V O L U M E 25, NO. 7, J U L Y 1 9 5 3
1113
for his helpful criticism. hppreciation is also extended to (4) Kolthoff. I. AI., and Robinson, C., Ibid., 45, 169 (1926). R. c., E~~~~~~ i ~ lllinois b ~ ~~ , ~of ~ ~~ ~ h,--hicago. ~ ~ ~ iI ~] I . ,~ ~~(5) Kouba, ~ ,. D. , L., ~ Kicklighter, ~ ~ and Becker, w. W., -&SAL. CHEM., 20, 948 (1948). for his kindness in submitting several compounds required for (6) J I ~ K -4, ~ F,, ~ ,them, fievs,,51, 304-5 (1962). this investigation. (7) Meites, Louis, and RIeites, Thelma, ASAL. CHEM.,20, 984 (1948). LITERATURE CITED
(1) Butts, P. G., Neikle, TV. J., Shovers, John, Kouba, D. L., and
Becker, W. W., AXAL.CHEX.,20,
947 (1948).
(2) Knecht, E., and Hibbert, E., “New Reduction Methods in Volu-
metric Analysis,” New York, Longmans, Green and Co., 1925. (3) Kolthoff, I. hI.. Rec. trap. chim., 43, 7 6 8 (1924).
(8) Siggia, Sydney, “Quantitative Organic Analysis via Functional Groups,” p. 8 2 , New York, John Wiley 8: Sons, 1949. (9) Van Duin, C. F., Chem. Weekblad, 16, 1111 (1919). (10) Zimmerman, R. P., and Lieber, E., - 4 s ~CHEM., ~. 22, 1151 (1950).
R E C E I V Efor O review NorenibPr 6, 1952.
Accepted JIarch 12, 1953.
Effect of Cuprous Iodide on the lodometric Determination of Iron in Presence of Sulfate EDWARD W. K41MMOCK’ AND E R N E S T H. S W I F T Gates and Crellin Laboratories of Chemistry, California Institute of Technology, Pasadena, Calif. determination of ferric iron can be made in Tthe iodometric presence of high concentrations of sulfate if a t least 200 HE
mg. of cuprous iodide are added dissolved in the potassium iodide. In the iodometric determination of copper the necessity for the use of a solu1)le thiocyanate can be eliminated by the addition of an equimolal or greater quantity of ferrous iron. The presenee of moderate concentrations of sulfate cause the iodometric determination of iron to be less satisfactory in that higher concentration> of acid and of iodide are required, and the end points are less permanent ( 5 , 6). I n a study of the iodometric determination of ropper in sulfate-hydrogen sulfate buffers ( 4 ) , the authors observtld that an additional quantity of iodine -toichiometrically equivalent to any ferric iron present was liberated and that stable end points were obtained. Hahn and Windish (3) have reported the catalytic effect of small amounts of cuprous iodide on the ferric iron-iodide reaction, and during the progress of the work reported below) Hahn ( 2 ) proposed the use of cuprous iodide in the iodometric determination of iron, but in neutral solutions and in iodide concentrations sufficiently high, at least a t the beginning of the titration, to dissolve the cuprous iodide precipitate. Recently a procedure was described by Brasted ( 1 ) for the iodometric codetermination of copper and iron in which use is made of nitric arid ~olutionsand of sulfamic acid to eliminate nitrous acid. The present work was undertaken in order to investigate the effect of copper on the iodometric determination of iron in sulfate-hydrogen sulfate solutions similar to those used for the iodometric determination of copper. I n the course of this work additional observations were made on the effect of iron on the iodometric determination of copper. R E A G E h T S AND S O L U T I O N S
A copper solution, approximately 0.1 F (volume formal), was prepared from reagent grade copper sulfate pentahydrate, and was standardized against a standard thiosulfate solution by the procedure of Hammock and Swift (4). -4ferric chloride solution, approximately 0.1 F in ferric chloride and 0.1 F in hydrochloric acid, v,-as prepared from ferric chloride hexahydrate and hydrochloric acid. This solution was tested for ferrous iron and standardized against the same standard thiosulfate by the procedure outlined by Swift (6). The thiosulfate solution was prepared and standardized against Bureau of Standards primary standard potassium dichromate. The normality of the thiosulfate was checked at different times throughout the qeries of experiments, and no noticeable change was detected. -4hydrogen sulfate-sulfate buffer solution 1.80 F in ammonium sulfate and 0.060 F in sulfuric acid was prepared; the p H of this solution waa found by means of a glass electrode pH meter to be approximately 2 . 1
Present addresr. John Rluir College, Pasadena, Cahf.
Cuprous iodide was prepared by m i h g cupric sulfate with potassium iodide in a hydrogen sulfate-sulfate buffered solution, removing the excess iodine with thiosulfate, Jvashing the precipitate, dissolving the cuprous iodide in a saturated solution of potasqium iodide. removing all traces of iodine, and then reprecipitating the cuprous iodide by adding Lvater. The process was repeated if a portion of the cuprous iodide gave an iodine color with starch when dissolved in a solution of potassium iodide. The cuprous iodide thus prepared was washed with water, ethyl alcohol, and ethyl ether, then dried under a vacuum desiccator, and stored in a dark glass-stoppered bottle. At no time during this investigation did the cuprous iodide show an iodine color when dissolved in potassium iodide. (Subsequently the cuprous iodide did give an iodine color after it had inadvertently been heated to approximately 80” C.) PROCEDURE
Appropriate volumes of the copper and ferric solutions were pipetted into a 250-ml. “iodine flask,” 50 ml. of the acid sulfatesulfate buffered solution were added, the volume was made up to 100 ml., and 5 grams of potassium iodide dissolved in 5 to 10 ml. of water were added. The resulting mixture was titrated with the standard thiosulfate solution until the iodine color was indistinct, starch indicator solution vias added, and the titration was continued just to the disappearance of the starch-iodine color; this is designated the iodide end point. hpprouimately 5 grams of potassium thiocyanate were added, and if the starch-iodine color reappeared, the titration was continued to the disappearance of this color; this is designated the thiocyanate end point. Calibrated burets and pipets were used. The titration values shown in the tables are the result of a t least duplicate measurements; where the deviations are not shown, these did not exceed 0.1%. T I T R A T I O N O F C U P R I C 4ND F E R R I C SOLUTIONS
A series of titrations was made to determine the effect of varying the ratio of ferric iron and cupric copper in the iodometric determination of these substances. Hammock and Swift ( 4 ) reported that the presence of 60 mg. of iron in the iodometric determination of approximately 300 mg. of copper did not affect the
Table I. T i t r a t i o n s of Solutions w i t h Various C o n c e n t r a t i o n s of Iron(II1) a n d Copper(I1) Thiosulfate Used, MI.
Thiosulfate Calculated,
Cu(I1) ’ Fe(II1) end point end pointa MI. 4 185 0 000 43.71 3 765 0 413; 43.65 3 347 0 827 43.61 2 510 1 6.54 43.51 1 673 2 482 43 40 43.31 0 8365 3 309 0 418 3 732 43.26 0 000 4 136 43.23 a Average of 2 or more titrations. Maximum deviation was of t h e order of 2 parts per thousand. b A barely perceptible pink color, presumably the ferric thiocyanate complex, was visible after removing t h e starch-iodine color. This color increased in intensity a s t h e quantity of iron was increased.
ANALYTICAL CHEMISTRY
1114 difference between the iodide end point and the thiocyanate end point'. I n the present series of experiments, the quantities of the copper and of the ferric iron were separately varied from none t o approximately 4 millimoles. The data obtained from this series of titrations are presented in Table I and show that a quantitat,ive titration of both cupric copper and ferric iron is obtained provided approximately 1 millimole of copper is present; in addition, the difference between the iodide and thiocyanate end points becomes smaller as the quantity of ferric iron is increased. EFFECT O F CUPROUS IODIDE UPON THE TITRATIOX OF FERRIC IRON
The data of Table I indicate that when the quant,ity of iron becomes roughly equivalent to the quantity of copper, the addition of thiocyanate causes no difference in the position of the end point. -41~0, the data suggests the possibility that the titration of ferric iron in an acid sulfate-sulfate buffered solution could be made possible by the addition of cuprous iodide to the solution to be titrated. The results of a series of titrations made to study this possibility are shown in Table 11. The cuprous iodide was dissolved in and added with the concentrated potassium iodide solution. The data of Table I1 show that lvith quantities of cuprous iodide greater than 200 nig. (approximately 1 millimole), sharp end points were obtained, and they were stable for 2 t o 5 minutes. .Ilso, little or no iodine color appeared upon the addition of 3 to 5 grams of potassium thiocyanate a t or near the iodide end point, and the iodide and thiocyanate end points were identical. Vnder the conditions of these titrations ferric iron can be determined iodometrically in the presence of high concentrations of sulfate provided that a t least 200 mg. of cuprous iodide are present. Experiments showed that approximately 150 mg. of cuprous iodide were required to saturate 100 nil. of a solution containing the quantities of sulfate-hydrogen sulfate buffer and of iodide taken in the above procedure; larger quantities viere required when iron(I1) or copper(I1) were initially present. This indicates that the dissolved cuprous iodide is acting as a honiogeneous catalyst by means of the copper(1)-copper(I1) couple and Table 11. Effects of Various Quantities of Iron and of Cuprous Iodide on Iodometric Titration of Ferric Iron Ferric Iron Taken CUI Used, lI1. Nilli-' d d d e d . Iodide Thiocyanate moles Mg. end point end pointa 43.23 43.23 500 Group Ib 4 . 3 1 6 43,22 43.22 400 43.23 43.23 300 43.23 43.25 200 Premature and recurring 100 end pointd Group2
3 519
300 200 150 100 50
;;g
36 87 36 88 36 93 36 93 36 91 36 91 No stable end pointd
Thiosulfate, Calculated 42 23 42 23 42 23 42 23 42 23
Stability of E n d Point, Min.
g: 5c 5
;fi g
;:
36 88 36 88 36 88
3 to 5 1to2
26 22 26 22 26 23 jC 26 23 5 26 24 26 24 3 to 4 26 25 26 23 26 25 26 23 2 26 23 26 23 S o stable end pointd 26 23 Group 4 1 229 32 10 32 10 32 10 jC 32 14 32 14 32 10 2 to 5 32 14 32 14 32 10 1 E n d point recurs nithin 30 seconds a U on addition of potassium thiocyanate a perceptible pink color appeared): addition of thiosulfate had no effect upon this color which was not SO intense as to mask the starch end point. b Each volume of Group 1 is average value for three titrations made by the senior author with the maximum deviation being approximately 2 parts per thousand. T'olurnes shown for Groups 2 3 and 4 were obtained from titrations by students of the senior author, k i t b the true value being unknown. Volumes given are the average of two or more titrations, with the maximum deviations being of the order of 2 parts per thousand. e Or longer d Addition 'of thiocyanate after the calculated end pointpauseda pink color of sufficient intensity to interfere with detection of starch iodine color. Group 3
2 510
300 200 150 100 50 300 200 I50 100
Table 111. Effect of Ferrous Iron on the Iodometric Titration of Copper (Coppel taken: 4 183 millimoles: calculated voluine of thiosulfate, 43.71 i d . ) Thiosulfate Added, MLa F ~ ~ S H I ) ~ I S O I ) Added. ~~H?O Iodide Thiocyanate llillimoles end point end pointb 43 69 2.55 43 72 43 70 3.83 43 73 43.70 5.10 43.72 43 70 6.38 43.73 43.72 7.66 43.72 43.72 10.20 43.72 43.72 12.75 43.72 Volumes are average of 3 to 5 titrations with maximum deviations not exceeding 0.04 nil. 6 Upon addition of the KSCX a faint pink color developed which became iiiore intense with larger quantities of iron; the color developed slowly and was not so intense a s t o obscure the starch end point.
that adequate cuprou; iodide should be added to saturate the solution, but that more than this quantity would not cau2e further i.nprovement. This is in agreement with the data of Table 11. EFFECT OF FERROUS IRON ON THE DETERMINATION OF COPPER
.Isis seen from Table I, when copper is determined iodometrirally in the presence of ferric iron equivalent to or greater than the copper, both elements are titrated, and the difference between the iodide and the thiocyanate end points essentially disappears. T h a t the reduction of the ferric iron is substantially complete is shown by the lack of significant color after addition of the thiocyanate. The series of experiments shown in Table I11 was made to study the effect upon the difference between the iodide and thiocyanate end points of adding various quantities of ferrous iron to the copper(I1) solution before the titration. The ferrous iron was added as ferroammonium sulfate hexahydrate [Fe(NHa)z(S04)2.6H20]and before the addition of the potassium iodide. The ferrous salt gave a slight color when treated with 50 ml. of of the acid sulfate-sulfate buffered solution and 5 grams of potassium thiocyanate; blank titrations made with 1-, 5 , and 10-gram samples of the ferrous sulfate gave a correction value of 0.05 ml. of thiosulfate per gram of ferrous salt. The data of Table I11 sholv that the addition of sufficient ferrous iron minimizes the difference between the iodide and thiocyanate end point,s, and indicate that, where advantageous for any reason, ferrous iron can be substituted for the thiocyanate in the iodometric determination of copper. Experiments in which the addition of the ferrous sulfate was withheld until a t or near the iodide end point did not yield satisfactory titrations. I t seems possible that with sufficient iron(I1) initially present, less iodine is present during the formation of the cuprous iodide and, therefore, less iodine is adsorbed by the precipitate. Experiments have shown that if equal volumes of a copper(I1) solution are added slowly to two sulfate-hydrogen sulfate buffered iodide solutions and one solution contains iron(I1) the resultant iodine color in that solution is much less intense. ACKNOWLEDG4IENT
The authors wish to express their appreciation for the cooperation of the students of the senior author. Working with unknown solutions, they obtained a considerable number of the values given in Tables I and 111. LITERATURE CITED 1\1)B r a s t e d , R. C., ASAL.CHEM., 24,1040 (1952). (2) H a h n , F. L., Anal. Chim. Acta, 3, 85 (1949). (3) H a h n , F. L., a n d W i n d i s h , H . , Ber., 56B,598 (1923). (4) H a m m o c k , E. W.,and Swift, E. H., ANAL. CHEY., 21, 975 (1949). ( 5 ) Kolthoff, I. hl.,Pharm. Weekblad., 58, 1510 (1921). (8) Swift, E. H., J . Am. Chem. SOC.,53, 2882 (1929). RECEIVEDfor review January 5, 1953. -4ccepted March 9, 1953. Contribution 1767 from t h e Gates and Crellin Laboratories of Chemistry.
