Effect of formation of insoluble polysilver salts on argentometric titration of mercaptoacids

ANALYTICAL

Effect

BIOCHEMISTRY

2, 158-167

of Formation Argentometric

(1960)

of Insoluble Polysilver Titration of Mercaptoacids

E. M. CARR From the Research Laboratory Received

Salts

on

M. BIT-ALKHAS

AND

oj the Z’oni Compnny, May

Chicago, Illinois

2, 1960

INTRODUCTION

The two most widely used volumetric reagents for the determination of the thiol group are standard solutions of iodine and of silver nitrate. The iodimetric titration, first employed by Klason and Carlson (l), is based on the oxidation of the thiol to the corresponding disulfide. The titration with silver involves formation of non-ionized silver mercaptide (2). In 1946, Kolthoff and Harris introduced amperometric techniques in titration of thiols with silver nitrate in ammonia buffers and were able to perform titrations in the milligram range with excellent accuracy (3). In a later publication, they demonstrated that titration with silver nitrate in an ammoniacal medium is applicable, in general, to all monofunctional mercaptans, that is to all mercaptans containing the thiol group as the only functional group (4). They found, however, that the course of reaction of such mercaptans with iodine depended on whether the thiol group was primary, secondary, or tertiary. Because of its simplicity, good accuracy at low levels, and apparent selectivity t.oward the thiol group, the method of Kolthoff and Harris seemed well suited for biochemical work. Benesch and Benesch adapted it, in 1948, to the determination of thiol groups in materials of biological importance (5). Further in 1955, Benesch, Lardy, and Benesch introduced tris (hydroxymethyl) aminomethane (Tris) as the complexing agent for silver in place of ammonia (6). This enabled them to carry out determinations in neutral media. Because of the susceptibility of thiols to air oxidation in alkaline solutions, the new complexing agent offered an important advantage. Nevertheless, the selectivity of ammoniacal silver toward thiol groups, in the presence of other functional groups which are present in materials of biological interest, was questioned by Cecil in 1950 (7). Kolthoff and Stricks observed that when cysteine is titrated amperometrically with silver nitrate in ammonia buffers, an excess of silver ammine is tied with silver cysteinate to form a higher complex (8). The 158

ARGENTOMETRIC

TITRATION

OF

MERCAPTOACIDS

159

complexing of silver was quite pronounced when the concentration of titrated cysteine exceeded 5 x lo-” ilf. Cecil and i%Phee examined conditions for the potentiometric titration of certain thiols with silver nitrate (9). The pH range had to be controlled carefully; otherwise, high results were obtained, and were attributed to the format-ion of higher complexes between the silver mercaptide and additional silvci reagent. In 1958, Burton published a crit,ical review of argentometric methods as applied to cysteine, glutathione, and other substances (10). In numerous cases, the observed consumption of silver was markedly higher than can be accounted for by the stoichiometric relation RS- + Ag+ ----fRSAg The errors occurred regardless of whether Tris or ammonia was used to complex silver. In view of these findings, Burton concluded that amperometric t.itration of thiols with silver is not universally applicable. This paper presents evidence that formation of salts of the general formula AgSRCOOAg is responsible for high results observed when argentometric titration is applied to substances possessing both thiol and carboxylic acid functions. Since many biologically important thiols belong to this class of compounds, high results reported by Cecil, Burton, and others (11) may be caused by format.ion of analogous complexes. ESPERIMEKTAL Four mercaptoacids were analyzed by selected volumetric methods. The compounds were: mercaptoacetic acid, ,&mercaptopropionic acid, mercaptosuccinic acid, and a,&-dimercaptoadipic acid. These substances were titrated by several independent methods: iodimetric, alkalimetric, and argentometric. In addition, the following experiments were performed. A solution of the mercaptoacid was introducd into a moderate excess of ammoniacal silver. Under these conditions, three of the four mercaptoacids (dimercaptoadipic acid was an exception) formed precipitates. From a determination of the amount of residual silver in the solution by back titration with potassium thiocyanate, the silver consumed in forming the precipitate could be calculated. The composition of the precipitate was ascertained by the determinations of silver, sulfur, and sulfhydryl content, and by infrared spectroscopy. MATERIALS :Vlercaptoacetic acid (72% solution in water) was distilled at 107”C, at a pressure of 16 mm Hg; lit., 107-108”/16 mm (12a). Since highly concentrated solutions of this acid contain some thiolactone, the concentrate, prior to use, was diluted with water and allowed to stand over-

160

CARR

AND

BIT-ALKHAS

night to open the thiolactone ring. It was established separately that only negligible amounts of disulfide are formed under these conditions. Dithiodiacetic acid was recrystallized from ethyl acetate-benzene, m.p. 102-103°C; lit., 10%109°C (12b). /3-Mercaptopropionic acid was distilled at 106”C, at a pressure of 9 mm Hg; lit., 114-115.5”C at 13 mm (12~). Mercaptosuccinic acid was recrystallized from water, m.p. 152153°C; lit., 151°C (12d). cY,cr’-Dimercaptoadipic acid was recrystallized from ethyl acetate-benzene, m.p. 186-188°C; lit., 188°C (13). METHODS

Volumetric determinations were performed, in most cases, on 0.1 M aqueous solutions of the mercaptoacids. An exception was made in the case of dimercaptoadipic acid because of its limited solubility in water. For this reason, solid dimercaptoadipic acid was weighed out for each individual titration. To facilitate comparison with other mercaptoacids, the results were recalculated to refer to a hypothetical 0.1 M solution (20.82 g/l) of dimercaptoadipic acid. Amperometric titrations were performed using conditions and apparatus recommended by Kolthoff and Harris (3). The 0.1 M stock solutions were diluted lo- or 20-fold. Aliquots of the dilute solutions were introduced into the titration vessel and were titrated with 0.01 M silver nitrate solution. No potential was applied between the rotating platinum electrode (r.p.e.) and the mercuric iodide reference cell except in the titration of mercaptopropionic acid. In the latter case, in the absence of applied potential, an anodic current was observed which was possibly caused by the oxidation of the sulfhydryl group at the r.p.e. This anodic process obscured the end point. To minimize this undesirable effect, titrations of mercaptopropionic acid were conducted at the potential of -0.1 v applied to t,he r.p.e. Under these conditions, the anodic current was absent and the end point was detectable with good precision. The results of the amperometric determinations were calculated back to refer to the concentrations of the original 0.1 III stock solutions. Iodimetric titrations of the thiol groups were performed in weakly acid solutions (pH = 3-4) using 0.1 N iodine solution and starch indicator. These conditions had to be modified in the cases of mercaptopropionic and mercaptosuccinic acids. The former showed a tendency to become oxidized beyond the disulfide stage while the latter, a secondary mercaptan, could possibly enter into a side reaction with iodine and form a sulfenyl iodide (4). To suppress these undesirable reactions, titrations of these two mercaptoacids were performed in strongly acid media (pH = 1) and in the presence of 2-3 gm potassium iodide. The latter was introduced into the solution immediately prior to titration. Alkalimetric titration of the carboxylio groups were performed with

,~R~;ESTO\I~:‘TI~I(’

TITRATION

OF

MERCAP’I’OACIW

161

0.1 LV sodium hydroxide. The end point was detected potentiometrically with a Beckman model G pH meter. The behavior of mercaptoacids and their quantitative analysis in the presence of an excess of silver diammine complex was examined in the following manner. Fifty milliliters of standard 0.1 d4 silver nitrate solution, 1 gm ammonium nitrate, and 2 ml of cont. ammonia were introduced into a loo-ml volumetric flask. Then, an aliquot of the solution of the mercapt,oacid, containing 0.5-2.0 mmoles of the latter, was introduced under continuous swirling. The swirling was continued for about one more minute. Water was added to bring the volume to the loo-ml mark, and the contents were mixed thoroughly. The amount of silver remaining in the solution was determined in the following manner. The suspension of the precipitated material was filtered, and the filtrate was collected in a dry container. An aliquot of the filtrate (25 ml) was acidified with 25 ml of 1 N nitric acid and titrated with standard 0.1 N potassium thiocyanate solution to the ferric alum indicator end point according to the usual Volhard procedure. This procedure gave quantitative result,s when applied to solutions of known concentration of mercaptoacetic acid if the content of the latter was calculated from the stoichiometry of the equation HSCHzCOOH

+ 2Ag+

+ .4g,SCzHzO?

+ 2H+

The same procedure was applied to solutions of mercaptopropionic acid with the following minor changes. Only 1 ml of cont. ammonia was used. The precipitate began to form after several minutes. The suspension was allowed to stand for 2 more hours and shaken occasionally to make the precipitation complete. The concentration of mercaptopropionic acid was calculated assuming that 2 moles of silver were equivalent to 1 mole of the mercaptoacid. The composition of the disilver salt of mercaptoacetic acid obtained above was verified by the following procedure. The fresh precipitate was collected on a sintered-glass filter, washed with water containing 10% methanol, and dried to constant weight over phosphorous pentoxide. To establish the thiol content, the precipitate was dissolved in dilute acetic acid, an excess of potassium iodide was added to precipitate the silver ion, and the resulting mixture was titrated immediately with 0.1 N iodine solution. Another portion of the precipitate was dissolved in dilute nitric acid and titrated for silver ion by the Volhard method using standard 0.1 N potassium thiocyanate. The sulfur content was obtained by destructive oxidation in a Parr bomb by the usual sodium peroxide treatment. The oxidation products were leached out with hot water, and the insoluble silver residue was filtered off and washed. The filtrate was combined with the washings and acidified with dilute hydrochloric acid

0.‘09974.1001 0.0861-0.0871(4) 0.0994-o. 0999 0.0967-O. 0973

Iodimetric

(4) (4)

(3)

OF MERCAPTOACIDS

TABLE 1

found,

Method

PROCEDURES: Concentration

Alkalimetric

VARIOUS

0.1038-O. 1038 (2) 0.0855-0.0860 (4) 0.0994-O. 1000 (3) 0.0980.0981(4)

BY

OF RESULTS

0.12ooXl. 1300 (2) 0.0888~. 0920 (4) 0.1020-0.1080 (4) 0.1150X). 1220 (4)

Amperometric

moles/P

COMPARISON

0.0996-o. 0.0850-O.

solution,

it EVRR

did not depend

-d -e

and

1013 (4)b 0858 (3)<

back-titration

Precipitation

a Range of observed concentrations is given. Figures in parentheses denote number of determinations. * The amounts of mercaptoacetic acid taken were in the range 0.5-2.0 mmole. In this range, the result of a determination on the size of the sample taken for analysis. c The amounts of mercaptopropionic acid taken for analysis were in the range 0.42-1.26 mmole. d Results were not satisfactory from a quantitative view. From the determination of the silver ion remaining in the established that the precipitate contained about 1.8-2.0 moles of silver per mole of mercaptosuccinic acid. e No precipitate was formed between the mercaptoacid and an excess of ammoniacal silver.

Mercaptoacetic Mercaptopropionic Mercaptosuccinic Dimercaptoadipic

blercaptoaoid titrated

TITRATION

z %

it Y e 2

Ez D

ARGENTOMETHIC

TITRATION

OF

MERCAPTOACIDS

163

to a pII of about 3. The sulfate content was then determined in the usual manner by precipitation with barium chloride. The following procedure was used to see whether the precipitate contained ammonia, either present as a silver diammine ion or as an ammonium ion. The precipitate was freshly prepared by introducing mercaptoacetic acid into an excess of ammoniacal silver according to the method already outlined. The precipitate was collected, without washing, on filter paper, and blotted to remove most of the mother liquor. The blotted precipitate was then divided into two portions. Portion 1 was weighed and introduced into a Kjeldahl distillation apparatus to determine the total ammonia content of the moist precipitate. Portion 2 was weighed and dried to determine the amount of mot,her liquor contained in the precipitate. Ammonia bound by the precipitate was then calculated as a difference between the total ammonia and the known ammonia content contributed by the mother liquor. In a separate experiment, it was established that the presence of silver ion does not interfere with the quantitative recovery of ammonia in the Kjeldahl distillation. RESULTS

The various

analytical

determinations TABLE

DETERMINATION OF MERCAPTOACETIC AND BACK-TITRATION WITH VARIABLE

are assembled in Tables l-3. 2

ACID (1.000 MMOLE) BY PRECIPITATION THIOCYANATE IN THE PRESENCE OF

AMOUNTS

OF DITHIODIACETIC

Dithiodiacetic acid introduced (mmolea x 100)

ACID

Mercaptan found (mmoles)

0.00

1.00 3.00 5.00 Average

1.0005 1.0015 0.9955 1.0015 0.9998

DISCUSSION

Mercaptoacetic acid was studied in more detail than the other compounds since it is believed that it is representative of the behavior of mercaptoacids. As can be seen from Table 1, the results of the amperometric determinations were high in each case and out of line with the results of the other procedures. The magnitude of this discrepancy was significant and not explainable in terms of the uncertainty associated with the detection of the end point. In contrast, excellent agreement was obtained between the iodimetric method and the method based on precipitation followed by back-titration with thiocyanate. The alkalimetric titrations gave somewhat higher results, but the discrepancy is very

164

CARR

AND

BIT-ALKHAS

TABLE COMPOSITION

Type

of analysis

Silver Sulfur (-SCH,COO-) NHab

OF THE PRECIPITATE AND AN EXCESS

3

FORMED BETWEEN OF AMMONIACAL Per cent

found

70.09 10.17 29.77a 0.2

MERCAPTOACETIC SILVER

ACID

Per cent calculated for Ar&HzOB

70.54 10.48 29.44 o.oc

a Determined by the iodimetric titration of the thiol group. b Determined on a freshly prepared precipitate which was not washed. Excess moisture was removed by blotting (see Me&&). The figure was corrected for the liquid content of the precipitate (under 5%) and the ammonia content of the mother liquor. c Ammonia content of the hypothetical compound AgSCH&OOAg(NH& is lO.O’%.

likely caused by the presence of small amounts of disulfide or perhaps other impurity of an acidic character. In the course of an amperometric titration, the initial color of the precipitate was yellow. Upon addition of more silver, the color gradually faded until it appeared white. A similar process took place when the mercaptoacetic acid solution was introduced into an excess of ammoniacal silver. In the first few seconds, a yellow color was observed. Agitation of the suspension for about one more minute brought about a sudden change of color from yellow to white. It would seem that the yellow color is characteristic of silver mercaptide of mercaptoacetic acid, while the disilver salt is, apparently, white. With the back-titration method, it was our experience that if, for any reason (such as insufficient swirling), the precipitate failed t,o change from yellow to white, the disilver salt was not formed quantitatively and, as a result, the stoichiometry corresponded to a precipitate containing a mixture of mono- and disilver salts. The precipitate formed in the precipitation and back titration procedure is stoichiometrically the disilver mercaptoacetate, AgSCH&OOAg. This conclusion was confirmed by the direct analysis of the precipitate, as shown in Table 3. Infrared spectra of the precipitate also substantiate this view. The presence of a band at 6.35 ,L rather than at 5.9 p indicates that the carboxylic group is in the ionic form (14~). For example, absorption of the carboxylic group in its non-ionized form in acetic acid occurs at 5.80 ,u, while the carboxylate ion in silver acetate was found to absorb at 6.30 p. Furthermore, the absorption band at 3.9 p, characteristic of the -SH group, is missing, as would be expected if a mercaptide had formed (14b). The amount of ammonia held by freshly formed precipitate, as seen from Table 3, is insignificant. Consequently, silver diammine or ammonium ion cannot be the counterions of the mercaptide

ARGENTOMETRIC

TITRATION

OF

165

MERCAPTOACIDS

or carboxylate groups, and hence silver ion is the only remaining candidate for salt formation. The view that a definite compound has been formed is further strengthened by the fact t.hat the composition of the precipitate remained constant over a fourfold change in the amount of the mercaptoacid (see Table 1, footnote b). Had some other phenomenon been responsible for the consumption of silver (such as adsorption or coprecipitation), the amount of silver tied with the precipitate could be expected to depend on the ratio of concentrations of the silver diammine ion and mercaptoacetic acid. It appears that the carboxyl group of a mercaptoacid is endowed with special properties as regards combination with silver so that, it is not at all comparable to other carboxylates. The difference in behavior of mercaptoacetic acid and its disulfide suggests a plausible structure for the disilver salt of the former. We recall first that the silver ion has a coordination number of 2, but that the bonds tend to be colinear (15). The stability of the group sg+ . . . -oc0

is evidently not very strong for an ordinary carboxylate, such as that in dithiodiacetic acid, for, as the data in Table 2 show, the latter does not remove silver from the silver diammine ion. Furthermore, when silver ion was introduced into an ammoniacal solution of dithiodiacetic acid under conditions which were employed for the amperometric determination of mercaptoacids, there was not indication that the silver diammine ion formed a complex with dithiodiacetic acid. In a mercaptoacid, however, double linkages of the following type become possible:

j.\

/

, ,’

\

‘C-CH,-S’

Ag-0

/

C-CH?-

\

Ag-0

‘?l

/

“‘C:-(:H&

,f Ag- ( )

Ag-0 ‘\

/

9

/ L--s \

\

-hAg-

166

CARR

AND

BIT-ALKHAS

The likelihood of the silver ion being removed from solution is now much greater because both mercaptide and carboxylate groups are involved in bonding. Furthermore, conditions for the buildup of an infinite lattice present themselves, and hence the extra stabilization due to crystallization becomes available. Bonds such as pictured in (I) are not possible in an ordinary carboxylic acid nor in the disulfide, dithiodiacetic acid. In the former no sulfur is present; in the latter an-SS-is available but the affinity of the disulfide group toward metal ions is exceedingly weak. Although complete analytical data for the precipitate of disilver salt were obtained only for mercaptoacetic acid, the behavior of the other mercaptoacids fits the scheme pictured in (I). Amperometric titrations invariably give high results, indicative of an uptake of the silver ion beyond that equivalent to -SH group alone, and under proper conditions a precipitate will form containing a high content of silver. It is of interest to note that an attempt to prepare a trisilver salt of mercaptosuccinate, which has one sulfhydryl and two carboxylic groups, was unsuccessful, although a disilver salt is apparently obtained (Table 1, footnote d). Evidently, the second carboxyl does not have sufficient affinity to hold a third silver. Such behavior is fully in accord with the structure pictured in (I), since the second carboxyl does not have a matching mercaptan group to supply the additional affinity needed to hold the additional silver ion. Many biological substances containing sulfhydryl groups also contain carboxylate ions. It seems likely that in these cases, excess silver ion may also be taken up if the carboxylate group is suitably placed with respect to the mercaptan. If these groups are attached to a macromolecule, lattice formation would not follow, since the groups holding silver would not be free to approach each other. Nevertheless, the amount of silver ion bound could exceed the equivalent amount of mercaptan and hence lead to the high results which have been observed in some amperometric titrations. SUMMARY

High results observed when mercaptoacids are titrated in ammoniacal media with silver nitrate are caused by the formation of insoluble precipitates containing silver in the forms of both the mercaptide and the carboxylate salt. In the presence of a moderate excess of ammoniacal silver, mercaptoacetic acid or p-mercaptopropionic acid can be precipitated quantitatively as a disilver salt. The readiness with which silver combines with the carboxylate ion in compounds of this type is postulated to be favored by the crystal structure of the resulting compound.

ARGENTOXIETRI~!

TITH.4TIOZI

OI:

MERC.4PTOAClW

The authors gratefully acknowledge :rtl\~ic~t~31id lwlp rstcntlrrl I. M. Klote of P\‘orthwstern Univcrsit!..

167

to them by Prof.

REFEREXCES P., .~ND CARLSON, T., ner. 39, 738 (1906). M. W., AND RYLASD, L. B., Znd. Eng. Chem. Ad. Ed. 8, 16 (1936). 3. KOLTHOFF, I. M., AND HARRIS, R. IS., Ind. Eng. Chon. Anal. Ed. 18, 161(1946). 4. KOLTHOFF, I. M., AND HARRIS, W. E., Ad. Chem. 21, 963 (1949). 5. BENESCH, R., AND BENESCH, R. E., Arch. Biochem. 19, 35 (1948). 6. BENESCH, R. E., LARDY, H., ASD BENESCH, R.. J. Biol. Chcm. 216, 663 (195.5). 7. CECIL, R., Biochcm. J. 47, 572 (1950). S. KOLTHOFF, I. M., .~ND STRICKS, W., J. dw. Chon. Sac. 72, 1952 (1950). 9. CECIL, R., .~ND MCPHEE, J. R., Biochem. J. 59, 234 (1955). 10. BURTOS, H., Biockim. et Biophys. Acta 29, 193 (1958). 11. SLUYTERMAN, L. A., BioclGm. et Biophys. Actcz 25, 402 (19573. 12. “Dictionary of Organic Compounds” (I. Heilhron and H. M. Bunbury, eds.). Oxford Univ. Press, New York, PIT.Y., 1953. (n) Vol. IV. p. 495; (h) Vol. IT, p. 144; (c) Vol. IV, p. 496; (d) Vol. IV, p, 497. 13. HREDYA, A., Ber. 71B, 289 (19%). 14. BELLAMY, L. J., “The Infrared Spectra of Complcs Molwulrx.” J. Wiley 6 Sons, New York, N. Y., 1958. (n) pp. 174-6; (b) p. 352. 1.5. WELLS, -4. F., Structtwal Inorpanir Chrmistry,” pp. 504-6.Oxford Univ. Press, London. 1945. 1. KLASOS,

2. TAMELE,