Complementary Use of Electrochemistry and Synthetic Redox

Laboratory Experiment pubs.acs.org/jchemeduc

Cite This: J. Chem. Educ. XXXX, XXX, XXX−XXX

Complementary Use of Electrochemistry and Synthetic Redox Chemistry in the Oxidation of Decamethylferrocene: An Integrated Advanced Laboratory Experiment William E. Geiger* Department of Chemistry, University of Vermont, Burlington, Vermont 05405, United States

Downloaded via TUFTS UNIV on July 13, 2018 at 17:30:36 (UTC). See https://pubs.acs.org/sharingguidelines for options on how to legitimately share published articles.

S Supporting Information *

ABSTRACT: An integrated pair of experiments has been developed to introduce undergraduate students to the ways in which electrochemistry can complement conventional synthetic chemistry involving redox reactions. Students first use cyclic voltammetry and linear scan voltammetry to characterize and quantify the products of the reaction between decamethylferrocene and the 4-nitrophenyldiazonium ion. In the following lab period, they synthesize the decamethylferrocenium ion as its tetrafluoroborate salt.

KEYWORDS: Upper-Division Undergraduate, Analytical Chemistry, Inorganic Chemistry, Organometallics, Laboratory Instruction, Instrumental Methods, Oxidation State, Reactions, Electrochemistry

■

OVERVIEW

Although electron-transfer reactions are of central importance in chemistry and biochemistry, they are frequently underemphasized in the undergraduate laboratory curriculum. Here we describe an experimental sequence which integrates the two most commonly employed experimental methods for studying electron-transfer reactions: electrochemistry and synthetic redox chemistry. Focused on the decamethylferrocene/ decamethylferrocenium ion redox couple, 10/+, the experiments have been performed for a number of years by third- or fourthyear undergraduates in an integrated analytical/inorganic laboratory at the University of Vermont. After first characterizing the electrochemical oxidation of 1 by cyclic voltammetry (CV), students then use CV and linear scan voltammetry (LSV) to identify and quantify the products of the reaction between 1 and the one-electron oxidant, 4-nitrophenyldiazonium ion ([4-NO2(C6H4)N2]+), eq 1, in an electrochemical cell. In a second lab period, the chemical oxidation is carried out in a simple flask and the tetrafluoroborate salt of 1+ is isolated. A total of six laboratory hours, over 2 days, is allowed for the combined procedures.

Upon completion of these experiments, a student should be able to (i) perform and evaluate a cyclic voltammetry experiment to determine the redox potential of a compound; (ii) choose, from a list of formal redox potentials, a sufficiently powerful reagent to perform a chemical oxidation of the compound; (iii) be aware of the difference between the thermodynamic and kinetic aspects of reversible vs irreversible redox reactions; and (iv) appreciate how cyclic voltammetry may be used for qualitative and quantitative analysis of the products of chemical reactions.

■

Chemically Reversible Redox Reactions

These experiments encourage a discussion of the fundamental differences between electrochemical and chemical approaches to carrying out the one-electron oxidation of a compound, A, to A+ (1 to 1+ in the present case). The electrochemical method is a heterogeneous process in which an electron is

1 + [4‐NO2 (C6H4)N2][BF4 ] + CH3CN → 1[BF4 ] + C6H5NO2 + N2 + “CH 2CN” © XXXX American Chemical Society and Division of Chemical Education, Inc.

CONCEPTUAL BACKGROUND

Received: January 19, 2018 Revised: June 21, 2018

(1) A

DOI: 10.1021/acs.jchemed.8b00021 J. Chem. Educ. XXXX, XXX, XXX−XXX

Journal of Chemical Education

Laboratory Experiment

donated to a solid “working” electrode from the molecule A in solution. The reaction proceeds in favor of A+ when the applied electrode potential, Eappl, is more positive than the formal potential, E°′(A0/+), of the A0/+ couple (eq 2). The chemical redox method involves a homogeneous reaction between two species in solution. It utilizes an oxidizing agent, Ox ([4-NO2(C6H4)N2]+ in the present case), which accepts an electron from A, giving A+ as well as the reduced form of the oxidizing agent, Red (eq 3). In the homogeneous reaction of eq 3, the difference in formal potentials of the two redox pairs, ΔE°′ = E°′(Ox/Red) − E°′(A0/+) (eq 4), may be considered to be the “cell” potential for the overall redox reaction. If the two couples are both reversible (i.e., Nernstian), E°′(Ox/Red) must be greater than E°′(A0/+) to push the reaction in favor of A+. A ⥂ A+ + e‐

when Eappl > E°′(A0/ +)

reaction medium. The initial CV measurements on A are used to characterize the electron-transfer stoichiometry of the oxidation of A (present case, one electron), and to determine the formal potential of the redox couple, E°′(A0/+). In addition, the CV scans allow one to probe the chemical stability (i.e., the longevity) of the oxidized product, A+, over an approximate 10 s time scale. A number of CV experiments of this type have been described in this Journal.2−8 Since CV scans are carried out using rather small working electrodes (typically, disks of 1−3 mm radius), they are used almost exclusively for analytical purposes and cause only minute changes in the concentration of A in the bulk of solution. Of course, the entire solution of A could be converted to A+ in a traditional “bulk electrolysis”9 by using a working electrode of much larger area. However, carrying out the bulk conversion to A+ with a homogeneous chemical oxidant is usually less technically challenging, in addition to allowing for shorter reaction times. One disadvantage of the homogeneous redox approach is that the second product generated in the reaction [Red in eq 3] must be accounted for. In the present experiment, the ultimate second product is nitrobenzene, which is identified through CV scans of the postreaction solution.

(2)

A + Ox F A+ + Red

(3)

ΔE°′ = E°′(Ox/Red) − E°′(A0/ +)

(4)

Equation 5, which is based on the Nernst equation, shows how ΔE°′ is related to the ratio of product, A+, to the starting material, A, in the equilibrated redox reaction. Ratios of 10:1 and 100:1 for [A]+:[A] are obtained for ΔE°′ values of 0.118 and 0.236 V, respectively, when the temperature is in K and ΔE°′ is in volts. | l [A+] o o logm } o o = 8.47ΔE°′ o (5) n [A] o ~ This calculation is the basis of the rule of thumb that, for a “complete” redox reaction, the oxidant should have a formal potential that is at least 200 mV positive of that of the compound, A, being oxidized.

The Two Redox Pairs: Decamethylferrocene and Nitrophenyldiazonium Ion

Note that the potentials in this paper are given vs the experimental reference electrode, Ag/AgCl in CH3CN/0.1 M [NBu4][PF6]. Literature references for these redox reactions have frequently reported the potentials vs other reference couples, most commonly the ferrocene/ferrocenium couple or the Hg(0)/Hg(I) couple of the saturated calomel electrode (SCE). Measurements in our laboratory show that the potentials measured vs Ag/AgCl in CH3CN/0.1 M [NBu4][PF6] may be converted to values vs ferrocene or the SCE by subtraction of 0.48 or 0.08 V, respectively. A discussion of some of the issues involved in converting reference potentials in nonaqueous media is available.1 The overall chemical reaction between decamethylferrocene and the nitrophenyldiazonium ion was given in eq 1, where the acetonitrile solvent is assumed to play the role of H atom donor in the reaction sequence. The oxidative half-reaction is the simple one-electron oxidation of decamethylferrocene (eq 6), which has a formal potential of −0.03 V vs Ag/AgCl under our conditions.

Irreversible Redox Reactions

Consider the more generic set of reactions in Scheme 1. Here, in addition to the redox reaction between A and Ox, we Scheme 1. Possible “Follow-Up” Reactions of the Initial Redox Products, A+ and Red, Produced from the OneElectron-Transfer Reaction of A with an Oxidant, Ox

1 F 1+ + e−

(6)

E°′ = − 0.03 V

The reductive half-reaction is more complex. It involves the oneelectron reduction of the nitrophenyldiazonium ion (eq 7) at a cathodic peak potential, Ep, of 0.08 V,10−12 accompanied by very rapid loss of dinitrogen to give the nitrophenyl radical (eq 8). The reaction in eq 8 corresponds to the follow-up reaction for Red going to Red′ in Scheme 1. The nitrophenyl radical then abstracts a hydrogen atom (likely from solvent), giving nitrobenzene as the final reduction product (eq 9). Thus, the oxidative half-reaction is reversible, whereas the reductive halfreaction is irreversible.

consider the possibility that either or both of the initial products, A+ and Red, might be kinetically unstable, undergoing a so-called “follow-up” reaction. In such a case, the progress of the redox reaction is influenced by a kinetic factor which becomes important if either kc(A) or kc(Red) is significant on the experimental time scale. In the reaction between 1 and [4-NO2(C6H4)N2]+, the follow-up reaction Red → Red′ indeed comes into play. Electrochemistry as Complement to Homogeneous Redox Chemistry

A broad discussion of the advantages and disadvantages of electrochemical vs chemical approaches to synthetic redox chemistry is available.1 Here, we emphasize how the two approaches can complement each other in developing a strategy for the chemical preparation of A+ and in analyzing the B

[4 ‐NO2 (C6H4)N2]+ + e− → 4 ‐NO2 (C6H4)N•2

(7)

4 ‐NO2 (C6H4)N2 → C6H4NO•2 + N2

(8)

C6H4NO•2 + CH3CN → C6H5NO2 + “CH 2CN”

(9)




Referring back to Scheme 1, the irreversibility of the reduction raises the likelihood that the electron-transfer reaction will proceed beyond the point predicted by employing the nominal potential of 0.08 V (eq 7). Noting that the difference between the cathodic and anodic potentials for the two half-cell reactions (110 mV) is less than the “rule of thumb” value for a “complete” redox reaction (≥200 mV, see above), we ask the student to compute the maximum yield of 1+ that would be expected were the reduction of [4-NO2(C6H4)N2]+ to be reversible with an ΔE°′ value of 0.08 V. The maximum yield of 88% for 1+, predicted by eq 5, may later be compared with that determined experimentally. More extensive discussions of the role of chemical irreversibility in promoting electron-transfer reactions are available.1,13

Figure 1. Cyclic voltammogram of 6 × 10−4 M decamethylferrocene (1) in CH3CN/0.1 M [NBu4][PF6] at 2 mm glassy carbon electrode (gce), scan rate 0.2 V s−1.

■

HAZARDS Safety goggles and gloves should be worn during the experiments, which should be carried out in a fume hood. The solvent vapors from the electrochemistry experiment should also be properly vented. Since dichloromethane, acetonitrile, and nitrobenzene are all toxic permeators, special attention should be paid to avoiding contact of these liquids with the skin. Dichloromethane is a suspected human carcinogen.

■

of different experimental options are available for obtaining the desired “S”-shaped quasi-steady-state LSVs. Whereas rotating disk, bead, or wire electrodes are among the most rigorous tools for obtaining steady-state scans,15−17 other procedures may be employed if rotating electrode equipment is not available. These include the possibility of using an ultramicroelectrode disk as the working electrode, as has been described previously in this Journal.5 A simpler, albeit less rigorous, strategy is to obtain very slow LSV scans (≤2 mV s−1) at the same 2 mm glassy carbon electrode that had been employed for the CV scans. As long as the convective forces in the unstirred electrochemical cell remain essentially unchanged, these relatively crude LSV scans can be used effectively. The scans shown in Figure 2 are typical of those

RESULTS AND DISCUSSION

Detailed experimental procedures are given in the Supporting Information (SI). The experiments have been carried out by students either individually or in groups of two. On the first day, the student does CV and LSV scans of a solution of 1, adds an equivalent of the [4-NO2(C6H4)N2]+ oxidant to the solution, observes the color change to the green of 1+, and obtains another sequence of CV and LSV scans, along with a visible absorption spectrum of the reaction solution. Part I.2 of the Experimental Section of the Supporting Information describes the experimental cell setup and details how the diagnostics of cyclic voltammetry may be used to examine fundamental aspects of the oxidation of decamethylferrocene. Students first make up a stock 50 mL solution of about 0.6 mM 1 in acetonitrile that contains 0.1 M [NBu4][PF6] as the supporting electrolyte. Half of this is transferred to an electrochemical cell and degassed with flowing nitrogen or argon.14 Students next carry out a series of CV scans at different scan rates (e.g., 0.1 to 1 V s−1) that allow them to determine the E1/2 (here, taken to be identical to E°′) of the 10/+ couple and whether or not it involves a diffusioncontrolled, one-electron, process that is chemically and electrochemically reversible (i.e., “quasi-Nernstian”). To this point, the experiment is simply following procedures described in detail in ref 2 and other earlier papers.3−8 An E1/2 value of −0.03 V vs Ag/AgCl is determined for the 10/+ couple under these conditions (see Figure 1). A key addition to the present experiment is that the student must also obtain an LSV in the 1/1+ potential region under “quasi-steady-state” (i.e., apporximately time-independent) conditions. This technique is superior to cyclic voltammetry in allowing one to determine the exact solution quantities of oxidized vs reduced forms of a reversible redox pair, 1+ vs 1, respectively, in the present case. Thus, the student obtains LSV scans before and after addition of the diazonium ion oxidant to the solution, using the results of the two scans to determine the extent of the reaction, i.e. the “percent yield” of 1+. A number

Figure 2. Linear scan voltammograms at 2 mm gce in region of decamthylferrocene/decamethylferrocenium couple before (solid line) and after (dashed line) addition of 0.95 equiv of [4NO2(C6H4)N2][BF4] to solution, scan rate 2 mV s−1. Anodic currents are negative.

obtained using this technique. Although there are small variations (“noise”) in the currents owing to the uneven hydrodynamics of the nominally motion-quiet cell, the curves nevertheless give the desired information about the progress of the reaction. First, one sees that the sum of the anodic (negative) and cathodic (positive) currents after addition of 0.95 equiv of [4-NO2(C6H4)N2]+ to the solution essentially matches that of the anodic current prior to its addition. This indicates that there has been essentially no loss of the 1/1+ redox pair in the reaction.18 Second, the fact that, after the reaction, the anodic current (for oxidation of 1) is only about 10% of the cathodic current (for reduction of 1+) means that the oxidation reaction has proceeded in very high yield.19 Additional information is then obtained in the CV scan20 taken after addition of the oxidant (Figure 3, solid line). In addition to showing the reversible reduction of 1+ at E1/2 = C




M−1 cm−1)22−26 with shoulders at higher energies (see Figure S1). The chemical synthesis of 1[BF4] (SI Experimental Section, part II), performed on the second day, is straightforward. Although column chromatography might be employed to separate the decamethylferrocenium salt from the reaction mixture, we have obtained good elemental analyses on the product using the evaporation/wash/decantation sequence described in the Experimental Section. Students characterize the product by optical and IR spectroscopies. The former confirms the identity of 1+ (by comparison with the spectrum obtained for 1+ on day one) but does not identify the counterion, [BF4]−. The latter may be confirmed by comparison of the IR spectrum of the solid product (Figure 4) with that of [NBu4][BF4] (Supporting Information), which was provided to the student. The presence of the tetrafluoroborate anion in the isolated solid is indicated by the strong absorption at 1054 cm−1.27 The absorptions near 2925 cm−1 arise from the C−H stretches of the methyl groups in 1+, and those at 1476 cm−1 and ca. 1387 cm−1 have been attributed to ring motions of the cyclopentadienyl ligand in 1+.28

Figure 3. Cyclic voltammetry scans (positive to negative potentials) (2 mm gce, 0.2 V s−1) of a solution after the reaction of 1 with 0.98 equiv of [4-NO2(C6H4)N2][BF4], showing the reduction of 1+ near 0 V and the reduction of nitrobenzene near −1.1 V. The CV of the dashed line was taken after addition of 1 μL of nitrobenzene to the reaction mixture.

−0.03 V, a second reversible reduction of approximately equal current is observed. The latter arises from the product of the reduction of [4-NO2(C6H4)N2]+ by 1. The experimental writeup given to the students tells them that this reduction product will be either nitrobenzene or 1,1′-dinitrobiphenyl, the E1/2 potentials of these two possible products being close to −1.1 V10−12 and −0.6 V,21 respectively. After confirming from the CV that the reaction product is very likely to be nitrobenzene (measured E1/2 = −1.09 V), the student adds a known amount of this compound (1 μL = 1.2 mg) to the solution. Replication of the CV scan shows an increase in the current of the nitrobenzene wave (Figure 3, dashed line), immediately confirming the identity of the reduction product. Furthermore, the student is instructed to use these data to quantify, in the lab write-up, the amount of nitrobenzene produced in the reaction. The combination of LSV and CV scans taken before and after the chemical reaction has thus determined both the identities and quantities of the two reaction products, namely, the decamethylferrocenium ion and nitrobenzene. The characterization of 1+ in the reaction solution is also carried out by optical spectroscopy recorded for a sample taken from the electrochemical cell. The decamethylferrocenium ion has a major absorption band at λmax = 778 nm (ε = 5.7 × 102

■

CONCLUSIONS Two of the most valuable instructional aspects of this experiment are (i) introduction of the idea that the chemical reversibility of a redox couple plays a significant role in determining its function in a redox reaction and (ii) demonstration of how cyclic voltammetry and other electrochemical and spectroscopic methods can be used to design a redox reaction and to monitor the reaction products. Regarding point i, it is useful to keep in mind that students have previously used thermodynamics (via standard potentials and the Nernst equation) almost exclusively to calculate quantities like cell potentials and equilibrium constants of redox reactions (e.g., reduction of Cu2+ by Zn). The present experiment brings a kinetic factor into play, namely, in the fact that the chemical instability of the one-electron product of the nitrophenyldiazonium ion reduction pushes the redox reaction toward completion. The chemical irreversibility of the reduction process make the nitrophenyldiazonium ion a

Figure 4. IR spectrum of 1[BF4] in KBr. D




stronger oxidant than would be the case if its reduction potential (Ep = 0.08 V) were actually a standard (reversible) potential. It is important for the student to write a balanced equation for the overall redox reaction, being cognizant of the fact that eq 10, which has often appeared in student reports, is not balanced in terms of the number of hydrogen atoms. The question is then raised: Where did the extra hydrogen atom (not a proton) come from? Here, one can only speculate, although solvent (see eq 9) and trace water are the most likely candidates as the H atom donor.

■

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. ORCID

William E. Geiger: 0000-0002-9152-8550 Notes

The author declares no competing financial interest.

■

1 + [4‐NO2 (C6H4)N2][BF4 ] → 1[BF4 ] + C6H5NO2 + N2

ACKNOWLEDGMENTS The author thanks Neil G. Connelly for helpful discussions and advice, and the National Science Foundation (most recently through CHE-1565541) for its general support of work in the area of organometallic electrochemistry.

(10)

The analytical aspects of the experiment (point ii above) offer an instructional opportunity to have the student consider the more general question of how one can characterize paramagnetic species that are not easily identified by NMR spectroscopy.29 Besides the voltammetry and optical and IR spectroscopy techniques included in this experiment, the instructor might consider expanding the characterization of 1[BF4] to include solution conductivity measurements to confirm that 1[BF4] is a 1:1 salt. Another option is the use the Evans-NMR method to determine the magnetic susceptibility of the decamethylferrocenium ion. These additional characterizations have been done in our course during an extra lab period. Regarding solution conductivity, interested students have performed that measurement on a 1 mM solution of 1[BF4] in acetonitrile, comparing the result with identical measurements on simple salts such as [NBu4][BF4] and neutral compounds such as decamethylferrocene.30 The degree to which students have grasped the intellectual goals of this and other experiments during the semester is assessed and graded through one of three types of required reports: a traditional lab report of approximately five pages (typically, three per semester), a 20 min ACS-type oral report to the class, or a “manuscript”-level report of 15−20 pages. The latter two options are particularly effective in assessing the depth of the student’s understanding of the material. The manuscript report is written in the format of a full paper that might be submitted to an ACS journal. The student is asked to provide the kind of literature background and detailed explanation that would be expected for a manuscript at such a level. The paper is critically evaluated by the instructor and returned to the student for a second draft, wherein he or she is expected to correct, broaden, and deepen the points raised by the instructor. In some cases, a third draft has been allowed. In this way, deficiencies in understanding the main points of the experiments can be identified and singled out for additional work and study. A grade for the manuscript report is based entirely on the quality of the final version.

■

Student handout (PDF, DOCX)

■

REFERENCES

(1) Connelly, N. G.; Geiger, W. E. Chemical Redox Agents for Organometallic Electrochemistry. Chem. Rev. 1996, 96, 877−910. (2) Van Benschoten, J. J.; Lewis, J. Y.; Heineman, W. R.; Roston, D. A.; Kissinger, P. T. Cyclic Voltammetry Experiment. J. Chem. Educ. 1983, 60, 772−776. (3) Carriedo, G. A. The Use of Cyclic Voltammetry in the Study of the Chemistry of Metal-Carbonyls. J. Chem. Educ. 1988, 65, 1020− 1022. (4) Gomez, M. E.; Kaifer, A. E. Voltammetric Behavior of a Ferrocene Derivative. J. Chem. Educ. 1992, 69, 502−505. (5) Ching, S.; Dudek, R.; Tabet, E. Cyclic Voltammetry with Ultramicroelectrodes. J. Chem. Educ. 1994, 71, 602−605. (6) Wheeler, J. F.; Wheeler, S. K.; Wright, L. L. Electrochemical Measurements in the Undergraduate Curriculum. J. Chem. Educ. 1997, 74, 72−73. (7) Brown, J. H. Development and Use of Cyclic Voltammetry Simulator to Introduce Undergraduate Students to Electrochemical Simulations. J. Chem. Educ. 2015, 92, 1490−1496. (8) Brown, J. H. Analysis of Two Redox Couples in a Series: An Expanded Experiment to Introduce Undergraduate Students to Cyclic Voltammetry and Electrochemical Simulations. J. Chem. Educ. 2016, 93, 1326−1329. (9) Bard, A. J.; Faulkner, L. R. Electrochemical Methods, 2nd ed.; John Wiley & Sons: New York, 2001; pp 417−430. (10) Elofson, R. M.; Gadallah, F. F. Substituent Effects in the Polarography of Aromatic Diazonium Salts. J. Org. Chem. 1969, 34, 854−857. (11) Delamar, M.; Hitmi, R.; Pinson, J.; Savéant, J. M. Covalent Modification of Carbon Surfaces by Grafting of Functionalized Aryl Radicals Produced From Electrochemical Reduction of Diazonium Salts. J. Am. Chem. Soc. 1992, 114, 5883−5884. (12) Delamar, M.; Désarmot, G.; Fagebaume, O.; Hitmi, R.; Pinsonc, J.; Savéant, J. M. Modification of Carbon Fiber Surfaces by Electrochemical Reduction of Aryl Diazonium Salts: Application to Carbon Epoxy Composites. Carbon 1997, 35, 801−807 Note that the Ep value of 0.08 V given in the text is the average of values reported in the references.10−12. (13) Bard, A. J.; Faulkner, L. R. Electrochemical Methods, 2nd ed.; John Wiley & Sons: New York, 2001; pp 36−39. (14) An alternative approach would be to add the sample of decamethylferrocene to the electrochemical cell only after the student has obtained “background” cyclic voltammograms on the degassed electrolyte solution (15) Rieger, P. H. Electrochemistry, 2nd ed.; Chapman & Hall: New York, 1994; pp 207−214. (16) Hammerich, O.; Speiser, B. In Organic Electrochemistry, 5th ed.; Hammerich, O., Speiser, B., Eds.; CRC Press: Boca Raton, 2016; pp 150−154.

ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available on the ACS Publications website at DOI: 10.1021/acs.jchemed.8b00021. Experimental Section including description of materials used and detailed laboratory procedures for day 1 (electrochemistry and redox reaction in electrochemical cell) and day 2 (synthesis of 1[BF4]), instructor notes, optical spectrum of 1[BF4], and IR spectrum of [NBu4][BF4] in KBr (PDF, DOCX) E




(17) Adams, R. N. Electrochemistry at Solid Electrodes; Marcel Dekker, Inc.: New York, 1969; pp 80−86. (18) This assumes that the diffusion coefficiences of 1 and 1+ are not significantly different. (19) Owing to the fact that the presence of unreacted [4NO2(C6H4)N2][BF4]+ in solution causes adsorption and electrode fouling at glassy carbon electrodes, the experimental procedure calls for a slightly substoichiometric amount of oxidant to be added to the solution of 1 (see Supporting Information). The student must take this into account when computing the percent yield of 1+ in the reaction. (20) In order to begin the CV scan at the benign potential, it is initiated at a potential that is positive of the E1/2 of decamethylferrocene (21) Pilard, J.-F.; Postec, S.; Simonet, J.; Mousset, G. Formation and Reactivity of Arylradicals from Cathodic Cleavage of Halonitrobenzenes. Electrochim. Acta 1998, 43, 3135−3140. (22) This extinction coefficient is the average of the three values given in refs 21−23. (23) Gale, R. J.; Singh; Job, R. Metallocene Electrochemistry. I. Evidence for Electronic Stabilization with Alkylated Cyclopentadiene: Electrochemical Synthesis of Decamethylferricinium Dication. J. Organomet. Chem. 1980, 199, C44−C46. (24) Frey, J. E.; Du Pont, L. E.; Puckett, J. J. Formation Constants of Radical-Ion Pairs and Charge-Transfer Complexes of Tetracyanoethylene with Group 8 Metallocenes. J. Org. Chem. 1994, 59, 5386− 5392. (25) Nelsen, S. F.; Chen, L.-J.; Ramm, M. T.; Voy, G. T.; Powell, D. R.; Accola, M. A.; Seehafer, T. R.; Sabelko, J. J.; Pladziewicz, J. R. Intermolecular Electron-Transfer Reactions Involving Hydrazines. J. Org. Chem. 1996, 61, 1405−1412. (26) A smaller ε value (394 M−1 cm−1) has been reported for the λ = 778 nm band for 1+ in aqueous solution. See: Carney, M. J.; Lesniak, J. S.; Likar, M. D.; Pladziewicz, J. R. Ferrocene Derivatives as Metalloprotein Redox Probes: Electron-Transfer Reactions of Ferrocene and Ferricenium Ion Derivatives with Cytochrome C. J. Am. Chem. Soc. 1984, 106, 2565−2569. (27) Macchioni, A.; Zuccaccia, C.; Clot, E.; Gruet, K.; Crabtree, R. H. Selective Ion Pairing in [Ir(bipy)H2(PRPh2)2][A] (A = PF6, BF4, CF3SO3, BPh4, R = Me, Ph): Experimental Identification and Theoretical Understanding. Organometallics 2001, 20, 2367−2373. (28) King, R. B.; Bisnette, M. B. Organometallic Chemistry of the Transition Metals XXI. Some π-Pentamethylcyclopentadienyl Derivatives of Various Transition Metals. J. Organomet. Chem. 1967, 8, 287−297. (29) The student will either know or be informed that paramagnetic compounds are not usually NMR-active (1+ is an exception to this rule, being NMR-active owing to its extremely fast electronic relaxation). See: Miller, J. S.; Calabrese, J. C.; Rommelmann, H.; Chittipeddi, S. R.; Zhang, J. H.; Reiff, W. M.; Epstein, A. J. Ferromagnetic Behavior of [Fe(C5Me5)2][TCNE]. J. Am. Chem. Soc. 1987, 109, 769−781. (30) Rieger, P. H. Electrochemistry, 2nd ed.; Chapman & Hall: New York, 1994; pp 109−127.

F


Complementary Use of Electrochemistry and Synthetic Redox

Recommend Documents