Gravimetric and Volumetric Determination of Germanium HOBART H. WILLARD AND C. W. ZUEHLKE', University of Michigan, A n n Arbor, Mich. DISCUSSION
A volumetric method for germanium was developed, based on the quantitative formation of potassium thiogermanate in acetate-buffered solution b y treatment with potassium sulfide, removal of the excess hydrogen sulfide b y carbon dioxide, and titration of the sulfur with iodine. The presence of large concentrations of sodium chloride leads to both a poritive and a negative error in the application of this method, depending on conditions. The negative error ir due to the adsorption of the sulfide ion on the free sulfur formed b y titration with iodine, while the positive error is probably connected with the formation of a higher sulfide. Germanium i s quantitatively precipitated b y 5,6-benzoquinoline as a complex oxalate. Although the compound cannot b e weighed directly, it can b e ignited to germanium dioxide, thus affording a good gravimetric method. H i g h concentrations of sodium chloride prevent this precipitation.
The success oT the thiogermanate method depends upon the stability of potassium thiogermanate in acid solution. It waa previously reported that B solution of potassium thiogermanate contains insufficient free sulfide ion to precipitate cobalt or nickel sulfide. The carbon dioxide treatment readily removes the excess hydrogen sulfide from solution without decomposing the thiogermanate. Further experiments were made by subjecting a solution of potassium thiogermanate to a rapid stream of carbon dioxide. Samples were withdrawn periodically and titrated with iodine to observe the decrease in sulfideion concentration with time. The results (Table I) show that the loss after 2 hours' treatment is negligible. I n order to attain complete reaction between hydrogen sulfide and germanium dioxide it is essential that a high concentration of sulfide ion be applied. The reaction is rapid but, nevertheless, incomplete when only the equilibrium concentration of a saturated solution of hydrogen sulfide is used. The data given in Table I1 were obtained by treating buffered solutions of germanium dioxide with a constant flow of hydrogen sulfide for varying periods of time. The excess hydrogen sulfide was removed and the solutions were titrated in the usual manner. Although the reaction is nearly complete after 5 minutes, quantitative combination is not attained even after 20 minutes. The data in Table I1 indicate that the hydrogen sulfidegermanium dioxide system behaves as though an equilibrium is rapidly established a little short of complete reaction. Prolonged treatment with hydrogen sulfide is without beneficial effect but quantitative results are obtained when a condition of supersaturation is maintained for a short period of time by the addition of potassium sulfide. Experiments concerned with the effect of temperature on the completeness of reaction are in agreement with this concept. Increasing the temperature decreases the solubility of hydrogen sulfide and, in general, decreases the amount of sulfur combined. Solutions of germanium dioxide buffered with acetic acid and potassium acetate were treated with hydrogen sulfide for 15 minutes at various temperatures. The amount of sulfur combined was determined in the usual manner. At 25" C. the reaction was 96% complete while a t 80" C. only 65% of the germanium reacted in 15 minutes. I n unbuffered solution only 70 and 11% reaction was obtained a t these two temperatures, respectively. The thiogermanate ion cannot be titrated directly with iodine to a satisfactory end point. Apparently the free sulfur formed
W
ILLARD and Zuehlke (.I) recently reported the preparation of potassium thiogermanate. This compound serves as the basis of a volumetric method for germanium. Germanium dioxide, in acetate-buffered medja, may be quantitatively converted to potassium thiogermanate by treatment with hydrogen sulfide or potassium sulfide: 2Ge(OH),
+ 5HaS + 2Khc
= K2GelSs
+ 2HAc + 8H20
The excess hydrogen sulfide is removed by passing a rapid stream of carbon dioxide through the solution, after which the sulfide ion is determined by treatment with an excess of standard iodine solution and back-titration with sodium thiosulfate solution:
+
KzGeZSb 512
+ 8Hn0
=
55
+ 2Ge(OH), + 8HI + 2KI
REAGENTS
Germanium dioxide solution, , approximately 2 grams of germanium dioxide dissolved in 1hter of water. This solution is stable and remains perfectly clear. Acetic acid, 2.5 molar. Potassium sulfide solution, approximately 8 grams of potassium hydroxide dissolved in 100 ml. of water and saturated with hydrogen sulfide. The solution was cooled to 0" C. during the saturation with hydrogen sulfide to avoid formation of thosulfate and polysulfide by atmospheric oxidation. Standard iodine solution approximately 0.1 N . Standard sodium thiosul!ate solution, approximately 0.1 N . PROCEDURE
A measured aliquot of the germanium dioxide solution was transferred to a 25-cm. (10-inch) test tube and diluted to 25 ml., and 20 ml. of potassium sulfide solutlon were added, followed by 15 ml. of acetic acid solution. The acid was allowed to run down the side of the test tube to promote better mixing of the solution and to avoid a violent effervescence of hydrogen sulfide a t the surface. The solution was allowed to stand for 5 minutes. Carbon dioxide was bubbled through the solution rapidly by means of a delivery tube extending to the bottom of the test tube. A mechanical stirrer was used to break up the bubbles of carbon dioxide to promote more efficient removal of the excess hydrogen sulfide. Twenty rmnutes' passage of the gas was found to be sufficient to give a negative test for hydrogen sulfide in the issuing gases by lead acetate test paper. The solution was transferred to a large beaker and diluted to 1 liter. A measured excess of standard iodine solution was added allowed to stand for 15 minutes, and back-titrated with standard sodium thiosulfate solution uslng starch solution as an indicator. The end point may be determined to *0.05 ml.
Stability of KZGerSs Solution to COZTreatment
Table 1. Time of Cor Paeaage Min.
0.1 N Iodine S o h -
Time of Con Paasage Min.
0.1 N Iodine 501u-
M1.
15 30 60 120
10.58 10.58 10.57 10.51
15 30 60 120
10.67° 10.40 10.39 10.34
a
tion Consumed
Apparently some exceas hydrogen sulEde still remained.
Table 11.
322
Rate of Formation of KsGezSs
Ma.
Time of HnS passage Min.
48.6 48.6 48.6 48.6
5 10 15 20
GeOn
* Present address, General Chemical Company, New York, N. Y.
tion Consumed M1.
UeOz Found At 25' C. At 40' C. Ma. Ma. 47.0 48.1 48.3 48.2
47.7 48.1 48.5
ANALYTICAL EDITION
May, 1944 Table 111.
Dotormination of Germanium by Thiogermanrte Method
GeOr Taken
GeOs Found
GeOn Taken
Mv.
Ma.
Ma.
GeOn Found
Mv .
52.2 52.2 52.2 52.2 49.8
52.2 52.1 52.2 52.1 49.8
39.9 29.8 20.0 10.0
39.9 29.8 20.0 9.9
in the reaction adsorbs some of the thiogermanate ion and thereby prevents complete reaction and leads to a fading end point. The extent of this adsorption was shown to be dependent upon the concentration of electrolyte present. By dilution to a volume of 1 liter before titration a true end point was obtained. This dilution serves not only to reduce the concentration of electrolyte but also the concentration of the thiogermanate ion in equilibrium with the free sulfur as it is formed. The amount of iodine consumed when the equivalent of 50 mg. of germanium oxide was titrated in a volume of 250 ml. rather than 1 liter was found to be about 2% less than the true value. This method is fairly simple, requires about 1.5 hours for a single determination, and is capable under favorable conditions of yielding precise results, as is shown in Table 111.
323
The low results in the sodium chloride solutions are unquestionably due to an adsorption error. The high results which arise when the sodium chloride is present during the sulfiding reaction are, on the contrary, not easily explained. This positive error may be connected in some manner with the formation of a higher sulfide of the type K2Ge2E$(I), but not enough information concerning these thiogermanates is available at this time to arrive at any satisfactory conclusion. "he thiogermanate method, under favorable conditions, is capable of yielding results precise to 0.1 mg. of germanium, owing to its low equivalent weight in .the thiogermanates (1 ml. of 0.1 N iodine solution is equivalent to 2.092 mg. of germanium dioxide, or to 1.452 mg. of elemental germanium). The scope of the method is seriously limited, however, by the sodiurp chloride error, since germanium is usually separated by distillation from constanbboiling hydrochloric acid. Undoubtedly many specific analytical situations do exist in which the method could be advantageously employed, particularly where the amount of germanium is low. Since germanium is frequently encountered as a minute constituent, the possibility of applying this method on the micro scale should not be overlooked.
Table IV. Effect of Sodium Chloride on Thiogermrnrte Method -GeOz Found
INFLUENCE OF SODIUM CHLORIDE ON THIOGERMANATE METHOD
Germanium is usually separated for analysis by distillation as the tetrachloride, or more probably, as a complex acid of the type HZGeCla. Since this distillation is most efficiently carried out from constant-boiling hydrochloric acid, the final determination of germanium must frequently be made in a solution containing large amounts of chloride ion. It has already been shown that high concentrations of electrolyte lead to a false end point in the thiogermanate method. Accordingly an investigation was undertaken to determine the effect of sodium chloride on the application of this method. These experiments may be divided into three series:
GeOr Taken
MU. 10.0
No NaCl
added MV. 9.9
20.0
20.0
29.8
29.8
39.9
39.9
49.9
49.8
Gramsof NaCl added
NaCl added after sulfiding
Me.
NaCl Added before Sulfiding Normal Inverted titration titration
Moa
MU.
20 30 20 30 20 30 20 30
24
INTERFERING SUBSTANCES
1. Varying amounts of sodium chloride were added to the solution just prior to titration. 2. Varying amounts of sodium chloride were added to the solution prior to the sulfiding reaction. 3. Varying amounts of sodium chloride were added to the solution prior to the sulfiding reaction and the sulfur retained was determined by an inverse titration technique. I n this procedure a solution containing an excess of iodine was diluted to about 800 ml. and the thiogermanate solution was added slowly with stirring. The sulfur was theieby formed in a solution containing very little unreacted thiogermanate ion and the sodium chloride was allowed to reach its highest concentration only a t the end of the reaction. The results of these experiments are given in Table IV.
I n cOnsidering the data collected in Table IV, except for a few irregularities, the following generalizations may be made. I n the absence of sodium chloride the results are precise within the limits of experimental error. When sodium chloride is added just before the titration a large negative error arises due to adsorption of the thiogermanate ion on the free sulfur formed. The magnitude of this error is proportional to the concentration of sodium chloride present and also the amount of germanium dioxide taken. When, however, the sodium chloride is present during the sulfiding process, the results are higher, apparently because of some kind of a compensating positive error. With small amounts of germanium dioxide this positive error predominates, leading to results which are above the theoretical. With large amounts of germanium dioxide the negative adsorption error predominates, leading to low results. When the inverted titration technique is used the results are generally higher than those obtained by the usual method, owing to the partial elimination of the negative adsorptirin error.
Metals which form insoluble sulfides a t a pH of 4.6 must be absent,. GRAVIMETRIC METHOD
Willard and Toribara (3) showed that tin forms a complex oxalate and assigned the formula KsSn2(C204)r to its potassium salt. Tchakirian (2) attempted to prepare the analogous compound of germanium but found that the potassium salt was too unstable to isolate. He did, however, establish the existence of trioxalatogermanio acid, HZGe(C2O&, by isolating it as the quinine and strychnine salts. I t was found that 5,bbenaoquinoline also forms an insoluble derivative with trioxalagogermnnic acid which shows promise as an analytical precipitant for germanium.
REAGENTS.Germanium dioxide solution, a proximately 2 grams of germanium dioxide dissolved in 1 liter ofwater. 5,bBenzoquinoline oxalate solution. Ten grams of 5,&benzoquinoline (may be obtained from the Eastman Kodak Company) were treated with 5 grams of oxalic acid dissolved in 50 ml. of water. The suspension was heated to promote solution of the base, filtered while hot, and diluted to 500 ml. EXPERIMENTAL.Measured aliquots of the germanium dioxide solution were diluted to 400 ml., treated with 5 grams of oxalic acid, and heated to promote the formation of trioxalatogermanic acid. Twenty-five milliliters of the reagent solution were added. Upon cooling to room temperature the derivative was precipitated in long crjvtalline needles. The solutions were allowed to stand overnight to ensure complete precipitation, and filtered. After washing with a dilute solution of oxalic acid and the reagent, the precipitate was ignited in a platinum crucible in a muffle a t 700" to 800" C. to a pure white residue of germanium dioxide. Precipitation was found to he quantitative, as shown by Table V
INDUSTRIAL AND ENGINEERING CHEMISTRY
324
Table V. Precipitation of Gnmsnivm with 5,6-Benzoquinoline Ge01 Found XL7.
Ge01 Taken
Mo.
Mg.
MU.
843 54.2
84.2 54.0
49.9 49.8
49.8
Ge01 Taken
GeOi Found
50.1
5,6Benaoquinoline would appear to be a superior precipitant for germanium for the fallowing reason?: The precipitation procedure is simple and yields a product of very high molecular weight. The crystalline precipitate is easy to filter and wash and there is accordingly little danger of contamination by foreign ions. TM: precipitate is readily converted to germanium dioxide. A specimen of 5,fi-benmquinoline trioxaJatogermanate was purified by recrystallization from water. Weighed samples were ignited to the oxide and from the loss in weight 8. ratio of 1.98 males of base for each mole of germanium was edouleted. B e es,uuse B definite composition was indicated, an attempt was msde to preoipitate the derivative in pure form and to weigh it directly. In all these attempts some of the reagent was copreeipitated, even when its concentration was reduced to the point where quantitative preoipitatian of the germanium was no longer obtained. The precipitate was also found to lose weight slowly i n s desiccator and to come to a constant value only after 30 hours. A concentration of 20 to 30 grams of sodium chloride in a vol-
Vol. 16, No. 5
ume of 400 ml. completely prevented precipitation of this derivstive. I t would appear that the germanium is firmly bound in B complex ion of the type &CIS-- under these conditions and t h e r e fore fails to form the oxalate complex which is necessary for precipitation. Of the other members of the fourth periodic group, titanium, zirconium, and tin are known to form complex oxalates. 5 , 6 Benzoquinoline was found to give a precipitate when solutions of these elements were treated by the same procedure as was used for germanium. The precipitate formed with tin resembles that of germanium very closely, while the products obtained with titanium and sirconium appeared to be much more insoluble and flocculent in nature. These compounds are being investigated. INTERFER~NGSUBSTANCES. All elements which form insoluble oxaltLtes when treated with oxalic acid must be absent. Titanium, zirconium, tin, and to a lesser extent, iron, form complex oxalates which are precipitated by the reagent.
(1942).
(4) Willard. H.H.,and Zuehlke. C. W.. Ibid., 65.1887 (1943).
FROM a disiartation avbmitted by C. W. Zuehike to the Graduate Bohool of the University of Miahigan in psrtid fulfillment oi the reqvirernents ior the demee of doctor of philwmhyin ehernintry
Support for Kjeldahl
Flasks
JACOUELINE FRONT, Mellon Institute, Pittrburgh, Pa.
BECAUSE
of their peculiar shape, with round bottom and elongated wide neck, Kjeldahl flasks do not fit into any rack ordinarily found in a chemical laboratory. Tall rectangular supports have been used with clamps holding the flssks, but such arrangements are unstable and working the clamps is ineonvenient and time-consuming. I t is considered desirable to have a support which will grasp the flask, hold it firmly, yet release it with one movement. Furthermore, it should hold the flash vertically, snd provide easy access to each f l a s ~ as , samples are measured into them.
The necks and lips of the flasks me not uniform in 8128; t h e r e fore, the holding device cannot be rigid if i t is t o accommodate these differences. To salve this problem in this laboratory, a holder of simple design has been made, The beveled openings are provided with B spring (see drawing). The fiberboard (or plastic) disk has s k places for flasks, slthough any number can he used. This disk is constructed so that i t rotates on a central shaft and each flask is equdly accessible. To stabilize the support, a heavy circular base of metal is employed, with a diameter of 35 em. (14inches) extending out as far as the widest part of the hanging flask. The upper portion of the neck of the flask is pushed into the disk opening. so that, when the flask drops down, i t catches on the bevel and is held firmly by thespring about 1.25 em. (0.5 inch) below thelip. Toremoveaflask from the rack it is merely necessary t o raise it straight up and slide i t out. This stand is very useful ' ' dding the flasks while >rsolid or liquid samare measured into I. It is substantial Igh to support several ml. Basks and their ents with no tend' to tip. Ammonia u..dlations and nitrogen digestions have beenfaciiitated considerably by this inexpensive a n d easily constructed support.
