986
ANALYTICAL CHEMISTRY
the pH to 11 with 1N sodium hydroxide solution. Digest the precipitate a t 60’ C. for 15 minutes and allow it to cool to room temperature. Filter the precipitate through a weighed mediumporosity sintered-glass crucible, wash several times with ammoniacal 50% alcohol solution, and dry a t 130” t o 140’ C. to constant weight (at least 2 hours). Calculate the actual weight of cadmium present in the sample analyzed by multiplying the weight of precipitate by the gravimetric factor 0.2109. Results of the precipitation of cadmium with this reagent in both the absence and presence of interfering ions can be found in Tables I and 11. CONCLUSIONS
The reagent is very easily prepared in good yields. The method described is fairly sensitive, the procedure is simple, and the precipitate formed has a comparatively small gravimetric factor. The precipitates are thermally stable and very insoluble in water and most organic solvents. The reagent is highly selective for cadmium and gives results that are accurate to somewhat better than +0.0002 gram in samples containing from 5 to 50 mg. of cadmium and in the presence of most metals except cobalt and nickel. These can be tolerated only in very small quantities, giving resulte which are accurate to 10.0004 gram. ACKNOWLEDGMENT
The authors gratefully acknowledge the support of this work by the iltomic Energy Commission.
Table 11. Precipitation of Cadmium
+
Metal Ion in ion G. Ba 0,100 0.100 0.100 0.100 0.100 Pb 0.100 Cr 0.100
Cadmium Taken, G. 0.0424 0.0424 0.0424 0.0424 0.0424 0.0424 0.0424 0.0424 0.0424 0.0424 0.0424 0,0424 0.0424 0.0420 0.0424 0.0424
2
2
0.100 0.100
A8
Bi Mn 0.100 Fe(II1) 0.100 Ca 0.100 cu 0,050 co 0.050 Ni 0.0~0 Ni Co 0 . 0 ~ 0each
+
Cadmium Found, G. 0.0426 0.0425 0,0422 0.0422 0.0432 0.0433 0.0420 0.0421 0.0422 0.0420 0.0415 0.0420 0.0428 0.0476 0,0447 0.0493
Error,
G. +0.0002 +0.0001 -0.0002 -0.0002 +o. 0008
+o. 0009 - 0 0004 +o. 0002
-0.0002 -0.0004 -0.0009 - 0 0004 +O 0004 +O 0052 +O 0023 4-0.0069
LITERATURE CITED
(1) Charles, R. G., and Freiser, H., unpublished research.
(2) Flagg, J. F., “Organic Reagents,” New York, Interscience Pub-
lishers, 1948. (3) Hillebrand, FV. F., and Lundell, 0. E. F., “Applied Inorganic Analysis,” Ken- York, John Wiley & Sons, 1929. (4) Siegfried, S., and bfoser, bl., Be?., 5 5 , 1089 (1922). ( 5 ) Welcher, F. J., “Organic Analytical Reagents,” Kew York, D. Van Nostrand Co., 1947. RECEIVED for review January 4 , 1952. Accepted March 17, 1952. tribution 847, Department of Chemistry, Cniverslty of Pittsburgh,
Con-
Primary Coulometric Determination of Iron(ll) and Arsenic( Ill) New Method f o r Current Summation WILLI.431 M. MAcNEVIN
AND
BERTSIL B. BAKER
Department of Chemistry, T h e Ohio S t a t e University, Columbus, Ohio This work had tw-o principal objectives: to carry out the primary coulometric oxidation of Fe” and As”‘ in acid solution at a platinum anode, and to investigate the potentialities of a new method of estimating the total number of coulombs of electricity used in a primary coulometric analysis. Iron and arsenic were successfully determined using a hydrogen-oxygen coulometer, with a relative accuracy of about 1%. It w-as shown that a plot of log current DS. time for theoxidation of iron gives a straight line and that from the values of the slope and intercept the number of coulombs used, and thus the amount of iron present, can be calculated with a relative accuracy of about 2%. The area of application of primary coulometric analysis has been extended to iron and arsenic. A new method for current summation obviates the necessity of running the electrolysis to completion and thus allows calculation of results after only 10 to 15 minutes instead of the 60 minutes usually required.
T
HE work reported here had two principal objectives. The first was to develop suitable procedures for the primary coulometric oxidation of iron(I1) and arsenic(II1). The second was to investigate the potentialities of a new method of estimating the total current (coulombs) used in a primary coulometric analysis. The determinations of iron(I1) and arsenic(II1) are of interest for several reasons. Neither element has been determined in a primary coulometric determination. Both reactions involve changes between higher valence states, and because the deposition of a solid phase is not involved, these reactions could not be used in classical electrogravimetric analysis. As the electrolytic oxidation of iron proceeds smoothly and rapidly, it was expected
to provide an ideal case for testing the new method of current integration. The electrolytic oxidation of arsenic(II1) in strong acid solution had been previously suggested (6) as a possible coulometric reaction. I n the early part of this work the oxidation of iron(I1) and arsenic(II1) was studied using a hydrogen-oxygen coulometer. I n the latter part the total current used was estimated from the slope and intercept of the curve of the log current us. time for the electrolysis. The expression used is mathematically derived from the following equation, developed by Lingane ( 2 ) for the current in an electrolysis in which a single reaction is proceeding a t 100% efficiency:
V O L U M E 2 4 , N O . 6, J U N E 1 9 5 2
it =
987
&-kt
where i t is the current a t any time t, io is the initial current, and k is a constant. According to this equation, a plot of log current us. time should give a straight line. Lingane ( 4 ) had shown that the current-time relationships in the controlled potential reduction of lead, copper, and picric acid do obey such an equation. Figure 1 shows a typical curve for the oxidation of iron when current is plotted against time. The plot of log current us. time shows that a straight-line relation is alm followed for the oxidation of iron(I1). In the work reported here, the Lingane equation has been integrated between time limits of zero and infinity to give an expression for the area under the curve which is the total coulombs used. Thus
Potentiostat. The potentiostat was of the design of Lingane ( 3 ) and the contacts on the galvanometer relay were adjusted,to provide control to &0.01 volt. Coulometer. A hydrogen-oxygen coulometer of the Lingane ( 4 ) type was used for measuring the total current used. The electrolyte was 0.5 M sodium sulfate and the following vapor pressures were used for correcting the barometric pressure readings: 24O, 22.2 mm.; 25', 23.3 mm.; 26", 24.7 mm.; 27", 26.2 mm.; 28", 27.8 mm. This hydrogen-oxygen coulometer has been carefully studied by Lingane ( 4 ) and found to be a precise instrument, giving an observed value of 0.1741 cc. of gas per coulomb as compared with a theoretical value of 0.1739 cc. per coulomb. The value 0.1741 was used in this work. mil
it = i , e - k t
m
i"
= -1 0 = -
k
2.303k
-0
-m
(using logarithms to base 10)
Figure 2. Electrolysis Cell
, 0 0
, ,
~
2
Figure 1.
4
, Y f 9 08 14 16 I8 20
6 8 IO I2 TIME I N MINUTES
Electrolytic Oxidation of Iron - 0.8 volt, 0.932 m e .
R u n 69,
TT-hen log current is plotted against time, io is the intercept and -k is the slope. Thus the result of a coulometric analysis may be calculated as soon as the electrolysis has run long enough to establish the slope and intercept of the curve, and a considerable saving in time is effected by obviating the necessity of continuing the electrolysis until completion. APPARATUS
Electrolysis Cell. The cell used (Figure 2) consisted of two 180-ml. electrolytic beakers sealed together near the base by a 40-mm. length of 20-mm. tubing. A filter disk of medium porosity fritted glass a t the middle of the connecting tube separated the anode and cathode compartments. The anode compartment was provided with a plastic coated magnetic stirring bar.
Recorder. Current values were recorded a t about 1-minute intervals by a Leeds and Sorthrup Model S 40,000 potentiometertype recorder which measured the voltage drop across a resistance in series with the electrolysis cell. Electrodes. The platinum gauze electrodes used were of the Slomin type. By geometric measurement the anode area was 75 sq. cm. and the cathode area 40 sq. cm. The actual area was larger because the surface had been sand-blasted. PROCEDURE
Seventy-five milliliters of 1 sulfuric acid was added to each compartment of the electrolysis cell. The potentiostat was set a t -1.4 t o -1.6 volts (us. S.C.E.), a value several tenths of a volt more negative than that to be employed in the subsequent analysis. [The sign of electrode potentials used in this paper follows the system described by Latimer (I).] The solution was electrolyzed for about 10 minutes to oxidize traces of impurities present. The current was then turned off and, in the determin% tion of iron, the solution out-gassed for 10 minutes with nitrogen or carbon dioxide. A 25-ml. aliquot of sample was then added to the anode compartment and a 25-nil. portion of 1 .RI sulfuric acid to the cathode compartment. Whenever the coulometer n a s used to measure the current, another 25-ml. portion of 1 M acid was added to the cathode compartment. This provided a slow seepage of acid toward the anode compartment and tended to prevent loss of anolyte. The additional acid was omitted whenever the slope-intercept method was studied, as its use is baaed on t h e assumption of constant volume. The potentiostat was then set a t the desired potential and the electrolysis begun. R E S U L T S AND DISCUSSION
Data for experiments on the primary coulometric oxidation of iron(I1) using a hydrogen-oxygen coulometer are given in Table
988 ---
ANALYTICAL CHEMISTRY ~
~~
Table I.
~
~ _ _
Oxidation of Iron
+
(Hydrogen-oxygen coulometer' 60 61 58 59 -1.0 -1.0 -0.9 -1.0 50 60 60 60 8.93 8.74 9.03 8 96
+
748 721 25.5 7.75
R u n No.
Anode ootential Length'of run min. €12 Oz collehted, ml. Barometric presaure Observed Corrected Temperature C. H? 02 (cor;ect&), ml. Iron found Me. Mg. Iron present Ale. hlg. Error Absolute, me. Relative, %
733 726 25,5 7.64
752 724 26.0 7.79
752 725 25 6 7.88
61C -1.0 57
5,03
62U -1.0 45 5.02
734 707 25.8 4.28
734 707 25.6 4.27
0.255
0.163 23.9
0.456 23,s
0.469 26.2
0.464 2.5.9
14.2
0.254 14.2
0.466 26.0
0.466 26.0
0.466 26.0
0 466 26.0
0.257 11 4
0.257 11.4
-0,003 -0.6
-0,010 -2.1
+0.003 +0.6
-0.002 -0.8
-0,002 -0.4
-0.003 -1.2
Table 11. Oxidation of Arsenic (Hydrogen-oxygen coulometer) 163 165 Run KO. 162 -1 0 --I 2 Anode potential -1 0 80 80 Length of r u n min. 60 41 4 43 6 43 8 H2 02 colle6ted. ml.
+
Arsenic present
lie. hIE.
Erro; Absolute, me. Relative, %
166 -1
2
60 I7 8
167 -1 2 64
8 82
preseived. Rate of stirring is such a factor, hut this was believed to be satisfactorily controlled. Although a nonsynchronous motor was used, the speed was checked with a stroboscope and the regulatiiig resistance adjusted to give the same revolutions per minute in successive runs. However, the position of the stirring bar in the beaker may have varied slightly from run to run and thus the effective stirring ma\ have differe?. More reproducible results might l i e obtained b>- using a synchronous motor to rotate a cvlindrical anode. However, this type of irreproducibility does not affect practical use of method. Unusual difficulties were encountered in apph ing the slope-intercept method to the determinxtion of arsenic. This prohlem will he reported i n a future publication. SU.MM4RY
The primxi\ coulomrtric determination of iron(I1) and arsenic(111)has been carried out by oyidation in acid solution a t a platinum anode. 4 relative accuracy of about lm0 has been 011tained using the hydrogen-ocygen coulometer.
Table 111. Oxidation of Iron, Slope-Intereept Method 2 26 84.7
84.,
-0 02 -0 8
-0 02 -0.8
2.26
2.26 84.7 -0.01
-0,4
0.904 33.9 - 0 006 -0.i
0.432 16 9 -0.007 -1.5
Run No. Potential Iron present, me. Log iL
69
64
-0.8 n , 466 2 198 158 0 0896 0.476
zc
k Iron calculated, me.
-0 0 2 219 0 0
8 932 340
71
74
-0.8
-0.8
0,466
0.932 2.473 297 0 0834 0.964
OF20 951
1 996 99 0 ,O X 7
0,473
FFTOY
I. A relative accuracy of 1% was obtained in the semimicro range of 14 to 26 mg. of iron. This is comparable to Lingaiie's results ( 4 )for the coulometric reduction of copper, bismuth, and lead a t a mercury cathode. More recently ( 5 ) he reported a higher accuracy of a fern tenths per cent in the coulometric determination of chloride, bromide. a n d iodide, and it is likely that, the relative accuracy of the iron tleterniination could be improvetl by using a larger sample. Results obtained in the det'erminxtioii of arsenic( 111)with the hydrogen-oxygen coulometer are giveii i n Table 11. Again a relative accuracy of around 1%was obtained. Table I11 gives typical results for the application of the slopeintercept method to the oxidation of iron(I1). Log currenttime curves are shown in Figure 3. The relative accuracy of t,hese determinations is about 2%. This could perhaps be somewhat improved by using a continuous rather than an intermittent type of recorder, but probably not beyond 1% relative error. This compares with a limit of accuracy of 0.1 to 0.5% using the coulometer method. Thus t'he slope-intercept method offers the advantage of reducing the time required for an average coulometric analysis from around 60 to about 15 minutes while st,ill retaining an accuracy sufficient for many purposes. Although no complete explanation can be offered, the curvep in Figure 3 do not follow Lingane's derived expression ( 8 ) for the
Absolute, me. Relative, 9
0.01
0.02 2
2
2
0.03 3
-% relating the area A under the log 2.30312 current-time curve, io. the intercept, and - k , the slope, has been derived. It may be applied to calculate the total current uqetl as soon as sufficicnt time has elapsed for evaluation of i, and / . An equation, ,4
=
22 21
DA -3 where D is the diffusion coefficient, A is the area Vd of the electrode, Vis the volume of the solution, and d is the thickness of the diffusion layer. According to this equation, all these slope, k =
curves should have the same slope and the intercept should therefore be directly proportional to the amount of iron present. This was not found to hold true in the present work. Runs 61 and 74 were duplicates of each other. as were runs 69 and 71. The slopes and intercepts for each pair of duplicates are not the same. Despite this variation, the area under each curve agrees with the theoretical. Apparently another variable, not recognized nor controlled, affected both slope and intercept in individual determinations. However, this variable was of such a nature as not to affect the progrese of n run once begun, so that the straight-line relationship hetn-ern log current a n < I time \\-as
01
0
2
4
6
8
10 12
14
16 18
20
TIME IN M I N U T E S
Figure 3. T i m e h g Current Curves for Electrolytic Oxidation of Iron
V O L U M E 2 4 , N O . 6, J U N E 1 9 5 2
989
Results for the determination of iron(I1) obtained by the slope-intercept method agree to within 270 of the true value. While slope and intercept have not been found reproducible for duplicate runs, the area under the curve agrees with the theoretical and may be quickly evaluated from the slope and intercept with an accuracy sufficient for many purposes.
LITERATURE CITED
(1) (2)
(3) (4) (5)
(6) ACKNOWLEDGMENT
This research was suppofied in part from funds by The the university for University Research Foundation Ohio aid in fundamental research.
Latimer, W. M., “Oxidation Potentials,” New York, Prentice Hall, 1938. Lingane, J. J., Anal. Chirn. Acta, 2, 584 (1948). Lingane, J. J., IND.ENG.CHEM.,ANAL.ED., 17, 332 (1945). Lingane, J. J., J. Am. Chem. SOC.,67, 1916 (1945). Lingane, J. J., and Small, L. A,, ANAL.CHEM.,21, 1119 (1949). MacNevin, W. M., and Martin, G. L., J. Am. Chem. SOC.,71, 204 (1949).
RECEIVEDfor review January 11, 1952. Accepted March 10, 1952. Presented before the Division of Analytical Chemistry a t the 121st Meeting of t h e AMERICANCHEMICAL SOCIETY,Buffalo, N. Y. From a thesis submitted to the Graduate School of The Ohio State University in partial fulfillment of the requirements for the degree of doctor of philosophy, J u n e 1951.
Mercurimetric Determination of Chlorides and Water-Soluble ChIorohydrins WILLIARI G. DOMASK AND KENNETH A. KOBE, The University of Texas, Austin, Tex.
A relatively simple method for analyzing chloride samples was needed, which would have a clear solution at the end point rather than a solution containing a precipitate, and would give highly reproducible results. A simple procedure has been developed which makes unnecessary a blank determination or correction for complexes that would otherwise affect results. Mercuric nitrate reagent is used to obtain a sharp end point in a clear solution, thereby permitting titrations which are easier to conduct than by methods that use silver nitrate. Complexes are accounted for by the technique of “calibrating” the reagent. The proposed method is routine and easy to conduct, and it permits rapid determinations with a high degree of precision. It can be employed readily by nontechnical personnel. ARIOUS modifications of the Volhard (10, 1 6 ) and Mohr v(6, 11) methods for chloride determinations have been applied successfully to mixtures containing inorganic and certain organic chlorides. Uhrig (16) suggested a method for chlorchydrin determination based on hydrolysis in the presence of sodium hydroxide or potassium hydroxide. Trafelet (14) presented an improved method for chlorohydrin samples, which, like the Uhrig method, is based on determination of the amount of inorganic chloride formed as a product of hydrolysis. The inorganic chloride content of a sample is obtained first by the Mohr method. Another portion of the sample is heated with sodium bicarbonate to hydrolyze selectively only the chlorohydrins; the solution is then neutralized and titrated as before and the chl,orohydrin chloride is determined by difference. A third portion of the sample is heated with sodium hydroxide t o hydrolyze all water-soluble organic chlorides, and the nonchlorohydrin organic chloride is determined by another difference calculation. Iiumerous aids and improvements have been suggested for argentometric methods for chlorides. These include filtration of the silver chloride precipitate ( 8 ) , use of a layer of organic solvent to remove the silver chloride from the aqueous layer (9, i g ) , use of nitrobenzene in a similar manner ( i ) ,and regulation of concentration of the ferric ion used as the indicator (13). MERCURIMETRIC METHODS
As pointed out by Kolthoff and Sandell (6),a reaction of great practical importance, which has not received the wide application it deserves, is that between halogen ions and mercuric ions to give soluble, slightly dissociated mercuric halides: Hg++
+ 2 C1-
HgC12
(1)
Dubsky and Trtilek ( 3 )and Roberta ( 7 ) described methods for chloride determinations in which mercuric nitrate is the reagent
and diphenylcarbazide and diphenylcarbazone are indicators. The method of Roberts includes the use of bromophenol blue as an indicator for adjusting the pH of the sample. Clarke ( 8 ) reported the mercuric nitrate-diphenylcarbazone method superior from the standpoint of stability of indicator and discernibility of end point. He also presented colorimetric and titrimetric methods for determination of chlorides in water, especially in trace quantities. In the titrimetric method, bromophenol blue is used to adjust the p H of the sample to approximately 3.6, and the sample is then titrated with 0.01 N mercuric nitrate solution. The mercuric ion reacts m-ith the chloride ion to form soluble and only slightly dissoGiated mercuric chloride. Any excess of mercuric ions forms a violet-colored complex with diphenylcarbazone. DEVELOPMENT OF MERCURIMETRIC METHOD FOR CHLOROHYDRINS
In the analysis of chlorohydrin samples it is generally desirable to use reagents that are approximately 0.1 AT. I n this concentration range, it was found that the diphenylcarbazone-bromcphenol blue mixed indicator recommended by Clarke ( 2 ) lacked the sharpness that was desirable for a high degree of precision. Hickman and Linstead (4)had found that xylene cyanole FF, when mixed with methyl orange, served to sharpen the methyl orange color change for base-acid titrations. Experiments showed that xylene cyanole FF, when added to the diphenylcarbazone indicator in optimum amount, gave excellent results in sharpening the end point without shifting the mercuric chloride equivalence point. When the mixed indicator containing diphenylcarbazone, bromophenol blue, and xylene cyanole FF is prepared as indicated below, it may be used both for pH adjustment and for obtaining a sharp mercurimetric end point. I n order to control the influence of pH, it is desirable to begin all titrations a t approximately the same pH. Bromophenol blue
