Table
VIII.
Analytical Precision Procedure on Alloys
Sample No
6 7 8 9 10
W.A. 68-10
of
v, % 4.00, 3.96, 4.92, 4.67. 3.41; 3.65. 4.18; 4.91, 4.57, 4.19, 3.96,
3.99 3.95 4.90 4.64 3.41 3.65 4.19 4.91 4.56 4.17 3.98
of one set of standards remained unchanged after 3 months. CONCLUSION
This procedure as outlined is useful as an analytical control method in tthe
production of pure titanium tetrachloride. It is rapid and no special equipment is necessary. KO solutions have to be made up and standardized, and errors accompanying their use are eliminated. The accuracy is sufficient for control purposes. Higher accuracy may be obtained by accurately weighing the sample in a glass ampoule. The range may be extended downward by using a larger sample and I-em. cells. I n the case of alloying amounts of vanadium, this procedure has a precision and accuracy compaiable m-ith volumetric methods. It coreis the range from 0.0 to 8% nuiadium. Technicians can run analyses TTith a minimum of training. A few simple precautions prevent the fading of the color caused by nioisture. The presence of chromium, which causes the only serious interference, is readily noticeable because of greenish
color. As the color is stable, time of measurement of absorbance is not critical. ACKNOWLEDGMENT
The authors wish to express their appreciation to Joe Mariner and John Reese for their cooperation in obtaining data. LITERATURE CITED
(1) Hillebrand, W. F., Lundell, G. E.,
“Applied Inorganic Analysis,” pp.
359-63, Wiley, New York, 1929.
(2) Sandell, E. B., “Colorimetric Determinations of Trace Metals,” p. 443, Interscience, Sew York, 1944. (3) Snell, D. S.,Snell, C . T , ‘Cplori-
metric Methods of Analysis, 3rd ed., vol. 11, p. 455, Van Nostrand, New York, 1949.
RLCEIVED for review July 11, 1956. Accepted November 12, 1956.
Amperometric Titration of Mercury(l1) with Tetraphenylarsonium Chloride OSCAR MENIS, ROBERT G. BALL, and
D. L. MANNING
Analytical Chemistry Division, Oak Ridge National laboratory, Oak Ridge, Tenn.
b The amperometric titration of mercury(l1) with tetraphenylarsonium chloride is carried out in a 0.4M nitric acid medium a t fixed potentials in the region of 0 to -0.5 volt vs. the saturated calomel electrode. The mercuric and tetraphenylarsonium ions react in an equimolar ratio to form a precipitate which has been determined by chemical analysis to be (CsH&AsHgCl2NOa. The reaction is believed to be essen2 CItially H g + + f (CsHj)dAs+ NO, -+ (C~H6)4AsHgC12N03. The substances that interfere can b e divided into three general classes, depending on whether the ion is reduced a t the titration potential and reacts with tetraphenylarsonium chloride. The degree of interference of each class has been determined. By this amperometric procedure, 0.025 to 2.30 mg. of mercury can b e titrated with a coefficient of variation of 3%.
+
+
was needed for the determination of mercury in either niicrogram or milligram quantities, in solutions of uranyl sulfate. The three most commonly used colorimetric methods are the dithizone, the dinaphtholthiocarbazone, and the diphenylMETHOD
carbazone methods (&?‘), but all require the extraction of the iiiercury reagent complex into an organic phase such as chloroform before nieasurement of absorbance. As these methods are subject to many interferenrcs, particularly from copper, procedures ( 6 , 7 ) by which mercury can be separated are necessary. These are rather involved. do not appear to tic casily adaptable to a general procedure for the dctermination of mercury, and did not seem suitable for the need of this laboratory. The polarographic deterininutiori of mercury, as revien ed by Kolthoff, did not appcar applicable ( 3 ) In a polarographic method, both nirrcurous and nicrcuric ions j-icltl \rcll-dcfined diffusion currents. As reduction of the mercuric ion in a nitrate medium starts at an applied potential of zero, i t would be difficult to measure the diffusion current, in view of the uncertainty of the residual current encountercd in different sample types. For this particular application. anipcrometric titration appeared to offer advantages over colorimetric or polarographic techniques. Because of the interference problem inherent in colorimetric procedures (5-?‘), an anipero-
metric method could be made more selective, would yield more precise results than direct polarographic measurements, and would not suffer from limitations such as reproducibility of capillary characteristics and rcversibility of the system. Ammonium thiocyanate was first investigated as a titrant for the amperometric titration of mercury. When 0.5 to 2 mg. of mercury in a IO-ml. rolume i n a nitric acid medium was titrated x i t h 0.01-Y aninionium thiocyanate a t -0.2 volt with a rotating platinum electrode us. the saturated calomel electrode (S.C.E.), the resulting titration curve revealed two end points, which corresponded closely to the formation of Hg(CiYS)- and Hg(CI\;S)2. The reactions m r e not stoichiometric in this concentration range; therefore, ammonium thiocyanate was not considered a suitable titrant for this determination. I n subsequent studies, tetraphenylarsoniuni chloride proved satisfactory for the amperoiiictric titration of small quantities of mercury. Willard and Smith (6) used tetraphenylarsoiiium chloride (TPAC) for the deterniination of relatively large quantities of mercury. They added a n escess of the reagent, precipitated the VOL. 29, NO. 2, FEBRUARY 1957
245
mercury as [(C&,)~4s]~HgCla, and determined excess reagent potentiometric titration with a standard solution of iodine. Lamprey (4) reported a conductometric determination based on the reaction used by Willard and Smith, but did not indicate its adaptability and limitations. The use of tetraphenylarsonium chloride for the amperometric titration of stannic ions and mercuric ions has been reported by Kolthoff and Johnson (2). They titrated from 14 to 35 mg. of tin and 30 to 80 mg. of mercury in a 3N hydrochloric acid medium a t an applied voltage of -0.3 us. S.E.C., and precipitated mercury as [(CGHS)&]2HgC14. They did not extend the method to the titration of smaller quantities of these ions. I n applying this method to the determination of 0.025 to 2.5 mg. of mercury it was found that optimum conditions for the amperometric titration differed considerably from those estahlished by Kolthoff and Johnson. Development of this procedure involved a study to determine optimum conditions relating to the supporting electrolyte, the applied potential a t which to conduct the titration, and the composition of the reaction product a t the end point. The precision and reliability of the new procedure were also established. REAGENTS
Mercuric nitrate solution, 2 mg. of mercury per ml. Dissolve 2.000 grams of reagent grade, triple-distilled mercury in 50 ml. of 81M nitric acid. Dilute to 1 liter with distilled Fvater. Tetraphenylarsonium chloride solution, 0.01M. Dissolve 5.5 grams of technical grade tetraphenylarsonium chloride, (C6Ha)&C1.HC1, in 200 ml. of water; then dilute to I liter. Standardize this solution by a potentiometric titration with a standard solution of iodine ( 8 ) .
sodium chloride media are shown in Tables I and 11. The precision in the amperometric titration of mercury is better in the nitrate medium than in the chloride solution. A coefficient of variation of 3% was obtained when the titrations were carried out in a 0.4M nitric acid supporting electrolyte over a range of 0.025 to 2.5 nig. of mercury. This was independent of the titration potential within the range of 0 to -0.6 volt us. S.C.E. A coefficient of variation of 11% was obtained, however, when the titrations were conducted in a 2M sodium chloride medium. The l o i ~ e rprecision for the titrations in the chloride medium is due, in part, to the relatively high and variable residual current which causes the measured current (diffusion minus residual current) t o be small and erratic. As a result, the lower limit of mercury that could be titrated in a chloride medium was about 0.5 mg. I n nitric medium, however, the residual current is both small and constant. Consequently, the lower limit of mercury was extended to about 0.02 mg. For this reason, a 0.4N nitric acid medium is recommended for use in the amperometric titration.
an aliquot of the test solution that contains 0.02 to 3 mg. of mercury, and add 0.4M nitric acid until the total volume is 15 ml. Adjust the applied voltage to -0.1 volt us. S.C.E. Add increments of titrant ranging from 0.01 to 0.1 ml., depending upon the amount of mercury being titrated. Stir the solution for 1 minute after the addition of each increment; then, after the turbulence in the solution has ceased, measure the diffusion current. Plot volume of titrant against current. Determine the end point of the titration by the extrapolation method, as illustrated in Figure 1. RESULTS
The results from the amperometric titration of mercury with tetraphenylarsonium chloride in nitric acid and
Table 1. Amperometric Titration of Mercury(1l) with Tetraphenylarsonium Chloride
Conditions -0 5 Potential us. S.C.E., volt Medium, HNOa, S 0 4 Titrant, (CaHb)4AsC1 HC1 (TPAC), M 0.0119
EXPERIMENTAL
968 794 589 497 0 386 0 196 0.0992 0.0957' 0,0510 0.0315 0.0103
0 0 0 0
2 30 1 84 1 38 1 15 0 920 0 460 0.230 0.248 0 124
2.31 1 90 1 41 1 19 0 921 0 468 0.237 0.236 0.126
0.074 0.078 0.025 0.025
1.oo 1 03 1 02 1 03 1 00 1 02 1.03 0.95 1.01 1.05 102
Selection of Applied Potential. The polarography of mercury(I1) and of the tetraphenylarsonium ion in nitrate and chloride media was investigated to determine the characteristics of their current-voltage curves. I n a nitric acid medium, the diffusion current of mercury was essentially parallel, with the voltage axis between +0.1 and -0.8 volt us. S.C.E. I n a sodium chloride medium the diffusion current, which was dependent upon the concentration of chloride, started to rise a t or near a zero applied potential and become parallel Fyith the voltage avis between -0.1 and -0.8 volt us.
Xoendpoint 0.012 . . . .. X = 1.01 moles TPhC/mole Hg
s v
= *00.026 = k3yO
Samples 8 throiigh 12 titrated with tetraphenylarsonium chloride solution.
0.0123M
APPARATUS
Buret. Micro-Metric syringe type, fitted with a 5-ml. Yale hypodermic syringe 5-Y, available from the MicroMetric Instrument Co., Cleveland, Ohio. One scale division is equivalent to 0.00267 ml. Polarograph. ORNL high sensitivity, Model Q 1160 ( I ) , equipped with dropping mercury electrode and saturated calomel reference electrode. Stirrer. Thomas magnetic stirrer and Teflon-covered stirring bar. Titration Cell. A 20-ml. glass cell fitted with a ground-glass ball joint was connected t o the electrode and buret assembly. It was connected to the saturated calomel electrode with a potassium nitrate-agar salt bridge.
**\\
I
l
l
&*%?!i*-*=*=*=.
PROCEDURE
Transfer into the polarographic cell
246
ANALYTICAL CHEMISTRY
Figure 1.
Amperometric titration of mercury
-
S.C.E. The tetraphenylarsonium ion did not exhibit a diffusion current in either nitrate or chloride medium. The amperometric titration of mercury can, therefore, be carried out at any potential within these limits. Effect of Deaeration. T h e deaeration step can be eliminated when milligram quantities of mercury are titrated at a n applied potential of -0.1 volt. Under these conditions t h e contribution of t h e diffusion current of oxygen t o t h e residual current is negligible. When microgram quantities of mercury were titrated, however, t h e method was more reliable when t h e residual current was kept a t a minimum through a deaeration step. Effect of Chloride. The effect of t h e chloride concentration in t h e supporting electrolyte on t h e stoichiome t r y of t h e reaction was investigated. Initially the titration was conducted in a 2M sodium chloride medium similar t o the one used by Willard and Smith for quantitative precipitation of mercury. Table I1 s h o w that the mole ratio of tetraphenylarsonium chloride to mercury at the end point was 1 t o 1. Additional tests in the absence of a chloride supporting medium (Table I) revealed that the mole ratio of tetraphenylarsonium chloride to mercury in the reaction product was also 1 to 1.
Table II. Amperometric Titration of Mercury in Sodium Chloride Medium Gonditions Potential cs. S.C.E.. volt -0 5 %Tedium,NaC1, JI ’ 2 Titrant, (C6Hj)4AAsC1 €IC1 (TPXC), W 0 0119 Mole TPAC, ~Ierc~b Yk, . Ratio, 111. Added Found TPAC/Hg 1.040 2.30 2.48 1.08 0.960 2.30 2.29 1.00 1.10 0.G71 1.38 1.60 0 534 1.15 1.27 1.11 0 436 1.15 1.04 0.91 0 387 0 92 Q 92 1 00 0 336 0 92 0 80 0 87 0 310 0 69 0 74 1 08 0 235 0 46 0 56 1 22 N o end point 0.23 X
s v
= = =
1.05 moles TPAC/mole Hg 2ro.11 2r11yo
Further additions of titrant in excess of that required for a 2 t o 1 mole ratio of reagent to mercury showed no evidence of a second end point. Apparently, the composition of the product a t the end point is independent of the chloride ion concentration of the supporting electrolyte up to 2111. In order to determine the effect of larger concentrations of chloride, as
Tabie 111.
Mole Ratio of Reactants a s Function of Chloride Concentration
Medium Electrolyte Molarity 2 SaCl 2.5 3
HC1
3
Mo!e Ratio, TPSC/Hg
Reactants, Mmole TPAC Mer cur y 0.0112 0,0118 0.0258 0,0124 0.0258 0.0124 0.0127 0.0062 0.505 0.263 S o end point 0.124
the medium for titrating mercury, additional tests were carried out under essentially the conditions employed b y Kolthoff and Johnson ( 2 ) . Larger quantities of mercury (50 mg.) reacted in a mole ratio of 2 moles of tetraphenylarsonium ion to 1 mole of mercury, in agreement \yith their findings. However, for lower concentrations (2.5 mg. of mercury) the end point could not be detected in a 3M hydrochloric acid medium. When a 3M sodium chloride medium was substituted, the mercury reacted with tetraphenylarsonium chloride in a 2 to 1 mole ratio. These tests indicate that depending upon the chloride concentration of the medium small quantities of mercury can react with tetraphenylarsonium chloride in either a 1 to 1 or 2 to 1 mole ratio (TPAC/Hg). iinalysis of the precipitate obtained by reaction of equimolar quantities of niercuric and tetraphenylarsonium ions in 0.4M nitric acid revealed that i t was composed of 54% tetraphenylarsonium ion, 28y0mercuric ion, 11% chloride, and 8% nitrate, corresponding to the formula (C6H6)&HgCl2YO3. Therefore the reaction followed by the indicator electrode during the titration ib apparently:
iiithough the compound is soluble under the conditions of this titration, i t is apparently not dissociated, as the break in the titration cwve a t the end point is relatively sharp. Lorn yields of the order of 10 to 50% were obtained in gathering the precipitate for analysis. T h a t the chloride ion plays a role in the reaction was also established. K h e n titrations were carried out with a tetraphenylarsonium reagent prepared in the nitrate form, no reaction was noted. The tIyo chloride ions present in the reagent are therefore essential for this reaction. Effect of Nitric Acid. As a nitric acid medium was used in this titration, t h e effect of its concentration as supporting electrolyte mas studied. The titration of 10-5 mole of mercury in the presence of 0.02 to 1.531 nitric acid
Table IV.
0.95
2.08 2.08 2.01 1.92 ..
Effect of Interfering Ions
Tolerance Potential within of 2% Error Interference Titration, Mmoles/ Class Ion V us. S.C.E. Mmole H g 1 Cd++ -0.3 5 2
uo*++
3
Sn(I\’)
0.0
-0.4 -0.3
90
10 5
yielded results that were independent of the acid variation. I n concentrations above 1.5M nitric acid, dissolution of the falling mercury drop caused increasingly positive errors in the mercury titration. At a nitric acid concentration of 2 X , mercury dissolved at such a fast rate that no end point in the titration was found. Interfering Ions. Previous authors (8, 9) have listed the cations and anions that are known to react with tetraphenylarsonium chloride: Anions include permanganate, chromate, molybdate, periodate, perchlorate, iodide, bromide, and thiocyanate. Cations that form halide complexes and, consequently, precipitate with tetraphenylarsonium chloride include bismuth, tin, titanium, zinc, cadmium, and iron. I n general, the ions that interfere with the amperometric titration of mercury can be divided into three classes. 1. Ions that produce no diffusion current at the potential of the titration, but react Kith tetraphenylarsoiiium chloride. Examples: perchlorate, nitrate in very high concentrations, and cadmium, the last at a potential below -0.G volt us. S.C.E. 2. Ions that exhibit a diffusion current at the potential of the titration (potential more negative than zero applied voltage) but do not react with tetraphenylarsonium chloride. Examples: uranyl ion and dissolved oxygen. 3. Ions that exhibit a diffusion current under the conditions of class 2 and also precipitate with tetraphenylarsonium chloride. Tin is a typical example.
The degree of interference of a typical ion from each class was stcdied (Table
IV) * Cadmium(I1) presents a possible source of error in the determination of VOL. 29, NO. 2, FEBRUARY 1957
247
mercury by reacting with the tetraphenylarsonium chloride titrant. However, as mercury forms a more insoluble precipitate with tetraphenylarsonium chloride than does cadmium, its solubility product is exceeded first and consequently the cadmium does not interfere when as much as 5 mmoles of cadmium is present per millimole of mercury On the other hand, larger amounts of cadmium repress the solubility of the cadmium-tetraphenylarsonium chloride preciDitate, by virtue of the common ion effect, t o the point where both cadmium and mercury precipitate upon addition of the tetraphenylarsonium chloride titrant. The uranyl ion typifies the second type of interference. Elimination of interference is largely dependent upon the ability of the instrument to suppress the diffusion current of the nonreacting ion and yet measure the change in the diffusion current due to the removal of the ion being titrated. With the polarograph used in this study, the current due to 0.1 mmole of uranium can be suppressed in the titration of 0.01 mmole of mercury at a potential of -0.4 volt. At an applied voltage
of zero, 0.9 mmole of uranyl ion per 0.01 mrnole of mercury can be tolerated. A t this potential, the uranyl ion does not exhibit a diffusion current; therefore, a large quantity can be tolerated. In the presence of more than 0.9 mmole of uranium per 0.01 mmole of mercury, the titration of the mercury with tetraphenylarsonium chloride is erratic. The effect of stannic ion was studied as representative of the third class of interfering ions, which reacts with the titrant and also produces a diffusion current a t the titration potential. A mole ratio of stannic ion to mercury of 5 to 1 can he tolerated without effect. A t higher mole ratios, the error in the mercury determination is pronounced. The interfering ions that were tested form more insoluble tetraphenylarsonium chloride precipitates than other ions of that class. Therefore, the interference from other ions will probably be of the same order or less than that of the ion tested. ACKNOWLEDGMENT
The authors acknowledge the assistance of AI. A. M a r k in the preparation of this report.
LITERATURE CITED
Kelley, M. T., bliller, H. H., ANAI.. CHEM.24, 1895 (1952j Kolthoff. I. M.. Johnson, R. A., J Electrochem. Soc. 98, 231 (1951). Kolthoff, I. &I Lingane, .,,, J. J., “Polarography, 2nd ed., vol. 11, pp. 511, 577, Interscience, New York, 1952. (4) Lamprey. H., thesis, University of Michigan, 1935. (5) Rudden, C. J., “Analytical Chemistry of the Manhattan Project,” let ed., pp. 399-402, NcGraw-Hill New York. 19.50. (6) Sandell, E. B., ‘ Colorimetric Determination of Traces of Metals,” 2nd ed., pp. 441-52, Interscience, New York, 1950. (7) Snell, F. D., Snt41, C. T., “Colorimetric Methods of Analvsis.” 3rd ed., pp. 63-78, Van Xostcand, S e w York, 1953. (8) Willard, H. H., Smith, G. SI.,ISD. ENG. CHEK., ANAL.ED. 11. 186 (1939). (9) Ibid., p 269. RECEIVEDfor review August 27, 1956. Accepted November 17, 1956. Division of Analytical Chemistry, 130th hleeting, ACS, Atlantic City, N. J., Se tember 1956. Work carried out under Eontract KO. W-7405-eng-26 at Oak Ridge National Laboratory, operated by Union Carbidc Nuclear Co., a division of Union Carbide and Carbon Corp., for the Atomic Energy Commission.
Precipitation of Zinc Phosphates from Solutions of Sodium Ortho-, Pyro-, and Triphosphate 0.T. QUIMBY
and
H. W. McCUNE
Miami Valley laboratories, The Procfer and Gamble Co., Cincinnati 3 I , Ohio
,This study, undertaken to evaluate zinc precipitation methods of determining pyrophosphate and triphosphate, involved precipitation of zinc orthophosphate, pyrophosphate, and triphosphate, alone and in mixtures; coprecipitation of triphosphate with pyrophosphate at pH 3.8 was determined with radio-tagged phosphates. The triphosphate interference with pyrophosphate precipitation in the Bell method often causes the results for both species to be significantly in error; nevertheless, it may be possible to modify this method to yield trustworthy rapid determination of pyrophosphate in such samples as commercial triphosphate. Increasing the pH from 3.8 to near 5 will bring about nearly complete precipitation of orthophosphate, pyrophosphate, and triphosphate, but gives no promise of
248
ANALYTICAL CHEMISTRY
providing a trustworthy indirect method for one of the three species in their mixtures. Only in special casese.g., intermediate pH’s and low orthophosphate level-will it be practical to precipitate all of the pyrophosphate and triphosphate without contamination by orthophosphate. A new phase Zn?PzOj.3Hz0 was encountered.
S
published methods for the determination of triphosphate and pyrophosphate involve the use of zinc ions for the precipitation of one or both of these phosphates (9, 4, 6, 7 , 16, 19, 66, 27, SI). I n some of the supporting experiments on which the methods are based, and in the studies of condensed phosphate precipitations by other metal ions such as manganese(I1) ( I Y ) , and EVERAL
tris(ethy1enediamine) cobalt(111) (61), there is evidence that pyrophosphate and triphosphate are harder to separate from each other than orthophosphate or trimetaphosphate. Recent work ( $ I ) , as well as the present study, indicates that pyrophosphate and triphosphate usually coprecipitate. Furthermore, large amounts of pyrophosphate inhibit the precipitation of triphosphate and vice versa. It is therefore pertinent to examine the basis for these methods. Common ring phosphates (trimeta and tetrameta) were not studied because their zinc salts are soluble in water. EXPERIMENTAL
Preparation of tagged and inactive sodium phosphates has been described
