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Metallic Copper + Concentrated Nitric Acid
I adapted this lecture experiment from the "Coin-Operated Reaction" developed by Ron Perkins. I have made video records of a few lecture demonstrations, not to replace doing them in class, but rather so students can go back and review the experiments later in the year.

I wanted an interesting system with which to make observations on the first day of class. I also wanted to give students an opportunity to explain what they saw in a logical manner. This system is rich with ideas such as gas laws, air pressure, complex ions, oxidizing acids, oxidation reduction, and gas solubility in water.

I have the gas bubble into an aquarium since I have no hood in my classroom. The first time I tried this set-up, I was surprised when the gas cooled and the water was pushed back into the reaction flask. Here is a three-minute streaming video version of the lecture experiment (Real Audio plug-in required). Note: there is sound, but we are not saying anything. We took this movie with a digital camera and then our Stage Crew chief (James Jontz) made the 700 mb video into a 7 mb streaming video using his Macintosh (technically unexplainable by me).

The gas is bubbled into the aquarium so it won't get into the air. I also place a wet cloth over the opening of the aquarium.

What You Do...

To a 500 mL florence flask, add about 50 mL of concentrated nitric acid. Place a coil of copper wire into the acid and stopper with a one-holed rubber stopper fitted with a long tube. The end of the tube is placed into a large container filled with water. (I use a plastic aquarium.)[1]

As soon as the Cuo contacts the HNO3(conc) the red-brown NO2(g) forms.

What You See...

Many changes occur during this demonstration. When the copper wire (or use two pre-1982 pennies) is added to the colorless nitric acid, the solution turns green and a large amount of red-brown gas is formed. The air being displaced by the gas formation can be seen bubbling through the water. The flask gets VERY warm. When enough gas is formed, it bubbles through the water (keep the liquid stirred so most of it will dissolve. The gas that makes it to the top is noxious.

When the water siphons back into the flask, the blue Cu(H2O)42+ forms.

Later...

The gas in the flask begins to cool and therefore contracts. (I am reading my Ira Remsen story and allow students to notice the change.) As the pressure inside the flask decreases, the outside air pressure begins to push the water back toward the original flask. In addition, the red-brown gas dissolves in the water. Eventually, the water rushes into the flask, the solution turns characteristic blue, and the red-brown gas disappears as it is dissolved.

The Set Up...

Equipment:

The glass tube is bent in such a way as to connect the top of the flask with the bottom of the water container. The water in the container can be stirred by hand or with a magnetic stirrer. It needs to be stirred, however, or else the NO2 gas collects above the liquid (as it did when I took this picture...whew!).

The Chemistry...

Oxidation of copper metal with a strong oxidizing agent, conc. nitric acid.
In a classic experiment, copper metal is turned into copper(II) ion while the nitrogen(V) in the nitrate ion becomes nitrogen(IV) in the nitrogen dioxide gas.

Charles' Law
As the temperature from the reaction warms the gas, it expands. Later, as it cools, the gas contracts.

Nonmetal oxides are acid anhydrides (also link to acid rain)
Although the nitrogen dioxide gas is noxious and toxic, it dissolves readily in water and make the solution acidic. This can be shown by adding a little indicator to the water and making the water slightly basic before the copper is added to the acid.

Air pressure
As the pressure in the flask is decreased as it cools, the outside pressure pushes the water up the tubing toward the flask. The nitrogen dioxide gas is not pulling the water in.

Descriptive chemistry--copper solutions are green and blue
The colored solutions come from complexes of copper(II) ion in solution. Aqueous copper ion is blue, Cu(H2O)42+ The green must be copper surrounded by nitrates(?)

Discussion...

During this demonstration I read the reminiscence by Ira Remson quoted in Bassam Shakhashiri's demonstration book.[2]

While reading a textbook of chemistry I came upon the statement, "nitric acid acts upon copper." I was getting tired of reading such absurd stuff and I was determined to see what this meant. Copper was more or less familiar to me, for copper cents were then in use. I had seen a bottle marked nitric acid on a table in the doctor's office where I was then "doing time." I did not know its pecularities, but the spirit of adventure was upon me. Having nitric acid and copper, I had only to learn what the words "act upon" meant. The statement "nitric acid acts upon copper" would be more than mere words.

All was still. In the interest of knowledge I was even willing to sacrifice one of the few copper cents then in my possission. I put one of them on the table, opened the bottle marked nitric acid, poured some of the liquid on the copper and prepared to make an observation. But what was this wonderful thing which I beheld? The cent was already changed and it was no small change either. A green-blue liquid foamed and fumed over the cent and over the table. The air in the neighborhood of the performance became colored dark red. A great colored cloud arose. This was disagreeable and suffocating. How should I stop this?

I tried to get rid of the objectionable mess by picking it up and throwing it out of the window. I learned another fact. Nitric acid not only acts upon copper, but it acts upon fingers. The pain led to another unpremeditated experiment. I drew my fingers across my trousers and another fact was discovered. Nitric acid acts upon trousers. Taking everything into consideration, that was the most impressive experiment and relatively probably the most costly experiment I have ever performed... It was a revelation to me. It resulted in a desire on my part to learn more about that remarkable kind of action. Plainly, the only way to learn about it was to see its results, to experiment, to work in a laboratory.[3]

Safety and Disposal...

The solution is highly acidic. I pour it out into a large beaker or battery jar and add excess sodium carbonate. The carbon dioxide bubbles indicate neutralization and the resulting copper carbonate precipitate is filtered, placed in a baggie and thrown away. The neutralized filtrate can be disposed of as you would any simple salt solution. Procedures may vary from location to location.

References:
[1] This demonstration is based on one shown by Ron Perkins called the "Coin Operated Demonstration"
[2] Shakhishiri, B.Z. "Chemical Demonstrations Volume 1--A Handbook for Teachers of Chemistry"; The University of Wisconsin Press: Madison, Wisconsin, 1983
[3] Getman, F.H. "The Life of Ira Remsem"; Journal of Chemical Education: Easton, Pennsylvania, 1940; pp9-10.
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