caused by triiodide formation (1). Potentials calculated assuming that iodine is the product of oxidation agree well with those calculated from the data of Latimer ( 7 ) . The oxidation of bromide or reduction of bromine in 0.05M sulfuric acid is similar to the iodide-iodine case except that, as found by Lingane and Anson (8), the reactions are much less reversible. The reactions a t stationary electrodes do give straight lines on plots corresponding to Equations 8 and 9 and the slopes are reasonable. However, a definite discrepancy exists betnTeen formal potentials calculated for oxidation and those calculated for reduction, again most noticeable a t short electrolysis times. I n the cases discussed here, the last
term in Equations 8 and 9 was never greater than 1 mv.; thus the convergent nature of the diffusion makes very little difference. If the logarithmic charge term is not made dimensionless as mentioned above, other terms appear in the charge-potential relationship in cases where 5 # y. These terms contain D, TO, and t to a more important extent and make it necessary to know these parameters more closely. Making the logarithmic charge term dimensionless has the disadvantage of introducing a concentration term, but this datum is . often readily available.
(2) Booman, G. L., Morgan, E., Crittenden, A. L., J . Am. Chem. SOC.78, 5533 (1956). (3) Carslaw, H. S., Jaeger, J. C., “Conduction of Heat in Solids,” p. 280, Oxford Press, London, 1948. (4) Delahay, P., [(New Instrumental
Methods in Electrochemistry,” Chap.
3, Interscience, New York, 1954. (5) Hush, N. S., 2. Elektrochem. 61, 738 (1957). \ - - - . I -
(6) Kolthoff, I. M., Tomsicek, W. J., J . Phys. Chem. 39,945 (1935). (7) Latimer, W. M., “Oxidation Poten-
tials,
2nd ed., Prentice-Hall, New
York, 1952. (8) Lingane, J. J., Anson, F. C., ANAL. CHEM.28,1871 (1956). (9) Stackelberg, M. von, Pilgram, M., Toome, V., 2. Elektrochem. 57, 342 (1953).
LITERATURE CITED
(1) Beilby, A. L., Crittenden, A. L., J . Phys. Chem. 64, 177 (1960).
RECEIVED for review November 9, 1959. Accepted March 7, 1960.
Si muIta neo us PoIa rogra phic Dete r mina ti o n of Lead a nd Azide Ions of Lead Azides in Aqueous Media JAMES 1. BRYANT and MARYLAND D.
KEMP
0.S. Army Engineer Research and Development laboratories, Fori Belvoir, Va. b A polarographic method has been developed for the simultaneous determination of lead and azide ions of lead azide in aqueous solutions. Its sensitivity allows safe low concentrations of lead and azide ions to b e effectively determined in a fraction of the time required by classical analytical methods. The polarographic method was standardized b y gravimetric determinations.
A
in the study of the properties of azides is the lack of an accurate and convenient method for the analysis of solutions of the more unstable salts and hydrazoic acid. Gravimetric and volumetric methods are time-consuming and generally dangerous, and often give low results (1). Gas volumetric methods do not give accurate results when very dilute solutions or very small samples, which are required for purposes of safety, are to be analyzed. Haul and Scholz (3) made polarographic studies of the azide ion of lead azide and reported the wave height to be concentrationdependent from 4 X to 10-3N. Later Vasiliev (7) determined the lead content of lead azide polarographically and calculated the azide concentration, assuming a stoichiometric relationship. The purpose of this study was to develop an effective and convenient MAJOR DIFFICULTY
758
ANALYTICAL CHEMISTRY
method for routine determination of lead and azide content of lead azide simultaneously with the polarograph. The purity of samples used was checked by gravimetric determinations. I n 1934 Revenda (6) showed that in the presence of depolarizers which have a great affinity for mercury ions, anodic depolarization at the dropping mercury electrode gives well defined polarographic waves. Later Kolthoff and Miller (4) extended this investigation in a more analytical manner and reported that within certain concentration ranges the polarographic step was concentration-dependent. These investigators worked with chloride, bromide, and iodide solutions among others, and found that upon the electro-oxidation of mercury in the presence of these anions the metal formed a difficultly soluble mercurous salt a t the surface of the drop. Because of their halogenoid character, similar electrode reactions should be expected for the azides. This expected similarity has been found to exist. Thus, if this property of an electro-oxidized dropping mercury electrode is utilized here, the mercury goes into solution as mercurous ion, making possible the formation of a precipitate of difficultly soluble mercurous azide adsorbed on the surface of the mercury drops. Because of the resultant depolarization, an anodic current develops and its magnitude is proportional to the
rate of diffusion of the azide ion to the electrode surface, which in turn is dependent upon the difference in the azide ion concentration in the bulk solution and a t the mercury drop solution interface. The limiting current is reached when the azide ion concentration a t the mercury drop surface becomes effectively zero. Then id =
KCNS-
(1)
where C N ~ is- the concentration of azide ion in the bulk solution. The equations for the electrode reaction for the azide ion are Hg+Hg+ Hg”
+ ?Js-
+
E
+ HgN;
(2)
(3)
and the net electrode reaction is Hg
+ N a - 4 HgNa +
(4)
The cathodic reduction of the lead ion is of general familiarity and is not discussed here. EXPERIMENTAL
A 10-3N stock solution of lead azide was prepared by dissolving 0.2914 gram of lead azide in O.1N potassium nitrate and diluting to 2 liters. The lead azide used was precipitated by mixing equal volumes of soIutions of 2N doubly recrystallized sodium azide and 1N certified lead nitrate. The precipitate was successively washed with distilled water, alcohol, and ether and
I
i /
ANODIC
I
I
+OAO t0.30 +020 1010
CATHODIC
I
I
I
I
I
I
I
-
000 -010 - 0 2 0 .OM -040 - O M -060 - 0 7 0 E M F vs. S C E , VOLTS
Figure 1. Continuous polarogram of lead and azide ions azide in 0.1 N potassium nitrate
of lead
finally dried in a vacuum desiccator for 24 hours. Smaller concentrations of lead azide down to 2 X 10-4N were prepared by appropriate dilutions of the stock solution. Because of the unstable nature of crystalline lead azide and its sensitivity to light, all operations were carried out on fractional gram quantities and in the absence of actinic light. A Leeds & Northrup Electro-Chemograph Type E was used to obtain polarograms of the solutions. The reference electrode was a Hg-Hg2S04 electrode acidified with sulfuric acid. The use of a saturated calomel electrode was discontinued when distorted azide waves, believed due t o interference of chloride ion, were obtained. The results obtained with Hg-Hg2SOa electrode were computed in terms of the calomel electrode for convenience. Electroanalysis v a s carried out in a conventional H-cell which was thermostated a t 25' k 0.25" C. The dropping mercury electrode mas of the usual type and its drop time was determined at the halfwave potential of each ion investigated. Before analysis, solutions were deaerated with nitrogen and during electroanalysis maintained under a hydrogen atmosphere. Because the time of electrolysis was short (both ions approximately 6 minutes), the solutions were not buffered, and no measurable change of p H could be detected during electrolysis. When the solutions were allowed to stand in the cell longer than 30 minutes, measurable decreases in pH could be detected. This change of pH, due to diffusion of hydrogen ions from the reference electrode, caused the formation of hydrazoic acid, which on volatilization decreased the azide ion concentration of the solution. Therefore, in all cases it was necessary to carry out the electrolysis immediately after placing the solution in the cell before effective diffusion could take place. Good azide waves could be obtained without deaeration and maxima suppressor; however, to obtain reproducible lead waves of good form, use of maxima suppressor, rigid deaeration, and maintenance of hydrogen atmosphere during electrolysis were suppressor, required. As maxima Triton X-100 was added to the solutions
to give a cell concentration of 0.002%. Because improper disposal of azide salts can be dangerous, a vessel containing saturated ceric ammonium sulfate solution was attached to the suction apparatus used to empty the cell to receive the solution after electroanalysis. Thus the azide ion was immediately destroyed and the mercury was recovered periodically for reuse. Determinations were made of solutionsof Pb(NJ2 from to 2 X 10-4N. To obtain information concerning the stability of standard solutions, a timeconcentration study was made using two solutions stored under different conditions. The first (Solution A) was a t all times protected from actinic light and large temperature variations. The second (Solution B) was stored in a conventional manner and as a result a t times exposed to strong artificial and sun light and possible temperature variation of 15" to 20". Gravimetric Determination of Lead and Azide Ions. The method for the azide ion was substantially the same as that revien-ed by Dennis and Isham ( 2 ) . Two 50-ml. aliquots were taken from the 10-0.47 lead azide stock solution and the silver azide was precipitated by adding 1.0 ml. of 5% silver nitrate to each aliquot. The solutions were allowed to stand overnight in a dark room and the precipitated silver azide was collected on a Gooch crucible and washed freely with distilled water. (It is essential that the water be free of nitric acid, because silver azide dissolves in dilute nitric acid.) The precipitate was dried a t 100' C., ignited a t 240' C. (AgN3 explodes a t 273' C.), and weighed as AgN3. The lead ion was determined using the original freshly precipitated lead azide salt. The salt was dissolved in 5% acetic acid and after the addition of 2 ml. of sulfuric acid heated t o sulfur trioxide fumes, and precipitated as sulfate. The precipitate was dried a t 100" C a n d i g n i t e d a t 600" C. (6). Asin the polarographic determinations, because of the unstable nature of lead azide, all operations were carried out in the absence of actinic light and additionally behind explosion-proof shields.
L
r040
E M F
+020
-030
vs.
t0lO
S C E ,
000
VOLTS
Figure 2. Method of measuring azide diffusion current RESULTS AND DISCUSSION
The results of the gravimetric determinations are shown in Table I.
Table 1.
Gravimetric Determinations
of Azide and Lead Ions AgNa Formed, Gram Theor. Exptl. 0.0075 0.0075
Sample Wt., Gram 0.3014 0,1467 0.1480
0.0075 0.0076
Differ-
ence, Gram
0,0000 0,0001
PbSOa Formed,
Gram Exptl.
Theor. 0.3138 0.1627 0,1541
0.3133 0.1523 0.1538
0.0006 0.0004 0.0003
Figure 1 shows a continuous polarogram of azide and lead ions. Contrary t o the double chloride ( 4 ) and azide (3) waves previously reported, the waves observed in these studies showed only a single diffusion step. Kolthoff and Miller (4) report a combining of the double waves of the bromide ion by the addition of gelatin to bromide solutions. The first of the double waves is explained as due to the formation of a film of mercurous bromide on the surface of the mercury drop, which interferes m-ith the electrode reaction. Apparently, gelatin prevents the formation of a coherent film by being strongly adsorbed on the drop surface. I n this case, the same end result was achieved through the use of Triton X-100. The average half-wave potential of the azide ion was $278 mv. and that of the lead ion -381 mv. v s . S.C.E. Because a distinct plateau did not precede the azide wave (Figure l), the normal extrapolation method of determining VOL. 32, NO. 7, JUNE 1960
759
the wave height was not used. Instead, line AA’ (Figure 2) was drawn through the rising portion of the wave and line BB’ through and parallel t o the slope
Table
II.
preceding this wave. The intersection of AA’ and BB’ was taken as the point of inflection, P. Next, line CC’ was drawn through the residual current and
Polarographic Data for Lead and Azide Ions as Lead Azide in 0.1N Potassium Nitrate Solution
Pb(N& Diffusion Concn., C, Current, Meq./L. zd, pa. 0.2 0.4 0.5 0.6 0.7 0.8 0.9 1.0
id/c
0.84 1.67 2.07 2.54 2.85 3.32 3.72 4.20
4.20 4.18 4.14 4.23 4.07 4.15 4.13 4.20 Mean 4.16
Capillary Constant, Ellz, Mv. us. M2/3t1’6,Mg.2/3Sec.-1/2 S.C.E. Lead Ion 2.178 2.175 2.177 2.177 2.178 2.174 2.179 2.176 2.177 f 0.0013
-378 - 383 -376
-383 -379 -382 -383 -382 -381 f 25
E3/1MV.
31 30 36 33 34 29 ...
... 32 & 2.01
Bzide Ion Standard deviation
1.210/, 0.2 1.04 5.20 2.445 0.4 2.06 5.16 2.440 0.5 2.59 5.18 2.442 0.6 3.19 5.31 2.440 0.7 3.65 5.22 2.441 0.8 4.22. 5.28 2.440 0.9 4.67 5.19 2.441 1.0 5.21 5.21 2.441 Mean 5.22 2.441 f 0,0010 Standard deviation (idle). 0.9Syo (id/C).
273 279 277 276 275 276 277 274 276 f 1 . 4
57 61 58 55 56 57
... ...
57 f 01
Figure 3. Relation between diffusion current and lead azide concentrations
I
3-
DD‘ drawn parallel to it and through the point of inflection, P. The distance, h, parallel to the current axis was considered the diffusion current. The polarographic data for the lead and azide ions are giyen in Table 11. The standard deviation of i d / ‘ C values for the lead ion r a s l.2Y0 and that for the azide ion 0.98%. The capillary constant, m2/3t1/6, differed for the two ions studied, giving values of 2.777 f 0.0013 rng.*i3 sec.-1/2 for lead and 2.441 =t 0.0010 mg.2’3sec.-1/2 for the azide ion. The empirical half-wave potential was practically constant in both cases and showed variations less than & 1.5 mv. Curves of each ion were analyzed according to Tomes’ criterion of reversibility. from which E s 4 - Ellavalues should be approximately equal to 0.0564/n. From the analysis, it appeared that the electrode reaction of the azide ion (Equation 4) approaches complete reversibility, while that for the lead ion, in spite of good agreement of Ellz values, shows a considerable degree of irreversibility. The divergence of the latter may be due in part to the imperfect suppression of lead wave maxima. For comparison, diffusion currents as a function of concentrations of lead and azide ion &re s~ shown in Figure 3. Table I11 shows the results of the stability study of standard polarographic solutions with time and storage conditions. The diffusion current of both ions of the two solutions remained practically constant for 30 days. This result is significant, because it shows that unlike the crystalline salt, standard solutions of lead azide may be stored under normal laboratory conditions for one month and possibly considerably longer. The polarographic method provides it safe and effective technique for the simultaneous determination of both lead and azide ions of lead azide. Its accuracy, which is within 1%, exceeds that of most “wet” methods and compares favorably with other instrumental techniques. LITERATURE ClTED
I 2 CONCENTRATION
Table 111.
6
13 22 31
760
ME9 / L
Dependence of Diffusion Current of Lead and Azide Ion of Lead Azide on Time and Storage Conditions
Time, Days 0
OF PblN,L,
Diffusion Current, pa. Solution A Solution B Lead ion ilzide ion Lead ion Azide ion 4.21 4.20 4.22 4.28 4.18
ANALYTICAL CHEMISTRY
5.18
5.20 ~
5.22 5.22 5.11
4.25 4.22 4.21 4.20 4.20 ~
~~
5.22 5.21 . 5.22 5.20 5.29
(1) Audrieth, L. F., Chem. Revs. 15, 169217 11934). ( 2 ) Dinnis, ’L. >I., Isham, H., J . Am. Chem. SOC.29, 18 (1907). (3) Haul, R. A. W., Schole, E., 2. Elektrochem. 5 2 , 226-34 (1948). (4) Kolthoff, I. M., Miller, G. S., J . Am. Chem. SOC.6 3 , 1405 (1941). (5) Revenda, J., Collection Czechoslou. Chem. Communs. 6,453 (1934). (6) Treadwell, F. P., Hall, W. T., “Analytical Chemistry,” pp. 58-59, W h y , Xew York, 1942. ( 7 ) Vasiliev, D. F., Trudy Komissii Anal. Khim. Otdel Khim. N a u k Akad. S.S.S.R. 2 (5), 90 (1949); Chem. Abstracts 44, 9300.
RECEIVEDfor review January 14, 1059, Accepted March 18, 1960.
