Stannous Chloride-Iodine and Zinc-Ferrocyanide Titrations Application of the Dead-Stop End Point D. F. SWINEHART University of Oregon, Eugene, Ore.
HE dead-stop end-point apparatus of Foulk and Bawden , 1 6 ( ' which they applied t o the iodine-thiosulfate and the iodine-arsenite titrations, has been found t o be applicable to several other titrations (1-4, 8). In this paper the characteristics of the end points for the stannous chloride-iodine and the zincferrocyanide titrations using this apparatus are described.
solutions yet another factor of 2 and still obtain a satisfactorlend point. In general, these titrations mere reproducible to 1 t o 2 parts per thousand. Representative results are shown in Table I. The precision of results was about the same a t all dilutions down to 0.005 M zinc sulfate and 0.0033 M ferrocyanide. Several hundred zinc analyses have been carried out in this laboratory using 0.0067 M ferrocyanide with consistently excellent results. The titration is rapid and no delay is involved except for the time required to heat the solutions to about 70" C. The end point is sluggish below 60 O C.
STANNOUS CH W R I D E I O D I N E TITRATION
A dilute solution of stannous chloride in 0.1 M hydrochloric acid may be titrated with a dilute iodine solution, using the deadstop apparatus to detect the end point, if proper precautions are taken t o eliminate atmospheric oxygen from the solutions. Xot m l y does the presence of oxygen introduce error due to the oxidation of stannous ion, but the end point does not function properly in the presence of oxygen. Ten-milliliter titrations of 0.005 N stannous chloride with 0.005 iodine were successfully carried out by working in an atmosphere of carbon dioxide or purified nitrogen. The stannous chloride solution was made up in a threenecked flask and the inert gas was huhhlc~dthrough it for about 2 hours before use. This tinit, \\-as required to stabilize the solution so that the normality did not change at, an appreciable rate. It was found that carbon dioxide directly from the tank was much better than nitrogen, even after the latter was purified by bubbling through alkaline pyrogallol or Fieser's solution (6). Using nitrogen, the stannous chloride solution never became stable, but the normality changed at, the rate of 2 t o 370 per hour even after long bubbling. After 2 hours' hubbling with carbon dioxide the solution became so stable that, Xvith continuous bubbling, the normality did not change appreciahly in a day or two. A sample for t'itratiori was removed by pipet and the titration was carried out, in a large test tube while the solution was stirred with a stream of carbon dioxide. The burets had fine capillary tips which were immersed in the solution during titration. With excess stannous chloride in the titration vessel and 50 mv. across the electrodes, the galvanometer (sensitivity 1.7 X ampere per mm.) showed its characteristic small, steady deflection. -4s iodine was added, near the end point, the galvanometer showed small deflections, returning each time t o its steady reading with stirring, A permanent additional deflection indicated the end point. The apparatus thus behaves as it does for the thiosulfate-iodine end point (6). Duplicate 10-ml. titrations agreed to a precision of 1 to 2 parts per thousand. ZINC-FERROCYANIDE TITRATION
The apparatus responds in a similar manner to the titration of zinc ion in dilute sulfuric acid and ammonium sulfate a t 60" to 80' C. A potential of 200 mv. was used across the dead-stop (Jlectrodes. Less than 150 mv. resulted in a less satisfactory and less sensitive end point. If the solution contains excess zinc, the galvanometer shows its characteristic small deflection and a sharp additional deflection occurs when a trace of excess ferrocyanide is added. The presence of ammonium sulfate was fouhd to be necessary in order to get a sharp deflection a t the end point. The reason for this is not apparent. Zinc solutions containing 2 grams of ammonium sulfate and 2 ml. of 6 .V sulfuric acid in 100 ml. of solution behaved satisfactorily. Using a galvanometer with a sensitivity of 1.7 X 10-8 ampere per mm., the end point of the titration of 0.01 M zinc sulfate with 0.0067 M ferrocyanide is sharp. I t is possible to dilute both
Table 1.
Representative Results
(Galvanometer sensitivity 1.7 X 10-8 ampere per mm. Potential across dead-stop electrodes 200 mv. Concentrations are approximate only) 0.01 M 0.0067 iM ZnSO4. hll. KIFe(CN)B, M1. Ratio 5.00 5.00 10.00 10.00 18.00 15.00 20.00 20.00 Av. ratio, 1.007. Av. deviation from mean. 1part
5.04 5.03 10.07 10.09 15.10 l5,07 20.16 20.15
1.008 1.006 1.007 1.009 1.007 1.005 1.008 1.008
per thousand.
Attempts to titrate 0.001 'II zinc sulfate with 0.00067 M ferrocyanide using a galvanometer with a sensitivity of 3 X IO-* ampere per mm. were less successful. The end point is no longer sharp a t these dilutions. The galvanometer deflection sets in gradually, presumably because of the finite solubility of the slightly soluble salt, KzZnt [Fe(CK),],. Titrations can be carried out with these solutions with soniewhat less precision (3 to 4 parts per thousand) by titrating t o an arbitrarily chosen deflection. Using a meter of sensitivity 1 t o 3 X 10-6 ampere per mm. instead of the galvanomet'er, 0.1 M zinc sulfate may be titrated with 0.067 J4 ferrocyanide. This has the advantage of great simplicity when a relatively inexpensive and rugged meter is used. It has obvious advant,ages over the use of an external indicator for the detection of the end point I t is easily as good as t,he end point using sodium diphenylamine sulfonate, and it does not involve t,he visual judgment of a color change. Because the stoichiometry of the reaction of zinc with ferrocyanide has been the subject of previous investigations-e.g., that by Kolthoff and Pearson (?')-this subject has not been investigated here. The advice of these authors that the conditions must be maintained constant if reproducible results are desired was followed. .Immonium sulfate vias added to the titration mixture because they advise it, and was later found t o be a necessary component. No study was made of interferences. Presumably the same interferences may be expected with this method as with other methods of detecting the end point. Reagent grade cheniicals and conduct,ivity water were used throughout the work. ACKNOWLEDGMENT
The author desires t,o acknowledge financial support from the Research Council of the Cniversity of Oregon. LITERATURE CITED
(1) Bishop and \Tallare. 1x11. ENG. CHEM., ANAL. ED., 17, 563 (1945).
380
V O L U M E 2 3 , N O . 2, F E B R U A R Y 1 9 5 1 (2) Brann and Clapp, J . Am. Chem. SOC.,51, 39 (1929). (3) Carter and Williamson, Analyst, 70, 369 (1945). (4) Clippinger and Foulk, IND.ESG. CHEM.,ANAL.ED.. 11, 216 (1939). (5) Fieser, * J . Am. Chena. Soc., 46, 2639 (1924). (6) Foulk and Rawden, Ibid., 48,2045 (1926).
38 1 (7) Kolthoff and Pearson. I m . ENG. CHEM.,ANAL. ED., 4 , 147 (1932). (8) Wernimont and Hopkinson, I b d . , 15, 272 (1943). RECEIVED March 6 , 19.50. Presented a t the 115th Xfeeting of the CHEMICAL SOCIETY, Ssn Francisco. Calif.
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Tetrahydroxy Cobalt(l1) Ion as a Qualitative Test for Cobalt SAUL GORDON AND JAMES 31. SCHREYER University of Kentucky, Lexington, K y .
HEX a cobaltous salt solution is added to an excess of satuwrated sodium or potassium hydroxide solution, the cobaltous hydroxide that precipitates locally dissolves upon stirring to produce a deep blue solution which contains tetrahydroxy cobalt(I1) ion ( 5 ) . In the course of studying the spectrophotometric properties of this cobaltous complex in strongly alkaline solutions, the authors realized that the blue solution formed is specific enough to be utilized as a confirmatory test for the qualitative :inalysis for the cobaltous ion. The formation of this complex ion may he represented by the following equations:
+ 2 0 H - +Co(OH)z(Hz0)44 + 2HzO Co(OII)P(HzO)a+ 2 0 H - +Co(OH)4(HzO)n-- + 2H20 Co(H20)6++
1tt.icht.I (4)suggested the use of the slight solubility of cobaltous hydroxide in concentrated alkalies for the qualitative separation of sni:ill amounts of cobalt from nickel, the latter forming an insoluble hydroxide. Donath ( 2 )investigated some of the chemical properties of this complex, postulating the presence of a cobaltite ion, Coot--. Alvarez ( I ) reported that 1 drop of a 1% cobalt salt solution produces a pronounced blue color when added to a boiling solution of potassium or sodium carbonate. Qualitative rrayeiits used as tests for cobalt were reviewed by the International Commission for Reactions and Reagents ( 6 ) , but the application of this colored complex of cobalt was not reported. Although this reaction has been known for a long time, it has not been modified for adaptation to any scheme of cation analysis, nor have its 1iniitat.ions of detection and sensitivity been determined. EXPERLMENTAL
The schtmes most frequently used separate cobalt, nickel, and very often zinc, as a subgroup of the Group I11 cations. Therefore, Hogness and Johnson’s (5) method of analysis for the zinc subgroup was modified t o include this test for the detection of cobaltous ion. To determine the usefulness of this test, the Group 111 cations-aluminun, chromium, cobalt, iron, manganese, nickel, and zinc-Fvere determined in many combinations on a semimicro scale. The test solutions varied in concentration from 0.3 t o 50 mg. of the metal ions. Bfter dissolving the cobalt, nickel, and zinc sulfides in 1 ml. of 6 ,If hydrochloric acid and 1 ml. of 6 AT nitric acid, heat the mixture to boiling, and then transfer it to a casserole. With a stirring rod coalesce the sulfur formed, remove, and discard it, Then evaporate the solution to dryness. Dissolve the residue in a buffer solution of 1 ml. of 2 M sodium bisulfate and 2 ml. of saturated sodium sulfate solution. Divide this solution into two equal portions, A and B. To portion A add 5 drops of 3 M ammonium acetate, and pass hydrogen sulfide into the cold solution for a t least 1 minute Heat in a low flame, gradually raising the temperature almost to boiling. The appearance of a white or very light gray precipitate of zinc sulfide indicates the presence of zinc. Evaporate portion B of the buffer solution in a casserole, almost to dryness. Cool and add 6 drops of water, dissolving as much solid as possible. Label this solution C. Add 4 drops of solution C to 1.5 ml. of saturated sodium hydroxide solution in a 4-ml. test tube. The formation of a blue recipitate at the top of the solution indicates the presence of c o h t . Tap the t u b gently to allow the solutions and precipitate to mix.
Centrifuge. A blue supernatant, especially near the upper burface of the solution (and usually a blue-tinged precipitate) confirm the presence of cobalt. This blue test for cobalt is best seen if com ared with an equal amount of water against a white backgrounz To 1 or 2 drops of the buffer solution C, add 1 drop of 15 M ammonia solution, 4 drops of water, and 2 drops of dimethylglyoxime reagent. The formation of a red precipitate indicates the presence of nickel. DISCUSSION
Any white precipitate formed when the buffer solution is a d d d t o the saturated sodium hydroxide solution is sodium sulfate. light green precipitate yhich might appear a t t,his point is nickel hydroxide. The blue precipitate formed is either a basic cohaltous salt or white sodium sulfate colored by the cobaltous hydroxy complex ion. When the mixture R i centrifuged the blue complex solution is readily discernible. The absence of a blue solution or a blue-tinged precipitate indicates the absence of cobalt, nithin the limits of sensitivity of the test. The reactions that take place when the buffered solution is added t,o the saturated sodium hydroxide solution, assuming that cobalt, nickel, and zinc are present, may be represented by the following equations :
+
+ 2H20 (pink, Co(OH)t(H,O)a + 2015- --+ Co(OH)4(HzO)z-- + 2H10 (deep blue) CO(HZO)B-+ 2 0 H -
+
Ni(H20)6++ 2 0 H - +
+Co(OH)2(HzO)44
Si(OH)n(H20)r 4
+
Zn(Hz0)4++ 3 0 H - --+ H[Zn(OH)a]2Na+
+
2H20 (light green
+ 3Hz0 (colorless)
+ HSO4- + OH- -+- IVa&O, 4 + H2O (white)
This blue complex ion of cobalt may be formed in eit,her sat,urated sodium hydroxide or saturated potassium hydroxide solutions. The saturated sodium hydroxide solution is more satisfactory because of the higher hydroxyl ion concentration, which produces a deeper blue color with the cobaltous ion. If the test i3 applied t o a solution of cobalt without using the buffer solution, the sensitivity is about the same. The advantage of using the evaporated buffer solution is t’hat the white precipitate of sodium sulfate aids in the centrifugation of any other precipitate present and accentuates the blue color of the complex of cobalt. Although this method of cobaltous ion detection is suitable for use on both semimicro and macro scales, it is not satisfactory as a spot test. For most satisfactory results with this cobalt, test, the hydroxyl ion concentration must be greater than 12 M , which is not practical in spot test analyses. The test gave satisfactory results in the hands of a class of 126 in elementary qualitative analysis; 95.3 and 98.4’% of them reported successful tests on their known and unknown Group 111solutions, respectively. SENSITIVITY
The smallest amount of cobalt which may be detected is about 0.05 mg. in 1 drop of solution added to 1.5 ml. of saturated sodium
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