place, the propionic acid is held strongly. Propionic acid molecules not finding a strongly interacting site move on rapidly. This phenomenon contributes to the peak broadening and tailing so greatly manifest with propionic acid. As the size of the sample of propionic acid is increased, the limited number of strongly interacting sites becomes saturated ; and the bulk of the propionic acid moves on more rapidly, thus resulting in the decrease in Io observed. Effects of this type have been reported by Ottenstein (9). Column Poisoning. Previous work ( I ) has shown that 0.1pl samples of the test compounds injected successively into the asphalt column did not change the value of the interaction coefficient. This suggests that no irreversible reactions of the test compounds, or column poisoning, are taking place. However, because of the reproducibility problem and extreme “tailing” encountered with propionic acid, the possibility of irreversible reaction of this test compound with the asphalt was re-examined. Propionic acid was injected repeatedly on the sample asphalt column, using sample sizes 0.1 and 2.0 111. When 0.1-p1 samples were used, the I , varied within the limits of the experimental error. When large 2 . 0 - 4 injections were made, the Io’s dropped from 122 to 95; however, when a 0.1-pl sample was subsequently injected, the previous value of 122 was obtained. These data show no significant “poisoning” by propionic acid. The explanation for the lower values of I , obtained when 2.0 pl of propionic acid were used was given in the previous section. Water Displacement. Water in solute samples in other
systems (IO) has a displacement effect on the solute, thus decreasing retention times. To investigate the possibility of a displacement effect of water as a test compound contaminant, the water was added to the polar test compounds triethylamine, formamide, and propionic acid. Any displacement effect present should be evident with these polar test compounds because of competition by the water for polar sites in the asphalt. The doped samples were compared with anhydrous test compounds by determining Io’s on the same column. The addition of water up to 50 mole percent produced no effect on the Io’s; however, when 80 mole percent water was added to propionic acid, the I , dropped from 120 to 114, indicating that excessive amounts of water may be undesirable. Water itself as a test compound moves through the column with no measurable retention. Trace amounts of water present in carbon tetrachloride solutions of asphalt have previously been shown by infrared studies (11) to hydrogen bond with asphalt at ambient temperatures; however, the near inertness of water in the GLC column probably results from low solubility and the high temperatures, rather than insensitivity to the polar groups present in asphalt. RECEIVED for review December 15,1969. Accepted March 2, 1970. Work reported in this paper was done under a cooperative agreement between the Bureau of Mines, U S . Department of the Interior, and the University of Wyoming. Mention of specific models of equipment or brand names of material is made to facilitate understanding and does not imply endorsement by the Bureau of Mines. ~~
(9) D. M. Ottenstein,J. Gas Chromafogr.,6,129 (1968).
~~~
~
(10) S.G. Perry, J . Chromatogr., 23,468 (1966). (11) J. C. Petersen, Fuel, 46, 295 (1967).
~~~
New Determination of Stoichiometry of the lodometric Method far Ozone Analysis at pH 7.0 A. William Boyd, Clive Willis, and Ronald Cyrl Physical Chemistry Branch, Chalk River Nuclear Laboratories, Atomic Energy of Canada Limited, Chalk River, Ontario, Canada
IN RECENT WORK on the yield of ozone in the radiolysis of oxygen ( I , 2), we compared the optical method of ozone determination using light at 254 nm, with the method of oxidation of iodide in buffered neutral solutions. This showed that for the conditions of our determinations one molecule of ozone liberated more than one molecule of iodine. This paper is a report on this stoichiometry. Previous work has shown that the amount of iodine liberated when ozone is added to iodide solutions is markedly dependent on pH. Ingols et al. (3), measured the iodine 1 Present address, Chemistry Dept., University of Toronto, Toronto, Ontario, Canada.
(1) A. W. Boyd, C. Willis, R. Cyr and D. A. Armstrong, Can. J. Chem., 47,4715 (1969). (2) C. Willis, A. W. Boyd, M. J. Young, and D. A. Armstrong, ibid., in press. (3) R. S. Ingols, R. H. Fetner, and W. H. Eberhardt, Adu. Chem. Ser., 21,102 (1959). 670
ANALYTICAL CHEMISTRY, VOL. 42, NO. 6, MAY 1970
formed on ozone addition to iodide solutions at pH 2 and pH 9. They found that about twice as much iodine was liberated at the lower pH than at the higher. At pH 9 one molecule of oxygen was formed per molecule of iodine liberated-Le., (Oz/Iz= 1) as expected from the classical equation 0 3
+ 2H+ + 21-
0 2
+ 1 2 + H20
(1)
However, there was a considerably lower ratio of oxygen to iodine liberated at pH 2. They also found that more iodine was liberated at pH 7 than at pH 9. Byers and Saltzmann (4), carried out similar measurements in iodide solutions at pH 7 and pH 14. They obtained an increase in the 13-/03 ratio on going to the lower pH (12 will exist in the complexed IB- form in iodide solutions). They assumed that at pH 7 the ratio of molecules of iodine to molecules of ozone was 1 :1. This assumption is not in accord with the conclusions of Ingols et al. (4) D. H. Byers and B. E. Saltzmann,Adu. Chem. Ser., 21,93 (1959).
Table I. Yields Observed in Oxygen Irradiations Optical method Iodide method Dose O3 formed O3formed Dose 1 3 - formed eV X 10'9 mole X 10-6 Dose X 10-25 eV x 1019 mole X 10-6 0.1893 0.0291 0.1017 0.4030 2.129 0.5716 2.193 0.2607 0.0382 0.1220 0.2893 0.0612 0.2141 0.6673 2.307 0.3712 0.1807 0.5803 0.7167 1.931 0.1926 0.6317 0.8473 2.015 0.4206 0.2680 0.8779 1.886 0.4688 0.8842 1.0175 2.349 0.4331 0.2741 0.9529 0.4886 0.2897 0,9850 1.034 2.116 0.5080 0,4804 1.542 2.142 1.088 0.3617 1.210 2.186 0.6280 1.373 0.6556 1.950 2.282 0.6170 1.408 0.7227 0.6590 2.248 1.447 2.002 1.988 0.7779 2.516 2.070 1.041 0.9057 2.923 2.212 1.038 2.133 0.8926 2.690 2.320 1.037 2.218 0.9178 2.900 4.757 2.059 9.796 0.7938 2.57 9.414 2.274 21.41 0.9210 2.72 0.8969 2.73 mean 0 3 formed mole eV-1 X 10-26 = 2.108 0.9163 2.78 Dose 0.8470 2.698 0.8626 2.95 Standard deviation = 0.128 0.9163 2.905 0,8450 2,784 0.9142 2.837 1.027 3.39 0.9588 3.180 O3 formed 3.22 f 0. l 6 ___ratio - 1.53 i 0.12 1.022 3.180 13- formed 2. l1 f 0 . l 3 1.022 3.380 1.022 3.19 mean 1 3 - formed mole eV-l X = 3.222 Dose Standard deviation = 0.157
Because of this disagreement and because we have been able to measure the concentration of ozone absolutely, we have re-examined the stoichiometry of the iodide method at p H = 7.0. Our results are in general agreement with the assumptions of Ingols et al. EXPERIMENTAL Ozone Production. The ozone was produced by irradiating gaseous oxygen both at 195 and 298 OK with single pulses from a n electron accelerator, a Febetron 705. Details of the irradiation procedures and dosimetry are given elsewhere (I, 2, 5). The oxygen was contained in quartz irradiation cells of approximately 100-ml volume fitted with a 2-mm bore stopcock. Stopcocks were greased with either Dow Corning high vacuum silicone lubricant or KEL-F No. 90 grease (3M Corp.). Neither of these lubricants appeared to affect the ozone determinations and in general the KEL-F grease was used. Normally about 2 X mole of ozone was produced from a single pulse irradiation of oxygen a t 700 Torr. The reproducibility of the irradiation pulses was & 2 % as determined calorimetrically (6). The yield of ozone was the same a t 195 and 298 OK. Optical Determination of Ozone. The irradiation cell was attached to a n evacuated 10-cm gas spectrophotometer cell. The oxygen in the irradiation cell was irradiated and, after waiting a suitable time (-10 min) for mixing, the gas was allowed to expand into the spectrophotometer cell. This was then removed and its optical density scanned as a function (5) C. Willis, A. W. Boyd, and D. A. Armstrong, Can. J . Chem., 47, 3783 (1969). (6) C. Willis, 0. A. Miller, A. E. Rothwell, and A. W. Boyd, Ada. Chem. Ser., 81,539 (1968).
13- formed Dose X 10-26 3.498 3.196 3.497 3.103 3.280 3.276 3.476 3.400 3.210 3.345 2.974 3.411 3,234 3.227 3.014 3.160 3.238 2.953 3.044 3.034 3.185 3.420 3.170 3.295 3.103 3.301 3.317 3.112 3.307 3.121
of wavelength in a Perkin-Elmer 450 spectrophotometer. The concentration of ozone was determined a t 254 nm using the extinction coefficient E = 135 cm-l S.T.P., determined absolutely by DeMore and Raper (7) who measured the ozone by its decomposition to oxygen. It has been confirmed more recently by Griggs (8). No difference was determined in the ozone concentration on standing up to 30 min in the light after irradiation. Iodide Determination of Ozone. The irradiated cell was cooled by liquid nitrogen and a few milliliters of iodide reagent were added through the stopcock. The cell and reagent were shaken until they reached room temperature. The cell was washed out with iodide reagent and diluted to 250 ml. An aliquot of this solution was measured in a 1-cm cell in a Beckman D.U. Spectrophotometer. An unirradiated sample of oxygen gave negligible optical density compared to a n irradiated sample. The iodide reagent was made by dissolving 20.0 grams KI, 27.2 grams KH2P04,and 28.4 grams Na2HP04in pure water and diluting to 2000 ml, to give a p H = 7.0. A calibration curve for iodine was made using a solution of Is- previously standardized with pure dry arsenious oxide. Materials. The oxygen used was Matheson C.P. grade gas. All chemicals were reagent grade. RESULTS AND DISCUSSION
The results of the optical and the iodide methods are compared in Table I. As can be seen, the average concentration of ozone formed per unit dose, as measured optically, is 2.108 i 0.128 X mole eV-l: the amount of I B - liberated is 3.222 + 0.157 X 10-25 mole eV-1. The ratio of these, then, (7) W. B. DeMore and 0. Raper, J . Phys. Chem., 68,412 (1964). (8) M. Griggs, J . Chem. Pliys., 49, 857 (1968). ANALYTICAL CHEMISTRY, VOL. 42, NO. 6, MAY 1970
671
*
gives the stoichiometry of the ozone-iodide reactions as 1.53 0.12 molecules of iodine liberated per molecule of ozone. This result is not in agreement with the ratio of 1 :1 obtained by Saltzmann and Gilbert (9). Our ratio is based on the absolute determination of the amount of ozone by the optical method. This method is also the basis for our yield of ozone in the electron pulse radiolysis of oxygen of G(O3) = 12.8 f 0.6 molecules per 100 eV and this value is in good agreement with two independent measurements, G(03) = 13.8 0.7 (10) and G ( 0 3 )= 13.0 d= 0.6 (11). If the stoichiometry of one iodine molecule liberated for one ozone molecule absorbed assumed by Ingols ef a/. (3) for pH 9.0 can be applied to the alkaline conditions used by Byers and Saltzmann (4) at pH 14, then the results of these latter workers give a stoichiometry for neutral pH of 1.54 molecules of iodine liberated per molecule of ozone absorbed. This is in excellent agreement with our observed value. Discussion of other values which show the enhanced yields of iodine are given in the papers of these authors (3, 4), and need not be repeated here. Any mechanism which explains this enhanced yield must explain the important observations of Ingols et a/: more iodine and less oxygen are liberated as the pH is lowered. Although the mechanism advanced in the following discussion is somewhat speculative, it does correlate these observations. Alder and Hill (12) found that UV absorption of ozone decreases faster than the decreasing ability of the same solution to liberate iodine. They concluded that a decomposition product of ozone, the hydroperoxyl ion, HOz-, liberates iodine from iodide in the absence of ozone. Other workers have arrived at similar conclusions (see Reference 3). It has been observed (13, 14, that the ozonide ion 0 3 - is formed in the decomposition of ozone in alkaline solution. It would also be expected to be formed by reaction with iodide, i.e., 1 0 3 I-+o3-12 (2) 2
*
+
+
(9) B. E. Saltzmann and N. Gilbert, Amer. Znd. Hyg. J., 20, 379 (1959). (10) J. A. Ghormley, C. J. Hochanadel, and J. W. Boyle, J . Chem. Phys., 50,419 (1969). (11) G. M. Meaburn, D. Perner, J. LeCalve, and M. Bourene, J . Phys. Chem., 72, 3920 (1968). (12) M. G. Alder and G. R. Hill, J . Amer. Chem. SOC.,72, 1884 (1950). (13) J. Weiss, Trans. Faraday SOC.,31, 668 (1935). (14) G. Czapski, A. Samuni, and R. Yelin, Israel J . Chem., 7, 167 (1969).
672
ANALYTICAL CHEMISTRY, VOL. 42, NO. 6, MAY 1970
Presumably, under alkaline conditions, the second step of the iodine liberation would then be of the form
+ I- + HzO
03-
+
20H-
+ + 211 2 0 2
(3)
This gives a stoichiometry of one to one in agreement with Equation 1. Now, if the ozonide ion existed in equilibrium with HO3, H+
+
a HO3
03-
(4)
then as the pH is lowered, less 03-ion will be available for Reaction 3. Furthermore, if H 0 3 reacts with iodide via Reaction 5 rather than via Reaction 3 HO3
+ 1-
+
HOs-
+ 2-1 Iz
(5)
then H03- which is the dissociated form of H203, could oxidize additional iodide to iodine.
+ Hz0 + 21 Iz
(6)
+ 21 Iz -
(7)
1 H202+1-+0H+OH-f-I~ 2
(8)
HO3-
+ I- + 2H+
HOz
+
+ I- + H+
OH
HOz +
H20z
+ I - + O H - + - I z21
-
(9)
Such a series of reactions would lead to enhanced yields of iodine and reduced yields of oxygen and would also show a pH dependence due to ionization of the H 0 3 (which must have apKvalue of about 7 to 8). The species H z 0 3and HOa- are well known from radiation chemistry studies (15)and in fact a stoichiometry for oxidation given by Reactions 6-9 has been reported (16,17).
RECEIVED for review January 9, 1970. Accepted February 25, 1970. (15) B. M. J. Bielski and H. A. Schwarz, J. Phys. Chem., 72, 3836 (1968) and references contained therein. (16) G. Czapski and B. H. J. Bielski, ibid., 67, 2180 (1963). (17) K. Sehested, E. Bjergbakke, 0. L Rasmussen, and H. Fricke, J. Chem. Phys., 31,3159 (1969).
