2254 by Hirano.* Interestingly enough, sodium carbonate forms both a heptahydrate and a decahydrate with the two transition regions near 30-34.5" for the crystalline solids. FujitaZ0has reported an anomaly in the viscosity of sodium carbonate solution (1% concentration) near 31" and (less pronounced) near 35". We do not have any simple explanation for these divergent results. It might be speculated, however, that at least for the viscosity measurements, anomalies might be caused by thin layers of solution a t the air-glass-solution interface (in the viscometer) becoming supersaturated (owing to local evaporation) and thus causing spurious effects, owing to decomposition of the crystalline, solid hydrates affecting the flow regime. Acknowledgment. The authors wish to thank Miss Irene Cooper for assistance with the viscosity measurements. The authors also wish to acknowledge the support of the office of Saline Water for this study. (20) K. Fujita, Bull. Chem. SOC.Jap., 32, 1004 (1959).
A n Attempted Kinetic Study of Chemiluminescent
Electron-Transfer Reactions1&
by Robert Livingstonlb and Horst R. Leonhardtlt Department of Chemistry, University of Minnesota, Minneapolb, Minnesota (Received November SO, 1967)
It was demonstrated r e ~ e n t l y ~that - ~ the anaerobic oxidation of radical anions of aromatic hydrocarbons is, in many cases, accompanied by strong chemiluminescence. These reactions are rapid, and there have been no reports published of direct studies of their kinetics. I n the hope of obtaining some kinetic information, we have investigated several of these reactions using a fast, straight-flow reactor.
NOTES Radical anions of 9,lO-diphenylanthracene were dissolved in both solvents to serve as scavengers and simultaneously as indicators for oxygen and water. Solutions of radical anions of perylene and 9,10diphenylanthracene in dimethoxymethane were prepared according to HoijtinP by the reaction of elemental sodium with the hydrocarbons. The alkali metal was obtained by pyrolytic decomposition of sodium azide. I n the case of dimethylformamide, the radicals were prepared in dimethoxymethane and after evaporating the solvent dissolved in dimethylformamide. Apparatus. The straight-flow reactor was of conventional design. The reagent solutions were confined in 50-ml gas-tight syringes (Hamilton Co.). Their stainless steel plungers were equipped with Teflon tips and were propelled at identical speeds by a single motor-driven rack and pinion. The glass tubes which were sealed to the syringes were equipped with Teflonglass valves so arranged that the system could be evacuated or flushed with dry argon, the syringes filled with the reagent solutions from attached reservoirs, or the solutions forced through the mixer into the capillary reaction tube. Large differences in the thermal coefficient of expansion between Pyrex and Teflon accentuated by a transition point in the thermal expansion of Teflon at approximately 24" made it necessary to keep the syringes and valves near room temperature. I n low-temperature experiments the temperature of the reacting mixture was controlled by means of thin-walled, glass-capillary helical heat exchangers situated in the leads close to the mixing chamber. The effectiveness of the temperature control was tested with a thermocouple placed in the outlet of the helices. Mixing was achieved in the lower part of the vertical reaction tube which was made of 0.5-mm bore capillary glass tubing, and which was sealed to the leads as described by Chance.8 The mixing times were a function of solvent, flow rate, and temperature. Under the conditions employed, mixing times ranged from 5 7 X 10-4 sec (dimethoxymethane, 0.42 ml/sec, sec (dimethoxyroom temperature) to 51 X methane, 0.30 ml/sec, -75").
Experimental Methods and Materials Materials. Perylene was purified chromatographically according to Sangster and 1rvine;b 9,lO-diphenyl(1) (a) This work was supported by the U. S. Army Research Office (Durham). (b) Department of Chemistry, San Diego State College, anthracene, sodium azide, and benzoyl peroxide were San Diego, Calif. recrystallized; tetracyanoethylene was purified by (2) D. M. Hercules, Science, 145, 808 (1964). ~ E. A. Chandross and F. I. Sonntag, J . A m . Chem. Soc., 86,3179 sublimation; and HgC12, SbCI,, SnCL, and K ~ z O (3) (all Baker Analyzed grade) were used without further (1964). purification. (4) K.S. V. Santhanam and A. J. Bard, ibid., 87, 139 (1965). (5) R. C. Sangster and J. W. Irvine, J . Chem. Phys., 24, 697 (1966). Dimethoxymethane was refluxed over calcium hy(6) P. Balk, G.J. Hoijtink, and J. W. H. Schreurs, Rec. Trav. Chim., dride and was distilled. Oxygen and the last traces of 76, 813 (1957). water were removed by degassing the predried liquid (7)L. H. Gillespie, D. 0.Saxton, and F. M. Chapman, Machdne over lithium aluminum hydride using standard freezeDesign, 32, 156 (1960). thaw cycling. Dimethylformamide was distilled under (8) F. J. W. Roughton and B. Chance, "Technique of Organic Chemistry," 8. L. Friess, E. S. Lewis, and A. Weissberger, Ed., vacuum, degassed by repeated freezing and thawing, Vol. 8, Part 2, Interscience Publishers, Inc., New York, N. Y.,1963, and treated with molecular sieve (30-40 mesh, Type 4A). p 710. The Journal of Physical Chemistry
NOTES
2255
Table I [anion], M
Oxidant
HgCb (CN)ZC=C (CN)2
1.0 2.5
x x
10-4
10-4
3.2 X 5 . 2 x 10-4
The relative intensity of chemiluminescence was measured by conducting a fixed fraction of the light emitted from the mixer through a light pipe to an RCA 1P21 photomultiplier. The output of the photomultiplier was fed to a Keithly Model 410 recording micro-microammeter,
Results and Conclusions The systems studied were the perylene (singly charged) anion in dimethoxymethane with HgC12, SnC14,SbCI,, benzoyl peroxide, and tetracyanoethylene; the perylene anion with benzoyl peroxide in dimethylformamide ; and the 9,lO-diphenylanthracene anion with HgClz in dimethoxymethane. The concentrations M. to 5 X of the reactants ranged from 1 X I n all these systems bright chemiluminescence was observed which visually appeared to be identical with the fluorescence of perylene and diphenylanthracene, respectively, in agreement, with measurements of chemiluminescence spectra in similar systems by Weller and Zachariasse,g and by Chandross and Sonntag.'" The light was produced within a very short zone which strongly suggests that mixing was the rate-limiting process. Therefore, only upper limits could be obtained for the half-lives of the reactions, which were as follows: in dimethoxymethane, 7 X loU4sec a t room temperature, 2 X 10u3sec at 0", and 1 X 10-2 sec a t -75" ;in dimethylformamide, 3.5 X lod3sec at room temperature. If we assume that the first kinetic step is a bimolecular reaction between anion and oxidant, these upper limits of mixing times and the (lowest) concentrations of reagents may be used to calculate limiting, minimum values for the second-order rate constants. With this assumption the following results, which are independent of the chemical nature of the reactants, are obtained: in dimethoxymethane at room temperature and -75", 5 X 10' and 3 X loe M-' sec-I, respectively; and in dimethylformamide at room temperature, 9 X lo6 M-' sec-l. Upper limits for the rate constants can be obtained from the theory of diff usion-controlled reactions. Using the Debye relation, k = 8RT/3q,11 one gets the following order-of-magnitude values : in dimethoxymethane k,, 2 X 10'0 M-1 sec-I a t room temperature, 1.5 X 10'" M-1 sec-' a t 0", and 2 X lo9 -M-l sec-' at -75"; in dimethylformamide, at room temperature, k,,, = 1 X 1010 M-I sec-l. Values for the viscosity q were either taken directly from ref 12 and 13, or calculated according to the relation q(T) = (Noh/V',) exp(3.8Tb/T),14 where N O is
-
1.00 0.80
0.02 0.007
50
114
Avogadro's number, h is Planck's constant, V , is the molar volume, and Tb is the boiling point of the compound. Comparing maximum and minimum values of the second-order rate constants, it appears likely that the reactions are diffusionally limited. A similar but unpublished conclusion for the system rubrenedimethylformamide has been attributed to Lansbury, Hercules, and Roe.15 The results of a few measurements of the relative intensity of chemiluminescence of the reactions of perylene anion with HgClz and with tetracyanoethylene in dimethoxymethane are given in Table I. The intensities, in arbitrary units, were normalized to the HgClz reaction at room temperature. The actual concentrations of perylene anion in the mixing chamber were probably somewhat lower than the listed values, since some of the reagent was always lost by reaction with water adsorbed on the walls of the "dried" glass tubing, etc. The reduction of intensity which is observed at the lower temperature is due in part to the increase in length of the mixing zone which causes the fraction of light collected by the light pipe to drop, and to the decreased rate of flow entering the reaction tube. However, the two effects cannot account entirely for the intensity change. Lengthening of the mixing zone reduces the intensity by less than a factor of 5, and the flow rates change only by a factor of 1.4; Le., the intensity at -75" should be reduced by less than a factor of 7 if there were no other effects. The last column of Table I shows that the experimentally observed reduction is markedly greater. This appears to indicate a decrease in quantum yield with decreasing temperature for the two systems studied. Unlike the reactions with the other oxidants, the oxidation of the perylene anion by peroxydisulfate ion in dimethylformamide is relatively slow. Its half-life time depends on concentration and temperature but is (9) A. Weller and K. Zachariasse, J . Chem. Phys., 46, 4984 (1967). (10) E. A. Chandross and F. I. Sonntag, J. Am. Chem. SOC.,8 8 , 1089 (1966). (11) P. Debye, Trans. Electrochem. SOC.,8 2 , 266 (1942). (12) "Beilstein's Handbook of Organic Chemistry," Vol. I, 2nd ed, Springer, Berlin, 1941, p 638. (13) R. A. Robinson and R. H. Stokes, "Electrolyte Solutions," Butterworth and Co., Ltd., London, 1959, p 317. (14) R. B. Bird, W. A. Stuart, and E. N. Lightfoot, "Transport Phenomena," John Wiley and Sons, Inc., New York, N. Y., 1966, p 29. (15) 8. W. Feldberg, J . Am. Chem. Soc., 88, 390 (1966); see ref 12. Volume 'YW,Number 6 June 1968
NOTES
2256 of the order of magnitude 1 sec. The kinetics of this reaction are being investigated with the aid of a slower straight-flow reactor.
Ion Diffusion into Fused Silica from Molten Salts
by Kurt H . Stern Institute for Basic Standards, National Bureau of Standards, Washington, D. C. ZOZSg (Received December 14, 1967)
The first quantitative study of ion diffusion into glass was reported by Schulze more than 50 years ago.' He immersed a soda glass in molten AgN03and analyzed both phases. His results were consistent with 1 : l exchange of sodium ions in the glass for silver ions in the melt. Since Schulze's time, many diffusion studies have been reported2 which support the model of silicate glasses as cation exchangers. Since the anionic silicate network excludes foreign anions, exchange of (monovalent) cations must be 1:l to preserve electrical neutrality. All of these reported studies have been carried out with glasses containing relatively high concentrations of alkali metal cations (most frequently sodium), usually 10-20 wt % of metal oxide. Consequently, such studies were carried out at temperatures lower than 500°, Le., below the softening point of the glasses. I n most instances diffusion profiles in the glass as a function of melt and glass composition, temperature, and time were measured, but 1: 1 exchange was assumed. I n order to aid in the interpretation of electrical conductivity data of Vycor and fused silica immersed in molten salts3 and of membrane potential studies, it seemed of interest to investigate the ion-exchange problem for fused silica. Since commercially available grades contain ionic impurities only in the parts per million-parts per billion range, weight-change techniques are excluded. However, activation analysis and spectroscopic methods are adequate for the measurement of glass composition. Experiments are reported for two temperatures, 570 and 890". Only silver halides were used at the lower temperature, alkali halides and AgCl at the higher. The glass was General Electric Type 204,4which is prepared by vacuum fusion of quartz, and has the following nominalmaximum impurities5 (in ppm) : LizO,2; Na20, 7; KzO, 6; MgO, 4; CaO, 23; A1203,65; Ti02, 5 ; Fe203,9. For the alkali metal ions, these figures correspond to micromole fractions of 8,13,and 8, respectively. Thus the exchange of all the alkali metal ions in a 1-g piece for silver ions would result in a weight changeof -50 fig. Slices of 13-mm 0.d. tubing were used in the experiTh,e Journal of Physical Chemistry
ments. They were suspended from Vycor hooks in the appropriate molten salt, contained in A1203crucibles, Experiments were carried out in a helium-filled drybox with a maximum 100 ppm of H2O vapor. Before analysis samples were freed from adherent salt by immersion in H20, dilute NH40H, and dilute HF. Extreme care was taken to avoid surface contamination by handling. A preliminary weight-change experiment in AgCl and AgCl with 1-2 mol % additions of NaCl and KCl yielded weight changes of -1 mg, both positive and negative, after immersion up to 60 hr at 570". With 2% LiCl in AgC1, a weight loss of 1.3 mg (0.1 wt %) after 20 hr was accompanied by surface pitting. These observations may be accounted for by reaction of the glass with LiCl to form volatile SiC14 and/or with the impurity Liz0 to form soluble lithium silicates. I n order to investigate Na+-Ag+ exchange quantitatively, 1.4-g slices from the same tube were immersed in AgCl and AgBr for 48 hr a t 590". The sample immersed in AgCl was etched and had lost 3 mg. The sample in AgBr had changed neither in appearance nor in weight, The above samples were analyzed by activation analysis as follows : The radioactivation products chosen for the analyses were silver-110 (hlZ = 24 sec), chlorine-37 (LIZ = 37 min), and bromine-82 (& = 36 hr). Sodium was determined at the same time using the activation product sodium-24 ( t l l z = 15 hr). The silver analyses required a 30-sec irradiation in the pneumatic tube at the Naval Research Laboratory reactor (thermal neutron flux = 8 X 10" neutrons cm-2 sec-l), while the other three elements were all determined from the same 5-min irradiation in the pneumatic tube. The various y rays produced from the activation products were resolved by a NaI(T1) scintillation detector in conjunction with a 400-channel pulse-height analyzer. Each sample was counted long enough to obtain good counting statistics. Half-lives were followed to verify identification. The results are shown in Table I. The constancy of Ag Na is particularly noticeable. It shows that the Ag+-Na+ exchange is 1 : l . It is surprising, however, that K + and Li+, which are present in concentrations comparable to Na+, apparently do not exchange. If mobility were the critical factor, one would expect the faster3 Li+ to exchange preferentially; if the cations occupy different sized
+
(1) G. Schulze, Ann. Phya., 40, 336 (1913). (2) K.H.Stern, Chem. Rev., 66,355 (1966). (3) K.H. Stern, J. EZectrochem. SOC.,112, 208 (1966). (4) Certain commercial materials are identified in this paper in order to adequately specify the experimental procedure. In no case does such identification imply recommendation or endorsement by the National Bureau of Standards. (6) General Electric Fused-Quartz Catalog, 1964.
