In the Classroom edited by
JCE DigiDemos: Tested Demonstrations
Ed Vitz
Kutztown University Kutztown, PA 19530
Electrochemical Polishing of Silverware: A Demonstration of Voltaic and Galvanic Cells
submitted by: Michelle M. Ivey* and Eugene T. Smith Harriet L. Wilkes Honors College, Florida Atlantic University, Jupiter, FL 33418; *
[email protected] checked by:
Garry McGlaun Department of Chemistry, Gainesville College, Gainsville, GA 30503
James H. Niewahner Department of Chemistry, Northern Kentucky University, Highlands Heights, KY 41076
Stainless steel flatware is used more often than flatware made of sterling silver or plated silver, not only because of its lower cost, but also because it does not need to be polished to maintain its beauty. Over the years, many different polishes, solutions, and devices for polishing silverware have been marketed to the public. One such device uses a metal plate that, when submerged in hot water with the addition of an “activator”, will “magically” polish any silverware that touches the plate (1, 2). In all likelihood, the metal plate is aluminum and the activator is baking soda (sodium bicarbonate). This chemistry has previously appeared as a classroom activity where tarnished pieces of silver were generated by exposing the silverware overnight, or longer, to one of the following: powdered sulfur (3), a hardboiled egg (3), an egg yolk (4), mayonnaise (3), mustard (3), or a rubber band (4). This chemistry was also demonstrated in an article entitled “All Natural Cleaning” in Good Housekeeping magazine (5). What makes this demonstration so attractive to the students is the fact that silver can be “polished” using two items routinely found in the kitchen: aluminum foil and baking soda. Here we offer a significant improvement by incorporating an additional step to instantly tarnish the silverware as part of the demonstration before the tarnish is electrochemically removed. Tarnish on silverware is a result of exposure to pollutants in the atmosphere. The rate at which flatware will tarnish will depend upon the geographical location. The Clean Air Act Amendments of 1990 has resulted in a decrease in SO2 emissions in most states (6), and thus silver will tarnish with a much slower rate than in years past. As sulfur-containing species in the atmosphere decrease, other compounds will play a larger role in the corrosion of silver. For example, Bouquet et al. found that silver chloride was more abundant than silver sulfide in tarnish in Paris (7). As a result, the chemical composition of pieces that have tarnished naturally may vary widely and may yield less reproducible results when used in a demonstration. Another difficulty is the rate of atmospheric tarnishing; in some areas it may take several years for a noticeable golden patina to develop. When using naturally tarnished silver pieces to demonstrate electrochemical conversion back to metallic silver in a large lecture hall, the difference in the golden patina and the polished silver may not be dramatic enough for students in the back row to appreciate. Tarnishing by exposure to sulfur powder or sulfurcontaining foods is messy and can take several days to develop a dark patina. A layer of tarnish can be generated by immersing a utensil in a solution of (NH4)2S2 (8); however, it is less obvious 68
that this process is a redox reaction. Therefore, we have incorporated an additional step where an electrolytic cell, powered by a 9 V battery, is used to generate a layer of Ag2S on the silver surface, which can then be removed using the aluminum–baking soda system to convert it back to silver metal. This improvement offers several advantages: (i) The black colored Ag2S layer is much more dramatic than a tarnish layer that is generated over a few months or years of atmospheric exposure, (ii) the layer is more reproducible, containing only silver sulfide, (iii) silver pieces of any size can be tarnished easily, and (iv) both electrolytic and galvanic cells are employed in the same demonstration. This demonstration shows how a battery and a graphite electrode can be used to facilitate the formation of a layer of tarnish (Ag2S) when a piece of silverware is immersed in a sodium sulfide solution. The second half of the demonstration leads students to use their knowledge of electrochemistry to utilize household items such as aluminum foil and baking soda to electrochemically polish the silverware by reducing the silver sulfide to form silver metal. Procedure Electrolytic Cell To Tarnish Silver To construct the electrolytic cell, patchcords are used to connect an untarnished silver utensil, such as a fork, to the (+) terminal of the 9 V alkaline battery. The fork serves as the anode. The cathode for the reaction is a ¼ in. graphite rod, connected to the (−) terminal of the battery. Initially, only the fork is placed in the sodium sulfide solution (0.1 M is more than sufficient) to show that nothing happens. Only when the graphite electrode is placed in the solution will a reaction occur. The silver that is submerged in the solution turns black almost immediately, and bubbles can be seen being generated at the carbon electrode, owing to the formation of hydrogen gas.
2H2O 2e
H 2OH 2
(1)
The carbon electrode only needs to be immersed for a few seconds before the eating utensil is impressively tarnished. The fork is then removed from the solution and rinsed with water. When performing this part of the procedure, it is important to note that once the utensil has been placed in the sodium sulfide solution, it should not be taken out of solution to show that nothing has happened. This is because as soon as the silver
Journal of Chemical Education • Vol. 85 No. 1 January 2008 • www.JCE.DivCHED.org • © Division of Chemical Education
In the Classroom Table 1. Some Useful Electrochemical Potentials Eq
E 0/V
Reaction
i
O2 4Η 4e
ii
2H 2e
iii
Ag2S 2Η 2e
iv
Ag2S 2e
2Ag S2
‒0.691
v
Zn2 2e
Zn
‒0.7681
vi
2H2O 2e
vii
Al3 3e
2H2O
H2
1.299 0.000
2Ag H2S
H2 2OH
Al
‒0.0366
‒0.8277 ‒1.662
Note: Potentials from ref 7.
surface coated with sodium sulfide solution is exposed to air, the oxygen in air is a sufficient oxidizing agent to oxidize the silver to silver sulfide. This effect will also occur at the part of the utensil that is at the air solution interface, so it is best to not let it sit in the sulfide solution for very long. Electrochemically Polishing the Tarnished Utensil To electrochemically polish the silverware, half-fill a 2000 mL beaker with water and heat on a hotplate. The water must be sufficiently deep that the tarnished area of the fork will be completely submerged. However, it is important to have the beaker only half-full, because as baking soda is added, gaseous CO2 will be produced according to eqs 2 and 3, where Keq is the equilibrium constant for reaction 3 and Kb1 is the dissociation constant for the bicarbonate ion and is related to the first acid dissociation constant for carbonic acid, H2CO3 by Kb1 = Kw∙Ka1.
HCO3 (aq) H2O(l)
H2CO3(aq)
Kb1
Keq
H2CO3(aq) OH (aq)
CO2(g) H2O(l)
(2) (3)
When the water is at or near boiling (95–100 °C), fold up a ~12 in. square piece of foil and use tongs to submerge it in the water. If a lower temperature is employed, the electrochemical polishing will be slower. Slowly add ~10 cm3 of baking soda for every 100 mL of water, taking care that the evolving CO2 does not cause water to overflow the beaker. A small beaker is a convenient measure for the solid sodium bicarbonate. Use the tongs to place the tarnished fork in the beaker containing the aluminum foil, making sure that the fork is firmly in contact with the foil. The tarnish will disappear within 1–2 minutes. Use the tongs to remove the fork from the beaker (Caution: it is hot), use water to cool it so that it can be handled, and then use a paper towel to wipe off any silver sulfide particles that may be stuck onto the fork. The tarnish has been electrochemically converted back to solid silver.
If the demonstration is to be performed multiple times, it may be necessary to add more baking soda to the beaker of hot water. As silver is polished, the aluminum ions react with the hydroxide ions in solution to produce solid aluminum hydroxide, which appears as a white powder. This will cause the reaction shown in eq 2 to shift to the right according the Le Châtelier’s principle, producing more carbonic acid, which will then result in CO2 gas evolution, as shown in eq 3. Even when silver is not being polished, carbon dioxide will continue to come out of solution, because the partial pressure of CO2 in the atmosphere is small (9). A good rule of thumb is to add baking soda until the addition of more causes quite a lot of fizzing. One consequence of using a voltaic cell to tarnish the silverware, and then converting it back to silver metal is that upon closer inspection of the item, it is possible to tell where the utensil was tarnished and then reconverted, probably owing to some roughening of the surface in the process. This is not really a problem for the demonstration, because the effect is fairly subtle, but it is probably best not to use expensive silverware for this demonstration. While this effect can be rectified by using silver polish to smooth out the surface, scratched up silverware purchased at yard sales or Internet auctions is a more practical approach. (Look for very tarnished pieces so that you know they are actually silver. Such pieces are usually relatively inexpensive as well.) Hazards During the process to remove the tarnish, care must be taken with the boiling water, especially when adding the baking soda. The sodium sulfide solution should be collected for disposal as hazardous waste. Discussion Voltaic Cell To Oxidize Silver to Silver Sulfide Initially, the students are asked, “If silverware is placed in a solution of sodium sulfide, will it tarnish?” The species present in solution are Na+, S2−, and H2O. If silver is to be oxidized to silver sulfide as shown as the reverse of eq iv in Table 1, then there must be some species that is reduced. The Na+ cannot be reduced to sodium, and therefore the only species that can react is water. Combining eqs iv and vi in Table 1, we obtain the overall balanced equation:
2Ag S2 2H2O
Ag S H 2OH 2 2
(4)
Because hydrogen is not present significantly in the atmosphere and the concentration of hydroxide is small, this reaction is spontaneous under atmospheric conditions, however the evolution of hydrogen has a large energy of activation, which can be overcome by supplying a large enough overpotential using a 9 V battery, as described above in the demonstration section. Electrochemical Conversion of Silver Sulfide to Metallic Silver Initially, the students are shown an advertisement for a product used to “magically” polish tarnished silverware (1, 2). It shows what appears to be a metal plate, which is immersed in hot water with the addition of an activator, which is a white powder that dissolves in the water. The picture shows that when
© Division of Chemical Education • www.JCE.DivCHED.org • Vol. 85 No. 1 January 2008 • Journal of Chemical Education
69
In the Classroom
the silver item is touched to the plate, any of the tarnish below the waterline is removed. The students are then challenged to try and discern the identity of both the metal plate and the activator, given their knowledge of general chemistry. The first question to ask is “Other than the fact that we have been talking about electrochemistry recently, why might we think that this is an example of electrochemistry?” Some of the typical answers may be: (i) Silver changes its oxidation number from +1 in Ag2S to 0 in metallic silver. (ii) Nothing happens until the item touches the plate, which suggests that an electrical circuit must be completed for the reaction to occur. (iii) Only the part that is under water will react, which is similar to the behavior in standard galvanic cells, such as a Cu∙Zn cell in which the zinc electrode dissolves and the copper electrode grows, both of which occurs at the surface of the electrodes immersed in the solution. Therefore, the next questions are “What could the polishing plate be made of ?” and “What properties should it have?” (i) It appears to be metallic. For the rest of the discussion, assume that the plate is a pure metal. (ii) It must not react with both air and water. (iii). All products of the reaction should be non-toxic, given that this is being used to polish eating utensils. (iv) The plate should be relatively inexpensive to produce. (v) It needs to be something that is easily oxidized. Figure 1 summarizes the logic process used to determine the identity of the metal plate. From Table 1, we can see that Ered° for silver sulfide is ‒0.691 V, and therefore, a metallic species, M, that has a reduction potential more negative than this is required to reduce the silver sulfide back to silver metal in the presence of 1 M sulfide; somewhat less negative potential is required at low sulfide concentrations.
Mn ne
M(s)
Eredp 0.691 V
yes
Li, Na, Mg, Al, Cr, Zn, Mn, K, Ca, Ba, La
Li, Na, K, Ba
no
stable in air yes
Mg, Al, Cr, Zn, Mn, Ca, La
Mg, Ca, La
no
stable in H2O yes
Al, Cr, Zn, Mn
Cr, Mn
no
non-toxic yes
Al, Zn
(5)
The students are asked to use the table of reduction potential found in their textbook (10) to suggest possible candidates for M, which are then discussed. These tables are selective, and therefore other candidates may arise depending on the textbook used. The first ones that can be eliminated as reactants are the alkali metals and alkaline earth metals, because these react (often violently) with air or water. Thus Mg, Na, Ca, Ba, K, Li are eliminated. For the same reason, La can also be rejected, as it reacts with water to produce hydrogen gas. Manganese can be eliminated because in some forms, especially powders, is poisonous if inhaled, and chromium because its ions are toxic [the Cr(VI) ion was showcased in the movie Erin Brockovich]. Therefore, we are left with Zn∙Zn2+, with a reduction potential of ‒0.7618 V, and Al∙Al3+, with a reduction potential of ‒1.662 V. Because a more negative potential indicates that there is more of a thermodynamic driving force for the metal to reduce the silver sulfide to silver metal, aluminum is a better candidate for the metal plate. Aluminum also has the advantage of being less expensive per pound than zinc (11) and more readily available. So the last question is “What is the identity of the ‘activator’ and what is its role?” One of the most important considerations is that the compound must be completely safe to use on eating utensils, and so the students are encouraged to think of food or food ingredients that are white powders. Some possibilities include salt, flour, sugar, baking soda, and baking powder. Flour can be eliminated because it will not dissolve in water, and salt because salt will corrode silver. To choose between the other possibilities, the role of the activator must be examined. 70
Eredp 0.69 V
Zn
no
lowest cost yes
Al
Figure 1. Flow chart for determining the identity of the metal plate.
If equations iv and vii from Table 1 are combined, the overall balanced reaction between aluminum and silver sulfide is
3 Ag2S 2Al
6 Ag 2Al 3 3S 2
(6)
Aluminum sulfide decomposes in water (9), and therefore we would not expect this compound to precipitate. No additional species show up in this equation, but the presence of the activator definitely accelerates the reaction, probably by removing Al3+ ions from solution by forming a compound and precipitating. Therefore sugar can be eliminated. Baking powder is essentially baking soda with the addition of acidic salts, which is undesirable because acidic conditions can result in the formation of hydrogen sulfide by the reaction of sulfide ions with protons. Baking soda (sodium bicarbonate) will make the solution slightly basic promoting the formation of aluminum hydroxide (Ksp = 1.9 × 10‒33) (12), as well as inhibiting the formation of hydrogen sulfide.
Journal of Chemical Education • Vol. 85 No. 1 January 2008 • www.JCE.DivCHED.org • © Division of Chemical Education
In the Classroom
So now we have identified the metal plate as being, in all likelihood, aluminum, and the activator as sodium bicarbonate. We prefer to use a box of baking soda, rather than using research grade sodium bicarbonate, because this emphasizes to the students that the chemistry really is being performed with everyday items, rather than “chemicals”. The students enjoy this demonstration because it is chemistry with a real-life application, and because they can use their knowledge of electrochemistry to solve the mystery to determine the possible composition of the cleaning plate and the “activator”. Note Added in Proof Measurements of the second dissociation of H2S show that in solution, sulfide exists as HS–, not S2– (13). In this manuscript, sulfide is represented as S2– in electrochemical potentials to be consistent with general chemistry textbooks and the CRC Handbook of Chemistry and Physics (9). Literature Cited 1. Museum Precious Metals Cleaning Plate at Hammacher Schlemmer. http://www.hammacher.com/publish/65593.asp# (accessed Sep 2007). 2. Hammacher Schlemmer. Sky Mall, Summer, 2006, p 46. 3. JCE Editorial Staff. J. Chem. Educ. 2000, 77, 328A.
4. Borgford, Christie L.; Summerlin, Lee R. Chemical Activities; American Chemical Society: Washington, DC, 1988, p 164. 5. De Jong, Michael. All Natural Cleaning. Good Housekeeping, May, 2006, pp 58–59. 6. Butler, Thomas J.; Likens, Gene E.; Stunder, Barbara J. B. Atmos. Environ. 2001, 35, 1015–1028. 7. Bouquet, Simone; Bodin, Camille; Fiaud, Christian. C. R. Acad. Sci. Paris 1993, 316, 459–464. 8. Feinstein, H. I. J. Chem. Educ. 1976, 53, A34. 9. CRC Handbook of Chemistry and Physics, 71st ed.; Lide, David R, Ed.; CRC Press: Boca Raton, FL, 1991. 10. Zumdahl, Steven S.; Zumdahl, Susan A. Chemistry, 6th ed.; Houghton Mifflin Company: Boston, 2003. 11. Current Primary and Scrap Metal Prices. http://www.metalprices. com/ (accessed Sep 2007). 12. Whitten, Kenneth W.; Davis, Raymond E.; Peck, M. Larry; Stanley, George G. Chemistry, 8th ed.; Thomson Brooks/Cole: Belmont, CA, 2007. 13. Myers, R.J. J. Chem. Educ. 1986, 63, 687.
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