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C. J. HALLADA, G. A. TSIGDINOS, AND B. S. HUDSON
Molybdovanadophosphoric Acids and Their Salts. 11. investigation of Solution Properties by C. J. Hallada, G . A. Tsigdinos, and B. S. Hudson Climax Molybdenum Company of Michigan Research Laboratory, Ann Arbor, Michigan
48105
(Received J u n e 94, 1968)
Conductivity and pH measurements were made on the heteropoly acids H4 [PMoIIVO~O] and HS [PMoloVtO4o] at 25” in aqueous solution to establish the stability and acid strength of the compounds. The conductivity data were treated using the Onsager-Fuoss theory to find values of A0 and a. The data show that H4[PhfollV040] and H5[PMO~OVZO~O] are strong 1:4 and 1:5 electrolytes and that H5 [PMo~oVZO~O] is the more stable compound in solution. Hydrolytic instability of [PMosVaOto]precludes a firm conclusion about the strength of this acid.
Introduction Studies of the solution properties of heteropoly acids have been somewhat sparse despite the general interest in these compounds for many years. Deterrents to such studies have primarily been the instability of the compounds and the uncertainty of the composition of the compounds. The studies that have been performed have dealt primarily with heteropoly tungstates and have used polarographic, viscosity, light-scattering, and sedimentation-velocity techniq~esl-~ to determine the nature of the species in solution. Recently the preparation of a series of niolybdovaiiadophosphoric acids was reported,s and modifications of these preparative techniques were used in this laboratory to produce 1 3 4 [PPIoIIVO~O]~ 34H20, Ha [PMolOVz0401.32H20, and H6 [ P X ~ O ~ V ~3O4 ~HO~]o . I~n preliminary work, these acids appeared to be stable in aqueous solution; therefore, an attempt has been made to define better the stability and electrolytic nature of these acids in aqueous solutions. Conductance and pH measurements have been used to establish the hydrolytic stability, the electrolyte type of the acids, and the strength of the acids. The heteropoly acids as a group are known to be strong electrolytes, but neutralization leads to some degradation and makes it difficult to define the acid strength with certainty. The Onsager-Fuoss conductivity theory has been derived for symmetrical electrolyte^.^^^ For completely unassociated electrolytes the theory may be written in the form
-
A = Ao -
SC‘”+ EClog c
+ JC
where A is the measured equivalent conductance at concentration c , Ao is the equivalent conductance at infinite dilution, c is the concentration in equivalents per liter, S is the Onsager limiting law slope (a function of no,temperature, dielectric constant of the solvent, viscosity of the solvent, and valence type of the electrolyte), E is a constant defined by the same variables as T h e Journal of Physical Chemistry
S, and J is a constant defined by the same variables as S and E but including the closest distance of approach of the ions, d. To find values for Ao and J, values of
A‘ are calculated A‘
E il
-j-
SC‘” - EC log c
=
Ao
+ JC
Highly dissociated unsymmetrical electrolytes have been shown to behave in a manner similar to that predicted by the Onsager-Fuoss theory or an empirical equation of the same form.g Reasonable values of the conductivity at infinite dilution and the closest distance of approach have been found when the Onsager-Fuoss theory has been applied to data from unsymmetrical, strong electrolytes, although A’ us. c plots do begin to curve downward at higher concentrations.1° It has also been shown that acids obey the conductivity theory despite the different transport mechanism involved for protons.lOrll Therefore, conductivity data can help establish the stability, the electrolyte type, and the mobility of the anion. When association occurs, it is difficult to define the degree of association and the closest distance of approach of unsymmetrical electro(1) M. T. Pope and G.
N. Varga, Inorg. Chern., 5, 1249 (1966).
(2) T. Kurucsev, A. N. Sargeson, and B. 0. West, J . P h y s . Chem., 61, 1569 (1957).
(3) J. 6. Johnson, K. A. Kraus, and G. Scatchard, ibid., 64, 1697 (1960).
(4) M. C. Baker, P. A. Lyons, and S.J. Singer, J . A m e r . Chem. SOC., 77, 2011 (1955). (5) M. P. Courtin, F, Chauveau, and P. Souchay, Compt. Rend., 258, 1246 (1964).
(6) G. A. Tsigdinos and C. J. Hallada, Inorg. Chem., 7, 437 (1968).
(7) R. M. Fuoss and L. Onsager, J . P h y s . Chem., 61, 668 (1957). M, Fuoss and F. Accascina, “Electrolytic Conductance,” Interscience Publishers, New York, N. Y . , 1959. (9) H. S. Harned and B. B. Owen, “The Physical Chemistry of Electrolytic Solutions,” 3rd ed, Reinhold Publishing Corp., New York, N. Y., 1957, p 207. (10) G. Atkinson, M. Yokoi, and C. J. Hallada, J . A m e r . Chem. SOC., (8) R.
83, 1590 (1961). (11) H. 0. Spivey and T. Shedlovsky, J . P h y s . Chem., 71, 2165 (1967).
MOLYBDOVANADOPHOSOPHORIC ACIDSAND THEIRSALTS
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lytes from conductivity measurements. The associated species itself is conducting until a neutral species is formed-and this species is probably formed only in very concentrated solutions. However, pH data can help to clarify whether or not ion association is occurring. More information regarding the structure and chemical behavior of these compounds can be found in the previous papera6
Experimental Section The compounds used in this work were prepared as described previously.6 The number of waters of hydration can vary considerably in these compounds, but a series of analyses indicated that the chemical formulas were H4 [ P ~ ~ o ~ ~ V34Hz0, O ~ O ] Hg . [PMoI~VZO~O] 32H20, 350 and Hg[P%fogV3040].34Hz0 within two or three waters of hydration, This amount of uncertainty is well within the uncertainty attending the theoretical treat3251 I , , I , ment of data obtained on electrolytes of this type. 0 0.02 0.04 0.06 0.08 0.10 0.12 Attempts were made to prepare the E a + salts by ion CONCENTRATION EQUlV./I)1/2 exchange and by neutralization. Both techniques led Figure 1. Phoreograms for molybdovanadophosphoric to sufficient degradation of the anion so that the comacids. Straight lines are limiting law slopes. pounds could not be used for conductance work. The conductance was measured on solutions prepared trations, the A value for Hg[Ph!t09V304~]became quite from several stock solutions. The stock solutions of high, indicating an increased hydrogen ion concentraH4[ P ~ ! T . O ~ ~ V and O ~ H6 ~ ] [ P M O ~ ~ V were ~ O ~made ~ ] from tion due to hydrolysis. (The data on H ~ [ P M o s V ~ ~ ~ O ] two preparations of these two acids. The measureare not included in this paper.) ments were made by the weight dilution technique in a flask-type cell containing shiny platinum electrodes and Results and Discussion were calibrated with standard potassium chloride Conductivity Results. The conductivity data obsolutions. During the measurements, the solutions tained on H4 [P~LIOI~VO~O] and H5 [PMolbVz040]in water were maintained at 25.00 h 0.01", and concentrations are given in Table I, and phoreograms are shown in Figwere varied froin about 0.2 to 12 mequiv/l. Resistance measurements were made with an Industrial Instruments Model RC-18 conductance bridge at both 1 and Table I : Equivalent Conductance of 3 kc. However, since the frequency seemed to have Molybdovanadophosphoric Acids no effect, most measurements were made at 3 kc H a [PBl~iiVOaol-HE [PM~loVzOio]only. loa,, Potentiometric measurements of the H+ ion activity equiv/l. A equiv/l. A were made with a Beckman Model 101900 research p H 3.59 463.4 2.28 415.2 meter. The range of the data was limited by uncer5.39 449.9 4.44 400.2 tainties of the measurement at high acid concentra6.17 451.2 6.61 392.6 tions (pH
