ANALYTICAL CHEMISTRY
1058 Table 11.
Estimation of Thorium in Monazite Sands
Monazite Designation Source Travancore U-1 Travancore u-2 Brazil u-3 Brazil u-4 Travancore u-5 Bra$ U-6 v-1 1 v-2 ? v-3 Brazil Brazil Travancore
Per Cent ThOz (Average) kolumetric Iodate 2 precipitation tations Hexamine 9.33 9.87 9.93 9.15 9.13 9.21 6.31 6.09 LBO 5.97 5.95 8171 8.66 8.88 6.13 5.99 6.07 8.75 8.68 5.77 6.07 8.62 8.90 5.99 6.54 6.35 6.43 8.84 8.90
1 precipi-
cases. The accuracy of the iodate procedure so demonstrated is supported by the fact that data may be obtained with it in less than 8 hours as compared with the 2 days required for the hexamine. The pyrophosphate procedure of Carney and C a m p bell ( 2 ) is even more time-consuming. ACKNOWLEDGMENT
The authors wish to express their appreciation to the Office of Naval Research for support received during this investigation. LITERATURE CITED
Banks, C. V., and Diehl, H., ANAL.CHEM.,19, 222 (1947). Carney, R. J., andcampbell, E. D., J. Am. Chem. SOC.,36, 1134 ( 19 14).
Average results obtained for a number of monazite sands of varying thorium contents are summarized in Table 11. Data for both single and double iodate precipitatims are included for certain samples for purposes of comparison. Included also are data for the same samples as obtained by the hexamethylenetetramine (hexamine) procedure of Ismail and Harwood (4),which yields quantitative results but is excessively tedious because of the multiple precipitations required for clean separations. Data obtained by the volumetric iodate procedure do not check the hexamine data exactly, but the agreement is very good in almost all
Chernikhov, Yu. A., and Uspenskaya, T. A . , Zavodslzaya Lab., 9, 276 (1940).
Ismail, A. M., and Harwood, H. F., Analyst, 62, 185 (1937). Justel, B., Die Chemie, 56, 157 (1943). Meyer, R. J., and Speter, M., Chem.-Ztg., 34, 306 (1910). Moeller, T., and Fritz, N. D., SCI.Rept. N6ori-71, T.O. VII, University of Illinois, Jan. 23, 1948. Moeller, T., Schweitzer, G. K., and Starr, D. D., Chem. Rev.,42, 63 (1948).
Spacu, G., and Spacu, P., Bull. Sect. sci. acad. romaine, 26, 295 (1945)
I
RECEIVED December 26,
1947.
Quantitative Determination of Hydrazine R. A. PENNEMAN AND L. F. AUDRIETH, University of Illinois, Urbana, Ill. Methods recorded in the literature for the quantitative determination of hydrazine were surveyed and those which promised both speed and accuracy were studied experimentally during the analysis of hundreds of hydrazine samples. Excellent and reproducible results were obtained using either the direct iodine or the direct iodate method with solvent or with indicator. The direct acid titration of free hydrazine with 0.5 N hydrochloric acid to either the methyl red or methyl orange end point, followed by
I
N COKNECTION with studies involving the chemistry of
hydrazine, the need for a rapid and convenient method for the quantitative determination of the free base made it necessary to check recorded procedures experimentally, especially because no summary of methods had appeared in the literature since 1924 (1,9). Procedures for quantitative determination of hydrazine are based on its behavior either as a weak base or as a reducing agent. Consequently, available methods involve either direct titration with standard acids (6, 14) or oxidation to nitrogen by such oxidants as iodine (1, 9 ) , bromine (f), permanganate (9), bromate (9, 10,f6),iodate (9, IS),ferricyanide ( 5 ) ,chloramine (fd), or hypochlorous acid (1). Both types of procedures must be carried out under carefully controlled conditions if accuracy and reproducibility are to be achieved, first, because hydrazonium hydroxide is a weak base (Ka = 8.5 X lo-') and secondly, because oxidation often leads to formation of appreciable quantities of ammonia and hydrazoic acid, in addition to nitrogen ( 3 ) . Methods which make use of iodine, iodate, and permanganate may involve direct titration or these reagents may be used to effect oxidation of hydrazine, followed by determination of a product of the reaction or of the excess of standard oxidant. The ferricyanide procedure is likewise an indirect method in which the ferrocyanide equivalent to hydrazine is determined ceriometrically.
iodate oxidation, is a useful combination where ammonia and/or other basic constituents are also present. The use of micropipets is recommended for the analysis of concentrated hydrazine solutions. Prompt acidification of hydrazine samples prior to analysis avoids absorption of carbon dioxide or moisture or loss of hydrazine by air oxidation. Because of oxidation of free hydrazine by atmospheric oxygen, methods involving the determination of hydrazine in alkaline solution are not recommended.
Inasmuch as a direct titration is better in principle, and more convenient where analyses must be made frequently, the indirect iodate (f), iodine (f), and ferricyanide (5) methods were not checked, even though excellent accuracy is claimed for them (1, 9). The bromine and hypochlorous acid methods were not evaluated, because the reagents are of an objectionable nature. Because of the similarity between the iodate and the bromate methods, the latter was not checked. The experience gained through hundreds of analyses has shown that the direct iodate procedure and a combination acid-iodate method developed by the authors are the most suitable, especially where solutions containing free hydrazine are to be analyzed. The latter procedure was found especially useful where ammonia or some other base is present in solutions containing hydrazine. RECOMMENDED TECHNIQUES FOR HANDLING HYDRAZINE SOLUTIONS
As solutions containing hydrazine are subject to oxidation bg atmospheric oxygen (1-4, '7, f4), fume badly, and absorb carbon dioxide, weight burets or especially constructed pipets should be usled for all sample weighings. The hydrazine should be added to water containing a slight excess of acid, followed by dilution in avolumetric flask to approximately 0.025 M . Where iodate oxidation is to follow, hydrochloric acid should be used. Such acid solutions are stable indefinitely.
1059
V O L U M E 20, NO. 11, N O V E M B E R 1 9 4 8 In the acidimetric titration of free hydrazine, recently boiled water (free from oxygen and carbon dioxide) should be used for dilution of samples. With sufficient care, the use of micropipets yielded results which were accurate to within 0.1 to O.27,, especially where very small amounts of concentrated hydrazine solutions were analyzed. Micropipets, 115 to 580 microliters in volume, were constructed from capillary tubing by blowing two tiny reservoirs and drawing the capillary down a t both ends. Using a 1-nil. hypodermic syringe, the pipet was filled with the liquid to be analyzed. As the openings of the capillary were extremely fine, the pipet could be wiped and weighed with no additional precaution against hydrazine loss. The results in Table I show that the reproducibility of filling and transfer of the contents is excellent. In ?fleeting discharge of the pipet, the tip is placed beneath the surface of 50 ml. of boiled water and the contents are expelled by depressing the plunger of the syringe. The tip is moved to another section of the liquid, filled, and emptied twice. Finally, the tip is washed with a few drops of distilled water and quantitative transfer is complete. Before re-use, the pipet is dried by connecting it t,o the laboratory vacuum line and drawing air through it.
Table I.
NzH,+
0.3992 S HC1
Sample Gram 0.5908 0.5909 0.5927 0.5908
M1. 22.96 22.94 23.03 22.95
0.1004 M KIOa for 100/250 Aliquot ’MZ. 32.48 32.51 32.57 32.51
XZH4 Obsd.
“8
Obsd.
70
7%
44.22 44.26 44.21 44.26
2.92 2.87 2.92 2.88
after titration with standard acid can be titrated by the iodate method. The difference in the two titers gives a measure of the amount of other basic components that are present.
+ H20 +NIH~OH+ H +
The hydrolysis constant, K =
li
Ka
=
is in fact the dia-
sociation constant of the hydrazonium ion acting as a weak acid. This calculated value checks the value determmed by Gilbert (81, who found K N ~ H=~ + 1.02 X lo-* By employing the usual simplifying assumptions it is possible to calculate the pH of solutions containing varying concentrations of hydrazoniuni ion. The calculated values are compared with measured pH values of solutions of hydrazonium hydrochloride. (These values nere determined by Walter L-. Fackler, Jr., whose assistance is gratefully acknowledged.) Conon. of N I H b + , Moles/L.
.Methyl red t o color of p H 4.5 buffer
Table 11. Analysis of Synthetic Sample Containing 44.227, Hydrazine and 2.8770 -4mmonia by Combination Acid and Iodate Method
+ H + +.K2Hs+
Hydrazine is a weak monoacid base Kith a dissociation constant a t 25” C. of Kb = 8.5 X lo-’ (11). Hydrazine salts, therefore, undergo appreciable hydrolysis, as indicated by the conwntional Arrhenius type equation:
Indicator Methyl orange to color of p H 4 buffer
99.88 100.00 99.90 AV. 99.93 99.94 99.83 .I\,. 9 9 . 8 8
DIRECT ACID TITRATION (6, Z4)
N2Hc
Titration of 0.2 M Hydrazine with 0.5 A- Hydrochloric Acid
7% NpH4
Calcd. p H 5.46 4.96 4.62 4.46 4.32 4.12 3.96
Measured p H of NxHsCl S o h . 6.01 5.36 4.88 4.72 4.58 4.33 4.17
These considerations make it obvious that no one indicatoi even if buffered, will give an accurate end point in the titratioii of hydrazine by standard acid. This is especially true where the concentrations of hydrazine (and of standard acid) are relatively low or are varied, as may be the case in carrying out routine analytical determinations. The curves of Figure 1 represent pH changes of solutions of varying hydrazine concentrations \then titrated nith acids of varying normality. For solutions 0.1 111 a i t h respect to hydrazine, titration should be effected by using acid where normality is a t least 0.4 AT. An appreciable indicator error may result if a more dilute acid is employed. Cnder these circumstances, the true inflection point falls in the lower (more acid) range of methyl red, and in the upper (more basic) range of methyl orange. Results of analyses using both indicators and titrating with 0.5 S hydrochloric acid appear in Table I. The samples, weighing about 0.6 gram each, mere taken using the micropipet technique described.
6 5
I,
4
1
el Figure 1. i. 0.1
C.
0.1
pH TS. Milliliters of Acid for Hydrazine Titrated with Hydrochloric Acid N acid and 0.05 M N2H4. B. 0.5 N acid and 0.1 . M NIH,. N acid and 0.025 M NzH,. D. Useful pH ranges of methyl red and methyl orange as indicators
To check the accuracy of this method, a synthetic mixture was prepared containing known amounts of ammonia and hydrazine. In a weight buret, 25.4587 grams of 52.677, hydrazine were mixed with 4.8630 grams of 17.86y0 ammonia. The contents of the buret were thoroughly mixed and four samples taken using the micropipet technique. The results are recorded in Table 11. DIRECT IODATE, USING SOLVEKT (9)
SzHI
+ KIOI + 2HC1-
In the presence of ammonia or other base, the acid method will not indicate the true amount of hydrazine. However, combination of the acidimetric with an oxidimetric method on any one sample, in that order, yields a procedure for determining both total base and hydrazine-for example, a hydrazine solution
+ IC1 + N) + 3HzO
--+ i X 2 + 4 H + + 4e+ IO3- + 6” --+ 3H20 + + C1IC1 + e -
.K2H4
5e1/212
COMBINATION ACIDIMETRIC AND IODATE METHOD
KC1
In the presence of concentrated hydrochloric acid (4 A T or above) hydrazine can be titrated directly with standard iodate bolution. The addition of iodate is continued until the iodine color is discharged. Actually, the initial reaction involves reduction of iodate to iodine, which is subsequently oxidized by additonal iodate to iodine chloride, resulting in the disappearance of the iodine color. Partial equations can be written representing these various steps. I/J2
4
A few milliliters of chloroform or carbon tetrachloride are added to dissolve the iodine; the end point is reached when the solvent layer is decolorized. A sample of Merck’s “suitable for microanalysis” hydrazine sulfate, S,H4.H2SO4, was recrystallized and dried a t 110’ C. and then titrated using a 0.4000 N (0.1000 44) solution of Baker’s potassium iodate. Duplicate analyses were made, the second sample containing 0.5 gram of ammonium chloride per gram of
1060
ANALYTICAL CHEMISTRY
Table 111. Effect of Hydrochloric Acid Normality on Reaction between Hydrazine and Iodate Normality of HC1 12
No iodine color in solvent, aqueous solution yellow Appearance of faint iodine color in solvent, aqueous solution yellow Definite iodine color in solvent, aqueous solution yellow Deep iodine color in solvent, aqueous solution brown
7.2 6.8
Table IV.
w
DIRECT IODINE METHOD (9)
Remarks
Nz evolution observed immediately.
8.3
KMO*
until the iodine color begins to lighten. The iodate analyses presented in Table I1 were obtained using this method.
Hydrazine in Samples I2
7%
KIOa %
Difference Based on Iodate
96.9 99.4 99.4 98.2 99.4 99.4 99.3 97.4 .. 99.1 99.4 10.04 ... 16:b 11.63 ... 12.00 13.39 ... 13.55 19.68 ... 19.75 a Order of addition: NaOH.-KRlnOa, NzH4,HzSOd, KI. b Order of addition: SzH4, h a O H , K.\InOd, KI, HzSOd.
% -3.9 -1.25 -2.05 -0.36 0.2b -3. l b -1.2b -0.35b
hydrazine sulfate. Results of these test runs indicated 100.1 and l O O . O ~ , purity of the hydrazine sulfate. These experiments confirmed the literature claims (1) that the presence of ammonium salts does not interfere with this procedure. As the end point depends on the appearance of iodine and its subsequent disappearance by oxidation to iodine monochloride, it is necessary to keep the normality of hydrochloric acid within certain limits, preferably between 3 and 5 in order that free iodine niay be formed. At high concentrations of hydrochloric acid, the reduction of iodate by hydrazine presumably stops a t the IC12-stage, since no free iodine is formed This undoubtedly results from the stability of the IC1,- ion in the presence of high concentrations of chloride according to the reaction: IClA-+ ICl(aq.) C1-; K = 6 X The importance of this requirement was demonstrated by progressively diluting 100 ml. of concentrated hydrochloric acid (12 S ) , containing potassium iodate and an excess of hydrazine, with water and shaking with 5 ml. of carbon tetrachloride (Table 111). The influence of hydrochloric acid concentration on this reaction seems to have escaped attention by previous workers.
+
Reagents Required. Concentrated hydrochloric acid, chloroform or carbon tetrachloride and standard potassium iodate. Procedure. In a glass-stoppered flask containing the solution of the sample place 207, more than an equal volume of concentrated hydrochloric acid (12 S)and 5 ml. of carbon tetrachloride. Add standard iodate solution until the aqueous layer begins to change from a dark brown color to a light yellow. At this point add the iodate dropwise and shake the solution vigorously after the addition of each drop. When the iodine color is compJetely discharged from the solvent laver, the end point has been reached. Final acid normality should be between 3 and 5 . DIRECT IODATE WITH INTERNAL INDICATOR (13)
The procedure is the same as for the direct iodate method using solvent, except that a water-soluble dye is used instead of the solvent. Amaranth and Brilliant Ponceaux 5R having British Colour Index numbers 184 and 185, respectively, are satisfactory for this purpose. (These dyes are known under the Sational .hiline Co. names of Kool Red, 40F, and Brilliant Scarlet, 3R.) -10.27, aqueous solution of either indicator is made up; 3 to 5 drops are sufficient to give a distinct end point in 260 ml. of solution. The dyes are not affected by hydrochloric acid, iodine, or iodine monochloride under conditions of the titration, but are readily destroyed by a trace of excess iodate in 3 to 5 S hydrochloric acid at temperatures above 30" C. The heat of dilution of concentrated hydrochloric acid to 5 to 6 S is sufficient to raise the temperature above this permissible minimum. The addition of the indicator is delayed until the end point is approached-that is,
NzH4
+ 212
4HI
+ hTz
Hydrazine can be titrated directly with standard iodine solution if the pH is maintained between 7.0 and 7.4 and if the last drops of iodine solution are added a t intervals of a few seconds. At a pH lower than 7 the reaction is quantitative bur very slow. If the pH is greater than 7.5 the method gives lon results. When 0.1 S iodine solution is used, 1 drop will give t o 200 ml. of the solution a perceptible yellow color that is permanent for several minutes. Against a good light and a whitr background the end point is easily discernible. The presence of ammonium salts was found t o h a v e no effect. triplicate analyses of a hydrazine sample, the last one containing 0.5 gram of ammonium chloride required 32.95, 33.00, and 32.99 ml. of a standard iodine solution. In experiments made by the authors, no reagent blank wa.c subtracted, since 150 ml. of water containing 0.5 gram of potassium iodide (amount present after a normal titration) and adjusted to pH 7.0 to 7.2 with sodium bicarbonate required but 0.02 ml. of standard iodine solution. Reagents required Standard iodine, standard sodium thiosulfate, and sodium bicarbonate. Procedure. A hydrazine sample is diluted to approximately 150 ml. in a 600-ml. beaker containing pH meter electrodes and a stirrer. Solid sodium bicarbonate is added until the pH has been adjusted to 7.0 to 7.2 and standard iodine solution is added. Additional sodium bicarbonate is added whenever necessary t o maintain the pH in the desired range. When the yellow color begins to linger for a fraction of a second, the rate of iodine addition is decreased to 1 drop every 5 seconds. At the equivalence point the color will persist for several minutes. Hydrazine is readily oxidized by air a t a pH above 7 . In carrying out this procedure it is desirable to eliminate this possible source of error, by introducing purified nitrogen gas into the container over the solution during the titration. INDIRECT PERBlhNGUiATE METHOD ( I )
+ 4MnOn-+
3N2H4
++
4?*In02 3Se 4 0 H -
+ 4H20 (alkaline solution)
Direct titration of a hot 1 S hydrochloric acid solution of hydrazine with permanganate is not recommended, as the end point fades rapidly (9). Kolthoff (9) did, hon-ever, achieve excellent results by oxidizing hydrazine in alkaline solution with excess permanganate and back-titrating the ewess in acid solution, using potassium iodide and thiosulfate. In applying this method to free hydrazine, the authors found that this procedure is subject to erratic variations, amounting frequently to 1% and occasionally to as much as 3 to 5%. The deviations on duplicate titrations were often as large as 0.5 to 1%. In comparison with results obtained by the iodate and iodine methods, the errors using the potassium permanganate method m r e always negative-Le., gave loJyer values for hydrazine content. I n checking this method, Kolthoff used standard solutions of hydrazine sulfate, whereas hydrazine solutions of widely differing concentrations, some containing acid and others only free hydrazine, mere used by the authors. I n an effort to determine the cause of the observed discrepancies, the order of addition of reagents was varied-that is, hydrazine was added to alkaline permanganate rather than permanganate to the alkaline hydrazine solution. As iodine is subject to oxidation to iodate in alkaline solution, the addition of potassium iodide was delayed until after acidification. Hoa ever, the method still gave results that frequently deviated 1% or more. These discrepancies can, however, be accounted for by the fact that alkaline solutions of hydrazine are susceptible to air oxidation, even over relatively short time intervals between preparation of samples and reaction with Permanganate. The pro-
V O L U M E 20, N O . 11, N O V E M B E R 1 9 4 8 cedure is therefore not recommended where free hydrazine is to be determined. Typical results are presented in Table IV. ACKNOWLEDGMENT
The authors wish to acknodedge the assistance of Paul Mohr,
E. A. Brown, and H. A. Gaarder in checking independently the analytical procedures described in this paper. Acknowledgment is also made to the Western Cartridge Company Division of Olin Industries, Inc., for a grant of funds to the University of Illinois to provide the services of part-time analysts in facilitating this investigation. LlTERATURE CITED (1) Bray and Cuy, J . Am. Chem. SOC., 46, 858 (1924). (2) Brown. E. A, thesis, University of Illinois. 1947.
1061 Browne and Shetterly, J . Am. Chem. SOC.,31, 783 (1909). Cuy and Bray, Ibid., 46, 1786 (1924). Dernbach and Mehlig, IND.ENG.CHEM.,ANAL.ED., 14, 58 (1942).
Gilbert, J . Am. Chem. SOC.,46, 2650 (1924). Ibid., 51, 2744 (1929). Ibid.., 58. ~~,1605 ... (1936). Kolthoff, Ibid., 46, 2009 (1924). Kurtenacker and Wagner, Z . anorg. Chem., 120, 261 (1922). Schwarzenbach, H e h . Chim. Acta, 19, 178 (1936). Singh and Rehmann, J . Indian Chem. Soe., 17, 169 (1940). Smith and Wilcox, IXD.ENG.CHEW,AN.~L.ED., 14, 4 9 (1942). Stolle, J . prakt. Chem., (2) 42, 525 (1890) ; (2) 66, 332 (1902). Saebelledy and Madis, Ber. ungar. pharm. Ges., 13, 368 (1937). ~~
>
~
-
-
~
.
RECEIVED February 2, 1048. Abstracted from a portion of a thesis submitted b y R. A. Penneman t o the Graduate College of the University of Illinois in partial fulfillment of the requirements for the Ph.D. degree.
ANALYSIS OF ALIPHATIC PER ACIDS FRASIC P. GREENSPAN LTD DONiLD G. RIAcKELLAR Riiffalo Electro-Chemical Co., Znc., Buffulo 7, S. Y . A new method for the analysis of per acid solutions is based upon the use of ceric sulfate as a titrant for the hydrogen peroxide present, followed by an iodometric determination of the per acid present. Data are presented showing the results of analyses of sample per acids by the proposed method compared to the potassium permanganate-thiosulfate method of D'Ans and Frey. The new method gires higher, more reliable values.
C
OMMERCI-kL introduction of peracetic acid has stimulated renewed interest in the per acids as bleaching agents, polymerization catalysts, and oxidants in organic synthesis (4, 5, 7 , 10) In the course of extensive work on peracetic acid and other aliphatic per acids, the need arose for an accurate and rapid method of determining aliphatic per acids in the presence of hydrogen peroxide. Hydrogen peroxide is usually found associated with the per acids in aqueous solutions as a result of (1) preparation of the per acid from hydrogen peroxide (8) and (2) hydrolysis of the per acid. Where per acids are prepared by the reaction of an acyl anhydride with hydrogen peroxide, particularly in the absence of a mineral acid catalyst, diacyl peroxide will be present Preparation of per acids by interaction of concentrated hydrogen peroxide and the aliphatic acid, as used by the authors, does not give rise to diacyl peroxides. In a simple and convenient method of analysis of per acid mixtures used by D'Ans and Frey ( 2 , 5) hydrogen peroxide present is titrated in the cold with a standard solution of potassium permanganate to the conventional pink end point. Potassium iodide is then rapidly added and the iodine liberated by the per acid is titrated with a standard thiosulfate solution. This procedure Ras based upon a method of Baeyer and Villiger (1) for hydrogen peroxide-persulfuric acids analyses. Extensive use of the D'Ans and Frey method by this laboratory indicated several inherent deficiencies. The last portion of the permanganate titration was comparativelv sluggish, and this interfered with a sharp and reproducible end point. Variation in time taken before addition of the potassium iodide solution resulted in varying thiosulfate titers for per acid content, and reflected on the accuracy of the method. Furthermore, failure to add the potassium iodide solution immediately a t the permanganate end point gave a rapid development of a deep red-purple color, presumably resulting from the oxidation of Mn-* to a higher valenced manganese compound. This striking and intense color change has been developed by this laboratory into a qualitative test for peracetic acid. D'Ans and Frey ( 2 , originally noted the
oxidation of N n + " by per acid a t the end point and hydrolysis of the per acid as possible sources of error but concluded that such errors were insignificant if one titrated rapidly in the cold. A new analytical procedure for the analysis of aliphatic per acid solutions has been developed, based upon the use of standard ceric sulfate for the initial hydrogen peroxide titration followed by an iodometric determination of the active oxygen present as per acid. Ceric sulfate has been found to satisfy the requirements of a stoichiometrical reaction with the hydrogen peroxide in the presence of the per acid and nonreactivity with the per acid. Further, the C e + + formed does not react with the per acid as does M n + + associated with the use of potassium permanganate. Ceric sulfate has been previously recommended for hydrogen peroxide analyses by Furman and Wallace (6). Diacyl peroxides, if present, do not interfere with the respective hydrogen peroxide and per acid titrations. Such peroxides react very slowly with cold aqueous hydrogen iodide. Where interest lies in a determination of the diacyl peroxide content, this may be obtained by heating the solution being analyzed for 10 minutes on a steam bath after the completion of the hydrogen peroxide and per acid titrations (9). EXPERIMENTAL
Reagents. Ceric sulfate (ammonium tetrasulfatocerate) 0.1 A' in 0.05 h' sulfuric acid. Potassium iodide solution, 10%. Sodium thiosulfate, 0.1 A '. Ferroin indicator (0-phenanthroline-ferrous ion). Procedure. The sample of per acid is accurately weighed and placed in a 500-ml. Erlenmeyer flask containing 150 ml. of 570 sulfuric acid and sufficient cracked ice to maintain a temperature of 0 ' to 10' C. An adequate sample is chosen, when possible, to give approximately 40 ml. of thiosulfate titration. Three drops of ferroin indicator are added and the flask contents are titrated with 0.1 A- ceric sulfate to the disappearance of the salmon color of the indicator. Ten milliliters of the 10% potassium iodide solution are then added and the liberated iodine is tirated with 0.1 N sodium thiosulfate. Starch indicator is added near the end point for the thiosulfate titration. Calculations. = ml. of ceric sulfate X N X 17 2 2 10 X sample weight
