Copper and Nitric Acid Chemistry Demonstration


Copper and Nitric Acid Chemistry Demonstration
The copper in nitric acid chemistry demonstration is a dramatic color change chemical reaction.

The copper and nitric acid reaction is a dramatic color change chemistry demonstration. The reaction illustrates several chemistry principles, including exothermic reactions, redox reactions, coordination complexes, oxidation, oxidation states, and the metal activity series. Here are instructions explaining how you perform this demonstration safely, with a look at its chemical reactions.

Materials

You only need two common chemicals. The most important part of the reaction is the choice of reaction vessel. The reaction produces heat, so use a study glass container.

  • 5 g copper
  • 40 ml concentrated nitric acid (HNO3)
  • Water
  • 1-liter flask (Erlenmeyer, boiling flask, or Buchner flask)
  • Clamp stand
  • Bowl (optional)

The original demonstration uses a copper penny, but modern pennies are zinc plated with a thin layer of copper. A better choice is a piece of copper wool or some copper shavings. The reaction works fine with copper wire, but is not as dramatic because the wire has less surface area.

A smaller version of the demonstration uses a bit of copper, a small volume of nitric acid, and a borosilicate glass test tube.

Perform the Copper and Nitric Acid Chemistry Demonstration

Nothing could be easier! Set up and perform the demonstration inside a fume hood.

Copper and Nitric Acid Reaction
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  • Pour the nitric acid into the flask.
  • When you are ready for the reaction, add the copper.

Initially, the nitric acid attacks the copper, turning the liquid green and releasing heat and reddish brown nitrogen dioxide vapor. Eventually, even the liquid turns brown.

  • Add water and dilute the solution.

Diluting the acid changes the conditions. The liquid changes color into a bright blue, while the vapor changes from reddish brown to colorless.

A Look at the Chemistry

If you look at the metal reactivity series, copper is pretty unreactive. It’s even considered a noble metal by some chemists. It resists oxidation by hydrochloric acid (HCl), yet readily reacts with nitric acid (HNO3). This is because nitric acid acts both as an oxidizer and an acid. Copper reacts with nitric acid, forming aqueous copper nitrate, nitrogen dioxide gas, and water.

Cu(s) + 4HNO3(aq) → Cu(NO3)2(aq) + 2NO2(g) + 2H2O(l)

The reaction immediately produces heat (reaching 60 to 70 degrees C) and releases deeply-colored nitrogen dioxide gas. The green color comes from copper(II) ions forming a coordination complex with nitrate ions. Diluting the concentrated acid with water changes the liquid color to blue as the water displaces the nitrate ions, leaving only aqueous copper(II) nitrate. The water reacts with nitrogen dioxide and forms nitric oxide.

3Cu(s) + 8HNO3(aq) → 3Cu2+(aq) + 2NO(g) + 4H2O(l)+ 6NO3(aq)

The concentration of the acid affects its oxidizing capacity. For example, copper does not react with dilute sulfuric acid (H2SO4), but a similar reaction occurs in concentrated sulfuric acid:

Cu + 2H2SO4 → SO2 + 2H2O + SO42− + Cu2+

Containing the Copper and Nitric Acid Reaction

A few simple revisions contain the reaction and improve both the safety and dramatic effect of the copper and nitric acid chemistry demonstration. You can perform this variation of the copper and nitric acid reaction out in the open, but it’s still a good idea to separate the set-up from the audience using a safety shield.

  1. Add nitric acid to a round-bottomed borosilicate flask. Clamp it into position on a stand. Ideally, use a borosilicate flask and place a bowl beneath the flask in case the glass leaks or breaks.
  2. Fill an Erlenmeyer (conical) flask with water and clamp it into position near the round flask.
  3. Stopper the round flask (acid) and loosely plug the conical flask with glass wool. The glass wool prevents the escape of nitrogen dioxide into the outside air. Insert glass tubing the ends reach the bottoms of each flask. (Don’t use plastic tubing.)
  4. When you are ready for the demonstration, add the copper to the borosilicate flask and fit the stopper and tube onto it.

Initially, the liquid in the round flask turns green and evolves reddish brown nitrogen dioxide. After about a minute and a half, the reaction slows and cools. The pressure reduction from the cooling draws water in from the conical flask. This dilutes the nitric acid and also reacts with the nitrogen dioxide gas, forming a fountain. Finally, the liquid in the round flask turns blue as copper nitrate forms.

Safety and Disposal

  • Only perform this demonstration if you are a chemist or chemistry educator and have access to proper safety gear and a fume hood. Nitric acid is a corrosive strong acid, while nitrogen dioxide is a toxic reddish-brown gas. Wear gloves, goggles, and a lab coat. Perform the open demonstration under a fume hood.
  • Please choose sturdy glassware for this demonstration. The initial reaction produces heat, so there is a risk of glassware breakage. For this reason, a boiling flask is ideal. Alternatively, use a Buchner flask.
  • After the demonstration, neutralize the dilute nitric acid using any inorganic base, such a baking soda, sodium hydroxide solution, or potassium hydroxide solution. The neutralization reaction also produces some heat. Afterward, you can safely wash the liquids down the drain with water.

References

  • Cotton, F. Albert; Wilkinson, Geoffrey (1988). Advanced Inorganic Chemistry (5th ed.). New York: John Wiley & Sons. 769-881.
  • Shakhashiri, Bassam Z. (1985). “Properties of Nitrogen(II) Oxide”. Chemical Demonstrations: A Handbook for Teachers of Chemistry Volume 2. The University of Wisconsin Press. ISBN: 978-0299101305.
  • Shakhashiri, Bassam Z. (1985). “Coin-Operated Red, White, and Blue Demonstration: Fountain Effect With Nitric Acid and Copper”. Chemical Demonstrations: A Handbook for Teachers of Chemistry Volume 3. The University of Wisconsin Press. 83-91. ISBN: 978-0299119508.
  • 163-166.Summerlin, Lee R.; Borgford, Christie L., Ealy, Julie B. (1988) Chemical Demonstrations: A Sourcebook for Teachers Volume 2 (2nd ed.). American Chemical Society. ISBN: 978-0841215351.