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Explanation of the Copper in Nitric Acid demonstration

What is supposed to be a beautiful demonstration of the reaction of nitric acid, resulting in three flasks of approximately equal volume but one red, one "white" (clear, actually) and one blue, results in an explosion, but, fortunately, without injury.  This demonstrates MANY points, but none so clearly as the need to practice proper safety precautions whenever working with chemicals.

OK, this is a tough one to explain, but here goes.  We have three flasks; the first has a little bit of concentrated nitric acid, and is connected by tube to a second flask, filled almost entirely with a weak solution of nitric acid and a little phenolphthalein indicator (phenolphthalein is colorless in an acidic solution but pink in a basic solution, so this solution is colorless).  The second flask is connected by tubing to a third flask, which is open to the atmosphere and is about half full of sodium hydroxide solution.  The tubing in the second and third flask only extend about halfway down the flask.

Now, two pennies (made of copper) are dropped into the first flask (with the concentrated nitric acid) and the stopper to the flask is replaced.  Immediately, a thick reddish brown gas fills the free space in the flask (nitrogen (IV) oxide) and the nitric acid becomes a green solution (copper (II) nitrate is green).  This reddish brown gas flows through the tube into the second flask (filled entirely with nitric acid); the gas forces the liquid in the second flask to flow into the third flask.

The sodium hydroxide solution in the third flask is strong enough to completely neutralize the acidic solution entering the flask, so the liquid becomes green as the flask fills.  As soon as the liquid level in the second flask is below the glass tubing (about halfway down the flask), the liquid stops flowing; now the second flask is half filled, and the third flask is almost entirely filled with a pink liquid.  Once the liquid stops flowing, the reddish brown nitrogen (IV) oxide gas fills the space above the liquid in the second flask.

As the copper is depleted in the first flask, the nitrogen (IV) oxide is no longer being formed.  In the second flask, the nitrogen (IV) oxide begins to dissolve in the water (forming nitric acid), creating a vacuum.  This draws liquid from the third flask back into the second flask.  The nitric acid in the second flask is now strong enough to neutralize the sodium hydroxide from the second flask, so the phenolphthalein again becomes clear.  The liquid in the second flask is now clear, and the flask is almost entirely filled with liquid.

What was SUPPOSED to happen at this point is the nitrogen (IV) oxide in the first flask should have begun dissolving in the liquid (as well as cooling), which should have created a vacuum, drawing the liquid from the second flask into the first flask.  The water, when it mixes with the copper (II) nitrate, then, should have formed a blue copper (II) hydrate complex.  The net result should have been, then, one half-filled red flask, one half-filled colorless (white) flask, and one half-filled blue flask.

Unfortunately, a plug in the tubing did not allow the gas from the first flask to continue to flow into the second flask (where it would have dissolved).  Because of this, there was pressure buildup in the first flask as gas was evolved, but could not escape (because of the plugged hose).  Thus, not only did the water NOT travel back into the first flask like it was supposed to have, but the pressure buildup in the first flask blew off (with great force) the rubber stopper in the first flask.  

This demonstration is based on that by Bassam Z. Shakhashiri, Chemical Demonstrations: A Handbook for Teachers of Chemistry, Volume 3 (The University of Wisconsin Press, Madison, Wisconsin, 1989), pp. 83-91.  The explanation is much more in-depth in this textbook, and the demonstration is far more impressive in person; I would highly recommend the purchase of this extraordinary set of books for anybody interested in this or other chemical demonstrations.  As with any demonstration, be sure to follow the directions very carefully and observe all applicable safety precautions.  This web site does not purport to be a site designed to convey directions for chemical demonstrations, and is not liable for any injuries or damages sustained by those who would attempt to re-create this demonstration without proper training, supervision or instructions.  Special thanks to Lori Weismantel for her help in creating this demonstration and to Jeremy Javers for his subsequent work to ensure that the NEXT time we try this demonstration, the ending will be much safer!