tetroxide (4,11). Precipitation with dimethylglyoxime in acid solution will separate palladium froni platinum (2, 4 ) and from most other elements, including nickel. Additional work is needed to test the reliability of the separation of small amounts of palladium from larger amounts of nickel, cobalt, and iron. Composition of the Complexes. Application of the mole ratio method (12) shorvcd the existence of two complexes: a red complex, n i t h absorbance maximum a t 548 nig and a side plateau a t 517 mu (curve A of Figure I ) , in which the palladium to reagent ratio is 1 to 2, and a yellow complex, Rith absoibance maximum a t 466 mp (curve C of Figure l ) , in n hich the palladium to reagent ratio is 1 to 1. Data for the mole ratio plots were taken in two series. I n one series the concentration of the reagent was held constant (3.0 X 10 4.M in the final solution) and the palladium concentration n as varird; absorbance measurements were made a t several wave lengths, including the LT ave length of the absorption peaks of the two species (Figure 2). In the other series the palladium concentration vas held constant (1.5 X 10-6JI in the final solution) and the concentration of quinoxaline-2.3-dithiol n as varied (Figure 3). The evidence for the 1 to 1 and the 1 to 2 complexes iq clearly shown; the
slight deviations in the positions of some of t h r breaks in the curves from the integral mole ratios could be caused by uncertainty of the exact purity of the solid quinoxaline-2,3-dithiol reagent.
Table V.
Tolerance for Foreign tons
(Palladium concentration, 0.4 and 1.60 p.P.m.) Tolerance, Foreign Ion P.P.hI. Platinum( IV) 0.3 Rhodium( 111) 8 “ RutheniumlITI’i Osrnium(1V) ’ 0.2 Iridium( IV) 10 Sickel(11) 0.05 Cobalt(11) 0.02 Iron( I1j 4 Iron(II1) 1 GoldlIII) 10 10 >20 8
>20
15 10
17 70
200
50 > 100 60
Acetate Sulfate 50 Phosphate a Sitrite caused decrease in absorbance, probably through complex formation with palladium.
ACKNO WLEDGMENl
Preliminary n ork on thr palladium quinoxaline-2,3-dithiol color reaction was performed in this laboratory by Harold T). Powell in 1957, while cmployed on Sational Science Foundation Grant SSF GI889 to the senior author. LITERATURE CITED
(1) Ayres, G. H., Alsop, J . H., ANAL.
CHEX 31, 1135 (1959).
R., Ibid., 25, 980 (1953). (3) Cheng, K. L., I b i d . , 26, 1894 (1954). (4) Hillebrand, W.F., Lundell, G. E. F., Bright, H. A,, Hoffman, J. I., “Applied Inorganic Analysis,” 2nd Pd., pp. 356, 372, )$?ley, Sew York, 1953. ( 5 ) Morrison, D. C., Furst, A , , J . Org. Chem. 21, 470 (1956). (6) Sielsch, W., 2. anal. Chem. 142, 30 (1954). ( 7 ) Rice, E. FY.,AX.kL. CHEM. 24, 1995 (1952). (8) Ryan, D. E., Analyst 76, 310 (1951). (9) Skoog, D. .4.,Lai, hI., Furst, A . , ANAL.CHEM.30,365 (1958). (10) Sogani, N. C., Bhattacharyya, S. C., I b i d . , 29,397 (1957). (11) Steele, E. L., Yoe, J. H., Ibid., 29, 1622 119571. (12) Yoe, J. H., Jones, H. L., ISD. ENG. CHEM.,ANAL.ED. 16, 11 (1944). (13) Yoe, J. H., Kirkland, J. J . , ANAL. CHEJI.26,1335 (1954)
( 2 ) Ayres, G. H., Berg, E .
RECEIVEDfor review July 24, 1959. ilccepted August 31, 1959.
Reactivity of Oxidizing Agents with Potassium Iodide Reagent AUBREY P. ALTSHULLER, CAROL M. SCHWAB, and MARIBEL BARE Robert A. Tuft Sanitary Engineering Center, Public Health Service, U. S. Department o f Health, Education, and Welfare, Cincinnati 26, Ohio
b The reactions of iodide ion with peracetic acid, succinic acid peroxide, cumene hydroperoxide, di-tert-butyl peroxide, hydrogen peroxide, ozone, nitrite ion, and tert-butyl nitrite in the 10+ to 10-sM range have been investigated at various pH values. The effect of adding sulfamic acid along with the phosphoric acid used for acidification was studied also. Reaction rate curves were obtained for most of the reactions studied.
T
HE PRESCWE of oxidizing materials in the atmosphere has been detected in man\- localities by the use of iodide solutions (4,10-13, 1 5 ) . The iodide ion in solution is readily oxidized by many oxidizing agents; consequently, the term “oxidant” 13 used t o represent the
net result of atmospheric reactivity with iodidc. I n the Los Angeles area, the major part of the oxidant usually is ozone, but the nature of the remaining oxidizing matcrial is not dcfinitely knonn ( 7 , 1 1 , 12). Although somr work has bwn tione on the reactivity of srveral organic- reroxidcs in the niicrogram rang(. nitli iodide ion ( S ) , it was f P l t that it noiild be useful to investigate thrl reaction of 6 X 10-’M (1% by weight) iodide ion with ozone, hydrogen peroxide, peracetic acid, succinic acid peroxide, cumene hydroperoxide, di-tert-butyl peroxide, nitrite ion, and tert-butyl nitrite in the to lo-5.V range undcr varying conditions of pH. As sulfamic acid has a decided effect in quenching the reactivity of nitrogen dioxide with iodide, particularly in 1M sodium hydroxide solution
( d ) , the effect of sulfamic a d \vas drtermined on all of the oxidizing materials studied. Reaction rate data were obtained to decide the conditions under which the determinations would br least deprndrnt on time of standing befow analysis. EXPERIMENTAL DETAILS
The cumenc hydroperoxide, peracetic acid, succinic acid peroxide, di-tertbutyl peroxide, hytlrogrn prroxide, tert-butyl nitrite, and inorganic nitrite were all diluted to about 0.5 pmolt per ml., and 0.25 to 2 ml. of these solutions were diluted up to 10 ml. with l.007, potassium iodide solutions. Consequently, t,he concentrations of reactants were 10-5 to 10-4.11 in prroxides and nitrites and 6.0 X 10-*.\1 in iodide ion. The ozone was produced by a small ultraviolet ozone generator and was passed a t VOL. 31, NO. 12, DECEMBER 1959
0
1987
Table 1.
Triiodide Ion Yield from Reactioins with Iodide Ion at Various pH Values
PH Initial Final 7 1.2 10 1.0 0.9 12 13 1.2 14 4.0 7 1.3 10 1.05 12 0.9 13 1.15 14 4.2 7 5.1 10 2.5 12 1.7 13 2.8 7 6.6 10 10.0 12 11.4 13 12.3
Acidifying Medium 76%. HIP04 ( f sulfamic acid)
36%",PO4
3.6%HaPO4
No acidification
Absorbance per pmole Peracetic Succinic acid Cumene acid peroxide hydroperoxide 2.70 2.05 0.85 1.80 1.00 1.60 1.90 1.20 1.65 1.55 1.15 1.65 1.35 1.85 0.55 2.75 1.90 0.85 1.65 2.00 0.85 1.65 2.00 1.15 1.65 1.65 1.05 1.45 0.50 1.35 2.65 1.85 0.60 1.60 1.80 0.45 1.75 0.50 1.55 1.55 1.55 0.40 2.65 2.15 0.85 0.45 0.40 0.20 0.20 0.25 0.10 0.00
0.00
0.00
RESULTS A N D DISCUSSION
Table 11. Wect of Sulfamic Acid on Reaction of Nitrous Acid with Iodide Ion to Form Triiodide
PH Initial Final 7 1.2
Absorbance Ratioa 2.5
Without sulfamic acid/with sulfamic acid. Ten minutes after acidification with 36y0 HaPOl saturated with sulfamic acid. Table 111. Relative Triiodide Yields from Reaction of Ozone with Iodide Ion at Various pH Values
PH No. of Relative Initial Final Runs Absorbance" 7 1.2 54 1.00 io i.0 47 0.67 12 0.9 58 0.70 13 1.2 56 0.67 14 4.0 35 0.63 a Ten minutes after acidification with 36Y0 HaPO, saturated with sulfamic acid.
various flow rates into the 1.00% iodide solutions. The initial pH values of the iodide solutions varied nominally from 7 to 14. The solutions were made up a t about pH 7 with a phoshate buffer, a t pH 10 with a glycine guffer, and a t pH 12, 13, and 14 using 0.01, 0.1, and 1M sodium hydroxide. W'hen these solutions were acidified with either 36 or 3.6% phosphoric acid solution, the final pH values ranged from 1 to 5 units. The phosphoric acid solutions were saturated with sulfamic acid. The analytical determinations were made with a Beckman DU spectrophotometer using 1-cm. Corex cells. The absorbance of the triiodide ion formed was determined a t 353 mp.
1988
ANALYTICAL CHEMISTRY
The absorbaxxe of the triiodide ion formed per pmole of oxidizing material is listed in Table I for the reactions of peracetic acid, succinic acid peroxide, and cumene hydroperoxide with iodide ion. The absorbances were measured after standing for 10 minutes following addition of the oxidizing agent both with and without acidification. The reaction rate is greatly diminished in alkaline solution. The reactivities for peracetic acid, succinic acid peroxide, and cumene peroxide with iodide ion were only 15 to 20y0 at pH 10 of what they were a t pH 7. The amount of triiodide ion formed drops off a t pH 12 to about 10% of that a t pH 7 and to nearly 0 at pH 13. A real difference exists between the absorbance of iodide solutions a t an initial pH of 7 to which peracetic acid is added and then acidified, and iodide solutions at higher initial pH values which are also acidified. Even without acidification, iodide solutions a t initial pH of 7 containing peracetic acid reach the same absorbance despite the differences in final pH of 1.1 and 6.6 in the two instances. It appears that the initial excess of hydroxyl ion suppresses the initial stages of the reaction. As has been suggested regarding the ozoneiodide reaction in highly alkaline medium ( 2 ) , it may be that the peracetic acid (as well as the cumene hydroperoxide and succinic acid peroxide) rapidly converts the iodide to hypoiodite, which then reacts very slowly to form iodate, which in turn oxidizes iodide to iodine. The di-tert-butyl peroxide-iodide solutions were studied under the same experimental conditions, but no increase in absorbance above the blank was observed even after 3 hours. Iodide solutions containing as much as 10 to 20 pmoles per ml. of the di-tert-butyl per-
oxide showed only about a 0.1 unit increase in absorbance after 24 hours. This is in agreement with previous results (3). Iodide solutions a t room temperature apparently are essentially unreactive with dialkyl peroxides added in micromole concentrations. Sulfamic acid has no decided effect on cumene hydroperoxide, peracetic acid, or succinic acid peroxide, nor on the ozone-iodide and hydrogen peroxideiodide reactions. Small effects are generally due to the increased acidity of some of the solutions acidified with the phosphoric acid-sulfamic acid compared to those acidified with phosphoric acid alone. The effect of sulfamic acid on reducing the reactivity of iodide ion with nitrite ion by removal of nitrous acid was also briefly investigated (Table 11). The reduction in reactivity is most pronounced in reaction mixtures initial a t pH 14, but the reactivity towards nitrous acid of iodide ion is reduced a t all pH values used by addition of sulfamic acid. The reaction of nitrous acid with sulfamic acid arpears to be as rapid or more so than the nitrous acid-iodide reaction under all the exrerimental conditions employed in this work. If the solutions are not acidified, the nitrite ion itself does not appear to react LTith the iodide ion. Similarly, tert-butyl nitrite does not react with iodide solutions maintained a t pH 7. Gast.ous nitrogen dioxide drops rapidly in reactivity with iodide with incrpasing pH (6). When sulfamic acid is added to the acidifying medium, the reaction of the nitrous acid formed with iodide ion has been shown previously to be greatly inhibited when the ozone is absorbed in 1M sodium hydroxide solutions of iodide ( 2 ) . However, the sulfamic acid will not eliminate any direct reaction of nitrogen dioxide with iodide which may occur (8). The reactivity of iodide ion with ozone drops off rapidly with increasing initial pH (Table 111). It should be noted that the relative absorbance for solutions at initial p H values of 10, 12, and 13 are only slightly greater than those a t p H 14. The ratio of the reactivity of the solutions a t initial pH values of 7 and 14 is in agreement with previous results (a). The changes in the concentration of triiodide ion with time as the iodide ion reacts with peracetic acid, succinic acid peroxide, and cumene hydroperoxide are shon-n in representative curves (Figures 1 to 3). The changes in the triiodide ion concentration are represented by the changes in absorbance at 353 mp. The initial rates of formation of triiodide ion are rapid, particularly for the reactions of peracetic acid and succinic acid peroxide. The relative rates of the reactions in their initial stages are in the following order : peracetic acid-iodide ion
0
I
1
I
I
1IK.N.
Figure 1. Typical reaction rate curves for reaction of peracetic acid with iodide ion
greater than succinic acid peroxideiodide ion much greater than cumene hydroperoxide-iodide ion. The data on the ozone-iodide reaction indicate that ozone doesnot react with iodide ion to form triiodide ion as rapidly as peracetic acid or succinic acid peroxide. The rate of formation of triiodide ion is relatively independent of initial pH between 7 and 13. I n most of the reactions, if the initial pH is 14 or if the reactions occur in alkaline solution, a very sharp decrease in the rate of formation of triiodide ion occurs after the initial stages of the reactions. Similar data were obtained for the hydrogen peroxide reaction with iodide ion a t initial pH values of 7 and 14. The concentration of triiodide ion stopped increasing in these measurements between 30 and 60 minutes. Because hydrogen peroxide reacts very slowly, if a t all, in alkaline medium to form triiodide ion ( I ) , no data were obtained for unacidified solutions. The decrease in absorbance observed in the later stages of the reactions of iodide ion with peracetic acid, succinic acid peroxide, and cumene hydroperoxide in neutral and alkaline solutions may be the result of successive reactions of the oxidizing agent with the triiodide ion to form a molecular or ionic species which absorbs less a t 353 mp than does the triiodide ion. I n alkaline solution, hydrogen peroxide has been shown to react rapidly with triiodide to convert it back to iodide (1). Either the same or similar reactions must be destroying the triiodide ion in the reactions in neutral and especially in alkaline medium discussed above. Not only ozone, but also peroxy acids, acyl peroxides, hydroperoxides, hydrogen peroxide, nitrous acid, alkyl nitrites, and nitrogen dioxide all will react rapidly with iodide ion in acidic medium even a t very low concentrations (1 to 10 y per ml. of solution). Only dialkyl peroxides are unreactive. I n neutral or basic medium, hydrogen peroxide, nitrites, and alkyl nitrites are unreactive or relatively urireactive with iodide ion. I n a weakly acidic medium, sulfamic acid will greatly reduce the reactivity of nitrogen dioxide and iodide ion. How-
ever, even after the elimination of a number of oxidizing agents, the more active organic peroxides will still react with iodide ion as rapidly or more rapidly than ozone does. The reactions of iodide ion with peracetic acid and succinic acid peroxide a t pH 7 appear convenient for analytical use as the rate curves are fairly horizontal from 10 to 60 minutes compared to the corresponding curves a t p H 10, 12, and 13 (not shown on Figures 1 and 2). However, the corresponding horizontal portion of the rite curve for the iodide ion-cumene hydroperoxide reaction does not occur until about 3 hours after the reaction starts, although it continues horizontal after that time for at least 4 hours.
.
0
I
2 ll*.I*.
.
..
1
1
6
Figure 2. Typical reaction rate curves for reaction of succinic acid peroxide with iodide ion
The rate curves for the reactions of peracetic acid and succinic acid peroxide with iodide ion in solutions initially a t p H 14 but subsequently acidified are much flatter after the initial few minutes of the reaction than are those a t lower initial and final pH values. These results iadicate that analytical measurements of this reaction can be made conveniently a t any time between about 10 minutes and several hours. The corresponding rate curve for the cumene hydroperoxide-iodide ion reaction does not level off for several hours, However, the slope of the curve is small, so that measurements are possible within the same time period as for peracetic acid and succinic acid peroxide.
ultraviolet absorption results are higher than the chemical results. At night the ultraviolet values are often the higher ones. At least two active species besides ozone probably are involved in these determinations. One or more are responsible for the high chemical results; the others yield the high ultraviolet values. Peroxyacyl nitrite hydrolyzes in aqueous solution to give an acid solution and it oxidizes iodide ion readily (9). The ultraviolet spectrum of peroxyacetyl nitrite apparently is known qualitatively and then only down to about 2800 A. (16). If it is assumed that peroxyacyl nitrites do not absorb strongly compared with ozone in the 2600-A. region but react with the chemical reagents used, then the former may possibly be responsible for the data showing the high chemical results. Peroxy acids are formed a t higher concentrations (1000 p.p.m.) of hydrocarbons and nitrogen dioxide (9), but there are only uncertain indications of such a species in reactions conducted in the 1- to 10-p.p.m. range (1’7). Ozonides have been detected both in nitrogen dioxide-lefin (17) and olefin-ozone reactions (16) in the 1- to 10-p.p.m. range. Their reactivity in the chemical procedures does not seem to be well established. Some data in the literature (14) indicate that ozonides do not show an absorption peak near 2600 A. and absorb weakly to below 2200 A. It is not certain whether or not one of these species could be present in sufEcient quantity to give the higher chemical values for “oxidant.” I I
I
I
0
I
1 IIK,
I
3
I
M.
Figure 3. Typical reaction rate curves for reaction of cumene hydroperoxide with iodide ion
INTERPRETATION OF ATMOSPHERIC OXIDANT RESULTS
I n the studies of atmospheric oxidizing power, results have been obtained by automatic instruments using 20% potassium iodide at p H 7 , phenolphthalein, and ultraviolet absorptian (6, 7, 12). All of these methods have been calibrated directly or indirectly with respect to ozone. Yet on some days, particularly a t midday, the results from the chemical methods are appreciably higher than those from the ultraviolet absorption, while on other days, the
Nitrogen dioxide is usually low when ozone is high; furthermore, part of what appears to be nitrogen dioxide by the modified Griess procedure may actually be peroxyacyl nitrite (16). The portion which is nitrogen dioxide reacts with iodide a t p H 7 in the ratio of 5 p.p.m. of nitrogen dioxide for each p.p.m. of ozone (4). The total effect of the nitrogen dioxide would not explain the larger portion of the discrepancy in results for the days on which the chemical values for “oxidant” are higher VOL 31, NO. 12, DECEMBER 1959
1989
than the ultraviolet values for “ozone.” As nitrogen dioxide does not absorb in the same range as ozone, it could not be responsible for the high ultraviolet va1uc.s for “ozone.” The higher ultraviolet values could be explained if a sufficient concentration of reducing agent were present to reduce the chemicd oxidant readings. Sulfur dioxide would be a likely suspect. Data on sulfur dioxide concentrations have not been published along with t,hr “oxidant” v:ilues. so the possible importance of the sulfur dioxide is difficult to evaluate. Several chemical specirs. and especially peroxyxyl nitrite, may be responsihlc for the hig!ier results for “oxidantsJJ obtained chrmicnlly than by ultraviolet analysis. There is no dtafinite evidence as to the species rcsponsible for high ultral-iolpt results for ‘Lozonc.JJ
X. E., J . .4ir Pollution Control .lssoc.
LITERATURE CITED
(1) Bray, K. C., Liebhafsky, H. A,, J . A m . Chem. Soc. 53, 38 (1931). ( 3 ) Byers, D. H., Saltzman, B. E., A m . Ind. Hyg. Assoc. J . 19, 25 (1958). (3) Cadle, R. D., Huff, H., J . Phys. Chem. 54, 1191 (1950). ( 4 ) Cholak, J., Schafer, L. J., Yeager, D. W., J . Air Pollution Control Assoc.
5.227 (1956). (5) ‘Effenberger, E., 2.anal. C h e m 134, 106 (19511. ( 6 ) Faith, W. L., Hitchcock, L. B., Xciburger, M.,Renzetti, Y . A . , Rogers, L. H., Second Tech. Progress Rept., KO.12, Air Pollntion Foundation, Los Angeles, Calif., November 1955. ( 7 ) Faith, W. L., Renzetti, S . A,, Rogers, L. H., Third Tech. Progress Rept., KO. 17. .4ir Pollution Foundation. Los Angeles, Calif, Xarch 1957. (8) Fay, H., Nohr, P. H., lIcDanie1, P. K., A m . I n d . Hyg. d s s o c . Quart. 18, 19 (1957). (9) Hanbt, P. L., Stephens, E. R., Scott,
5,219 (1956). (10) Littman, F. E., Renolirl. 11. \V.! ~\N.IL. CHEX.2 5 , 1180 (1953). ( i l ) Littman, F. E., Narynowski, C. R.,I b i d . , 28, 810 (1956). (12) Renzetti, S . -4..Romariovalq. J. C., J . A i r Poiliction Control -4ssor. 6 , 154 (1956). (13) Richards, L. ll., Zbirl.j 5 , 216 (1956). (14) Riechc, *4.,Koch, I
