T H E ELECTROMETRIC DETERRIISXTIOK OF SMALL AhlOUKTS O F F E R R I C IROK* BY J. F. K I S G A S D R. K . W A S H B U R S E
During some work which is in progress in this laboratory, it became necessary to determine quantitatively small amounts of iron in solution in the ferric condition. Gravimetric methods were not used because they are too inaccurate for small amounts of iron. The ordinary volumetric methods could not be used because of the great errors caused by the poor end points when dilute solutions are used. The usual colorimetric methods were tried but were found to be unsatisfactory because of the intrinsic inaccuracy of the colorimetric method, the difficulty of comparing slight differences in colors and of correctly reproducing these colors. Hostetter and Roberts’ have developed a very exact method for the direct electrometric determination of small quantities of ferrous iron in which they are able to realize an accuracy of one part in a thousand with a total of 0 . 0 0 0 j 6 grams of iron in a volume of solution of from I O to 35 cc. This method could have been adapted to our use if a satisfactory way could have been worked out for the reduction of a small quantity of ferric iron in a small volume of solution. Hostetter and Roberts have also suggested the use of stannous chloride for electrometrically determining small quantities of ferric iron but we found very dilute solutions of stannous chloride difficult to store and with them reproducible end points hard to obtain, Adam2 finds that the direct electrometric titration of ferric iron with stannous chloride gives unsatisfactory results. Hendrixson and T’erbeck3have suggested the use of titanium salts as reducing agents in electrometric titrations and have shown how such solutions may be standardized against potassium permanganate and potassium dichromate directly or through ferric alum solutions. The titanous solutions which they used were approximately 0 . I normal concentration. In some work with titanous salts on the direct electrometric titration of numerous cations and anions singly and in mixtures, Kolthoff and Tomicek4 have determined ferric iron, working with approximately 0 . I normal solutions. The value of an accurate method for the direct determination of small quantities of ferric iron is obvious. \Then iron samples are put into solution for analysis, the iron is usually partly or completely in the ferric state. Ferrous iron is more easily oxidized to ferric iron than is ferric iron reduced to ferrous iron and the excess of the oxidant is more easily removed than the excess of the reductant. Any process for determining small quantities of iron in which
* Contribution from the Thompson Chemical Laboratory, Killiams College. Hostetter and Roberts: J. Am. Chem. Soc., 41, 1337 (1919). H. R. A4dam:J. S.African Chem. Inst.. 8, 7 (192j). TI7. S.Hendrixson and L. hl. T’erbeck: J. Am. Chem. Soc., 44, 2382 (1922). I. 11.Iiolthoff and 0. Tomicek: Rec. Trav. chim., 43, 798 (1924).
ELECTROMETRIC DETERMISATION OF FERRIC IRON
I 689
the entire process is carried out continuously in one container is desirable because of the elimination of loss by transfer and because of the possibility of keeping down the volume of the solution, TYe have worked out the details of a method by which we are able to determine as little as 0.0003 to 0.0005 grams of ferric iron by the direct electrometric titration with titanous sulfate with an error of less than 0.1 percent.
FIG.I An electrometric titration apparatus for titrating small volumes of heated solutions which are protected from the action of oxygen of the air.
We have developed a simple electrometric titration apparatus which can be easily assembled out of standard laboratory equipment. We have been able to store titanous sulfate solutions for this work as dilute as 0.001normal to 0.0003 normal.
1690
J. F. K I N G A S D R. N . VAEHBCRNE
The Apparatus The apparatus used for the electrometric determinations is shown in Fig. I where : h is a three-liter round-bottom Pyrex flask for storage of the titanous sulfate solutions. a is the inlet from the carbon dioxide or hydrogen generator. B is the mercury seal b a rubber stopper c a cork stopper d paraffin e mercury f metal collar of seal g clamp t o support mercury seal V h e n sealing was. paraffin or collodion were used a t B and D, it was found that the slight jarring of the burette in using caused leaks which admitted enough air t o oxidize the titanous solution very rapidly. C is a well ground stop-cock for filling the burette. h is a capillary tip touching the side of the burette. D is a mercury seal similar t o B. E is a storage reservoir for the saturated potassium chloride solution for the calomel cell. F is the calomel cell. G is a IOO cc round-bottom long-neck Pyres flask which was used as the titrating flask. i is a rubber nipple with holes punched through and serves as a cap to keep out the air. j Carbon dioxide gas inlet k Carbon dioxide gas outlet 1 platinum electrode shown in Fig. z m capillary extension tip to burette n turned up tip from calomel cell. H is a shaking device t o keep the solution constantly stirred. o is a bearing p rubber tension fastened here q asbestos rest for flask Constant readings could only be obtained on the potentiometer when a vigorous and uniform stirring was maintained by the shaking device. I parts of a n electric gong with a n extension soldered on the clapper a t r. A switch in series with “I” makes it possible t o start and stop the shaker for the addition of portions of the titanous sulfate solution during the titration, X rheostat in series with the gong controls the speed J of shaking. K’ires connected to the potentiometer which in our work was a Leeds and Northrup Type I< potentiometer.
1691
ELECTROMETRIC DETERRIINSTION O F FERRIC IRO?;
The electrode shown in Figure 2 was designed t o support the very fine platinum mire (0.008; mm in diameter) and to keep the mercury used for making contact with the platinum wire above the hot solution in which the electrode was used. With a n ordinary platinum wire electrode trouble was frequently experienced due t o slight cracks developing where the wire was sealed in with sealing glass which alloved the hot acid solution t o come in contact with the mercury. This electrode was found t o be very satisfactory, t o give constant and reproducible readings. It was cleaned and stored in hot chromic acid and washed with distilled water. Since the solutions were heated when titrated. a small electric fan was used near the burette t o keep the air in circulation around the burette during the titration so that there would be no volume change in the solution in the burette. Readings over a period of time were made to establish the fact that the fan served this purpose.
I/
Preparation and Storage of Dilute Titanous Sulfate Solutions Knecht and Hibbert' and Thornton and Chapman* recommend the storage of titanous solutions in an atmosphere of hydrogen gas. Thornton and Chapman report their titanous sulfate solutions remaining constant over a reriod of 42 days. With the dilute solutions which we used it was necessary t o give special attention t o the question of storage. TTTehave used both hydrogen and carbon dioxide and with the precautions given below h a r e found that our solutions kept well enough for use with occasional restandardization. The standard FIG.2 titanous sulfate solutions were prepared as follom : An electrode Sulfuric acid solution was boiled for about two hours in a offineplatinthree-liter round bottom Pyres flask t o free it from dissolved osygen. The solution was made of such a strength that solutions. after boiling it was 2 0 percent (by weight) in sulfuric acid. Xfter boiling, the flask was stoppered and hydrogen was passed through the solution as it cooled. The hydrogen, prepared in a IGpp generator with C. P. zinc and sulfuric acid containing trace of copper sulfate, was purified by passing through four wash bottles, ( I ) 2 0 5 KOH, ( 2 ) alkaline permanganate, ( 3 ) two bottles of alkaline pyrogallol. To the sulfuric acid solution was then added the correct volume of LaMotte Chemical Company's 1 5 7 stock ~ titanous sulfate solution t o make an approximate 0.001normal solution. K i t h hydrogen bubbling through the solution, the flask was then fitted with its mercury seal and attached t o the titrating apparatus which itself had been freed from air. The solution thus prepared was compared with a standard ferric iron solution. The standard ferric iron solution was made from a
ig
E. Knecht and E. Hihlier SeiT- Reduction 1Iethotls in Volumetric Analysis." IT-. 11.Thornton, Jr. and E. Chapman. J. Am. Chem. Soc., 43, 91 (1921).
1692
J. F. KING AND R. N. WAEHBURNE
weighed amount of clear crystals of C. P. ferrous ammonium sulfate from a freshly opened bottle. Enough of the salt was taken t o make 500 cc of a solution ten times as strong as the desired standard. The weighed salt was dissolved in a small amount of air free 2 0 percent sulfuric acid solution and oxidized with 0 . I normal potassium permanganate. This oxidized solution was made up to j o o cc with the special sulfuric acid. so cc of this solution was then diluted t o 500 cc for use in standardization. The dilutions were carried out in calibrated apparatus in a thermostat. The following data gives a record of the change in strength of a titanous sulfate solution made as above: On Feb. 3rd a titanous sulfate solution was prepared and when titrated electrometrically against a standard ferric iron solution, it 11-asfound that 13.04 cc of the titanous sulfate solution was equivalent to z j cc of 0.001normal ferric iron solution. That the solution thus stored was effectually protected from oxidation can be seen from the following tests: 2j
cc of
0.001 normal
ferric solution equivalent Date
Feb. 9th
cc titanous sulfate
solution 1 4 .I4
16
14.I 5
25
14.34 14.46 14 6 j 14.75
Mar. 7 23 26
During the entire time that these tests were made this solution was being drawn upon for use in analyses. Effect of Exclusion of Air during Titration Although Thornton and Chapman found that they vere able t o obtain consistent results titrating with titanous solutions exposed to the air when the standardization was made under the same conditions as the analysis, with our more dilute solutions and using the electrometric end point, a series of experiments showed that it would be necessary to protect the solutions during titration from the oxygen of the air. Curves A, B and C in Figure 3 show the results of the titration of a 0.01 normal ferric iron solution with a 0.02 normal titanous sulfate solution in an open flask. Curve D in the same figure is a typical curve resulting from the titration of the same solutions in a flask with the solution protected from the air with a blanket of carbon dioxide. (Analyses of several samples of the carbon dioxide gas from a tank of the gas used, taken a t different times as the tank was being emptied, showed an oxygen content of less than 0.0 j percent by volume.) Three such titrations as shown by curve D gave results as follows: cc of ferric solution 25 25
25
cc titanous solution 11 61 1 1 .j 9
11.60
ELECTROhIETRIC DETERMIXA4TION O F FERRIC IRON
I693
Since the results with the 0 . 0 2 normal titanous sulfate solutions showed that air must be excluded from the flask during titration, we used the carbon dioxide blanket in all the work with the more dilute solutions.
Effect of Acidity I n their work on the electrometric titration of ferrous iron by potassium dichromate, Hostetter and Roberts found that over a range of 1 7 % t o 67%
43
3Q 35
31
27
23 19
R I/ 6
7
8
9
10
//
12
/3
FIG.3 Curves to show the effect of excluding air from the flask during titration.
sulfuric acid, the potential change at the end point increased with increasing acidity. With 0.001normal dichromate solutions, they found that a very sharp end point could be realized if both the iron solution and the dichromate solution were made up with 2 j percent sulfuric acid. I n a series of titrations of ferric iron with titanous sulfate solutions, we studied the effect of increasing the acidity of both solutions from 5% to 2 0 7 ~in sulfuric acid. A 0.001normal ferric solution made up with j% sulfuric acid was titrated with a titanous sulfate solution also made up with jyc sulfuric acid. The titration was run in duplicate with results as follows: cc ferric solution
cc titanous solution
25
13.54
25
13.54
1694
J. F. KIh'G A S D R. N. WASHBURTU'E
The shape of the curves from which the end points were read off is shown by curve A on Figure 4. S e s t a 0.001normal ferric solution was made up with 2 0 7 sulfuric acid and this was titrated with a titanous sulfate solution made up with 2 0 5 sulfuric acid. The titration n a s macle in triplicate with the following results: c r ferric solution 25
( c titanous solution 14 3 2
25
14 3 2
5'
14
33
The shape of the curves from which the end points were read off is shown by curve B in Figure 4.
FIG.1 Curves to shon- the effect of increasing the acidity of the solutions.
W t h the 2 0 5 sulfuric acid solution, the addition of 0 . 0 ; cc of the titanous sulfate solution caused a drop in voltage of 0 . 1 2 0 volts while with the jcc sulfuric acid solution the addition of 0 . 0 ; cc of the titanous sulfate solution caused a drop of 0.080 volts. TI-e therefore made use of 2 0 5 sulfuric acid in all the rest of the titrations. Effect of the Increase in Temperature The time required to reach equilibrium after each addition of titanous sulfate solution was noticably greater n ith the dilute solutions. For instance, when working with 0.0003 norninl titanous sulfate solution, in one experiment and exems of the titanous sulfate solution n-as added t o the iron solution over the usual period of time. Then the voltage change n-as taben over a period of ten minutes as follons:
ELECTROMETRIC DETERMIXATION O F FERRIC IRON
Time in minutes 0 I
2
3
5
volts
Time in minutes
volts
0.3j6 0.341 0.322 0.302 0.213
6 8 9
0.187
1695
0.170
0.16; const ant
IO
Knecht and Hibbert and Kolthoff and Tomicek were successful in titrating with solutions of titanous salts more concentrated than ours a t temperatures
23
24
25
26
27
V O L UM€ O f T/TANOUS 5UL PHATE 5 O L UT/ON /N CUR/C C€NT/T/M€T€RS
FIG.5 Curves to show the effect of increuee in temperature.
above room temperature. K e ohtaiced sharp and reproducible curves with a 0.0003 normal titanous sulfate solution nhen the 2 0 percent sulfuric solution of ferric iron was heated to the boiling point (about 106 degrees Centigrade). I n Figure j are given two curves t o show the effect of temperature. Curve €3 was made from data obtained when the titration n a s carried out at rcom temperature n ith solutions approximately 0.001 normal and curveXn-asniade from data obtained by titrating the same solutions a t the boiling point. In the hot solutions not only are the end 1-oints sharrer hut the drop in potential of 0.140volt is caused by the addition of 0 . 2 cc of the reagent as against in the solution a t rooin temperature of 0.110by the addition of 0 . ; cc of the reagent.
I 696
J. F. K I N G AND R. N. WASHBURKE
Accuracy of Method I n order to test out the accuracy of the method, three series of experiments were carried out with different solutions. Each titration was run in duplicate and the results given in the following table are the average of the two. Column I cc of iron solution used 2 grams of iron in this volume of solution corrected for the iron impurity in the reagents (determined by running a “blank”). 3 cc of titanous solution used (read off the voltage-volume graphs). 4 grams of iron equivalent to this volume of titanous solution (Le. “grams of iron found”) The first duplicate analysis in a series is in this case used as the one to “standardize” the titanous sulfate solution. 5 ratio of the volume of ferric solution to the volume of titanous solution. 6 percent error considering the first analysis as the standard.
TABLE O F RESULTS cc. of iron
solution
Licorrected” grams of iron taken
cc. titanous
solution
“grams of iron found”
ratio of volumes
percent error
Series I 2j.00
0.001407
49.98 3. 3 4 . 9 8 4. 1 4 . 9 8
0,002812 0.001966
5.
0.0021j3
2j.21
6. 1 9 . 9 8 7. 2 9 . 9 8
O.OOI720
20.13 30.20
8. 2 5 . 0 0
o.0004203 0 0003360
I. 2.
0.0008429
16.30 32’64
0.001407
22.82
9.72
0.I
0 . 0 0 I 968
533 1.531 I . j32
0.0008399
1.539
0.37
0,002Ij
3 19 0 . 0 0 2 579
0.9916 0.9925 0.9927
0.07
0.0004203 0.0003361 0.000;0;9
1.022
0,0028I 5
1’
0.I
Series I1 2 j
.oo
0 0 0 2 581 ~
0 ,O O I 7
0.08
Series I11 9. 1 9 . 9 8 IO. 2 9 . 9 8
~
0.0005042
24.46 19.56 29.44
1.022
I ,019
0.03 0.3
I n Figure 6 are reproduced the curves which resulted from the six analyses giving the data in Series I1 of the table. No. 5 is the average of Curve C and D S o . 6 is the average of Curves A and B S o . 7 is the average of Curves E and F These curves are reproduced for the purpose of showing the sharpness of the end-points in the analyses with a dilute titanous sulfate solution according to the method which was developed.
ELECTROMETRIC DETERMINATION O F FERRIC IRON
I697
Summary I. An electrometric method has been described for the determination of small amounts of iron. h procedure for satisfactorily storing titanous sulfate solutions as 2. dilute as 0.0003 normal for periods of several weeks has been given.
FIG.6 Typical curves resulting from the titration of a small amount of ferric iron w t h an approximate 0.001 normal titanous sulfate solution.
3 , A simple electrometric titration apparatus has been described where small volumes of a reagent may be shaken, heated, protected from the air by an inert gas and electrometrically titrated. 4. h platinum electrode with fine platinum wire has been designed for use in hot solutions. j. Results of our work show that it is possible t o determine as small an amount of iron as 0.0003 grams in the ferric condition by direct electrometric titration with an accuracy of one part in a thousand. T~allzamsdoicn,.Mass.
