Because of the difficulty and cost of making a conventional gold electrode, many workers resort to using some form of platinum electrode, even though a gold electrode would be more adequate. T h e examples of potentiometric and voltammetric measurements given here indicate the possible usefulness of t h e new type of sensors. They can perform many of t h e tasks of conventional electrodes and can be constructed of many different types of metals or
mixtures of metals while requiring much less metal per electrode than was previously possible. Furthermore, because the outer surface is glass rather t h a n metal, they are much less susceptible to poisoning t h a n conventional metal electrodes. Received for review March 5 , 1973. Accepted M a y 7 , 1973. One of us (G.J.L.) acknowledges the financial support of Owens-Illinois in the form of a graduate fellowship.
Determination of Organic Peroxides in Low Concentration by a Biamperometric Method F. C. Montgomery, R. W. Larson, and W. H. Richardson' Department of Chemistry. California State University. San Diego, Calif. 927 75
Various chemical methods have been developed for the determination of organic peroxides ( I , 2) b u t t h e iodometric method is by far the most widely used. Typically, acetic acid is used as the proton source in t h e reaction of the peroxide with iodide t o give iodine or triiodide. T h e use of acetic acid is perfectly suitable for the determination of peroxides in high concentration; however, it is unsatisfactory for determination of peroxides in low concentration because of large blanks. T o circumvent the problem of large blanks encountered with acetic acid, Hartman and White (3) used citric acid as the proton source in a mixture of tert-butyl alcohol a n d carbon tetrachloride. By this procedure, the liberated iodine was titrated with sodium thiosulfate solution t o a visual end point with the aid of starch indicator. It became necessary for us to develop a n analytical procedure for the determination of small quantities of 1,2dioxetanes (four-membered cyclic peroxides). Although the method of H a r t m a n and White held promise, the visual end-point determination would prove too insensitive for our purposes. We report here a biamperometric procedure with sodium thiosulfate and potassium iodate as titrants and with the peroxide decomposition carried out in a tert-butyl alcohol solution containing potassium iodide and citric acid. T h e method is suitable for conveniently measuring low concentrations of peroxides and should prove useful in kinetic studies or where hazards or natural occurrence dictate analyses of small amounts of peroxides.
EXPERIMENTAL Apparatus. The circuit diagram (Figure 1) for the biamperometric titration ( 4 ) apparatus is a modification of that given by Potter and White (5, 6). The electrodes (A, A) are placed in the solution to be titrated, which is contained in an open 100-ml beaker and stirred magnetically. To whom correspondence should be addressed. Peroxides," Vol. 0 ,D. Swern, Ed., Wiley-lnterscience, New Y o r k , N . Y . , 1971, p 579. ( 2 ) R. M . Johnson a n d I . W. Siddiqi, "The Determination of Organic Peroxides," Pergamon Press, New Y o r k , N . Y . . 1970. (3) L. Hartman a n d M. D. L. White, Anal Chem., 24, 527 (1952). (4) J. T. Stock, "Arnperornetric Titrations,"Interscience Publishers, New ( 1 ) R. D. Mair and R. T . Hall, "Organic
Y o r k . N . Y . , 1965, C h a p . 4 and 5. ( 5 ) E C . Potter and J. F. White, J. Appl. (1957), ( 6 ) Ref. 4 , p 582. 2258
Chem.
Reagents. tert-Butyl hydroperoxide (Lucidol) was purified by azeotropic separation and distillation ( 7 ) . Benzoyl peroxide (Lucidol), lauroyl peroxide (Lucidol), tert-butyl alcohol (MC/B, reagent), acetic acid (DuPont, reagent), benzene (MC/B, reagent), citric acid (MC/B, reagent), potassium iodide (MC/B, reagent), sodium carbonate MC/B, reagent), potassium iodate (MC/B), reagent), and sodium thiosulfate pentahydrate (MC/B, reagent) were used as supplied. Doubly distilled and boiled water was used to prepare the titrants. The citric acid solution was prepared by adding 75 ml of tertbutyl alcohol to 15 grams of citric acid, and the mixture was heated with stirring for about 10 min to dissolve the acid. A potassium iodide solution was prepared from 10 grams of potassium iodide, 0.5 gram of sodium carbonate, and 10 ml of doubly distilled water. Both solutions were purged with nitrogen by means of a fritted glass filter-stick. Weighed amounts of the peroxides were diluted in benzene to known concentrations. The sodium thiosulfate titrant solution ( c a 1.8 X 10-4N) was obtained by diluting 1.8 ml of 0.1N sodium thiosulfate solution to 1.0 1. with boiled doubly distilled water and 50 ml of tert-butyl alcohol. The potassium iodate titrant solution (1.05 X 10-4M) was prepared mmol) to 500 ml of solution by dissolving 11.2 mg (5.23 X with boiled doubly distilled water. Benzene solutions of approximately 0.50M (1.ON) tert-butyl hydroperoxide, benzoyl peroxide, and lauroyl peroxide were prepared for analysis by a conventional method. These solutions were diluted to between approximately 1.1 X 10-4 to 1.1 X lO-3M for analysis by the biamperometric method. Procedure. The sodium thiosulfate solution (ca 1.8 X 10-4N) was standardized in the following manner. A 100-ml beaker, containing a magnetic stirring bar, was charged by means of pipets with 5.0 ml of citric acid solution and then 0.5 ml of potassium iodide solution. This solution was diluted with 15 ml of boiled doubly distilled water and the electrodes (A, A) were placed in the solution and energized. The sample was titrated with standard iodate solution until current flow was detected by the pH meter. The amount of added iodate solution was recorded and then 2.00 ml of sodium thiosulfate solution was added. The current flow decreased and the solution was again titrated with iodate solution until current flow was again detected. The amount of added iodate solution was recorded and this procedure was repeated for additions of 3.00, 4.00, 5.00, 6.00, and 7.00 ml of thiosulfate solution. Prior to the initial titration, 1.00 ml of benzene was added to simulate titrations of the peroxide solutions. The results of the standardization titration are given in Figure 2. The normality of the thiosulfate solution is calculated from Figure 2 with the aid of the titration reactions (Equations 1 and Z), where
(London). 7, 309
( 7 ) P. D . Bartlett (1958).
ANALYTICAL CHEMISTRY, VOL. 45, NO. 13, NOVEMBER 1973
and
R. R.
Hiatt,
J. Amer.
Chem. Soc..
80, 1398
Table I. Comparison of the Biamperometric (Method A) and the Classical Titrimetric (Method B) Procedures for the Analysis of Peroxide Solutionsa % PurityC
__
0.518 1.065 x 1 0 - 4 1.065 x 1 0 - 3
mmol titratedb 0.518 1.065 x 10-4 1.065 x 10-3
0.500
0.500
5.00 x 10-4
5.00 x 10-4
0.500
0.500
5.00 x 10-4
5.00 x 10-4
Peroxide
Mb
t-C4HgOOH
(CsH5COz)z ( C i iH23COz)z
Method A
Method B 97.5 f 0 . 4
...
97.5 f 2.5 97.2 f 0.7
% Difference, method ( A - B) ...
... ... 100.6 3= 0.2
... 99.4 f 0.1
0.0
-0.3 ...
-1.2
... 95.8 f 0.5 ...
92.3 f 1.2
, . .
-3.5
a Benzene solutions. Calculated from the weight of peroxide sample employed. Calculated as the percentage of mmol of peroxide observed by Method A or B to mmol of peroxide in the sample based on the original weight of peroxide. Each entry is given as the average of two measurements. -
2.501
B
0'
2.00
4.00 NaS,O,, ml
6.00
8.00
Figure 2. Standardization of the sodium thiosulfate solution by the biamperometric method with standard 1.05 X 10-4M potassium iodate solution. From this plot and Equation 3, the normaiity of the thiosulfate solution is calculated to be 1.77 X 10-4M
4,
A Figure 1. Circuit diagram for biamperometric titration includes 0.010-in. diameter platinum wire electrodes sealed in glass (A, A ) ; two 1.5-V D-cell flashlight batteries ( B ) ; resistors R1 (1.5 k U ) , R2 (1.0 k!]), R3 (2.5 k Q ) and R4 (50 MR); and a Beckman zeromatic pH meter Each electrode is composed of a 30-mm length of platinum wire, sealed in a 100-mm length of 7 - m m 0.d. capillary tubing, so that 25 m m of the wire protrudes. The end of the platinum wire, in the glass tubing, is silver soldered to 20-gauge copper wire iodine is given as I2 rather than 13- for simplicity (8).The molarity, 10;516H' = 31, 3H20 (1)
+ + I 2 + 2sy0:j'-=
+
21-
+ s,o,'-
( 21
which is equivalent to the normality, of the thiosulfate solution is calculated from Equation 3,
flushed with nitrogen for about 3 min. The flask was stoppered and allowed to stand for 20 min in a dark cabinet. The contents were then poured into the 100-ml titration beaker containing a magnetic stirring bar, and the flask was rinsed with two 5 m l portions of boiled doubly distilled water. The electrodes were placed in the solution and energized. Enough standardized thiosulfate solution was added via a buret to cause the current flow to decrease and reach a minimum, and then an excess of 2-4 ml of thiosulfate solution was added. The standard iodate solution was added from a buret until one drop ( c a 0.02 ml) caused a permanent deflection of the pH meter. The net ml of standardized thiosulfate solution was obtained from Equation 5 ,
Net ml S,O,'-
= {[Total ml S?O,?-] [xs ml S,0J2-]} - [ml S?03-- blank]
(51
where the [xs ml S2032-] term was obtained by converting the ml of iodate titer to ml of thiosulfate with the aid of Figure 2. The blank was obtained by repeating the procedure with the peroxide omitted. From the net ml of thiosulfate and the normality of this titrant, the mequiv of peroxide in the aliquot was obtained. The higher concentration solutions (ca., 1.ON) of the peroxides were analyzed by a classical iodometric method, titrating with 0.1Nsodium thiosulfate to a visual end point (9).
RESULTS AND DISCUSSION where [(ml IO3-)/(ml S 2 0 3 2 - ) ] is the slope of Figure 2 and [IOs-] is the molarity of the standard iodate solution. Aliquots (1.00 ml) of the 1.1 X to 1.1 X 10-3M peroxide solutions were allowed to react with the potassium iodide solution, according to Equation 4,
+
+
+
ROOR 212Hf = I? 2ROH (4) in the following manner. A 50-ml 3 Erlenmeyer flask was charged with 5.0 ml of the citric acid solution and then 0.5 ml of potassium iodide solution. The flask was flushed with nitrogen for about 2 min and then stoppered. A 1.00-ml aliquot of the peroxide solution was pipetted into the flask and again the flask was (8)
I . M . Kolthoff and E. B. Sandell, "Textbook of Quantitative Inorganic Analysis."The Macmillan Co.. New York. N . Y . , 1952. p p 587, 590.
Results of the biamperometric method (Method A ) a t low peroxide concentrations and of the classical titrimetric method (Method B) a t high peroxide concentration are reported as per cent purity of the peroxide. A comparison of these two methods on this basis is given in Table I. In general, the agreement between the two methods is good and some of the difference can be attributed to error in dilution. T h e error between measurements for Method A is highest with the 1.065 x lO-4M solution of tert-butyl hydroperoxide, b u t this could be reduced by using lower concentrations of the titrants. (9) W . H.
Richardson, J. Arner. Chern. S o c . . 8 7 , 247
(1965)
ANALYTICAL CHEMISTRY, VOL. 45, NO. 13, NOVEMBER 1973
2259
Although large blanks are encountered with Method B in determinations of peroxides a t low concentrations, Method A circumvents this difficulty. Blanks for Method A were usually in the range of zero to 3 x 10-5 mequiv of the thiosulfate titrant. Furthermore, there is no need to maintain a nitrogen atmosphere over the solution during titration. In addition, the iodide-citric acid tert-butyl alcohol solution has been allowed to stand under nitrogen for as long as 75 min without increasing the blank. It was unnecessary for us to determine smaller amounts of peroxides and this is no doubt the case in many other applications. However, the biamperometric method could be refined to determine even smaller quantities of peroxides with some simple changes. For example, a smaller titration vessel could be used so t h a t the iodine would be in
higher concentration after transfer from the peroxide-iodide reaction vessel. In addition, the concentrations of the titrants could be lowered. In summary, the present biamperometric procedure offers a convenient method of determining small (to mmol) quantities of peroxides, which is not possible by classical methods ( e . g . , Method B). With simple modifications, even smaller amounts of peroxides could be determined without difficulty. Received for review April 5 , 1973. Accepted J u n e 29, 1973. The authors thank the donors of the Petroleum Research Fund, administered by the American Chemical Society, and the U.S. Army Research Office (Durham) for support of this work.
Determination of Free Cyanide in Ferro- and Ferricyanides J. M. Kruse and L. E. Thibault ltek Corporation. 10 Maguire Road. Lexington. Mass. 021 73
Ferrocyanides have found widespread commercial usage both as photographic bleaching agents and as electrostatic conversion washes. A major problem with the use of ferrocyanides, as well as the closely related ferricyanides, is the formation of hydrogen cyanide. Although the complex cyanides do not appear to be toxic to fish (1-3), the liberated hydrogen cyanide is toxic not only to fish b u t also to putrefraction bacteria and thus raises a disposal problem. The determination of free hydrogen cyanide present in solutions of ferro- and ferricyanides is therefore of importance not only to obtain a measure of the amount present a t any one time, but also to measure the rate a t which free cyanide is formed from the complex cyanides under a variety of conditions. Because the free cyanide in solutions of ferro- and ferricyanide is generally present a t low concentrations, a colorimetric method of measurement appeared desirable. However, because the complex cyanides are themselves colored, and also to prevent any interaction between the reagents and the complex cyanides, a separation of the free cyanide from the complex cyanides was desired. In addition, the measurement of the cyanide should not upset the total system-i. e . , liberate cyanide from the complex cyanides during the analysis. Although extractive procedures for the measurement of free HCN have been described in natural and waste water (4, 5 ) , one of these is carried out under acid conditions which upset the system, while the other was applied to nearly neutral water samples without p H control. Also, a gas chromatographic procedure based on sampling of gaseous HCN above the liquid sample has been described (6). The authors bubble air through the sample, trap the HCN to achieve a sufficiently high concentration, and then measure the trapped (1) C. J. Terhaar et a i . . Photogr. Sci. Eng.. 16, 3 7 0 ( 1 9 7 2 ) . (2) P. Doudoroff, Sewage ind. Wastes. 28, 1020 ( 1 9 5 6 ) . (3) P. Doudoroff, G . Leduc, and C. R. Schneider, Trans. Amer. Fish. SOC..95, 290 (1966) (4) H. A . C. Montgomery, D. K . Gardiner. and J. G G . Gregory, Analyst / L o n d o n ) . 94, 284 (1969). (5) J. M . Kruse and M . G . Mellon, Sewage ind Wastes. 24, 1254 (1952) (6) C. R. Schneider and H . Freund, Anal. Chem., 34, 6 9 (1962).
2260
HCN. The method appeared to give good results, although the procedure appeared quite complex and time consuming. As a simpler procedure for the routine measurement of HCN, the separation of the free HCN from the sample by means of diffusion in Conway microdiffusion cells appeared attractive. The diffusion step requires a minimum of manipulation, can serve to concentrate the free cyanide in the final solution from the levels present in the original sample, required no special apparatus, and could be carried out a t a controlled pH. The final measurement could then be carried out by titration or colorimetry. For rate studies, a final colorimetric measurement was necessary, but the initial concentrations of cyanide in some electrostatic conversion washes were so high t h a t titration of the separated HCN was possible, and even direct measurement in the sample with a specific ion electrode should be possible (7, 8). Initial experiments were conducted to measure the transport of hydrogen cyanide as a function of p H and time. As shown by the d a t a in Table I, complete recovery of free cyanide could be obtained a t p H 7 or less in diffusion periods of five hours or less. At p H values >9, recovery decreased; while a t p H values below 5, the decomposition of complex cyanides was more significant. Thus, p H 7 was used for the diffusion studies.
EXPERIMENTAL Apparatus and Reagents. Microdiffusion Dishes Standard Conway Microdiffusion cells, such as A . H. Thomas Catalog No. 4472-R10 or equivalent, proved satisfactory. Buffer A pH 7.0 KHzP04-NaOH buffer was used. This buffer must be free of formalin, as the formaldehyde will react with any free HCS. Coior Reagents The 3-methyl-l-phenyl-5-pyrazolone,the dimer, and chloramine T reagents were Eastman organic chemicals. The pyridine-pyrazolone reagent was prepared by adding to 125 ml of hot water, stirexcess 3-methyl-1-phenyl-5-pyrazolone (7)
(8)
Orion Research Inc., Bulletin 94-06, 11 Blackstone St., Cambridge, Mass. B. Fleet and H. von Storp, Anal. Chem.. 43, 1575 (1971).
ANALYTICAL CHEMISTRY, VOL. 45, NO. 13, NOVEMBER 7973
