4066
turn becomes less r e ~ t r i c t i v e . ~Thus ~ the enhanced intensity a t the longest wavelength transition shifts the apparent absorption maximum to 502 mp for AODXA, while for AO-NaPLG the absorption maximum is displaced to 495 mp with a broadened, longer wavelength limb (Figures 1 and 2 ) . It should be noted that the position of the maximum intensity of the a band has been shown to vary with different polymers.2b The changes in rotational strengths with D I P of three negative partial Cotton effects in the AO-DNA system (Table 11) are what would be expected from the spectroscopic studies of A0.5 A t D I P 0.02, the strongest is the negative Cotton effect a t 500 mp which would correspond to the 0 --t 0 vibrational level of the ~LI, band of unbound AO. The rotational strength of this component, becomes overshadowed by the positive one a t longer wavelengths as DIP increases. It appears reasonable. therefore, to conclude that for AO-DNA the positive partial Cotton effect is associated with the ‘L, band, and the three negative partial Cotton effects, whose positions agree quite well with those found by Zanker, et a1.,5130for free AO, are related to the vibrational levels of the ‘Lb transition. The less orderly variation of the partial Cotton effects with DIP for AO-NaPLG makes an interpretation of the origin of its optically active transitions more dif-
NOTES
ficult. For AO-NaPLG at DIP 0.0001 the same argument used for AO-DNA may be applied to all but the fourth partial Cotton effect (Table I1 and Figure 9). This transition could be assigned to one of the higher order vibrational levels of the ‘La band,30with the qualification that some of the intermediate levels were not resolved in the present work. Alternatively, the positive Cotton effect at 441 mp may be attributed to a transition of A 0 dimer, since a small fraction of bound dye could exist as aggregates. At, higher DIP ratios the situation is complicated. Whether or not the positive partial Cotton effect at 522 mp (Figures 7 and 8) belongs to the ‘La transition of monomeric bound A 0 is uncertain. Since A 0 can bind to NaPLG in a number of ways the dye, being bound in different environments, could give rise to partial Cotton effects a t different wavelengths. Finally, we wish to point out that the role of water has been ignored for the present. Also the problem of whether the ground-state interaction between bound monoprotonated A 0 and polymer sites modifies the electronic configuration of the dye remains unresolved. ’
Acknowledgment. The authors are indebted to Dr. E. Charney of this laboratory for useful discussions and wish to acknowledge the assistance of Mr. R. Shrager of the Computation and Data Processing Branch.
NOTES
The Reaction of Sulfur Dioxide with Active Nitrogen by A. Jacob8,’&R. A. Westbury,lb and C. A. Winklerla Departments of Chemistry, iMcGill University, Montreal, Quebec, Canada, and Marianopolis College, Montreal, Quebec, Canada (Received M a y IS, 1966)
Since the reactions of active nitrogen with hydrogen sulfide and with sulfur have been studied previously,2 it was of interest to investigate the corresponding reaction with sulfur dioxide, in which sulfur possesses a formal positive charge. Previous studies have indicated that SO2 was not decomposed by active nitrogen.3v4 The Journal of Physical Chemistry
The apparatus and methods were essentially similar to those used in many earlier studies from this laborat ~ r y . ~The , ~ system was of the conventional fastflow type in which active nitrogen was formed by using either a condensed electrode or a microwave discharge. The discharge was operated for at least 60 min in the “poisoned” system and for 2 hr in the “unpoisoned” (1) (a) hlcGill University; (b) Marianopolis College. (2) (a) R. A. Westbury and C. A. Winkler, Can. J . Chem., 38, 334 (1960); (b) J. A. S. Bett and C. A. Winkler, J . Phys. Chem., 59, 371 (1955). (3) K. D. Bayes, D. Kivelson, and S. C. Wong, J . Chem. Phys., 37, 1217 (1962). (4) J. J. Smith and W. J. Jolly, I n o r g . Chem., 4, 1006 (1965). (5) P. A. Gartaganis and C. A. Winkler, Can. J . Chem., 34, 1457 (1956). (6) E.M.Levy and C. A. Winkler, ibid., 40, 686 (1962).
4067
NOTES
system before each experiment of 100-see duration. The reaction vessel was a straight, Pyrex-glass tube of 32-mm i d . , with i*. fixed reactant jet 17 cm below the discharge. The pressure in the system was 2 torr with a flow rate of molecular nitrogen of 190 X mole sec-‘. Nitrogen (Linde “bond dry”) was used after it had passed through a copper furnace a t 420’ to remove possible traces of oxygen. During experiments in the “unpoisoned” system, a liquid-air trap was used to remove traces of moisture from the gas. Commercial YO of 99% purity (Matheson Co.) was frozen at liquid-nitrogen temperature and freed from nitrogen by evacuation; S O 2 and N203 were removed by distillation of KO through a silica-gel column a t -78”. Anhydrous ammonia and sulfur dioxide (Matheson Co.) were used after three bulb-to-bulb distillations during which only the middle fraction was retained. Excess S H 3 or SO2 was trapped a t liquid-air temperature. A Kjeldahl distillation was used to follow destruction of YH3, and iodimetry was used to follow SO2destruction. Experiments were first made in an “unpoisoned” system, using either a condensed electrode or a microwave discharge. The active-nitrogen flow rates, estimated by the gas-phase NO “titration,”’ were 2.35 X and 1.83 X mole sec-l, respectively. S o destruction of SO2 was detectable. Neither was it possible to detect any destruction of NH3 by active nitrogen under these conditions. In a system (‘poisoned” with water vapor (-5 X mole sec-l) in the nitrogen stream through the microwave discharge, the active-nitrogen flow rate was increased only slightly (15%), and neither SO2 nor NH3 was detectably decomposed. When the nitrogen stream through the condensed discharge contained a similar relatively large amount of water vapor, the active-nitrogen flow rate was increased to 21 X mole sec-l. The SO2 reaction was then accompanied by the familiar blue glow, associated with the reaction, N 0 -+ S O * . The glow extended from the sulfur dioxide inlet jet to the cold trap. The amount of SO2 decomposed increased markedly with SO2 flow rate (Table I) and a pale yellow solid collected in the cold trap. When this was warmed to room temperature, an oily film remained. This residue was acidic and was probably sulfuric acid. Ammonia was also decomposed with the same experimental conditions (Table I), but the maximum extent of its decomposition never exceeded about one-fifth the N atom flow rate. When the water vapor supply was reduced to 0.1 X low6mole sec-l, with a concomitant decrease in the active-nitrogen flow rate, the decomposition of SO2 did not give rise to the blue glow, nor was the yellow after-
+
Table I:
Reactions of SO2 and NH3 with Active Nitrogen”
so2
so2
flow
reacted
NHa flow
9.0 9.5 14.0 17.0 47.5
7.5 9.2 11.8 13.2 16.2
3.8 3.9 5.4 5.5 11.3
NHs reacted
3.8 3.9 4.1 4.1
4.1
a System “poisoned’” with HzO vapor. N atom flow rate mole sec-1, by N O “titration”; ratio of this value to 21 x maximum HCN yield from CzHa reaction was 1.4. All quantities are (moles sec-1) x lo+.
glow extinguished, even a t high SO2 flow rates. Almost no yellow solid was collected. The data are shown in Table 11, together with corresponding results for the reaction of NH3 under the same conditions. (A blank experiment in the absence of ammonia indicated that no NH3 was formed by reaction of active nitrogen with the trace of H 2 0present.)
Table I1 : Reactions of SOz and NH3 with Active Nitrogen“ Activenitrogen flow
10.5 10.5 10.5 11.8 11.8
,902 flow
SO2 reacted
9.0 11.3 14.8 7.0 23.4
1.9 1.7 1.8 2.0 2.0
Activenitrogen flOW
7.9 10.5 11.8 13.5 13.5
N Ha
NHs flow
reacted
10.4 7.5 22.6 8.0 5.6
1.5 1.7 2.0 2.2 2.5
System “poisoned” with HzO vapor. ,411 quantities are N atom flow rate by NO “titration.” Ratio (moles sec-1) X of this value to maximum HCN yield from CzHa reaction was 1.4.
The relative behavior of SO2 and NH3, in their reactions with active nitrogen, was studied independently with a second apparatus, in which the active nitrogen was formed in a condensed discharge. The observations reported above were fully confirmed. It was also found that an increase of reaction temperature to 200” had no effect on the limiting extents to which either NH3 or SO2 reacted (cf. a similar earlier observation for the NHI reaction, ref 8). mole sec-I), With a small amount of H2 (-0.5 X instead of H 2 0 vapor, in the nitrogen flow before it entered the condensed discharge, the active-nitrogen (7) F. Kaufman and J. R. Kelso, J . Chem. Phys., 27,1209 (1957). (8) G. R. Freeman and C. A. Winkler, J . Phys. Chem., 59, 371 (1955).
Volume 70, Number 12 December 1966
4068
NOTES
flow rate was 13.3 X mole sec-I. (Under these conditions the active-nitrogen flow rate obtained with the microwave discharge was 2.43 X mole sec-l.) The SO2 reaction occurred with the formation of little solid, with the results shown in Table 111. Also included in the table are data for the NH3 reaction under similar experimental conditions.
Table In: Reactions of SO2 and NHa with Active Nitrogen“ so2 flow
6.8 10.4 11.3
18.8
SOP reacted
NHs flow
NHa reacted
1.3 1.1 1.6 1.4
11.3 16.2 20.0 22.5
1.1 1.3 1.2 1.2
System “poisoned” with H2. N atom flow rate 13.3 mole sec-1. All quantities are (moles sec-1) x lo4. a
x lo-*
It is evident from the data that ammonia and sulfur dioxide are decomposed to approximately the same extent by active nitrogen from a condensed discharge. Failure to decompose either SO2 or NH3 by active nitrogen from the microwave discharge might be due simply to the low N atom concentration produced. Alternatively, there might be a qualitative difference in the active nitrogen produced by the two methods. The relatively larger extents of SO2 and NH3 reactions in the system “poisoned” with water vapor, compared with the system “poisoned” with H:,, might also be due to a larger concentration of active species, although it is possible that it reflects some interference from oxygen atom reactions. The persistence of the yellow nitrogen afterglow during the SO2 and the NH3 reactions, even for very high reactant flow rates (much above the value corresponding to the maximum destruction of either SOz or NH3), indicates that these reactions involve a species other than N atoms in the 4S state. Previous work has indicated that an excited nitrogen molecule, either IS:, (32u+)9-11 or K:, (52,+),12 might be responsible for the ammonia reaction. It seems likely that the same species is also responsible for the SO:, reaction. Attempts to study the reaction of active nitrogen with mixtures of SO:, and NH3 were unsuccessful. When the two gases were mixed, a very rapid gas-phase reaction occurred and a yellow-brown water-soluble solid was formed. It was probably H:,NS0:,NH413 or (NH4)zS:,05.14(Sulfur dioxide is a strong electron acceptor and forms a variety of molecular complexes with electron donors.) The Journal of Physical Chemistry
Acknowledgment. Acknowledgment is gratefully made to the American Sulfur Institute for financial assistance during this investigation. ~~~
(9) A. N. Wright and C. A. Wikler, Can. J . Chem., 20, 5 (1962). (10) A. N. Wright, R. L. Nelson, and C. A. Winkler, ibid., 40, 5 (1962). (11) H. B. Dunford, J . Phys. Chem., 67, 258 (1963). (12) K. D. Bayes and G. B. Kistiakowsky, J . C h a . Phys., 32, 992 (1960). (13) K. A. Hofmann and U. R. Hofmann, “Anorganische Chemie,” 11th ed, Frieder, Vieweg, and Sohn, Braunschweig, Germany, 1945, p 176. (14) T. Hata and S. Kinumaki, Nature, 203, 1378 (1964).
Homogeneous Chemical Kinetics with the Rotating Disk Electrode by P. A. Malachesky, L. S. Marcoux, and R. N. Adams Department of Chemistry, University of Kansas, Lawrence, Kansas 66044 (Received June 19,1966)
Homogeneous chemical reactions coupled to electron transfers are of great importance in studies of organic electrode processes.’I2 The rotating disk electrode (RDE) and the rotating ring-disk electrode have been shown t,o be of value in investigations of electrode rea c t i o n ~ . ~However, ~* the capabilities of the RDE have not been fully exploited. Theoretical treatments, with experimental verific& tion, have appeared for the ECE electrolysis mech& chronoamperomenism (1) for chron~potentiometry,~ tryJ6cyclic v~ltarnmetry,~ cyclic chronopotentiometry,* and polar~graphy.~We wish to present an approximate treatment of the ECE mechanism for the R D E which appears applicable to first-order rate constants (1) T. Mizoguchi and R. N. Adams, J . Am. Chem. Soc., 84, 2058 (1962). (2) D. Hawley and R. N. Adams, J . Electroanal. Chem., 8, 163 (1964). (3) Z. Galus and R. N. Adams, J . Am. Chem. SOC.,84, 2061 (1962). (4) V. G. Levich, “Physicochemical Hydrodynamics,” Prentice-Hall, Inc,, Englewood Cliffs, N. J., 1962. (5) A. C. Testa and W. H. Reinmuth, J . A m . Chem. SOC.,83, 784 (1961). (6) G. S. Alberts and I. Shain, Anal. Chem., 35, 1859 (1963). (7) R. S. Nicholson and I. Shain, ibid., 37, 178 (1965). (8) H. B. Herman and A. J. Bard, J . Phys. Chem., 70, 396 (1966). (9) R. S. Nicholson, J. M. Wilson, and M. L. Olmstead, Anal. Chem., 38, 542 (1966).
