaction with what I call “hand held thermite” or “thermite balls”.  I can make the same reaction on a smaller scale with two rusty cannon balls, one covered with aluminum foil.  With enough activation energy, provided by hitting the balls together at a high speed, sparks fly and there’s a loud pop.  The high speed video doesn’t do justice to the reaction without the sound effects, but you get to see the classic facial expressions I make when I do the demo.  I never realized how much I flip my hair until I started this blog project.

Ready for the hand held thermite reaction.

Tuesday, December 6, 2011

The Exploding Pumpkin Demo

Here's my version of the Exploding Pumpkin Demo.


Thanks to my friends who inspired me to try this demo by posting a video of a another chemistry teacher doing this on my facebook page.  My version is a small scale, "cute", look at the reaction of calcium carbide and water.  The mixture produces acetylene gas, which is flammable.  The pumpkin gets filled with the acetylene gas, if you wait long enough, and makes a nice explosion.  This demo was a good way to say goodbye to Halloween

Monday, December 5, 2011

Heat of Dilution of Sulfuric Acid

Measuring heat of solution.

My thoughtful students are contemplating their data.
Calorimetry is not exactly a flashy subject.  Heat exchange into and out of water does not make for a showy demo.  However, today I did a demo for my honors class of the heat of dilution of concentrated sulfuric acid, Demo 1.6 in volume 1.  To my surprise, this demo offered an element of excitement that I was not expecting (why am I surprised that the demos from Shakhashiri are exciting?).  I decided to run three trials at the same time, dissolving 10, 20, and 30 mL of concentrated sulfuric acid in 100mL of water.  I used one of my Lab Quests from Vernier with 3 temperature probes so we could watch the temperature readings graphically projected on the big screen.  I asked for three volunteers; the hands went up immediately.  Three boys came up to pour in the acid and help monitor the reactions.  I counted to three and then they all poured at the same time.  The exciting part was the 30 mL sample, which produced enough heat to boil the water and overflow out of the cup all over the bench.  It was a surprise to us all to see this happen.  The spike in the temperature readings told the whole story of the heat produced from diluting sulfuric acid in water.  (Do as you "oughter", add acid to water.)  We followed up with the calculation of the heat produced in each cup, which led to the conclusion that this is a linear relationship.  The demo ended with me pouring baking soda onto the spilled acid, producing the classic “volcano” reaction on the bench.  Welcome to calorimetry, kids.

"What, no flames today?"

Taking his temperature with the Vernier probe.

Saturday, November 5, 2011

Making Water

The base is in the 1-L volumetric flask.

Adding the acid to the base
More than 1 liter here!
Shakhashiri has all the answers.  I like to do the demonstration of a neutralization reaction with equal volumes of hydrochloric acid and sodium hydroxide.  The reaction produces salt and water.  If you mix 500 milliliters of 2-molar solutions of the acid and base, you will produce 18 mL of water (1 mole).  I learned about this demo at ChemEd, a great conference for chemistry teachers.  In previous years, I used a 1-liter graduated cylinder, my favorite demonstration glassware, for this reaction.  Although it always worked, the additional volume of water was very difficult to see.  Shakhashiri has a description of this same demo in volume 3, Demo 9.17, with the solution to my problem.  He uses volumetric flasks!  Of course!  Today, when I combined the 500 mL of acid and base, both made up in 500 mL volumetric flasks, the additional water was easily visible in the neck of a 1-liter volumetric flask.  I added some universal indicator to the acid and base solutions to give the demo some color.  This was an extra bonus because the two solutions formed layers when I poured them together the 1-Liter flask.  We all enjoyed the interesting effect of the blue and pink layers.  When I mixed the flask, the volume increase caused pressure to build up in the flask and the stopper popped off a couple of times.  Overall, the Demo was a great way to show the products of a neutralization reaction, including the often-overlooked water molecules that are formed.
Releasing the pressure.
Admiring the interesting layers.

Saturday, October 29, 2011

The “Mini Grain Elevator” Demo

Take a look at this one!

I have wanted to do this demo ever since my second year of teaching.  I have read the instructions many times, found in Volume 1, Demo 1.41 of Shakhashiri’s Chemical Demonstrations.  Every time I have read them, I would think, “I’ve got to try this, but it seems scary.”  The set-up always gave me pause.   I put it off until I “had more time”; just another excuse to delay plunging into this experiment.

The Flinn Kit
At the end of last year I was looking through the Flinn catalog for my summer supply order.  The kit for the Mini Grain Elevator caught my eye.  I decided to get the kit and give it a try; surely a kit would be easier and the experts at Flinn have tested it.   I was still too scared to do it by myself, so I dragged my ever-willing colleague into the fun.  We set up the quart sized paint can with a pipette full of lycopodium powder. Just light the candle, secure the lid of the can with a hammer, and then squeeze the pipette into the sealed paint can.  Simple, right?  The instructions for the kit say, “Stand back as far as possible as you squeeze the pipette.”  This didn’t put us at ease. On the first try we got a big bang, the lid flew off, it was exciting.  But the pipette rig was not reproducible.  Attempts two and three were duds, nothing happened and the candle went out.  I decided to go back to “the bible” and, with the support and enthusiasm of my colleague, together we spent the afternoon building our Mini Grain elevator rig.

Here are a few pictures from the construction of our Mini Grain Elevator.   We bent a funnel for the inside of the can, punched a hole in the paint can, scrounged around the science building for a rubber bulb, and put the whole thing together in a 1 gallon paint can.

This rig was awesome.  It made a loud explosion, the lid flew off the can, flames shot out of the top of the can, and the candle flew out of the can.  Attempts two and three were equally impressive.  Making the rig was worth the effort! My colleague and I spent a couple of days setting off the Mini Grain Elevator for all of our classes and for any other visitors to the lab.   It was really fun every time we ignited the can.  The physics guys next door and the biology teacher down the hall came in to watch it go off several times too!  My husband even made the trek from the next building to watch the famous Mini Grain Elevator explode.  My reputation as a closet pyro was confirmed with this demo.
Bending the stem of the funnel.
Second attempt.  We broke the first funnel.

I used my awl to make a hole in the paint can.
A look at the inside of the can.

Attaching the funnel to the can.
Loading the lycopodium powder.

Friday, October 21, 2011

Demo Show for Pack 26

This week I gave a demo show for my sons’ Cub Scout Pack. 

Pouring carbon dioxide onto a candle.
I started off the show with the Genie in a Bottle trick, a ChemFax by Flinn Scientific (Publication No. 91200).  I wrapped a 2-liter plastic bottle in aluminum foil to make the Genie’s bottle.  The “Genie” was a reaction of 30% hydrogen peroxide and sodium iodide, the catalyst in the decomposition reaction of hydrogen peroxide that produces water and heat.  The sodium iodide crystals were suspended in a small sack at the top of the bottle held in place by a rubber stopper.  When I took the stopper off the bottle, the Genie escaped in a big puff of steam, and left behind a shrunken plastic bottle. 

For my second demo, I asked one of the Tiger Cubs to come up and fill up a beaker with Styrofoam packing peanuts.  The peanuts “disappeared” as fast as he could put them into the beaker.  I encouraged him to try harder, but the peanuts kept shrinking before his eyes.  What he didn’t know is that I had poured some acetone, a clear and colorless organic solvent, into the beaker before the show.  The small amount of acetone quickly dissolved the packing peanuts as he dropped them in the beaker.

 Blue water for displacement.

Next I did a water displacement demonstration with baking soda and vinegar, one of the many reactions that produce carbon dioxide gas.  I like doing this demo for kids because it’s a different look at the classic “volcano reaction”.  Taking the advice of Dr. Shakhashiri, I used large apparatus to make the reaction visible to the audience.  I collected the gas in a 2-liter bottle, using a 5-gallon fish tank for the water displacement.  I had a volunteer mark the water level on the tank before the reaction so we could see any change.  I put food coloring in the water in the bottle so there was another visible change for the kids to observe during the demo.  Once we filled the bottle with carbon dioxide gas, I used it to put out a candle.  I love to show kids that gases can pour, even when we can’t see them.

Heating the water in the can.

Ready to flip it into the ice water bath.
The finale was the classic “Can Crushing Demo”.  I heated a small amount of water in an empty soda can until steam was visible.  I quickly inverted the can in an ice water bath.  The can crushed with a shocking smash.   I learned how to do this demo during my summer workshop at Longwood College.  I was hooked on chemistry demos after performing this one for my class that fall.

The can crushed in the ice water.

Thanks to Pack 26 for inviting me for a fun night of science.

Friday, October 7, 2011

Oooo That Smell…

This sign was posted by my physics colleague today.

I really don’t try to make my colleagues hate me, I swear.   

Today we did the classic Iron and Sulfur experiment.  This is a full assault to the senses, most notable the olfactory.  In this experiment we compare the physical properties of iron, sulfur, and a mixture of iron and sulfur.  Then we put the mixture in a test tube and heat it vigorously, which causes a chemical reaction between the two elements. 
The iron and sulfur reaction is very exciting to watch.
The reaction takes a few minutes to get going.  At first, the reaction mixture smolders a bit; some of the sulfur comes off as a gas, and condenses on the sides of the test tube as a red and yellow coating.  Then a red glow begins to form in the test tube, which spreads until it looks like lava inside the test tube.  It actually produces enough heat to sustain the molten-looking reaction for several seconds without the Bunsen burner flame.  

Notice the puff of sulfur coming out of the test tube.
Plunging the test tube in water is all part of the fun.

 For the finale, the students plunge the test tube into cold water.  The reaction mixture gives off a big puff of steam, accompanied by a loud hissing sound, and some of the water boils around the black product pellet.  The hot glass cracks, releasing their sample of iron sulfide into the beaker of water.  At this point, everyone in the room is too excited about the flames, the molten product, and the breaking glass to notice the strong sulfur smell.  Of course, the students in the neighboring physics and biology labs are not amused when the odor works it way throughout the building.  

 Sorry guys, but this one is totally worth it!


Saturday, October 1, 2011

A Visit with Dr. Shakhashiri

What a thrill to meet Dr. Shakhashiri again!
I’m still giddy from meeting Dr. Shakhashiri again after all these years.  I am so thrilled to have had the chance to thank him for changing my life.  Okay that sounds really dramatic, but it’s true.  Without that workshop at Longwood College in the summer after my first year of teaching, I would have quit teaching chemistry.

Dr. Shakhashiri is dropping dry ice into these 1-liter grads.
Today I enjoyed another inspiring demonstration show from Dr. Shakhashiri at the 100th year celebration of the Connecticut Valley Section (CVS) of the ACS.  You can see where my love of the 1-liter graduated cylinder originated!  After every demonstration he performed, I turned to my colleague and said “We’re totally doing that one this year!”

 Watching Dr. Shakhashiri perform his demos is not a passive activity.  He engages every member of the audience with a steady stream of questions.  At the beginning of the talk Shakhashiri said, “Chemistry is the science of the familiar”.  He reinforces this statement with his demonstrations by relating the chemistry to everyday experiences.  He is a master at drawing out all the chemistry that his audience already knows but they didn’t know was chemistry.

Dr. Shakhashiri pours boiling water over dry ice to make fog.
Today feels like the right moment to tell the story of my second year of teaching.  I returned to my school in rural Virginia armed with new ideas and filled with determination to “do it right” this time.  I opened the year with an exploding hydrogen balloon on the first day.  My students were excited to come to class on day two.  I implemented a new micro scale lab curriculum in my classes, another part of the workshop at Longwood.  I tried some new classroom management techniques that I learned from long conversations with the other experienced teachers at the workshop.  I maintained contact with several of these teachers my second year to help me through the rough spots.  It was a completely different experience; I fell in love with teaching that year.  

Check out Dr. Shakhashiri's website.

Friday, September 30, 2011

Are You Feeling Dense?

My beautiful density column.
Density is an introductory concept that comes up several times throughout the year in chemistry class.  Every year I have this “what, you mean you don’t know this” moment with my students.  I know that they have measured density many times throughout their school careers.  Eighth grade science could be renamed “An In-depth Look at Density”.  But my chemistry students give me the blank stare when I start talking density.

I have a set of classic density demos that I enjoy doing for my kids to investigate the relative density of solids, liquids, and gases.  This year I decided to stretch my repertoire with a multilayer density column. I used Demo 9.2 from Shakhashiri’s Volume 3 as my inspiration, along with two of my colleagues who do variations of the density column. I decided to use glycerin, water, antifreeze, corn oil, and ethanol for my column.  Adding a little food color to the water and the alcohol gave me five nicely colored layers.
 Liquids and objects for the column.
Look for the square piece of plastic in each beaker.

My plan was to construct the density column in a 1-liter graduated cylinder (my favorite piece of glassware for demos) with an object floating on each layer.  I did some tests of the liquids and objects in beakers to determine what would sink and what would float.  If you look closely you can see that the plastic square sinks in the antifreeze but floats on the glycerin.  Perfect!  An object at the intersection of the bottom two layers.  I tried a plastic piece from my molecular model kit; it floats on the water but sinks in vegetable oil.  I left my oak sample at home, but I expected it to float in the oil but sink in the alcohol.  Top it all off with a cork that floats on the alcohol. 

I’m feeling pretty confident as I begin to construct my column.  I grabbed the oldest sample of glycerin from the stock room for my demo because I don’t need to use it for a chemical reaction.  It turned out that the old glycerin has a great amber color, which made the column that much more interesting.  First I poured in 200 mL of the golden glycerin (density = 1.25 g/cm3).  Next came the antifreeze, made mostly of ethylene glycol (density 1.11 g/cm3) with a great neon green color.  As I poured in the third liquid, water with some red food coloring (density = 1.0 g/cm3) it mixed with the antifreeze.  Wait a minute; the antifreeze was soluble in water.  Of course!  When you take one look at the structure of ethylene glycol, with an –OH group on each end, then it’s obvious that this polar molecule will feel right at home in an aqueous solution. 
 Here's the antifreeze dissolved in the red colored water.  No good.

So I started again without the antifreeze.  Glycerin, red colored water, then vegetable oil (density = 0.91 g/cm3), and I topped it off with some blue colored ethanol (density = 0.789 g/cm3).  Beautiful!  The layers kept to themselves very nicely.  Then I had a lot of fun dropping objects into the column.  Sadly, the oak piece floated on top of the alcohol.  The black rubber stopper dropped to the bottom but the white one stopped on top of the glycerin.  The visual impact of this demonstration is worth 1,000 notes on the topic.

Here’s a run-down of my favorite density demos that I use every year.
Diet Coke floats, regular Coke sinks.  Constant volume, different mass.
The floating egg!  The egg is resting on a layer of brine, tap water is on top.

Soap bubbles float on a layer of carbon dioxide gas.  This one is really exciting to see.

Friday, September 23, 2011

Don’t Go Changin’ …

The beginning of the year is wide open for fun demonstrations of introductory concepts.  I like to capitalize on this opportunity to keep my kids engaged and have some fun with demos that do unexpected things.  The topic of the day was physical and chemical properties.  I decided to try to boil water in a paper cup to demonstrate the high heat capacity of water, a characteristic physical property of water.  This demo is described in Demo 9.4 Boiling Water in a Paper Cup:  Heat Capacity of Water found in volume 3 of Shakhashiri’s Chemical Demonstrations, A Handbook for Teachers of Chemistry.  First I put an empty paper cup in the flame.  It did just what everyone expected, it burned quickly.   I filled the second cup half way with water.  When I placed it in the burner, the top part of the cup caught fire and burned off, but the bottom portion that was in contact with the water remained unaffected.
The top portion of the cup is burning while the bottom half stays intact.

I heated the water in the paper cup to the boiling point and then proceeded to boil it all off in the paper cup. 

The water is boiling in the Dixie cup!

I was very excited to see the water boiling in the cup.  As a bonus, I got the chance talk about the phase change happening to the water.  The paper cup provided a great visual display that the water, not the cup, absorbed the energy from the burner.
 The "after" photo of the cup.

How do you narrow down the possibilities for demonstrating chemical properties?  There are so many great demos to choose from for this one, but in the spirit of this blog project I decided to try something new.  In my email inbox I found my inspiration from Flinn Fax!, Vol. 11-4; an excellent publication from Flinn Scientific Inc. with demonstration ideas that are timely for the curriculum and the season.  “The Carbon Soufflé, A Sweet Exothermic Reaction” was on the cover.  I did this reaction in year two of my teaching career, I think.  Let’s just say it’s been a long time since I’ve done this one.  I poured 18 molar sulfuric acid (highly concentrated acid) on sugar (sucrose).  It worked great!  After a short delay, the reaction began to boil, it produced heat, and the carbon soufflé emerged from the beaker.  You can watch the video of the reaction to see how cool it is and to hear the voices of my students as they watch this unexpected event.  I made one mistake in my explanation, which you will hear on the video.  The reaction produces carbon and water.  The carbon column grows from the beaker as water vapor forms gas bubbles in the hot carbon.  At one point in the demo I said that carbon dioxide is also formed in the reaction, but after further reading, I realized that this is not the case.  Sugar is a carbohydrate made up of carbon, oxygen, and hydrogen in an interesting double ring structure.  The sulfuric acid causes a dehydration reaction of the sugar, which removes hydrogen and oxygen from the sugar molecule to form water and carbon.  In the end, I had a large carbon “popsicle”. 

Friday, September 16, 2011

Simple Distillation is Not So Simple

 Simple distillation is a deceptive name.  This week I conducted a demonstration of a simple distillation for my honors class as part of the Foul Water Lab from the ChemComm curriculum.   (I don’t use ChemComm, but I like this lab from the book.)  I used Demo 9.10 Separating Liquids: Fraction Distillation found in volume 3 of Shakhashiri’s Chemical Demonstrations, A Handbook for Teachers of Chemistry as a reference for the set-up and discussion. 

Here's the simple distillation apparatus.
The distillation demo was the final step in a water purification lab conducted by my Honors students.  They were challenged to purify a water sample that contained oil, coffee grounds, garlic powder, and salt.  I used the “Foul Water Lab”, which is an excellent introduction to separating mixtures, chemistry lab techniques, and the physical properties of matter.  The students did a series of purification steps:  separating the oil from the water, sand/gravel filtration to remove large particles, charcoal absorption to remove odors, and a distillation to remove the salt (and the rest of the impurities that they didn’t get out before).  I conducted the final step of the experiment as my demo this week because my lab tables don’t have sinks (very annoying) and we don’t have enough condensers (they are useless without sinks, by the way). 

My colleague suggested that I distill Cherry Coke first to introduce the technique, something I had never tried.  This was a lot of fun because the kids always love testing something they might actually eat or drink.  I decided to set up the apparatus during class (some of the groups were still finishing up the filtration, so half the class had some down time).  Most of the students had never seen a distillation.  I handed the tubing to a couple of boys and asked them to connect it for me, while I attached clamps to ring stands.  I asked for more helpers to hold the round bottom flask and condenser in place while I clamped it all together.  The set-up was a distant cousin to the fractional distillation apparatus in my college organic chem lab, with glassware that all fits together and bars in the back of the hood for clamps.  After a lot of adjusting, I managed to assemble the apparatus with the Coke in the flask, when I realized that I had forgotten to add the boiling chips.  So I loosened some clamps, lowered the round bottom flask, added the boiling stones, and manipulated the apparatus back together.  My students watched, gathered around the front bench, a few lending a hand to help me steady everything as I readjusted.   At the same time I explained how the system works to separate liquids with different boiling points.  Finally we were ready to light the Bunsen burner.  (As a side note, the first time they see me light the burner is a big deal every year.  Soon enough the magic is gone and this becomes just part of the regular lab routine.)

The first thing we observed was rapid bubbling as we watched the carbon dioxide leaving the solution.  I bubbled the gas through limewater to confirm the presence of CO2, another suggestion from my colleague.  Very cool so far.  Then we started collecting a liquid.  I wish that I had remembered to put a thermometer in the flask to note the boiling point of the distillate, but I didn’t think about that until it was too late.  It’s been a really long time since I’ve done a distillation.  It all started coming back to me as we were boiling the Coke.  Maybe I should have spent a little longer reading Shakhashiri, because a thermometer is clearly labeled in the diagram!  We collected a clear, colorless liquid form of the cherry/vanilla flavoring.  The strong odor of cherry coke was overpowering in the small sample we collected.  The second liquid we collected was also clear and colorless with a faint smell of cherry coke; this time I think we collected water with a small amount flavoring contaminate.  After we had about 100 mL of distillate, I removed the pot and let the students observe the smell of the liquid left behind, flat and flavorless Coke.  They enthusiastically passed around the three samples and commented on their odor and color.

 Distilling the foul water sample, the last step in the purification lab.

Next we distilled their foul water samples.  I combined several groups’ samples into the pot because we were only going to do this once.  The brownish, electrically conductive water sample in the pot started to boil vigorously as the bell rang ending the period.  No!  My kingdom for 20 more minutes.  We were already collecting distillate, so I decided to let it continue to distill after the period ended.  I saved the purified water sample into an Erlenmeyer flask and saved the remaining pot liquid.  During the next class period, I was a bit nervous to test the conductivity of the solution.  What if it didn’t work?  How would we explain the presence of salt in our purified sample?  I decided to let the anticipation build by showing the class a control test of the conductivity of distilled water (no light) and salt water (bright light).  We were all pumped up when the light did not come on when we submerged the electrodes in our purified water. The only disappointment was the lingering garlic odor in our water sample.  The purification would have been better with a fractional distillation set-up rather than a simple distillation (a second condenser above the pot makes all the difference), possibly separating out the organic substance in the solution responsible for the lingering odor.  However, we did the job we set out to accomplish; the salt was definitely gone from the final product.

 Compare the sample of untreated foul water to our purified sample.  Looks nice, but don't smell it!
 Here's the distillate we collected and the liquid that was left in the pot at the end of the distillation.