The Chemical Educator, Vol. 10, No. 2, Published on Web 02/03/2005, 10.1333/s00897040867a, © 2005 The Chemical Educator

Spectrochemical Series of Cobalt(III). An Experiment for High School through College

Adam R. Riordan, Ariane Jansma, Sarah Fleischman, David B. Green, and Douglas R. Mulford*

Department of Chemistry, Pepperdine University, Malibu, CA 90263, Douglas.Mulford@pepperdine.edu

Received July 7, 2004. Accepted September 2, 2004.

Abstract: We present a laboratory project suitable for high school though upper-division college courses that has been designed to allow students to qualitatively investigate or quantitatively generate the spectrochemical series for octahedral cobalt(III) complexes. This experiment is useful for concrete, visual demonstration of the structural and electronic properties of ligand-to-metal bonding. Water, ammonia, nitrite, oxalate, glycinate, ethylenediamine, 1,10-phenanthroline, carbonate, and cyanide are used as ligands to form the cobalt complexes. Each complex has a different color and l_max. This allows students to derive the spectrochemical series by visual inspection or UV–visible spectroscopy. At the high school level, the series can be generated from visual inspection of the colored compounds by use of the transmission color wheel. At the general chemistry level, the series can be generated from spectroscopic data using the l_max of lowest energy. In advanced courses students can calculate crystal field splitting values, D₀, and the ordering of the spectrochemical series can be justified from a metal–ligand interaction point of view.

Introduction

Concepts of bonding and the electronic and orbital interactions and changes that occur upon bonding are often difficult for students at all levels in the chemistry curriculum to grasp. These difficulties, and the resulting misconceptions, have been the subject of many papers [1–14]. The bonding between metals and ligands is most often explained at the general chemistry level through crystal field theory (CFT). (A brief review of CFT is presented in the supporting materials.) In more advanced classes, the more powerful ligand field theory (LFT), which incorporates both CFT and molecular orbital theory, is commonly used. It is often difficult for students to grasp these seemingly abstract theories without concrete evidence.

Electronic and orbital changes occur when a ligand binds to a transition metal. These electronic changes are manifested in the splitting of the metal dorbitals during complex formation and are often macroscopically seen as a change of color from the free metal to the complex. The magnitude of this splitting, designated D₀ in an octahedral electronic environment, can be measured spectroscopically and depends upon the size and charge of the metal, the ionic nature of the ligand, and according to LFT, the s and p donor/acceptor characteristics of the ligand. It is said that a “stronger” ligand field will produce a larger D₀, and when several ligands are compared, the spectrochemical series can be produced.

Because the spectrochemical series is a fundamental topic of inorganic chemistry and provides a link between the qualitative concepts of chemical bonding and the quantum mechanical changes that occur during bond formation in transition metal complexes, it provides the logical experimental foundation and validation for CFT and LFT. Because of its predictive power for a wide range of metal–ligand interactions, many instructors require their students simply to memorize the spectrochemical series. Rarely are students required to generate the series for themselves, and few experiments exist in the literature which allow students to generate a spectrochemical series for more than a few coordination compounds [15–24].Not only does the synthesis of metal complexes with direct observation of the concomitant color changes provide a satisfying laboratory experience, but experimental generation and verification of the spectrochemical series results in a much deeper understanding of these concepts.

A rigorous analysis of the spectroscopic data to determine the ligand field splitting parameter demands the use of the Tanabe–Sugano (TS) energy-level diagram specific for the metal–complex geometry under study. Despite this requirement, many formally and informally published experiments ignore the use of TS diagrams in the quantitative determination of D₀ or provide the reduced data on relatively simple systems and leave it as an academic process to calculate D₀. This experiment represents the first literature-published example the authors have found that allows for an experimental demonstration of this process.

In the investigation described in this paper, the ordering of ligands into a spectrochemical series in terms of their crystal field strength arises both visually (from the color of the complex in solution) and spectroscopically (by quantitative measurement of the d-orbital splitting). An advantage of this experiment is that it is scalable for high school through upper-division college levels. For introductory levels, the highly colored complexes that result from varying the identity of the ligand bound to a common metal can be used to generate a spectrochemical series by visual inspection. This provides strong physical evidence for the involved electronic effects. With the wide availability of scanning and diode-array UV–visible spectrophotometers, the ligands can be quantitatively ranked by the lowest energy wavelengths of maximum absorption (l_max). Being a direct function of D₀, the l_max of lowest energy can be used to accurately generate the spectrochemical series. It should be noted, however, that contrary to some previously published experiments [22, 23, 25] for d²–d⁹ metal complexes, D₀ cannot be equated to the l_max of lowest energy but must be calculated using a Tanabe–Sugano diagram. In addition, in classes in which metal-to-ligand bonding interactions are discussed in detail, the spectrochemical series can not only be produced, but the ranking of ligands can be justified/explained in terms of ligand-bonding interactions such as s- and p-acceptor/donor.

In this experiment, an extensivespectrochemical series can easily be produced within the time constraints of a normal laboratory class. To clearly demonstrate the crystal field splitting of a wide variety of ligands, this investigation uses a common metal throughout with a single coordination geometry, and all species contain only a single type of ligand. Cobalt(III) is used because it adopts an octahedral geometry for all compounds synthesized in this experiment. There have been other experiments [15–24] that produced limited or mixed metal series, but these previous experiments are inadequate to clearly demonstrate the electronic and orbital changes that occur when a ligand binds to a transition metal.

This laboratory investigation utilizes adapted time-efficient synthetic procedures allowing students to gain both a clearer understanding of the origin of color in certain molecular species and of the structure and bonding information these colors can provide. In addition, these syntheses can be used as a vehicle to discuss many aspects of coordination chemistry including synthesis, equilibrium interactions, solvation effects, formation constants, LeChâtelier’s principle, coupled reactions, and redox chemistry.

Experimental

All solutions are analyzed in situ so isolation of the solid complexes is not necessary. Though some of these synthetic procedures have been known for some time [26], they have been altered in the current experiment to facilitate completion in a standard three-hour laboratory session. If necessary, syntheses can be spread over more than a single laboratory period.

Hazards. Cobalt(II) nitrate hexahydrate (Co(NO₃)₂^·6 H₂O) [CAS 10026-22-9] and cobalt(II) chloride hexahydrate (CoCl₂^·6 H₂O) [CAS 7791-13-1]: Contact with cobalt solutions should be treated by washing with soap and water. Chronic exposure to the cobalt ion and its compounds can result in decreased pulmonary function, contact sensitivity and dermatitis, and respiratory distress. Co(NO₃)₂^·6 H₂O is an oxidizer; contact with strong reducing agents should be avoided.

Ethylenediammine [CAS 000107-15-3] is harmful if swallowed, inhaled, or absorbed through the skin. The pure liquid has a strong odor and should be handled in the hood. Hydrogen peroxide, 30% (H₂O₂) [CAS 7722-84-1] is an oxidizer that can cause mild burns on skin. Care should be taken to avoid contact. In the event of direct contact, wash the affected area immediately with cool water.

Hydrochloric acid (HCl) [CAS 7647-01-0] and nitric acid (HNO₃) [CAS 7697-37-2] are strong acids that can cause severe burns on skin. Care should be taken to avoid contact. In the event of direct contact, wash the affected area immediately with copious amounts of cool water.Nitric acid is a strong oxidizer; contact with certain organic reagents may initiate strong exothermic reactions.

Glacial acetic acid (HC₂H₃O₂) [CAS 64-19-7] is a corrosive liquid with a strong odor. Contact with skin can cause burns and prolonged inhalation can cause lung tissue damage. In the event of contact, wash the affected area immediately with cool water. Glacial acetic acid should be handled in the hood.

Aqueous ammonia (NH₃) [CAS 7664-41-7] has a pungent, suffocating odor. Ammonia gas produced from the concentrated solution is extremely destructive to tissues of the mucous membranes and upper respiratory tract. As with skin contact with any base, the solution produces a slippery sensation and can cause mild to severe burns unless washed off immediately.

Sodium nitrite (NaNO₂) [CAS 7632-00-0] is a stable salt, incompatible with strong oxidizing agents. The salt is used as a food preservative but should be considered toxic if swallowed.

Experimental Conditions. All chemicals were of reagent grade or better and were used as purchased without further purification. In all synthetic procedures, the cobalt(II) complex is formed first then oxidized to the cobalt(III) complex. All UV–visible spectra were recorded without delay after synthesis from 250 to 700 nm for the complexes as synthesized unless otherwise noted. Deionized water was used as the reference solution for all spectra.

Hexaaquacobalt(II) [Co(H₂O)₆]²⁺ stock solution.Dissolve 0.613 g of Co(NO₃)₂^·6 H₂O (2.1 mmol) in 500 mL of deionized water to produce an approximately 0.0042 M stock solution. (Note that the UV–visible spectrum of this complex is not taken.)

Triglycinatocobalt(III) [Co(gly)₃]. Add 0.75 g of sodium glycinate (syn. glycine, sodium salt) (7.7 mmol) to 10 mL of stock [Co(H₂O)₆]²⁺ solution. Add 5 mL of 3% hydrogen peroxide to oxidize the cobalt. Allow the reaction to proceed for approximately 5 min to completely form the purple complex.

Trioxalatocobaltate(III) [Co(ox)₃]^3–. Add 1.55g of K₂C₂O₄^·H₂O (8.4 mmol) to 10.0 mL of the stock [Co(H₂O)₆]²⁺ solution. Add 10 mL of 3% H₂O₂ to oxidize the metal and heat gently to 30 to 40 °C with stirring (10 to 15 min) until blue-green color is seen.

Tricarbonatocobaltate(III) [Co(CO₃)₃]^3–. Dissolve 2.99 g Co(NO₃)₂^·6 H₂O (10.3 mmol) in a solution of 50 mL of water and 0.5 mL 30% H₂O₂. Combine this solution all at once with a solution of 16.9 g NaHCO₃ (269 mmol) in 50 mL H₂O and 0.5 mL 30 % H₂O₂. Use a flask large enough to allow for rapid gas evolution and stir for approximately 1 min until bubbling stops. Dilute 1 mL of the resulting solution to 25 mL prior to obtaining the UV–visible spectrum. Save the remaining solution for the preparation of hexaaquacobalt(III), [Co(H₂O)₆]³⁺ [27].

Hexaaquacobalt(III) [Co(H₂O)₆]³⁺. Add 20 mL of the solution containing tricarbonatocobaltate(III) a few milliliters at a time to 80 mL of 4 M HNO₃. Addition of the [Co(CO₃)₃]^3– solution to the nitric acid should be slow enough to prevent bubbling over. Hydrochloric acid must be avoided because the cobalt(III) ion can be reduced by chloride ion. Obtain the UV–visible spectrum. Save this solution for the preparation of tris(1,10-phenanthroline)cobalt(III), [Co(phen)₃]³⁺.

Tris(1,10-phenanthroline)cobalt(III) [Co(phen)₃]³⁺. Add 0.115 g 1,10-phenanthroline (0.5 mmol) to 10 mL of the solution containing hexaaquacobalt(III). Add to this 10 mL of 3% H₂O₂ and stir for 5 min.

Hexanitrocobaltate(III) [Co(NO₂)₆]^3–. Dissolve 0.38 g Co(NO₃)₂^·6 H₂O (1.31 mmol) in 3 mL of water. Dissolve 3.07 g NaNO₂ (44.5 mmol) in 5 mL of water. Combine the two solutions and add 1.5 mL glacial acetic acid. Note that the solution will bubble vigorously. Dilute to 50 mL, stopper, and shake vigorously for 2 to 3 min. Dilute 1 mL of the solution to 50 mL with water prior to recording the UV–visible spectrum.

Hexaamminecobalt(III) [Co(NH₃)₆]³⁺. In a flask, dissolve 9.74 g Co(NO₃)₂^·6 H₂O (33.5 mmol) and 5.33g NH₄Cl (99.7 mmol) in 50 mL of water. Add 17 mL of concentrated aqueous ammonia and suspend 0.13 g activated charcoal in the solution. Vigorously bubble air through the solution with a fritted-glass gas-dispersion tube until oxidation is complete (approximately 3 to 4 hours). Bubbling with pure oxygen reduces the oxidation time to approximately 1 h. Isolate the resulting charcoal-contaminated precipitate by vacuum filtration using Whatman #1 (or other fine porosity) filter paper in a Büchner funnel. Dissolve the solid in 50 mL of 0.24 M HCl and heat gently for 10 min. Remove the activated charcoal from the hot solution by vacuum filtration in a Büchner funnel. Dilute 1 mL of the solution to 10 mL with water prior to recording the UV–visible spectrum.

Hexacyanocobaltate(III) [Co(CN)₆]^3–.Add 2.55 g CoCl₂^·6 H₂O (10.7 mmol) to approximately 75 mL of water and bring to a boil. With stirring, add a solution of 1.56 g (23.9 mmol) KCN in 30 mL