(11, which would have required an O’Zn ratio of 4.0. The BIPY and P H E N complexes have rather similar spectral curves, showing a pronounced shoulder with maximum absorbance a t about 425 mp. In a single test, successive increments of EtnZn were added to 25 ml. of a toluene solution containing excess P H E N plus enough EtzZn to give a definite color. The plot of absorbance us. EtzZn added was linear over the range tested (0.4 to 3 mg. of Et2Zn) with an indicated absorptivity of about 14 liters/gram-cm. Spectrophotometric determination of EtrZn in this range should present no difficulty. Thermometric titration with P H E N provides a simple, rapid, and precise method for determining the net EtPZn content of a solution which may also contain its oxidation or hydrolysis
products. Titration with Oxine gives a precise measure of the total zinc content in Et2Zn solutions which contain its oxidation products, but is not reliable if hydrolysis products are present. Under favorable circumstances oxine gives, in the same titration, a measure of both total zinc and EtzZn. The use of oxine as titrant in a hydrocarbon system may be of interest as a possible means for the analysis of other organometallic compounds. ACKNOWLEDGMENT
The authors thank John Boor, Jr., for samples of purified diethylzinc and for helpful suggestions and encouragement. LITERATURE CITED
(1) Abraham, M. H., Chem. Ind. (London) 1959,p. 750.
(2) Bamford, C. H., Newitt, D. M,, J . Chem. SOC.1946, p. 688. (3) Bock, H., 2. “Vaturforsch 17b, 426 (1962). (4) Coate;; G. E., “Organometallic Compounds, 2nd ed., Wiley, New York, 1960. (5) Everson, W. L., ANAL. CHEM.36, 854 (1964). (6) Everson, W. L., Ramirez, E. >I., Ibzd., 37, 806 (1965). (7) Haurowitz, F., Mikrochemie 7, 88 (1929). (8) Herold, R. J., Aggarwal, S.L., Neff, W., Can. J . Chem. 41, 1368 (1963). (9) Novak, K., Chem. Prumysl 12, 551 (1962). (10) Pajaro, G., Biagini, S., Fiumani, D., Angew. Chem. (Int.. Ed.) 2, 94 (1963). (11) Thompson, H. W., Kelland, N . S., J . Chem. SOC.1933, p. 746. (12) S’ogel, C. H., Monatsber. Deut. A k a d . Wiss. Berlin 2, 115 (1960). RECEIVEDfor review January 8, 1965. Accepted ;\larch 29, 1965.
Polarographic Behavior of Molybdenum(V1) in Acidic Chloride Media J. J. WITTICK’ and G. A. RECHNITZ Department o f Chemistry, University of Pennsylvania, Philadelphia, Pa.
b The polarographic behavior of molybdenum VI has been investigated in the concentration range of 1 10-3 to 5 I O - 6 ~ M~(VII in supporting electrolytes of widely varying hydrogen ion and chloride ion concentrations. On the basis of polarographic, coulometric, and spectrophotometric evidence, the three waves present in dilute hydrochloric acid are ascribed to the reduction of two species of molybdenum(V1) in slow equilibrium, to produce molybdenum(V) and (111) on the first and third waves, on the one hand, and molybdenum(1V) at potentials corresponding to the second wave, on the other. A reversible, 2-electron wave for the reduction of molybdenum(V), formed by reoxidation of molybdenum(lll), is reported and discussed in critical detail along with other secondary reactions resulting from the disproportionation of molybdenum(lV) and the reoxidation of molybdenum(lll), respectively.
x
x
portion of our knowledge of the chemistry of molybdenum in hydrochloric acid rests upon conclusions drawn from polarographic data. Thus, the catalytic reduction of perchlorate ( I ?‘), nitrate (16), and oxalate (34) in the presence of molybdenum is attributed to certain oxidation states of molybdenum on the basis of SUBSTANTIAL
816
ANALYTICAL CHEMISTRY
polarographic investigations which demonstrate the enhancement of characteristic reduction waves. Polarography has also been employed as the basis of arguments concerning the nature of the isopoly acids which are characteristic of Mo(V1) in acidic media (12, 32) and to explain the complex and controversial color changes which accompany the quinquevalent and trivalent oxidation states (14, 21) in solution. I n all of these diagnostic attacks upon the structural or mechanistic problems of molybdenum chemistry, it is obvious that a reliable interpretation of the polarographic behavior of molybdenum(V1) is essential. Attempts to interpret the polarographic waves exhibited by Mo(V1) in hydrochloric acid began in 1941 with a rudimentary examination by Holtje and Geyer (20). Carritt (6) used controlled potential electrolysis to identify polarographic reduction products and postulated the existence of two species of Mo(V1) in sluggish equilibrium to account for the observed polarographic waves. Although his work was done a t relatively high concentrations of molybdenum where serious problems with maxima were encountered, he furnished the first clues to the difficulties inherent in molybdenum polarography. h comprehensive re-examination of the system was subsequently reported by Guibe and Souchay (14) and included polarographic data on the lower oxidation
states of molybdenum in partial contradiction to the results of Carritt. H o w ever, Guibe and Souchay gave no data regarding capillary characteristics or maximum suppressors, and their conclusions based on relative wave heights were unsupported by any of the usual tests for diffusion control. In a recent investigation, Haight ( I 7 ) concluded that one of the characteristic reduction waves of Mo(V1) involves the production of Mo(1V). Both Carritt and Guibe attributed this wave to the reduction of Mo(V) to Mo(IIIj, however. The present investigation represents an attempt to furnish a definitive explanation for the polarographic behavior of Jlo(V1) and thereby to resolve the above differences by the use of controlled potential coulometry and through the extension of measurements to new concentration ranges under conditions which permit the isolation of critical experimental variables. EXPERIMENTAL
Apparatus. A Sargent Model XV Polarograph, having a pen speed of 1 second for full-scale deflection and equipped with a micro range extender for high sensitivity work, was used to record all polarograms. Maximum suppressors, including agar, have been shown to distort the Present address, Merck, Sharp & Dohme Research Laboratories, Rahway, N. J.
"7.'.
r d t bridge from reference electrode
+'
H 03 V
fb-
t-I 1
Figure 2. Over-all polarographic behavior of Mo(VI) in acidic chloride media ot 30' C. D-
Figure 1 .
Polarographic cell
polarographic waves of hlo(V1) (6) and were, therefore, avoided during this investigation. The polarographic cell, shown in Figure 1, was designed to permit complete and convenient deaeration of the test solutions with minimal contamination. Gas inlets A and B were supplied through a Y connection from the same source of vanadium (I I)-scrubbed , presaturated nitrogen. With the cell empty, inlet B was clamped off and the stopper a t C was put in place. The test solution was placed in the dropping mercury electrode (D.M.E.) compartment and deaerated for 15 to 20 minutes. Nitrogen pressure prevented the solution from passing through t,he sintered-glass frit into the side arm. After the oxygen was removed, inlet B was opened, allowing the t'est solution to flow up the side arm until the levels were the same in both sides of the cell. The stopper a t C was then removed, and the D.5L.E. and the salt bridge from the reference electrode were inserted as shown. X flexible polyethylene tube filled with KCI-saturated agar was used as the salt bridge from a saturated calomel reference electrode (S.C.E.). The resistance of the completed cell, as measured with a Serfass h C bridge, was 1900 ohms. Mercury from the D.M.E. was collected in the bulge a t D to avoid blocking the sintered-glass frit. The cell proved sturdy, convenient to use, and easy to clean. A three-compartment cell, similar to that described by Meites (28), was used for controlled potential electrolysis. The mercury pool working electrode had a n area of approximately 50 sq. em. and was separated from the platinum foil counterelectrode by two sintered-glass frit,s to minimize diffusion. h Leeds and Northrup Model 1199-31 S.C.E., the tip of which was positioned close to the surface of the mercury working elect,rode, was used as the reference elect,rode. Provision was made to bubble oxygen-free nitrogen through the solution in the work electrode compartment during elec-
Origins for all curves morked a t 0 volt V.I S.C.E. and 0 p a , H .t concentrations, molar A, B, C, D, E. 5.0, 2.5, 1.0, 0.5, 0.1, respectively CI - concentrations, molar F, G, H, I , 1. 5.0, 2.5, 1.0, 0.5, 0.1, respectively Mo(VI) concentration. 2.5 X 10 -4M
trolysis; stirring was accomplished mechanically from above. The potential of the working electrode was controlled by a N7enking Model 6431 R potentiostat and was continuously monitored by a Wenking Model PRV 12 electrometer. The observed electrolysis current was monitored by recording the I R drop across a precision resistor with a Photovolt Model 43 recorder. A Perkin-Elmer Model 202 spectrophotometer was used for all optical measurements. Reagents. Reagent grade chemicals and conductivity water were used throughout. Potentiometric titrations were carried out using G . F. Smith standardized ceric sulfate in 1 5 HzS04. A stock solution, 0.50.1.1 in Mo(VI), was prepared from reagent grade NazMo04 2 H 2 0 by dissolving a n appropriate amount of the solid in conductivity water. The solution was standardized both by the classical FbXlOO4 method (19) and by reduction to Mo(II1) with mercury in concentrated HC1 and subsequent oxidimetric titration with ceric sulfate ( I f , $3). Other Mo(V1) solutions were prepared from the stock solutions. Sodium chloride was used t'o adjust the chloride ion concentration. Procedure. .ill polarographic and controlled potential experiments were carried out a t 30.0' + 0.1' C., except where otherwise noted. T h e polarograms are maximum current, i, envelopes corrected for residual current and also corrected for JR drop where necessary. No R C damping was used. At the higher sensitivities, a linear compensating voltage was applied with the micro range estender attachment to keep the entire polarogram on scale. The capillary used had a drop time, t, of 5.14 secondsand a n m value of 1.48 mg. set.-' (m2/3t1'B = 1.70 mg. sec.-1'2) measured at -0.80 volt us. S.C.E. in 0.1M HCI and a t an effective mercury column height, h, of 58.5 cm. Half-wave potentials,
were not measured to better than +10 mv. because of the poor definition of the polarographic waves. For the same reason, limiting currents, ill were somewhat arbitrarily defined. The controlled potential electrolyses were accomplished by pre-electrolyzing 145 ml. of supporting electrolyte and then adding 5.0 ml. of deaerated molybdate solution. During all operations involving reduced or deaerated solutions, great care was taken to prevent contamination by atmospheric oxygen. To obtain the net electrolysis current the observed current values were corrected for the pre-electrolysis or background current. The current-time curves were integrated graphically to obtain the number of coulombs passed during the electrolysis. RESULTS
Over-All Polarographic Behavior of Mo(VI). Polarograms were obtained
for concentrations of Mo(V1) from 1 X lO-3M to 5 X 10-5X and, in a few cases, to 5 X 10-6J.1. The H + and CIconcentrations were both varied from 0 . l N to 5.0M; the C1- concentration could not be varied independently at high H + concentrations because the anions of the other common strong acids are either reduced catalytically or change the nature of the polarograms through complexation. Typical polarograms in the various supporting electrolytes are given in Figure 2 for a Mo(V1) concentration of 2.5 X lO-4M. Figure 3 shows the effect of varying the Mo(V1) concentration a t constant 0.1JI HC1. The relative wave heights of waves I and I11 decrease as the HC1 or C1- concentration increases and as the Slo(V1) concentration decreases; the opposite effect is observed for wave 11. At a hlo(V1) concentration of 5 X 10-5~.1in VOL. 37, NO. 7, JUNE 1965
81 7
0.1M HC1, waves I and IT1 have all but disappeared; and, if the C1- concentration is increased to 1.OJ4, these two waves vanish even at hlo(V1) concentrations as high as 1 X 10-5M. At moderate HCl and/or C1- concentra-, tions, maxima are observed in the polarograms for the higher Mo(V1) concentrations. Distortions, probably due to incipient maxima formation and/or to the presence of polymeric species, begin to appear in the region between waves I and I1 a t Mo(V1) concentrations of 1 x 10-~1.;rin 0.1M HCI, and wave I11 is almost completely masked by the hydrogen discharge wave. Since wave I merges with the anodic mercury wave, its ElI2 cannot be measured directly. The Ellzof wave I1 becomes more positive as the HCl or C1- concentration is increased, until the wave also merges with the anodic wave of mercury. Table I lists those values of the Ell2 of wave I1 which can be estimated with some reliability. The Ell2 of wave I11 shifts in a more complicated manner with changes in supporting electrolyte (see Table 11). A few values of the E , / 2of electrolytically prepared Mo(V) are included for comparison and are discussed below.
u -04 POTENTIAL
-06 -08 -10
-02
' 0
VI
S C E ,VOLT
Figure 3. Effect of concentration on polarograms of Mo(VI) in 0.1M HCI at
30" C. Mo(VI) concentrations, molar A. 5 x 10-4 B. 1 x 10-4 C. 2.5 X D. 1 x 1 0 - 6
E. 5
x
10-6
Table I.
HCl concentration,
M
0.1
Half-Wave Potentials of Wave II at Volt us. S.C.E.
30" C.
C1- concentration, M 0.1
0.5
-0.23
-0.20 -0.12
0.5 1.0 2.5 5.0
1.0
-0.18 -0.11
-0.08
2.5 -0.14
5.0 -0,OFP
-0.08 -0.07 -0.06'
-0.06a -0.050 b b
Approximate values due to proximity of anodic mercury wave Wave I1 merged with anodic mercury wave.
a
b
Table
II.
Half-Wave Potentials of Wave 111 and Mo(V) at 1 denum and at 30" C.
X 10-4M"Molyb-
Volt us. S.C.E. HCI. concentration, M 0.1 0.5
1.0 2.5
5.0
PoIarographic wave Wave I11 MO(V)
C1- concentration, M 0.1
0.5
1.0
-0.66 -0.66
-0.70
-0.71 -0.68 -0.60a
-0.52"
-0.550
-0.5a
Wave I11 Mo(V) Wave I11 Mo(V) Wave I11
-0.64 -0.62
2.5 -0.6Sa
5.0 -0.54a -0.50 b
b
-0.34
-0.51 -0.4O
Mo(V) Wave I11 MO(l-)
a
b c
Obtained at 1 X 10-S.M Mo(V1) due to poor wave definition at 1 X 10-4M. Wave I11 no longer present. Wave for hIo(T') merged with anodic mercury wave.
81 8
ANALYTICAL CHEMISTRY
b b c
Effect of Mercury Column Height and Temperature. The effect of
mercury column height on the limiting current for each of the three waves was determined in 0.1M HCl a t a Mo(V1) concentration of 1 X 10-4M for values of h between 28.5 and 88.5 cm. The limiting current for wave I was measured a t -0.10 volt, that for wave I1 was taken as i i a t -0.40 volt minus ii a t -0.10 volt, and that for wave I11 was taken as il a t -0.80 volt minus il a t -0.40 volt. The selection of -0.15 volt for il on wave I and/or -0.30 or -0.50 volt for ii on wave I1 did not affect the results. The values of the exponent, z, in the relationship:
ii = kh.
(1)
where k is a proportionality constant (%), were found to be 0.35, 0.53, and 0.27 for waves I, 11, and 111, In other supporting respectively. electrolytes the values for waves I and I11 decreased somewhat. The linear plots of il us. h1I2 passed through the origin only for wave 11. The temperature coefficients of the limiting currents (ZZ), measured between 3" and 47" C., were 1.8, 0.9, and 4.0 % per degree centigrade for waves I, 11, and 111, respectively, in 0.l.V HCl. These results indicate partial kinetic control for waves I and I11 and diffusion control for wave 11. No significant change in the Ellz of wave I1 was observed in the temperature range covered, while wave I11 shifted approsimately 50 mv. in a positive direction for the same temperature increase. The changes brought about by the temperature variation were reversible, and the polarograms were unchanged when solutions were allowed to stand for an hour with the D.M.E. immersed (mercury flowing) a t either 3" or 47" C. Tests for Adsorption. The "swayback" appearance of wave 11 gave rise to the possibility that some adsorption process is involved in the electrode mechanism of this wave (29). Although the application of Equation 1 gave evidence for diffusion control (z = 0.5) rather than adsorption control (z = l), two other tests were carried out to show the absence of adsorption phenomena. First, plots of drop time us. potential of the D.M.E. ( E d e )were made for 1 X 10-4Jf Mo(V1) in 0.1M HCl and for 0.1M HC1 alone and were found to be not significantly different (29). Second, there was no indication of abnormal current-time behavior for individual mercury drops a t potentials corresponding to wave I1 (29). Since the polarograph had a relatively fast pen speed and was used without RC damping, distortions in the currenttime curves for the individual drops would have been a t least qualitatively evident.
O
a
i
P
A
- a5
POTENTIAL
VI.
L -10
S C.E.. VOLT
-024-
-2.0
-LO
0
+IO
+20
Figure 4. Application of fundamental wave equation to wave I1 with three waves present 0.1 M HCI, 5 X 10
Mo(VI)
Application of Fundamental Wave Equation. The fundamental equation
for polarographic waves a t 30" C. (25) :
where n is the electron change and id is the maximum diffusion current, was applied to wave I1 under the solution conditions, where all three waves were present and where wave I had completely vanished. Figures 4 and 5 show the pertinent polarograms and the EI,us. log ( i / i d - i) plots. I n both cases the agreement of the experimental points with the theoretical line for a reversible, 2-electron process is excellent.
logarithm of the electrolysis current, i,, is plotted against the electrolysis time, t,, shows the course of the electrolysis. For comparison, the linear log i, us. t, plot, which would be expected (8, 26) for a 1-electron, diffusion-controlled process under the experimental conditions used here, is also shown in Figure 6 3 . From a knowledge of the slope of such plots and the value of log i, a t t, = 0, n can be calculated (8, 27). Using this technique, as well as that of graphical integration under the i, - t, curve, n was found to have a value of 1, corresponding to the reduction of Mo(V1) to Mo(V). Titration of an aliquot of the product solution and comparison of the polarogram (14) and spectrum (2, 16) obtained after reduction with literature information confirmed Mo(V) as the reduction product. Wave 11, as well as wave I, disappears during the electrolysis even though the species responsible for wave I1 is electroinactive a t -0.10 volt. WAVE11. Almost immediately after the beginning of bulk electrolysis a t -0.40 volt (plateau of wave 11) in 0.1 M HCl, a film formed on the mercury electrode surface, rapidly became thicker, and was finally dispersed throughout the solution. The film formation caused irregularities and noise in the current-time curve and, therefore, cast doubt upon the validity of the values of n (1.5 to 2.3) which were obtained. The current decayed, nevertheless, a t a rate approximating that of a diffusion-controlled process. After filtering off the precipitate, analysis of the reduced solution indicated that 35 to #yo of the Mo added initially was present in solution as Mo(V) and that the remaining 60 to 65y0 had precipitated. These percentages did not vary more than a few
Half-Wave Potentials of Mo(V).
The half-wave potentials given in Table I1 for Mo(V) were obtained by' reduction of Mo(V1) a t -0.10 volt, in 0.1M HCl followed by dilution with the appropriate, oxygen-free electrolyte, or by direct electrolysis on wave I (see below) in supporting electrolytes where this was possible. Application of Equation 1 to the wave for Mo(V) in 0.1M HCl showed partial kinetic control similar to t h a t of wave 111. Controlled Potential
0
-02
-0.4
- o'22
-0.8 -1.0
r
Electrolysis.
WAVE I. The results of controlled potential electrolyses a t -0.10 volt (on wave I) in 0.1M HCl are explained by reference to Figure 6. Similar results were obtained for other supporting electrolytes in which wave I appeared. Figure 6,A and C, show, respectively, the polarograins and the ultraviolet absorption spectra of the starting solution and the solution after completion of the electrolysis; and Figure 6,B, where the
-0.6
POTENTIAL v t S C E .VOLT
N E I = -0.17VOLT
g
-o.141
- .-
ni:, -2.0
//
,
-10
,
0
LOG
,
,
t 1.0
, t20
(h)
Figure 5. Application of fundamental wave equation to wave II in absence of wave I 0.1M H+, 1 . O M C I - , 1 X 1 O 4 M M o I V I )
POTENTUL n S.C E.. VOLT
4 2
Wavelenplh, m y
Figure 6. Controlled potential reduction on wave I in
0.1M HCI Mo(VI) concentrotion, 5 X 10 -'M
per cent (relative) even though the Mo(V1) concentration was changed from lO-3M to 10-5M. The amount of precipitate formed decreased, however, as either the HCl or C1- concentration was increased, only a trace being found if the C1- concentration was increased to 1.OM and none a t all in 0.5M HCl. A t supporting electrolyte concentrations greater than these and a t reduction potentials positive enough to avoid the foot of wave 111-Le., a t -0.40 volt in 5.OM C1-, 0.1M H+complete conversion of Mo(V1) to Mo(V) took place, and a linear log i, us. t, plot was obtained corresponding to a value of n = 1. Larger samples of the precipitate were obtained by successive electrolysis of lO-3M Mo(V1) solutions in 0.1M HCI, and portions were analyzed for AMocontent by oxidation to Mo(V1) with HNOa and precipitation as PbMOO4. The average oxidation state was determined by oxidation 'with excess Ce(1V) and subsequent titration of the unreacted Ce(1V) with standardized Fe(I1). The results showed that the precipitate consisted of 56y0 Mo with an average oxidation state of +4.7. A spectrographic analysis detected no mercury in the precipitate, which was shown to be amorphous by x-ray techniques. WAVE111. Reduction on wave I11 was carried out a t -0.80 volt in 0.1M HC1 and yielded the typical results shown in Figure 7. The electrolysis current (Figure 7,B) decayed rapidly a t first but became almost constant after an hour and decreased only very slowly thereafter. Taking the almost constant current as the background current, a linear log i, us. t, plot corresponding to n = 1 was obtained for VOL.
37, NO. 7, JUNE 1 9 6 5
819
MO(VI)
/
1.0 -
Mo MO
POTENTIL VI. S.C E., VOLT
U i
+-
I
05 0.5 0
W
0:
05
V
/ [ 0 I -- I 5 ,
0
-02
-04
I
-06
I
,
-0 8
/
-10
POTENTIAL v s S C E , VOLT
Figure 8. Polarograms obtained during controlled potential oxidation of Mo(lll) in 0.1M HCI A.
Mo(illJ solution before oxidation 8. After 24 hours continuous oxidation a t -0.37 volt C. After 4 8 hours continuous oxidation a t -0.37 volt Mo concentration, 5 X 10 -4M.
Wavelength, m p
Figure 7. Controlled potential reduction on wave 111 in 0.1M HCI Mo(VIJconcentration, 5 X 10 -4M
the initial portion of the curve. The spectrum, polarogram, and a Ce(1V) titration of an aliquot of the solution confirmed the complete conversion of Mo(V1) to llo(V) a t this point in the electrolysis. After 24 hours of continuous electrolysis, the polarogram and spectrum shown in Figure 7, A and C, were obtained. A Ce(IV) titration established the product a t this time as 110(111). The two anodic waves-Le., a small wave a t about -0.44 volt and a larger but incompletely developed wave a t about -0.27 volt-were both found to be diffusion-controlled according to their dependence on the mercury column height. There was no difference between the behavior of ;\Io(V) produced by partial electrolysis on wave I11 and that produced by total electrolysis on wave I. Essentially the same results were obtained as the C1- concentration increased from O . l J 1 , except that the Ellz values of both anodic hlo(II1) waves became more positive and the height of the wave at -0.44 volt decreased, while the height of the other wave increased until, in 5.0.21 C1- and 0.1M H+, a single wave, having an E l z of -0.17 volt, was observed. In 5M HC1, a t potentials more negative than -0.7 volt, extensive reduction of hydrogen ion occurs; but a t all more positive potentials, electrolysis proceeds a t rates approaching diffusion control, first to ?rIo(V) and finally to ~ l o ( 1 1 1 ) . The log i, us. t, plots were curved and no anodic waves were observed for SIo(II1). Reoxidation of Mo(II1). Solutions of electrolytically prepared hlo(II1) in 0.1.21 HC1 were slowly reoxidized t o M o ( V ) , probably by reaction with 820
o
ANALYTICAL CHEMISTRY
H+, if stored under oxygen-free conditions out of contact with mercury. The absorption spectra of such solutions were characterized by the appearance and growth of the Mo(V) absorption maxima a t 256 and 295 mp upon standing. The corresponding polarograms showed a mixed cathodicanodic wave a t -0.44 volt rather than the completely anodic wave of Mo(II1). A similar, but slower, oxidation took place in 5M HC1; in 0.1M H f , 8M C1(LiCl), on the other hand, no reaction was observed even after a month a t room temperature. Mo(II1) solutions which were prepared chemically by reduction of hIo(V1) with mercury in 10M HC1 and which were then diluted to 0.1V HC1 with oxygen-free water, behaved like the electrolytically prepared solutions, except that they showed an apparent induction period before the onset of oxidation. I t would appear, then, that the relatively stable, red lIo(II1) species present a t high C1concentrations, generally considered to be hl0Cl6-~( I S , I8),is slowly hydrolyzed to a more reactive species which can be oxidized by hydrogen ion a t low C1concentrations. The possible existence of a reversible wave a t -0.44 volt for the Mo(V)/ Mo(II1) couple and its similarity to a wave found by Guibe and Souchay (14) and attributed to the reduction of Mo(1V) to l I o ( I I I ) , were investigated by carrying out a controlled potential oxidation of Mo(II1) in 0.1M HC1 at -0.37 volt (see Figure 8). Figure 8, A , B , and C, shows, respectively, the polarograms obtained before oxidation began, after 24 hours of oxidation, and after 48 hours. The electrolysis current was small and, except during the initial phase, decreased very slowly during the experiment. Changing the potential to -0.05 volt did not increase the rate of electrolysis significantly and resulted simply in an additional in-
LOG
(&) d-
Figure 9. Application of fundamental wave equation to polarographic wave produced on controlled potential oxidation of Mo(lll) in 0.1M HCI
crease in the height of the two cathodic waves shown in Figure 8, C. A plot of Ede us. log (i/& -i) for the cathodic wave at Ellz = -0.44 volt (in Figure 8, C) is shown in Figure 9 and indicates a reversible, 2-electron process. Diffusion control was established according to the mercury column height dependence of the limiting current. The same two cathodic waves were obtained upon exposure of electrolytically prepared Mo(II1) solutions to atmospheric oxygen, except that the relative heights of the waves were reversed (the more negative wave being the larger). The second wave resembled that of electrolytically prepared Mo(V), but its EliLa t about -0.70 volt is slightly more negative. The slow current decay during the electrolysis and the very poor definition of the anodic waves of Figure 8, B and C, point to the existence of a number of species in very slow equilibrium. The species responsible for the anodic wave a t -0.44 volt, which must also be the species oxidized by H+, accounts for approximately 10% of the total Mo(III), considering the wave height and the number of coulombs passed during the initial phase of the electrolysis. Mo-
lecular oxygen apparently oxidizes additional Mo(II1) species to produce the cathodic wave a t -0.70 volt; the product, in turn, must represent a more stable form of Mo(V) to which the hfo(V) species responsible for the wave at -0.44 volt is slowly converted. DISCUSSION AND CONCLUSIONS
The evidence given above leads to the conclusion that waves I and I1 represent the reduction of two species of Mo(V1) which are in sluggish equilibrium (6). On wave I, one of the species-say, Mo(V1)A--is reduced to Mo (V). Bard and Solon (3) have pointed out that a slow reduction at controlled potential (Figure 6, B ) , which gives the correct value for n, may indicate that the electroactive species is involved in a sluggish equilibrium. The partial kinetic character of wave I and the concurrent disappearance of wave I1 upon electrolysis on wave I strongly point to the same explanation. Kraus, Yelson, and Moore (25) have shown by their ion exchange experiments that, even at low concentrations of Mo(VI), two species exist in slow equilibrium a t HC1 concentrations below 1M. Wave I1 results from the reversible reduction of the second Mo(V1) species in sluggish equilibrium-say, Mo(V1)B to Mo(1V)-which is known to undergo rapid disproportionation to Mo(V) and Mo(II1) (4). The observed diffusion control of wave I1 is expected since, a t potentials on the plateau of this wave, both Mo(V1)A and Mo(V1)B diffuse to the electrode surface and are immediately reduced without significant relative changes in their equilibrium concentrations. The reversibility of the wave is established by the E d e -log (i/id - i) plots (Figures 4 and 5) and the constancy of its El,, with changes in temperature (24). That the wave represents the reduction of Mo(V1) rather than of Mo(V), as had been postulated by Guibe and Souchay (14), Carritt (6), and others (10, %0),is a necessary conclusion from the fact that there is no preceding wave a t low concentrations of molybdenum. Haight (17) based a similar conclusion on his observation, not further substantiated in his published work, that only one wave exists in 0.3M HCl a t a Mo(V1) concentration near 10-4M. It appears reasonable to attribute wave I11 to the reduction of’Mo(V) to Mo(III), because the of the wave is identical to that for the reduction of Mo(V) in 0.1M HC1 and varies in essentially the same manner as that of Mo(V) in other supporting electrolytes (Table 11). The slightly more negative Eliz values shown by wave I11 are probably due to differences in the nature of Mo(V) as it exists in the bulk of the solution and a t the electrode surface immediately after reduction from Mo
(VI). The relatively large shift in the
Ellzvalues with changes in temperature, the absence of an anodic wave a t the same Ellz for the reduction product, and the drawn-out appearance of wave I11 indicate its irreversibility. The problem of deciding which Mo(V1) species is responsible for wave I11 may be resolved by considering the behavior of the reduction products of waves I and 11. The Mo(V) reduced a t potentials on wave I11 can be supplied only by direct reduction from Mo(V1) or by secondary reactions involving the disproportionation of Mo (IV). However, if the disproportionation of the Mo(1V) produced on wave I1 is fast enough t o be polarographically significant, its effects should have been apparent on wave I1 itself, since the resulting Mo(II1) would have been reoxidized electrochemically. Since this was not the case, Mo(VI)A, the species responsible for wave I, must be the ultimate source of the Mo(V) which is reduced on wave 111. The similar variations observed in the relative wave heights of waves I and I11 under changing conditions are not inconsistent with this conclusion, but the partial kinetic character of the waves prevents any direct correlation with the amount of Mo(V1)A present in solution. There are two possible causes for the partial kinetic character of wave 111, both of which involve regeneration of Mo(V). The reoxidation of Mo(II1) by H + is probably the cause of the partial kinetic control observed for the polarographic wave of Mo(V) itself and would, therefore, contribute to that of wave 111. Also, Mo(II1) is known to react almost instantaneously with Mo (VI) in dilute HC1 and a t a somewhat slower rate in more concentrated acid to produce Mo(V) (4,6). The results of controlled potential electrolyses on waves I1 and I11 may now be explained in the light of the above conclusions. Of the two disproportionation products of the Mo(1V) produced by reduction on wave 11, Mo (V) appears to be stable; while Mo(II1) can react in several ways, all of which lead to Mo(V). Thus, Mo(II1) has been shown to be reoxidized by Mo(VI), by H+, and even electrochemically a t potentials more positive than the first anodic wave of Mo(II1). While all of these processes can occur in dilute HC1, provided that the Mo(II1) formed by disproportionation is the active species, only the first process would be important a t high C1- concentrations. If the reoxidation process is fast compared to diffusion, a diffusion-controlled reduction to Mo(V) would be expected for electrolysis on wave 11. This is exactly what is observed in supporting electrolytes where no precipitate forms and in which the Xlo(V) reduction wave does not interfere-Le., in 5M C1-, 0.1M H +
a t -0.40 volt. In 0.1M HC1, however, the reoxidation process is interfered with by precipitate formation. Making the reasonable assumption that only two species are involved, the average oxidation state of this precipitate, +4.7, indicates that one of the species is either Mo(V1) or Mo(V), while the other is either Mo(1V) or Mo(II1). Mo(II1) does not appear to be a likely component of the precipitate, since it is oxidized by Mo(V1) and can co-exist with Mo(V) in this solvent with no apparent interaction. It seems likely, therefore, that Mo(1V) is one of the components of the precipitate. Since no precipitate forms if electrolysis is carried out on wave 111, the other component must be one which is present a t the potentials of wave I1 but not a t those of wave 111. On wave 11, the reactions
+ eMo(VI)B + 2eMo(V1)A
F?
Mo(V)
(3)
Mo(IV)
(4)
occur a t the electrode surface. On wave 111, since there is no indication that Mo(1V) is further reduced a t any attainable potential, the following reactions occur:
+ 3eMo(V1)B + 2e-
Mo(V1)A
F?
Mo(II1)
(5)
Mo(1V)
(4)
the disproportionation products of Mo (IV) being present a t the potentials of both waves. The only species missing on wave I11 is seen to be the species of Mo(V) produced by reduction of Mo (V1)A; thus, this may be the other component of the precipitate. Controlled potential reductions on wave I11 and on the single wave found in 5M HC1 reflect the effects of Mo(V) regeneration in the nature of their log i, us. t, plots and in the complete reduction of Mo(V1) to Mo(V) before the appearance of any Mo(II1). In the case of wave 111, all the Mo(II1) formed is rapidly reoxidized by Mo(V1) until the Mo(V1) is completely reduced; only then does Mo(II1) appear in detectable amounts. This sequence of processes, and the slower reoxidation of Mo(II1) by hydrogen ion, give rise to the results of Figure 7, B. Essentially the same processes take place in 5M HC1, except that the reactions with Mo(V1) and hydrogen ion are retarded so that the net production of Mo(II1) proceeds a t a rate which is close to diffusion control but which nevertheless causes curvature in the log i, us. t. plots, Species of Molybdenum VI in Acidic Chloride Media. Although agreement is not complete on this point, it seems t h a t , a t Concentrations low enough to avoid polyacid formation (below lop3or 10-4LLf)1the molybdate anion, Mooa,+ is converted to monomeric molybdic acid, HzMoOa,both VOL. 37, NO. 7, JUNE 1965
0
821
hydrogens of which are reported to be of about the same strength ( K , = K z = lop4) (SO,31). As the HC1 concentration is raised, the hydroxyl groups of molybdic acid, which can also be written MO(OH)~,are successively replaced by chloride ions. Although the presence of appreciable concentrations of uncomplexed molybdenyl ion, MoOnf2, has been apparently ruled out (f), chloro complexes containing the molybdenyl group seem to be favored. Mo02C12is probably an important species a t moderate concentrations of HC1 (2 to 7 M) (7, 9 ), together with anionic species such as Mo02C13- and M O O ~ Cwhich ~~-~ prevail in more concentrated acid media (1, 25). Polymers such as HMo206+ in dilute acid (7) and (M0O2Cl2)2 in more concentrated HC1 (9) have been postulated in the to lO+M Mo(V1) concentration range. Guibe and Souchay (14) have argued that wave I in dilute HC1 is due to the reduction of a chloride-containing species, since the wave is not affected by the presence of 0.1M HC104 in 0.2M HCl, whereas it vanishes completely in 0.3M HC104. The disappearance of wave I at low concentrations of Mo(V1) indicates that the chloro species must be at least a dimer, while the concurrent increase in the height of wave I1 would denote the reduction of a less polymerized, probably monomeric species. Since the height of wave I1 increases a t the expense of wave I as the C1- concentration is increased, it would also
appear that Mo(V1)B contains more chloride than Mo (IV) A. LITERATURE CITED
( 1 ) Aveston, J., Anacker, E. W., Johnson,
J. S., Inorg. Chem. 3, 735 (1964). (2) Babko, A. K., Get’man, T. E., Zh. Neorgan. Khim. 4, 585 (1959); C. A . 53, 21338g (1959). (3) Bard, A. J., Solon, E., J . Phys. Chem. 67, 2326 (1963). (4) Bergh, A. A., Haight, G. P., Jr., Inorg. Chem. 1 , 688 (1962). (5) Busev, A. I., Li, G., Zh. Analit. Khim. 15, 191 (1960); C. A . 54, 1627911 (1960). (6) Carritt, D. E., Ph.D. thesis, Harvard University, Cambridge, Mass., 1947. (7) Chauveau, F., Compt. Rend. 242,2154 (1956). (8) Delahay, P., “New Instrumental
Methods in Electrochemistry,” Chap. 12, Interscience, New York, 1954. (9) Diamond, R. M., J . Phys. Chem. 61,
75 (1957). (10) El-Shamy, H. K., Barakat, M. F.1 Egypt J . Chem. 2, 101 (1959); C . A . 54, 2047a (1960). (11) Furrnan, N. H., Murray, W. M., Jr., J . A m . Chem. SOC.58, 1689 (1936). (12) Grasshoff. K.. Hahn., H.., Z . A.nul. Chem. 186, i32 (’1962). (13) Guibe, L., Souchay, P., Compt.. Rend. 244, 780 (1957). (14) Guibe, L., Souchay, P., J . Chim. Phys. 54, 684 (1957). (15) Haight, G. P., Jr., Acta Chem. Scand. 16, 659 (1962). (16) Haight, G. P., Jr., J . Inorg. Nucl. Chem. 24, 663 (1962). (17) Ibid., p. 673. (18) Hartmann, H., Schmidt, H. J., Z . Physik. Chem. (Frankfurt) 1 1 , 234 (1957). (19) Hillebrand, W. F., Lundell, G. E. F., ~
Bright, H. A., Hoffman, J. I., “ A p plied Inorganic Analysis,” 2nd ed., p. 313, Wiley, New York, 1953. (20) Holtje, R., Geyer, R., Z . Anorg.
Allgem. Chem. 246, 258 (1941). (21) Jakdb, M., Ogorzatek, M., Sikorski, H., Roczniki Chem. 35, 3 (1961); J . Electroanal. Chem. 3, A83 (1962). (22) Kolthoff, I. M., Lingane, J. J., “Polarography,” Vol. 1, pp. 90-3, Interscience, New York, 1952. (23) Ibid., p. 192. (24) Ibid., p. 202. (25) Kraus, K. A., Nelson, F., Moore, G. E., J . Am. Chem. SOC. 77, 3972 (1955). (26) Lingane, J. J., “Electroanalytical Chemistry,” 2nd ed., pp. 222-9, Interscience, New York, 1958. (27) Ibid., p. 459. (28) Meites, L. J., Record Chem. Progr. 22, 81 (1961). (29) Reilley, C. N., Stumm, W., in
“Progress in Polarography,” P. Zuman, ed., Vol. 1, Chap. V, Interscience, New York, 1962. (30) Rohwer, E. F. C. H., Cruywagen, J. J., J . S. African Chem. Inst. 16, 26 (1963). (31) Schwarzenbach, G., Meier, J., J . Inorg. IVUCZ. Chem. 8 , 302 (1958). (32) Schwing, J., Compt. Rend. 254, 4018 (1962). (33) Smith, G. F., “Cerate Oxidimetry,” G. F. Smith Chemical Co., Columbus, Ohio, 1942. (34) Zahnow, E. W., Robinson, R. J., J. Electroanal. Chem. 3. 263 (1962). (35) Zuman, P., in “Progress in Polarography,” P. Zuman, ed., Vol. 2, pp. 596-9, Interscience, New York, 1962.
RECEIVEDfor review March 2, 1965. Accepted April 14, 1965. 12th Anachem Conference, Detroit, Mich., October 21 to 23, 1964. Financial support by the Laboratory for Research on the Structure of Matter is gratefully acknowledged.
Correlation of C, Terms with Pore Size Distribution in Gas Liquid Chromatography N. C. SAHA and J. C. GlDDlNGS D eparfmenf of Chemistry, University of Ufah, Salt Lake Cify, Utah
b
An attempt has been made to correlate the measured Ci term of gas liquid chromatography with physical properties and the structure of the support; in particular, with the support’s pore size distribution. Experimental values are compared with an earlier equation based on the very simple hypothesis that the depth of a pore can be characterized by the apparent diameter as measured by mercury penetration into the support. Using experimental D I values obtained here and elsewhere, the results are found satisfactory for Chromosorb W and Gas Chrom S, where agreement is usually within a factor of 2 or 3. In the case of Chromosorb P and G the calculated values are roughly 50 times too low. Possible reasons for
822
ANALYTICAL CHEMISTRY
the discrepancies are discussed. To a good approximation, CI is proportional to R(l - R)/DI as predicted by theory.
T
HE
liquid phase nonequilibrium,
C I J term, which ordinarily has a
dominant effect on column efficiency at high velocities, assumes a value determined by the configuration of the stationary liquid. Most evidence indicates that the bulk of liquid seeks out the smallest pores and cavities of the support; thus liquid configuration would be identified with that of t l e small pores. Based on this hypothesis, two approximate equations have been proposed by Giddings ( 5 , 6) to relate Ci to pore size distribution. These
equations are tested here with the help of extensive chromatographic data and pore size determinations. The agreement with theory is satisfactory in some cases and unsatisfactory in others. Pore size distributions obtained from mercury injection [as first suggested by Baker, Lee, and Wall ( I ) , and used here] give a gross picture of pore size but do not provide the geometrical details. I n view of this it was felt at the outset (6) that agreement with either equation within a factor of 3 or 4 would be gratifying. (Since C I increases with “depth” squared, this is an error of 2 or less in the actual dimension of the liquid units.) This kind of agreement has been achieved in some cases, and represents a step forward in the independent prediction of plate height
