eentration exceeded 6.55iv. Consistent and satisfactory recovery was obtained below a concentration of 3.55N. Some results obtained with this modified procedure are shown in Table 11. Usually, low recoveries were obtained with these substances with the regular persulfate oxidation method, due either
to their low solubility in water or the heterocyclic rings they contained. Satisfactory results were obtained with the modified Procedure. This method will be of little value, because of sampling difficulties, for compounds, such as cholesterol, which cannot be solubilized with concentrated sulfuric acid.
LITERATURE CITED
(l) J.r N.~ ANAL.CHEM.26, 1503 (1954). (2) Van Slyke, D. D., Folch, J., J. Bid. s.j
Chem. 136,509 (1940).
R~~~~~~~for review J Accepted March 28, 1957.
~ 25, 1957. ~
Cerimetric Titration of Iron Using a Mixed Indicator WALTER R. HEUMANN and BRANKO BELOVIC Department of Chemistry, University of Montreal, Montreal, Canada
b 1,I 0-Phenanthroline ferrous sulfate is a very accurate indicator for the cerimetric titration of iron, but its abrupt color change lacks a preceding warning signal. A method is proposed whereby a warning is provided by an additional indicator of lower redox potential than that of the equivalence point, which produces a color change somewhat earlier than the main indicator.
tentials of indicators, which would meet the requirements: E = Eo
N
TEE
CERIMETRIC
TITRATION
Of
iron in dilute sulfuric acid 1,10phenanthroline ferrous sulfate or ferroin is an outstanding indicator as far as accuracy, stability, reversibility, and color are concerned. Its only disadvantage is the complete absence of any warning prior to the abrupt color change, which does not take more than 1 drop of titrant for 40 to 50 ml. of 0.1N cerate solution. Diehl and Smith (8) proposed the “titration thief” to overcome this difficulty. It occurred to the authors that a convenient warning would be given by the color change of an indicator of lower standard redox potential than that of the equivalence point of the titration. The color change of this warning indicator, which precedes that of the main indicator, must not mask the color change of the main indicator ferroin. The intensity of its colors must permit the use of only a small amount, so as not to affect the accuracy of the titration. It was assumed that a warning starting a t a point where about 99% of the iron is oxidized would be satisfactory. That means that in a titration requiring, say, 40 ml. of standard cerate solution, 0.4 ml. or 8 to 10 drops of titrant would have to be added after the warning occurred. The redox potentials involved in this titration are shown in Table I. Using the Nernst equation, one can calculate as follows the standard po1226
ANALYTICAL CHEMISTRY
(1)
12
Introducing the standard potential of iron from Table I and supposing that 99% of the iron is oxidized, we obtain the potential where the warning should start: E = 0.68
I
[oxid.] + 0.059 - log ___ [red.]
99 = 0.80 + 0.059 log 1
(2)
As the perceptible color change of any redox indicator occurs within a potential range of 0.08 to 0.10 volt and the standard potential, as a rule, lies in the middle of the color change, the standard potential of the wanted indicator should be about 0.84 to 0.85 volt. Table I1 gives the only three indicators falling in the neighborhood of this potential value, which are thoroughly described in the literature. To investigate the suitability of these indiTable 1. Redox Potentials in 1 N Sulfuric Acid Type of Poten-
System Fe+++/Feif Ce+d/Ce+++ Titration Ce+4/Feii Ferriin/ferroin
Potential Eo E,
tial, volts
Eesui1. Eo
1.06 1.06
References
0.68
1.44
\
,
(4)
Table II.
cators, potentiometric titrations of ferrous iron were carried out with a standard cerate solution in the presence of the indicators, using a Beckman Model H-2 p H meter with a platinum indicator and a calomel reference electrode. The acid concentration throughout the titrations was maintained a t 1.V in sulfuric acid. Forty milliliters of O.lOOOLV ferrous sulfate solution in 1N sulfuric acid was diluted with 200 ml. of 1N sulfuric acid and the indicator to be investigated was added. The titration was carried out with a 0.1OOON solution of ammonium sulfatocerate in 1N sulfuric acid. Each indicator was first tried alone and then in a mixture with ferroin. Figure 1 shows the relation between the zones of color change of the indicators and the percentage of titration completed. The N-methyldiphenylamine-p-sulfonic acid sodium salt, prepared according to Knop and Kubelkova-Knopova (6),gave a rather early warning, starting when 96.670 of the iron was oxidized. This was in accordance with its somewhat low standard potential of 0.80 volt. The main disadvantage of this indicator is that the reddish hue of its oxidized form, which remains in the solution after the ferroin has changed from reddish orange to pale blue, somewhat masks the color change of the ferroin. The 3,4,7,8-tetramethyl-l,10-phenanthroline ferrous sulfate or tetramethyl-
Redox Indicators
Eo
Indicator N-Methyldiphenylamine-p-
sulfonic acid
3,4,7,%Tetramethyl-l,lO-phe-
nanthroline ferrous sulfate Diphenylaminesulfonic acid
- Color Change by Oxidation From Colorless
To
Reddish violet
in 1N
&SO4,
Volt 0.80
Reddish orange Pale blue, almost 0.81 colorless Violet 0.84 Colorless
References (5) (1, 8 )
(6)
~
~
Percentage of reaction Figure 1. Zones of color change of warning indicators shown on cerium(1V)-iron(1l) titration curve
a. N-Methyldiphenylaminesulfonicacid h. 3,4,7,8-Tetramethyl-l,lO-phenanthroline ferrous sulfate c.
Diphenylaminesulfonic acid
ferroin showed a warning zone closer ta the end point than the preceding indicator. It started when 98.170 of the iron was oxidized. This indicator, however, was unsuitable because its color change is practically the same as that of the ferroin, the reddish orange colors of the reduced forms of both indicators being almost identical. The diphenylaminesulfonic acid, the only one of the three indicators commercially available, gave the warning zone closest to the end point. It starts when 99.1% of the iron is oxidized and just reaches its end at 99.94%-i.e., short of the end point and the abrupt color change of the ferroin. Its color change does not interfere with that of
the main indicator. It goes from colorless to dark violet, whereas the ferroin changes from reddish orange to a pale blue which is almost colorless. The color change of the first indicator is distinct in the presence of the orange color of the ferroin, the intermediate mixed color being a dark wine-red. The whole range of the combined color changes goes thus from reddish orange over dark wine-red to violet, and the latter change is abrupt and easily perceptible. The warning actually starts even earlier than the first persistent coloration by the warning indicator, because the temporary color changes prior to the persistent one are also distinct. The warning thus develops
gradually and extends over a t least 1 to 2% of the total volume of titrant solution. Results were identical with the sodium and barium salts of the indicator. The diphenylaminesulfonic acid sodium salt is suggested as an auxiliary or warning indicator in the proportion of 2 to 3 drops of 0.01iM solution (270 mg. of sodium salt per 100 ml. of water) for each 100 ml. of titrated solution, together with 1 drop of 0.01M solution of ferroin. It can be easily calculated that this additional amount of indicator does not significantly affect the accuracy of the titration. \Then the titration is done as described above-Le., to reach a total volume of almost 300 ml. a t the end point-6 to 8 drops or about 0.3 ml. of warning indicator will have been used. As this quantity requires 0.06 ml. of 0.1N cerate solution for its color change, the error due to the warning indicator amounts to 1 part in 667 parts or to 0.15%. The error caused by the ferroin is about one sixth of this value according to the proportion in which the two indicators are used. The combined indicator error amounts thus to approximately 0.18%. For work of high accuracy this error can be taken into account. LITERATURE CITED
Brandt, W. K., Smith, G. F., ANAL. CHEM.21, 1313-19 (1949). Diehl, H., Smith, G. F., “Quantitative Analysis,” p. 249, Wiley, Kew York, 1952. (3) Ibid. , p. 275. (4) Hume, D. Ii., Kolthoff, I. M., J . Am. Chem. SOC.65, 1895-7 (1943). Knop, J., Kubelkova-Knopova, O., 2.anal. Chem. 122, 183-201 (1941). Sarver, L. A., Kolthoff, I. M., J. Am. Chem. Soc. 53, 2903 (1931). Smith, G. F., “Cerate Oxidimetry,” p. 22, G. Frederick Smith Chemical Co., Columbus, Ohio, 1942. Ibid., p. 25.
RECEIVEDfor review August 6, 1956. Accepted January 31, 1957.
Determination of Sulfur in Nickel by the Evolution Method C. L. LUKE Bell Telephone Laboratories, Inc., Murray Hill,
b The evolution method for the determination of sulfur in nickel has been improved by using platinic chloride to accelerate the solution of the metal in hydrochloric acid.
A
method for the determination of sulfur in nickel is needed in certain phases of vacuum tube developRAPID
N. J.
ment. Attempts to develop an oxygen combustion-iodometric titration method using an induction furnace with automatic sulfur determinator were not successful because no way was found to eliminate the high and variable blanks due to the ceramic crucibles. I n considering other possible methods, it seemed probable that an evolution
method might be used, if means were found for speeding up the solution of nickel in hydrochloric acid and it could be demonstrated that all the sulfur in nickel is present as sulfide (1). Dissolution of nickel in hydrochloric acid can be greatly accelerated by the addition of a small amount of platinic chloride. During the dissolution VOL. 29, NO. 8, AUGUST 1957
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