Friday, April 8, 2011

A Chemist in the Kitchen

The following is entirely true.  No names or ingredients have been changed since the original, well, incident, occurred a few Thanksgivings ago.

I was living in Chicago in the 90’s, first in Hyde Park (home of the wonderful Seminary Coop Bookstore and 57th Street Books) and, later, in Printer’s Row, which is just south of the Chicago Loop.  As it was nearly Thanksgiving and I was living 1000 miles from my family (in Boston) at the time, I was pleased to receive a Thanksgiving Day invitation to (relatively) nearby Champaign-Urbana.  I had been to Champaign a few times before and, well, there it is.  At any rate, I gratefully accepted the invitation of the hostess and I volunteered to bake banana bread for the occasion. How hard can it be?  I’m a chemist, after all.  Since this was weeks before Thanksgiving, I felt I had ample time to get the job done.  Over the next few days, I researched banana bread recipes and converged on a consensus recipe and made a shopping list as I did so.  Here was my list:

Very ripe bananas
Baking powder

Some of these items I had on hand, but some I did not so, about a week before Thanksgiving, I went to the grocery store with my list and acquired the necessary ingredients.  I was most concerned with the quality of the bananas as that is, after all, the key ingredient.  That said, a week would be more than enough time to bring the bananas to the “ready” state.  I had everything well in hand.

OK, now fast-forward a little bit to, um, 11:45 PM on the eve of Thanksgiving, and you find me in my kitchen, looking over the recipe, which read “put flour, baking powder, and baking soda in a bowl.”  Baking soda?  That wasn’t on my list!  What am I supposed to do now?  I should note that, at this point in the story, most people, knowing that I am a chemist guessed [incorrectly] that I went to the lab to get some baking soda.  Baking soda is just a chemical and whether you buy it in the grocery store or from the Aldrich Chemical Company, it’s the same powdery white, odorless salt that dissolves instantly in water.  It sounds a lot like iocane powder which, you may recall, was a critical element (though not an element) in The Princess Bride.

So I looked around my apartment and asked myself, “OK, what is the purpose of baking soda in baking?”  Baking soda is only sodium bicarbonate, NaHCO3.  It’s weakly acidic and weakly basic at the same time.  When it reacts with acid, it makes carbonic acid, H2CO3, which rapidly decomposes to form carbon dioxide, CO2, and water, H2O.

You can see this in action if you dissolve a tablespoon of baking soda in about ½ cup of water and pour in a few tablespoons of vinegar (the active ingredient in vinegar is acetic acid and, in stores, you can by about a 3% solution – not very concentrated, so won’t hurt you, but it has a strong odor).  You’ll see a gentle bubbling. You might want to try this in a sink.  And it wouldn’t hurt to put on a pair of safety glasses – sunglasses would work too - so that you don’t splash any vinegar in your eyes.  But even if you do, flush with water and you’ll be fine.  Want more bubbles?  Add vinegar directly to a tablespoon of baking soda in a glass. Try a little at first, just to see what you’re getting yourself into…

So the point of this excursion is to remind/show you that the purpose of baking soda is to generate CO2, which makes the bread rise.  Baking powder actually contains a little bit of baking soda too – I didn’t know this at the time – as well as cream of tartar (derived from tartaric acid, which is found in wine as well as bananas), the purpose of which is to release a little acid and, in turn, CO2.  But let’s return to my banana bread story.  

So, by now it’s Midnight and I am driving to Champaign in 10 hours.  By this time I had been thinking about the purpose of the baking soda and noticed, in my medicine cabinet a bottle of, wait for it, Tums!  OK, so Tums doesn’t have baking soda in it.  Rather, the active ingredient in Tums is magnesium carbonate, MgCO3, but it still reacts with acid and makes CO2 (don’t believe me? Pulverize a pellet and add some vinegar to it…).  So, I took a single, mint-flavored, Tums and crushed it with the back of a spoon in a sturdy mug (I didn’t have, and still don’t, a mortar and pestle in my kitchen – it would have been really handy… my birthday is in October, by the way) and added it to the mix.  I continued with the recipe, baked the banana bread, and drove down to Champaign for a Thanksgiving dinner that couldn’t be beat.  The banana bread looked and tasted like the real thing and nobody was the wiser.  Some time later, I confessed to the hostess who graciously replied, “I guess the Tums helped to settle our stomachs after eating so much turkey.”  Thanks, Jeanne.

I have since made banana bread by the “standard” recipe but, a few years ago, I did a side-by-side comparison of Tums vs. baking soda.  In a quality-control experiment conducted by students in Haverford College Chem 101b, the banana bread baked with baking soda rose noticeably higher, but was otherwise (and taste-wise) indistinguishable from the Tums variety.  Maybe more Tums next time?  We’ll see.

Try this at home!

Friday, February 4, 2011

A Chemist in Winter

I was in a drug store recently, and the following conversation (almost) ensued:

Clerk:      Can I help you?
Me:         Yes, I’m looking for the hot packs.
Clerk:      Aisle 4, look in the back.
Me:         Yeah, I was just there.
Clerk:      You couldn’t find them?
Me:         I saw the metal ones, but I was hoping for the kind where you crack open the inner pouch and the whole bag just gets hot.  No rubbing required.
Clerk:      Does thermodynamics work differently for you than everybody else?

Let me explain…

Cold and Hot packs work, in many cases, on the same general principle.  In each case, a chemical reaction occurs – typically when you break the inner pouch – that either releases or absorbs heat.  Depending on the direction of heat flow, in or out of the bag, the pack will become cold or hot.

First, cold packs…
Cold packs have two separate compartments. One contains a granular solid, and the other is just a liquid – water, in fact.  The solid is an ionic compound called ammonium nitrate. (“Ionic” because it is composed of two different charged chemicals) When you break the inner bag, water mixes with the ammonium nitrate and, in this process, the salt dissolves.  When salts dissolve in water, the ammonium ions (NH4+) are completely separated from the nitrate ions (NO3-) and lose all memory of their history as they become surrounded with the water.  Think of it like a mosh pit at a rock concert. 

Positive charges (here the ammonium ions) and negative charges (the nitrate ions) are attracted to each other in the salt form, and the positive and negative ions are each attracted to the water molecules in the solution, but there is an energy “trade-off” that occurs.  It turns out that the attractive forces between the ions aren’t completely offset by the new attractive forces between the water and the ions. The salt dissolves but, in the process, absorbs a little heat from the surrounding water.  We call this process endothermic or “heat absorbing.”  The salt, in dissolving, absorbs heat from the surroundings (in this case, the water) and the bag becomes cold.  Ta-da!

Now hot packs…
Hot packs work by a very similar process with salt (usually calcium chloride, CaCl2) in one compartment and water in the other.  However, when you break the inner pouch, the salt dissolves in an exothermic or “heat releasing” reaction.  The excess heat is released to the surroundings (in this case, the water) and the bag becomes hot.  Same general reaction (dissolution of a salt), but the heat flows out instead of in.

Other kinds of hot packs?

But there is another kind of hot pack you can buy that consists of iron filings, charcoal, salt, and sawdust in a sealed pouch that is sold in a sealed envelope.  To activate the hot pouch, you open the envelope and expose the pouch to air.  This leads to an exothermic (heat releasing) chemical reaction between iron and oxygen (in the air) to produce iron oxide (Fe2O3), also known as “rust.”

Try this at home!

Friday, January 28, 2011

Practical Friday - Smoke Detectors!

Everybody has seen smoke detectors (if not, maybe you should talk to your landlord...) but not many people know, really, how they work.  Should we be afraid of them?  No, although I wouldn't suggest taking one apart.  Here's why.

Ionizing radiation – how do smoke detectors work?

Most smoke detectors rely on the natural radioactivity of the element, Americium that ejects "alpha" particles (small charged particles with 2 protons and 2 neutrons).  These particles are ejected with sufficient energy that, when they collide with other molecules (like gas molecules - don't forget that we're talking about smoke detectors), they knock electrons out of the other molecule. Radiation like this is called "ionizing" radiation, because the collision leads to the formation of ions.  Chemists and physicists love symbolic notation, so the equation below shows what is going on when Americium-241 decays.  On the left side of the equation is the symbol for Americium-241, the radioactive isotope with 95 protons and 146 neutrons (note that neutrons and protons add up to give 241, which is the mass number of this isotope).  Anyway, the result of the decay even is shown to the right of the arrow, and leads to the formation of Neptunium-237 and an alpha particle (2 protons, 2 neutrons).

O.K., so what?  Is this bad?  Well… yes.

The thing is that the alpha particles, once ejected have enough energy to damage cells (which is why you should never, ever, take apart a smoke detector).

Is there no GOOD news?

1. Alpha particles can’t pass through paper (also, by the way, don’t ever take apart a smoke detector)

2. Sometimes ionizing molecules is good.  Oh Yeah?  How?


Q:   How do they work?  What does this have to do with science?

A:   Inside the detector is a small sample of Americium, a radioactive element that decays by, wait for it… alpha emission.  The alpha particles emitted are constantly colliding with - and ionizing - gas molecules in air (nitrogen, oxygen, etc) and these form ions that are attracted to + and – charged plate in the detector (generating the background current in the device).  When smoke comes along, it soaks up these ions and changes the current, which turns the alarm ON!

Q:  Are they safe? 

A:  Yes. Remember that a particle can’t go through paper.

Q:  Really?

A:  Yes, but don’t take them apart because, after all, there is a little piece of Americium in there…

Tuesday, January 25, 2011

What is Practical Friday?

It was my wife's idea, really.  I was teaching at Haverford College (general chemistry, among other classes) one semester and, like many chemistry teachers, I used in-class demonstrations to drive home the "point" of the lecture.  This would continue but Practical Friday, we thought, would be something special, done weekly, where I take a step outside of class to point out the greater value of chemistry to science, if not to society.  What followed was a combination of whacky (What happens when you want to bake something but don't have baking soda?), resourceful (How do you measure the pH of a solution if you're stuck on an island with only...wait for it... cabbage?), and athletic (Practical Friday: Winter Olympics Edition!).

See you Friday!