lodometric Determination of Cobalt H. .4. LAITINEN kVD L. W. BURDETT', University of Illinois, lirbana, Ill. The object of this investigation was to find a simple and convenient volumetric method for the determination of cobalt in complex compounds. The sample is prepared by ignition to the oxide and fusion with pyrosulfate. The solution is treated with potassium bicarbonate and hydrogen peroxide to form a carbonato complex of cobalt(Il1) which
reacts with iodide in acid solution to yield cobalt(I1) and iodine. The iodine is titrated with thiosulfate, using a visual starch or an amperometric end point. In the absence of interfering metals, cobalt can be determined accurately. The interference of moderate amounts of iron and of small amounts of nickel can be prevented by the addition of fluoride.
I
N CON?;ECTION with the syiithesis of various cobalt(II1) complexes,asimpleand rapid method for the determination of c-obaltin these compounds was desired. Lack of interferences was of secondary importance because of the freedom from other metals. One of the simplest of the existing procedures was first investigated by Job ( 6 ) , who oxidized the cobalt with hydrogen peroxide in the presence of potassium bicarbonate to give a soluble green carbonato complex. This he reduced with ferrom pyrophosphate, the excess of which was titrated with potassium permanganate solution. Metal ( 7 ) also formed this soluble green complex as a step in his procedure. However, as a subsequent step he decomposed the complex by boiling in sodium hydroxide solution to give cobalt (111) hydroxide. After the excess peroxide had been removed hy boiling, the solution was made acid in the presence of potasYium iodide, which was oxidized as the precipitate dissolved. Engle and Gustavson (6) found that 0.5 to 2 hours were required to dissolve the cobaltic hydroxide completely, depending on the amount of cobalt in the sample, and therefore used a carbon dioxide atmosphere to avoid air oxidation of iodide. They oxidized the cobalt directly to the hydroxide, avoiding the formation of the carbonato complex. Difficulty in dissolving the last trares of cobaltic hydroxide was reported by Willard and Hall
investigation of the nature of' the voniplex is under way in this laboratory. This same green complex qerved as a basis for a colorirnetric procedure by Ayres ( 1). AMPEROMETRIC TITRATIOK APPARATUS
A Fisher Elecdropode was used as a source of potential as well as a means of current measurement. STANDARD SOLUTIONS
Standard Thiosulfate Solution. This solution was prepared from C.P. sodium thiosulfate and was standardized against a standard potassium iodate solution made from previously dried analytical reagent grade potassium iodate which was used as a primary standard. Standard Cobalt Solution. This solution was prepared from carbonatotetramminecobaltic nitrate, for which a method of DreDaration is described bv Walton ( 8 ) . The comuound was twice recrystallized from aqueous solution by the addition of an equal volume of ethyl alcohol, washed with alcohol, and then dried a t 90" C. for 1 hour. Its purity was verified by ignition t o CorOc as described by Brintzinger and Hesse (3). Thp purity, calculated from five ignitions, was 99.95 f 0.07%. The standard solution was prepared by ignition of the complex to Co804,followed by pyrosulfate fusion. The melts from several such fusions were dissolved, transferred to a volumetric flask, and diluted to the mark. Other solutions were prepared by dissolving C.P. grade salts without further purification.
(9 1.
Job's method has a number of objectionable features. The large volume employed (150 ml.) requires a large amount of potassium bicarbonate (30-gram excess) in order to give a stable complex. This large excess must be carefully neutralized, to prevent loss due to effervescence. The large volume would probably make complete removal of the excess peroxide difficult, as the concentration of cobalt, which catalyzes the decomposition, would be decreased. An excess of bicarbonate before addition of the hydrogen peroxide is desirable, so that the carbonato complex will form immediately, rather than cobaltic hydroxide which must be dissolved by addition of excess bicarbonate. The use of a carbon dioxide atmosphere complicates the method and should be avoided. An iodometric procedure would be simpler and more direct than the addition of ferrous iron followed by a titration of the excess. Only one stable standard solution would be involved, rsther than two solutions of lesser stability. An increased number of interferences would no doubt result from the uae of iodide as a reducing agent. For the present application, however, this is of no consequence. The exact formula of the green carbonato coniplev in solution has not been determined. Job assigned it the formula KZCoO3, while Durrant ( 3 ) considered it to be (KCO2-0)zCo-0(4),in a more recent investigation, isoC O ( O C O ~ K ) ~Duval . lated from the green solution a green solid which he called cobalt (111)-triscarbonatocobaltiate, Co [Co(C08)a1, on the basis of determination of cobalt and carbon dioxide. In solution, however, it would appear plausible to assume all of the cobalt to be - CO(CO3)2-. An in the same form, possibly C O ( C O ~ ) ~ - -or 1
EXPERIMENTAL
Details of Method Finally Adopted. To 20 to 25 ml. of solution in a 250- to 300-ml. Erlenmeyer k k containing from 1.5 to 250 mg. of cobalt in sulfuric acid solution, add sufficient sodium or potassium bicarbonate t o neutralize the acid, then add 5 grams in excess. Add 5 ml. of 30% hydrogen peroxide and after effervescence has subsided somewhat, wash down the sides of the flask with distilled water from a wash bottle t o ensure complete oxidation of the cobalt. This also serves to wash down any hydrogen peroxide which was splattered upon the walk of the
Table I.
Present address, Standard Oil Co. (Indiana), Whiting, Ind
1268
Analysis of Standard Cobalt Solution by Iodometric Method
Normality of Cobalt Solution
Cobalt Present
Cobalt Found
Mg.
dug.
Relative Error
7% -0.07 0.0 0.08 - 0 , 1.5
0.01
0.001
a
14.72
'
1.473
Amperometric titration of iodine
14.74 14.73 14.75 14.73 1.485 1.491 l.503Q 1.444a 1.49Ia 1.~503~ 1.497a 1 .497a
0.0 0.13 0.07 0.20 0.07 0 1 2 -2 1 2 1 1
8 1 9 0 2 0 6 8
V O L U M E 23, NO. 9, S E P T E M B E R 1 9 5 1
1269
flask by the effervescing solution, in order t o allow it to undcxrgo decomposition. Heat the solution gently until effervescence due to the decomposition of the hydrogen peroxide ceases. (If the cobalt concentration is 0.05 ,Mor higher, standing a t room temperature for 5 minutes will destroy the excess hydrogen peroxide. If the holution is heated too long or too strongly, some decomposition of the complex may occur.) Add several small portions (0.5 gram) of potassium (or sodium) bicarbonate a t intervals during the heating period. (ThiP step has proved effective in keeping the complex stable.) Cool the solution with a cold water (or ice) bath, dilute to 100 ml., and add 5 grams of potassium iodide. Carefully neutralize the solution 1% ith 1 to 1 hydrochloric acid, adding the acid dropwise while swirling the solution until the effervescence (.eases, then add 10 ml. in excess and titrate the liberated iodine with standard thiosulfate solution. Results obtained by the above procedure in the analysis of htandard cobalt wlutionp are given in Table I.
1- is the original volume of the solution, X is the volunic of reagent added, and i is the measured rurrent. This correction is unnecessary if the concentration of the thiosulfate is ten or more times the concentration of the iodine solution being t i t r a t d .
The amperometric method was employed only for ver>-dilute solutions. Some results obtained are given in Table I. The tendency to high results with 0.001 Jf cobalt solutions i!: caused by the incomplete decomposition of the last trace of hyc11,ogen peroxide. The decomposition is complete with 0.005 11f cwbalt' solution. Some other results arv given in Table I1 in coniicvtion with application of the method in the presence of iron. .\ t>.I)ical titration curve is shown in Figure 1. Interferences with Method. Any element which will oxidize iodide to iodine after treatment with hydrogen peroxide and xidification under the conditio1i.Qof the cobalt determination will interfere with this method unless it has previously been rendered inactive by complex formation or some other means. Among these interfering elements are iron, chromium, manganese, copper, antimony, molybdenum, tungsten, and vanadium. Effect of Presence of Iron. It was found that the interfwence o f iron could be largely prevented bj- the addition of 2 grams of rodium or potassium fluoride aiid 2 grams of sodium acetate t o the cobalt solution prior t o the addition of the bicarbonate. 17nder these conditions a light colored crystalline precipitate formed, the exact nature of \vhic.h is still uncertain. As the addition of fluoride to an acidic ferric iron solution gave the same precipitate, it is assumed to be FvF3.41/2H,0,which is a yellow csrystalline compound only slightly soluhle in cold water. This precipitate dissolved if the solution w:is heated longer than had been done in the procedure previously outlined. More satisfactory results were obtained by following the usual procedure and allowing the precipitate to remain, but this made the starch iodine end point somewhat more difficult to see when anslyzing dilute solutions-e.g., 0.001 M . Somewhat high results were obtained in the analysis of both 0.01 and 0.001 M solutions, although in the case of 0.01 M solutions, the error was only of the order of 0.4% or less in the presence of a 3-fold excess of iron (Table 11). The error with the 0.001 M solution, in the presence of a 20-f-!d excess of iron, was of the order of 4 to 5%. Smaller amounts of iron gave correspondingly lower errors.
Table 11. Analysis of Cobalt Solutions in Presence of Varying Amounts of Iron Figure
1.
Titration of Iodine Thiosulfate
w-ith Sodium
Using rotating platinum electrode 111 ml. of 1.13 X 10-4 .M iodine with 9.74 X 10-4 M sodium thiosulfate in dilute hydrochloric acid solution approximately 0.1 M in sodium chloride. Applied e.m.f.. -0.2 volt DS. S.C.E.
Amperometric Titration of Liberated Iodine. With 0,001 M thiosulfate solutions, the end points were not so easily distinguishable as with solutions of higher concentration. It was especially difficult to distinguish the end point of a solution containing various amounts of iron. Therefore, the amperometric titration of the iodine with thiosulfate solution, using a rotating platinum electrode, was investigated and found to give very sharp end points when using an applied potential of -0.2 volt (its. the baturated calomel electrode). Khen the amperometric titration was employed, the method was carried out exsctly as when the iodine-starch end point was used, up to the point of the actual titration. At that point the iodine solution was transferred from the 250-ml. Erlenmeyer flask to a 250-ml. beaker, after which it was titrated. The total volume of the solution a t the end of the titration was measured by means of a graduated cylinder, and this volume, less the volume of thiosulfate added, was taken as the volume of the iodine solution at the beginning of the titration. This value ~ 1 necessary 3 in order to make a correction for the volume where rhange according to the formula: i,,,,. = (' + V ij
Cdbalt Present
Iron Present
Cobalt Found
MQ.
,MU.
.MU.
?
14.72 14.72 14.72 14.72 1.472 1.472 1.472 1,472 1.472 1,472 1.472 1.472 1.472
50
14.76 14.75 14.77 14.77 1.426 1.480 1.497 1.472 1.438 1.538 1.544 1.47Za 1.47ga
0.2i 0.20 0.33 0.33 -3.2 0.54 1.7 0.0 -2.4 4.3
50 30 30 10 10 10
20 20 30 30 50 50
a Precipitate filtered off prios peroxide.
to
Relative Error
4.9
0 0 0 34
heating of solution to r c m o v c cxcess ~
. .
______.
-
Good accuracy and precision were obtained by filtering off the precipitate by suction, using a sintered-glass crucible of medium porosity, prior to the heating of the solution. Under these conditions 1.5 mg. of cobalt were determined in the presence of 50 mg. of iron with an accuracy and precision of 0.35% or better. These results are also given in Table 11. Effect of Presence of Nickel. Sickel, unexpectedly, also interferes. When the solution is heated to remove excess peroxide it becomes more basic anti a precipitate of what was apparently nickel hydroxide forms. The presence of this precipitate causes high results to be obtained possibly because of the fomla-
ANALYTICAL CHEMISTRY
1270 tion of some nickel dioxide from the nickel hydroxide in the presence of hydrogen peroxide. The modified method employed in the presence of iron also gave improved results in the presence of nickel, as the nickel was not readily precipitated, because of complex formation with the fluoride. The values given in Table I11 were obtained using this modified procedure. Results were still high, however; 1.5 mg. of cobalt were determined with an accuracy of 3.7% in the presence of 50 mg. of nickel. Amperometric titration of the iodine with thiosulfate was made in each of the analyses listed.
Table 111. Analysis of Dilute Cobalt Solutions in Presence of Varying Amounts of Nickel Cobalt Present
Me.
Kickel Present MC7.
Cobalt Found
Relative Error
Mu.
%
I t is concluded that a preliminary separation of cobalt is necessary if appreciable quantities of nickel are present. Determination of Cobalt in Cobalt Complexes. The application of the method to the determination of cobalt in complex salts requires first the complete decomposition of the complex salt, in order that the carbonato complex can be formed. The only satisfactory procedure investigated involved the destruction of the complex by careful ignition, followed by fusion of the resulting oxide prior to analysis by the iodometric method. Weigh out a sample large enough to give a suitable titration into a porcelain crucible and decompose it by careful heating with a gas burner. Heat the sample from the top by carefully pointing the burner toward the crucible until all the visible material is decomposed. (This is to avoid loss of sample by being blown out of the crucible as decomposition gases are evolved.) Gently heat the bottom of the crucible to complete the decomposition. Fuse the resulting oxide with potassium pyrosulfate and allow the melt to cool somewhat. Loosen the melt by adding a mall amount of water and gently heating with a gas burner with a low flame. Transfer the contents of the crucible to a 250-ml. Erlenmeyer flask and complete the dissolution of the sample. Complete the analysis as described above. Results obtained with this method in the analysis of a, sample of carbonatotetramminecobaltic nitrate are given in Table IF-. The maximum deviation from theoretical is O.l5%, while the average deviation is only 0.05%.
Table IV. Weight of Sample
Analysis of Carbonatotetramminecobaltic Nitrate Weight of Oxide
Theoretical Weight of Oxide
A preliminary separation of cobalt is necessary if large amounts of iron, or significant amounts of nickel, are present. An amperometric titration of the liberated iodine with thiosulfate solution provides a convenient method of determining the iodine content in dilute solutions. In the acidification step a relatively large excess of potassium iodide is employed to avoid loss of iodine through volatilization. Hydrochloric acid is employed rather than sulfuric, as is specified in many iodometric methods, on the theory that any chlorine produced by oxidation of chloride ion by the trivalent cobalt would react with iodide to give an equivalent amount of iodine rather than being evolved, as would be the case Rith most of the oxygen liberated from a sulfuric acid solution. Table V shows the results of a number of experiments made on dilutions of a single cobalt sulfate solution to determine the requirements for, and extent of, hydrogen peroxide removal. Complete removal of hydrogen peroxide is attained in solutions as little as 0.005 M in cobalt, as shown by the agreement between the resultg obtained with dilute and concentrated solutions. Solutions which were 0.01 M in cobalt gave high results when titrated after standing for 10 minutes a t room temperature, but gave the correct result after heating. Solutions which were more than about 0.05 M in cobalt concentration did not even require heating for peroxide removal, if they were allowed to stand a t room temperature for 10 minutes before titration.
Table V. Effect of Various Conditions on Removal of Hydrogen Peroxide from a Solution Containing the Carbonatocobaltic Complex Concentration of Cobalt Solution, M 0.1
Conditions Employed for Peroxide Removal Sample heated till effervescence ceased
Cobalt Found, Me. 2.472 2.477 2.478
0.1
Sample allowed to stand at room temperature for 5 minutes after complex formation
2.482 2.478 2.480
0.1
Sample allowed to stahd at room temperature for 10 minutes after complex formation
2.477 2.478 2,475
0.1
Sample allowed to stand at room temperature for 30 minutes after complex formation
2.477 2,475 2.478
0.1
Sample allowed to stand a t room temperature for 18 hours after complex formation
2.478 2.479
0.01
Sample allowed to stand at room temperature for 10 minutes after complex formation
0.2586 0.2571
0.01
Sample heated till effervescence ceased
0.2175 0.2473 0.2473
0.005
Sample heated till effervescence ceased
0.1236 0.1236 0.1239
Purity from Iodometric Analysis
G.
G.
G.
%
0.8231 0.6769 0.7543 0.8239 0.7965
0.2684 0.2181
0,2662 0.2181
100.15 loo. 00 99,98 100.08 100,Ol
Table V also shows the stability of the complex, as satisfactory results were obtained after the sample had stood for 18 hours after complex formation. LITERATURE CITED
Decomposition of cobalt complexes by wet oxidation with sulfuric and nitric acids or perchloric and nitric acids was unsuccessful, because it was difficult or impossible to remove the h a 1 traces of nitrogen oxides which interfere with the iodometric procedure. DISCUSSION AND SUMMARY
The procedure as outlined was found to be rapid and simple, especially when applied to samples containing no metallic ions other than cobalt, such as the various complexes of cobalt.
(1) Ayres, G. H . , Rept. S e w England Assoc. Chem. Teachers, 42, 143-7 (1941). (2) Brintainger, H., a n d Hesse, B . , Z . anal. Chem., 122, 241 (1941). (3) Durrant, C. H . , J. Chem. SOC.,87, 1781 (1905). (4) D u v a l , C., Anal. Chim. Acta, 1, 201 (1947). (5) Engle, J., a n d Gustavson, G . , J. Ind. Eng. Chem., 8, 901 (1916). (6) Job, A., A n n . chim. phus., 20, 214 (1900). (7) Metal, G., Z . anal. Chem., 53,537 (1914). (8) Walton, H . F., “Inorganic Preparations,” p . 91, N e w York, Prentice-Hall, 1948. (9) Willard, H . H . , a n d Hall, O., J . A m . Chem. SOC.,44, 2237 (1922). RECEWED March 21, 1951.
