Trace Chloride Determination by Rate Controlled Coulometric Titration Mlchael J. Zetlmelsl' and Davld F. Laurence Tretolite Division, Petrolite Corporation, 369 Marshall Avenue, St. Louis, Mlssouri 63 I 79
A new type of coulometrlc Instrument whlch uses feedback from the sensor to the generator In such a way that the current Is a decaylng exponentlal functlon of the concentration of chlorlde Is descrlbed In detall. Mathematlcally, H Is shown that for thls Instrument dE/dt Is a constant at a good dlstance from the equlvalence polnt and equal to one hatl the InHlal constant value In the vlclnlty of the equlvalence polnt. Uslng conventlonal cell components wlth the Instrument, trace chlorlde determlnatlons are performed automatlcally on samples wHh concentratlons as low as 2.8 X lo4 M. A technlque requlrlng some manual control Is descrlbed for samples down to 5.6 X 10-e M. The preclslon of the data Is wlthln a few percent of the relathre standard devlatlon. The devlatkns from the known values are no worse than 5 % for the samples down to 2.8 X lo-' M. The relatlve accuracy gets worse for the more dilute samples. Common Interferences to argenthnetrlc analysk are obviated by means of preoxldatlon wlth KMn04.
Since the time of Lingane's ( I , 2 ) original work on coulometric determination of chloride in the early fifties, determination of chloride by coulometric procedures has become widely accepted. The lower sensitivity level for C1- has been extended into the nanoequivalent region by the use of special instrumentation and cells. The use of feedback circuita, such that silver ions are produced only as needed, dates back to a t least 1960 (3). The method involved use of 70% acetic acid as electrolyte to suppress Ksp levels. Work at the nanoequivalent level was reported by Ladrach et al. ( 4 ) ,but success of the method depended on precise knowledge of the end-point potential which changed appreciably with small amounts of water added to the acetic acid electrolyte. Bishop aad Dhaneshwar employed differential electrolytic potentiometry to determine C1- at the subnanogram level at about 20% accuracy and with special manipulation of the sample and the use of a 0.5-mL microcell (5). The use of organic solvents as electrolyte has been widely investigated, and seems to be gaining universal acceptance (6, 7 ) . Cedergren also has reported data in the nanoequivalent range, but his method at these lower levels entails making a curve each day to obtain a correction for the drift a t the electrodes ( 7 ) . The lowest concentration of chloride in the samples analyzed by Bishop and Cedergren was 5 X M. McCracken has developed an instrument in which a feedback signal which is proportional to the first time derivative of the indicator electrode potential continuously varies the electrolysis current. Data in the 1-100 pequiv range are reported (8). In this paper, we would like to report C1- analysis with a new type of instrument which uses feedback from the sensor to the generator in such a way that the generator current decays exponentially as needed. As pointed out by the other authors, the benefits of this type of instrumentation arise from a better approximation of equilibrium at the end point, avoidance of overshoot, and ease of performing cumulative titrations in the same solvent. The range of data that we have investigated extends down to a lower sensitivity of 0.2 ppm C1- in water or about 6 X lo4
M. The method is fully automatic for concentrations down to 1 ppm, and requires a minimum of manual control for samples more dilute than this. No standardization or blank correction is made and macro titration cells with conventional components are employed. Minor drift of the electrodes is no problem. Moreover, multiple titrations can be performed in the same electrolyte, each titration taking only a few minutes to perform. In short, this technique is a rapid, simple, and automatic method of routinely analyzing for C1- a t trace levels.
PRINCIPLE OF THE INSTRUMENT Whereas other coulometric instruments have worked on either a constant current or a constant potential principle, this instrument works on the principle of constant rate of change of potential with time. The instrument supplies a current which is proportional to the antilog of the voltage difference between the cell potential and some bias voltage which is selected to be near the equivalence point. When a large amount of C1- is present, the cell potential is far removed from the bias voltage, and a large current passes. As the C1- is titrated and its concentration is reduced, the cell potential increases and the titration current level exponentially decreases. Consequently, the titrator approaches the equivalence point slowly near the end of a titration, thereby reducing the possibility of overshoot. Moreover, since current levels are low near the end of a titration, only very slight errors are encountered when Ehiasis set considerably different from the equivalence point potential. A block diagram of the apparatus is given in Figure 1. A differential input amplifier receives the voltage signals from the reference, E , and indicator, E", electrodes respectively and provides an output voltage equal to the difference between the two, E" - E = E. This potential difference will be directly proportional at all times to the logarithm of the C1- concentration through the Nernst equation as follows: E = Eo - (RT/nF)In [Cl-1. A comparator receives the voltage output from the amplifier and subtracts from it some arbitrary reference voltage, Ebias. This output of the comparator, E Eblas,is applied to an antilogarithm converter in such a way that the output of the antilogarithm converter is proportional to the comparator output, i = i, exp ((-nF/RT) ( E - Ebias)) where i, is a small preset current of a few microamps. The current output is passed through the generator section of the cell. The total amount of current passed during a determination is registered in the integrator of the instrument. The experimental relationship between current at the working electrode and cell potential brings about a situation in which silver ions are generated according to the demand for them by the presence of excess chloride ions. Moreover, the chloride is titrated at a constant rate as reflected in the derivation presented in the following paragraphs. In these determinations, we are concerned with the coulometric generation of silver ions which then titrate the chloride ions according to the following equation: Ag+ + C1AgCl. In order to explain how the instrument can process the chloride at a constant rate, it is convenient to consider first the situation that exists a t a good distance from the equivalence point; i.e., chloride decreasing in concentration
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ANALYTICAL CHEMISTRY, VOL. 49, NO. 11, SEPTEMBER 1977
1557
MU L t VOL TS/MlNU TE
+
MILL /AMPERE SECONDS
CATHODE
Flgure 1. Block diagram of instrument
on a 1:l basis with the silver ion being produced. For this situation, we can write the following Nernst equation:
As described by MacInnes (9), a more complicated expression is needed to describe adequately the concentrations of silver and chloride in the vicinity of the equivalence point of this type of titration. Exactly at the equivalence point, of course, the following holds: [Ag'] = [Cl-] and [Ag+][Cl-] =
where [Cl-1, = initial chloride ion concentration, [QIFV] = amount of charge that has passed, F = Faraday constant, and V = volume.
Kp.
[Cl-], - Q = ex$-(E-nf FV RT
- E")]
Equation 1 can be differentiated with respect to time:
But, dQldt = i, and, substituting Equation 2 into 3
Let (RT/nfl 1/FV = K. Built into the instrument is Equation 5:
(5) SO:
Because silver ion is the generated species, it is convenient in describing the situation near equivalence to discuss a point infinitesimally past equivalence, and look at the cell potential as a function of silver ion concentration. At this hypothetical point infinitesimally past equivalence, the chloride ion concentration will be infinitesimally less than and the silver infinitesimally greater than
6.
The situation of being slightly past equivalence is brought about by adding silver ions coulometrically. Some of these silver ions stay in solution as excess and some cause more chloride to precipitate. Consequently, the increment in the equivalents of charge, AQ/96 500, can be described aa the s u m of the silver that stays in solution as excess and that which precipitates more AgC1. The part of the silver that precipitates will, of course, precipitate C1- in the same amount and consequently that part of the silver can be equated to the amount of Cl- that precipitates, A[Cl-], [AQIF = A(equivdents Ag+)] [A(equivalents Cl-)]. But, [A(equivalents Ag+)]/ V = A[Ag+] and [A(equivalents Cl-)]/V = A[Cl-] where V = volume. Therefore, AQIF = V(A[Ag+] + A[Cl-]) and
+
AQ - A [ Ag'] A [ C1-] = -
FV
(9)
Substituting Equation 9 into 8
dt -
(7)
Equation 7 says that, at a considerable distance from the end point, d E l d t is constant for any given value of Ebias. Moreover, the value of dE/dt can be made greater or smaller by changing the value of Ebias. This value, however, should be chosen as close to the equivalence point as possible. 1558
ANALYTICAL CHEMISTRY, VOL. 49, NO. 11, SEPTEMBER 1977
Solving the quadratic equation:
A[Ag*l = -G+ AQ/(2FV) Taking the first derivative of Equation 11with respect to time, and rearranging:
400
In the vicinity of the equivalence point AQ has become very small, and the second term in the parentheses of Equation 12 approaches zero. Consequently,
350
12.5
-
or
We can now write a Nernst equation describing the cell potential as a function of [Ag+].
RT E = E”’ + -In [Ag’] nF
Since [Ag+] = [Cl-] a t equivalence, we can write 2oo‘
b
I d
i b
Flgure 2. Plots of E , In(i),
and finally,
RT
Comparing, Equation 19 with Equation 7 , we see that dE/dt, the rate of change of potential with time at equivalence, is the value of dE/dt a t a good distance away from equivalence. Consequently, we can begin a titration and determine dE/dt at the start of a titration. When the starting value of dE/dt has decayed to half the original value, the determination is complete. There is no need for a description of the situation much beyond the equivalence point since the generation of silver ions essentially stops a t equivalence owing to the fact that E - &!bias is zero a t this point and i = i,. Proving that the potential changes at a constant rate with time is also proof, of course, that the chloride is disappearing at a constant rate with time. Considering the first derivative of the Nernst equation for [Cl-]
dE- - -R T d In [Cl-] nF
dt
dt
a t 25 “C and for a 1-electron change
dE _dt
-
-0.026
d In [Cl-] dt
=
-0.059
d log [Cl-] dt
ii
+
Q 6‘
I
I b ’-2 5
Time (each unit = 9375 sec 1
Substituting Equations 14 and 17 into 16:
d E 1 dt 2
5
(21)
This says that for each 26-mV change in potential per unit time, the [Cl-] ion correspondingly decreases by a factor of “e” (59 mV per factor of 10) in the same period of time. A plot of actual experimental data for E vs. t and In i vs. t is given in Figure 2. The consequent decrease in [Cl-] is
In[Cl-] vs. time
calculated in the same figure. As we can see from Figure 2, E is a linear function of time except at the beginning and end of the titration. The former deviation is due to time lag at the sensor electrode which prevents the sensor from instantaneously seeing the change in C1- due to the current that has passed. The latter deviation is due, of course, to the fact that the current becomes very small when E > In other words, there is no longer any significant demand for current. The deviation from linearity in the beginning of the C1- curve is due, of course, to the deviation in the potential curve. The important point to note from these curves is that the current line is parallel to the C1line. This parallelism is reflected in the following equations describing the In (i) and In [Cl-] functions with potential
nF In (i) = In (i,) - -(E - Ebbs) RT nF In [Cl-] = --(E - E”) RT
(22)
Since E is a linear function of time and In (i) and In [Cl-] are inverse linear functions of time, In (i) and In [Cl-] are inverse linear functions of E. This is reflected in Figure 3 where actual data for In (i) vs. E is plotted and In [Cl-] vs. E is calculated. The failure of the two lines to superimpose is explained by the difference in the constants of the two equations. Also, the deviation from linearity in the current line is explained once again by time lag in the sensor electrodes. If there were no time lag, the In (i) and In [Cl-] lines would be parallel from the beginning. In performing a titration, an is set and current begins to pass according to Equation 5. The passage of current causes decay in the C1- ion concentration, causing an increase in cell potential. When the cell potential has reached the value of Ebias,the current is equal to i,, and at this point the current automatically and very rapidly decays to a value approaching ANALYTICAL CHEMISTRY, VOL. 49, NO. 11, SEPTEMBER 1977
1559
T i m e (each unit = . 9 3 7 5 s e c . ) 2,51
y
j
y
5
p
$i
7
1 ‘p
lag a t the electrode. Some knowledge of the standard Sshaped curve is helpful in deciding this setting. If the break is very sharp, a setting anywhere along the sharp portion, and preferably below the midpoint of the break can be used. If the nature of the standard curve is not known, one can set a very high Ebias, let the machine pass current until the potential levels out, then back off until the current nulls to i,. One or two quick pre-titrations of this type can determine Ebim. Because of the decaying nature of the current, relatively large errors in the end-point setting will lead to very small errors in the Cl- determination as long as a value on the steep portion of the S-shaped curve is chosen.
EXPERIMENTAL
-2.5j
250
L
I
300
350
Cell Potential, E h v )
Flgure 3. Plots of In(/), In[Cl-] vs.
cell potential and time
zero. An Ebiss setting of 350 was used for the data generated in Figures 2 and 3. T o get an idea of how closely these curves approach ideality, as shown in Figure 3, we can calculate the current line using the i, value which is set in the instrument. The lack of correspondence is explained by the approximately half second time lag between what the sensor sees and what the working electrode is doing. This lag causes an insignificant error at the minute current levels a t the end of a titration. Moreover, this error could be made even less by slowing the titration down by choosing a smaller Ebim or lowering the value of i,. Using the approximate calculated concentration of C1at the end of the titration, a Ksp of about would be calculated, in close agreement with published values of K,, for AgCl in 90% acetone (10) which was the approximate composition of this electrolyte. In short, the advantage of this method is that relatively large currents pass a t a good distance away from the equivalence point, and decrease to very small values as the equivalence point is reached. The error due to time lag at the electrodes and the resulting end-point overshoot is minimized. Consequently, titrations in nonaqueous media can be performed without worrying about poor transport properties and endpoint overshoot. Moreover, since the titrations essentially stop a t the endpoint, cumulative titrations can be performed. In using a nonaqueous electrolyte and adding aqueous samples, the end-point potential will change slightly from sample to sample due to changes in the Ksp. It should be clear to the reader that the ideal E b l m setting is at, or a little before, the equivalence point potential. Slight anticipation of the equivalence point in the setting helps compensate for time 1560
ANALYTICAL CHEMISTRY, VOL. 49, NO. 11, SEPTEMBER 1977
Apparatus. The instrument used in these determinations is the Petrolite “Iontrak Analyzer”, available from Petrolite Instruments, 4411 Bluebonnet, Suite C, Stafford, Texas 77477. The data reported in this paper were determined using two basic cells, designated cell A and cell B. Cell A consisted of a 100-mLBerzelius beaker with the electrodes mounted in a Teflon cap, Part No. 180751, available from Beckman Instruments. The generator cathode was a platinum wire in a glass sleeve with fine frit, and a planar spiral of silver wire about 1 cm in diameter was the anode. The anode spiral was placed about 2 mm from the frit. A silver wire was the indicator electrode and the Orion Model 90-02-000 electrode, or a low impedance custom-made glass electrode from Markson was the reference. Cell B consisted of the Brinkmann cell cap, EA874, and the titration reservoir No. EA875-50 or EA875-20. A glass sleeve with fine frit was mounted in a plastic sleeve No. 20-21-570-4,available from Brinkmann. A hole was drilled into the side of the sleeve to accommodate the silver anode of the same type as used in cell A. Platinum was once again the cathode. Brinkmann No. EA242 silver electrode was the indicator and the Brinkmann double junction electrode was the reference. Alternatively, the Beckman glass electrode in a porous Vycor test tube, filled with water and a few drops of 35% HCIOl was used as a reference. The internal portion of the Brinkmann double junction reference was filled with a saturated solution of LiCl in methanol or KCl in water. The external chamber was filled with titration electrolyte. It was our experience that a few simple cell maintenance procedures should be followed at least daily. The anode should be cleaned by dipping in concentrated nitric acid solution for a few seconds and then rinsing thoroughly in distilled or deionized water. The isolation frit should be soaked in nitric acid overnight, and rinsed thoroughly before use. The indicator should be wiped with a soft tissue when precipitate begins to flocculate at its surface. The solvent should be titrated with standard solution twice after any such manipulation of cell components. Reagents. The electrolyte was acetone plus 5 drops of 35% HC104. Cell A used 50 mL of solvent whereas cell B used 20 mL. The isolation sleeve was filled about half way with acetone and one drop of 35% HC104was added. Other organic solvents could doubtlessly be used as long as the K,, of AgCl in them is low (5-7). The same electrolyte can be used for up to 20 titrations, or until the acetone becomes appreciably wet. Standard C1- solution, prepared by accurately weighing 0.0850 g dry, reagent grade NaCl into 100 mL deionized water, was pre-titrated several times in fresh electrolyte before samples were titrated. Standard C1solution was added to the titration vessel with a Hamilton microsyringe, CB-700-200. Preoxidation of Interferences. Common interferences such as Sz-, SH-, CN-, and SCN- were conveniently handled by means of preoxidation with KMnOL under acidic conditions. The sample with interferences was added to the cell, and 7-10 additional drops of acid were added. With the electrodes out of the solution, 170 KMn04was added until a pink color persisted. The excess MnOL was discharged by adding 1% ascorbic acid until the color disappeared. Ascorbic acid was used instead of the more conventional oxalic acid because the latter reacted with Ag+ in acetone. One precaution with the preoxidation procedure is that the acetone must have a few percent water for good results. The addition of extra aqueous acid takes care of this problem. Instrumental Procedures. For samples with 1 ppm C1- or more, the analyst can leave the machine unattended during a
Table I. Determinations of C1- in Aqueous Samples Using Cell A Cl- taken,
a
P€!
No. of determinations
Mean, p g
50.4 50.4 50.4 50.4 25.2 25.2 25.2
7 9 10 10 60 6 29
52 52 49.7 48.3 25.95 26 25.2
SD = Standard Deviation.
@sa
RSD,
SD,
1.71 1.52 0.75 1.49 0.57 1.14 0.25
Double junction Double junction Double junction Double junction Glass electrode Glass electrode Glass electrode
3.3 2.9 1.5 3.1 2.2 4.4 1.0
Ag/AgCl Ag/AgCl Ag/AgCl Ag/AgCl
RSD = Relative Standard Deviation.
Table 11. Determination of C1- in Aqueous Samples Using Cell B C1- taken, No. of Pg determinations Mean, fig SD, P g
RSD, %
Reference electrode Double junction SCE Double junction SCE Double junction SCE Double junction SCE Double junction SCE Double junction SCE Double junction SCE Double junction SCE Glass electrode in porous Vycor test tube Glass electrode in porous Vycor test tube Glass electrode in porous Vycor test tube Glass electrode in porous Vycor test tube
53 53 53 53 53 10.3 10.3 5.15 10.3
15 6 6 9 12 5 5 6 6
54.3 52.4 55.25 53.37 54.63 10.1 10.09 5.4 10.2
0.75 0.38 0.94 1.05 0.13 0.48 0.25 0.25 0.43
1.3 0.7 1.7 2 0.24 4.9 2.4 5.7 4.2
10.3
10
9.8
0.52
5.3
10.3
11
10.26
0.68
5.7
9
5.10
0.14
2.7
5.15
Reference electrode
%b
titration. The only setting to be determined is Eblas,which, for concentrated samples, can vary over the range, 10G200 mV. can be determined by running a few trial auto-titrations of C1standard and finding the setting for where the current nulls to zero at the end point. A potential close to the equivalence point potential can also be chosen for Ebias, if that potential is known. Using the Orion reference electrode, Eblmsettings in the 270-350 mV range were found optimum. With the Brinkmann electrode, using nonaqueous filling solutions,potentials of 375-425 mV were found optimum. With the glass electrode, a setting of -125 to -150 mV was used. For samples with much less than 1 ppm, observation of the change in dE/dt was used to indicate the end point. Dilute solutions require addition of such large samples that we end up with high K,, values and small potential displacements. Consequently, we are fairly close to equivalence even at the beginning of a titration, and we must titrate very slowly or we will overshoot. After adding sample and letting the electrodes equilibrate, the titration is run with a bias voltage sufficient to give an initial rate of 10-20 mV per minute. This initial rate of change, dE/dt, is the maximum, constant rate at which the instrument settles. When the rate has decayed to half the initial value, the titration is complete. Several pretitrations were performed before data collection began. RESULTS AND DISCUSSION Representative data of C1- analyses using cell A are given in Table I. To put these data in the context of a field sample, it is possible to use up to a 10-mL aqueous sample in 50 mL of electrolyte. Consequently, the data presented here represent about 5 ppm C1- (1.4 X M) in the sample. For samples where greater sensitivity is required, a smaller cell such as the one discussed below with Table 11, or the rate end-point method can be used. The data in Table I were generated using either a double junction Ag/AgCl reference electrode filled with aqueous KC1 interior, and aqueous NH4N03 exterior solutions or a glass electrode. Whereas the glass electrode was noisier than the double junction Ag/ AgCl electrode, it had the advantage of absolutely eliminating the possibility of C1- leak.
Table 111. Determination of Cl- in Aqueous Samples Using Cell B and Various E,, Settings No. of deCltertaken, minpg ations 53 53 53 53
5 2 4 4
Mean, pg
SD
RSD, %
Ebias
55.4 54.8 55.5 55.5
1.82 3.3 0.7 1.58
3.2 6 1.2 2.8
350 400 450 500
In Table 11, data are reported from cell B and two different reference electrodes. Comparing the deviations from the known values and the precision of the data from the two cells for C1- values of about 50 pg, it is concluded that this cell is slightly better than the 50-mL cell. Moreover, because of the smaller volume and consequent higher C1- concentration for a given C1- assay, it was possible to analyze samples down to 5 pg C1-. When the series at 5 pg C1- was performed, a total of 5 mL of water had been added to the cell. Therefore, these results would represent aqueous samples with about 1ppm C1- (2.8 X M), about a factor of 5 more sensitive than the data in Table I. The deviations from the known values are no worse than 5% for any of the data in Tables I and 11,even the less concentrated samples. Tables I and I1 contain many sets of multiple determinations under identical conditions. The reason for this repetition was to convince ourselves of the reproducibility of our methods. A period of at least a day, and in some cases several months, elapsed between the data gathering represented in each line of Tables I and 11. Both the double junction SCE and a glass electrode were used with cell B with comparable results. A 17-mm porous Vycor test tube filled half way with water and then two drops of 35% HC104 offered an aqueous bridging solution between the glass electrode and the electrolyte. This configuration ANALYTICAL CHEMISTRY, VOL. 49, NO. 11, SEPTEMBER 1977
1561
Table IV. Determination of C1- Using dE/dt Mode on Aqueous Samples Such That Titration Cell Contains 33.3% H,O-66.6% Acetone ugCI-lmL taken
Sample size. m L’
determinations
p gcl-lm L
SD
1.04 0.207 0.202 0.207
50 50 25 25
14 8 12 15
1.08 0.204 0.25 0.232
0.02 0.036 0.008 0.009
No. of
Table V. Determination of C1- in Presence of S 2 - , SH-, CN-, and SCNc1taken, No. of Mean, RSD, Mg determinations Pg SD % 50.1 25.2 50.2 50.2 50.2
35 60 8 8 13
50.8 26 53.65 51.43 51.43
1.25 0.57 3.85 3.12 1.97
removed some of the electrical noise from the glass electrode, and prevented dehydration of the membrane. The porous Vycor junction offered very low impedance and low leak rate. As pointed out above, the leeway in setting Ebim is a function of the total potential displacement. From Equation 5, it is obvious that Ebias must be greater than the cell potential at the start of a titration and less than, or equal to, the cell potential at the end in order for the current to approach zero. Consequently, the range of possible choices for E b i m is greater for a larger potential break than for a smaller one. To illustrate the invariance of the results when the break is large, the data in Table I11 were generated. The potential break for this determination is 250 mV. As we can see, the results were the same for &ias settings over a 150-mV range. The only caution is that the first titration after changing E b i a s should not be used, because this titration sets the finishing point of a determination. Whether that point is exactly the same as the equivalence point is not important after the first titration because it will take exactly the amount of C1- in the next sample to bring us back to the finishing point. The end-point can also be located by observing a decrease in dEJdt. This method depends on the fact that potential is a linear function of time as described above. As the end point of a titration approaches, the change of potential with time approaches one half the original value, and we can use this point as an indication of the end point of a determination. The way that we actually monitor the linear E vs. t region is by observing the region where dEJdt is constant, and then continuing to pass current until dEJdt decays to one half of its initial value. Data generated using this mode of operation are summarized in Table IV. The enhanced sensitivity comes from the fact that we can analyze samples that cause a very small
1562
ANALYTICAL CHEMISTRY, VOL. 49, NO. 11, SEPTEMBER 1977
RSD,
Mean. ~~
2.47 2.20 7.18 6.07 3.83
~~
SZ0 0 106.2 106.2 106.2
%
1.54 17.57 6.51 3.85
Interferences present, pg SHCN 0 0 0 0 297
0 0 52 52 52
SCN0 0 0 51 51
potential displacement, 10-20 mV. Note that the sample size used in Table IV was 25-50 mL. A 100-mL Berzelius beaker was used for the smaller samples and a 200-mL beaker for the larger. Such large aqueous samples give us an appreciable amount of C1- ( ~ pg) 5 even for samples of very low concentration such as 0.2 ppm. These data demonstrate another order of magnitude sensitivity over the data in Tables I and 11,and concentrations comparable to those reported by Bishop and Cedergren (5, 7). The precision and accuracy of these determinations is not as good as for the more concentrated samples in the auto-mode of operation. Interferences common to argentimetric analysis such as S2-, SH-, CN-, and SCN- have been successfully handled by means of preoxidation with KMn04. Data from these analyses are reported in Table V.
ACKNOWLEDGMENT The authors acknowledge the many enlightening discussions with R. R. Annand of Tretolite Division, Petrolite Corporation, and J. M. Victor of Petrolite Instruments. LITERATURE CITED J. J. Llngane and L. A. Small, Anal. Chem., 21, 119 (1949). J. J. Llngane, Anal. Chem., 26, 630 (1954). D. M. Coulson and L. A. Cavanaugh, Anal. Chem. 32, 1245 (1960). W. Ladrach, F. Van de Craats, and P. Gouverneur, Anal. Chim. Acta, 50, 219 (1970). (5) E. Bishop and G. Dhaneshwar, Anal. Chem., 38, 726 (1964). (6) T. S. Propokov, Mikrochim Acta, 401 (1968). (7) A. Cedergren and G. Johansson, Talanta, 18, 917 (1971). (6) J. E. McCracken and J. C. Guyon, Chem. Instrum., 3 , 311 (1972). (9) D. A. MacInnes, “The PrinciDles of Electrochemistry”, Dover Publications, New York, N.Y., 1961, p 314. (10) G. J. Janz and R. P. T. Tomkins, “Nonaqueous Electrolytes Handbook”, Vol. 2, Academic Press, New York, N.Y., 1973, p 123. (1) (2) (3) (4)
RECEIVED for review April 21, 1977. Accepted June 13,1977.
