ANALYTICAL CHEMISTRY
306 chloride in a different cell. This impurity is present in the records of both the radioactive and inactive compounds and so is not radioactive but is probably distilled hy liquid nitrogen from the stopcock grease. The deviation of the two curves which are evident a t the right of each deep hand is the displacement of the band due to the isotope effect. I t is most pronounced in the large band a t 13 microns, where the band due to carbon 14 overlaps the normal band and produces broadening. By using a gas cell or a dilute solution of the radioactive carbon tetrachloride in a solvent transparent to the infrared in the 13-micron region, it would be possible to resolve the two bands and obtain a quantitative value for the isotope ratio.
Young and N. R. 1Ieeks of the Infrared Department of this laboratory for their help in building cells and obtaining spectra.
ACKNOWLEDG.MENT
RECEIVED .lpril 28, 1949. Presented before t h e DiI-ision of Physical and CHEmcAL Inorganic Chemistry at the 115th Meeting of the .I\IERICAN SOCIETY, San Francisco, Calif.
The authors wibh to acknowledge their indebtedness to C. W.
LITERATURE CITED (1) Beamer, W. H., J . Am. Chem. Soc., 70, 3900 (1948).
(2) Dauben, W. G., Reid, J. C., and Yankwich, P. E., AXAL.CHEM., 19, 828 (1947). (3) Kiederl, J. B., and Niederl, V., “Micromethods of Quantitative Organic Analysis,” 2nd ed., p . 108, S e w York, John Wiley 8: Sons. 1942. (4) Regier, R. B., AXAL.CHEM.,21, 1020 (1949).
.~
~~
(5) Roberts, J. D., Bennett, W., Holroyd, E. W., and Fugett, C. H.,
Ibid., 20, 904 (1948).
Determination of Chloride Modification of Volhard Method ERKEST H. SWIFT, G. AIIRON ARCAYD, RALPH LUTWACK, A N D DiLE J. RlEIER California Institute of Technology, Pasadena, Calif. Calculations have indicated and experiments confirmed that stable end points can be obtained in the presence of the silver chloride precipitate when chloride is determined by the Volhard titration, provided the ferric iron concentration is made approximately 0.20 volume formal at the end point. Alternative procedures for the titration have been developed. Confirmatoryanalyses indicate that an accuracy of *0.1 mg. of chloride can be attained when the titration is made with 0.1 formal standard solutions and when the final volume of the titrated solution is approximately 100 ml.
T
H E desirability of having a volumetric method for the determination of chloride in acid solutions has led to the utilization of the original Volhard ( 8 ) titration of silver. An excess of silver ion is added, and the excess is back-titrated with thiocyanate, ferric iron being used as indicator. Drechsel ( 2 ) and later Rosanoff and Hill (6) pointed out that this procedure is inaccurate because of the metathesis of the silver chloride to silver thiocyanate which occurs in the vicinity of the end point. Kolthoff and Stenger (4)discuss this effect and show calculations indicating that under the usual conditions as regards the concentration of ferric indicator, volume, etc., it will be necessary to add an excess of 1.60 ml. of 0.1 F thiocyanate before a permanent end point is obtained; they state that “experimentally it is found that this excess amounts to 2.5 ml. of 0.1 F thiocyanate.” Various expedients have been used to minimize this effect. Rosanoff and Hill ( 6 ) recommended that the silver chloride precipitate be filtered before the titration, and Xolthoff and Stenger (4)confirm the accuracy of this procedure. I n order to obviate the filtration, Rothmund and Burgstaller ( 7 ) suggested thorough coagulation of the precipitate or the addition of a layer of an organic solvent such as ether; Caldwell and Moyer ( 1 ) recommended the addition of nitrobenzene and this procedure has been used extensively. ITo.il-ever, in none of these latter methods do equilibrium conditions prevail and therefore the procedures have to be closely followed in order to obtain reproducible values. Advantage does not seem to have been taken of the fact that an increase in the concentration of the ferric ion used as indicator would cause a decrease in both the thiocyanate and chloride ion concentrations a t the end point and therefore in the error of the titration under equilibrium conditions. The indicator concentration which would theoretically eliminate this errnr ran be calculated as follows:
For the titration to be free of error, the total equivalents of d v e r present when the end point is taken should be equal to t h e sum of the total equivalents of chloride and of thiocyanate present] or 2: equiv. Ag = 2 equiv. C1
+ 2 equiv. SCN
(1)
At the end point of the titration
+ equiv. h g S C S + equiv. h g + (2) + equiv. C1(3) equiv. AgSCS + equiv. SCS- +
2 equiv. h g = equiv. .IgC1
Z equiv. C1 = equiv. .IgC1 3
equiv. S C S =
equiv. FeSCS”
(4)
+ equiv. S C Sequiv. - + FeSCS--
(5)
Therefore Equation 1 can be reduced to Equiv. A I g += equiv. C1-
Previous experiments had shoivn that a perceptible color !vas obtained upon adding 0.10 ml. of 0.01 volume F potassium thiocyanate to 100 ml. of a solution n-hich was 0.013 F in ferric nitrate and 0.6 F in nitric acid. The dissociation constant of the Fe(SCS) complex has been recently determined by Frank and Oswalt (3) to have the value 7 . 2 5 X 10-3, from which one calculates that the F e ( S C S ) + + concentration in the above solution was approximately 6.4 X 10-6 molal. Making use of this value and taking approximate values for the solubility products of silver chloride and silver thiocyanate to be 10-lo and lo-’*, respectively, one calculates that for no error in the titration the molal. From this thiocyanate concentration should he 7.2 X value one calculates that if the ferric iron concentration were made 0.64 molal, a theoretically perfect titration would be possible under equilibrium conditions. If one is n-illing to tolerate an error of -O.l%, and the assumption is made that 2.5 millimoles of chloride are taken and that the final volume is 100 nil., one calculates that the thiocyanate should be 2.3 X lo-? molal and that therefore the ferric iron concentration can be 0.2 molal. ++
As a result of these considerations an experimental study of the effect of increasing the ferric iron concent,ration has been made
307
V O L U M E 2 2 , NO. 2, F E B R U A R Y 1 9 5 0 Table I. Effect of Ferric Concentration on Potassium Thiocyanate Required for Detection of End Point SCN Eul,t.
So.
Fe(lL‘OS)I Formaiity
KSCS Formality
Calculated MoLi1ty
So precipitate present
I
0 . .5 0.3 0.2
4
0.1
(i
0.2 0.1
?.I .,.2 1.8 3.5
x x x x
10-5
10-8 10-8 10-6
AgCl precipitate present 3 . 5 x 104 . 1 x lo-‘
1 . 0 x 10-7 1 . 3 X lo-’ 6.7 X lo-* 2 . 5 x 10-7
minutf,s Experiments made at temperatures as low a i 5’ C. did not qhoTv an appreciable increase in the sensitivity of detection of the end point. I n titrations in which an equivalent concentration of ferric sulfate was substituted for the nitrate the solution had a distinct greenish-yello\v color 1% hich made the end point lrs1 diatinct, and nrgative error\ of about 0.5% were obtainccl.
Table 11. 1 . 3 X IO-’ 3 . 4 x 10-7
and the results of this study together with t x o alternative procedures are presented below. REIGENTS AND APPARATUS
Ileagent grade chemicals were used in all cases. Volume measurements n-ere used exclusively, the apparatus was especially calibrated, and appropriate temperature corrections ivere made. The silver nitrate and sodium chloride solutions were prepared from the appropriate weight of dried crystals, and by dilution to volume. The potassium thiocj-anate solution was standardized against the silver nitrate solution. The ferric nitrate solution was prepared from the approxiniate weight of the reagent product and the concentration was determined iodometrically. These solutions, as well as the nitric acid, were tested for chloride. DETERMIKATION O F OPTIMUM ACID AND FERRIC CONCEXTRATIONS
Preliminary experiments indicated that in solutions approximately 0.5 F in ferric nitrate the nitric acid had to be approxitnately 0.7 F in order to remove the color caused by partial hydrolysis of the iron. A4ccordingly,in order to provide a reasonable factor of safety, all experiments were carried out in solutions approximately 1 F in nitric acid. I n the course of these experiments it was noted that, even in 1 F nitric acid, solutions which were more concentrated than 0.1 F in ferric nitrate had a slight residual color, grayish tinged with purple, which v-as not diniinished by increasing the acid concentration. This color, probably that attributed by llabinorritch and Stockmayer ( 5 )to hydrated ferric ion, is perceptible n-ith 0.15 F , and distinct with 0.3 F , ferric nitratc, solutions, and made it desirable to determine at what fcrric iron concentration the dctcction of the thiocyanate color was mo:t sensitive. T h r rrmlts of tsperimcnt~for this purpose are tabulated in Table I. In the second column are shown the total-that is, the volunii, iornial-concent,r:itioii~ of the ferric nit,rate. The third colunm ? h o w the formal conccntrations of the t,hiocyanate calt~ulntetl from the volume of the solution and the volume of 0.00214 F potassium thiocyanate added to produce a perceptible color. The fourth column s h o w the calculated volume molal concmtrations of the thiocyanate ion; these values are approxim o t e only, for no effort has been made to evaluate the activities of the various specire. In experiments 5 and 6 a silver chloride precipitcite of average size rvas present, together with a small excess of i.hloride, added to prevent motathesis of the precipitate to silver thiocyanate. Titrations of chloride were then made with various concentrations of ferric nitrate, and as a result of the data obtained, which are shon-n in Table 11, an indicator concentration of 0.2 F \vas adopted for future experiments. Experiments showed that somewhat better values were obtained by back-titrating with the silver solution to the disappearance of the color than by using the appearance of the color; scarcely more time is required, as the solution has to be vigorously shaken until the color remains permanent in the latter case. Solutions Jl-hich were titrated to the appearance of the color showed no evidence of fading, even when allowed to stand, with frequent shaking, for periods up to 30
Effect of Ferric Concentration on Accuracy of Titration of Chloride
Expt, F ~ ( S O , J ~ 40. Formality la 0.30 b 2a b 32% b 4 5a b 6a b
0.25 0.20 0.20 0.15 0.10
0.05 0.05
Wt. of Chloride, Mg. Taken Found 104.7 105.1 104.8 126.3 126.4 126.5 104.7 104.6 104.7 104.7 63.26 63.12 98.70 98.40 98.30 98.68 98.01 97.80
Error LIg.
i0.4 10.1 +0.1 +0.2
-0.1 0.0 -0.14 -0.30 -0.40 -0.67 -0.88
70
+0.38 +0.10 +0.08 + O . 16 -0.10 0.00 -0.22 -n.30 -0.40 -0.G8 -0.90
The follon.ing procedure was thrn used to make a series of analyses of chloride solutions.
To 50 nil. of the chloride solution there was added M standard solution of silver nitrate, approximately 0.1 AT, until, after vigorous shaking, a “clear point” was observed, then 2 to 3 ml. in excess were added. Then 17 ml. of 6 F nitric acid were added, followed by 10 ml. of 2.2 F ferric nitrate. The solution was titsated with an approximately 0.1 F st,andard solution of potassium thiocyanate to an easily perceptible pink color which was stable after 1 minute of vigorous swirling. The mixture was then back-titrated with the silver nitrate to the disappearance of the pink color. Comparison was made with a previously titrated solution containing a slight excess of silver nitrate or a prepared mixture of the same composition with a slight excess of silver nitrate. The final volume was between 105 and 110 ml. The room temperature varied from 24’ to 27 a C. Ten analyses were made in which the quantity of chloride taken was varied from 40 to 200 mg.; the volume of approximately 0.1 F silver nitrate used ranged from 12 to 45 ml. and that of the thiocyanate from 0.5 to 5 ml. Stable end points were obtained, t’he average error mas, without, regard to sign, 0.06 mg. of chloride, and the maximum error was 0.11 mg. of chloride. Results of confirmatory analyses of solid samples made by essentially this procedure are shown in Table 111. ALTERh-ATIVE PROCEDURE
The obvious disadvantage of having to make two back-titrations, and of having to shake a full minute in order to attain stable conditions in the back-titration with t,hiocyanate,led to an investigation to ascertain if a small amount of thiocyanate could be added t o the acid chloride solution Containing the ferric nitrate, and the solution then titrated with the .ilver nitrate to the disappearance of the color.
A series of five prelimiriary experiments as made in which to the chloride in about 50 nil. of water there were added 17 ml. of 6 F nitric acid and 10 ml. of 2 F ferric nitrate, followed by either 0.04 ml. of 0.1072 F or 1.00 ml. of 0.01072 F potassium thiocyanate. The solutions were then titrated r i t h silver nitrate to the disappearance of the indicator color. The solutions were shaken vigorously. Comparison solutions were used as in the procedure given above; the final volume was approximately 100 ml. Again stable end points were obtained, the average error without regard to sign \vas 0.08 mg. of chloride, and the maximum error was 0.22 mg. of chloride. The possibility of error arising from decomposition of the thiocyanate during the titration was investigated by preparing solutions for the titration and then letting them stand for various times, There was no detectable diminution in the intensity of
ANALYTICAL CHEMISTRY
308 Table 111. Expt. No.
Confirmatory Analyses
KSCN, MI. MI. 1. Analysis of Solid Unknown, 36.26% C1
W t . of Sample, Mg.
AgNOs,
r: CI
Found
Procedure h 1 2 3
433.6 361.7 361.7
4 5 6
453.6 453.6 361.7 361.7
(0.1260 36.93 29.47 29.48
F)
(0.01072 F ) 1.00 1.00 1.00
36 24 3 6 2; 36 27
(0.1072 F ) 0.93 0.98 1.25 1.26
36.30 36.29 36.27 36.27
Procedure B
7
37.69 37.71 30.47 30.47
2.
Analyses of Solutions
Procedure .4. 25 .OO ml. of S a C l solution taken, 1.00 ml. of 0.01031 P KBCS added AgK03 NaCl Formality M1. used Formalit)Found Calculated 1 2 3 4 5 6 7 8 9 10
25.34 24.84 25.43 25.65 23.91 25.33 23.15 22.15 23.58 23.55
0.1035 0.1035 0.1035 0.1035 0.1107 0.1035 0.1107 0.1107 0.1107 0.1107
0.1045 0.1028 0.1053 0.1052 0.1059 0.1048 0.1025 0,0981 0.1044 0.1043
0.1017 0 1028 0.1051 0.1060 0.1058 0,1047 0,1026 0.0981 0.1044 0,1048
the color during these periods, which extended for as long as 2 hours, and the data from the titrations show no detectable effect of standing. This alternative procedure is much more rapid than the first one, because it is possible to add the thiocyanate by means bf a 1-ml. pipet. Less shaking is required to obtain a permanent end point, probably because adsorption of thiocyanate on the silver thiocyanate precipitate is minimized by the low thiocyanate ion concentration prevailing in the presence of the excess ferric ion. Should the end point be overrun, an additional portion of thiocyanate can be added and titrated; where rapid titrations are desired this procedure can be used to advantage. In view of the above results the two following procedures are recommended: Procedure A. That quantity of the substance to be determined which will require between 20 and 40 ml. of 0.1 F silver nitrate should be contained in about 30 ml. of a solution which also contains approximately 100 millimoles of nitric acid. Add
10 ml. of 2 F ferric nitrate (chloride-free) and 1.00 ml. of standard 0 01 F potassium thiocyanate by means of a 1-ml. pipet. Titrate, while vigorously swirling the mixture, to the disappearance of the color of the ferric thiocyanate complex. Compare the solution with one which has been similarly titrated, but to which an excess of 1 to 2 drops of the silver nitrate has been added. The final volume should be approximately 100 ml. Procedure B. That quantity of the substance to be determined which will require between 20 and 40 ml. of 0.1 F silver nitrate should be contained in about 30 ml. of solution. Add 1 i ml. of 6 F nitric acid and titrate with 0.1 F silver nitrate to a clear point, then add 1 to 2 ml. in excess. (If this titration can be made in a neutral solution a better clear point will be obtained; the acid should then be added.) Bdd 10 nil. of 2 F ferric nitrate and titrate with 0.1 F potassium thiocyanate until a distinct pink color is obtained which is permanent after 1 minute of rapid swirling. Back-titrate with the 0.1 F silver nitrate until the color just disappears. Compare the solution with one which has been similarly titrated to an excess of 1 to 2 drops of silver nitrate. The final volume should be approximately 100 ml. COYFIRMATORY ANALYSES
Analyaes were made by the above procedures of a commercial analyzed sample of a “soluble chloride” and of solutions for student analysis prepared by weighing out appropriate weights of dry reagent grade sodium chloride and diluting to volume in a 2-liter calibrated volumetric flask. The solid sample was analyzed gravimetrically by precipitation of silver chloride; the average of six determinations gave a mean of 36.26% chloride. The data obtained are shovn in Table 111. ACKNOW LEDGJIEST
The authors are indebted to Paul Farrington for assistance in the experimental work and to Y. C. Tang for the preparation and analysis of the chloride solutions of Table 111. LITER4TURE CITED
Caldwell and Moyer, IND. EX. CHEM.,.$SIL. ED., 7, 38 (1935). Drechsel, J . prakt. Chem., 15, 191 (1877). Frank and Oswalt, J . Am. Chem. SOC.,69, 1321 (1947). Kolthoff and Stenger, “Volumetric Analysis,” Vol. 11,2nd ed., pp. 259-61. New York. Interscience Publishers. 1947. Rabinowitch and Stockmsyer, J . Am. Chem. Soc., 64, 335 (1942). Rosanoff and Hill, Ibid., 29, 269 (1907). Rothmund and Burgstaller, 2. anorg. Chem., 63, 330 (1909). Volhard, J. prakt. Chem., 117, 217 (1874). RECEIVEDMarch 10, 1949. Contribution 1302 from the Gates and Crellin Laboratories, California Institute of Technology.
Radiometric Determination of Gold and Rhenium EDWARD D. GOLDBERG AND HARRISON BROWK Institute f o r Nuclear Studies, University of Chicago, and Argonne National Laboratory, Chicago, I l l . Methods have been developed for the determination of gold and rhenium using slow neutron activation analysis. On samples of iron meteorites weighing approximately 0.5 gram, concentrations of gold and rhenium in the range of a few parts per million are determined with a precision of 15% or better.
C
ONTINUIXG work on the development of analytical procedures utilizing the neutron pile, the pile irradiation method, as previously described (Z), has been extended to gold and rhenium. The procedures have been applied specifically to iron meteorites. Weighed amounts of meteorite and the element whose concentration is sought are simultaneously irradiated in the thimble of the Argonne heavy water pile. A known amount of inert carrier of the element desired is added to a solution of the radioactive meteorite and to an aliquot of the irradiated standard. The activated atoms of the element along with the carrier are chemically separated from the meteorite. The ratio of the
amount of carrier recovered to the amount added gives the chemical yield of the separation. A comparison of the decay curves of the meteorite and standard samples gives the concentration of the element in the meteorite. Absorption curves along with decay curves of both samples are compared and used as the criterion of completeness of separation. SAMPLE PREPARATION
The specimens of meteorites subjected t o analyses were of two forms, metal and metal powder. The metal powders re-
