Table 1.
Sominal Composition of Alloy
Determination of Zirconium in Titanium (Confinued)
Zirconium Added,
% 10.00
Ti-lT5rl
0.50
Fe-1 .75 Cr-3,OO 0-0.50 1\2)-0. 04 6-0.03 w-0,02
Zirconium Found, % p-Bromomandelic acid p-Chloromandelic acid 10.18 10 10 10. 1.5 10 23 10.14 10 12 Av. 10.16 10 15 0.53 0 48 0.53 0 48 0.50 0.50 0.48 0.52 0.50 0.52 0.50 0 51 0.50 0.52 0 51 0.50 Av. 0.505 0 505
1.50
3.00
5.00
1,54 1.50 1.48 1.48 1.48 1.49 1.49 1.48 1.48 rlv. 1.49 3.01 3.02 3.02 3.02 Av. 3.02 5.06 5.04 5.04 5.06 5.09 5.06 5 04 5 06 5.03 Av. 5.05
1 52 1 1 1 1
50 53 48
53 1 51 1 48
1 50 1 50 1 51
3 01
3 03 3 02 3 02 3.02 5 02 5 03 5 03 5 5 5 5 5 5 5
04
04 04 05
06
05 04
Wash the precipitate 10 to 12 times with distilled water, transfer to a platinum crucible, char slowly, ignite a t 1000° C., and weigh. The zirconium oxide can also be weighed by transferring it to a tared weighing dish.
Zryc
=
Zr02 weight X 0.7403 X 100 -
sample weight
RESULTS
Results obtained on commercial titanium alloys are given in Table I. Zirconium was added in the form of a standard zirconyl chloride solution along with the dissolving acid, before dissolution of the sample. The zirconium solution was standardized by the phosphate, cupferron, and mandelate methods. Bny hafnium present is included in the results. LITERATURE CITED
(1) Klingenberg, J. J., Papucci, R. A., A s - 4 ~ CHEM. . 24. 1861 11952). (2) Papucci, R. A., Fleishman; 0. AI., Klingenberg, J. J., Ibid., 125,
1758 (1953). (3) Papncci, R. A , , Klingenberg, J. J., Ibid., 27, 835 (1955). 14) Scott. W. W,.“Standard Methods of Chemical Analysis,” 5th ed., Tan Sostrand, New York, 1939. (5) Tour, Sam, and Co., 44 Trinity Place, New- York, K. Y . , Rept. R9141 (November 1951). (6) Willard, H. H., Diehl, H., “Advanced Quantitative Analysis,” Van Nostrand, Sew7 York, 1943. ~I
RECEIVEDfor review May 17, 1957. Accepted February 3, 1958.
Coulometric Titrations with Mercury(1) and (11) Determination of Sulfide EDWIN P. PRZYBYLOWICZ’ and L. B. ROGERS Department of Chemistry and laboratory for Nuclear Science, Massachusetts Institute of Technology, Cambridge 38, Mass.
)The coulometric determination of sulfide was undertaken because the extremely low solubility product reported for mercury(l1) sulfide appeared to offer an excellent means for testing the ultimate sensitivity of coulometry. Mercury(l1) could be generated in alkaline solution at 100% efficiency by using mercury-coated gold electrodes. By titrating a t 80”C., effects attributed to adsorption could be minimized. Amounts from 0.24 to 1.1 1 mg. were determined in 100 ml. of solution with amounts ; an error of about 3 ~ 0 . 2 7 ~ down to 0.060 mg., with an increasing The lower limit of error up to 2%.
1064
ANALYTICAL CHEMISTRY
the determination was probably set b y the grid-current of the continuousindicating potentiometer as evidenced by behavior in the end point region characteristic of a polarized system.
I
the sensitivity of a potentiometric end point-Le., the magnitude of the potential jump a t the end point-is a function of the solubility product of the precipitate formed. As sulfide fornis a very insoluble precipitate with niercury(I1) ion, it was thought that a coulometric titration in which mercury(I1) ion was N A PRECIPITATIOS TITR.4TIOK,
generated could be used in the titration of very small amounts of sulfide with good precision and accuracy. The potentiometric titration of mercury(I1) ion with sulfide was first reported by Pinkhof (14), who used a mercury indicator electrode to titrate various metals M ith sodium sulfide. Mercury(I1) ion was titrated in a solution which contained a small excess of potassium iodide; the results had a constant 2% negative error which was attributed to adsorption. Kolthoff and 1 Present address, Research Laboratories, Eastman Kodak, Rochester 4,N. Y.
T’erzijl (8) obtained good results in the titration of sulfide with mercury(I1) chloride when the solution m s approximately 0.1F in sodium hydroxide. The maximum change in potential was sharply defined and the inflection point n a s found a t -0.200 volt (us. N.H.E.). A 0.001F sulfide solution was titrated with a n accuracy t o 2 to 3%. The addition of “inert” divalent salts gave negative errors as high as 27,. The present procedure, although limited in scope, can be used for very accurate titrations of small amounts of sulfide. Certain factors d-hich limit the use of potentiometry in titrations of small amounts of sulfide have been studied. EXPERIMENTAL
Reagents. A O.1F sodium sulfide solution mas prepared from hydrogen sulfide gas (Matheson Co.). After 500 ml. of deaerated distilled water had been purged with hydrogen sulfide for approximately 1 hour, 2.0 grams of sodium hydroxide pellets (Mallinckrodt AR grade) mere added and t h e solution was thoroughly mixed. T h e sulfide solutions were stored in amber bottles under a nitrogen atmosphere. Care was taken to exclude all traces of oxygen by using special self-sealing Buna X rubber caps (courtesy of Firestone Industrial Products, Soblesrille, Ind.) t o stopper the bottles. The rubber caps were held tightly in place by sere\$- caps punched with holrs. Solutions prepared and stored in this manner Jvere stable for about 1 week. Samples were withdrawn by means of a calibrated 1-ml. hypodermic syringe mounted in a holder which contained a n adjustable screw in line with the syringe plunger. The screw acted as a stop for the plunger, thus determining the volume of the syringe. This syringe assembly is commercially available from I\lacalaster Bicknell Co., Cambridge, Mass. The syringe was calibrated b y Iveighing the amount of distilled water delivered. Five determinations gave a mean value of 0.9998 ml. with a standard deviation of 0.0011 ml. The sulfide solutions were standardized by the iodate-thiosulfate method (18). I n this standardization procedure sulfite as well as sulfide was titrated. Apparatus. T h e apparatus has been described (16). I n t h e present study mercury-coated gold electrodes were used. I n order t o be able t o heat t h e solution conveniently while t h e cell was on a magnetic stirrer, a n immersion heater of about 180 watts was constructed by coiling 6 inches of Chrome1 C resistance wire (B. and S. No. 28) around the outside of a 3-mm. borosilicate glass tube. This assembly was then placed in a 7-mm. tube which had one end sealed off, and the tube was filled with sand to provide more efficient thermal conduction. The heater was operated from a Variac variable transformer a t about 20 volts. A 100-ml. sample could be heated t o SO” C. in about 10 minutps.
Table
I.
Current, Ma.
20.00 10.00
10.00
Elapsed Time, Sec. 165.6 331.0 165.6
XO. of Trials 3
2
6 3
Std. Dev.,
Amount Taken,
Amount Found,
AV.~ Error:
Mg.
Mg.
Mg.
c ,c
0.0003 0.0002
1.114 1.114
1.097 1,100
-1,52
0.0008
0,0012 138,8 4 0,0006 1.096 160.1 Data not conected for sulfite content of sample. 5.00
a
Typical Titration Data for Sulfide Using Recommended Procedure
The effect of the rate of stirring on the potential Tvas investigated, using a mercury-coated gold mire (0.2 X 0 . 5 em.) sealed into a polystyrene rod (1 0 x 7 5 em.) a t right angles to the axis of the rod and approximately 3 mm. from the end. The electrode was connected by means of interchangeable gears to an 1800 r.p.m. synchronous motor, so that it could be rotated at fixed speeds between 150 and 3600 r.p.m. in increments of 150 r.p.m. The potential of this electrode, as a function of the speed of rotation, was measured us. a saturated calomel reference electrode (S.C.E.) in a solution which was 0 1F in sodium hydroxide and about 1 X lO+F in sulfide. Titration Procedure. A 100-ml. portion of 0.1F sodium hydroxide was deaerated with prepurified nitrogen for 10 minutes M hile t h e solution was being heated t o 80” C. and during t h e course of t h e titration. T h e mercurycoated gold electrodes were then cathodized for 30 seconds a t 10 ma. in the titration medium, following n hich a small amount of sulfide, approximately 0 05 t o 0 15 mi. of a 0 OOlF solution, was added to bring the potential down to #bout -0 5 volt (tis. S.C.E.), a point on the steepest part of the potentiometric curve. When this potential was seriously “overshot,” a small amount of niercury(I1) ion was generated to bring the potential back to -0 5 volt. Finally, 1 000 ml. of a standard sulfide solution was added to the titration cell and mercury(I1) generated until the potential reached the same value it had just before the sample was added. Ordinarily, the solution was stirred with a magnetic stirrer a t a rate of 700 r.p.m., as determined using a stroboscope, because, at slower speeds, the potential readings in the vicinity of the end point were not reproducible. RESULTS
Titrations of sulfide were studied using generating currents from 20 to 0.50 ma. I n the region TI here the precision of the titration was to better than 0.27,. the results ivere consistently low b y about 1.3%, as shown by Table I. This error was shown to be due to the presence of sulfite, hich \vas titrated along with the sulfide in the standardization procedure, \$ hereas. in the coulometric procedure, only the sulfide reacted. By analyzing polarographically (7’) the stock solution? of sulfide, it n as found
0.55il
0.5505
0,2350 0,0606
0.0584
0.2305
-1.23 -1.19 -1.93 -3.66
that they contained between 1.25 and 1.56% sulfite, which may have been produced b y oxidation of sulfide. If a correction of -1.40yG is applied to the quantitative data in Table I, the results for the larger amounts of sulfide are accurate to within 0.2%. Several titrations were carried out a t higher generating currents to ascertain the usefulness of this procedure for larger amounts of sulfide. A t 50 ma., the recorded potentiometric curve indicated the apparent current efficiency of the process was less than 100%. The electrode became visibly coated 11 ith mercury(I1) oxide, even before all of the sulfide was titrated. If the titration was stopped prior to the end point. the oxide coating slowly turned black, indicating that metathesis t o the less soluble sulfide v a s taking place. This slow metathesis resulted in potential. which, in the vicinity of the end point, drifted for about 30 minutes before reaching a stable reading. Thus, for larger amounts of sulfide than those reported in Table I a larger generating current could be used successfully onljwhen the electrode area was made proportionately larger, so as to gire n current density smaller than 2 ma. per sq. em. The alternative of using smaller generating currents for longer periods of time is less attractive. The titration of smaller amounts of sulfide was studied a t a generating current of l ma. Recorded potentiometric titration curves Jvere relatively flat, with the breaks in potential which OCcurred 10 to 20% past the theoretical equivalence point. Furthermore, the potentials in the vicinity of the break were not stable, but drifted in a cathodic direction. This phenomenon )vas considered to be due to the adsorption of mercury(I1) ion b y the freshly precipitated mercury(I1) sulfide (19). If one considers that some sulfide had to be added to the hydroxide solution t o adjust the potential before adding the sample, all solutions had a relatively constant amount of mercury(I1) sulfide present in addition to that resulting from the sample. When the amount of sulfide taken for analysis ]vas large, the amount of mercury(I1) ion adsorbed on this mercury(I1) sulfide was insignificant compared to the amount of generated mercury(I1) ion; for smaller VOL. 30, N O . 6, JUNE 1958
1065
amounts of sulfide, the amount of adsorbed mercury(I1) ion was significant. htanual coulometric titrations could be carried out at a generating current of 1 ma., but it was necessary to wait 1 t o 2 minutes in the vicinity of the end point for the potential to become stable. Even when these results were corrected for the sulfite content of the stock solution, they were 2% low (Table I). At generating currents of less than 1 ma., it was impossible t o obtain a suitable potential break near the equivalence point. Attempts to eliminate the negative blank of the solution b y prolonged cathodization of the mercury-coated gold electrodes succeeded in removing the blank, but did not improve the sharpness of the potentiometric titration curves (compare curves A and B, Figure 1). However, the addition of gelatin (0.1% b y weight) improved the s h a r p ness of the equivalence point break somewhat (C, Figure l), indicating that adsorption was at least partly responsible for the flatness of the potentiometric curve. DISCUSSION
I n t h e absence of sulfide, the anodic dissolution of mercury in alkaline solution causes the formation of mercury(I1) oxide (yellow). T h e presence of anions which form less soluble precipitates t h a n mercury(I1) oxide favors the formation of those precipitates. Polarographic studies of the anodic wave of sulfide ion have shown (6) that even in 0.1F sodium hydroxide the anodic sulfide wave obeys the reaction: Reaction Product.
Hg(s)
+ HS- + OHHgS(s)
=
+ H20 + 2e-
(1)
Llercury(1) sulfide is unstable and, if formed, disproportionates to mercury and mercury(I1) sulfide. I n attempting to determine whether or not one could distinguish between the formation of mercury(1) and (11), volumetric titrations of 100 ml. of a 0.018' solution of sodium sulfide in 1F sodium hydroxide were carried out, using 0.18' mercury(1) nitrate and 0.1F mercury(I1) nitrate as reagents. These titrations were followed potentiometrically using a mercury-coated gold electrode us. a saturated calomel reference electrode. Potential readings during the course of each titration became stable shortly after each portion of the titrant was added. From the virtually identical potentiometric curves that were obtained, it was evident that one could not differentiate between the two valences in that way. Electrode Type. When a t t e m p t s were made t o generate mercury(I1) ions in stirred alkaline sulfide solution from a mercury pool anode, the elec1066
ANALYTICAL CHEMISTRY
I
I
I
I
40
80
120
I I 1 160 200 240 Time (Seconds)
I
I
280
320
Figure 1. Titration curves of 1.64 peq. of sulfide in 0.1F sodium hydroxide a t 80" C. using a current of 1.096 ma. A. Solution potential - 0 . 0 9 0 volt before sulfide sample was added. Electrodes not pretreated prior to titration B . Solution potential brought to -0.300 volt before sulfide was added. Electrodes precathadized for 100 seconds at 1.096 ma. C. Solution potential brought ta -0.300 volt before sulfide sample was added. Titration carried out in presence of 0.1% gelatin. Electrodes precathodized for 100 seconds at 1.096 ma.
trode became coated with a mercury(11) sulfide layer which prevented t h e further precipitation of sulfide in the bulk of the solution by blocking the transport of mercury(I1) ions from the electrode surface to the body of the solution. I n some cases, the current in the generating circuit dropped to zero several seconds after the titration had been started, indicating that the film formed on the electrode was an effective insulator. This same type of phenomenon was observed by Revenda (17), who first studied anions that give anodic polarographic waves, using a dropping mercury electrode. 1Iajer (11, 12) later reported that, in the case of bromide and chloride, the insoluble film on the mercury electrode increased the resistance of the cell by lo6t o lo7ohms. Kolthoff and Miller (6) later reinvestigated Revenda's work and actually measured an increase of 25,000 ohms in the cell resistance which accompanied the formation of a n insoluble mercury(1) bromide film on the mercury electrode. I n contrast to the above findings, which were all obtained on micro dropping mercury electrodes, the present authors have found that when larger stationary electrodes are used for the coulometric generation of mercury(1) or (11), only sulfide forms a film n hich appears to be impermeable. By using a larger surface area, the difficulties of film formation were circumvented in the case of the halides, a n obserration confirmed by the recent nork of DeFord and Horn (3). I n the presence of sulfide, mercurgcoated silver and mercury-coated gold electrodes behaved in a fashion similar to the mercury pool electrodes, but the film did not adhere as readily to the
I
1
amalgam electrode, presumably because of differences in the surface tension of the mercury. Mercury-coated electrodes Ti-ere employed in this study because they were more convenient to handle than the mercury pool. Gold n-as preferred to silver for use as the supporting metal. Though gold(I1) sulfide is knov-n to be very stable ( I O ) , no free
I
- 120
I
I
I
I
I
-80
-40
0
40
W
I
I l2U
I
I 200
160
Time (Seconds)
Figure 3. Coulometric titration of 6.20 peq. of sulfide in 0.1 F sodium hydroxide a t 30' C. 10.00-ma. generating current not interrupted after once started at zero time
3
L
m 0
-I
I 2
u - 0.4 - 0.5 E (Volt
VS.
S C E.)
Figure 5. Relation of log r.p.m. to potential for mercury(ll) sulfide end point Broken line, predicted slope for m =I 0.8
E ( V o l t vs. S.C.E.)
Figure 4. Effect of rate of electrode rotation on potential for mercury(I1) sulfide end point
energy data are a\-ailable concerning its stability. However, it was assumed that because gold metal has a much more noble oxidation potential than mercury, no significant amount of its sulfide could be formed during a titration. Electrode Pretreatment. hlercurycoated gold electrodes were pretreated b y either cathodizing or anodizing the mercury-coated electrodes in 0.1F sodium hydroxide a t 10 ma. for 30 seconds. Results obtained a t 60" C. are shown in Figure 2. Anodizing in 0 . W sodium hydroxide produced a coat-
ing of yellow mercury(I1) oxide on the electrode surface. This layer remained even after the electrode had been rinsed with mater, although presumably some oxide was washed off. When the electrode was placed in 0.1F sodium hydroxide and the solution was deaerated] the oxide film slowly dissolved (10) according to the reaction: HgO(s)
+ OH-
HHg02-
(2)
\Then a given amount of sulfide was introduced into the cell, the reactions HHgOz-
+ HS-
=
HgS(s)
+ 20H-
(3)
+ HS-
=
HgS(s)
+ OH-
(4)
and HgO(s)
could occur to make the subsequent analytical results for sulfide appear lorn. When the indicator electrode was anodized and rinsed before use in a titration of sulfide (Figure 2, A and B ) , the re-
corded potential did not follow the course of the reaction (although visibly a reaction was occurring). This suggested that Reaction 4 \vas very slowl so that the indicator electrode did not respond to the sulfide ion concentration. From this evidence it is obvious that the best possible conditions for the coulometric determination of sulfide involve cathodization of both indicator and generator electrodes prior to the titration. No significant difference n a s observed between electrodes precathodized in the presence or absence of sulfide. Effect of Temperature. Figure 3 shows a curve typical of those recorded a t room temperature. Point d corresponds to the point a t which the sulfide sample n a s added t o t h e titrating medium. T h e decrease of potential to its equilibrium value was very slow and u a s attributed to the fact that a film of mercury(I1) sulfide had to form on the mercury surface before the electrode could properly respond to the sulfide concentration in solution. Once the electrode had reached its equilibrium value, the generation of niercury(11) ions was started. The remaining portion of the curve, recorded during the coulometric generation of niercury(11)ion, does not represent the potentiometric curve expected from calculations based on the solubility product. From the solubility product it can be shown that a t a generating current of 10 ma., the slope, A E j A t . of the potentiometric curve a t the inflection point should be 588 mv. per 0.1 second. I n addition, the inflection point occurred rvhen the solution was anywhere from 120 to 130% titrated. When an exact stoichiometric amount of mercury(I1) ion was generated to react with the sulfide in solution and then the generation discontinued, the potential x a s observed to drift slonly for about 45 to 60 minutes before it came t o a steady value. This indicated that there was some slow step in the titration which ruled out determination of sulfide by continuous titration and automatic recording of the resulting potentiometric curve. Because the kinetics of precipitation of mercury(I1) sulfide in basic solution were rapid, it was concluded that the slow step observed in the coulometric titration was due either to the film formation previously observed, or to adsorption of mercury(I1) ions on the freshly precipitated mercury sulfide. By carrying out the titration a t an elevated temperature and by pretreatment of the electrodes, the effects of both factors were counteracted. Effect of Stirring. I n t h e vicinity of t h e end point (97 to 103% titrated), the potential could be changed as much as 200 mv. merely by changing the rate of stirring of t h e solution (Figure 4). This phenomenon, which had previously been encountered in VOL. 30, NO. 6, JUNE 1958
1067
working with dilute solutions in the bromine-bromide system (16), indicates that the electrode is polarized. Hence, one must conclude that the grid current, though only of the order of 10-12 ampere, is sufficiently large to affect the concentrations of mercury(I1) and sulfide present in the layer of solution adjacent t o the electrode surface. Because the extent of polarization that can result from such a small current is generally thought to be negligible, i t seems worth while to show how this effect can be accounted for from the standpoint of a potentiometric titration a t constant current. If a potential-measuring device such as the Leeds & Xorthrup p H indicator is used, a grid leakage current of about lo-'* ampere is continually passed through the indicator electrode system. If the direction of current flow favors the oxidation of mercury a t the indicator electrode, mercury ions are produced which can react with sulfide. For a classical potentiometric titration of sulfide with mercury(I1) ion in 0.1F sodium hydroxide, it can be shown that the theoretical equivalence point is reached when (HgS) = (HS-) = 1.2 X 10-27 molejliter
(5)
using a value of the solubility product for mercuric sulfide of 1.6 X 10-54 (4) and a value for the ionization of bisulfide of (1.3 X lO-I3) (9). Considering this equivalence point concentration merely as an order of magnitude, it is evident that in the vicinity of the equivalence point the rate of change of potential as a function of added reagent should be very large. At the same time, one can calculate, as shown below, that a concentration of about 10-14F sulfide ought to react completely with the mercury(I1) ion generated by a current of ampere. Thus, the total flow of a substance having a diffusion coefficient, D, through a diffusion layer of thickness, I , to an electrode of area, -4, is given by Fick's first I ~ Kof diffusion : AD
'
f = -(C
1
- C,) =
i nF
(6)
where n is the number of electrons involved in the electrode reaction, F is the Faraday, and i is the current. When the current, i, is approximately equal to the diffusion current, the concentration of bisulfide ion a t the electrode surface, CO, becomes vanishingly small compared to concentration C, so that the last equation reduces to Zd
AD nF-C 1
(7)
The value of C can be estimated by 1068
ANALYTICAL CHEMISTRY
assuming the diffusion coefficient for bisulfide to be equal to that for chloride (6),2.03 X sq. em. per second at 25'; the diffusion layer thickness to be 0.025 cm. ( 6 ) ; the electrode area to be 1 sq. cm; and the indicator current to ampere. At a concenbe 4.2 X tration of 2.6 X 10-14F, bisulfide will reach the electrode at the same rate as mercury(I1) ions are generated. The maximum change of potential should occur a t a concentration of bisulfide of 2 X 10-14F instead of 1.2 X 10-*'F, which one would calculate for an equilibrium measurement. By using the mnvimum change of potential as the equivalence point, a slight negatiye bias is introduced, but this error is negligibly small even in titrations of micromolar amounts of sulfide. Hon-ever, because the thickness of the cliffusion liyer is a function of stirring, the estent of polarization and the stoichiometry for the ma\-imum change in potential will change accordingly. This has been reported for the bromine - bromide system (16), though in the latter case the flow of current through the indicator circuit worked in the opposite direction. One can elucidate further the effect of stirring, knon-ing that the diffusion current of an electrode has been shown to be related to the rate of rotation ( I S ) by the general equation id
= kr"
(8)
where k is conqtant, m is a constant whose value is reported ( I , 2 ) to be between 0.6 and 0.8, and r is the rate of rotation of the electrode. When this is substituted into the Ilkovic equation, one obtains a general expression of the form
where K" represents a collection of constants, i is the constant grid current passed through the indicator electrode, and each of the other symbols has its usual significance. When Icr" is large compared to i, a plot of the logarithm of the speed of rotation us. potential becomes a straight line, as shown in Figure 5. Constants k, m, and i have not been quantitatively evaluated, but the slope of the curve should be RTmInF. Unfortunately, E varies only slightly with r when the latter is large, so the error in determining this slope is rather large. From three points obtained at higher speeds of rotation, the slope was found to be 30 mv. per tenfold change in rate of rotation. The dotted line in Figure 6 has a slope of 24 mv., which corresponds to a value of 0.8 for m. Since completion of this study,
Strange (20) has reported a nonlinear relationship between the logarithm of the sulfide (or cyanide) concentration and the potential on a silver electrode which he attributed to concentration polarization. Higher stirring rates produced the expected linear relationship. For the general case, polarization will usually be most evident in the region of the equivalence point because of the susceptibility of the potential to change by stirring. At points far removed from the equivalence point, stability will be achieved, xhich, however. should not be considered as evidence for absence of polarization (16). Furthermore, it is obvious that the niagnitudes of the exchange current-Le., concentration of the ions and the rate the constant of their reaction-and current density-Le., grid current and size of the indicator electrode-will determine whether or not polarization phenomena will be encountered. Unfortunately, agreement between different studies, even when conducted b y the same investigator, may be complicated by the fact that the exchange current observed for a given solid electrode can easily change between run>, because it is usually gravely affected by impurities in the solution and by pretreatment of the electrode. At the present time, only qualitative agreement can be shown with the expect&tion of finding that the mercury half cell, which has the larger exchange current, exhibits polarization a t loner concentrations than the less rapidly reversible bromine half cell. As a second example, it is worth while to consider the behavior observed for a mercury indicator electrode in acid solution. Unlike the behavior reported above, the titration of iodide with mercury(1) yielded only barely detectable changes of potential in the end point region when a stirrer was alternately started and stopped. If the reaction rates under the two conditions are assumed to be the same, the smaller change for iodide is the result of its larger solubility. Titrations of the more soluble bromide and chloride showed no effect of stirring on the potential. ACKNOWLEDGMENT
One of the authors (E.P.P.) wishes to thank the Eastman Kodak Co. for a fellomhip; the other wishes to acknowledge partial support of the U.S. Atomic Energy Commission. LITERATURE CITED
(1) Brunner, E., 2. physilz. Chem. 47, 56 (1904). (2) Ibid., 56, 321 (1906). (3) DeFord, D. D., Horn, H., .ISIL. CHEII.2 8 , 797 (1956).
(1) Goates, J. R., Cole, .4.R., Gray, E. L., J . Am. Cheni. Soc. 73, 3596
(1951). Kolthoff, I. M., Lingane, J. J., “Polarography,” 2nd ed., Interscience, S e w York, 1952. Kolthoff, I. lf,,Miller, C. S., J . Bm. Cheni. SOC.6 3 , 1405 (1941). Thid., p. 2818. Kolthoff, I. hl., Verzijl, E. J. A. H., 12ec. trav. chim. 42, 1055 (1923). Iiury, J. \V., Zielen, J. G., Latimer,
If-. M., J . Electrochem. Soc. 100,
168 (1953). Latimer, K , &I., “Oxidation Poten-
tials,” 2nd ed., Prentice-Hall, New York, 1952. i l l ) Maier. V.. Collection Csechoslov. C‘hem. Communs. 7. 146 11935). (12) Ibid., 9, 360 (1937). (13) Nernst, W., llerriam, E. S., Z . phys. Chem. 52, 235 (1905). (14) Pinkhof. J.. dissertation. Amsterdam, ’1919. (15) Przybylowicz, E. P., Rogers, L. B., ASAL. CHEM.28, 799 (1956). (16) Purdy, W. C., Burns, E. A,, Rogers, L. B Ibid., 27, 1988 (1955). (17) Revenda, J., Collection Czechoslov. Chevi. C o ? m i i o ~ s 6, . 453 (1934). .
I
\
,
~
(18) Scott, ]IK. , , ‘.Standard Methods of Chemical Analysis,” 5th ed., Van Xostrand, Sew York, p. 1443. (19) Spaulding, G. H., MciYabb, W.II., J . Franklin Inst. 237,207 (1944). (20) Strange, J. P., .ZX.~L.CHEX 29, 1878
(1957).
RECEIVEDfor review May 22, 1957. ..iccepted February 6, 1958. Abstracted from a thesis submitted by E. P. Przybylowicz t o the llassachusetts Institute of Technology in partial fulfillment of the requirements for the degree of doctor of philosophy in September 1956.
Dete rminuti o n of Cysteine with Ferricyanide by Amperometric Titration with Two Polarized Electrodes H. GEORGE WADDILL and GEORGE GORIN Deportment o f Chemistry, Oklahoma State University, Stillwoter, Okla.
b Semimicro amounts of cysteine can be determined by direct titration with standard potassium ferricyanide (0.1 to 0.001M) in phosphate buffer of p H 7. Two platinum electrodes connected to a 1 00-mv. direct-current source and to a sensitive galvanometer are used to detect the end point, which is indicated by the onset of a current that increases linearly with the volume of excess reagent. Amounts of cysteine mmole between 1 and 1.5 X can b e determined with a precision The procedure is straightwithin 1%. forward and the stoichiometry is not sensitive to external conditions. The apparatus is simpler than that required for other electrometric methods. Precision also compares favorably with other methods. As other common amino acids do not interfere, the method can be applied to amino acid mixtures.
C
is of great importance in biochemistry, and many methods have been developed for its deterniination (2, 3). I n the past the most popular method involved titration with iodine. Ho.c\ever, this gives a somewhat variable stoichiometry unless the conditions are carefully controlled (3, 12). Recently, methods have been proposed which make use of certain metal ions, and in which the end point is determined amperometrically with a rotating platinum electrode and nonpolarizable halfcell (Q-11). The method described here utilizes potassium ferricyanide as reagent and two polarized platinum electrodes for detecting the end point. Kendall and Holst (’7) and Mason ( I S ) have used ferricyanide for the titration of cysteine. I n the latter case the end point was indicated b y the appearance of ferricyanide TSTEIXE
color, but in the former the method of end point detection was not specified. Because the color of ferricyanide is rather pale, the end point cannot easily be discerned in dilute solutions, and large blanks may be needed. Mirsky (15) used an excess of ferricyanide and determined the amount reduced to ferrocyanide colorimetrically by converting the latter to Prussian blue; the precision n a s within 5%. dnson ( 1 ) used the same method in several investigations; he determined, more accurately than had been done theretofore, that the stoichiometry a t pH 6.8 \\as 1 to 1; that, at that p H value, other oyidizable amino acids interfered little or none; and that the reaction n.as subject t o catalysis b y certain metal ions, among which copper(I1) was quite effective. Use of tmo polarized electrodes for detecting the end point of titrations was initiated by Foulk and Ban-den (6) in connection with the reaction of iodine with thiosulfate. They named the end point “dead stop” because, when thiosulfate is used as titrant, the galvanometer readings decrease rapidly and come to a stop m-hen the end point is reached. Hon-ever, the name is inappropriate when an electroactive species is used as the titrating reagent because the end point is then indicated b y the onset of current; this is the case in the method under discussion. The current flows n hen both ferrocyanide and ferricyanide ions are present, and the latter can accumulate only after the end point. The magnitude of the current depends principally on the concentration of these ions, the potential applied, and the size of the elwtrodes. The response of the galvanometer a t the end point, therefore, varies. Indeed, it is an advantage of the method that its sensitivity can be adjusted b y
.Q I E
Figure C. E. G.
1.
Circuit
1.5-volt dry cell Electrodes Galvanometer
R1. R2.
S.
diagram 100,000 ohms Decade box Switch
altering these factors. For a more complete discussion of the principles underlying the method, reference may be made to discussions by Delahay ( 4 ) and Kolthoff (8). APPARATUS
The electrical part of the apparatus consisted of a 1 5-volt dry cell-e.g., Eveready, S o . 6-a 100,000-ohm resistor, a 10,000-ohm decade resistance box, and a galvanometer with sensitivity of 0 10 pa. per division (G-11 Laboratories, Inc., Catalog KO. 570-211), connected as shown in Figure 1. Kot shown is the damping resistance of the ga!vanometer, which was connected across its terminals. Because the cell discharges slowly through the high resistance of the circuit, this can he left connected a t all times, except perhaps when a long period of disuse is contemplated. The electrodes were pieces of 22-gag~ platinum wire about 0 5 em. long, fuscd in the end of soft glass tubes, inside which connection was made in t h t uaud way with the aid of mercury. The electrodes were mounted in a S o . 00 twoVOL. 30, NO. 6 , JUNE 1958
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