A limited amount of work was also done with ammonium nitrate. As shown in Figure 1, the sensitivity of peak halfwidth to changes in water content is very much less than for the perchlorate. This is presumably due to the different base strengths of the anions in DMF. This affects the proton exchange equilibrium as shown by the equations
+ HzO %NH3 + H30+ H30+ + Nos- % HNOa + Hz0 Ha0+ + clod- % HC104 + H20 NH4+
(1)
(2) (3)
For a given water content, Reactions 2 and 3 will come to equilibrium at different concentrations which will in turn influence the exchange rate in Reaction 1. The results with ammonium nitrate show that although the method depends on proton exchange with the ammonium ion, independent calibration curves will be required if the procedure is to extend to salts with various anions.
The observed sensitivity of peak half-width to the type of anion appears also to provide a basis for developing a new and powerful method for measuring relative base strengths of anions in nonleveling solvents. Perchlorate is known to be a weaker base than nitrate. It gives a broader peak (Figure 1) which indicates a faster exchange rate in Reaction 1. This in turn results from the fact that Reaction 3 lies farther to the left than Reaction 2. Thus, either the slope or the ordinate at some convenient abscissa in Figure 1 is a parameter for measuring relative base strengths. The method can obviously be generalized to a series of ammonium salts of various anions dissolved in nonaqueous solvents of sufficiently high dielectric constant to ensure adequate solubility and dissociation.
RECEIVED for review June 9, 1969. Accepted September 8, 1969. This work was supported by the Office of Naval Research, Power Program, under Contract Nonr-4008(07).
Thiocyanate Induced Adsorption of Zin,c Ion at the Mercury Electrode George Lauer and R. A. Osteryoungl Science Center, North American Rockwell Corp., Thousand Oaks, Calg. 91360
ANIONinduced adsorption of various white metal ions at the mercury electrode has been reported in a number of recent papers ( I ) . We report here on the thiocyanate induced adsorption of zinc, The adsorptive behavior of zinc ion is distinctly different from the previously studied metals in a number of ways. Two qualitative theories rationalizing the induction effect of anions on the adsorption phenomenon have been presented ( I , 2). The data which we present here fit neither completely and appear to give credence to portions of both theories. We also report on the coadsorption behavior of zinc and cadmium. We believe that this is the first quantitative study of competitive adsorption to be reported in the literature.
COMPUTER
1
11'
1 1-L
REFERENCE
1 7
COUNTER ELECTRODE
INDICATOR ELECTRODE
I
1 AN+bOG DIGITAL CONVERTER
Figure 1. Block diagram of instrumental arrangement. A l , A2, and A3 are analog devices, Type 201, operational amplifiers
EXPERIMENTAL
Solutions were made up using triply distilled water and A. R. grade chemicals without further purification. The computer system has been described previously (3). A block diagram of the interconnections is shown in Figure 1. All the data reported here are for 25 "C maintained by circulating water from a temperature bath through a jacketed cell. Measurement of Adsorbed Species. The quantity of adsorbed species at a given potential was measured using the technique of chronocoulometry (4). A Kemula-type hanging mercury drop electrode was used throughout. This elec1 Present address, Department of Chemistry, Colorado State University, Fort Collins, Colo.
(1) F. C. Anson and D. J. Barclay, ANAL. CHEM., 40, 1791 (1968). (2) G . W. O'Dom and R. W. Murray, J. Electroanal. Chem., 16, 327 (1968). (3) G. Lauer and R. A. Osteryoung, ANAL. CHEM.,40 (lo), 30A (1968). (4) D. H. Christie, R. A . Osteryoung, and F. C. Anson, J . Electroanal. Chem., 13, 236 (1967). 1882
0
trode was modified by fitting a 3-inch diameter head with markings at 18" intervals to the micrometer screw; this modification greatly increases the precision with which one can extrude a drop of a given volume. In most cases an electrode area of 0.032 cm2 was used. The chronocoulometric technique requires that the electrode be initially at a potential Ei where no faradaic current flows-i.e., well anodic of the polarographic wave. The potential is then stepped to a point where the faradaic current is controlled solely by diffusion-Le., well on the polarographic diffusion plateau. This potential is maintained for a period of time, T , and then stepped back to the original potential. The integral of the current, Q,is measured as a function of time. During the forward step, 0 < t < T , Q is given by
Q = *nFD"2CoAt u 2 + Q d l + nFp
6
(1)
where Qdl is the charge required for the double layer,
ANALYTICAL CHEMISTRY, VOL. 41, NO. 13, NOVEMBER 1969
Q ~=I JE;Cdiff(E) dE
(2)
2-,
.
/ 02
01
c4
0 3
[ Zn'']
05
06
mM
Figure 2. Adsorption isotherm of Zn(I1) in 0.5F NaSCN-OSF N a N 0 3 E = - 300 mV
DS.
SCE
nFr is the charge required to reduce the species adsorbed at Ei and the other terms have their usual significance. For t > T we define the quantity Qr = Q ( T ) - Q(t). Q, is given by Qr
=
[
2nFD"2CoA l + -
z/n
6 =
"Zr]
6
+ aonFAr +
2/; + G -
dt
Qdc
(3)
Supporting electrolyte: 0.5M NaSCN-0.5MNaN03
(4)
during the measurement and the specific charge q m &C/cm2) was determined using the relation
and 2nFD"'CoA
Qc
=-
d7 l/n
Figure 3. Adsorption of Zn(I1) as a function of electrode potential [Zn(II)] = l.OmM
(5)
and al and a. are constants determined from the approximation as described previously (5). These were determined for each run by least squares regression analysis. They are functions only of r and T and are very close to unity and zero. The quantity of electroactive species adsorbed at the initial potential is given by
where I , and Ib are the intercepts of the Q us. r1'2 and Q,us. 6 plots, respectively. The slopes and intercepts were determined by a least squares regression fit of the data to the proper function of time. In all cases, 100 points were taken on the forward step and 200 on the reverse. In a majority of the experimental runs the data parameter, Q, was sampled at 400-psec intervals. The data reported are the result of ensemble averaging at least three individual runs prior to the regression analysis. Standard deviations for the data in all cases were 0.1 of the value or better. The electronic charge on the mercury electrode, qm, at a given potential was measured at a dropping mercury electrode as described previously (6). In many of the solutions, the quantity of adsorbed material, n F r , was sufficiently great that the adsorption is diffusion controlled on a dropping mercury electrode, and therefore, the influence on q m of the adsorbing species appears as a "faradaic" component. T o overcome this problem the solution was agitated with argon (5) F. C. Anson, J. H. Christie, and R. A. Osteryoung, J. Electroanal. Chem., 13, 343 (1967).
+ kt.
Q/(t)2'3= 0.8515 m2'3qm
(7)
This expression assumes that, in a stirred solution, the faradaic current (due to impurities) is limited by the area only--i.e., concentration polarization is absent. The validity of this relation was determined empirically, by checking the goodness of fit of the data to the equation. Values of qm in solutions of concentrated electrolyte only, where the diffusion problem would be nonexistent, were measured using the procedure of Reference (5) and Equation 7 above, and gave essentially identical values.
RESULTS AND DISCUSSION Zinc is strongly adsorbed at mercury in the presence of thiocyanate ion. A typical adsorption isotherm is shown in Figure 2. There is no faradaic current in the region -200 mV to -900 mV cs. SCE. As has been previously shown (7), the adsorption of Zn(I1) in this medium is strongly potential dependent. A typical plot nFr us. potential is shown in Figure 3. The shape of this plot is unusual in that the quantity of adsorbed species reaches a very sharp maximum. Such behavior has not been previously reported, although O'Dom and Murray ( 2 ) have reported the existence of a very gentle maximum in the quantity of material adsorbed. In order to gain a better understanding of the nature of the adsorbed species, the specific charge on the electrode qln,was measured both in the supporting electrolyte and in the solution containing zinc species, using the technique previously de(6) G. Lauer and R. A. Osteryoung, ANAL.CHEM., 39, 1866 (1967). (7) R. A. Osteryoung and J. H. Christie, J. Phys. Chem., 71, 1348 (1967).
ANALYTICAL CHEMISTRY, VOL. 41, NO. 13, NOVEMBER 1969
1883
I
Table I. Correspondence of Double-Layer Charge Measured by Two Different Techniques 0.5M NaSCN-0.5N NaNOa 1mM Zn, 0.5mM Cd A Q (DPSCC)a E, mV us. SCE pC/cm A q m 2 b pC/cmz 200 225 250 270 290 325 350 375
29.8 30.0 28.9 27.3 23.6 21.4 20.8 20.0
30.2 29.8 28.2 27.1 22.7 21.4 20.9 19.9
Value obtained as double layer charging contribution from double potential step chronocoulometry from E< to 900 mV. * Obtained by subtraction of the charge at 900 mV in the absence of ZnZ+ and Cd2+from the charge at E; obtained in the presence of ZnZ+ and Cd2+. 0
scribed (Equation 7). Figure 4 shows a plot of qm us. potential for the two solutions. Cathodic of the adsorption peak in Figure 3 the charge on the electrode is uniformly more positive by approximately 2 pC/cm2 in the zinc containing solution. However, in the potential region of the peak, the charges are equal and there is a sharp drop in qn at the point where the sharp drop in adsorption takes place. This indicates that there should be a significant peak in the differential double layer capacity at this point. As a check for self consistency, the values of the charge required to charge the double layer, Qd1, were measured using double potential step chronocoulometry and were checked against the difference in charge on the electrode at the initial and the final potentials. The data are given in Table I. The qm value at -900 mV was taken in the absence of Cd(I1). Figure 5 shows the differential double layer capacity for this solution as a function of potential obtained using a standard
I
06V 04 200
I
I
300
400
500
- E rnV Y S
I
I
I
600
700
830
900
SC E
Figure 5. Differential double layer capacity on mercury as a function of potential Solution: 1mM Zn(I1); 0.5M NaN08; 0.5M NaSCN
capacitance bridge and a dropping mercury electrode. There is clearly a peak at 300 mV us. SCE as expected from the qn cs. E plot. Anson et al. (5) have thoroughly investigated the thiocyanate induced adsorption of cadmium. As the polarographic waves of Cd(I1) and Zn(I1) are well separated, it is possible to study the coadsorption phenomena of these two species. This is accomplished by first stepping from the initial [anodic of the Cd(I1) wave] to a point on the Cd diffusion plateau anodic of the Zn(I1) wave. The chronocoulometric data for this step give the quantity of cadmium species adsorbed. A step is then taken between El and a potential well on the Zn(I1) diffusion plateau giving the sum of cadmium and zinc species adsorbed; the quantity of adsorbed zinc is then obtained by difference. Figure 6 shows the interaction phenomenon observed. The concentration of Cd(I1) was
i
25 i
-IC
E=-400mV
C
4:' //
c.i
$
E l e c t r o l y t e Alone
tIC
4 '
0
5
With ImMZnLJI)
x With l r n M Z n ( 8 I ) O S m M
Cd
(II)
+2c
I
1
300
400
I
I
500
600
-EvsSCE
I
I
700
800
I
E
1
1
0 2
I [ZnZt]
Figure 4. Charge on mercury electrode as a function of potential in 0.5M NaN03-0.5M NaSCN 1884
I
I
06
04
,
! r,8
mv'
mM
Figure 6. Adsorption of Zn(I1) and Cd(I1) in 0.5M NaN03-0.5MNaSCN,[Cd(II)] = 1 mM
ANALYTICAL CHEMISTRY, VOL. 41, NO. 13, NOVEMBER 1969
5c
4c
3c
+ N
N
E \
0
p
EO
L Y
I20
I
I
200
303
40C
-EvsSCE
-E
VI
Figure 8. Adsorption peak behavior of [Zn(II)] as a function of the concentration of NaSCN at unity ionic strength [Zn(II)] = 1mM
S C E mV
Figure 7. Adsorptive behavior of Zn(I1) as a function of [Cd(II)] in 0.5M NaN03-0.5M NaSCN, [Zn(II)] = 1mM
kept constant and the concentration of Zn(I1) was varied. It is interesting that the sum of Zn(I1) and Cd(I1) adsorbed is approximately equal to the quantity of adsorbed Zn(I1) by itself (see Figure 2). This equality, however is only approximate and appears to hold only for the dilute solutions of zinc. Even at a Cd/Zn concentration ratio of 1 O : l the zinc effectively competes for available adsorption sites on the mercury electrode. At smaller ratios the quantity of adsorbed cadmium goes essentially to zero at all potentials except a very narrow potential range as discussed below. The adsorptive behavior of cadmium in the presence of zinc in thiocyanate medium was investigated over the whole
E us. SCE 100 125 150 175 200 225 250 275 300 325 350 375 400 450
potential region available between mercury oxidation and the foot of the cadmium wave. The results are shown in Figure 7. The quantity of adsorbed cadmium is essentially zero over the entire potential range except in the potential region where the zinc adsorption drops precipitously. Similar behavior is observed at other SCN- concentrations. The presence of cadmium in the solution does affect the adsorption of Zn(1I) in the vicinity of the adsorption maximum. The effect on zinc adsorption as the concentration of cadmium is increased is also shown. As the cadmium concentration increases the maximum disappears; however, the quantity of adsorbed zinc at potentials on either side of the maximum appears to be invariant with added cadmium. O’Dom and Murray (2) and Anson and Barclay ( I ) have
Table 11. Adsorptive Behavior of Zinc as a Function of Concentration and Potential 0.05M NaSCN 0.95M NaN03 O.lmN Zn 0.2mM Zn 0.4mM Z n 0.75mM Zn qm No Zn E nFr(zn) 4m nFr(Zn) qm nFr(Zn) qm nFr(Zn) qm 22.3 20.3 18.4 16.9 16.0 14.1 12.8 11.5 10.3 9.27 7.33 6.16 4.48
100 150 160 170 180 190 200 200 220 230 250 300 350 400
7.23 8.59 8.89 8.33 8.93 8.83 8.54 9.58 8.40 8.11 8.28 6.70 5.63 4.29
21.1 17.3 16.7 16.0 15.3 14.8 14.3 13.7 13.3 12.8 11.8 9.53 7.53 5.45
14.0 15.8 15.6 15.8 15.8 16.3 15.8 15.4 15.4 14.8 14.2 12.4 10.3 7.92
20.3 16.9 16.0 15.8 15.2 14.6 13.9 13.6 13.4 13.7 11.6 9.75 7.64 5.76
21.3 22.3 22.5 22.3 22.6 21.7 21.7 21.1 21 .o 20.8 19.8 17.6 14.7 11.7
18.8 16.7 16.1 15.9 15.4 14.6 14.4 13.7 13.3 13.0 12.2 10.0 8.03 6.00
24.9 25.8 26.1 26.4 25.7 27.5 25.5 25.6 25.2 24.7 23.9 21 .o 18.1 14.3
ANALYTICAL CHEMISTRY, VOL. 41, NO. 13, NOVEMBER 1969
19.4 16.5 16.1 15.5 15.1 14.8 14.3 13.9 13.5 13.0 12.2 10.2 8.16 6.25
1885
proposed theories of anion induced adsorption. These theories differ primarily in that one considers that the adsorption is induced by adsorbed anions which remain adsorbed on the surface in the presence of the adsorbed metal species whereas the other considers a competition between adsorbed anions and metal complex for available sites on the mercury. The behavior of zinc in thiocyanate cannot be said to be sufficiently understood to add strength to either argument. The salient difference in the zinc system compared to those previously studied is the very sharp peaks observed in the amount adsorbed at anodic potentials. The sharp maximum may, perhaps, be due to the formation of ZnHg(SCN)4--i.e., there may be a light oxidation of the mercury at positive potentials in the presence of the adsorbed zinc and thiocyanate. At the potential where reduction of the ZnHg(SCN)r occurs, to leave the adsorbed zinc-thiocyanate, one could expect a high degree of disorder which would allow the cadmium to occupy some of the available sites, as observed. The potential of the peak is a function, however, of the concentration of SCN- in the solution as shown in Figure 8 and the shift, qualitatively, is in accord with what might be expected if a mercury-zinc-thiocyanate species formed at the surface. There is a distinct cathodic shift with increasing SCN- concentration; this shift does not correlate with the concentration of any of the thiocyanate complexes of zinc using published stability constant values (8).
The difference in charge on the electrode, qm, we believe to be real. (See Table I.) However, the change in the charge (8) R. E. Frank and D. N. Hume, J. Amer. Chem. Soc., 75, 1736
(1953).
qm produced is much less than the quantity of adsorbed ma-
terial. This has been employed as a n argument by Anson and Barclay to indicate that the adsorbed species is neutral. However, Mohilner (9) pointed out that the relation
must be valid. Plots of qm 1;s. log C and nFr cs. E should have the same slope, as indicated by Equation 8 above. From Figure 3, a value of ZF[bI’/bEj‘ of about 5.3 x 10-5 C/V may be calculated. This would lead to a value of the term aQ/blogC, of about 1.6 X C per decade change in concentration of adsorbate in solution. Data in Table I1 were also used to construct such plots; subject to a considerable experimental scatter, values of about 4 X 10-5 CjV were obtained for both plots. Thus, the small change in qm obtained does not necessarily indicate the adsorption of a neutral species. Other arguments presented by Anson and Barclay, however, are in accord with the absorption of a neutral species. The data presented here do not conclusively distinguish between theories. ACKNOWLEDGMENT
We acknowledge the patience and skill of Richard Carpenter who obtained much of the data with enthusiastic resignation to his fate.
RECEIVED for review June 9, 1969. Accepted September 5, 1969. (9) D. Mohilner, Colorado State University, Fort Collins, Colo., private communication.
Determination of Chlorine in Selenium by a Distillation-Atomic Absorption Procedure Wladislaw Reichel and Laszlo Acs Canadian Copper Refiners Ltd., Montreal East, Quebec, Canada
CHLORINE, probably in the form of interstitial chloride ion, has a modifying role in the crystalline growth of selenium. In selenium destined for semiconductor application, the determination of chlorine, both in high purity and chlorinedoped products, is becoming increasingly important. A versatile and accurate procedure was sought to cover chlorine concentrations between several hundred ppm and the lowest attainable detection limit. In the method of T. E. Green ( I ) , based on dissolution of selenium in nitric acid in the presence of excess silver ions, the separation of the precipitated silver chloride and spectrophotometric determination of silver as sulfide results in consistently low recoveries, probably because of volatilization of chlorine during dissolution and/or the solubility of silver chloride in selenious acid- nitric medium. Because these losses were near 20 Hg C1 per 5-gram sample, another approach was needed, particularly for low chlorine concentrations. (1) T. E. Green, unpublished papers, as described by I. M. Kolthoff and P. J. Elving, “Treatise on Analytical Chemistry,” Part 11, Vol. 7, p 198 (1961). 1886
The separation of hydrogen chloride by distillation was chosen to avoid interference from impurities in the sample and to overcome the problem of silver chloride solubility in selenious acid. Chlorine, distilled as hydrogen chloride, was collected in a solution containing excess silver ions. The precipitated silver chloride was filtered off and redissolved in an ammoniacal solution. A number of published procedures describe the determination of silver in an ammoniacal solution (2). However, the sensitivity, specificity, and rapidity of atomic absorption spectrophotometry (3-6) in measuring silver ion concentration lead to the selection of this technique. (2) I. M. Kolthoff and P. J. Elving, “Treatise on Analytical Chemistry,” Part 11, Vol. 7, p 380 (1961). (3) V. L. Ginzburg, D. M. Livshits and G. I. Satarina, Anal. Abstr., 13, 15 (1966). (4) U. Westerlund-Helmerson, Atomic Absorption Newsletter, 5, 97 (1 966). (5) F. M. Tindall, ibid.,p 140. (6) J. B. Ezell, Jr., ibid., 6, 84 (1967).
ANALYTICAL CHEMISTRY, VOL. 41, NO. 13, NOVEMBER 1969
