In the Laboratory
Spectator Ions ARE Important! A Kinetic Study of the Copper-Aluminum Displacement Reaction Sabrina G. Sobel* Department of Chemistry, Hofstra University, Hempstead, New York 11549-1510 *
[email protected] Skyler Cohen† Roslyn High School, Round Hill Road, Roslyn Heights, New York 11577 † Hofstra University Summer Science Research Program participant and 2009 graduate from Roslyn High School.
Metal displacement reactions are used often as classic examples of simple redox reactions. They can illustrate the electrochemical series and relative activities of metals (1). The copper-aluminum displacement reaction, in particular, typically appears as a general chemistry laboratory experiment (2). In the experiment herein, aqueous copper(II) chloride is reacted with household aluminum foil. The reaction is exothermic and can be completed by a student in about 20 min. The chemical equation of the reaction (eq 1) can be simplified to a net ionic equation (NIE, eq 2) that excludes the chloride ions because they are spectator ions: ð1Þ 3CuCl2 ðaqÞ þ 2AlðsÞ f 3CuðsÞ þ 2AlCl3 ðaqÞ 3Cu2þ ðaqÞ þ 2AlðsÞ f 3CuðsÞ þ 2Al3þ þ ðaqÞ
ð2Þ
The NIE implies that the chloride ions are unimportant to the course of reaction and that they can be ignored in calculations of expected mass of copper and percent copper recovered. However, interestingly, the chloride spectator ions accelerate the rate of reaction as compared to the same reaction with sulfate spectator ions. Aluminum Corrosion Aluminum corrosion is an important industrial concern (3). Studies have indicated that aluminum corrodes more quickly in seawater or when in contact with copper, either metallic or as Cu(II) in solution. Because Cu(II) has a positive reduction potential and Al(III) has a negative reduction potential, a spontaneous electrochemical reaction can be established easily, promoting aluminum oxidation. The “chloride-acceleration effect” is well established and is the major reason for accelerated aluminum corrosion in seawater. However, comparison of this effect with other halides has not been reported. Aluminum natively forms a tough, impervious oxide (Al2O3) layer on its surface upon natural exposure to air. This impedes the redox reaction until the oxide layer is removed. Under acidic (pH < 4) and basic (pH > 9) conditions, the oxide layer dissolves spontaneously (3) and accelerates aluminum corrosion as well. Kinetics of the Cu-Al Reaction Several groups have investigated the kinetics of the Cu-Al displacement reaction in the presence and absence of chloride 616
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ions (4-6). Because this is a biphasic reaction, the kinetics is more complicated than standard first-order homogeneous reactions. To keep reaction conditions constant, researchers used a rotating disk of Al during reaction. The accepted mechanism of reaction involves a five-step process: (i) diffusion of Cu2þ to the Al surface; that is, mass transfer; (ii) adsorption of Cu2þ ions; (iii) electron transfer from Al to Cu2þ; that is, surface reaction; (iv) desorption of Al3þ ions; and (v) diffusion of Al3þ ions away from the surface. Usually, step i or iii is considered to be the ratedetermining step. Preceding processes that must occur before copper deposition include oxide dissolution, copper nucleation on the surface, and increase in surface roughness, which are all dependent on the Al surface area available. These preceding processes cause the reaction to have an induction time in which reaction of copper is slow. The integrated rate law (5) is lnð½CuðIIÞ0 =½CuðIIÞÞ ¼ ðkA=V Þft - tind ð1 - θ0 Þð1 - expð - t=tind ÞÞg
ð3Þ
in which A is the surface area of aluminum, θ0 is the fraction of bare surface of Al not covered by Al2O3, [Cu(II)]0 is the initial concentration of the copper solution, V is the volume of copper solution, and tind is the induction time. When t is large, [1 exp(-t/tind)] approaches one, and the best fit to experimental data assumes θ0 = 0 (5). Assuming θ0 = 0 is reasonable since household aluminum foil has been exposed to air since manufacture; no special precautions have been taken to prevent formation of the Al2O3 surface layer. With these simplifying assumptions, the integrated rate law becomes first-order: ð4Þ lnð½CuðIIÞ0 =½CuðIIÞÞ ¼ ðkA=V Þft - tind g Plotting ln([Cu(II)]0/[Cu(II)]) versus t for the reaction region results in a straight-line fit in which the slope equals kA/V, and the x intercept can be used to calculate tind. Experiment Materials Nanopure water was used for all solutions. Copper(II) salt solutions (∼0.2 M) were prepared using CuCl2 3 2H2O (Fisher Scientific, 99%), CuBr2 (Mallinckrodt, 99%), or CuSO4 3 5H2O (Fisher Scientific, 98.4%). America's Choice Aluminum foil,
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Vol. 87 No. 6 June 2010 pubs.acs.org/jchemeduc r 2010 American Chemical Society and Division of Chemical Education, Inc. 10.1021/ed1001703 Published on Web 04/08/2010
In the Laboratory
Figure 1. Typical results for the reaction of the CuCl2 and CuBr2 solutions with aluminum foil.
Figure 2. Typical results for the reaction of the CuSO4 solution with aluminum foil. Table 1. Effect of Spectator Ion on Reaction Kinetics
not heavy duty, was used. The household aluminum foil (not heavy duty) is composed of 97.6-98.3% Al plus other minor components (7, 8) and has a thickness of 0.013 mm. The surface area of the aluminum foil has to be computed appropriately, considering that both sides of the aluminum foil react. Data Collection Two 3-h lab periods were needed to conduct this experiment. The CuCl2 and CuBr2 runs took 20 min each including preparation and the CuSO4 run took 130-150 min, depending on the number of data points desired. The UV-vis spectra can be obtained in the following lab period if UV-vis samples in the cuvettes were capped. Aluminum foil (∼0.4 g) was weighed and cut into approximately 0.5 in. squares. The copper(II) salt solution (100.0 mL of ∼0.2 M CuSO4 3 5H2O or 50.0 mL of ∼0.2 M CuCl2 3 2H2O or 50.0 mL of ∼0.2 M CuBr2 and 50.0 mL of H2O) was placed in a water-jacketed beaker. The beaker was attached to a water circulating bath set to 20.0 °C. Aluminum foil was added to the copper(II) salt solution, and the solution was stirred. Aliquots (1.00 mL) of the reacting solution were removed at designated intervals and filtered through a 0.45 μm Millipore filter into 1.00 mL plastic UV-vis cuvettes (1.00 cm path length). UV-vis spectra (500-900 nm) of the samples (including time zero) were collected on a Cary 300 UV-vis spectrophotometer. A water background blank was used. Hazards The Cu(II) salts are harmful by ingestion and are skin and eye irritants. Typical Results Analysis of the copper(II) chloride-aluminum foil reaction using the previously discussed kinetic model (5) worked well (Figure 1). The reaction region, 2-5 min, yielded a good linear fit, with k = 0.13(1) cm/min and tind = 1.22 min. The reaction of copper(II) bromide and aluminum was also studied (Figure 1). The copper(II) bromide reaction was comparable to that of chloride, but slightly slower with k = 0.106(3) cm/min and tind = 1.87 min. Copper(II) iodide could not be compared because it spontaneously decomposes to copper(I) iodide and iodine (9). For comparison, the reaction of copper(II) sulfate and aluminum was completed (Figure 2). In this case, the reaction is about 100 times slower, clearly demonstrating the halide-acceleration
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k/(cm 3 min-1)
Spectator ion ClBr
0.13
-
SO42-
tind/min 1.22
0.11
1.87
4.7 10-4
7.91
effect. The rate constant k was found to be 4.7(7) 10-4 cm/ min, and the induction period (tind) was 7.91 min. Discussion The kinetic model developed for a rotating aluminum disk reaction with a copper(II) solution (5) works well for the simpler conditions of household aluminum foil with a copper(II) solution. This simpler system may be more relevant to real-world applications in which less-than-pure aluminum samples are exposed to high chloride concentration and high ionic strength environments such as seawater ([Cl-] = 0.55 M, I = 1.1 M) (10). The data from the reactions described here show that bromide also accelerates the rate of the Cu-Al displacement reaction. The effect was found to be significant, speeding up the reaction by about 100 relative to a nonhalide spectator ion and shortening the induction period (Table 1). Therefore, a more general term, the “halide-acceleration effect”, should be used instead of the chloride-acceleration effect. The mechanism used to speed up the reaction is unclear. It seems reasonable to propose that the halides are interacting with the aluminum ions, possibly through formation of complex ions. The aluminum ion is a hard, strong Lewis acid that complexes more strongly with smaller halides (AlF4-, log K = 15.12) (10), which is consistent with the greater acceleration effect of chloride as compared to bromide. On the basis of the accepted mechanism, if step iii is the rate-determining step, then halide complexation of aluminum ions may be enhancing the favorability of electron transfer from aluminum to copper(II) ions. Unfortunately, the only thermodynamic information available for aluminum-chloride complexation is for the addition of one chloride and that value is low (log K = -1.0) (11). Halide complexation may be enhanced within the dissolving Al2O3 layer through an unidentified mechanism. Studies of the adsorption of chloride and sulfate indicate that these ions adsorb and interact differently with the surface aluminum oxide layer (12). More sulfate adsorbs on the Al2O3 surface (1.75 10-15 ions/cm2) than chloride (0.5 10-15 ions/cm2), but chloride ions are adsorbed more irreversibly,
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with about 70% of adsorbed ions incorporated into the oxide layer, as compared to 25% of sulfate ions. It is reasonable to propose that chloride ions accelerate dissolution of the Al2O3 layer through irreversible adsorption even though the overall anion density for both anions is comparable. Another factor to consider is the pH of the solution created by dissolved copper(II) ions. According to the literature (3), the Al2O3 layer dissolves spontaneously in solutions of pH lower than 4. To test the possible influence of pH on rate of reaction, the pH of each starting copper(II) salt solution was measured. The pH of the 0.1935 M CuCl2 solution was 3.443, and that of the 0.1937 M CuSO4 solution was 3.899, both below the pH 4 threshold for spontaneous Al2O3 dissolution. In addition, a 0.0387 M CuCl2 solution with a pH of 4.154 also exhibited accelerated reaction with aluminum. Therefore, the pH of the solution is not a dominant factor in creating the halideacceleration effect. An extension of this experiment could be to test this conclusion by buffering the solution with a pH 5 buffer. Conclusion This set of reactions nicely demonstrates that spectator ions influence the rate of the biphasic copper(II)-aluminum metal displacement reaction, showing a significant halide-acceleration effect. The published kinetic model (5) works well to describe the simple reaction of household aluminum foil with aqueous copper(II) salt solutions. This series of experiments can demonstrate to students that spectator ions cannot always be ignored, in a way that connects fundamental chemistry to real-life situations (aluminum corrosion). It also points out the threshold of knowledge and a place where more research is needed, emphasizing the constantly evolving nature of the scientific enterprise. Students and colleagues find this lack of explanation stimulating. Hypotheses fly thickly around when they are discussing it.
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Literature Cited 1. Kotz, J. C.; Treichel, P. M.; Townsend, J. R. Chemistry and Chemical Reactivity; 7th ed.; Thomson Brooks/Cole: Belmont, CA, 2008; Chapter 20. 2. Wagner, R. S.; Strothkamp, R.; Ryan, D. Ideas, Investigation, and Thought: A General Chemistry Lab Manual, 3rd ed.; Whittier Publications: Island Park, NY, 2000; pp 75-84. 3. Corrosion of Aluminum and Aluminum Alloys; Davis, J. R., Ed.; ASM International: Materials Park, OH, 1999; Chapter 2. 4. Annamalai, V.; Hiskey, J. B. Trans. Soc. Min. Eng. AIME 1978, 30, 650. 5. Wei, W. Y.; Lee, C.; Chen, H. J. Langmuir 1994, 10, 1980–1986. 6. Chen, H. J.; Lee, C. Langmuir 1994, 10, 3880–3886. 7. http://www.alcoa.com/mill_products/catalog/pdf/specialties/en/ foil_EN.pdf (accessed 12/30/2008) 8. Corrosion of Aluminum and Aluminum Alloys; Davis, J. R., Ed.; ASM International: Materials Park, OH, 1999; Table A-1, p 257. 9. Margolis, L. A.; Schaeffer, R. W.; Yoder, C. H. J. Chem. Educ. 2001, 78, 235–236. 10. CRC Handbook of Chemistry and Physics, 56th ed.; Weast, R. C., Ed.; CRC Press: Cleveland, OH, 1975-1976. 11. Martell, A. E.; Smith, R. M.; Motekaitis, R. J. NIST Standard Reference Database 46: Critically Selected Stability Constants of Metal Complexes; Standard Reference Data Program, National Institute of Standards and Technology: Gaithersburg, MD, 2004; v. 8.0 12. Kolics, A.; Polkinghorne, J. C.; Thomas, A. E.; Wieckowski, A. Chem. Mater. 1998, 10, 812–824.
Supporting Information Available Instructor notes for the laboratory experiment; student handout. This material is available via the Internet at http://pubs.acs.org.
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