V O L U M E 2 7 , NO. 8, A U G U S T 1 9 5 5
1331
per thousand. A number of other cooperative samples have been analyzed in a standards certification program with the NationaI Bureau of Standards and other laboratories with excellent checks and equally good precision. CONCLUSION
The tetraphenylarsonium method described by Potratz and coworkers is highly recommended for the determination of cobalt in stainless steels and other ferrous samples. The method is rapid and precise, being equally as good as the usual reference procedure. Exceptionally good precision is obtained and the procedure has been applied to a wide variety of materials such as the Inconels, nickel, chromium, and boron carbide.
LITERATURE CITED
(1) Affsprung, H. E., Barnes, N. A., and Potratz, H. A., ANAL,CHEM., 23, 1680 (1951). (2) American Society for Testing Materials, Philadelphia, Pa., “ASTM Methods for Chemical Analysis of Metals,” 1950. (3) Lundell, G. E. F., Hoffman, J. I., and Bright, H. il., “ChemicaI Analysis of Iron and Steel,” Wiley, New York, 1931. (4) Potratr, H. A , , and Rosen, J. M., ANAL.CHEM.,21, 1276 (1949). (5) Sandell, E. B., “Colorimetric Determination of Traces of Metals,’ ’ 2nd ed., Interscience, Sew York, 1950. RECEIVED for review February 2, 1955. Accepted March 23, 1955. T h e Knolls Atomic Power Laboratory is operated b y t h e General Electric Co. for t h e Atomic Energy Commission. T h e work reported here was aanied o u t under contract No. W-31-109 Eng-62.
Volumetric Determination of Mercury and the Use of Mercury Salts as Primary Acidimetric Standards SANTI R. PALIT and G. R. SOMAYAJULU lndian Association for the Cultivation o f Science, Calcutta 32, lndia
A rapid and accurate \olumetric method for the determination of mercury consists of converting the mercury into neutral mercury(l1) oxide, which is held in solution by complexing with a small quantity of acetamide or a large quantity of urea. The mercury is estimated by the acidimetric determination of the alliali liberated when potassium iodide or sodium thiosulfate is added to this mercury(I1) oxide complex. Because mercury(1) ion disproportionates quantitatively into mercury(I1) ion and metallic mercury when alkali is added in the presence of an amide, it can be estimated by determining the mercury(l1) ion in solution after such disproportionation using the method described. The accuracy with which mercury can be estimated with this method has prompted the recommendation of mercury salts as reliable primary acidimetric standards.
D
working in this laboratory, has developed a method
As for( lalkalimetric )? estimation of mercury utilizing the well
known solubility of mercury(I1) oxide in acetone. T h e method consists of conversion of the mercury(I1) salt into mercury(I1) oxide, which remains in solution in presence of acetone, followed by the liberation of an equivalent quantity of alkali by a large 4KI = excess of potassium iodide or sodium thiosulfate (HgO K*HgI, 2KOH). The method gives excellent results in favorable cases but has a number of drawbacks: I t is not applicable in the presence of a few interfering ions such as halide and phosphate. A precipitate of basic mercury(I1) salt is sometimes formed on addition of alkali or acetone, and this often takes an inordinately long time for complete dissolution in the aqueous acetone medium. The amount of potamium iodide or sodium thiosulfate added to liberate the alkali is required in large excess of the theoretical amount, and the amount of acetone required is much greater than the amount of mercury present. The present paper describes a n improved method in which all the above drawbacks have been almost completely eliminated, and i t has been possible to use mercury(I1) salts as primary standards in acidimetry.
+
+
THEORETICAL
The method depends on the fact t h a t any amide can prevent precipitation of mercury(I1) oxide when strong alkali is added t o a mercury(I1) salt. The amide can be present only in a very
small quantity-for example, a feFv milligrams of acetamide or formamide (0.5 mole of the former or 1 mole of the latter per mole of mercury(I1) oxide) are sufficient to prevent precipitation of mercury(I1) oxide from 10 ml. of 0.LV mercury(I1) salt, solution. Though any unsubstituted amide or imide would presumably be suitable, for various reasons acetamide or urea has been chosen for the present method, even though the latter is required in a comparatively large excess. The present method is similar t o t h a t of Das, except that a slight excess of acetamide or a large excess of urea (roughly 20 times the weight of mercury) is used in place of acetone to keep the mercury(I1) oxide in solution. Urea or acetamide is far superior to acetone, as not a trace of solid is precipitated by an excess of alkali and a bare excess of iodide or thiosulfate is required in the second st,age. The method gives excellent results when no interfering anions such as chloride or phosphate are present. I n the presence of chloride ions, which is rather common, the first end point is much less sharp than is usual in strong acidstrong base titrations owing to a progressive liberation of alkali by the chloride. Das has recommended using a mixed indicator which, though slightly better, falls short of the ideal. The difficulty, hovever, has been completely surmounted by conducting the first neutralization with alcoholic alkali in a 75% ethyl alcohol medium when the end point becomes very sharp. The second titration, after liberating t.he alkali by iodide or thiosulfate, is extremely sharp in all cases. GENERAL PROCEDURE
In Absence of Interfering Ions. To about 20 ml. of a solution containing about a millimole of mercury are added 4 to 5 grams (about 20 to 25 times the weight of mercury) of urea and a few drops of phenolphthalein indicator solution. Alternativelj-, 70 mg. of acetamide ( 5 ml. of a freshly prepared 1.4% solution) per millimole of mercury can be used in place of urea. The solution is now made alkaline by adding a slight excess of dilute sodium hydroxide solution. A dilute acid solution (roughly O . 1 N perchloric acid or nitric acid) is added drop by drop from a buret t o a sharp decolorization of the phenolphthalein. About 12 ml. of a neutral (roughly 0.5K) solution of potassium iodide or sodium thiosulfate (10 t o 20y0 in excess of the theoretical) is now added, when the solution again becomes vividly pink owing to liberation of alkali. The solution is titrated against a standard 0.1N acid t o an extremely sharp phenolphthalein end point. Hydrochloric acid is undesirable for the preliminary neutralization owing t o interference by the chloride ion. However, in the final titration with the standard acid all acids, including weak acids such as acetic acid, are suitable.
1332
ANALYTICAL CHEMISTRY
Some typical results are given in Table I for salts prepared from analytical grade mercury(I1 J o d e , as they are not available in a n analytically pure state. Mercury(I1) sulfate is available in analytically pure quality and can be estimated by the above procedure, provided it can be brought in solution. T o facilitate dissolving, 1 or 2 millimoles of mercury(I1) sulfate are treated with a slight excess of 0.1N alkali followed by 20 ml. of a 25y0 urea solution. On gentle warming all the mercury(I1) sulfate is dissolved and is ready for the usual analysis. Urea and acetamide have both been used in the proportiow mentioned above, and have been found to give equally satisfactory results. Potassium iodide and sodium thiosulfate function equally satisfactorily; potassiuni bromide can also be used with complete success, provided a n escess of about twice the weight of mercury present is used. At,tempts have been made t o dispense n-it,hiodide or thiosulfate by using a double indicator method. The accuracy attainable is very low, owing to a very grndiial change of pH, as ascertained by potentiometric tit'ration, if further acid is added after discharge of the phenolphthalein color. Direct Titration with Alkali. -1s the first end point with alkali is sharp, i t seems t h a t a direct titration with alkali is possible. I n fact, Fernandez and coworkers ( 2 ) have described such a method using acetone as the coniplesing agent. An investigation of this possibility was not made with the present method, a3 such a method would not have many applications-most mercury salts available are stoichiometrically inesact, and severnl of them require the addition of acid for their dissolution. However, a solution of mercury(I1) acetate obtained by dissolving the salt in water and filtering out the small amount of undissolved residue gives almost identical values by direct alkali titration ansl by acid titration as described previously.
of urea and 3 drops of 1% phenolphthalein solution are added t o the solution. T h e solution is now made slightly alkaline with alcoholic sodium hydroxide, and a dilute acid solution is added dropwise from a buret until the alkaline color just disappears. Ten milliliters of neutral (approximately0.W) potassium iodide or .uodium thiosulfate solution are added and the liberated alkali is titrated against a standard acid to a n estremely sharp phenolphthalein end point. Methanol can be used with equal success in place of ethyl alcohol in this titration. Sometimes a white precipitate forms on addition of urea t o solutions containing mercury(I1) salts. This is generally due to a n insufficiency of urea and can be rectified by adding more, Some typical results are s h o w in Table I1 which indicate the precision of the method.
Table 11.
Estimation of >lercury(LI) Chloride
Standard Acid Perchloric !Sa&Os
used)
knion
Erro, c,
.kcdate
0 1366 0 1794 0 1830 0.4816
n . 1368
0.22.56
n . 3360
n. 1794 0.1834 n 4817
Error,
% -0.11 +O.li 0.00
+o
n8
+0 15 0 no + 0 22 cn.n2
Acetic
Table III.
Estiniatioii of IIercur? ([I) and Mercury([) Sulfate Using Standard Nitric i c i d
Estimation of >Iercury(II) Salts Using Standard Nitric Acid IIgO, Grain Taken Found
Found, Gram 0 .ogna n.179n
Hydroclilnrii
Conipoiind IIercliry(I1) sulfate
Table?.
Taken, Grain 0 0909 0.1787 0 , 22.56 n 3331
lIercury(1) 3ulfate
Taken, Gram n 0906 0 1148 0 1181
n n
2042
n
1764 1798
n
9x1
n znzi
Foiind, Grain 0 ognij 0 1147 0 1189 o 2037 n. 1764 n.1801 0.2020 0.2162
Error,
c"c
+n
16 -0 n8 +0 17 -0 25 0.00
+o. 17 -n.nz +a n3
\-itrate
Perchlorate
o.n;ioi
n , 1234 0.3663
n
o.500
n . 1234 n. 2680
-0.21)
0.00 -0.11
For other salts such as sulfate and nitrate, the alkali titration values are slightly higher, presumably on-ing to some basic salt remaining undissolved on treatment rvjth wat,er. T h e discrepancy between the acid and the alkali titer of a mercury salt can be used as a measure of the stoichionietric inexactness of the qalt. Interfering Ions Present. I n the presence of other cations thc mercury can be separated quantitatively by methods similar to those outlined b y Das, except t h a t acetamide or urea is used iii place of acetone. The substitutions have the advantagc of avoiding the occasional precipitation of basic mercury( 11) snlt. Of the interfering anions-halide, thiocyanate, thiosulfate, cyanide, and phosphate-chloride ions are common, and the folloning modification works satisfactorily in their presence. The met'hod i.q illustrated by an analysis of mercury(I1) chloride which is available in the analytical reagent grade. About 0.2 gram of niercury(I1) chloride is dissolved in 20 nil. of a 1 to 3 mixture of water and alcohol (or rectified spirit I. After all the mercury salt is dispolved (1 or 2 minutes), 5 gram.
Phosphate can be removed easilv by carrying out the preliminary neutralization with barium hydioxide in place of sodium hydrovide until all the phosphate is precipitated. The filtrate c3n then be titrated as usual. This is a n improvement over the method used by Das which involves two steps: removal of the phosphate by a n excess of barium nitrate and subsequent removal of the latter as sulfate. The interference by chloride is much less than t h a t by bromide or iodide, and with a little practice fairly accurate values can be obtained by utilizing the normal procedure, particularly if a mixed indicator-phenolphthalein (3 parts) and l-naphtholphthalein (1 part)-is used as recommended by Das. INTERICTION B E T W E E S AMIDE AND MERCURY
That mercury(I1) oside is not precipitat,ed in the presence of a variety of substances including amides is well known and is referred to by Mellor ( 0 ) and Sidgwick ( I T ) . According t o Schoeller and Schrauth (IO), mercury(I1) oside dissolves in a solution of an amide such as acetamide to give (CHaCO NH)2 H g : imides, especially cyclic imides such as succinimide, react very easily. Ley and others (6-8) have shown t h a t the complex is on]>-slightly ionized and reacts only slowly with potassium iodide. Sidgwick (12) mentions t h a t the merc,ury-benzamide complex has the formula (CsHsCOSH)2H g and is so stable t h a t i t can he crystallized from hot potassium hydroside solution without decomposition. The formation of a urea-mercury(I1) nitrate com-
V O L U M E 27, NO. 7, J U L Y 1 9 5 5
1333
plex is also well known, and i t has the formula 0 H . H g S H C O S H Hg.NO3 ( 3 ) . Though the isolated compounds may have the above formulas, it has been observed t h a t 1 mole of acetamide can prevent precipitation of 2 moles of mercury and formamide is only roughly half as powerful, whereas dimethylformamide is practically devoid of a n y such power. Compounds of the types CHaCOS(1lgOH)z and HCOXHIlgOH are apparently formed, if only in solution. It is also possible to 11-rite six-membered cyclic donoracceptor strucntnres:
/O - - -Hg R.C
/O - --Hg\
‘SH-H’ I
?O
,CK,-CI
HSH-H,’
,o
0-Hg’
I1
grade mercury(I1) chloride, mercur? (11) sulfate, and mercurj ( I ) sulfate. Mercury(I1) ouide, which has been recommended as a primary acidimetric standard, n as used as a reference standard, tlie method having been checked by Kolthoff and van Berk ( 4 ) The result8 are shown in Table 11- nliich shows that mercury salts ale useful for standardizing both meak and strong acids
Table I\’. Standardization of Acids Using llercury(I1) Oxide, Mercury(I1) Chloride, JIercury(I1) Sulfate, and Mercury(1) Sulfate .4cld HgO, S HpC12, .Y IIgSOa, dV H F ~ Q OS ~, y-itric I’rrcliloric Hydrochloric -4cetic
0 09790
0.07626 0 07630 0,1030
0 09780
0 07623 0 07610 0 1030
0 09779 0.07608
0 09782
o.ioss
0 1030
0.07610
n 07619
Such formulas nierit serioris consideration in Tiem- of the unique properties of mercurj--e.g., niercuration, addition to tloable bond-as contrast,ed to any other metal. ESTIMATIOS OF 1\IERCCRY(I) S.ILTS
lIercury(1) salt can be e h n i a t e d by the present method because on addition of alkali and amide t o a mercury(1j salt in soliition as well as in the solid state, niercury(1) disproportionates into metallic mercury and niercury(I1) ions. T h e black precipitated mercury can be redissolved rind analyzed hj- tlie present procedure or the mercury(I1) ion remaining in solution can be thus analyzed. The above procedure has been tested by using analytically pure Inercury(1) sulfate. T h e latter is dissolved by treating 1 niillimole n-it,h a slight excess of 0.lA- sodium hydroxide solution followed b y 10 nil. of 25% urea solution. Tlie solution is \t-ariiied for a few Ininutes to complete the reaction. The black precipitate of mercury which sett,les down is filtered off and washed out, and the filtrate is analyzed volumetrically for the iiiercury(I1) salt. T h e results are shown in Table 111. Mercury(1) chloride is t h o dissolved by tlie same method, b u t the filtrate contains chloride ion and so a n acid titration by the usual niethod using the mixed iiiclicator in place of phenolphthalein is possible b u t yields sonien.hat less accurate results. Tlie residue of mercury can be dissolved and titrated, but t h a t method hap not, been tested. AIERCVRY SALTS A S PRIMARY ST.4VDARDS I S .ICIDIJIETRY AND ALKALIMETRY
From the high degree of accuracy attainable with the present method of estimating mercury it is evident t h a t if mercury Falts of the analytical reagent grade which are stable and stoichiometrically exact could be obtained, they could be utilized for standardization of both acids and alkalies. This method, unlike t h a t involving the use of carbonates and borax ( Ka?B40i.10H20), would require few special precautions. For standardizing strong alkali a known amount of the niercurp salt in the absence of interfering ions could be titrated against the alkali using pheiiolphthalein as indicator in the presence of urea or acetamide. .4lcoholic alkali could be conveniently standardized using mercury(I1j chloride. For etnndardizing a n acid, the alkali liberated by the addition of a slight excess of iodide or thiosulfate during the second stage of this method of est,imating mercury could be titrated. Unfortunately stoichiometricallj- exact and analytically pure mercury salts are not available except mercury(I1) chloride, niercury(I1) sulfate, and niercury(1) sulfate. Mercury(11) chloride could be used for standardizing an acid by the special method of using alcoholic medium as already described. Mercury(I1) sulfate and mercriry(1) sulfate are satisfactory for the purpose, as they could he used by following the usual procedure after dissolving them b\- methods already described. T o test the theory, four acid solutions-nitric, perchloric. hydrochloric, and acetic-were standardized against analytical
Clilwide ions interfere n-it!] tlie first neutralizat,ion brit do not affect the final titration Ivitli acid; thus hydrochloric acid can also be standardized by this method. T h e most convenieiit method, however, for routine laboratory purpose is to use a stock solut,ion of the easily dissolrecl mercury(I1) acetate and use i t as the routine alkali solution of knon-n strength after standardizing it against a standard acid or potassiuni biphthalnte solution. Such stock solutions seem to be indefinitely stable when kept in the dark. Use of a solution of sodium hydroxide which is free of carbonate a i d is stored in a ‘keasoiied” bottle (5) would no doubt be satisfactory, but ae no such special precautions need be taken with mercury salts the latter maj. be used with greater ease. LITERATURE CITED
Das, 11. S . ,Ax.4~.CHEK, 25, 1406 (1953). Fernandez, J. B., Snider, L. T., and Rietz, E. G., I h i d . , 23, 899 (1951). Iiarrer, P., “Organic Chemistry,” 4th ed., p. 232, Elsevier, Sew Tork, 1950. Kolthoff, I. hl., and Berk, L. H. v a n , 2. anal. Cheni., 71, 339 (1927). Iiunzler, J. E., ASIL. C H E M . , 25, 93 (1953). Leg, H., and Kissel, H., BFT.,32, 1357 (1899). Ley, H., and Schaefer, IC., Ibid., 35, 1309 (1902). Ley, H., and Schaefer, K., 2. p h y s i k . Chem., 28, 385 (1899). YIe!lor, J. W., “Comprehensire Treatise on Inorganic and Theoretical Chemistry,” Vo!. IT’, p. 784, Longmans, Green, London, 1923. Schoeller, W., and Schrauth, V., Ber.. 42, 784 (1909). Sidgwick, N. V.. “CheniiLa! Elements and Their Compounds,” vol. I, D. 319, Clarendon Press, Oxford, 1950 ( 1 2 ) Sidgwick,-K. V,, “Organic Chemistry of Nitrogen,” new ed., p. 142, Clarendon Press, Oxford, 1945
