Electrochemical investigation of surface phenomena at a mercury

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ANALYTICAL CHEMISTRY, VOL. 51, NO. 12, OCTOBER 1979

Electrochemical Investigation of Surface Phenomena at a Mercury Electrode in Vitamin B,2a Solution C. L. Schmidt and H. S. Swofford, Jr.’ Department of Chemistry, University of Minnesota, Minneapolis, Minnesota 55455

The adsorptlon of vitamin BlZa and Its reduced forms on a mercury electrode are lnvestlgated using cyclic voltammetry, chronopotentlometty,and dlfferentlal doublelayer capacitance measurements. The parent B12aIs shown to undergo two one-electron reductions In bulk solution. Vitamin B,, is shown to be the surface actlve form and to undergo a pH dependent rearrangement on the electrode surface. The relationshlp between the reductlon of adsorbed BIZ, and the catalytic dlscharge of hydrogen is also discussed.

Vitamin B12 or closely related compounds are known to function as coenzymes in both methyl transfer and hydride transfer reactions ( I ) . In both cases, the reduced states of the vitamin (Le., either vitamin BlZror B,,,) are believed to be involved in the reaction mechanisms ( 2 , 3 ) . Vitamin B12ais the only cobalamin known to be reduced from Co(II1) to Co(1) via two one-electron steps at a mercury electrode under polarographic conditions; all other cobalamins investigated are reduced via a single, irreversible two-electron transfer step (4). Thus, voltammetric studies of vitamin Blb could provide valuable information for understanding the biochemical involvement of Blz compounds in general. There have been several voltammetric studies of the B1&/Blzrand/or B12/Blb couples which have appeared in the literature (4-10). The electrochemical behavior of the Bla/B1& couple seems to be well understood. However, voltammetric studies of the BlZa/Blzrcouple on a mercury electrode have produced eratic results (such as a dip in the polarographic wave, an “impurity” wave, distorted cyclic voltammograms, etc.). Adsorption of the vitamin on the mercury electrode surface has often been cited as a possible source of this behavior. However, an understanding of the surface activity of the vitamin is far from complete. T o date, only one study has dealt specifically with this matter (11). In the present paper, a thorough investigation of the adsorption of the vitamin from a solution of vitamin BlZais carried out using cyclic voltammetry, chronopotentiometry, and differential double-layer capacitance measurements. A complete study of the voltammetric behavior of the vitamin Bla/BlZr couple (including the effect of adsorption on the reduction of B12J will be presented in a later paper.

EXPERIMENTAL Equipment. All voltammetric experiments conducted at sweep rates slower than 500 mV/s were performed using a PAR Model 174 Polarograph. Chronopotentiometric and voltammetric experiments a t fast sweep rates were performed with the aid of a multipurpose, three-electrode electrochemical instrument constructed in house (12). Slow sweep experiments (less that 300 mV/s) were recorded on a Moseley 2D-2M X - Y recorder which had a full scale response time of 0.5 s. The fast experiments were recorded on a Tektronix type 564 storage oscilloscope and then were photographed with a Polaroid camera attachment (Tektronix type C-12). Tektronix type 3A72 plug-in modules were used for both the X and Y inputs of the scope for the recording of cyclic and linear sweep voltammograms. For chronopotentiometric and differential dou0003-2700/79/035 1-2026$01.OO/O

ble-layer capacitance measurements, a Tektronix type 3B3 time base module was used for the X input of the scope. A schematic diagram of the apparatus used for the measurement of differential double-layer capacitance is shown in Figure 1. The capacitance bridge, tuned amplifier, and signal generator were all manufactured by the General Radio Corporation and have model numbers 1615-A,2465, and 1311-A,respectively. The output of the tuned amplifier was fed directly into the storage oscilloscope described above. The scope was triggered internally by the signal generated when a drop fell from the dropping electrode. A Leeds & Northrup type K-3 potentiometer was used to apply the potential to the dropping mercury electrode in connection with the double-layer capacitance measurements. The DME which was used for the differential double-layer capacitance measurements was constructed using a 0.1-mm bore capillary which was drawn down to an outside diameter of 0.5 mm (to reduce the shielding of the mercury drop by the capillary) (13). This capillary gave a convenient drop time of about 9 s when using a 40-cm head of mercury. In all experiments, the reference electrode (SCE) was isolated from the cell by a bridge containing the same electrolyte as the cell. Oxygen was removed from all experimental solutions by purging them with nitrogen for several minutes; an atmosphere of nitrogen was then maintained above all solutions during the course of the experimental work. Prior to entering into the cell, the nitrogen was first passed over hot copper turnings (to remove any residual oxygen) and then bubbled through water to minimize evaporative losses from the cell. Reagents. Vitamin B1%was purchased from the Tridom-Fluka Chemical Company in the form of hydroxocobalamin hydrochloride. Chloride was removed by anion exchange on Sephadex A-50 which had been charged with 0.0075 M NaOH. The effectiveness of this column toward the removal of chloride was M NaCl solution through checked by running 50 mL of a 3 X the column (Le., the approximate concentration of vitamin BlZa to be put through the column). The resulting solution was then acidified with HNO,; no precipitate was formed upon the addition of 0.1 M AgNO, (Le., the concentration of chloride ion was effectively reduced). The concentrations of all stock solutions were determined spectrophotometrically by conversion to the dicyano complex in 0.1 M KCN.

RESULTS AND DISCUSSION Studies by Cyclic Voltammetry. Vitamin BlZain bulk solution is well known to be electroactive at a mercury electrode; hence, it is reasonable to expect that the adsorbed species may be electroactive as well. Cyclic voltammetry is a particularly useful tool for this type of investigation since voltammetric peaks arising from the oxidation or reduction of adsorbed materials have many diagnostic characteristics and are easily distinguished from diffusion controlled processes (14). By using a dilute solution of the adsorbing material and a relatively fast sweep rate, it is possible to minimize the current flow due to diffusion controlled processes relative to that observed for surface processes. Typical results for such a cyclic voltammetric experiment for vitamin B12aare shown in Figure 2. This figure represents a “steady-state’’ cyclic voltammogram (Le., the condition reached after many cycles). Attention is focused on the pairs of peaks at +0.2 V and -1.6 D 1979 American Chemical Society

ANALYTICAL CHEMISTRY, VOL. 51,NO. 12, OCTOBER 1979

2027

Table I. Effect of Vitamin B,,, Concentration on Height of Surface Peaksa i, for i, for i, for i, for anodic cathod- anodic cathodpeak at ic peak peak at ic peak at 0.2 - 1.6 V, at - 1.6 0.2 v, PA V,PA V, P A [B,Z,l,M PA _0.3 0.2 1.0 x 0.2 0.5 4.0 3.0 3.0 4.0 x 10-5 1.6 4.8 3.0 9.0 x 10-5 3.0 2.0 4.8 3.0 1.7 x 10-4 3.0 2.4 4.8 3.3 x 3.0 3.0 4.8 2.8 3.0 5.7 x 10-4 3.0

4

a

x

NaC10,; sweep rate = 1.5 V/s; HMDE area = 2.9 cmz.

0.1 M

Figure 1. Differential capacitance apparatus

'O l

0.3

a1

'01

-03

05

-07

-09

-I

-13

-15

- 1 7

1.2

16

2'0

2 4

2 8

32

.OLTS V 5 SCE

Figure 2. Steady-state cyclic vokammograrn of

M B12ain 0.1 M NaClO at pH 6.5 (unbuffered). Sweep rate, 3.0 V/s; HMDE area, 2.9 1 X IO- crn2

V vs. SCE. The remaining small peaks are due to the behavior of bulk vitamin Biz,; this becomes obvious upon increasing the vitamin B12aconcentration. The observed peaks have several important characteristics which are immediately obvious. First, both pairs of peaks show only a small peak separation as opposed to a minimum peak separation of 56/n mV observed for a diffusion controlled process. Second, the peaks are nearly symmetrical in shape with an average half-width of about 80 mV. The actual individual peak widths have been observed to vary from 63 to 90 mV depending on the particular experimental conditions (i.e., concentration, sweep rate, the presence of foreign substances, etc.). The probable cause of these variations will be discussed later. However, the half-widths agree reasonably well with the theoretical value of 90 mV predicted for a reversible one-electron oxidation or reduction (14). Therefore, the cathodic peak a t +0.2 V would appear to represent the reduction of vitamin BlZato vitamin BlZr,and the cathodic peak at -1.6 V represents the subsequent reduction of vitamin BlZrto vitamin B12s. The dependence of the four peak heights on the bulk vitamin B1%concentration is shown in Table I. Again, the peak height values were obtained from steady-state cyclic voltammograms. The height of all the peaks, with the exception of the anodic peak at -1.6 V, becomes independent of the bulk concentration of vitamin Bl%as it is increased above 9 X M. This behavior is again what would be expected for an adsorbed electroactive species and represents the point where the electrode surface has attained maximum coverage. The anodic and cathodic peaks a t +0.2 V both reach the same constant peak height. The cathodic peak observed at -1.6 V

LOG

SWEEP RATE

Figure 3. Log of peak current (PA) as a function of log sweep rate (rnV/s) for the cathodic surface peak at -1.6V (lo-' M BIza in 0.1 M NaCIO,; HMDE area, 2.9 X cm2; pH adjusted w b NaOH or W104). Electrode was poised at -1.0 V for 2 min prior to sweep. Upper curve: pH 5.5; slope at upper end is 0.96;slope at lower end is 0.50. Lower curve: pH 10.2; slope is 1.00

reaches a much larger value at this pH, while the anodic peak at -1.6 V does not attain a limiting height at this experimental pH and sweep rate. These latter two observations would seem to be somewhat inconsistent with the proposed model. If n = 1for both observed sets of electron transfer reactions (when full electrode coverage is attained), the peaks would all be expected to have nearly identical heights. This is clearly not the case. Thus, a more careful study of the pair a t -1.6 V was warranted. The dependence of observed peak height on sweep rate was extremely useful in elucidating the nature of these apparent discrepancies. The anodic and cathodic peaks a t +0.2 V exhibit a linear dependence of peak height on sweep rate as is expected in the case of adsorption. The height of the cathodic peak at -1.6 is also linearly dependent on the sweep rate at pH 10.2 as is illustrated in Figure 3. However, at pH 5.5, a linear dependence on the sweep rate is approached only at fast sweep rates, and the peak current (i,) approaches values obtained a t pH 10.2. At slow sweep rates, the peak height is dependent on the square root of the voltage sweep rate indicating that a diffusion controlled reduction is primarily responsible for the observed current. The pH dependence described above suggests that the surface reduction of vitamin BlZrto BlZs(i.e., the cathodic surface peak observed at -1.6 V) is intimately tied to the catalytic reduction of hydrogen which has been previously observed in solutions of vitamin BIZ(5, 15, 16). At high pH (Le., pH 10.2), the current due to the catalytic evolution of

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ANALYTICAL CHEMISTRY, VOL. 51,

NO. 12, OCTOBER

1979

Table 11. Effect of pH, Sweep Rate, and Vitamin B,,, Concentration on the Peak Height of the Anodic Surface Peak at -1.6 Vapb pH 6.5 6.5 6.5 6.5 6.5 6.5 6.5 6.5

[B,,,], M 1.0 x

loT6

9.0 x lo-’ 3.3 X 5.7 x 5.7 x 5.7 x 5.7 x 10-4 5.7 X

sweep rate, V/s

0

3.0 6.0

6.0

IJ~.S/V

0 0.8

2.4

2.0 2.0 2.0

6.0

1.5

02 0.030

initial polarization potential, volts vs. SCE 0.3 0.2

i IV,

i, A

3.0 3.0 3.0

Table IV. Presence of Surface Peaks as ;Function Initial Polarization PotentiaW

12.1 2.3

0.1 0.0 - 0.1

-0.2 -0.3 -0.4

1.6 0.67

0.5

inverted peak 0.043

-0.5 - 0.6

0.022 2.0 10.2 5.7 X 0.044 0.087 2.0 10.2 5.7 x a The underlined values represent parameters which are being varied. 0.1 M NaClO,; HMDE area = 2.9 x lo-’ cmz.

- 0.7 - 0.8

PH 6.5 6.5

10.2 10.2

peak cathodic at 0.2 V anodic at 0.2 V anodic at - 1 . 6 V anodic at -1.6 V

3.6 3.9 3.7 3.7

X X X X

loT6

rrnax,

mol/cm 3.7 x 4.0 x

3.8 X 3.8 x

lo-“ lo-’’ 10.’’ lo-’’

hydrogen is extremely small with respect to the current arising from the reduction of adsorbed vitamin BlZr,and the linear sweep rate dependence is observed. At pH 5.5, the catalytic hydrogen current is roughly a factor of ten greater than the resulting current from the surface reduction when a sweep rate of 20 mV/s is used; thus, diffusion limits the current. As the voltage sweep rate is increased, the surface reduction current begins to overwhelm the catalytic hydrogen current since the latter increases with only the square root of the voltage sweep rate. I t should be noted that a somewhat different process is apparently involved in the catalytic reduction of hydrogen below pH 4.5. However, the catalytic hydrogen problem is not the primary focus of this paper; insights gleaned during the course of the primary study which are pertinent to this problem will be presented later. Discussion of the catalytic hydrogen process in the present context is directed toward establishing a more complete understanding of the observed surface electrochemistry for vitamin Blz. The anodic peak a t -1.6 V exhibits even more unusual behavior than its corresponding cathodic counterpart. Data for this peak at various concentrations, pHs, and voltage sweep rates are given in Table 11. The last column in this table (at the right) represents the ratio of the observed anodic peak current to the voltage sweep rate. This ratio would be expected to remain constant (for a well behaved surface process) with a maximum value of about 2.0 (Le., the same size as the other three observed surface peaks). It is the last three lines of the table that are of greatest significance. When the pH is increased from 6.5 to 10.2, the behavior of the peak current changes from a very unusual “inverted” peak to the expected ratio of 2.0. It is not coincidental that the expected behavior occurs when the contribution to the current from catalytic hydrogen evolution is minimal (i.e., pH 10.2). However, a complete explanation of this current inversion phenomenon and the data must await the presentation of additional experimental results. The maximum areas for the four surface peaks as determined by graphical integration are presented in Table 111. The units are reported as coulombs of charge per square centimeter of electrode surface. The areas of the pair of peaks at -1.6 V were determined a t pH 10.2 to minimize any

a

3.0 V/s.

1.0 5.2 8.0

8.0 8.0 8.0 8.0 8.0 8.0 8.0

-1.50 - 1.60 - 1.70

0.6 0.0 0.0

-1.1 -1.2

-1.3 -1.4

area, C/cm

0.0

-1.45

0.9

- 1.0

Table 111. Area of Surface Peaks

height of surface peaks at + 0 . 2 V, PA

8.0 8.0 8.0 8.0 8.0 8.0 7.8 5.0

-

of

M B,,, in 0.1 M NaC10, at pH 6 . 5 ; sweep rate =

Initial polarization time was 2 min.

contribution from catalytic hydrogen current. The areas of the peaks at +0.2 V remain unchanged as the pH is decreased; however, the peaks become obscured by the anodic limiting current a t the higher pH values. The maximum areas of all the peaks are equivalent within the experimental error of the method. This is further evidence that the number of electrons transferred in each step (i.e., peak) is identical. If it is assumed that n = 1as is suggested by the peak half-widths, it is possible to calculate the maximum surface excess (r-) of the adsorbed vitamin. The average value of 3.9 x lo-” mol/cm2 reported in Table I11 translates into an area of 430 A2 per adsorbed molecule. This value corresponds roughly to the projected geometrical area estimated for a vitamin BlZrmolecule adsorbed in a flat orientation (17). All studies discussed thus far support the adsorption of the vitamin onto the mercury electrode surface. However, we have yet to establish over what range of potentials the vitamin is adsorbed. Such information was obtained by using a dilute solution of vitamin BlZa,and poising the electrode a t the potential of interest. After allowing the adsorption process to come to an equilibrium, the potential was rapidly scanned through one cycle. The extent to which diffusion can supply more material on the time scale of the potential sweep is exceedingly small, so that the observed peaks are representative of the amount of material adsorbed at the preselected initial potential. For a lo4 M vitamin BlZasolution, adsorption equilibrium was determined to be well established with 90 s (Le., the peaks had attained their maximum size.) Thus, an initial polarization time of 90 s was used in all such studies. The results from these experiments are shown in Table IV. It can clearly be seen that the vitamin completely covers the electrode surface (at the concentration used) from +0.2 V to -1.6 V. Adsorption does not occur to any detectable extent at potentials which are either positive of +0,2 V or negative of -1.6 V. This strongly indicates that only the Co(I1) form of the vitamin (vitamin BlZr)is adsorbed. The adsorption of vitamin BlZris not, however, an irreversible process. Initial polarization of the electrode in the adsorption region h e . , between +0.2 V and -1.6 V) followed by polarization for a few seconds at +0.3 V or -1.7 V prior to the scan yields a cyclic

ANALYTICAL CHEMISTRY, VOL. 51, NO. 12, OCTOBER 1979

voltammogram which does not show any surface peaks. Therefore, these results can be summarized as follows: VB12,(solution) e + VBIZ,(surface) E”’ = 0.20 V (1) VB&urface) + e + VB12,(solution) Eo’ = -1.6 V (2) These results now allow a much more complete discussion and explanation for the data presented in Table 11. First, the situation that leads to the behavior observed at pH 10.2 is relatively simple to explain. As the potential is swept negatively past -1.6 V, the vitamin BIZron the surface is reduced to vitamin Blzsand then desorbed (Equation 2). On the return sweep, the concentration of vitamin BlZs at (or near) the electrode surface is sufficient to fully cover the electrode surface as it is re-oxidized to vitamin BlZr. Since the contribution from catalytic hydrogen evolutior. is very small, ideal surface behavior is observed. At pH 5.5, however, an analogous situation does not exist. Here, as the adsorbed vitamin BIZr is reduced to Blzs,the catalytic hydrogen process (which occurs simultaneously) produces a much larger current than at pH 10.2. Upon desorption of the vitamin BlZr,the catalytic hydrogen process ends abruptly (although hydrogen ions continue to diffuse to the electrode surface). As re-oxidation and re-adsorption occur on the return sweep, the catalytic hydrogen reduction can begin again. When the sweep rate is very slow, the catalytic hydrogen reduction current (diffusion process) exceeds the surface oxidation current and a net cathodic current is observed, thus explaining the existence of the inverted peak observed at -1.6 V. As the sweep rate is increased, the current due to the surface process increases much faster than the current due to the diffusion process (Le., catalytic hydrogen) and a normal surface peak is observed. The effect of concentration can be seen in the first four lines of Table 11. Eventually the bulk concentration of vitamin B,, becomes sufficiently low such that the electrode surface cannot completely cover on the return sweep, and a surface peak of diminished size is observed. Therefore, depending on the combination of experimental condition: pH, sweep rate, and bulk vitamin B12aconcentration, any of the following cases are observed at -1.6 V in cyclic voltammetry. 1. A “normal” pair of surface peaks (Equation 2) of equal height. 2. A totally irreversible reduction peak B,,,(surface) + e B,,,(solution). 3. A surface reduction peak on the negative voltage sweep and an inverted peak on the positive return sweep. Indeed, all three of these cases have also been observed by Tackett and Ide (15). However, these authors did not understand or appreciate the nature of the peaks, and they offered no reasonable explanation for the inverted peak. Furthermore, they stated that the anodic peak at -1.6 V was present only in freshly prepared solutions. We were not able to observe any effect of aging of the solution in our work. In fact, the behavior described above can be reproduced using solutions over one year old! On the basis of the arguments presented above, one would also expect to observe cases 1 and 2 for the pair of peaks at +0.2 V as well (except that the behavior of the anodic and cathodic peaks would be reversed). Case 1 (a normal pair of surface peaks-Equation 1)is observed most often. However, if the surface is allowed to cover at a potential between +0.2 V and -1.6 V using a low concentration of vitamin BlZa(ca. lo4 M), a subsequent scan at a relatively slow sweep rate (ca. 20 mV/s) shows the size of the cathodic peak to be greatly diminished with respect to the fully developed anodic peak. This effect of desorption at +0.2 V does not seem to occur until considerable lower concentrations of bulk BlZa(or slower sweep rates) are reached relative to the situation which produces a similar effect at -1.6 V. This may indicate that the kinetics of the desorption process are slower at +0.2 V.

+

2029

The position of all the surface peaks relative to their corresponding peaks for the reduction of the bulk vitamin is also consistent with the idea that vitamin BlZris the surface active form of the parent material. If the oxidized form of the vitamin were adsorbed as strongly as the reduced form, the peak potential for the surface peak would be expected to occur at the E” for the bulk material (19). On the other hand, stabilization of the reduced material relative to the oxidized form by stronger surface adsorption of the reduced form would effectively facilitate the reduction process and shift the peak potential positive of E o for the bulk redox process. This is exactly what is observed in the case of the set of peaks at +0.2 V (Equation 1). Alternatively, when the oxidized form is adsorbed more strongly than the reduced form, the reduction process is less favorable and a negative shift of the surface peak potential would be expected; this is observed for the reduction of adsorbed vitamin BlZrto bulk vitamin Blzsat -1.6 V (Equation 2). Finally, the small variations in the surface peak half-widths with respect to the theoretical value of 9O/n mV can now be discussed. The most complete and significant theoretical work on the nature of surface peaks has been published by Laviron (14). It is from his work that the theoretical half-width of 9 0 / n mV was derived. This derivation is straightforward, but it does assume that both the oxidized and reduced forms are strongly adsorbed on the electrode and that a Langmuir isotherm is followed in the adsorption process. The former condition is definitely not met in this instance, and adherence to the latter condition is also questionable. Laviron has shown that adherence to other isotherm forms can cause the peak half-widths to become significantly narrower or wider than the 9 0 / n mV value (although other characteristics remain unchanged). However, the consistency of peak areas, and the very reasonable value calculated for rmax are very strong evidence that the number of electrons transferred in each instance is indeed one. In addition, the peaks exhibit the most nearly ideal behavior at fast sweep rates, when the effect of desorption is minimized. The investigations discussed in this section raise questions as to why such seemingly obvious phenomena have not been previously reported. Part of the answer is most certainly that relatively dilute solutions m u s t be used before t h e surface peaks become significantly large uith respect to the diffusion peaks. Since it is advantageous to use more concentrated solutions when studying the electrochemistry of a soluble and diffusing species (to minimize the effect of charging current), the surface peaks (which are independent of concentration) are easily obscurred. Perhaps an even more important fact is that t h e surface peaks at f0.2 V a r e totally obscured by the current resulting from the oxidation of the mercury electrode when even small concentrations o f complexing or precipitating agents are present. In fact, the surface peaks at +0.2 V are not observed when hydroxocobalamin-hydrochloride is used as the source of vitamin B12aunless the chloride is first removed by anion exchange. Virtually all previous studies have been carried out using either hydroxocobalamin-hydrochloride, a method of vitamin BIZasynthesis involving chloride ion, or a supporting electrolyte containing chloride ion or EDTA. Thus, the surface peaks at +0.2 V have been missed in past published work. Finally, the surface related phenomena at -1.6 V have been observed (15), but they have never been recognized as such because of their association with the catalytic evolution of hydrogen. Studies by Chronopotentiometry. Chronopotentiometry experiments were carried out to further verify the conclusions drawn from the cyclic volammetric experiments. A typical

ANALYTICAL CHEMISTRY, VOL. 51, NO. 12, OCTOBER 1979

2030

-7

I 0

0.2 TIME

04 ISECONDSI

0.6

0.0

Figure 4. Chronopotentiometryof lo-’ M B,28in 0.1 M NaCiO, at pH 5.0. Applied current, 2.0 pA cathodic; HMDE area, 2.9 X cm2. Electrode was allowed to rest in solution for 90 s prior to scan

Table V. Chronopotentiometric Data for Adsorbed Vitamin B,,, on the HMDEaib $ / A , first

i, @ A 8.0 4.0

2.0 1.0 0.5

i r / A , second i T 1 l 2 / Asecond ,

transition, C/cm

transition, C/cm

transition, C/s *I1 .cm

4.4x 5.4 x 4.4 x 3.4 x 10-6 4.4 x 1 0 - 6

8.0 x 8.9 X 1.0 x 1.1x 10-5 2.2 x 10-5

4.8 X lo-’ 3.4 X 2.6 x 2 . 3 x 10-5 2 . 2 x 10-5

loF6M B,,, in 0.1 M NaC10, a t pH 5 . 5 . Electrode was allowed to rest in solution for 90 s prior to scan. chronopotentiogram is shown in Figure 4. Once again, a very dilute solution of vitamin B1, was used (lo4 M) to minimize any contribution to the transition times (7) which might result from diffusing material. To ensure the establishment of adsorption equilibrium, the open-circuited electrode was allowed to “rest” in solution for 90 s prior to the application of current. As demonstrated earlier, vitamin BlZais not adsorbed when electron transfer is allowed to occur (Le., a closed circuit condition) because it is either immediately reduced to vitamin BlZr,or desorbed if the potential is positive of +0.2 V. However, vitamin BlZais adsorbed on an open circuited electrode in this case as is obvious from Figure 4. The adsorption of vitamin Blk at open circuit was also verified by performing cyclic voltammetric experiments on an electrode which had been allowed to equilibrate with vitamin B,, under conditions similar to chronopotentiometry. The cyclic voltammetric studies showed that an electrode can be fully covered with vitamin B12aproviding the bulk concentration is in excess of M. If in Figure 4 the first transition (ca. +0.2 V) represents the following reduction, VBI2,(surface) e + VB,,,(surface) (3)

+

the product, i T , for each recorded transition would be expected to remain constant as the current is varied. Data for this reduction process at various applied currents are shown in Table V. The precision with which these short transition times can be measured is not great; the transition time is short with respect to the time required to charge the double layer. However, the product, i ~ is, constant within the experimental error expected for the method. Furthermore, the average value for i r / A is 4.4 X lo4 C/cm2 which compares favorably with the peak area determined via cyclic voltammetry (see Table 111). Assuming vitamins B12aand B12roccupy the same area in the adsorbed state, the two values should be equal since they both represent the total coulombs of charge required to carry out the reduction per unit area of electrode surface. The second transition shown in Figure 4 (ca. -1.6 V) would be expected to be longer than the first transition time owing

to the contribution from the catalytic hydrogen process occuring at pH 5.0. In addition, the behavior of this transition would be expected to lie somewhere intermediate between the ir behavior predicted for reduction of adsorbed substances and the behavior predicted for strictly diffusion controlled reductions. The last two columns of Table V show that this “mixed” character is indeed observed. At small values of the applied current, the product i d z approaches a nearly constant value because the diffusion controlled component predominates. As the applied current is increased, it is ir which begins to approach a more nearly constant value. Although there is still considerable contribution to the transition from the catalytic hydrogen process ( i is~ still nearly twice that observed for the first transition), the data definitely reflect the surface component in this process. This surface character again illustrates the adsorption of vitamin BlZron the electrode surface. Experiments by Differential Double-Layer Capacitance. Both cyclic voltammetry and chronopotentiometry have shown that vitamin BIzr is strongly adsorbed on a mercury electrode over a wide range of potentials. Adsorbed substances often cause distortions in polarographic waves (increased irreversibility), but the results discussed thus far cannot fully account for the unusual polarographic behavior observed for the reduction of vitamin BlZa to BlZr. The measurement of differential double-layer capacitance is not nearly as useful as cyclic voltammetric or chronopotentiometric data in determining the oxidation states and surface concentrations of adsorbed substances. However, it can be an extremely valuable tool for monitoring significant changes in the structure of the double layer as a function of time or applied potential. Such events are often signaled by abrupt changes in the electrode capacitance (including capacitance peaks in the case of desorption). The operation of the capacitance bridge used in conjunction with the fine-tipped capillary was checked by comparing double-layer capacitance measurements in 1 M KC1 with the values obtained by D. C. Grahame (13). The agreement was excellent. The pH of the vitamin B1%test solutions was adjusted by the addition of HC104 or NaOH. The use of buffers was avoided to circumvent any complications due to changes in the nature of the electrolyte with pH. Employing the fineM solution of tipped capillary described earlier, a 3 X vitamin BlZareached adsorption equilibrium within the first 5 s of drop life. This fact was indicated by the constancy of the capacitance per unit electrode area for times longer than 5 s. Therefore, the capacitance measurements were made at a point 6 s into drop life to ensure equilibrium conditions and, further, to minimize the effects of any faradaic current flow resulting from the reduction of bulk vitamin Bl,. During the determination of the optimum equilibration time, the capacitance was occasionally observed to come to balance as many as three times prior to the establishment of final adsorption equilibrium. It is not uncommon to observe two balance points (the first balance point can occur as the surface is still covering). However, the appearance of three balance points is very unusual (also providing much insight in this case) and will be discussed later in this section. The data recorded for the double-layer capacitance measurements are shown graphically in Figure 5. A plot of data taken at pH 10 is not included in Figure 5 since it is essentially identical to those data taken at pH 7 . Measurements were not made at extremely negative potentials because of the presence of a large faradaic current due to the catalytic reduction of hydrogen ion. The suppressed capacitance values for the solutions containing vitamin B12arelative to the curve for the supporting

2031

ANALYTICAL CHEMISTRY, VOL. 51, NO. 12, OCTOBER 1979

:I :-\ 0

2

4

6

5

TIME

,

,

0

);

2

,

,

4

,

,

6

,

,

,

a

ISEC,

Figure 6. Differential double-layer capacitance as a function of time into drop life (3 X M B12ain 1.O M NaCIO,). (A) pH 6.5, Eapp = -0.3 V; (B) pH 6.5, Eapp = -0.20 V; (C) pH 2.8, +Eapp = -0.20 V; (D) pH 2.8, Eapp= 0.05 V

03

01

-01

-0.3

-a5

VOLTS

VS

-07

-os

-in

-13

SCE

Flgure 5. Differential double-layer capacitance at DME as a function of applied potential. (A)1.0 M NaCIO,; (0)3 X lo-' M B12ain 1.0 M NaCIO, at pH 7.0; (0)3 X

M BlZa in 1.0 M NaCIO, at pH 2.7

electrolyte clearly illustrate the adsorption of the vitamin on the mercury electrode. Again, it should be noted that although vitamin Blzais the form present in bulk solution, it is vitamin BlZrwhich is adsorbed. The curve a t p H 7 agrees well with the capacitance curves obtained by Imhoff (11). However, Imhoff did not observe the peak a t +0.2 V (probably because virtually all of his work was carried out in a KCl electrolyte), nor did he do any experimental work using solutions of low PH. In light of the discussion in the preceding paragraphs of this section, the peak a t +0.2 V on the pH 7 curve (Figure 5) must result from desorption. For this same curve, an abrupt change in the capacitance is also observed around -0.25 V. Since vitamin BIZ*has been shown to fully cover the electrode in the voltage range from +0.2 V to -1.6 V (see Table IV), this change in capacitance is strong evidence that a restructuring of the adsorbed material occurs a t or near -0.25 V. T h e capacitance curve for the pH 2.7 solution exhibits several significant differences from its pH 7 counterpart. The most obvious of these is the disappearance of the peak a t +0.2 V. Cyclic voltammetric studies of the adsorption process in this pH range show that vitamin BIzris still adsorbed over the same range of potentials ($0.2 V to -1.6 V) and is desorbed following its oxidation to vitamin B1%. Thus, it would appear that the desorption process i s too slow to produce a peak on the time scale of this experiment (input frequency = 500 Hz). [It is interesting to note that the half-width of the cyclic voltammetric peak for the oxidation of vitamin Blz,(surface) to vitamin Bl,(bulk) is markedly reduced in solutions of low PH.1 The abrupt break in the capacitance curve for the p H 2.7 solution occurs at about -0.05 V as compared to -0.25 V for the p H 7 solution. Again, this type of behavior is indicative of a restructuring of the adsorbed material. The similarity in appearance to the curve for the pH 7 solution may even suggest that the same type of restructuring occurs in both solutions. It should also be mentioned that there is another (reproducible) abrupt change in capacitance a t -0.3 V for the p H 2.7 solution which could represent another change in the double-layer structure. However, the net change in capacitance is small and there is little other experimental evidence to support this speculation.

Further support for the restructuring of the adsorbed material a t -0.25 V (in the p H 7 solution), and at -0.05 V (in the p H 2.7 solution) is gained through an investigation of capacitance as a function of drop life a t various applied potentials. This unusual behavior (noted earlier in this section) is shown in Figure 6. A blunt capillary was used in these investigations simply because it is much less troublesome to use experimentally than the fine-tipped variety. Thus, measured capacitances may differ by 3% to 5% from their true values, but the curves in Figure 6 will still be qualitatively correct. Because of a faster rate of mercury flow from the blunt capillary, the final capacitance values shown on the plots in Figure 6, A and C, are slightly higher than the final equilibrium values reported in Figure 5 (Le., adsorption equilibrium was not quite attained before the end of drop life). A capacitancetime dependence was also noted by Imhoff. However, he did not investigate this phenomenon as a function of applied potential, or solution pH. Figure 6 illustrates the importance of both pH and applied potential as independent, experimental parameters. The plots in Figure 6, B and D, show that if the mercury electrode is poised a t a potential which is positive of the abrupt break in the capacitancepotential curve for each respective solution, the capacitance quickly attains an equilibrium value. However, when the electrode is poised at a potential negative of the break in the capacitance-potential curve, the capacitance rapidly reaches a metastable value which, after a period of time, decays sharply to the true equilibrium capacitance. This behavior is illustrated for each solution by curves A and C in Figure 6. The experiment was performed at several other potentials anodic and cathodic of the abrupt break with identical results. These results indicate a distinct difference in the double-layer structure on either side of the break in the capacitancepotential curves. The correlation of these capacitancetime curves with cyclic voltammetric data on a growing mercury drop further supports a change in the double-layer structure. The results of such experiments using the same blunt capillary and solutions as described above are shown in Figure 7. Although scans were taken at various times in drop life, only those scans that were initiated a t 5 and 7 s into drop life are included for the sake of clarity. Approximately 1s was required for each entire scan to be completed. Only the cyclic voltammograms of the pH 7 solution are shown as the results for the p H 2.7 solution are completely analogous. When the initial potential is positive of the break in the capacitance-potential curve (see Figure 7B), a normal cyclic voltammogram is obtained a t any time in drop life. The

2032

ANALYTICAL CHEMISTRY, VOL. 51, NO. 12, OCTOBER 1979

04

8

i

, 2.3

I

,

, C.’

.

‘rCLTS

.

.

-0.1

. -03

.

1

-0.5

V S SCE

Figure 7. Cyclic voltammetry on a DME (3 X M B,2a in 1.0 M NaCIO, at pH 6.5). Scans were initiated in the positive direction: sweep rate, 3.0 V l s . Inner curves were initiated at 5 s into drop life, outer curves at 7 s. (A) Initial potential = -0.3 V; (B) initial potential = -0.2

v

increase in peak size is simply due to the increased surface area because of the growth of the mercury drop. However, when the initial potential is even slightly negative of the break in the capacitance-potential curve (see Figure 7A), a very narrow, spike shaped peak begins to develop at about, 6 s into drop life. The time of the appearance of the spiked peak correlates well with the time of the occurrence of the break in the capacitancetime curves (compare with Figure 6, A and C). I t is somewhat difficult to determine whether a spiked peak such as this is due to a faradaic surface process (Le., an electron transfer reaction), or is simply due to a charging spike associated with the restructuring of the double layer; the characteristics of these two types of peak would be expected to be quite similar. However, there are two diagnostics which suggest that the peak arises from a charging phenomenon. First, the peak is extremely narrow (half-width of less than 30 mV). I t is not impossible for surface faradaic peaks to be this narrow, but the occurrence of such peaks is not common. Second, the peak is symmetrically shaped which would indicate fast and reversible electron transfer kinetics for a faradaic process (14). Under these circumstances, a cathodic peak would be expected a t the same potential on the return sweep. None is observed. Thus, the case for a restructuring or reorientation of the adsorbed material is a strong one. The exact nature of this restructuring is a very complex problem which has not been resolved. One possible explanation could be a shift from a flat orientation of the adsorbed vitamin B12r(positive of the break in the capacitance-potentia1 curve) to an edge-on orientation (at the more negative potentials). However, this type of reorientation should be accompanied by an increase in rmax (the maximum surface excess) which was never observed (14). Another possible explanation is that the “restructuring” represents the expulsion of water molecules and/or ions from the adsorbed layer. This behavior could cause a substantial decrease in the observed capacitance as is actually seen, and would probably cause the adsorbed layer to become relatively resistant to electron transfer or to penetration by any diffusing species. At least one similar case has been reported in the literature (18). Moreover, the po-

tential a t which this expulsion would occur would probably be affected by the protonation of the benzimidazole nitrogen, thus explaining the pH dependence of the break in the capacitance-potential curve (the benzimidazole nitrogen of adsorbed vitamin B12, would be expected to protonate somewhere between pH 3 and 5 ) . This is again consistent with observations in the literature (14). The Catalytic Hydrogen Process. As mentioned earlier, the catalytic hydrogen process is not a primary focus of this paper. However, several observations were made during the course of this investigation which may be of use in the ultimate elucidation of this problem. Above pH 4.5, the catalytic hydrogen process has been shown to occur simultaneously with the reduction of vitamin B12,(surface) to vitamin B12,(bulk). In DME polarography, the catalytic hydrogen current appears as a peak a t -1.6 V rather than a wave. Again, this reduction occurs as vitamin B12,(surface) is reduced to vitamin B,,(bulk); yet, a high concentration of vitamin B1%is present at or near the electrode whenever the electrode potential is negative of EoB12r,B1, and no catalytic hydrogen activity is observed. Both of these observations suggest the importance of adsorption in the reduction process. As the pH is decreased below 4.5, the process responsible for the production of the catalytic hydrogen current appears to change. In cyclic voltammetry a new, totally irreversible peak appears a t approximately -1.1 V. The peak height is directly proportional to the hydrogen ion concentration in unbuffered solutions. In fact, this peak represents the quantitative reduction of diffusing hydrogen ion. In buffered solutions the current flow is so great that hydrogen bubbles immediately form on the electrode and prevent any meaningful measurements. In DME polarography a t pHs below 4.5, the catalytic hydrogen process appears as an irreversible wave with a half-wave potential of about -1.1 V. The wave height is directly proportional to the hydrogen ion concentration in unbuffered solutions, and height studies indicate that the process is strictly diffusion controlled. As before, the process terminates a t --1.6 V with the desorption of vitamin B12r. The dependence of the catalytic hydrogen process on the adsorption of vitamin VIZris not too surprising since there are many instances in which only the adsorbed form of a molecule is catalytically active. Virtually all substances which facilitate the catalytic reduction of hydrogen ion contain a base which is capable of protonation. In such instances, the mechanism for the reduction is often believed to proceed via the following mechanism (19). B

+ H 3 0 +* B H + + H 2 0

+ e + BH. 2 BH- + 2B + Hz BH’

(4-36) (4-37) (4-38)

This same mechanism may explain the behavior of the catalytic hydrogen process in the presence of adsorbed vitamin B12r. Vitamin Bl, is known to exist almost entirely as the base-on form in bulk solution when the solution pH is above 2.9 (8). When vitamin BIZ,(surface)is reduced to BIzswhich exists exclusively in the base-off form, the vitamin BIZsmay reside on the electrode surface long enough to produce the observed catalytic hydrogen current. This would account for the behavior above p H 4.5. Below pH 4.5, the change in the nature of the catalytic reduction process may arise from the release of the base from the cobalt and its subsequent protonation. In bulk solution, substantial protonation of the base does not occur until p H 2.9 is reached. However, it is possible that the base is not bound as tightly in adsorbed vitamin BIZ*and protonation is

ANALYTICAL CHEMISTRY, VOL. 51, NO. 12, OCTOBER 1979

possible at considerably higher pH. In the pH region above 4.5, there is no evidence that the catalytic hydrogen process involves t h e cycling of t h e vitamin B12r/B12s couple as suggested by Lexa and Saueant. The catalytic hydrogen process is an interesting phenomenon and bears a remarkable resemblance to the so called "BrdiEka currents". The hypothesis in the preceding paragraph could possibly be tested by studying the electrochemical behavior of a cobinamide (which does not contain the benzimidazole group). The problem is certainly worthy of further investigation.

SUMMARY Vitamin BlZris shown to be the only form of the vitamin adsorbed on a mercury electrode (in a closed circuit configuration). Vitamin BlZrappears to be adsorbed in a flat orientation and fully covers the electrode surface from +0.2 to -1.6 V vs. SCE. The observed electrochemical activity arises from the following reactions.

+ le + le

B12,(solution) B,,,(surface)

B12,(surface) B12,(solution)

E"' = 0.2 V

E"' = -1.6 V

The adsorbed layer undergoes a rearrangement at -0.25 V in a pH 7 solution which occurs at ever more positive potentials as the solution pH is decreased. This rearranged layer is suggested to result from the expulsion of water molecules and

2033

ions from the adsorbed layer. The catalytic hydrogen process is shown to be dependent on the adsorption of vitamin B12ron the electrode surface. Above pH 4.5, the catalytic hydrogen process occurs simultaneously with the surface reduction of vitamin BlZr.Below pH 4.5 the process shifts to more positive potentials and probably involves protonation of the benzimidazole nitrogen.

LITERATURE CITED (1) H. P. C. Hogenkamp, Ann. Rev. Biochem., 37, 225 (1968). (2) H. A. Barker, Biochem. J., 105, 1 (1967). (3) H. Rudiger, Eur. J . Biochem.. 21, 264 (1971). (4) H. P. C. Hogenkamp and S.Holmes, Biochemistry, 9(9),1886 (1970). (5) H. Diehl and B. Jaselskls, J . Am. Chem. Soc., 76,4345 (1951). (6) R. L. Birke et al., J . Electroanal. Chem., 52, 237 (1974). (7) H. B. Mark, Anal. Lett., 912, 203 (1976). (8) D. Lexa and J. M. Saveant, J . Am. Chem. Soc., 98(9),2652 (1976). (9) H. B. Mark et al., J . Am. Chem. Soc., 98(9),2469 (1976). (IO) D. Lexa et al., J . Am. Chem. Soc., 99(8), 2786 (1977). (11) D. W. Imhoff, Ph.D. Dissertation, Ohio State University,Columbus, Ohio,

1966. (12) R. B. Fulton, Ph.D. Thesis, University of Minnesota, l968. (13) D. C. Grahame, J . Am. Chem. Soc., 71, 2975 (1949). (14) E. Laviron, J . Nectroanal. Chem., 52, 355-395 (1974). (15) S.L. Tacked and J. W.Ide, J . Electroanal. Chem., 30, 510 (1971). (16) H. Diehl, R. Sealock, and J. Morrison, Iowa State Coll. J . Sci., 24,433 (1956). (17) D. C. Hodgkin, Proc. R. SOC.London, Ser. A , 6485 (1965). (18) E. Laviron and L. Roullet, Bull. SOC. Chim. Fr., 5077 (1968). (19) S.Mairanovsky, "Catalytic and Kinetic Waves in Polarography", Plenum Press, New York, 1968,p 261.

RECEIVEDfor review October 26,1978. Accepted July 23,1979.

Determination of Stibine in Air with Pyridine-Silver Diethyldithiocarbamate Scrubber and Flameless Atomic Absorption Spectrometry J. B. Cross Phillips Petroleum Company, Research and Development, Bartlesville, Oklahoma 74004

A method to quantitatively measure stibine (SbH3) In air is reported. The stibine is removed from the air and trapped In a pyridine-silver diethyldithiocarbamate solution and the antimony content of the solutlon determined by flameless atomic absorption spectrometry. The sensitivity of the method is suitable for the measurement of stibine concentrations well below the threshold llmitlng value of 0.5 mg of stlblne/m3 air recommended by the American Conference of Government Industrial Hygienists.

The use of antimony compounds as metal passivation agents in hydrocarbon processing catalysts has created a need for an analytical method to measure the toxic antimony compound stibine (SbH3)in air ( I ) . The formation of stibine in industrial working environments can present serious toxicological hazards unless the necessary safety precautions are taken (2). To the author's knowledge, the measurement of stibine in air has not been reported previously; however, it is known to behave similarly to other metal hydrides, e.g., arsine, with respect to the complexing agent silver diethyldithiocarbamate (SDDC) ( 3 ) . SDDC dissolved in pyridine is commonly used as a scrubber solution for arsine in both colorimetric and

atomic absorption determinations; consequently, the reagent was investigated as a scrubber for stibine. For a solution to be an effective scrubber (1)it must provide quantitative recovery for the species of interest, ( 2 ) the solution formed must be stable, and (3) the species of interest must be easily measurable. Other desirable features of a scrubbing solution are that the reagents be free of contamination and that the solution have a long shelf life. The pyridine-SDDC scrubber solution meets these requirements and is used as the trapping reagent for arsine in the ASTM procedure for arsenic (3). Formation of an antimony-SDDC complex similar to the aresenic complex made this solution a likely candidate for trapping stibine. The American Conference of Government Industrial Hygienists recommended threshold limit value (TLV) for stibine in air is 0.5 mg per m3 of air ( 4 ) . A usable method for monitoring stibine should be capable of measuring levels lower than the TLV. The method described below provides this analytical capability. The method uses a pyridine-SDDC scrubber for stibine followed by analysis of the scrubber solution for antimony by flameless atomic absorption.

EXPERIMENTAL Reagents. The pyridine-SDDC scrubbing solution is prepared by dissolving 0.5 g of reagent grade silver diethyldithiocarbamate

0003-2700/79/0351-2033$01.00/0 0 1979 American Chemical Society