Use of Electrochemical Concentration Cells to Demonstrate the Dimeric Nature of Mercury(1) in Aqueous Media Deepta Bhattacharya and Dennis G. peters' Indiana University, Bloomington, IN 47405
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Most textbooks of eeneral chemistrv cite the fact that m~:rcuryrl~ ex~stsa s the dimeric species HE,.' in aqueous media, although no evidence is quoted t o indicate how this unique behavior is revealed experimentally In advanced texts and reference hooks on inorganic chemistry (1.2 I,one finds that the dimeric nature of m e r c u ~ \Ilhas been established in a number of ways, including
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the demonstration of the diamagnetism of mereury(1)both in solids and in solution the detection of discrete H ~ ions ~ in X-ray ~ + crystallographic studies of mereury(1)salts t h e observation of a strong line in the Raman spectrum of mercury(1) nitrate solutions that is attributable to a mercury-mercury stretching vihration t h e measurement of electrical conductivity of solutions of mercury(1)salts that indicates the presence of ~ g z " t h e treatment of chemical equilibria involving mercury(1) that demands its existence as a dimeric cation Unfortunately, each of these approaches is too sophisticated and time-consuming to serve as the basis for a convenient student laboratory experiment designed to verify the existence of HgzZ+. In this paper we describe a simple and attractive way to demonstrate that HgzZ+isindeed the form of mercury(1) in a n aqueous solution. The method is based on the construction of a series of electrochemical concentration cells contaming mercuryrl! and the measurement nnd analysis of the variation of the cell voltages as a functiun of the ratio of mercury(1) concentrations in each half-cell. I n addition, this 3-h exoeriment satisfies a continuing need in the general chemistry curriculum for good labo&ry experiences in electrochemistrv, and it provides an opportunity to revisit the important topic of preparing and diluting solutions with the aid of volumetric glassware. Methodology In constructing the concentration cells needed for the desired measurements, we have followed closely the simple, elegant, and highly recommended approach described by Craig, Ackermann, and Renfrow (3), in which a minigalvanic cell is assembled from a small glass vial containing 2 M ammonium nitrate solution (serving a s the salt bridge) and from a pair of medicine-dropper tubes (serving a s the two half-cell comoartments and d i ~ o i n ninto the &-bridge solution) plugged with cotton, L i n t k i n g aporooriate electrodes. and filled with solutions of electroactive species. We have used such minigalvanic cells with enorinous success for several years in a laboratory course for freshman honors students, and copies of our experiment are available upon request. Indeed the only alterations we have made in the design of the minigalvanic cell pictured in the paper by Craig and coworkers (3)are the use of the bottom 4.5-cm portion of a n 8-dram vial a s the container for the salt-bridge solution and ensuring that the cotton IAuthor to whom correspondence should be addressed.
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Journal of Chemical Education
plugs extend slightly out of the tips of the medicine-dropper tubes. Reagents and Apparatus
A stock solution of mercury(1) was prepared by dissolution of 2.806 g of commercially available (Mallinckrodt) mercurous nitrate monohydrate in 0.16 M nitric acid in a 100-mL volumetric flask (solution A). A small amount (about 0.5 m1.1 ufelemental mercury was then added to the flask. The nitric acid is essential to dissolve the solid and to prevent the formation of insoluble basic salts of mercury, and the elemental mercurv ensures the stability of merby tencury(1). A1.6 M solution of k t r i c acid was fold volumetric dilution of the concentrated reagent with distilled water (solution B). This solution was used to control the ionic strength of diluted rnercury(1) solutions in one set of measurements. Two mercury-coated gold wire electrodes were fabricated. We began the construction of these electrodes by inserting and centering 7-cm lengths of 0.051-m-diameter gold wire inside of 5-cm lengths of 1.5 to 1.8-mm-0.d. glass capillary tubes (Kimble, 34500 Kimaxdl). One end of each electrode was heated in a Bunsen flame until a tight seal was obtained. During this process some of the gold is melted so that a small sphere forms a t the end of the wire. Then the sealed end of each electrode with its protruding gold wire was immersed in elemental mercury until the exposed wire was completely and evenly coated with mercury, a s determined with the aid of a binocular microscope. This step is crucial to the success of the measurements, and erratic results are obtained if the mercury coating is incomplete or if it appears to consist of islands of mercury. Once the gold wires are properly coated with mercury, the electrodes can be used repeatedly and can be stored between uses in a capped glass vial. The glass capillary tube that sheaths the gold wire prevents the entire wire from being coated with mercury, thereby providing a n electrode that is convenient and safe for use by students. All voltage measurements were made with a Beckman Model 3 1 digital pH-millivolt meter. Procedure Five concentration (minigalvanic) cells should be constructed a s follows.
composition of Solutions for Concentration Cells
Cell Representations
Cell
If mercury(1) exists as Hg+,the concentration cells can be represented as
1 2 3
4
Composition of Solution for Lefl Half-Cell
Composition of Solution for Right Half-Cell
A
A A
10.00 mLof A + 17.44 mL of B, diluted to 100 mL 10.00mLof solution for cell 2 + 17.44 mL of B. diluted to 100 mL 10.00 mL of solution for cell 3 + 17.44 mL of B, diluted to 100 mL
A
where CLand CRdenote the molar (analytical) concentrations of Hg+in the left and right half-cells. If mercury(1)exists as HgzZ+,the concentration cells can be represented as
A
where 0.5C~and 0 . 5 C ~are the molar (analytical) concentrations of Hgz2+in the left and right half-cells and are onehalf of the concentrations based on the assumption that mercury(1) is a monomeric species (Hg+). Half-Reactionsand Overall Cell Reactions
For each of these cells, the molar (analytical) concentration of mercury(1)in the right half-cell has a fxed value of C;for the particular stock solution of mercury(1) specified earlier, C is 0.1 M if mercury(1) is Hg+ and 0.05 M ifmercury(1) is HgZ2+.On the other hand, the concentration of mercury(1) in the leR half-cells is changed systematically from C down to 104C. The table provides detailed information on preparing each of the solutions needed for the concentration cells. Pipets and burets should be used to deliver the appropriate volumes of solutions into 100-mL volumetric flasks. After the solutions in the table have been made, Pasteurtype disposable transfer pipets are used to add small portions of the desired mercury(1) solutions to two cottonplugged medicine-dropper tubes. A mercury-coated electrode is inserted into each half-cell compartment. The minigalvanic cell is assembled as described by Craig and coworkers (31, and the voltage of that cell is measured. For each of the four remaining concentration cells, we recommedicine-dro~~er mend that clean cotton-~lueeed . -. tubes he prepared and then filled with appropriate solutions listed in the table. The mercury-coated electrodes should be rinsed with distilled water and gently patted dry before being inserted into their respective half-cells. If the two stock solutiom (solutions A and B) specified in the table are available in advance, a pair of students can easily complete the necessary dilutions and the voltage measurements for all five concentration cells in a 3-h period. After the set of voltage measurements is obtained, the data can be treated as discussed later in this paper.
For the first concentration cell, we can write the following half-reactions and overall cell reaction. left half-cell (anode)
right half-cell (cathode)
overall cell reaction
For the second concentration cell we have the following half-reactions and overall cell reaction. leR half-cell (anode)
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Results and Discussion
Let us begin by proposing that rnercury(1) might exist in either of two states in an aqueous medium, that is, monomeric Hg+ or dimeric Hgz2+.The suggestion that mercury(1) could be monomeric is consistent with the nature of every other common metal cation with which the average student is familiar, whereas the idea that mercury(1) is dimeric affirms the statements found in textbooks. Then let us consider how to
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represent the concentration cells involving each form of mercury(1) write the half-reactions and overall cell reactions that take place in these concentration cells formulate the Nernst equation for each scenario ultimately show that mercury(1)is dimeric
right half-cell (cathode)
overall cell reaction H g p ( 0 . 5 ~2 ~) ~ g p ( 0 . 5 ~ ~ ) Nernst Equations for the Concentration Cells
On the basis of the overall cell reaction and on the presumption that the temperature is close to 25 OC (298 K),we can formulate the Nernst equation for the first concentration cell as
whereas the Nernst equation for the second cell is
]
0.059 Vmole log ~=u=0-( 0.bCR
2 mol C
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cally that Hgz2+is the state of mercury(1) in an aqueous solution. Truthfully, the experimental measurements in the figure demonstrate only that the charge on the mercury(1) species is +2; if the experiment had been done with solutions of mercury(I1)(i.e., HgZt)the results would have been identical with those shown in the figure. Thus, it is important to remind the student that the ability to formulate mercury(1)as Hgz2+depends on our knowing that the oxidation state of mercury is +1 in the original compound (mercurous nitrate monohydrate) from which the various solutions are pre ared. Finally, the student should note that dimeric Hg2I: shows no tendency to dissociate into monomeric Hg'. Shown in the fimue is one set of voltaee measurements (those symbolized by circles) acquired
