Element of the Month - Nitrogen

"In honor of the International Year of the Periodic Table this series of articles details the Element of the Month project developed by Stephen W. Wright (SWW), Associate Research Fellow at Pfizer Inc., and Marsha R. Folger (MRF), chemistry teacher (now retired) at Lyme – Old Lyme High School in Connecticut. Read The Element of the Month - An Introduction for an overview of the project and links to the other articles in the series." - Editor 

The third element highlighted in our Element of the Month program is nitrogen. While the relevance of nitrogen to the students’ lives is generally obvious to them, it is likely that the students have little familiarity with nitrogen despite the fact that they are usually aware that they walk in an atmosphere that is about three quarters nitrogen.

Occurrence in Nature

Students will respond that nitrogen is to be found as the uncombined element in the air, which is correct. In fact nitrogen is the most abundant free element on Earth. We ask how much of the atmosphere is nitrogen to remind them that nitrogen forms about 78% of the atmosphere, a point to which we will refer later. We then ask where else nitrogen is to be found in nature, and generally a little more prodding is required to provoke the answer that nitrogen is found in proteins.

Figure 1: Nitrogen containing products

Uses

We open this part of the discussion with the conclusion that nitrogen’s primary use then is that it is essential for life as a constituent of amino acids and therefore proteins. In fact, the “amino” term in the name “amino acid” is a nod to the nitrogen content. It follows, therefore, that some of nitrogen’s other uses are directed toward getting the element incorporated into proteins since humans lack the ability to synthesize amino acids. Nitrogen is the primary component of fertilizers. What else do we use nitrogen compounds for? We solicit ammonia, explosives and pyrotechnics, and certain plastics, such as nylon and Kevlar, as answers. We have on display items that contain nitrogen compounds (see figure 1). These can be items such as a pair of stockings, a container of fertilizer, a jar of protein supplement powder, a can of tuna, a bottle of household ammonia, a bottle of an ammonia-containing cleaner such as a glass cleaner, and if permitted, a box of highway flares. For safety and convenience, these containers may be empty and stored in a small box from year to year.

Figure 2: Ampule of nitrogen gas

Physical Properties

Clearly nitrogen is a gas at room temperature (figure 2). When asked, students will usually recall that liquid nitrogen is extremely cold. In fact its boiling point is -196 °C. We ask the students what such extreme cold would be useful for. Some possibilities include flash freezing of food or preserving tissue specimens. At this point, we announce that we have some liquid nitrogen and we proceed to do some cryogenics demonstrations to show what happens to items that contain water when frozen to such low temperatures. Wearing cryogenics gloves or ski gloves in addition to customary lab safety wear, we start with a few grapes, which are held in the liquid nitrogen until bubbling subsides and then placed on a block of wood and smashed with a hammer.¹ We repeat the process with a marshmallow, which generally produces much amusement. A tennis ball comes next, which can either be broken with a hammer or thrown on the floor. Students will be more surprised to see that a balloon touched to the surface of liquid nitrogen becomes deflated.² They will quickly understand why as we engage them to explain what they are observing as the experiment proceeds. Further, this process can be repeated. Lastly we freeze a carnation or two and crush the frozen blooms in the palm of a gloved hand. It is possible to hammer a nail with a frozen banana but this demonstration takes some time for the banana to freeze and requires some practice to establish the technique of hammering and the choice of appropriate nails and wood.³

Figure 3: Production of ammonia from nitrogen and hydrogen gases

Chemical Properties

Nitrogen is inert and does not support combustion. We show that a burning splinter of wood and a birthday candle are extinguished when lowered into a jar filled with nitrogen. We emphasize the strength of N-N triple bond and explain that’s why most nitrogen is found uncombined in the air, because any nitrogen that does end up as the element is essentially in an energy “hole” from which it is difficult to escape. The inert nature and low cost of nitrogen are useful when you want an oxygen free atmosphere around something, for example when welding or packaging salad mixes. On the board, we show the equation for converting nitrogen to ammonia, and emphasize that energy is needed to make the reaction go forward (figure 3). We note that the reaction is carried out by chemical factories to make ammonia for fertilizers, and also by certain bacteria, and that’s how nitrogen gets from the atmosphere into the food chain.

Video 1: Ammonia Fountain* 

Ammonia

For these experiments we have been fortunate to have access to a small cylinder of anhydrous ammonia gas. We emphasize that ammonia is a gas, with a boiling point of -33 °C, not a liquid as found in the supermarket. It has a pungent and familiar odor. We show how amazingly soluble ammonia gas is by passing ammonia gas from the cylinder (noting that it is a colorless gas) into an empty, stoppered 500 mL Erlenmeyer flask. The outlet is passed into a 1 L beaker of water containing a little phenolphthalein. When ammonia begins to pass into the beaker, the water will turn pink showing that the solution has become alkaline. Then we shut off the ammonia supply and watch as the pink water is pushed back from the beaker into the Erlenmeyer flask, practically filling the flask. The ammonia fountain if a favorite demonstration of many chemistry instructors (see video 1 above).⁴ Why does that happen? We emphasize the very high solubility of ammonia in water and note that 1 mL of water will dissolve nearly 1300 mL of the gas at room temperature. Most students will have never observed the phenomenon of a highly water soluble gas before. We note that the water solution of ammonia in the Erlenmeyer flask is the usual form of ammonia that consumers encounter.

Figure 4: The “Poor Man’s ammonium fountain set up

Many teachers would like to perform the ammonium fountain demonstration, but don't have the glassware shown in video 1. See what some call the "poor man's ammonium fountain" in figure 4. 

 

Figure 5: The reaction of ammonia and hydrochloric acid results in a cloud of ammonium chloride.

Being a base, ammonia combines with acids. We write the equation for the reaction of ammonia with hydrochloric acid on the board, and then place a few milliliters of diluted ammonium hydroxide on some paper toweling in an evaporating dish, and a few milliliters of diluted hydrochloric acid on some paper toweling in a second evaporating dish (alternatively, the two pieces of paper toweling may be placed side by side in the same evaporating dish).⁵ A cloud of white ammonium chloride “smoke” will appear, and we explain that this is formed from the reaction between the two gases escaping from the two solutions (see figure 5).

Figure 6: Some of the nitrogen oxides found in nature

Nitrogen Oxides

We explain that, under appropriate conditions, nitrogen will form compounds with oxygen. At the board, we show the formulas for the three nitrogen oxides (figure 6). It’s NO₂ that often makes urban atmospheres hazy and yellow, and note that we’ll discuss NO₂ further later on. Students will usually have some familiarity with nitrous oxide, N₂O. We note that nitrous oxide is colorless but not odorless or tasteless – it actually has a rather sweet taste. We also note that is the only other gas besides oxygen that gives a positive glowing splint test, which was demonstrated last month. In fact nitrous oxide supports combustion fully as well as oxygen. Why might that be? We emphasize the stability of the N₂ molecule and can see that N₂O is almost N₂! If a source of nitrous oxide is available, the glowing splint test may be performed in a jar of nitrous oxide, demonstrating that N₂O supports combustion.⁶ 

We note that nitrogen - oxygen bonds are found in nitrates too. What products contain nitrates? We solicit gunpowder, fireworks and explosives for answers. We explain that nitrogen wants to form strong N-N triple bonds, not weak N-O bonds, thus nitrates easily lose their oxygen atoms to reducing agents, and this is how pyrotechnics and explosives work. We demonstrate the oxidizing power of nitrates by dropping small charcoal pieces, one at a time, into a small porcelain evaporating dish containing potassium nitrate that has been heated to melting over a gas burner.⁷ We shut off and remove the burner immediately prior to adding the charcoal pieces.

The combustion of nitrocellulose, used in modern propellants, may be demonstrated by igniting pieces of a broken ping pong ball.⁸ A video showing the combustion of ping pong balls may be found in Tom Kuntzleman's ChemEd X blog post, A Simple, yet Dramatic Chemistry Experiment with Ping Pong Balls.⁹

Nitric Acid

We always make sure that we have time remaining to discuss nitric acid and reenact Ira Remsen’s investigation of it. We explain that concentrated nitric acid is a very corrosive acid and powerful oxidizing agent. We show, by equations previously written on the board, that nitric acid converts toluene into TNT and cellulose into nitrocellulose, which are explosives that we just talked about.

Figure 7: Ira Remsen Equation

At this point we perform Ira Remsen’s copper / nitric acid experiment (see the equation in figure 7), reading aloud his description of the experiment but explaining that, for safety, we will perform the experiment in a 1 L flask in a fume hood and not on the teacher’s desk!¹⁰ We pour 5 mL of concentrated nitric acid into the flask and add a penny. If desired, the experiment may be projected by video camera to improve visibility. The immediate fuming of the liquid, the rich blue color of copper(II) nitrate, and the clouds of red –brown NO₂ gas never fail to impress the class. We dilute the reaction with a large volume of water in the hood to stop the reaction, then pour the mixture through a funnel to recover the penny, which may be shown to the class after washing and drying. In a bit of humorous theater, we produce an old pair of cotton pants bearing numerous acid holes and claim them to belong to Remsen himself. Lastly, we hand out copies of the paper below as the students leave the classroom. Readers can download a pdf of the script below. You can view a smaller scale version of this demo in video 2 below.

Video 2: A few drops of concentrated nitric acid are place on a pre-1982 copper penny.*

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The Story of Ira Remsen¹¹

“While reading a textbook on chemistry, I came upon the statement, ‘nitric acid acts upon copper.’ I was getting tired of reading such absurd stuff and I 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 peculiarities, but I was getting on and likely to learn. The spirit of adventure was upon me. Having nitric acid and copper, I had only to learn what the words ‘act upon’ meant. Then the statement, ‘nitric acid acts upon copper,’ would be something 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 possession. I put one of them on the table; opened the bottle marked ‘nitric acid’; poured some 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 greenish 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, which I had meanwhile opened. 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 and, relatively, probably the most costly experiment I have ever performed. 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.”

 

References and Notes

1. (a) Conant, James Bryant; Black, Newton Henry New Practical Chemistry; Macmillan: New York, 1940; pp. 267-268; (b) Summerlin, Lee R.; Borgford, Christie L.; Ealy, Julie B. Chemical Demonstrations: A Sourcebook for Teachers Volume 2, 2nd ed.; American Chemical Society: Washington, DC, 1988; pp 20-21.

2. A video showing the behavior of a balloon in liquid nitrogen may be found in Tom Kuntzleman's ChemEd X blog post, Chemical Mystery #4: The Case of the Misbehaving Balloon. 

3. The banana must be held in the hand protected by heavy, cold weather gloves. We use roofing nails which have large heads and pound them into a very soft piece of pine. Even so, the banana will usually shatter quickly.

4. Several versions (including the one shown) of the ammonia fountain can be found on the ChemEd X Video page. See also (a) Ford, Leonard A. Chemical Magic, 2nd ed.; Dover: Mineola, NY, 1993; pp 33; (b) Conant, James Bryant; Black, Newton Henry New Practical Chemistry; Macmillan: New York, 1940; pp. 282.

5. (a) Ford, Leonard A. Chemical Magic, 2nd ed.; Dover: Mineola, NY, 1993; pp 91; (b) Conant, James Bryant; Black, Newton Henry New Practical Chemistry; Macmillan: New York, 1940; pp. 285.

6. Conant, James Bryant; Black, Newton Henry New Practical Chemistry; Macmillan: New York, 1940; pp. 303.

7. Stone, C. H., J. Chem. Educ., 1943, 20 (4), 200.

8. Summerlin, Lee R.; Borgford, Christie L.; Ealy, Julie B. Chemical Demonstrations: A Sourcebook for Teachers Volume 2, 2nd ed.; American Chemical Society: Washington, DC, 1988; pp 105.

9. Tom Kuntzleman's ChemEd X blog post, A Simple, yet Dramatic Chemistry Experiment with Ping Pong Balls.

10. See Michael Morgan's ChemEd X blog post, My First Day Demonstration. 

11. The script was published in F. H. Getman, "The Life of Ira Remsen," Journal of Chemical Education 1940, pp 9-10.

Note that nitric acid is usually made and sold as a 70% solution in water (70 grams of HNO₃ in 100 grams of solution), containing about 16 moles or 990 grams of nitric acid per liter.