DEMONSTRATION pubs.acs.org/jchemeduc
Better Ira Remsen Demonstration David K. Dalby,*,† James H. Maynard,‡ and John W. Moore‡ † ‡
Riverside High School, 794 Hammett Bridge Road, Greer, South Carolina 29687, United States Department of Chemistry, University of WisconsinMadison, Madison, Wisconsin 53706, United States
bS Supporting Information ABSTRACT: A method for demonstrating and expanding the Ira Remsen copper penny and nitric acid experience has been developed. It involves additional chemistry, showing five transformations of copper from one form to another. KEYWORDS: First-Year Undergraduate/General, High School/Introductory Chemistry, Second-Year Undergraduate, Demonstrations, Physical Chemistry, Acids/Bases, Aqueous Solution Chemistry, Coordination Compounds, Metals, UVVis Spectroscopy FEATURE: Tested Demonstration
M
any versions of the classic Ira Remsen experience involving copper and concentrated nitric acid have been used as lecture demonstrations.119 Remsen’s original reminiscence from 150 years ago is included in the Supporting Information,29 and his biography can be found on the Internet.20 A new version is presented that makes the demonstration more visible and dramatic and also extends the chemistry involved. In most educational settings, this demonstration must be done in a hood to prevent the NO2 from posing a danger to the audience, but a hood does not allow for a large group to view the demonstration adequately. This version enables a large group to see the reactions in detail. It also expands the experience to include additional reactions that form colorful copper complex ions. This provides an excellent beginning chemistry activity that will excite students and capture their attention. It also can be used as an anchor demonstration and revisited many times as the students gain the cognitive tools needed to understand different aspects such as reaction of acids with metals, complex-ion formation, redox reactions, acid anhydrides, solubility, precipitation, hydration, equilibrium, and disproportionation. The demonstration shows visually what is meant when chemistry is defined as the study of matter and how one form of matter interacts with another.
Figure 1. Demonstration assembly.
aspirator with the output tube submerged in water in a large flask so that the red-brown cloud of NO2(g) cannot escape into the room. The copper and nitric acid are allowed to react for about two minutes, and then the aspirator is turned on, the stopcock in the funnel is opened, and water is added, resulting in a blue solution characteristic of aqueous copper(II) ions. Aqueous ammonia is added slowly to this solution. As the pH increases, the color of the solution changes from blue to green due to formation of a copper(II) complex of the nitrite ion. Then copper(II) hydroxide precipitates, and finally the solution turns
’ SUMMARY OF THE DEMONSTRATION The demonstration is carried out using the equipment shown in Figure 1. While a student reads Ira Remsen’s description of his encounter with a copper cent and nitric acid, four pre-1982 pennies are placed into a thick-walled filtration flask and concentrated nitric acid is added. The flask neck is closed with a funnelstopcock device, and the side arm is connected to an Copyright r 2011 American Chemical Society and Division of Chemical Education, Inc.
Published: October 05, 2011 1687
dx.doi.org/10.1021/ed101172g | J. Chem. Educ. 2011, 88, 1687–1691
Journal of Chemical Education the deep blue characteristic of tetraamminecopper(II) ions. In the course of the demonstration, five chemical transformations of the copper occur and can be distinguished by color changes that are shown in the detailed description of the demonstration later in this article.
’ APPARATUS The device used to add water to a 4 L vacuum flask is shown in Figure 2A. It can easily be fabricated from the schedule 40 plastic pipe fittings shown in Figure 2B, a short section of 1/2 in. plastic schedule 40 pipe that extends to a point near the bottom of the vacuum flask, and a #12 rubber stopper. A hole in the rubber stopper for the 1/2 in. plastic pipe can be drilled using a 7/8 in. spade bit and a drill press. Should a 4 L vacuum flask not be appropriate or available, any other size may be used by altering the size of the rubber stopper and the length of the 1/2 in. schedule 40 pipe. The name, cost, Lowe’s part number, and UPC code at the time of this writing for the plastic pipe fittings are listed in Table 1. The total cost to build this item is about $12 and requires only a drill press and bit, a saw to cut the pipe, PVC cement, and about 10 to 15 min of assembly time. To assemble, remove any burrs produced when cutting pipe with a saw by using sandpaper or utility knife. Then simply apply PVC cement and quickly screw or slide the parts together.21,22 PVC cement will set in as little as 20 s. Priming the connections for this apparatus is not necessary.
DEMONSTRATION
During the mid-1800s, copper pennies were substantially different from pennies today. Pennies were composed of 95% copper and 5% zinc and tin.23 Although the copper content was essentially the same as for pre-1982 pennies, the dimensions were different (Table 2 and Figure 3). Pennies minted in 1982 may be composed of 95% copper and 5% zinc or may be made of 97.5% zinc and 2.5% copper.23 Pennies from 1983 and later are 97.5% zinc and 2.5% copper.23 Pre-1982 pennies are easily obtained from coins currently in circulation in the United States. Should the time come that pre-1982 pennies are no longer in circulation, coins can be obtained from local coin shops, online shops, or Ebay. If U.S. pre-1982 pennies are not available, other forms of copper such as copper shot or copper wire can be used. For the demonstration as presented here, about 3 g of copper will be consumed.
’ DETAILS OF THE DEMONSTRATION AND CHEMICAL EQUATIONS While someone reads Ira Remsen’s reminiscence, add four pennies and 25 mL of concentrated nitric acid to the 4 L flask. Quickly place the water-addition device, with the valve closed, on the flask. Within a few seconds, NO2 will fill the flask with the characteristic red-brown color (Figure 4A). The colorless liquid becomes green due to formation of complexes of copper(II) with nitrate ions (Figure 4B), but the green color will not be readily visible to the audience. The exact structure of these nitrate complexes is not known. Nitrate ions can be monodentate or bidentate ligands,24 and there may also be water molecules coordinated to the copper(II) ions. Assuming two nitrate ions and four water molecules coordinated to the copper(II), the reaction can be written CuðsÞ þ 4Hþ ðaqÞ þ 4NO3 ðaqÞ þ 2H2 OðlÞ f ½CuðH2 OÞ4 ðNO3 Þ2 ðaqÞ þ 2NO2 ðgÞ
ð1Þ
green
When the reading of Remsen’s words is over, place water in the funnel, turn on the vacuum aspirator, immediately open the valve to prevent a vacuum from being applied to the flask, and add water through the funnel until the flask is about 80% full. NO2 reacts with water, disproportionating to form nitrous acid and nitric acid: 2NO2 ðgÞ þ H2 OðlÞ f HNO2 ðaqÞ þ Hþ ðaqÞ þ NO3 ðaqÞ
ð2Þ
Figure 2. (A) The addition device and (B) the components.
At this point, the gaseous NO2 will have been fully aspirated or dissolved in the added water according to the eq 2. Because the concentration of nitrate ions is much smaller after dilution, the
Table 1. Addition-Device Components part no.
part name
cost
UPC
Lowe’s part no.
1
4 in. 2 in. coupling
$5.44
0012871059513
23321
2
2 in. 1/2 in. bushing threaded
$1.80
0025528134846
51013
3
1/2 in. nipple threaded
$0.33
0025528118129
56001
4
1/2 in. valve threaded
$2.65
0054211517208
21481
5 6
1/2 in. adapter slip threaded 1/2 in. pipe, 10 ft section
$0.34 $1.53
0025528131661 0024599050512
23855 23966
7
any all purpose PVC cement 1688
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DEMONSTRATION
below.) The equation for the reaction that produces the green solution is
Table 2. Characteristics of Copper Pennies diameter/ thickness/ mass/ surface area/ type of penny
mm
mm
g
cm2
Cu (%)
1900s: pre-1982
18.9
1.4
3.1
6.5
95
mid-1800s
27.8
2.2
10.1
½CuðH2 OÞ6 2þ ðaqÞ þ NO2 ðaqÞ f ½CuðH2 OÞ5 ðNO2 Þþ ðaqÞ þ H2 OðlÞ green
blue
14
ð4Þ
95
At pH = 5 a cloudy, a light blue precipitate begins to form (Figure 5C). The solution was filtered before measuring the spectrum of the solution, and the green color persists up to at least pH = 5.6. Additional spectra were taken after subsequent additions of aqueous ammonia. At pH = 7.9 there was so little copper(II) in solution that the absorbance maximum could not be determined. The precipitation reaction is ½CuðH2 OÞ5 ðNO2 Þþ ðaqÞ þ 2OH ðaqÞ f green
CuðOHÞ2 ðsÞ þ 5H2 OðlÞ þ NO2 ðaqÞ
ð5Þ
light blue ppt
Figure 3. Pennies from (left) 1900s and (right) 1800s.
At pH = 8.5 the absorbance maximum had shifted to 620 nm and the solution was the deep blue color characteristic of tetraamminecopper(II) ions (Figure 5D). CuðOHÞ2 ðsÞ þ 4NH3 ðaqÞ f ½CuðNH3 Þ4 2þ ðaqÞ þ 2OH ðaqÞ light blue ppt
dark blue
ð6Þ
Figure 4. (A) Nitrogen dioxide filling the flask and (B) the green color due to nitrate complexes of copper(II).
green nitrato complexes of copper(II) are converted to blue [Cu(H2O)6]2+ (Figure 5A). ½CuðH2 OÞ4 ðNO3 Þ2 ðaqÞ þ 2H2 OðlÞ f green
½CuðH2 OÞ6 2þ ðaqÞ þ 2NO3 ðaqÞ
ð3Þ
blue
This is a good time to explain that the blue color is due to Cu2+(aq), that is, the [Cu(H2O)6]2+ ion. To demonstrate three more forms of copper, a rough titration of the diluted reaction mixture with aqueous ammonia is done. The aqueous ammonia, with a concentration of 13 mol/L, can be introduced through the funnel. (If ventilation is not adequate and ammonia vapors become too concentrated for comfort, a less concentrated solution can be used; the volume of water added needs to be decreased to accommodate the larger volume of aqueous ammonia that will be required.) An interesting aspect of this demonstration is that the solution becomes green after about 250 mL of 1.0 M NH3(aq) has been added. (This does not happen when ammonia is added to aqueous copper(II) nitrate.) The green solution appears in the pH range between 2 and 5. Although it is usually not a good idea to describe what will happen in a demonstration before it happens, at this point it would be useful to ask students whether they see any difference in color (some color-blind students may not be able to). The green color is shown in Figure 5B. (This green color is attributed to [Cu(H2O)5(NO2)]+ ions, based on the results of several experiments that are described in detail
If the solution is intensely colored, it is a good idea to back light the flask to demonstrate that the precipitate is gone. At pH = 8.7 the absorbance maximum was 610 nm. Further addition of 1.0 M NH3(aq) did not change this maximum but did increase the absorbance. To complete the demonstration a total of about 450 mL of 1.0 M NH3(aq) was required. The formation of copper(II) hydroxide precipitate and tetramminecopper(II) ion in solution has been described in detail by Johnson, McQueen, and Rodolfa.25 For the convenience of instructors, two videos of the demonstration, with and without equations for the reactions, may be seen on the Internet.26
’ IDENTITY OF THE CHROMOPHORE IN THE GREEN SOLUTION Understanding why a green solution appears after the addition of aqueous ammonia in this demonstration, but does not appear when aqueous ammonia is added to aqueous copper(II) sulfate or copper(II) nitrate, is a good puzzle for students who know some coordination chemistry and acidbase equilibrium chemistry. To support the hypothesis that a complex involving nitrite ions was involved, two experiments were performed. Interpretation of these experiments illustrates the relationship between color of a solution and visible absorption spectrum. In one experiment, the demonstration was carried out in a hood at one-fourth scale. The spectrum and pH were measured immediately after water was added to quench the reaction and after each of a series of 10 mL aliquots of 1.0 M NH3(aq) was added. Spectra and pH data are provided in a spreadsheet in the Supporting Information. Because some of the NO2(g) dissolves in the water, HNO2 is present as well as nitric acid. For HNO2, pKa = 3.15.27 Therefore, until the pH reaches about 2, nearly all of the HNO2 is un-ionized. That is, the equilibrium HNO2 ðaqÞ h Hþ ðaqÞ þ NO2 ðaqÞ 1689
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DEMONSTRATION
Figure 5. (A) Solution containing [Cu(H2O)6]2+ ions, (B) solution containing [Cu(H2O)5(NO2)]+ ions, (C) copper(II) hydroxide beginning to form, and (D) solution containing tetraamminecopper(II) ions.
lies to the left and the concentration of nitrite ions is quite small. From pH = 2 to pH = 4 the HNO2 is converted to NO2, and it is in this pH range that the solution changes from blue to green. Formation of a complex with nitrite is borne out by the spectra measured as a function of volume of ammonia added. At pH = 1.3 the absorbance in the green has begun to decrease, and by pH = 1.4 the absorbance in the violet and blue has begun to increase. The minimum absorbance is 440 nm at pH = 1 and changes to 550 nm at pH = 2.5. The minimum remains at 550 nm up to pH = 5.6; above this pH, precipitation of copper(II) hydroxide reduces the concentration of copper ions sufficiently that it is difficult to determine absorbance minima or maxima. This change in the spectrum agrees with the spectra reported by Kossiakoff and Sickman for nitrite complexes of copper(II).28 They report an absorbance maximum for [Cu(H2O)5(NO2)]+ at 750 nm (compared to 800 nm for [Cu(H2O)6]2+). For the demonstration reaction at pH = 5 the spectrum has shifted toward shorter wavelengths, but the maximum absorbance is about 785 nm, indicating that not all of the copper(II) is complexed with nitrite. On the basis of another experiment and the equilibrium constants reported by Kossiakoff and Sickman28 no more than a single nitrite ion is coordinate-covalently bonded to the copper(II) ion. To simulate the demonstration reaction after quenching with water, a solution of Cu(NO3)2 with the same concentration as the demonstration reaction was prepared (based on the mass of copper that dissolved and the volume after water was added). To this was added enough sodium nitrite to produce the concentration of nitrite ions that would have been obtained if all of the NO2 had dissolved in the water. (Based on the equations given in the previous section, there is 1:1 stoichiometry between Cu dissolved and HNO2 produced.) This simulates the maximum concentration of nitrite ions possible in the solution. Nitric acid was added to adjust to pH = 1 (which was the pH of the demonstration solution immediately after addition of water). Then 20 mL aliquots of 1.0 M NH3(aq) were added and the pH and spectrum were recorded for each addition. The spectra change in the same way as the spectra for the demonstration reaction. The concentration of nitrite is clearly larger in this simulation because the solution is green at a lower pH and the absorbance maximum shifts from 800 to 776 nm. Based on the concentrations of copper(II) and nitrite ions (0.084 M) and the formation constant, K1 = 16 (at ionic strength = 1), reported by Kossiakoff and Sickman, less than half of the copper(II) would have nitrite coordinated, even for the highest concentration of nitrite ion that could possibly have been present in the demonstration reaction. The formation constant for [Cu(H2O)4(NO2)2], K2 = 1.6, is much smaller, so only a very small concentration of [Cu(H2O)4(NO2)2] should be
present. On the basis of the spectra of the solution and the reported equilibrium constants, the equation for formation of the green solution is ½CuðH2 OÞ6 2þ ðaqÞ þ NO2 ðaqÞ h ½CuðH2 OÞ5 ðNO2 Þþ ðaqÞ þ H2 OðlÞ
ð8Þ
’ HAZARDS Both the liquid and vapors of aqueous ammonia are extremely irritating especially to the eyes. Dispense aqueous ammonia in a hood and be sure an eyewash is accessible. Aqueous ammonia is moderately toxic by ingestion and inhalation and a serious respiratory hazard. LD50 is 350 mg/kg.30 According to Flinn’s suggested disposal method (#1031,32), the solution of tetraamminecopper(II) should be neutralized and flushed down the drain with excess water. This method should be used only if the drain is connected to a sanitary sewer system. Check state and local regulations to be certain that procedures are in compliance. Concentrated nitric acid is corrosive, a strong oxidant, and toxic by inhalation and ingestion; avoid contact with bases, acetic acid, and readily oxidized substances. Gloves should be used to protect skin, and if concentrated nitric acid is spilled on the skin, it should be washed off immediately with a large quantity of water. Concentrated nitric acid will attack and destroy metals and most plastics. TLV is 5.2 mg/m3.33 Nitrogen dioxide forms acid on contact with moisture and is a poisonous, choking gas. The reaction of copper and nitric acid should only be done in a hood unless the NO2 produced is contained or reacted, as it is by the equipment used in this demonstration. ’ ASSOCIATED CONTENT
bS
Supporting Information Spectra and pH data are provided in a spreadsheet; Ira Remsen’s description of the experience. This material is available via the Internet at http://pubs.acs.org.
’ AUTHOR INFORMATION Corresponding Author
*E-mail:
[email protected]. Notes
Randy Sullivan (Department of Chemistry, University of Oregon, Eugene, OR 974031253) tested this demonstration. 1690
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’ ACKNOWLEDGMENT We thank UW—Madison undergraduate students John S. Putze and Anna M. Fleischman for their help with experiments and for preparing graphs of experimental data. We thank Riverside High School students Emily M. Henderson and Jacary O. Richardson for producing the YouTube video. ’ REFERENCES (1) Ramette, R. W. J. Chem. Educ. 1980, 57 (1), 68–69. (2) Bent, H. A; Bent, B. E. J. Chem. Educ. 1987, 64 (3), 249–251. (3) Hauben, M.; Rayner-Canham, G. J. Coll. Sci. Teach. 1996, 15 (5), 368–71. (4) Lang, M. C.; Showalter, D. L.; Shulfer, G. J. Yes, Virginia, Learning Chemistry Can Be Fun; University of Wisconsin: Stevens Point, 1992; pp 23. (5) Shakhashiri, B. Z. Chemical Demonstrations: A Handbook for Teachers of Chemistry; University of Wisconsin Press: Madison, 1985; Vol. 2, pp 163166. (6) Shakhashiri, B. Z. Chemical Demonstrations: A Handbook for Teachers of Chemistry; University of Wisconsin Press: Madison, 1989; Vol. 3, pp 8391. (7) Summerlin, L. R., Borgford, C. L., Ealy, J. B. Chemical Demonstrations: A Sourcebook for Teachers; 2nd ed.; American Chemical Society: Washington, DC, 1988; Vol. 2, pp 45. (8) Nitric Acid Acts Upon Copper. http://www.chem.umn.edu/ services/lecturedemo/info/acid_rain.html (accessed Sep 2011). (9) Copper and Nitric Acid. http://www.angelo.edu/faculty/kboudrea/ demos/copper_HNO3/Cu_HNO3.htm (accessed Sep 2011). (10) Metallic Copper and Concentrated Nitric Acid. http://www. chemmybear.com/demo.htm (accessed Sep 2011). (11) Demonstrating Nitric Acid Acts Upon A Copper Penny Experiment. http://www.instructables.com/id/Nitric-acid-acts-upona-copper-penny-experiment/ (accessed Sep 2011). (12) Group Observing Reaction in a Fume Hood. http://www. youtube.com/watch?v=azMYIe47Uwk (accessed Sep 2011). (13) Copper and Nitric Acid Plus Nitrogen Dioxide and Water. http:// www.youtube.com/watch?v=5tY1Dh30j08 (accessed Sep 2011). (14) Smoking Penny. http://www.youtube.com/watch?v=KlOEskpy2-0 (accessed Sep 2011). (15) Close Up View of Copper and Nitric Acid. http://www. youtube.com/watch?v=NT5VKw0D6ME (accessed Sep 2011). (16) Copper and Nitric Acid Fountain. http://www.youtube.com/ watch?v=FdMrS1w2H-w (accessed Sep 2011). (17) Penny Operated Fountain. http://www.kentchemistry.com/ KentsDemos.htm (accessed Sep 2011). (18) Copper Powder and Nitric Acid. http://www.youtube.com/ watch?v=OoddtIlN5jU (accessed Sep 2011). (19) Nitric Acid and Penny Demonstration in Fume Hood. http:// www.youtube.com/watch?v=IDTbwEVA6AU&feature=fvst (accessed Sep 2011). (20) Biography of Ira Remsen. http://en.wikipedia.org/wiki/Ira_ Remsen (accessed Sep 2011). (21) Gluing PVC Pipe. http://www.youtube.com/watch?v= Ep0kDsOXUI0 (accessed Sep 2011). (22) How To Glue PVC Pipe. http://www.youtube.com/watch? v=5_MsMEQ8Q9Y (accessed Sep 2011). (23) Composition of US Cents. http://www.usmint.gov/about_ the_mint/fun_facts/?action=fun_facts2 (accessed Sep 2011). (24) Addison, C. C.; Logan, N.; Wallwork, S. C.; Garner, C. D. Q. Rev. Chem. Soc. 1971, 25, 289322. 10.1039/QR9712500289. (25) Johnson, A. R.; McQueen, T. M.; Rodolfa, K. T. Species Distribution Diagrams in the Copper-Ammonia System: An Updated and Expanded Demonstration Illustrating Complex Equilibria. J. Chem. Educ. 2005, 82 (3), 408. (26) Better Ira Remsen Demonstration with Equations. http:// www.youtube.com/watch?v=HdNyunHuntw and a video without equations
DEMONSTRATION
http://www.youtube.com/watch?v=nHSPnpznBvA . The URLs are case sensitive, and the videos are unlisted (accessed Sep 2011). (27) Lumme, P.; Tummavuori, J. Acta Chem. Scand. 1965, 19, 617. (28) Kossiakoff, A.; Sickman, D. V. J. Am. Chem. Soc. 1946, 68 (3), 442–444. (29) Getman, F. H. The Life of Ira Remsen; Journal of Chemical Education: Easton, PA, 1940; pp 910. (30) Flinn Scientific Catalogue/Reference Manual: Flinn Scientific: Batavia, IL, 2010; p 62. (31) Flinn Scientific Catalogue/Reference Manual: Flinn Scientific: Batavia, IL, 2010; p 1129. (32) Old Glory - A Patriotic Colors Demonstration; Flinn Scientific Publication No. 438.00; Flinn Scientific: Batavia, IL, 2009. (33) Flinn Scientific Catalogue/Reference Manual; Flinn Scientific: Batavia, IL, 2010; p 136.
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