Characterization of the Unusual Product from the Reaction between Cobalt[ll) Chloride, Ethane-1,2-diamine, and Hydrochloric Acid An Undergraduate Project Involving an Unknown Metal Complex Nell F. Curtis Victoria University of Wellington, Wellington. New Zealand Robert W. Hay University of Stlrling, Stirling, Scotland, U.K. Donald A. House University of Canterbuly, Christchurch, New Zealand Graeme H. Searlel University of Adelaide. Adelaide, South Australia, Australia The preparations of known inorganic compounds are germane t o undergraduate practical inorganic chemistry courses. Usually the prepared compounds are characterized by chemical reactions, analyses, and physical measurements, with the results being interpreted in terms of the known formulations, structures, and properties of the compounds. Many experiments of this type are available (14). For advanced students investigative or project work is desirable. This approach can be provided by experiments in which a compound prepared by or provided for the student is of unknown formula (although usually the constituents would he known), and various analyses and measurements are completed to deduce the formula and strucrure. Such experiments provide a research-type experience with a meaningful objective, and the benefits of projects involving integrated sequences of measurements are recognized (5,61. For a project, the "unknown" compound should be of such complexity that a variety of analyses and measurements is required to solve the problem. Metal complexes lead readily to this kind of experiment due to their severul constituents and thediversitv ofstrucruresand ~ r o ~ e r t i ehut s , thisdiversity requires pr%jecta to he tailored the individual compounds. Relatively few project-type experiments based on "unknown" metal complexes seem t o he available, however (7-15), and most of these are based on cohalt(II1) complexes. We here descrihe such an experiment involving the product from the reaction between cobalt(I1) chloride, ethane1.2-diamine (en) and concentrated hvdrochloric acid. This blue compound has been characterized as ethane-1,2-diammonium tetrachlorocohaltate(1I) chloride, (enH& [CoCL]C12, which contains the lahile tetrahedral complex ion [CoCld]2- (16,17). The compound can he prepared readily by students from directions provided, hut its composition is not immediately evident since its formulation has two unusual aspects. The en is protonated and serves as an inert countercation, and this function of en may he less familiar t o students than its role as a ligand. secondly, the compound is a double salt and contains both coordinated chlorides (four) and ionic chlorides (two). However, the complex hydrolyzes instantly in water to give pink Coaq2+and six C1- ions, and the resulting solution contains a number of ionic charges per mole (12) which is in excess of students' experience with
' Author to whom correspondence should be addressed.
Students may consider that some cobait(ili)could be formed by aerial oxidation, although such oxidation would be im~robabieunder the acidic conditions oithe synthesis.
simple complexes. These various properties make this a novel and interesting problem. We outline below the directions that we provide to students for preparing the complex and for carrying out the analyses and physical measurements. The project is carried out by third (final) year BSc students in pairs and requires two laboratory days. The experimental work is apportioned, and the results are collated and discussed, by the pair. Preparation of the Complex
Prepare a solution of cobalt(I1) chloride 6-hydrate (7.0 g) in hot concentrated hydrochloric acid (60 mL, about 80 'C), using an effective fumehood. T o this stirred solution add dropwise a solution of ethane-1,2-diamine (3.0 g, 0.050 mol) in water (5 mL). Allow the solution to cool slowly to 0 OC, when the product will crystallize as blue plates. These are filtered off in a dry sintered-glass Hirsh funnel, washed with small portions of concentrated hydrochloric acid, ethanol and acetone, and dried at 80 "C. (Yield 8.0 g, 80%) Recrystallize this product from hot concentrated hvdrochloric acid, using dry glassware; to avoid premature c&tallization, use 80 mL of acid at about 80 Y! and filter the hot solution through a preheated sintered-glass Hirsh funnel. After cooling, remove the recrystallized product as above. (Recovery 6.4 g.) Informailon Provlded The synthesis indicates that the constituents of the comWe give the pound are en, C1- and Co, probably as CO(II).~ nitrogen analysis as (14.5 f 0.4)%, and state that the formula weight is in the range 300-500. This nitrogen proportion allows the formula weight to he deduced as 193 6 per en. Thus there are two en in the formula weight, which is therefore about 386.
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Analyses and Measurements Analysis for CobaR
This may be carried out by atomic absorption at 304.4 nm. Cobalt standards of 40, 60, 80, and 100 ppm are provided, and a solution of the "unknown" of about 70 ppm cobalt should he made up for the analysis. In preparing this unknown solution, one Co per formula weight has to he assumed. Alternatively, cobalt may he estimated gravimetrically as the tetrapyridine complex [Co(py)a](SCN)z (18), using about 0.3 g of the unknown compound. Volume 63 Number 10 October 1986
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Analysis for Total Chlorine
The followine nrocedure allows for the nossihilitv of chlorine being coo;dinated in an inert compleb. A sample of the compound (about 0.13 g) is heated in a beaker with NaOH (30 mL of 1M ) a t just below boiling for 5 min. The solution is cooled, HN03 is added (30 mL of 2 M), and the volume is adjusted with water (to about 120 mL). The chloride is then titrated votentiometricallv with standard 0.05 M AeNO?solution. Alternatively, chloride may be estimated gravimetrically as AgCl(18), using about 0.2 g of the compound. Analysis for Ethane- 1.2diamine
This is carried out by decomposing the complex with sodium sulfide, removing the liberated en by a continuous distillation with toluene, and titrating the en with standard 0.1 M HC1, as described previously (19). Magnetic Susceptlbilify
This can be measured by the Gouy method, and p,ncalculated. Diamagnetic corrections can be made (about 240 X 10-%.g.s. units mol-' (1.4) or 3000 X m3 mol-I (2,3)), but they have little effect on perf. Conductance in Water
From measurements of the conductances a t 25 OC of solutions of different concentrations (0.016, 0.008, 0.004, 0.001 M), the molar conductance a t infinite dilution is determined and the number of charge units per formula weight is deduced (2-4). Displacement of ti+by Cation-Exchange
Du~licatecolumns (1.26 X 6 cm wet-bed) of cation-exchange resin (Dowex ~ o w : x ~200-400 , mesh; completely in the H+ form) are v r e ~ a r e dand c o m ~ o u n d(about 0.2 e) is passed through each eolumn, and ali the effluent together with washings is titrated with standard 0.1 M NaOH. The number of moles of H+ displaced per mole of compound is calculated. pH Titration with NaOH
Titrate an aqueous solution containing about 0.2 g of the compound with standard 0.1 M NaOH, and hence calculate the number of moles of H+ released per mole of compound. Visible Spectra
Visible absorption spectra run in water and dimethyl sulfoxide are provided or can be recorded by the students. Results and Dlscusslon An earlv observation is the instant color chanee in water. irorn bludto pink. This should suggest a lahile ciirnplex and that robaltlll) is involved therefore rather than cobalt(I11).' Students should then realize that the pink solution must contain Coaq2+.Depending on their past chemical experience, students may associate the blue color with tetrahedral rr.,,ri.l2,-" -.*, . The results for the Co, CI, and en analyses should he calculated as mole of component per gram of compound, and the mole ratio will then emerge as Co:Cl:en = 1:6:2. Assuming one Co per mole, this stoichiometry gives a formula weieht of 392 which is within the ranee stated. (The actual formula weight is 395.9). The presence of additibnal hydrogen (four) in the formula will not he evident from the ahove analyses (H = 1.0% wlw). However, astute students may realize that the ahove stoichiometrv does not allow charee balance, so that further cations must be involved. Calculations to be made from the other measurements require the use of an approximate formula weight, andclose values are available from the nitrogen analysis provided and 900
Journal of Chemical Education
from the deduced stoichiometry as above. Formula weights can also be calculated from the students' individual analytical results, by using the stoichiometry values. The p.ff measured is 4.6 Bohr magnetons which is within the range observed for tetrahedral Co(I1) (4.3-4.8), and this excludes the possibility of a Co(I1)-Co(II1) mix.2 The conductance measurements a t 25 "C give ADO as 820 ohm-' cm2 mol-', which indicates that the complex in solution gives 12 charge units per formula weight. More than two types of ion must be present in the solution, therefore, and probahly in the solid as well. The markedly different visible 510 nm. r 5 mol-' dm3 cm-') and sDectra in water (X, . ., dimethyl sulfoxide (A, 680, ;590, plus fine structure) should suggest that a tetrahedral complex is present in the weakly coordinating nonaqueous solvent and presumably in the solid state also (16). This must he ICoCL12-, and the other constituents of the solid are deduced f& the stoichiometry as 2CI- and 2(enH2)2+. In water these give 2(enH2)2++ Cow2++ 6C1-, accounting for the conductance results. Insolubility precludes conductance measurements in nonaqueous solvents. The cation-exchange experiment shows that six moles of H+ are disnlaced ner mole of comnlex (exnerimental values . . are around 5.95), indicating that the complex givesrise t o six cation charees in solution. Two of these are accounted for as Coaq2+,anduthe other four charges have therefore to be associated with the en, that is as 2 ( e n H ~ ) ~The + . applied band on the resin gives a definite appearance of containing two components; the front of the band is the most intensely pink (Coq2+), SO that there is a second component present which binds more strongly tothe resin. With the other information, this second component may be deduced as enHz2+. The end point of the pH titration corresponds to four moles of Hf released per mole of complex. (Experimental values are about 4.1). This implies that 4H+ must be combined with a base in the comnound. that is as 2(enHd2+. . ". During the titration there is a progressive pronounced color chanee to oranee-brown. which should sueeest that a free liganld (en) is geing released from its protonated form to c o m ~ l e xwith Co(I1). he infrared spectrum of the rornpound is not particularIv iniorrnative (16. 17). It is relatively simvle with several sharp bands, including two a t 1600 A d 1580 em-' which, together with complex absorption over the region 3100-2000 cm-', support the assignment of a structure a i t h protonated amine present. The absence of bands beyond 3200 cm-' and near 1630 em-' support the assignment of a non-hydrated structure for the product. Exneriments of this tvve . can be made of variable difficulty, depending on how much information is available about the compound initially and on the amount of guidance provided for carrying out the measurements. The ahove outline of directions demonstrates the scone of the characterizations which are possible with this compound, but if most of these are within the student's previous exoerience. the directione for characterizing the compound c o k d be restricted. A
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The Corresponding Copper(l1) Complex The product of the analogous preparation using copper(I1) chloride is (enHz)[CuClr],which provides an alternative unknown complex. This is not a double salt, and it thus provides for a simpler project experiment, yet i t permits most of the above technioues to he annlied. The orenaration is similar to that of the cibalt(11) complex excep't &at lower hvdrochloric acid concentration is used because the coopkr(11) complex is less soluble. Addition of en (3.0 g i n 10 m ' water) to CUCIY~HIO e in 60 mL of 6 M HCI) a t 60 O C - (7.0 . gives brownish-yellow cryst&. (Yield 9.5 g, 86%).This product is recrvstallized from 6 MHC1 a t 85 OC (about 200 mL is required);giving recovery about 75%. For this complex, copper may be determined volumetri-
~
cally by the iodine-thiosulfate method (18). Dissolve aeparate samples of about 0.1 g in water, add 5 mL of 10%KI solution, and titrate the liberated iodine with standard 0.01 M NazSzO3solution. Alternatively, atomic absorption may be used for the copper analysis. The other analyses and measurements may be carried out as for the cobalt(I1) complex, except that magnetic susceptibility 1.9 B.M.) is not a useful measurement in this instance. The complex clearly dissociates in water, since it gives a pale blue solution. This contains CuW2+(the spectrum shows a single d-d band at 810 nm), and the conductance (Am0 590 ohm-' em2 mol-1 at 25 "C) indicates eight charge units per formula weight. Four moles of H+ are displaced from cationexchange resin per mole of complex (experimental value 3.981, and two cation components are evident since Cuaq2+ binds to the resin ahead of colorless enHz2+.The pH titration shows that two protons are released per mole of complex, and during the titration the color changes progressively to deep blue as the en is deprotonated and can then complex with Cu(I1).The yellow color of the complex ismaintained in dimethyl sulfoxide, and the spectrum in this solvent shows the yellow color to result from charge transfer extending into the blue. The d-d absorption is indicated t o be in the near
infrared, and this low energy is consistent with a tetrahedral complex [CuCl4I2-rather than square. Projects involving other "unknown" inorganic compounds are being prepared for publication. Literature Cited (11 Adam. D. M.: Ramor, J. B. '"Advanced Pmtieal Inorganic Chemiaw"; Wiiy: London. 1965. (2) Jolly, W. L. "The Synthesis aod Characterization of Inorganic Cmnpoundl": Prentia~Hall:Englewiood Cliffs, NJ. 1970. (3) Marr. G.; Rocketf, B. W. "Practical I m r p n i c Chemistry"; Van Nwtrand Reinhold: London, 1972. (4) Bell, C. F. "Synthesis and P h y a i d Studies of Inorganic Compounds": Pewmon: Oxford, 1972. (5) Hunt. G. R A. J Chom. Educ. 1916.53.53. (6) C8rtwight.H. M. J.Chem. Educ 1980,57,309. 171 Thielrnann. V. J. Chem. Edue. 1974.51.536. E ~ U C . lgn,52,63. isi D U ~ ~ E. L .P. J. (9) Testa, J.F.; Wyatt.J.L. J. ChsmEduc. 1915,52,50. (I) William,T. R.;Haynes,L. W. J,Chem.Educ. 1977.54,246. (11) Meek,E.G.School Sci.Reu. 1977.58.489. (12) Alexander, J. J.;Doney, J.G. J.Chem.Educ. 1978,55,2M. (13) Nathan. L. C. '"Laboratory Projecf in Phyaieal Inorganic Chemiaw"; BrooLs.Cole: Monterey, CA. 1981. (11) hehlin. J. H.; Kehl, S. B.; Darlington, J. A. J. Chem. Educ. 1382.59.1048. (15) Pickering, M. J. Chem. Educ. 1985.62.142. (16) Smith,H. W.;Stratton, W. J . Inorp. Chrm. 1977,16,1640. (17) Bsnares. M. A,; Angaao. A.; Rodrigua. E. Polyhedmn. 19% 3.363. (18) Vogel, A. I. .'A Terfbwk of Quantitative Inorwnic Analysis Including Elementary Inatrumrntal Analysis", 3rd ed.. Longmans: London, 1961, pp 358,460,52941. (19) Searle, G.H. J.Chem.Educ. 1985.62.882.
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Number 10 October 1986
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