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The system mercury oxide-boron trioxide-water
Notes 6. F. A. Cotton and C. S. Kraihanzel, J. Am. Chem. Soc. 84, 4432 (1962). 7. F. Gomrz Beltran, L. A. Oro, M. P. Puebla and R. de Val, Rev. Acad. Ciencias. Zaragoza 30, 99 (1975). 8. C. A. Tolman, Chem. Rev. 77, 313 (1977). 9. S. O. Grim, D. A. Wheatland and W. McFarlane, J. Am. Chem. Soc. 89, 5573 (1967).
421
10. S. O. Grim, P. R. McAllister and R. M. Singer, J. Chem. Soc. Chem. Commun. 38 (1969). 11. R. L. Keiter and J. G. Verkade, Inorg. Chem. 8. 2115 (1969). 12. B. E. Mann, C. Masters and B. L. Shaw, J. Chem. Soc. (At 1104 (1971),
1. inorg,nucl.Chem.Vol.43,p. 421 PergamonPressLtd..1981. Printedin GreatBritain
O022-190218110201A)42115C~2.0010
The system mercury oxide-boron trioxide-water (Received for publication 29 May 1980) The system HgO-B20~-H20 was briefly investigated by Orlova[l]. Aqueous solutions of Hg(NO3)2 and borax were mixed but only precipitates of basic mercury salts were obtained. This result was expected as it is a well known fact that only insoluble basic mercury salts could be obtained from slightly basic aqueous solutions. These compounds are Hg(OHG Hg2(OH)2 or HgO. In accordance with Orlova[l] and contrary to earlier literature our investigations have shown the same results. It was easy to show that the precipitates which were obtained from aqueous mercury salt-borax solutions did not contain boron. This was the reason why we worked with neutral and slightly acidic solutions. Solutions of soluble Hg(I)----and ill)--salts and H3BO3 were used. These mercury salts are HgCI2, Hg(NO3)2, Hg(OOCCH3)2 and Hg2(NO3)2. However, it was impossible to isolate mercury borates from these solutions. It is known that
larger concentrations of polyborate ions in aqueous solutions exist primarily between pH 6 and 9[2]. Unfortunately the available Hg(I)--and ill)---salts are very unstable in such media. It seems quite obvious from these results that it is impossible to isolate mercury borates from aqueous solutions. Institute of Inorganic and Analytical Chemistry RALF JANDA Free University of Berlin GERT HELLER D-IO00 Berlin 33, Germany
REFERENCES 1. O, D. Orlova and E. A. Gynner, Ukr Khim. Zh. 39, Nr. 124 (1973). 2. R. Janda and G. Heller, Z Natuffo~ch. 34b, 585 (1979L
J. inorg,nucl.Chem.Vol.43,pp. 421-422 © PergamonPressLtd.,1981. Printedin GreatBritain
0022-1~2/81/0201-042U$02.00]0
Study of the red peroxovanadium complex in acidic media by EPR (Received 2 November 1979; received for publication 8 February 1980) A large number of papers have been published on various aspects of the solution chemistry of Vanadium-peroxy complexes[I-6]. All of these investigations have assumed that the red species formed by Vanadium and HzO2 in an acidic media contains vanadium in its pentavalent state. Our EPR studies of these solutions present evidence, however, that tetravalent vanadium is responsible for the red species formed, an idea not hitherto reported. For many years, the formation of this red complex has been the most important spot test for the detection of vanadium because of its sensitivity. The limit of identification is 2.5 ~, vanadium with a limit of dilution of 1:20,00017]. Various workers have tried to characterize the complex; however, the equation currently accepted for representing the equilibrium involved with the red complex is as follows[l]: VOw.+ + H20: ~ VO(O.,)+ + H20 (red complex) Kj = 3.5 x liP. It is believed that H202 can then further combine to yield a 1 : 2 yellow complex with vanadium as seen in the following[l]: VO(O2)+ + H202~VO(O2)2 + 2H + (red species) (yellow species) 1:1 1:2
I=1.0
Ks=I.0
T=25°C.
One would, therefore, expect the formation of the yellow species to be more favorable at higher pH. Dean[6] has reported that there is some dependence upon the anion of the acid on the behavior of the complex, however the complex can be generally represented as above. We prepared the reddish complex using three different acids: HCI, H2504, and HCIOd. In each solution the concentration of total vanadium was approx. 0.1 M. Varying amounts of H~()2 was added to obtain the red color, in some cases up to 6%. All solutions were prepared at 0°C. In the HCIQ solution, it appeared qualitatively that the oxygen decomposition was much more vigorous than with the other acids. The HCI and H~SO4 vanadium peroxy solutions were quite stable. Each solution gave an absorbance maximum at 450 nm, the wavelength at which the Em~,,for this red complex has been attributed[l, 6]. An EPR study of each of the red solutions of different acids gave identical spectra. The spectrum obtained is shown in Fig. I. This spectrum is identical to that given by various tetravalent vanadium complexes[8a]. The value obtained for the g tensor is 1.97 which is in close accordance with that found in the literature for most tetravalent Vanadium complexes.[8b]. Each of the red solutions were then made alkaline until the red color just disappeared. The appearance of these solutions was yellow. The solutions exhibited no EPR signal. The action of hydrogen peroxide as a reducing agent in acidic
421
10. S. O. Grim, P. R. McAllister and R. M. Singer, J. Chem. Soc. Chem. Commun. 38 (1969). 11. R. L. Keiter and J. G. Verkade, Inorg. Chem. 8. 2115 (1969). 12. B. E. Mann, C. Masters and B. L. Shaw, J. Chem. Soc. (At 1104 (1971),
1. inorg,nucl.Chem.Vol.43,p. 421 PergamonPressLtd..1981. Printedin GreatBritain
O022-190218110201A)42115C~2.0010
The system mercury oxide-boron trioxide-water (Received for publication 29 May 1980) The system HgO-B20~-H20 was briefly investigated by Orlova[l]. Aqueous solutions of Hg(NO3)2 and borax were mixed but only precipitates of basic mercury salts were obtained. This result was expected as it is a well known fact that only insoluble basic mercury salts could be obtained from slightly basic aqueous solutions. These compounds are Hg(OHG Hg2(OH)2 or HgO. In accordance with Orlova[l] and contrary to earlier literature our investigations have shown the same results. It was easy to show that the precipitates which were obtained from aqueous mercury salt-borax solutions did not contain boron. This was the reason why we worked with neutral and slightly acidic solutions. Solutions of soluble Hg(I)----and ill)--salts and H3BO3 were used. These mercury salts are HgCI2, Hg(NO3)2, Hg(OOCCH3)2 and Hg2(NO3)2. However, it was impossible to isolate mercury borates from these solutions. It is known that
larger concentrations of polyborate ions in aqueous solutions exist primarily between pH 6 and 9[2]. Unfortunately the available Hg(I)--and ill)---salts are very unstable in such media. It seems quite obvious from these results that it is impossible to isolate mercury borates from aqueous solutions. Institute of Inorganic and Analytical Chemistry RALF JANDA Free University of Berlin GERT HELLER D-IO00 Berlin 33, Germany
REFERENCES 1. O, D. Orlova and E. A. Gynner, Ukr Khim. Zh. 39, Nr. 124 (1973). 2. R. Janda and G. Heller, Z Natuffo~ch. 34b, 585 (1979L
J. inorg,nucl.Chem.Vol.43,pp. 421-422 © PergamonPressLtd.,1981. Printedin GreatBritain
0022-1~2/81/0201-042U$02.00]0
Study of the red peroxovanadium complex in acidic media by EPR (Received 2 November 1979; received for publication 8 February 1980) A large number of papers have been published on various aspects of the solution chemistry of Vanadium-peroxy complexes[I-6]. All of these investigations have assumed that the red species formed by Vanadium and HzO2 in an acidic media contains vanadium in its pentavalent state. Our EPR studies of these solutions present evidence, however, that tetravalent vanadium is responsible for the red species formed, an idea not hitherto reported. For many years, the formation of this red complex has been the most important spot test for the detection of vanadium because of its sensitivity. The limit of identification is 2.5 ~, vanadium with a limit of dilution of 1:20,00017]. Various workers have tried to characterize the complex; however, the equation currently accepted for representing the equilibrium involved with the red complex is as follows[l]: VOw.+ + H20: ~ VO(O.,)+ + H20 (red complex) Kj = 3.5 x liP. It is believed that H202 can then further combine to yield a 1 : 2 yellow complex with vanadium as seen in the following[l]: VO(O2)+ + H202~VO(O2)2 + 2H + (red species) (yellow species) 1:1 1:2
I=1.0
Ks=I.0
T=25°C.
One would, therefore, expect the formation of the yellow species to be more favorable at higher pH. Dean[6] has reported that there is some dependence upon the anion of the acid on the behavior of the complex, however the complex can be generally represented as above. We prepared the reddish complex using three different acids: HCI, H2504, and HCIOd. In each solution the concentration of total vanadium was approx. 0.1 M. Varying amounts of H~()2 was added to obtain the red color, in some cases up to 6%. All solutions were prepared at 0°C. In the HCIQ solution, it appeared qualitatively that the oxygen decomposition was much more vigorous than with the other acids. The HCI and H~SO4 vanadium peroxy solutions were quite stable. Each solution gave an absorbance maximum at 450 nm, the wavelength at which the Em~,,for this red complex has been attributed[l, 6]. An EPR study of each of the red solutions of different acids gave identical spectra. The spectrum obtained is shown in Fig. I. This spectrum is identical to that given by various tetravalent vanadium complexes[8a]. The value obtained for the g tensor is 1.97 which is in close accordance with that found in the literature for most tetravalent Vanadium complexes.[8b]. Each of the red solutions were then made alkaline until the red color just disappeared. The appearance of these solutions was yellow. The solutions exhibited no EPR signal. The action of hydrogen peroxide as a reducing agent in acidic