In the Classroom edited by
Tested Demonstrations
Ed Vitz Kutztown University Kutztown, PA 19530
Cobalt Alums A Demonstration Experiment submitted by:
Claus E. Schäffer* and Poul Steenberg Department of Inorganic Chemistry, University of Copenhagen, The H. C. Ørsted Institute, Universitetsparken 5, DK 2100 Copenhagen Ø, Denmark; *
[email protected] checked by:
Wayne Wolsey and Erica Nostrom Department of Chemistry, Macalaster College, St. Paul, MN 55105-1899
In this demonstration two alums of cobalt are isolated, the dark-blue cesium cobalt alum, CsCo(SO4)2⭈12H2O, and the ammonium cobalt alum, isolated as a light-blue solid solution in (NH4)Al(SO4)2⭈12H2O. Both alums belong to the most common alum type, the α alum. History Alums belong to an isotypic series of cubic double salts with the general formula MIMIII(SO4)2⭈12H2O containing a trivalent hexaaqua ion [M(OH2)6]3+. A related series of isotypic double salts containing divalent hexaaqua ions are the Tutton salts, alternatively called schönites from the mineral schönite, K2Mg(SO4)2⭈6H 2O. Intermittently, these two classes of double salts containing sulfate anions have attracted considerable interest for more than a hundred years. The solid salts of these two series and their aqueous solutions have similar colors and roughly the same visible spectra. It was early realized, therefore, that they contained the hexaaqua complexes of the trivalent and divalent metal ions, respectively, a recognition that contributed to our understanding of the general existence of hexaaqua ions. These two classes of salts have also provided a source of aqua ions of the platinum metals, whose usual commercial sources are chlorides or chlorido complexes. CsIr(SO4)2⭈12H2O was described in 1904 by Marino (1), but a detailed preparative procedure was reported only recently (2). CsRu(SO4)2⭈12H2O(3) has only been known for about 15 years. The same is true for (NH4)2RuII(SO4)2⭈6H2O (4 ), even though the corresponding FeII salt, known as Mohr’s salt (5), was described more than a hundred years ago. CsMo(SO4)2⭈12H2O is of an equally recent date (6, 7 ). For several metals, the double salts also provide a possibility to isolate the less usual oxidation states (e.g., TiIII, VIII, and VII). For trivalent metal ions, there seems to be a connection between the tendencies to form alums and to undergo classical coordination chemistry. When the ions become large enough they do neither. Examples are TlIII, YIII, and LaIII and all trivalent lanthanides, none of which have so far been shown to form alums. At present, 14 trivalent metal ions are known to make double salts of the general formula MIMIII(SO4)2⭈12H2O. These are shown in boldface type below. Al Sc
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Ti
V
Cr
Mn Fe
Co
…
Ga
…
In
Nb Mo Tc
Ru
Rh
Ta
Os
Ir
W
Re
Possible extensions of the alum series could still be the elements shown in italic type in the preceding table. The univalent cations MI can also be varied. Thus, Na+, + K , NH4+, Rb+, Cs+, Tl+, and CH3NH3+ form one or more alums. The anion can be either sulfate or selenate (8). Since water is the most important solvent, aqua ions have always attracted special interest. Moreover, by far the majority of stability constants for complexes use the corresponding aqua ions as a reference system. For example, in the definition of the gross complexity constant β6III = [[CoIII(NH 3) 6]3+]/ ([NH3]6[CoIII]) = 1035 M᎑6, it is understood that CoIII represents [CoIII(OH2)6]3+. In this context it is worth mentioning that the gross constant for the analogous divalent system is βII6 = 105 M᎑6 which means that the standard electrode potential for the system [CoII(NH3) 6]2+/[CoIII(NH3)6]3+ (0.06 V) is 1.77 V lower than the one for [CoII(OH2)6]2+/ [CoIII(OH2)6]3+ (1.83 V). This latter number means that [CoIII(OH2)6]3+ is one of the strongest oxidizing agents known in aqueous solution. Structure The structures of alums were studied in 1935 by Lipson (9), who reported that all alums belong to the same cubic space group Pa3 but can be divided into three subclasses, α, β, and γ. The size of the univalent cation has a major influence on the alum subclass. Most alums have the α structure, whereas alums containing large univalent cations such as cesium more often have the β structure, and there is only one example of the γ structure: NaAl(SO4)2⭈12H2O. In all alums, regardless of subclass, three subunits can be visualized: [MIII(OH2)6]3+, [MI(OH2)6]+, and SO42᎑. All of the subunit central atoms are situated on the crystallographic threefold axis of the cubic unit cell. The sulfate groups are situated between the two different cations and the sulfur atom and one oxygen atom are on the threefold axis. In the unique γ alum, the sulfate oxygen on the threefold axis points toward the small univalent cation, Na. By contrast, the sulfate ion is inverted in the α and β alums so that three oxygens point toward the larger univalent cation, resulting in a total coordination number of 12 (6 from water molecules and 6 from two different sulfate ions). One difference between α and β alums is the positions of the water molecules coordinated to the univalent cation. In both cases, the six water molecules are situated in a ring perpendicular to the crystallographic threefold axis. In α
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alums, this ring is puckered, with O–MI–O angles in the range of 64.6–66.5° (10); together with the 6 oxygen atoms of the two sulfate ions, the 6 oxygens from coordinated water define a rhombohedrically distorted icosahedron (site symmetry D3d, trigonal antiprism). In the β alums, the oxygen atoms of the water molecules form a planar ring (close to a regular hexagon) with O–MI–O angles of 60°, which together with the 6 sulfate oxygen atoms define an approximately cuboctahedral array. The cesium cobalt alum would a priori be expected to be a β alum because of the large size of the cesium ion. Surprisingly, the cesium sulfate alums of group 9 (Co, Rh, and Ir), all with low-spin d6 hexaaqua ions, have turned out to be α alums (11). This leads us to inspect the coordination around the trivalent cations [MIII(OH2)6]3+, which also exhibit distinct differences for α alums and β alums. In the β alums, the plane of the coordinated water molecules contains the metal ion MIII (planar, approximately trigonal coordination of the oxygen atom), whereas in the α alums, this plane is tilted (pyramidal coordination) by as much as 15–35° (12). Within this interval, the water molecules coordinated to Co, Rh, and Ir have the largest angles of tilt (10, 11). This results in decreased ability to form a π bond toward MIII, consistent with the electronic configuration of MIII, which, being low-spin d6, has no π contribution to its molecular orbital stabilization energy. When exposed to an octahedral ligand field, the d orbitals split into two sets of orbitals designated by the symmetry species eg and t2g of the octahedral group Oh. The eg orbitals (dz and dx –y ) point toward the ligands and become σ antibonding; the t2g orbitals (dxy, dyz, and dxz) point away from the ligands and become π antibonding. In alums, Co, Rh, and Ir have the electronic subconfiguration t2g6, meaning that only the t2g orbitals are filled. An additional reason why the water molecules coordinate pyramidally—and thus cause a change to the unexpected α form—could therefore be that this sort of coordination of the water molecules optimizes their σ interaction by making it sp3-like around the oxygen atom and at the same time minimizes their π overlap, thereby partially avoiding closed-shell repulsion from the filled t2g orbitals of the central atom. 2
2
2
Chemistry of CoIII Normally, the hexaaquacobalt(III) ion is prepared by electrolyzing a cobalt(II) sulfate solution using a rotating anode with a high overvoltage toward oxygen (e.g., a platinum or gold electrode) (13). In the present more convenient experiment, [Co(OH2)6]3+ is made chemically by an instant oxidation of [Co(OH2)6]2+ with hydrogen peroxide, a process described previously in this Journal (14). A direct oxidation with hydrogen peroxide is not possible for two reasons. First, it is not thermodynamically feasible because the standard electrode potential of the cobalt system is higher (1.83 V) than that of the hydrogen peroxide system (1.77 V when H2O2 is an oxidant). Second, hydrogen peroxide would be oxidized by CoIII ions, because the standard electrode potential for H2O2 as a reductant is 0.70 V in acidic solution: 2Co3+ + H2O2 → 2Co2+ + O2 + 2H+
(1)
However, by analogy to the situation for the ligand NH3, CoIII can be stabilized relative to CoII by complexation with
the carbonate ligand and thereby it becomes possible to oxidize CoII quantitatively to CoIII with hydrogen peroxide, as described by reaction 2: 2Co(OH2)62+ + 10HCO3᎑ + H2O2 → 2Co(CO3)33᎑ + 4CO2 + 18H2O
(2)
The green solution formed has a catalytic effect on the selfdecomposition of excess hydrogen peroxide, a catalysis that is of fundamental importance to avoid reaction 1 when in the next step, the hexaaquacobalt(III) ion is formed by acidification of the solution: Co(CO3)33᎑ + 6H+ + 3H2O → Co(OH2)63+ + 3CO2 (3) To avoid decomposition of the metastable hexaaquacobalt(III) ion, whose stability decreases with decreasing acidity (15), a large excess of acid and an immediate precipitation of a stable cobalt(III) salt are required. When CsHCO3 is used in reaction 2 and sulfuric acid in reaction 3, one can isolate the hexaaquacobalt(III) ion as the solid salt CsCo(SO4)2⭈12H2O. Cesium is chosen as the cation because it forms the least soluble and most stable alums. When KHCO3 is used in reaction 2 and sulfuric acid saturated with aluminum sulfate in reaction 3, one can after addition of ammonium sulfate isolate a light blue mixture of ammonium aluminum alum and ammonium cobalt alum NH4AlxCoy(SO4)2⭈12H2O, x + y = 1. This salt, a solid solution of NH4Co(SO4)2⭈12H2O in the corresponding Al alum, is stable for years, whereas the pure blue NH4Co(SO4)2⭈12H2O evolves oxygen within hours at room temperature and forms the pink crystalline (NH4)2Co(SO4)2⭈6H2O (13): 2NH4Co(SO4)2⭈12H2O → (NH4)2Co(SO4)2⭈6H2O + CoSO4⭈7H2O + H2SO4 + 1⁄2O2 + 10H2O
The experiments described below demonstrate these two different isolations of [Co(OH2)6]3+. Procedures
Prerequisites Crystals of CsCo(SO4)2⭈12H2O prepared in advance by the method described below The saturated solution of Al2(SO4)3 described in experiment 2 Crystals of NH4Al(SO4)2⭈12H2O or KAl(SO4)2⭈12H2O
Experiment 1, Cesium Cobalt Alum CoSO4⭈7H2O (2 g, 7.1 mmol) is dissolved in water (6 mL) and H2O2 [1 mL (30%), 10 mmol] is added. The red solution is added all at once to an Erlenmeyer flask (250 mL) containing an aqueous solution of CsHCO3 (14 g, 70 mmol) in a total volume of 16 mL. The flask should be shaken continuously while adding the cobalt solution and afterward for another minute to ensure the catalytic decomposition of excess H2O2. The green solution is slowly added while shaking to another Erlenmeyer flask (250 mL) containing H2SO4 (25 mL, 5 M), and the formation of the greenish blue [Co(OH2)6]3+ can be observed. To ensure crystallization within 4–5 min, the supersaturated solution is seeded with a minute amount of CsCo(SO4)2⭈12H2O crystals. The product can then be isolated, for example in a basket centrifuge.
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Experiment 2, Ammonium Cobalt Alum in Ammonium Aluminum Alum CoSO4⭈7H2O (2 g, 7.1 mmol) is dissolved in water (16 mL), and H2O2 [1 mL (30%), 10 mmol] is added. The red solution is added all at once to an Erlenmeyer flask (500 mL) containing a suspension of KHCO3 (25 g, 250 mmol) in water (50 mL). The flask should be shaken while the solution is being added and afterward for another minute to ensure the catalytic decomposition of excess H2O2. The green solution is slowly added, while shaking, to an Erlenmeyer flask (1000 mL) containing sulfuric acid (200 mL, 5 M) saturated with aluminum sulfate, and the formation of the greenish blue [Co(OH2)6]3+ can be observed. Solid (NH4)2SO4 (2 g) is then added and the solution is shaken for 2 minutes, after which another portion of (NH4)2SO4 (8 g) is added. Soon light greenish-blue crystals of the mixed cobalt–aluminum alum can be observed at the bottom of the flask. By tilting the flask, the crystals can be seen from a distance of 25 feet. If a longer distance is required, a basket centrifuge will allow the crystals to be observed on the white background of the filter paper. Comments on the Demonstration Experiments If experiment 1 has not been performed in advance and no crystals of CsCo(SO4)2⭈12H2O are available, another (nonreducing) α alum can be used for the seeding. We have successfully used the two common alums KAl(SO4)2⭈12H2O and (NH4)Al(SO4)2⭈12H2O as well as the crystals from experiment 2. If Cs2CO3 instead of CsHCO3 is available, one can easily convert the former to the latter by bubbling CO2 through an aqueous solution of Cs2CO3. The conversion takes approximately 1 h and the buffer minimum (ca. pH = 8) should at least be reached. Otherwise there is a risk of basic cobalt(III) solids being formed. If only cesium sulfate is available, this salt can easily be transformed into the hydrogen carbonate by adding excess barium hydroxide to precipitate
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BaSO4 and then adding excess carbon dioxide to precipitate BaCO3. Hazards CoSO4 has been reported to be carcinogenic. Therefore, it will be necessary to consult the locally approved disposal procedures that should be employed for all wastes. Literature Cited 1. Marino, L. Z. Anorg. Chem. 1904, 42, 212. 2. Gajhede, M.; Simonsen, K.; Skov, L. K. Acta Chem. Scand. 1993, 47, 271–274. 3. Bernhard, P.; Ludi, A. Inorg. Chem. 1984, 23, 870–872. 4. Joensen, F.; Schäffer, C. E. Acta Chem. Scand. A 1984, 38, 819–820. 5. Mohr, F. Lehrbuch der chemisch-analytischen Titriermethode, 6th ed.; Braunschweig, 1886; p 207. 6. Brorson, M.; Gajhede, M. Inorg. Chem. 1987, 26, 2109–2112. 7. Brorson, M.; Schäffer, C. E. Acta Chem. Scand. A 1986, 40, 358–360. 8. Best, S. P.; Forsyth, J. B. J. Chem. Soc., Dalton Trans. 1990, 395–400. 9. Lipson, H. Nature 1935, 135, 912. 10. Beattie, J. K.; Best, S. P.; Skelton, B. W.; White, A. H. J. Chem. Soc., Dalton Trans. 1981, 2105–2111. 11. Armstrong, R. S.; Beattie, J. K.; Best, S. P.; Skelton, B. W.; White, A. H. J. Chem. Soc., Dalton Trans. 1983, 1973–1975. 12. Beattie, J. K.; Best, S. P.; Del Favero, P.; Skelton, B. W.; Sobolev, A. N.; White, A. H. J. Chem. Soc., Dalton Trans. 1996, 1481–1486. 13. Palmer, W. G. Experimental Inorganic Chemistry; Cambridge University Press: Cambridge, 1959; p 529. 14. Schäffer, S. C.; Schäffer, C. E. J. Chem. Educ. 1996, 73, 180–181. 15. Davies, G.; Watkins, K. O. J. Phys. Chem. 1970, 74, 3388–3392.
Journal of Chemical Education • Vol. 79 No. 8 August 2002 • JChemEd.chem.wisc.edu
