In the Laboratory
Synthesis and Base Hydrolysis of a Cobalt(III) Complex Coordinated by a Thioether Ligand
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Lee Roecker Department of Chemistry, Gettysburg College, Gettysburg, PA 17325 and Department of Chemistry, Berea College, Berea, KY 40404;
[email protected] Experiments relating to the study of inorganic reaction mechanisms have often been described in this Journal. For example, visible spectroscopy has been used to study the spontaneous, base-assisted, and metal-assisted hydrolysis of the pentaamminechlorocobalt(III) ion (1) and the base-catalyzed isomerization of the oxygen-bonded to the nitrogen-bonded linkage isomer of [Co(NH3)5NO2]2+ (2). The aquation of the trans-tetraamminedichlorocobalt(III) ion has been studied by 1H NMR spectroscopy (3). These experiments relate to topics commonly taught in an advanced inorganic course. An experiment is described that uses UV spectroscopy to study the base hydrolysis of a cobalt(III) complex coordinated by a thioether ligand:
H2 N H2 N
NH2 Co
CH3
3+
2+ ź
OH
S
H2N H2N
N NH2 H2
NH2 Co
OH S
dition of hydroxide can return to its protonated state or lose the chloride to produce a five-coordinate intermediate in an SN1 process. The five-coordinate intermediate quickly picks up water and undergoes a proton rearrangement to yield the hydroxo product. While many essentials of the reaction mechanism were described in the 1960s (6), study of the base hydrolysis of Co(III) complexes remains an active area of research (7). Experimental Procedure Synthesis The synthesis of the thioether complex (8) can be completed in a three-hour laboratory session and begins with the preparation of [(en)2Co(SCH2CH2NH2]I2, 1:
CH3
(1)
N NH2 H2
H2O H2 O
This experiment introduces students to the SN1CB base hydrolysis mechanism (CB is conjugate base) and provides them with additional experience in the collection and analysis of kinetic data. Also, the two-step synthesis illustrates two common strategies used in the preparation of cobalt complexes. Highlighting the concepts of lability and inertness, labile Co(II) is first oxidized to an inert Co(III) complex. A new ligand environment about the inert Co(III) center is then formed by an organic transformation of a coordinated ligand (4, 5). The SN1CB mechanism for the base hydrolysis of the pentaamminechlorocobalt(III) ion is illustrated in Scheme I. The conjugate base of the cobalt complex that forms upon ad-
H3 N
NH3 Co
H3N
Cl
+ OHź
K
H3 N
NH3
NH3
H3 N
Co
H3N
NH3
2+
OH NH3
NH3
+ H2O
NH2
loss of leaving group
NH3
Cl
Co
H3N
NH3
1+
H2O fast
H3 N H3N
k
2+
NH3 Co
NH2
NH3 5-coordinate intermediate
Scheme I. The SN1CB mechanism illustrated for the [(NH3)5CoCl]2+ ion.
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Co
OH2
OH2
OH2
Cl2
1. 1,2-diaminoethane 2. (NH2CH2CH2S)2 3. KI
H2N H2N
NH2 Co
S
I2
N NH2 H2
(2)
1
Air is excluded from the reaction solutions to ensure that the disulfide is responsible for the oxidation of Co(II) to Co(III). The reduction of the disulfide to the thiolate results in the bidentate thiolato ligand being trapped in the coordination sphere of the inert Co(III) ion. Students study the base hydrolysis of [(en)2Co(S(CH3)CH2CH2NH2]I3, 2. This complex contains a coordinated thioether ligand and is prepared by the reaction of 1 with iodomethane in DMF (9):
H2N 2+
OH2
H2N
NH2 Co
S
N NH2 H2 1
1. CH3I/DMF
I2
2. KI
H2N H2N
NH2 Co
CH3 S
N NH2 H2
I3
(3)
2
The products from both reactions are isolated as iodide salts. Starting with 2.0 g of cobalt(II) chloride, typical isolated yields of 1 are 60%. Isolated yields of 2 are 80% when starting with 0.5 g of 1. The complexes are characterized by UV–vis spectroscopy. The spectrum of each exhibits an intense ligandto-metal charge transfer (LMCT) transition at 282 nm (1: ε = 14000 M‒1 cm‒1; 2: ε = 8400 M‒1 cm‒1) and a d–d transition in the visible region. Comparing the peak positions and ratios (AUV/AVIS) to the literature values is a convenient way to confirm the presence of a Co(III)–S bond and that the reactions have gone to completion. Kinetics The Co–S bond in 2 breaks in strongly basic solution (eq 1). The resulting loss of absorbance at the 282 nm transition
Journal of Chemical Education • Vol. 85 No. 11 November 2008 • www.JCE.DivCHED.org • © Division of Chemical Education
In the Laboratory 0.8
Absorbance
Absorbance
0.2
ln(At Ainf)
1
0.4
0.6
0.4
y = 0.0077x 0.8135 R2 = 0.9999
2
3
0
100
150
200
250
300
350
400
Time / s
0.0
0.0 250
350
450
550
0
650
200
400
600
800
1000
Time / s
Wavelength / nm Figure 1. UV–visible spectrum of the thioether complex 2 in 0.01 M CF3SO3H. The spectrum at the left was obtained on a 1:50 dilution of the sample used to obtain the visible spectrum.
Figure 3. Typical absorbance change observed during the course of the Co–S cleavage reaction at 282 nm, [OH‒ ] = 0.050 M. The inset shows a plot of ln|At – Ainf| vs time.
0.6
0.020
0.5
Time: 1 min 2 min 3 min 4 min 5 min 10 min
0.4 0.3 0.2
y = 0.1487x 0.0002 R2 = 0.9993
0.015
kobs / sź1
Absorbance
50
0.2
0.010
0.005
0.1 0.0 250
275
300
325
350
Wavelength / nm
0.000 0.00
0.02
0.04
0.06
0.08
0.10
0.12
Hydroxide Concentration / (mol/L)
Figure 2. Repetitive scans from 250 to 350 nm in 0.050 M NaOH (I = 0.100 M, NaCl). Top to bottom, time = 1, 2, 3, 4, 5, and 10 minutes.
Figure 4. Plot of kobs vs [OH‒ ] for the base hydrolysis of 2 at 25 °C.
offers a convenient signal to monitor the kinetics of the Co–S cleavage (10). Students start the kinetic study by confirming that a large absorbance change occurs at 282 nm by obtaining repetitive scans from 350 to 250 nm in 0.050 M NaOH. Kinetic data are then collected at 25 °C from [OH–] = 0.010 to 0.100 M (I = 0.100 M, NaCl). Runs are monitored at each concentration until an infinite-time absorbance (Ainf ) reading is reached. Plots of ln|At – Ainf | versus time are linear; the slopes of these plots yield values of kobs at each base concentration. The second-order rate constant for the reaction is then obtained from the slope of a plot of kobs versus [OH–]. In a four-hour laboratory period, data can be collected for the highest four base concentrations (0.100, 0.075, 0.050, 0.025 M). Data collection for the 0.010 M NaOH solution can be started at the end of the period, at a later time, or omitted.
Experimental Results
Hazards 1,2-Diaminoethane is corrosive; iodomethane is a suspected carcinogen; diethyl ether is flammable; DMF is an irritant, toxic to the liver, and flammable. In addition, all of these are volatile. Cobalt(II) chloride is an irritant and is toxic. Cystamine dihydrochloride is hazardous in case of ingestion. The 1,2 diaminoethane and the iodomethane should be weighed in a fume hood and transported in closed containers to the reaction hood.
The preparation of 1 is successful if attention is paid to reaction stoichiometry and the exclusion of O 2 during its preparation. Key to making this a more timely preparation than that originally reported (8) is the use of potassium iodide to precipitate the cobalt thiolato complex. Students should be able to prepare and isolate the brown iodide salt in an hour. Preparing and isolating the orange thioether complex, 2, takes about 45 minutes when using iodomethane as the alkylating agent. The UV–vis spectrum obtained for complex 2 is shown in Figure 1; note that the UV region was scanned after a 1:50 dilution. The peak ratio of A282/A484 = 46 compares favorably to the literature value of 47 (9). Repetitive scans in the UV region of the spectrum are shown in Figure 2; these illustrate the absorbance change that accompanies the reaction. A drop of aqueous 2 is added to a thermostated solution of aqueous base, mixed, and absorbance data collected at 282 nm. The typical absorbance change observed during the course of a reaction is shown in Figure 3; the inset shows a plot of ln|At – Ainf | versus time. The linearity of the plot of ln|At – Ainf | versus time is consistent with a pseudo-firstorder decay of 2. To determine the reaction order with respect to [OH–], a plot of kobs versus [OH–] is made (Figure 4). The linearity of the data indicates that the reaction is also first-order
© Division of Chemical Education • www.JCE.DivCHED.org • Vol. 85 No. 11 November 2008 • Journal of Chemical Education
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In the Laboratory
with respect to [OH–]; the slope of line is the second-order rate constant, kOH. Examination of the dependence of kobs on [OH–] based on the SN1CB mechanism (Scheme I) leads to kobs = kK[OH–]/(1 + K[OH–]. Departure from a first-order dependence, however, is not observed because the low acidity (small K ) of the cobalt complex (6, 7, 11, 12) simplifies the expression to kobs = kK[OH–]. The second-order rate constant (kOH) that students obtain from their linear plot of kobs versus [OH–] is a composite value, kK. Pedagogical Value Both the synthetic and kinetic portions of this exercise have the potential to enhance student learning. Synthetically, this experiment has several attractive features. It provides another laboratory exercise that illustrates the concepts of lability and inertness concerning the 2+ and 3+ oxidation states of cobalt that students study in lecture. Preparation of the thiolato complex alerts students to the oxidizing ability of disulfides while preparation of the thioether complex illustrates the utility of organic chemistry in inorganic synthesis and refreshes students’ memories of SN2 reactions. Concerning technique, the experiment exposes students to the use of gas cylinders, recrystallization, and UV–vis spectroscopy. The complexes are interesting spectroscopically because, in addition to a d–d transition typical of Co(III) complexes, the second d–d transition is obscured by the intense LMCT transition typical of a bond between Co(III) and sulfur. Once characterization is completed, students have a better understanding of how the molar absorptivities of d–d and LMCT transitions differ and how these differences can be useful to characterize a material and, in this case, to follow the course of a reaction. The kinetics portion of this laboratory also has features that make it desirable to use in the undergraduate laboratory. This experiment helps students to refine their data collection skills. The large absorbance change that occurs during the reaction makes the experiment fairly forgiving towards students who take longer than desired to mix solutions and return them to the spectrometer. However, if they make a more severe mistake, it is easy to stop, clean the UV cell, and try again. The reaction is easy to initiate and, because the reaction is monitored at a maximum that has a large molar absorptivity, only a small quantity of cobalt complex is needed to complete the kinetic runs. Students are usually apprehensive about collecting data prior to the experiment—they do not quite understand how to do it, so I assist them with their first trial. After collecting a data set or two on their own, they are usually quite confident and their technique continues to improve throughout the experiment. Students pose a variety of questions regarding data analysis because this kinetics experiment is more involved than those they have completed in earlier courses. While the computer attached to the spectrometer has software that will calculate rate constants, it is of more instructional value to have students export the data into a spreadsheet to perform the required manipulations. I suggest that students start data analysis by following the instructions provided in the laboratory handout and require that they show me their results and spreadsheet from the first trial. Performance related feedback (13) is then provided once it is apparent what each student does not understand. Common concerns that are addressed include improper mixing
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of solutions (evidenced by non-exponential decay at the beginning of the run), the concept of an infinite-time absorbance, how to determine the number of data points that encompass four half-lives, and the best way to set up the spreadsheet. Once students understand how to analyze one set of data, they are usually able to analyze the remaining sets. Conclusions The kinetics portion of this experiment has been part of the third- or fourth-year advanced laboratories at Berea College for the past eight years (14). Using instructor supplied 2, students complete the experiment in one to two weeks. It has been a useful way to introduce an inorganic reaction mechanism to students. At Gettysburg College, the synthetic component was introduced into the laboratory that accompanies the advanced inorganic course and has been used for two semesters. The synthesis complements other preparations done in the laboratory and, similar to the experience at Berea College, it provides students with practice in the collection and analysis of kinetics data. Students at Gettysburg College complete the synthesis and the kinetics study over the course of two four-hour laboratory periods with an additional visit to complete the slowest reaction conditions. If desired, the synthesis portion of the laboratory can function as a stand-alone experiment. Literature Cited 1. González, G.; Martinez, M. J. Chem. Educ. 2005, 82, 1671–1673. 2. Jackson, W. G.; Lawrance, G. A.; Lay, P. A.; Sargeson, A. M. J. Chem. Educ. 1981, 58, 734–738. 3. Orvis, J. A.; Dimetry, B.; Winge, J.; Mullis, T. C. J. Chem. Educ. 2003, 80, 803–805. 4. Harrowfield, J. M.; Lawrance, G. A.; Sargeson, A. M. J. Chem. Educ. 1985, 62, 804–805. 5. Hamilton, D. E. J. Chem. Educ. 1991, 68, A144–A146. 6. Tobe, M. L. Acc. Chem. Res. 1970, 3, 337–385. 7. Jackson, W. G. Inorg. React. Mech. 2002, 4, 1–30. 8. Nosco, D.; Deutsch, E. Inorg. Synth. 1982, 21, 19–23. 9. Elder, R. C.; Kennard, G. J.; Payne, M. D.; Deutsch, E. Inorg. Chem. 1978, 17, 1296–1303. 10. Roecker, L.; Deutsch, E. Inorg. Chem. 1985, 24, 16–24. 11. House, D. A. Coord. Chem. Rev. 1977, 23, 223–322. 12. Dixon, N. E.; Jackson, W. G.; Marty, W.; Sargeson, A. M. Inorg. Chem. 1982, 21, 688–697. 13. Brooks, D. W.; Schraw, G.; Crippen, K. J. J. Chem. Educ. 2005, 82, 637–640. 14. Roecker, L.; Baltisberger, J.; Saderholm, M.; Smithson, P.; Blair, L. J. Coll. Sci. Teach. 2007, 36, 36–44.
Supporting JCE Online Material
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Notes for instructors including possible modifications
Journal of Chemical Education • Vol. 85 No. 11 November 2008 • www.JCE.DivCHED.org • © Division of Chemical Education
