. NOTES
Dec., 1957
1665
agreement with the observed ethanol-water data than in the case of the aqueous data reported previously.6 Divergencies between theory and observation, however, are still apparent at higher concentrations of HC1 and particularly LiC1. Evidently ion-pair formation in the resin phase is still somewhat favored. I n 0.31 mole fraction methanol (Fig. 6) these divergencies diminish, and finally when the mole fraction of ethanol is increased to 0.40 vanish as anticipated (Fig. 7), although difficulties now appear in the region of the adsorption peak. The adsorption curves in 0.20 mole fraction acetone solutions (Fig. 6) are similar to but fall somewhat below the corresponding ethanol curves, suggesting, inasmuch as the dielectric constants of these two media are very similar, specific solvation effects. Figure 8 shows the variation of the adsorption of tracer zinc(I1) on Dowex 1-X8 chloride at 24.8 A 0.2" with the per cent. by weight of ethanol in ethanol-water solution. Log D increases as the dielectric constant decreases up to about 35 weight yo ethanol, presumably due to the increase of Kz',remains fairly constant in the region 35 to 65 weight % ethanol, and then at 65 weight yoethanol, the same concentration at which desorption of adsorbed zinc(I1) species into ethanol-water solution ceases,29 abruptly increases. This increase is accompanied by an inversion of the LiCl and HC1 selectivities, the greater adsorption now being from HC1 solution. I n the resin phase where competition for water of solvation is keenest, ion dehydration, resulting in an inversion of the order of effective ionic radii and greatly enhanced ionpair formation, is occurring. Such abrupt transitions in the water content of the solvation spheres of cations in ethanol-water mixtures have also been observed in the case of iron(III),80although
at lower water concentrations. Adsorbed zinc(I1) complex anions do not desorb into ethanolic media above 65 weight Yobecause ethanolic solvation, unlike hydration, is insufficiently strong to disrupt the ionic associations in the resin phase. The present results and conclusions are similar to those of Gable and Strobe16 who explained the enhanced selectivity of cation exchange in methanol in terms of ion solvation and ion-pair formation. Solvation is also known t o play an important role in solvent extraction where, commonly, the acid of the metal halide complex extracts with a definite number of waters of h y d r a t i ~ n . ~ ~ , ~ ~ Conclusion On the basis of the results reported in this and the previous papers of the present seriessJ9 we conclude that the adsorption of metals on anion exchange resins from complexing media is an equivalent exchange phenomenon involving anionic complexes of the metal, similar to cation adsorption on cation-exchange resins or simple anions on anionexchange resins, complicated by (a) the as yet poorly understood physical chemistry of concentrated solutions of electrolytes, (b) the dependence of the concentrations. of absorbing species on the concentration of complexing ligand and (c) ionic association in the resin phase, and possibly in the external solution phase as well, of the anionic metal complexes with the cation of the supporting electrolyte. Acknowledgment.-The authors wish to express their gratitude for the valuable assistance of Professors C. D. Coryell and J. w. Irvine, Jr., o f ' M.I.T. and R. N. Diamond of Cornel1 University and for facilities and funds supplied by the United States Atomic Energy Commission through the Division of Sponsored Research of the Massachusetts Institute of Technology.
(29) R. A. Horne, R. H. Holm and M. D. Meyers, THISJOURNAL, 61,1655 (1957). (30) R. A. Horne, Ph.D. Thesis, Columbia University, 1955.
(31) H.Irving and F. J. C. Rossotti, J . Chem. SOC.,1938 (1955). (32) P.C.Yates, R. Laran, R. E. Williama and J. E. Moore, J . A m . Chem. Soc.. 71, 2212 (1953).
NOTES THE PHOTOCHEMlCAL OXIDATION OF ZINC AND CADMIUM SULFIDE BY M. CLAREMARKHAM, JOANBARRY,MARION IAVA AND
JANET HADDAD
Department of Chemistry, Saint Joseph CoUeoe, West Hartford, Conn. Received July 10, 1967
I n the past few years considerable attention has been focused on the effects of ultraviolet light on photoconducting oxides and sulfides.' Many data have accumulated to show that zinc oxide is a photocatalyst for the formation of hydrogen peroxide from molecular oxygen, but that there must be oxidizable organic substances present if the hydrogen peroxide is to accumulate to any ap'
(1) M. C. Markham, J . Chsm. ad., 89, 640 (1966).
preciable e ~ t e n t . ~ The J net reaction seems to be the oxidation of the organic material and the reduction of molecular oxygen, part of which is found as hydrogen peroxide, and the remainder in the oxidized organic product^.^ Recently two articles have appeared in THIS JOURNALwith the conclusion that some photoactive metallic sulfides are catalysts for the formation of hydrogen peroxide from oxygen and water.6p6 Our finding of considerable amounts of (2) T. R. Rubin, J. G . Calvert, G . T. Rankin and W. MacNevin, J . A m . Chem. SOC.,71, 2850 (1953). (3) M. C. Markham and K. J. Laidler, THISJOURNAL, 67, 363 (1953). (4) C. B. Vail, J. P. Holmquint and L. White, Jr.. J . Am. Cham. ~ o c . ,76, ea4 (1954).
1666
NOTES.
Vol. 61
both metal ions and sulfate ions in the filtrates, dark. However the most reactive samples were when suspensions of zinc and cadmium sulfides those freshly precipitated under the conditions were exposed t o air and near ultraviolet light, described above. made more detailed examination of these reactions Table I shows the comparison of several of the seem advisable. samples tested. Experimental Procedure Preparation of Cadmium Sulfide Samples.-The most active form of cadmium sulfide was prepared by the method of Stephens, et al.,b by precipitation from cadmium oxide dissolved in excess nitric acid, with sodium sulfide solution as the source of sulfide ion. This treatment yielded a fine bright yellow powder which was oxidized rapidly in the light and turned white after several hours irradiation in water or dilute phenol solutions (0.001 M ) , and turned brown in more concentrated phenol solutions (0.10 M). As much as 50% of an initial 0.2-g. sample went into solution during 10 hours of irradiation with near ultraviolet and violet light. Oxidation and solution in the dark are very slight. We have not been able to obtain this particular crystal form a second time. Newer samples of cadmium sulfide supposedly prepared in the same way have a more orange color, and are not very photo-reactive. Our most easily reproducible photo-active samples of cadmium sulfide are prepared by dissolving 0.1 mole of cadmium chloride in 500 ml. of water to which has been added 0.3 mole of hydrochloric acid, and adding slowly, with stirring, 0.1 mole of sodium sulfide dissolved in 500 ml. of water. After precipitation the total volume is 1000 ml., and 0.2 mole of hydrochloric acid has been used to neutralize the hydroxide formed on hydrolysis of sodium sulfide. The final concentration of hydrochloric acid is about 0.1 M . The precipitate of yellow-orange cadmium sulfide is filtered a t once, washed several times with water, and allowed to undergo preliminary drying in the dark, then dried further in a vacuum desiccator over fused sodium sulfide, and finally over fused sodium hydroxide to remove adsorbed hydrogen sulfide. Commercial samples of deep orange cadmium sulfide were also tested and found to react only very slowly. Preparation of Zinc Sulfide Samples.-The zinc sulfide found to be most active photochemically was prepared by dissolving 0.1 mole of zinc nitrate in 500 ml. of water and slowly adding 0.1 mole of sodium sulfide in 500 ml. of water. This zinc sulfide precipitate, which darkens on exposure to air and li ht, was dried in a manner similar to the cadmium sullde described above. A similar preparation carried out in more concentrated solutions, yielded a very h e white zinc sulfide, easily becoming colloidal on suspension in water. A commercial sample of zinc sulfide, supplied by the New Jersey Zinc Company, was also found to be hotochemically oxidized, but not so rapidly as the freshPy precipitated samples. This zinc sulfide does not darken in the light, but produced more hydrogen peroxide. Apparatus.-The light source was a Hanovia high pressure quartz mercury arc, used in all the experiments with a Pyrex envelope to retain all radiation below about 3300 A. The mercury arc was in a vertical position and about 6 in. from the reaction vessel, which consisted o!k%) ml. Pyrex test-tube containing 0.2 g. of the metal sulfide in 25 ml. of distilled water. The test-tube was surrounded by a Pyrex jacket through which distilled water from a constant temperature bath could be circulated at any desired temperature. Stirring was effected and oxygen supplied by rapidly entraining air previously filtered through distilled water. Identification of Products.-Hydrogen peroxide was analyzed by oxidation of iodide ion and titration of the liberated iodine with standardized sodium thiosulfate, using starch as the indicator. Zinc and cadmium ions were quantitatively estimated by means of a Patwin, research model, polarograph. Sulfate was determined gravimetrically as barium sulfate.
Results All forms of zinc and cadmium sulfide tested were found to be oxidized photochemically much more rapidly than the thermal oxidation in the (5) R. E. Stephens. B. Ke and D. Trivich, THISJOVRNAL, 89, 906 1955). (6) L. I. Groesweiner, ibid., 69, 742 (1955).
TABLE I COMPARISON OF THE PHOTOCHEMICAL OXIDATIONOF VARIOUSSAMPLES OF ZINCA N D CADMIUM SULFIDES 108 Mmoles of M +X+
Conditions of irradiation
Type of sulfide
produced
1. CdS, yellow-orange 1 hour, 25O, light from CdC12 CdS, yellow-orange 1hour, 25", dark from CdCl2 2. CdS, yellow-orange 1 hour, 25", from Cd(N0& light 3. CdS, yellow-orange 1 hour, 25", from Cd(Ac)2 light 4. CdS, bright yellow, 0 . 5 hour, 25O, light original ppt. from Cd( NO& 5. CdS, deep orange, 1 hour, 25", commercial Relight ' agent grade 6. ZnS, white, N. J. 1 hour, 25", Zinc Co. light 7. ZnS, freshly pptd. 1 hour, 25", from Zn( NO& light (darkens on exposure)
21.0
6.0 1.6 11.0 82.5
10.5
29.6 36.0
Table I1 shows the relative rates at different temperatures of photo-oxidation of the cadmium sulfide precipitated from the chloride, and Table I11 of the zinc sulfide precipitated from the nitrate in dilute solution. TABLE I1 COMPARISON OF RATESOF OXIDATION OF CdS, SAMPLE 1 IN TABLE I, PPT.FROM CdC12, AT VARIOUSTEMPERATURES Conditions of irradiation
1 hour, 25", light 1hour, 25", dark 1 hour, 35", light 1 hour, 35", dark 1 hour, 45O, light
Cd++ (moles ,
X 108)
21.0
6.0 23.0
7.0 24.0
60r-(mmoles X 108)
Hi02 (moles X 109
18.6
5
None detectable 26.0
None detectable 3.1
None detectable 20.0
None detectable 1.1
Tables I V and V show the relative amounts of products formed progressively with time. Table VI shows some data on the extremely reactive cadmium sulfide which we have not been able to duplicate. I n this table there are also included some tests on the effect of adding organic material. It was considered possible that occluded or adsorbed nitrate ion might be responsible for the smaller activity of the cadmium sulfide samples. Therefore a sample of cadmium sulfide precipitated from the chloride was irradiated in 0.10 M NaNOs,
Dec., 1957
1667
NOTES
TAnLE 111 loidal. This cadmium sulfide was peptized into COMPARISON OF THE RATESOF OXIDATION OF ZnS, SAMPLE particles a t least fine enough to go through a double layer of Whatman #42 filter paper without suction. 7 IN TABLE I, AT VARIOUS TEMPERATURES Zn++
SOP-(mmoles X 103)
H2Oa (mmoles X 103)
36.0
35.7
7.0
8.0
None detectable None detectable None detectable None detectable
Conditions of irradiation
(mmoles X 103)
1 hour, 25O,
light 1 hour,25", dark 1 hour, 45", light 1 hour, 40", dark
Only a very little of this sample dissolved in an hour's irradiation; the function of the light seemed to be the discharge (or counter-charge) of the surface charge of the colloid. Preliminary runs with anti.mony sulfide indicate a photochemical oxidation in this case also.
Discussion and Conclusions Analysis of the data makes it clear that the zinc and cadmium sulfides most active in forming hydro5.0 None detectgen peroxide on irradiation in aqueous suspensions able are also going into solution as sulfates much more TABLE IV rapidly than in the dark. An exception is the PHOTO-OXIDATION PRODUCTS OF CdS (SAMPLE 1, TABLE I ) zinc sulfide suspension which turns dark under these conditions. No hydrogen peroxide is devers'us TIMEOF IRRADIATION Conditions of Cd++ sod-HiOa tected, but sulfate is formed as usual. Apparently (mmoles (mmoles irradiation, (mmoles in this case the zinc is reduced instead of the 250 x 103) x 103) x 103) oxygen from the air. Either crystal structure or 15 min., light 9 . 8 Very slight, 3.0 particle size is the important factor-possibly both. but definite Two possibilities must be considered. (1) The 5.0 60 min., light 21.0 18.6 reaction effected by the light may be the formation 120 min., light 34.4 24.8 7.5 of hydrogen peroxide by way of free radicals from 180 min., light 50.0 52.5 8.8 oxygen and water as Grossweiner supposed in the 10 hours, light 104.0 117.0 23.7 case of mercuric sulfide. He was unable to detect any sulfate or increase of metal ion in his experiTABLE V PHOTO-OXIDATION PRODUCTS OF ZnS (SAMPLE 7, TABLE I) ments. The origin of the sulfate ion in the present case would then be the oxidation of sulfide by versus TIMEOF OXIDATION hydrogen peroxide. Zn++ S04-Conditions of (mrnoles (mmoles From the data above there is evidence against irradiation, 2t0 X 108) X 103) HsOa this interpretation. From Table I1 we see that at 1 hour, light 30.0 35.7 None detectable higher temperatures there is less hydrogen peroxide, 88.2 93.0 None detectable 4 . 5 hours, light but the amount of sulfate does not change appreci10.0 10.7 None detectable 4 . 5 hours, dark ably. If the oxidation of sulfide to sulfate by hydrogen peroxide were a thermal reaction we TABLE VI should expect to find more sulfate a t the higher PHOTO-OXIDATION PRODUCTS OF MOST REACTIVECdS temperature where more hydrogen peroxide has (YELLOW, SAMPLE 4, TABLE I) apparently undergone thermal decomposition. Conditions of Cd++ so&-HzOn irradiation, (mmoles (mmoles (mmoles The possibility that the photochemical decompo25' x 10') x 103) x 103) sition of hydrogen peroxide is responsible for the Suspended in HzO oxidation seems to be excluded by the fact that 30 rnin., light 82.5 82.5 1.0 hydrogen peroxide does not absorb at this wave 5 hours, light 250 ... 12.0 length . 0.001 M Phenol (2) The other possible explanation is that the 87.8 30 min., light 87.8 15.0 hydrogen peroxide is formed by a reaction that 450 5 hours, light 380 35.0 would amount to a reoxidation of reduced metal, 0.010 M Phenol and that the effect of the light is simply an electron 100.0 30 min., light 129.0 36.8 transfer from the sulfide ion to a neighboring zinc 5 hours, light 210.0 ion in the crystal. 196.0 90.0 37.0
0.100 M Phenol 5 hours, light
40.0
ZnS +Zn+lS-l 170.5
199.5
52.5
but there was practically no change in concentrations of the various products. Similarly, the addition of chloride ion to the cadmium sulfide precipitated from the nitrate effects no increase of activity. One of the cadmium sulfide samples precipitated from the nitrate has the peculiar property of being peptized in water when suspended in the dark, but the colloidal filtrate could be irradiated and the particles would coagulate again. Conversely a sample could be irradiated in the usual way and filter clear. The residue could then be suspended in water in the dark and the filtrate would be col-
Oxygen might then be able to remove the electron from the reduced zinc to form a radical type peroxide ion which evidently can be stabilized in water to form hydrogen peroxide, while more molecular oxygen would be consumed to complete the oxidation of the sulfate. I n the case of the zinc sulfide that darkens in the light, either surface conditions or crystal structure may prevent the rapid re-oxidation of the zinc. There would thus be less peroxide formed, but the sulfate ion would still be produced by independent oxidation as appears consistent With the data of Table 111, where zinc and sulfate ions are produced in roughly equimolar amounts in the apparent absence of
1668
NOTES
hydrogen peroxide. The darkening of this form of zinc sulfide in the light is reversible in the dark. The second explanation is consistent with the observation that the stoichiometric relationships among the products, in the case of cadmium sulfide in Table IV, remain fairly constant over an extended period of time. If the oxidation of sulfide were effected principally by the hydrogen peroxide one would expect an acceleration of the appearance of sulfate ion with time as the hydrogen peroxide concentration increases in the solution. Further work on the stoichiometry of the oxygen uptake is now in progress in our laboratories, to help decide the correct interpretation of these surface oxidations of photoconducting sulfides. HEATS OF COMBUSTION. VI. THE HEATS OF COMBUSTION OF SOME AMINO ACIDS BY TOSHIO TSUZUEI A N D HERSCHEL HUNT Department of Chemistry and Purdus Research Foundation, Purdue University, Lafayette, Indiana Received June $4, 1967
The heats of combustion of L-glutamine, L-glutamic acid, L-valine and L-leucine, have been determined by means of a non-adiabatic calorimeter. The method is essentially the same as that described by previous workers in this except that thermistors are used as temperature-sensing elements rather than thermocouples.
Vol. 61 TABLE I NITROQENCONTENTS OF AMINOACID SAMPLES Amino acid
GGlutamine >Glutamic acid >Valine >Leucine
Nitrogen content. % Theor. Found
19.17 9.52 11.96 10.68
19.01 9.56 12.13 10.63
The heats of combustion of the amino acids at constant volume and a t 25" for the reaction producing gaseous carbon dioxide, liquid water and gaseous nitrogen, are given in Table 11. The standard deviations were calculated in accordance with recommendations by Rossini and Deminglband the atomic weights used in the computations are those of the 1953 revision: 0 = 16, C = 12.011, H = 1.0080, N = 14.008. TABLE I1 HEATS OF COMBUSTION OF AMINOACIDS Amino acid
>Glutamine (s), C~HIOO~NS LGlutamic acid (s), C5Ho04N L-Valine (s), CsHnOJV' >Leucine (s), CaHlaOzN
Heat of oombustion, kcal./mole
614.80 & 0.15 537.01 f .I8 697.93 & .I3 853.08 f .06
The heat of combustion of glutamic acid was previously reported by Fischer and Wrede,6 by Emery and Benedict,' and by Huffman, Ellis and Wredeg determined the heat of combustion of valine. The heat of combustion of leucine was determined by Stohmann and Langbeinlo and by Fischer and Wrede.6 The authors wish to express their appreciation t o National Science Foundation for sponsoring thls research and to Drs. J. W. Amy and P. R. Marshall for their work in the major part of the apparatus modification.
The apparatus, which is described elsewhere*l4has been modified so that two thermistors (Western Electric Type 14B, 2240 and 2225 ohms at room tem erature) with a 60turn 500-ohm Resomax otentiometer (%ink Aviation Company, Model 400) and a%rown null indicator (MinneapolisHoneywell Company Model 104-WIG) in a Wheatstone bridge circuit can be used to measure the temperatures of the calorimeter and of the jacket, respectively. The thermistors were calibrated against a platinum resistance thermometer, which had been calibrated previously, and it was possible to estimate the temperature reading to less than (5) F. D. Rossini and W. E. Deming, J . Wash. Acad. Sci., 29, 416 0.0002 O . The calorimeter was calibrated with National Bureau of (1939). (6) E. Fischer and F. Wrede, Akad. Wiss. B e d . Sitzungsber., 687 Standards benzoic acid, Standard Sample 398;. I n order to promote ignition of the amino acid samples a small pellet of (1904). (7) A. G . Emery and F. G. Benedict, A m . J . Physiol., 28,301 (1911). the standard benzoic acid was placed in the stainless steel (8) H. M. Huffman, E. L. Ellis and 8. W. Fox, J . A m . Chsm. Soc., combustion capsule at an angle against one edge of the amino acid pellet. The weights of the amino acid sample 58, 1728 (1936). (9) F. Wrede. 2. physik. Chem., 76, 81 (1910). and of the benzoic acid were adjusted so that the amount (10) F. Stohmann and H. Langbein, J . prakt. Chem., [Z] 44, 383 of heat liberated would be comparable to the amount liberated in the calibration experiments. Exactly 1.1 ml. of (1891). water was placed in a Parr double-valved oxygen bomb of 0.358-1. capacity and the bomb was fdled with Linde U.S.P. oxygen to a pressure of 30 atm. absolute at 25". REACTION TIME DISTRIBUTION I N Proper corrections were made for the heat of stirring, the LAMINAR FLOW KINETIC heat exchanged between the calorimeter and the jacket, the ignition energy, and for the heat of formation of nitric acid. MEASUREMENTS The calorimeter system with 2800 g. of water had a water BYKENNETH A. WILDE equivalent of 3285 g., and the heat evolved in each determination was corrected to the value which would have been Rohm and Haas Company, Redslons Arsenal Resparch Division, liberated in a system of exactly that water equivalent. HunlsvtEEe, Alabama All the samples of amino acids were prepared by Dr. Receiced Mail 87, 1967 J. P. Greenstein of National Cancer Institute, Bethesda, Maryland, and they were said to be better than 99.9% pure. Since the Reynolds number in laboratory flow Each of the these amino acid samples was recrystallized once from water and ethanol, followed by complete drying in Z ~ ~ C U Osystems is usually well in the laminar range, one in 100" (50" for L-glutamic acid). Nitrogen contents of would expect to obtain the usual parabolic velocity the recrystallized samples were determined by the micro- profile across a circular cross-section. This profile Kjeldahl method and compared with the theoretical values \vi11 obviously result in a pronounced spreading of as shown in Table I. (1) A. J. Miller and H. Hunt, THIS JOURNAL, 49, 20 (1945). (2) M. V. Sullivan and H. Hunt, ibid., 68, 497 (1949). (3) G.M. Kibler and H. Hunt, ibid., 68, 955 (1949). (4) C. B. M i l a and € Hunt, I . ibid , 45, 1846 (1941).
reaction times, but the effect seems to have escaped attention by workers in the field. Bosworthl has calculated a distribution function for reaction times (1) R, C . L. Boaworth, Phil. Mag., 89, 847 (1948).
