In the Laboratory

JChemEd.chem.wisc.edu • Vol. 75 No. 6 June 1998 • Journal of Chemical Education 773

The Microscale Laboratory

The Microscale Synthesis and Electrochemistry of Low-Valent Mononuclear Complexes (h3–C3H5)Fe(CO)3X( X = I, Br, Cl )

Enrico Mocellin,* Richard Russell, and Mauro Ravera**AMC2 – Australasian Microscale Chemistry Centre, School of Biological and Chemical SciencesDeakin University, Geelong, Victoria 3217, Australia

We have previously outlined the value of unified micro-scale laboratory programs (1). Continuing with this theme, wenow describe the microscale synthesis and electrochemical be-havior of the mononuclear complexes (η3–C3H5)Fe(CO)3X(X = I, Br, Cl). This experiment is especially well suited tostudents who have completed a basic lecture series on organo-metallic chemistry and associated electrochemical techniquesand who have an understanding of linear sweep voltammetryat solid electrodes (2).

The reaction of diiron nonacarbonyl with allyl halides(3) affords mixtures of the stereoisomer η3 complexes 1 and2 as shown in eq 1.

Fe2(CO)9 + CH2=CH–CH2X → (C3H5)Fe(CO)3X + Fe(CO)5 +CO

(1)

CO

FeOC X

CO CO

FeX CO

CO

1 2

HC

CH2

CH2 HC

CH2

CH2

endo exo

X = I, Br, Cl

Experimental Procedure

General Methods

All materials were used as purchased without further puri-fication. Elaborate glassware-drying proceduresare not required.

For the electrochemistry a standard three-electrode cell is configured to allow the tip ofthe reference electrode (saturated standardcalomel electrode [SSCE] or silver/silverchloride [Ag/AgCl]) to be proximate to theworking electrode (hanging mercury dropelectrode [HMDE] or mercury amalgam ongold electrode [Hg/Au]) and platinum solidelectrode. The cell is completed with an aux-iliary electrode (platinum wire) and a two-way gas bubbler. Experiments are performed

in acetonitrile, dichloromethane, or tetrahydrofuran. Allmeasurements are carried out under N2 or Ar atmosphere inquiescent deoxygenated solvent solutions, which have beenpurified according to the standard literature procedures.Solutions (5 mL) are prepared 1 × 10{3 M with respect to theanalyte and 1 × 10{1 M in [Bu4N][PF6].

Preparation of (C3H5)Fe(CO)3X (X = I, Br, Cl)Attach a nitrogen inlet to the side arm of a 25-mL two-

necked round-bottom flask containing a magnetic stirring bar,and flush the apparatus for several minutes. Without discon-necting the inert gas tube, introduce into the flask Fe2(CO)9

1

(250 mg, 0.68 mmol), 5 mL of hexane, and the relative quan-tity of halide (C3H5X: X = I 64 µL, 0.70 mmol; X = Br 60 µL,0.70 mmol; X = Cl 61 µL, 0.75 mmol). Connect a watercondenser to the flask and, using Tygon tubing, attach an oilbubbler, thus also providing an outlet for the gas purge. Placethe flask in a sand bath on a magnetic-stirring hot-plate andheat. The suspension is heated at 40 °C for the indicated timesdepending on the allyl halide used (X = I or Br, 35 min, X =Cl, 40 min). CAUTION: allyl chloride boils 45 °C. Maintain avery slow flow of nitrogen during the reaction to help removethe CO, but avoid stripping the solvent. Cool the solutionto room temperature and filter through a short chromatog-raphy column (1 cm × 1 cm) of silica (80–230 mesh), to re-move FeX2. The separation is accelerated by applying a slightvacuum to the filter flask at the bottom of the column. Washthe silica with small portions of hexane until the silica iswhite. The solvent is removed from the relevant fraction un-der reduced pressure.2 Recrystallize the target products fromhexane. Yields: I (1:1) 67%, Br (7:1) 46%, Cl(25:1) 35%.

*Corresponding author.**On leave from Dipartimento di Chimica

Inorganica, Chimica Fisica e Chimica deiMateriali, Università di Torino, via P. Giuria 7,10125 Torino, Italy.

xelpmoC)oxe:odne(

RMN a

RI bH1 )d( H2 )d( H3 )m(odne oxe odne oxe odne oxe

I=X)1:1(

38.3 73.2 12.4 97.3 46.4 39.5 ,sv4302,sv0802sv9002,hs4102

rB=X)1:7(

05.3 95.2 34.4 21.4 89.4 15.5 ,sv5402,sv0902sv1102

lC=X)1:52(

43.3 47.2 35.4 63.4 02.5 23.5 c ,sv1502,sv6902sv3102

a Chemical shift (δ) ppm in CDCl3 downfield from TMS observed at 270 MHz; d = doublet,

m = triplet of triplets. b νCO (cm{1) in hexane. c Partially obscured by the major isomer.

Table 1. IR and 1H NMR Spectral Datafor Complexes under Study

In the Laboratory

774 Journal of Chemical Education • Vol. 75 No. 6 June 1998 • JChemEd.chem.wisc.edu

Discussion

The presence of three car-bonyl bands in the solution IRspectra (Table 1) of each com-plex is consistent with CSrather than C3V symmetry. Therelative position of the bands isa function of the electronegativ-ity of the halides and their abil-ity to influence the Fe–COback donations.

The 1H-NMR spectrumof each individual reactionproduct (Table 1) shows twodistinct A2M2X systems of a symmetrical π-allyl group in themajor endo and minor exo isomers. The relative ratio of theexo/endo isomer varies with the halogen X (4 ).

Electrochemical BehaviorBoth endo and exo isomers of all three complexes exhibit

reversible cyclic voltammograms, indicative of two one-electronredox couples. Students are asked to record CV responses atHg and Pt electrodes for all the compounds (see ExperimentalProcedure for the conditions), at different concentrations(from 1 to 10 mM for at least 3–4 different concentrations)and with at least a decade of variation of scan rate for eachconcentration (preferably two orders of magnitude). Theresultant tabulation of the relevant parameters for eachvoltammogram should contain peak potentials, Ep, peak cur-rents, ip, and half-wave potentials, Ep/2, where Ep

a is anodicpotential and Ep

c is cathodic potential (2).The electrochemical behavior of the complexes depends

on the electrode material (Fig. 1) (5). In Figure 1a, the CV atHg exhibits two cathodic peaks, A and B, of equal height,each of which exhibits an associated reoxidation peak (D andC, respectively) in the reverse scan. Analysis of the CV re-sponse (6 ) with the scan rate varying from 0.05 to 5.0 V s{1

indicates that:1. The formal electrode potential values, E°′, evaluated

from (Epa + Ep

c)/2, are independent of the scan rate;2. The ip

a/ipc ratio of both peak systems remains equal

to unity, suggesting full chemical reversibility for bothprocesses.

3. For the first peak couple (A/D), the ∆Ep = Ep(A) –Ep(D) value increases from 70 to 200 mV, suggesting anelectrochemically quasi-reversible process, while thesecond peak system, B/C is a reversible process

4. Plots of ip vs ν1/2 and ip vs depolarizer concentrationfor each peak is linear and passes through the originas postulated by the Randles–Sevcik equation.

5. The ratio of ip(B) and the oxidation peak current offerrocene, added in equimolar concentration as internalstandard, is 1, demonstrating that all the processesinvolve one electron.

6. The electrode potential of the second reduction stepis almost independent of the nature of X and doesnot change significantly if excess X- is added to thesolution (i.e., as [Bu4N]X salts), whereas the first peakis profoundly affected by this addition (Le ChâtelierPrinciple).

At a Pt electrode the CV looks very different (Fig. 1b).A cathodic peak E (at a potential near that of peak B, found

in the previous experiment) is observed, followed in the re-verse scan by a directly associated reoxidation peak F (half inheight) and by a broad peak G at a more positive potential.

On the basis of such results, E1CE2 redox mechanismscan be proposed (on Hg, E1 > E2; on Pt, E1 ≈ E2), as seen inthe proposed mechanism (Scheme I) involving the reductivecleavage of the Fe–X bond (peaks A to E) with the subsequentformation of the X { anion and [(C3H5)Fe(CO)3]• radical,which is further reduced to the stable [(C3H5)Fe(CO)3]

{ an-ion (peaks B to E).

(η3 -C3H5)Fe(CO)3 X [(η3 -C3H5)Fe(CO)3 X]

[(η3 -C3H5)Fe(CO)3 ] [(η3 -C3H5)Fe(CO)3 ]

+ e

E1

+ e

E2

-X-

[(η3 -C3H5)Fe(CO)3 ]

Scheme I: Proposed mechanism for the electrochemical andchemical reactions

The dramatic difference in the redox behavior of the(C3H5)Fe(CO)3X complexes at the Hg and Pt electrode, as seenin Table 2, provides unambiguous evidence for the involvement

Figure1. CV response of a CH3CN solution of η3-(C3H5)Fe(CO)3Br at room temperature, scan rate0.2 V s{1: at (a) Hg and (b) Pt electrode; (•) starting potential.

(a) (b)

aFor the reversible processes, the E°′ value [E°’ ≈ (Epc + Ep

a)/2] isreported. For irreversible processes, the Ep value is reported at a scanrate of 0.20 V s{1.

niseireScaehtfoslaitnetoPedortcelExodeR.2elbaTstnevloStnereffiD

xelpmoC )s(kaeP edortcelE)ECSsvV(E a

HC 3 NC HC 2 lC 2 FHTI=X D/A

C/BgHgH

{ 83.0{ 07.0

{ 12.0{ 56.0

{ 65.0{ 77.0

F/EG

tPtP

{ 76.0{ 71.0

{ 36.0{ 80.0

{ 87.0{ 32.0

rB=X D/AC/B

gHgH

{ 04.0{ 07.0

{ 52.0{ 66.0

{ 55.0{ 77.0

F/EG

tPtP

{ 27.0{ 51.0

{ 56.0{ 20.0

{ 87.0{ 42.0

lC=X D/AC/B

gHgH

{ 24.0{ 17.0

{ 23.0{ 56.0

{ 06.0{ 87.0

F/EG

tPtP

{ 08.0{ 31.0

{ 18.011.0+

{ 87.0{ 42.0

In the Laboratory

JChemEd.chem.wisc.edu • Vol. 75 No. 6 June 1998 • Journal of Chemical Education 775

of the Hg in the first reduction step. The interaction of(C3H5)Fe(CO)3X with the Hg surface (i.e., a partial inser-tion of Hg into the Fe–X bond) facilitates the first reductionprocess. In the absence of such an interaction (i.e., at the Ptelectrode), the first reduction is shifted to a more negativepotential (200–250 mV) and the corresponding peak is verynear (I), overlaps (Br), or bypasses (Cl).

The voltammetric behavior should be compared with thepolarographic behavior reported by Gubin and Denisovic (7 )to observe the details that the voltammetric techniques betterdemonstrate.

Notes

1. Iron nonacarbonyl is available from many dealers. However itis possible to produce it by photolysis.

2. The removed solvent could contain Fe(CO)5 and unreactedallyl, and must be destroyed in a solution of sodium peroxide in water.

Literature Cited

1. Mocellin, E.; Ravera, M.; Russell, R. A.; Hynson, T. J. Chem.Educ. 1996, 73, A99.

2. Evans, D. H.; O’Connell, K. M.; Petersen, R. A.; Kelly, M. J. J.Chem. Educ. 1983, 60, 290. Mabbott, G. A. J. Chem. Educ. 1983,60, 697. Kissinger, P. T.; Heineman, W. R. J. Chem. Educ. 1983,60, 703. Van Benschoten, J. J.; Lewis, J. Y.; Heineman, W. R.;Roston, D. A.; Kissinger, P. T. J. Chem. Educ. 1983, 60, 772.

3. Murdoch, H. D.; Weiss, E. Helv. Chim. Acta 1962, 45, 1927.4. Faller, J. W. W.; Adams, M. A. J. Organomet. Chem. 1979, 170, 71.5. Osella, D.; Ravera, M.; Kukharenko, S. V.; Strelets, V. V.;

Hosuecroft, C. E. J. Organomet. Chem. 1991, 417, 421.6. Bard, A. J.; Faulkner, L. L. Electrochemical Methods; Wiley: New

York, 1980. Brown, E. R.; Sandifer, J. R. In Physical Methods ofChemistry; Rossiter, B. W.; Hamilton, G. F., Eds.; Wiley: NewYork, 1986; Vol. 2, p 273.

7. Gubin, S. P.; Denisovich, L. I. J. Organomet. Chem. 1968, 15,471. Gubin, S. P. Pure Appl. Chem. 1970, 23, 463.