In the Laboratory
W
Reproportionation of Copper(I)
Jan Malyszko* and Maria Kaczor Institute of Chemistry, Saint Cross Academy in Kielce, 5 Checinska St., PL-25020 Kielce, Poland; *
[email protected] Cu(II) + Cu
2Cu(I)
(1)
is expressed by 2
Kr =
C Cu(I)
CCu(II)
(2)
The equilibrium constant, or specifically the reproportionation constant, Kr, can be determined either by a direct measurement of the equilibrium concentrations of Cu(I) and Cu(II) or from the appropriate formal potentials of single redox couples: Cu(I)|Cu(II), Cu|Cu(II), and Cu|Cu(I). The equilibrium shown in eq 1 is very sensitive to the nature and concentration of the background electrolyte. In the presence of noncomplexing salts, for example sulfates or perchlorates, copper(I) is very unstable and tends to disproportionate. Under standard conditions, the formal potentials (identical to the standard potentials) for the Cu(I)|Cu(II), Cu|Cu(II), and Cu|Cu(I) pairs are 153, 327, and 521 mV versus standard hydrogen electrode, SHE, respectively (1). The equilibrium constant for the reaction in eq 1 is reported to be 6 × 10᎑7 mol L᎑1 at 25 ⬚C (2). Consequently, only low equilibrium concentrations of Cu(I) exist under these conditions. The relative stabilities of Cu(I) and Cu(II) depend strongly on the nature of complexing anions or other ligands present in solution. The formation of complexes shifts the formal potentials of the M|Mn+ couples in the negative direction. Thus, the relative stability of +1 and +2 valency states of copper may be altered by adding a ligand that forms strong complexes with Cu(I) ions but interacts weakly with Cu(II) ions. Numerous ligands form stable complexes with Cu(I) in aqueous solutions, for example ammonia, halides, and pseudohalides. Cu(I) and chloride ions form a series of strong complexes CuClnn−1, with n ≤ 4 (3). For instance, the dominant species in a 1 M chloride solution are CuCl32− and CuCl43−. Conversely, Cu(II) and chloride ions form only very weak complexes (4). (It should be emphasized that chloride– Cu(II) complexes are green whereas chloride–Cu(I) complexes are in general colorless.) Consequently, in the presence of chloride ions, the Cu(I) state is considerably stabilized relative to Cu(II). Formal potentials of 449, 314, and 190 mV versus SHE were determined for the Cu(I)|Cu(II), Cu|Cu(II), and Cu|Cu(I) couples, respectively, in 1 M KCl solution (5). From these data, the reproportionation constant was calculated to be 2.4 × 104 mol L᎑1 at 25 ⬚C. The change in ordering of the formal potentials for copper going from standard conditions (1) to 1 M KCl solution (5) is shown in Figure 1. 1048
Reproportionation of copper(I) in the chloride medium can be considered as a particular kind of metal corrosion, also called spontaneous metal dissolution. It should be emphasized that such a corrosion process occurs when copper is immersed into an electrolyte containing its own ion of higher oxidation state, Cu(II), which acts as the oxidant. Corrosion processes of metals are of great importance in the environment. It is essential that students become more aware of these topics. However, easy to perform demonstrations are rarely described in the literature. In this article we report an easy and attractive set of laboratory experiments illustrating the dramatic change in the redox and, consequently, corrosion properties of copper. The experiment was designed to show spontaneous dissolution of copper metal in solution. For comparison, an analogous experiment should be performed in a solution containing dilute H2SO4 as the background electrolyte. The SO42− ion is not able to complex copper(I) ion. In this solution, the redox properties of the system under study are comparable to those characteristic of the standard solution (Figure 1). From the comparison of the two experiments, some questions naturally arise. The instructor should explain why copper(II) ion has practically no oxidizing power against metallic copper in sulfuric acid solution, whereas it is a powerful oxidizing agent in the chloride medium. These phenomena can be rationalized by considering the schematic set of formal potentials given in Figure 1 for the two different media. The exercise provides the opportunity to discuss basic chemical and electrochemical concepts such as the potentials of redox pairs and the direction of the redox process. Further discussion is recommended to help the students understand the relationships between the thermodynamic parameters considered (redox potentials and reproportionation constants) and the mechanism of electrochemical reduction of Cu(II) ions in different media.
500
E °' vs SHE / mV
Numerous metals are able to form compounds with more than one oxidation state. Compounds or ions at intermediate oxidation states can undergo auto-oxidation–reduction reactions in which the same substance is both oxidized and reduced. Copper is a typical example of an element that forms compounds at +1 and +2 oxidation states. The equilibrium constant of the following reaction
Cu(I)/Cu(II)
400
300
Cu/Cu(II)
200
Cu/Cu(I) 100
standard solution
1 M KCl
Figure 1. Formal potentials for the Cu(I)|Cu(II), Cu|Cu(II), and Cu|Cu(I) redox pairs in aqueous solutions under standard conditions versus 1 M KCl.
Journal of Chemical Education • Vol. 80 No. 9 September 2003 • JChemEd.chem.wisc.edu
In the Laboratory
Experimental Procedure Copper foil, 0.1-mm thick, was cut into 10-mm × 50mm rectangles. The foil was washed with nitric acid, rinsed with water and alcohol, and air dried. It was then weighed and suspended above a magnetic stirring bar in a volume of solution in a 100-mL beaker. Two reactant solutions were used to perform the exercise: (A) 3.0 M NaCl, 0.05 M HCl, and 0.5 M CuCl2; (B) 0.1 M H2SO4 and 0.5 M CuSO4. The quantity of copper dissolved was obtained by weighing the foil before and every two minutes during the experiment. The foil was removed from the solution, washed with distilled water, dried gently between sheets of filter paper, and weighed with an analytical balance of 0.1 mg precision. The stirring rate must be kept constant during the experiment. Results Direct evidence for corrosion of the copper foil in solution A was obtained by visual observation. Small pits were formed on the metal surface after a few minutes. The pits continued to develop over the surface so that the specimen became gradually smaller in size, until within 30 minutes the copper foil was entirely consumed. An analogous experiment was carried out with another copper sample immersed into solution B for 30 minutes. By weighing the sample before and after the experiment, we observed no mass loss. This result confirms that the spontaneous dissolution of copper did not occur under these conditions. The results of this experiment in solution A are presented in Figure 2. The decrease in mass of the copper foil in solution A is initially proportional to time. Deviation from linearity at later times is due to a gradual diminution of the electrode surface during the dissolution. The reaction progress can also be measured by the potential of the Cu(I)|Cu(II) couple at an inert electrode. For this purpose, a Pt wire and a saturated calomel electrode (SCE) were placed into the solution. The dependence of the potential on the logarithm of reaction time is shown in Fig-
ure 3. The plot is linear with a slope of ᎑60 mV兾log-unit. The result is not surprising because the Cu(II) concentration changes only slightly during the experiment. The original green solution turns gradually deep brown during the reaction. This change is associated with the formation of a Cu(II)–Cu(I) mixed-valence species (6). Laboratories that possess a rotating disk electrode (RDE) can easily conduct this experiment under conditions of high reproducibility. The theoretical background of this technique as well as the possibility of its application to corrosion phenomena have been published in this Journal (7, 8). Hazards No chemicals or procedures used by students present any significant hazards. The stock solutions of acids used in this experiment (1 M) are irritating and thus require proper eye and skin protection. Conclusions This experiment illustrates that the formation of complexes influences the relative stabilities of oxidation states. During the experiment, students observe the progress of the reaction in the chloride medium by visual examination of the metallic copper and by monitoring its weight loss. The decrease in the potential of the Cu(I)|Cu(II) redox pair and a distinct variation in the solution color give further evidence for the generation of Cu(I) ions. The experiment described was successfully carried out during a course in electroanalytical chemistry for three consecutive years and was well received by the students. Some students were surprised by the results, which indicate a significant influence of the medium on the properties of redox system. If the test solutions are available in advance, a pair of students can easily complete the necessary gravimetric and potentiometric measurements in one hour. This type of experiment can also be used in a general chemistry laboratory to familiarize the students with different kinds of redox re-
200
440
430
E vs SHE / mV
−∆m / mg
150
100
50
420
410
400
390 0 0
5
10
15
20
25
t / min Figure 2. Mass loss of Cu foil during spontaneous dissolution in the aqueous solution containing 0.5 M CuCl2, 3.0 M NaCl, and 0.05 M HCl.
2
3
4
5
10
15
20
t / min Figure 3. Changes in the electrode potential of Cu(I)|Cu(II) during the spontaneous dissolution of copper in an aqueous solution containing 0.5 M CuCl2, 3.0 M NaCl, and 0.05 M HCl.
JChemEd.chem.wisc.edu • Vol. 80 No. 9 September 2003 • Journal of Chemical Education
1049
In the Laboratory
actions. We suggest preceding the experiment by an explanation of the theoretical background of the reproportionation and disproportionation reactions so that the students can properly understand the results. Acknowledgments We would like to thank Zbigniew Galus and Marek K. Kalinowski, from the Faculty of Chemistry at the University of Warsaw, for suggestions that led to an improved manuscript. W
Supplemental Material
The theoretical background related to this experiment and detailed notes for the students are available in this issue of JCE Online.
1050
Literature Cited 1. Bertocci, U.; Wagman, D. D. In Standard Potentials in Aqueous Solutions; Bard, A. J., Parsons, R., Jordan, J., Eds.; Marcel Dekker: New York, 1985; Chapter 11. 2. Burgess, J. Metal Ions in Solution; Ellis Horwood: Chichester, England, 1978; Chapter 8. 3. Ciavatta, L.; Juliano, M. Ann. Chim. (Rome) 1998, 88, 71. 4. Arnek, R.; Puigdomenech, I.; Valiente, M. Acta Chem. Scand. A 1982, 36, 15. 5. Malyszko, J. Habilitation Thesis; Pedagogical University in Siedlce, 1978. 6. McConnell, H.; Davidson N. J. Am. Chem. Soc. 1950, 72, 3168. 7. Drok, K. J.; Ritchie, I. M.; Power, G. P. J. Chem. Educ. 1998, 75, 1145. 8. Nikolic, J.; Exposito, E.; Iniesta, J.; Gonzales Garcia, J.; Montiel, V. J. Chem. Educ. 2000, 77, 1191.
Journal of Chemical Education • Vol. 80 No. 9 September 2003 • JChemEd.chem.wisc.edu
