Determination of silver by potentiometric titration with sodium thiosulfate

MICROCHEMICAL

JOURNAL

Determination

RAY

16, 410-418

(1971)

of Silver by Potentiometric with Sodium Thiosulfate

Titration

E. HUMPHREY, ROSE MARY MANISCALCO, AND WILLIE HINZE

Department

of Chemistry,

Sam Houston Received

State University,

March

Huntsville,

Texas 77340

6, 1971

Procedures have been reported for the potentiometric determination of low concentrations of silver ion by precipitation as the sulfide, where the sulfide ion is derived from the hydrolysis or decomposition of an organic sulfur compound such as thioacetamide (2) or dithiooxamide (5). Thiosulfate ion reacts with silver ion to form a slightly soluble salt which also decomposes in aqueous solution to yield silver sulfide (7). A&O,

+ H,O -+ Ag,S + S0,2- -t 2 H’.

Although this reaction is well known, few analytical applications have been reported. It has been used as the basis for an acidimetric technique for thiosulfate in which an excess of silver nitrate is added and the acid formed titrated with standard base (I). Muller (6) reported, in a potentiometric study of the interaction of silver ion with thiosulfate ion, that a sharp break occurred on adding thiosulfate ion to a solution containing silver ion but that the potential breaks on adding silver ion to a thiosulfate solution were very small. Titrations were conducted at the 0.1-M level with reasonably accurate results but were not studied as a quantitative procedure, Apparently, there have been no reported analytical applications of the potentiometric titration of silver ion with sodium thiosulfate solutions. We have found that small amounts of silver can be titrated quite accurately with sodium thiosulfate solutions employing a potentiometric end point. This titration can also be used to standardize sodium thiosulfate solutions using silver nitrate as a primary standard. Titrations can be performed in the presence of a number of other metals with good results. Successful titrations were possible using various temperatures and different electrolytes or acids. Also the titration of thiosulfate with silver ion gave reasonable accuracy although this titration was not studied extensively in this work. MATERIALS

AND

METHODS

All chemicals used were the best available reagent grade. Distilled water was used for all solutions. Potential measurements were made 410

TITRATION

OF

SILVER

WITH

411

THIOSULFATE

with a Beckman Zeromatic pH meter and a Corning Model 10 expanded scale pH meter equipped with a Sargent silver electrode and a saturated calomel electrode. Silver nitrate solutions were standardized by potentiometric titration against pure sodium chloride. Sodium thiosulfate solutions were standardized in the usual way by use of potassium iodide and potassium iodate with starch solution as indicator. Titrant solutions were usually from 0.005 to 0.01 M. Amperometric titrations were conducted employing a rotating platinum electrode with a Sargent synchronous rotator connected to a saturated caIome1 electrode. Microburets were used to allow accurate measurements of volumes. The procedure involved simply adding an aliquot of a standard silver nitrate solution to a 400-ml beaker, adding a volume of 0.01 M EDTA solution if a second, interfering ion was present, and diluting to 100 ml with distilled water or an aqueous solution of a strong acid or salt. In some experiments ethanol or aqueous ethanol was used. The electrodes were inserted and addition of the standard sodium thiosulfate solution begun. A few titrations involved titrating thiosulfate with standard silver nitrate. The majority of the titrations were at room temperature, although both higher and lower temperatures were tried. RESULTS

Determination of Low Concentrations of Silver Accurate results were obtained for the determination of quite small amounts of silver ion. Data for these titrations is presented in Table 1. At concentrations of 0.2 ppm or greater the potential break at the equivTABLE 1 TITRATION OF Low CONCENTRATIONS OF SILVER u

Ag present (ppm) 2140 113 56.5 28.2 14.1

5.65 2.40 1.35 0.67 0.27 0.15

Ag found (ppm) h 2133

111 57.0 28.3 14.1 5.72 2.42 1.40 0.69 0.28 0.14

Millivolts c 480 290 270 255 220 200 130 120 105 70 60

‘1Solutions were 0.1 M in HClO,. h Average of three titrations which were in close agreement. c Estimated extent of potential change through the region of the end point,

412

HUMPHREY,

MANISCALCO, AND HINZE

alence point was large enough so that location of the end point was not difficult. Below this level the titration curve was not sufficiently sharp to allow accurate location of the end point. These results compare favorably with photometric procedures for silver at these low levels (3, 4). The accuracy of the titration is probably somewhat better. A few minutes are required to attain a stable potential reading in the vicinity of the end point. The potential breaks at the various levels of silver ion concentration are shown in Table 1. Titration curves for various levels are shown in Fig. 1 and for the lowest concentrations tried in Fig. 2. Titrations with Other Metal Ions Present In order to evaluate the extent of possible interferences due to the presence of other ions, titrations were performd with added amounts of certain other metal ions present. These findings are summarized in Table 2. No interference was noticed in the case of cadmium, cobalt(II), nickel(II), or zinc. Silver was accurately titrated in the presence of considerably higher concentrations of these ions. On the other hand, the presence of lead(H), copper( iron(II1) , and mercury(I1) leads to serious problems due to the interaction of these metals with thio-

NCI,S, 0,) Ml. Fro. 1. Potentiometric titrations for various concentrations of silver ion.

TITRATION

OF

SILVER

I

WITH

I

413

THIOSULFATE

I

I

20

Na, 5, O,,MI FIG. 2. Potentiometric titrations of very low concentrations of silver ion.

sulfate. In order ‘to avoid the interference of all of these except mercury(II), an EDTA solution was added to complex the metal ions. This procedure was used by Kalbus (5) in his study of the dithiooxamide titration of silver. The metal-EDTA complexes of these ions, with the exception of mercury (II), do not react with thiosulfate. The silver titration does not appear to be affected by the presence of EDTA. Mercury(I1) reacts with thiosulfate to form a yellow precipitate. Silver can not be determined by this procedure in the presence of this ion.

TABLE 2 TITRATION

Ag present (ppm) 10.8 10.8 10.8 10.8 10.8 21.6 21.6

OF SILVER

IN THE PRESENCE OF OTHER

Metal present (ppm) Zn(II)

63

Pb(II) Cd(B) Co(I1) Ni(I1) Fe(II1) CLl(I1)

52a 56 59 59 56 I’ 64”.”

METAL

IONS

Ag found (ppm) 11.0 10.8 10.9 10.9 11.0 21.9 21.8

a Solution contained 0.5 mmoles of EDTA. bAn end point could be obtained without EDTA present but the potential break was not as large as with EDTA.

414

HUMPHREY,

Standardization

of

MANISCALCO,

Sodium Thiosulfate

AND

HINZE

Solutions with Silver Nitrate

Titrations involving larger amounts of silver ion, where the concentrations of both the silver and thiosulfate solutions were of the order of 0.1 M, also showed reasonably good results. Several thiosulfate solutions were standardized by the use of potassium iodate and potassium iodide and these results compared with those obtained by titration of standard silver solutions. Agreement of these titrations was probably within the limits of experimental error. These comparisons are shown in Table 3. It appears that silver can be determined at this concentration level by the use of potentiometric sodium thiosulfate titrations. In these titrations it was necessary to agitate the silver electrode occasionally to remove silver sulfide which tended to cling to it. Eflects of Acids, Salts, and Temperature

Muller’s (6) study appears to be the only report in the literature concerned with the potentiometric titration of silver with thiosulfate. It was reported that an accurate titration could be made if sodium acetate was added to neutralize the hydrogen ion produced by the reaction of thiosulfate with silver. The suggestion was made that the hydrogen ion destroyed the thiosulfate leading to an error. In this work titrations were accurate when the silver ion was in distilled water only or when the solution was 0.1 M in perchloric acid, nitric acid, sodium acetate, or potassium nitrate. Hydrogen ion did not appear to have adverse effect on the titration. In fact, the time required to obtain a stable reading in the region of the end point was somewhat less in the acid soluions. The only effect noted when sodium acetate was added was a slight decrease in the magnitude of the potential break at the end point. The effect of TABLE

3

COMPARISONOF SILVER NITRATE AND POTASSIUMIODATE STANDARDIZATION OF SODIUMTHI~WLFATE SOLUTIONS Molarity Na2S203Solution

KIOs

1 2 3 4 5 6 7 8

0.0079 0.0083 0.0632 0.0953 0.0992 0.1045 0.1105 0.1250

&WA 0.0078 0.0082 0.0635 0.0956 0.0991 0.1046 0.1104 0.1251

TITRATION

OF SILVER WITH THIOSULFATE

415

the presence of ethanol, perchloric acid or sodium acetate on the poten tiometric curves is shown in Fig. 3. The presence of an inert electrolyte, such as potassium nitrate, at the 0.1-M level also had no effect on the results. It was also found that silver ion could be titrated in the presence of chloride ion. Silver sulfide formed at the expense of the chloride. The potential break was considerably smaller and the time required somewhat greater than for titrations where there was no chloride present. No titrations could be obtained in solutions containing ammonium hydroxide. Apparently the reaction did not occur under these conditions. Heating the solutions to 45-50” C speeded up the titrations of silver with thiosulfate considerably with no apparent effect on accuracy or magnitude of the potential break. The time required in the vicinity of the end point was reduced almost 50% at the level of l-50 ppm silver. The use of ethanol as the solvent for the titration and 50% aqueous ethanol which was 0.05 M in perchloric acid was also tried. The potential break was somewhat higher in ethanol only and also slightly higher

B

EtOH-HC104

C

H,O-HCIO,

D

H,O-NaAc

\\I

MV

Na, S, O,,MI FIG. 3. Potentiometric titrations of silver ion at approximately under various conditions.

the 30-ppm level

416

HUMPHREY,

MANISCALCO, AND HINZE

in the aqueous ethanol. There was no significant difference in the time required but the error in recovery of the silver ion was slightly greater. Titration

of Thiosulfate with Silver Ion

During the course of the addition of silver ion to a solution containing thiosulfate ion no precipitation occurs until 2 moles of silver are added for each mole of thiosulfate. Precipitation takes place quite rapidly at this point and only slightly high results are found if the appearance of a precipitate is taken as the end point. The error, compared to the potentiometric end point, is of the order of 2-3 % . The potentiometric curve is quite erratic showing a general decreasein potential until somewhat beyond the midpoint in the titration followed by an increase in potential up to the end point at which time a very sharp drop occurs (Fig. 4). A slight inflection is evident at the point where the ratio of silver to thiosulfate is 1: 1 in the titration conducted at a lower temperature, about 2-5” C. The curves resemble quite closely those reported by Muller. Reasonably accurate results were obtained in these titrations when the concentration of thiosulfate was in the range of 40-80 ppm. Results are shown in Table 4. Titrations were not attempted at lower concentrations. The time required to obtain readings was less at room temperature. At elevated temperature the potential break was approxi-

I--

\ I-

,-

,-

II”

40

20

Ag NO, ,M13’ FIG. 4. Potentiometric titrations of 5.5 and 110 ppm thiosulfate ion with silver nitrate at room temperature.

TITRATION

OF SILVER TABLE

TITKATION

Na&O,

OF SODIUM

WITH

THIOSULFATE

417

4

THIOSLJLFATE

WITH

SILVEX

NITRATE

NasSnO:i found (ppm)

present (ppmi

80 81 JO 41

79 39

maely the same as in the reverse titration while at low temperature the break was smaller. Amperometric Titrations Several amperometric titrations were carried out involving addition of silver ion to thiosulfate and the reverse at the 50-ppm level of concentrations. Results were less accurate than in the potentiometric titrations so that this work was not pursued further. DISCUSSION

The earlier study by Muller (6) was not concerned with the analytical applications of the titrations and the results indicated that the accuracy was not particularly good. It was stated that sodium acetate had to be present for a good titration of silver with thiosulfate. Our results show that the accuracy of the titration of silver with thiosulfate is quite good, even at the low parts per million level, in the presence of either acids or salts. The only significant disadvantage is the time required to obtain a stable reading and this can be shortened by heating the solution during titration. The application to the low concentrations of silver is made possible by the use of an expanded scale potentiometer. Results are somewhat comparable to those reported by Kalbus (5) for titrations using dithiooxamide. The time required is about the same, although the sensitivity of the thiosu!fate is probably lower as the potential break is less. However, sodium thiosulfate is available in relatively high purity whereas the dithiooxamide apparently is not. Results are also reasonably good at higher concentration levels as shown by the data on standardization of sodium thiosulfate solutions with silver nitrate in Table 3. This procedure is one of the few alternatives to an iodometric method for standardization. Our work indicated that reagent grade Na,S,O:,. 5H,O is of very good quality and that titrant solutions could be made up by weight and volume for all but the most exacting work. The titration of thiosulfate with silver ion does not appear to be very

418

HUMPHREY,

MANISCALCO,

AND

HINZE

useful although quantitative results can be obtained. The potentiometric curve (Fig. 4), although erratic, is reproducible at lower temperatures and is similar to that reported by Muller. The potential break at a stoichiometry of two silver ions to one thiosulfate is very large. Precipitation does not occur until the end point is reached, and it is possible to detect this point visually by the appearance of the insoluble silver sulfide with an error of only 2-3%. Stable potential readings are obtained in a matter of a few minutes in this titration. SUMMARY Silver can be determined in the low parts per million range by potentiometric titration with sodium thiosulfate solution. Sodium thiosulfate solutions can be standardized by potentiometric titration of silver nitrate. Some time is required to achieve a stable potential reading in the vicinity of the end point. Reduced temperatures are required to obtain a useful potentiometric titration on adding silver ion to a thiosulfate solution. Quantitative results can be obtained in this titration also. ACKNOWLEDGMENT The support of the Robert A. Welch Foundation of Houston, Texas is gratefully acknowledged. Willie Hinze was a National Science Foundation Undergraduate Research Participant during the summer of 1969. Thanks are also due to Paula Hickman, Charles Eastland, and Alvin Vaught for obtaining some of the data. REFERENCES 1. Bodnar, J., New and simple method for the titrimetric determination of thiosulfate, even in the presence of sulfite. Z. Anal. Chem. 53, 37-41 (1914). 2. Bush, D. G., Zuehlke, C. W., and Ballard, A. E., Volumetric determination of silver using thioacetamide Anal. Chem. 31, 1368-1371 (1959). 3. Cave, G. C. B., and Hume, D. N., Calorimetric determination of silver with p-dimethylaminobenzalrhodamine. Anal. Chem. 24, 1503-1505 (1952). 4. Dux, J. P., and Feairheller, W. R., Spectrophotometric determination of trace amounts of silver by dithiol. Anal. Chem. 33, 445447 (1961). 5. Kalbus, L. H., and Kalbus, G. E., Potentiometric determination of silver with dithiooxamide. Anal. Chim. Acta 39, 335-340 (1967). 6. Muller, V. E., Elektrometrische Verfolgung der Reaktion zwischen Thiosulfat und Silvernitrat. Z. Anorg. Al/g. Chem. 134, 202-207 (1924). 7. Przybylowicz, E. P., and Zuehlke, C. W., Silver. In “Treatise on Analytical Chemistry” (I. M. Kolthoff and P. J. Elving, eds.), Part II, Vol. 4, pp. l-63. Interscience, New York, 1966.