Element of the Month - Copper

Text: Element of the Month: Copper Cu

"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 

In February we return to the more familiar elements and discuss a metallic element in depth with an examination of copper. This allows us to get into a discussion of the activity series of the metals and demonstrate the reactivity of various metals driven by their oxidation potentials. The chemistry of copper is of course colorful, which adds to the enjoyment of these experiments.


The students will usually be well aware that copper is used in coinage, cookware, roofing, plumbing, and electrical conductors. We note that copper is frequently alloyed with zinc to form brass, a popular decorative alloy, and with tin to give bronze, which is harder and less easily corroded than brass or copper, and is frequently used to make parts for ships such as propellers. We mention that copper compounds are used as fungicides and algaecides. On display on the desk are a jar of pennies, a scrap of copper pipe and a few fittings, a scrap of copper electrical cable, a scrap of copper flashing, a piece of copper cookware, a brass doorknob, and a scrap of pressure treated wood.

Physical Properties

Students will frequently be familiar with copper as a reddish, moderately dense, soft metal. We observe that copper is both malleable and ductile, and explain that these terms mean that it is easy to shape by hammering or by drawing through a die. We note that not all metals are, with iron being obviously more difficult to shape. Copper is an excellent conductor of both electricity and heat.

Figure 1: Left to right: copper II sulfate solution, copper II sulfate solution plus a small amount of NH4OH precipitating Cu(OH)2, copper II sulfate solution plus an excess of NH4OH resulting in the dissolution of Cu(OH)2 and formation of Cu(NH3)4 dication.

Chemical Properties 

We start by stating that copper is a relatively unreactive, or difficult to oxidize, metal. This is why it’s so useful. We don’t want roofs, coins, pipes, and wires that corrode easily. We demonstrate a qualitative analytical test for copper ion by treating some cupric sulfate solution with 2 M ammonium hydroxide solution in a large test tube. We show the initial formation of a blue precipitate of Cu(OH)2, followed by the dissolution of the precipitate upon the further addition of ammonium hydroxide solution to form the dark blue complex ion Cu(NH3)42+ (see figure 1).1 We also demonstrate the flame color of copper using a 0.2 M cupric chloride solution (see Video 1). This solution gives a much brighter flame color than does cupric sulfate.2 We mention that copper forms many complexes, not just with ammonia, which is why copper chemistry is so colorful. For example, we show and compare samples of anhydrous and hydrated cupric sulfate. Then we add several drops of water to the anhydrous cupric sulfate to show the hydration of the anhydrous cupric sulfate to generate the blue Cu(H2O)42+ ion. We particularly note the color change of white to blue, which are coincidentally the school colors for our high school. Next we make up a 0.2 M cupric chloride solution by dissolving 17 g of CuCl2 dihydrate in 500 mL water containing 10 mL concentrated hydrochloric acid (the HCl prevents the possible partial hydrolysis of CuCl2 to produce insoluble Cu(OH)2). We pour off 200 mL of this solution into a separate flask and show that copper ion also forms a complex with chloride ion by adding a large excess of anhydrous calcium chloride (25 g). The additional chloride ion favors the shift of equilibrium from blue Cu(H2O)42+ to green CuCl42+ ion.3 The colors of copper chemistry make it fun, while the low reactivity of the metal makes it useful.

Video 1: Colors of Elements in a Flame - Copper(II) Chloride. Credits: F. M. Hastings, J. J. Jacobsen. Derived from Jacobsen, Jerrold J.; Moore, John W., Chemistry Comes Alive!, Vol. 2; Journal of Chemical Education, 2000, 77(5), 671. DOI: 10.1021/ed077p671.*


Next we ask the class if copper metal will dissolve in hydrochloric acid. We demonstrate the reaction of samples of copper, iron, zinc, and aluminum wire with 6 M hydrochloric acid in large test tubes. We first allow the wire samples to react at room temperature, then briefly heat each tube in turn by placing the tube in a large beaker of very hot water (Note: it is important to make sure the copper and iron wire samples are free from corrosion or tarnish, and that the iron wire is free of any grease or oil). We explain that copper metal will only dissolve in oxidizing acids such as strong sulfuric acid or nitric acid and we remind them of the Ira Remsen experiment that we performed in November, outlined in Element of the Month: Nitrogen (see Video 2).

Video 2: A few drops of concentrated nitric acid are place on a pre-1982 copper penny. Credits: Jacobsen, Jerrold J.; Moore, John W., Chemistry Comes Alive!, Vol. 3; Journal of Chemical Eduacation, 1999 (76) 9 1311. DOI: 10.1021/ed076p1311.*

Based on the reactions of the wire samples with the hydrochloric acid, we ask the class: Which metal is most reactive? Which is least reactive? Which is in between? We write a list on the board in order of reactivity, lowest to highest based on the student choices. We ask, "are there any metals that are less reactive than copper? Generally the students will soon suggest silver, gold, and perhaps platinum. We add these metals to our list on the board. Next we ask the class, "what would happen if we put a metal that is more reactive than copper into a solution of a copper compound?" We demonstrate this by dissolving 25 g of CuCl2 dihydrate in 500 mL of water in a beaker with rapid magnetic stirring, and then we add 5 g of 20 mesh aluminum granules to the beaker.4 We continue to stir rapidly, and note the obvious heat of the reaction. When the reaction subsides, we filter the mixture through large fluted filter paper and note the colorless filtrate containing aluminum chloride and the filter paper containing the precipitated copper metal. We then ask the class, "what about the reverse reaction, putting copper into an aluminum solution?" Most students will immediately offer that no reaction is to be expected. We place a sample of copper foil into some dilute aluminum sulfate solution in a large test tube, and ask the class if we should pause the Element of the Month activity until a reaction is observed. A chorus of “No!” results and we move on.

Figure 3: Nails before (left) and after (right) exposure to 0.2 M CuCl2 solution for about 10 minutes. 

We ask, "what might happen if we use a less reactive metal such as iron with our copper chloride solution?" We place three or four clean steel nails into some 0.2 M copper chloride solution in a beaker. “Four penny” common nails work well, but they must be cleaned free of grease using a towel and some paint thinner or turpentine. After a minute or two, we pour off the copper chloride solution and display the nails on a paper towel (see figure 3). They are covered with a red brown deposit of copper metal. Next, we ask what might happen if we go the other way and put a less reactive metal such as silver into a copper solution. The class will again correctly guess that no reaction is to be expected. We confirm this by placing a piece of silver foil into some cupric sulfate solution (we do not use cupric chloride solution because the chloride ion can tarnish the silver surface). Okay, we ask, "what might happen if we put copper metal into a silver solution?" By now the class is on task and will correctly offer that a reaction will occur. We place a piece of copper foil into some 0.1 M (or 5%) silver nitrate solution in a large test tube, and show the rapid growth of silver particles on the surface of the copper foil.5 Lastly we ask the class how we might make copper metal deposit on silver, in effect making the reaction go “backwards”? This of course can only be accomplished by putting energy into the system, in the form of electrical energy. We show that we can electroplate copper onto silver by using a beaker containing 0.2 M cupric sulfate solution acidified with a little dilute sulfuric acid, using a copper foil anode and a silver foil cathode with about 3 volts DC current applied.6 A brown deposit of copper quickly appears on the silver foil. Lastly, we ask the class what would happen if we run the cell backwards. We reverse the polarity of the cell, and the copper deposit seems to almost magically disappear from the silver foil.7

**more videos from the ChemEd X Collection are available by subscription.

Notes and References:

  1. See, for example, Engelder, C. J.; Dunkelberger, T. H.; Schiller, W. J. Semi-Micro Qualitative Analysis; Wiley: New York, 1936; pp 122.

  2. The signature color emission is brought about by excitation of the metal chloride; see, for example: (a) Lancaster, R.; Shimizu, T.; Butler, R. E. A.; Hall, R. G. Fireworks: Principles and Practice Chemical Publishing Co.: New York, 1972; pp 60; (b) Conkling, J. R. Chemistry of Pyrotechnics Marcel Dekker: New York, 1985; pp 155 – 166.
  3. This experiment is modified from (a) Summerlin, Lee R.; Ealy, James L. Jr. Chemical Demonstrations: A Sourcebook for Teachers; American Chemical Society: Washington, DC, 1985; pp 54 – 55; and (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 71 – 72. We use anhydrous calcium chloride instead of concentrated hydrochloric acid in order to avoid a corrosive mixture while simultaneously producing an increase in temperature and a high concentration of chloride ion.
  4. This is a modification of 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 202. Fifteen grams of 30 mesh granular zinc metal may be used if granular aluminum is not available. Alternatively three grams of aluminum foil, cut into small pieces about 3 mm square, may be used.
  5. The demonstration of silver “foliage” on a copper foil “tree” has been popular for many years, see: Ford, Leonard A. Chemical Magic, 2nd ed.; Dover: Mineola, NY, 1993; pp 82.
  6. See Shakashiri, Bassam A. Chemical Demonstrations: A Handbook for Teachers, Vol. 4; University of Wisconsin Press: Madison, WI, 1992; pp 212 – 223.
  7. In these electroplating experiments, it is important to use copper(II) sulfate and not copper(II) chloride as the electrolyte solution. When the cell is run “backwards” at the end of the demonstration to remove the copper layer from the silver foil, some silver metal will be oxidized to silver ion at the anode and this will lead to a muddy suspension of AgCl if copper(II) chloride is used. We save the electrolyte solution and use it from year to year.


Safety: Video Demonstration

Demonstration videos presented here are not meant as tools to teach chemical demonstration techniques. They are meant as a tool for classroom use. The demonstrations may present safety hazards or show phenomena that are difficult for an entire class to observe in a live demonstration.

Those performing the demonstrations shown in this video have been trained and adhere to best safety practices.

Anyone thinking about performing a chemistry demonstration should first read and then adhere to the ACS Safety Guidelines for Chemical Demonstrations (2016) These guidelines are also available at ChemEd X.

General Safety

For Laboratory Work: Please refer to the ACS Guidelines for Chemical Laboratory Safety in Secondary Schools (2016).  

For Demonstrations: Please refer to the ACS Division of Chemical Education Safety Guidelines for Chemical Demonstrations.

Other Safety resources

RAMP: Recognize hazards; Assess the risks of hazards; Minimize the risks of hazards; Prepare for emergencies



Students who demonstrate understanding can construct and revise an explanation for the outcome of a simple chemical reaction based on the outermost electron states of atoms, trends in the periodic table, and knowledge of the patterns of chemical properties.

*More information about all DCI for HS-PS1 can be found at https://www.nextgenscience.org/dci-arrangement/hs-ps1-matter-and-its-interactions and further resources at https://www.nextgenscience.org.


Students who demonstrate understanding can construct and revise an explanation for the outcome of a simple chemical reaction based on the outermost electron states of atoms, trends in the periodic table, and knowledge of the patterns of chemical properties.

Assessment Boundary:

Assessment is limited to chemical reactions involving main group elements and combustion reactions.


Examples of chemical reactions could include the reaction of sodium and chlorine, of carbon and oxygen, or of carbon and hydrogen.