NOTES
1444
Experimental The freezing point determinations were carried out in an apparatus similar to that described by RlcMullan and Corbett.? A double walled cell, containing 10 ml. of the investigated solution, was placed in a cooling bath of ethanol and solid carbon dioxide. The solution was stirred by a magnetic stirrer. A thermistor (Standard Telephones & Cables, Ltd. Type F 2 3 ) was used for temperature measurements. A Wheatstone bridge (100,000 ohm 0. Wolff, Berlin), a Leeds & Northrup 2430D Galvanometer and a I .5 v. dry bat1,ery were used for resistaqce measurements. The thermistoI was calibrated with an ethanol filled thermometer in the range -50 to -62". At the freezing point of the pure solvent ( -59.7"), this thermistor had a resistance of 72,520 oh?: and the change with temperature was 413 ohms per 0.1 Since the greatest depression of the freezing point was 0.7", At was assumed to be proportional to AE; the error caused by this assumption was always less than 0.8%, and was neglected. The thermistor possessed remarkable stability with no change of more than 8 ohms (0.002") occurring during a two month period. This stability might be due to the veIy low currents passed through it (0.02 mamp.). Temperature measurements had an error of less than 0.001', but the reproducibility of the freezing point measurements (read from cooling curves) was lower because of changes in the degree of supercooling, rate of stirring, etc. A t least three determinations of the freezing point which did not differ by more than 0.004" were used for the calculation of the freezing point of each solution. The rate of cooling was 0.2' per minutc (prior to precipitation of th3 perchloric acid hydrate). Perchloric acid of 41.7% was prepared by miving nppropriate weights of concentrated perchloric acid (70.677,) and water . Chromous perchlorate was prepared by electrolytic reduction of Cr(Cl0,)g in HClO, as described in a previous publicationP The dinucleai. chromic perchlorate was prepared by passing air through the chromous solution for 15 minutes. Hexaquo chromic perchlorate was prepared by reduction of CrOI in HC10, with Hz02and crystallization. Ferric perchlorate and thorium perchlorate were prepared from the chlorides by fuming with IICIO, until all traces of chloride were expelled. The perchlorates of magnesium, zinc and nickel were prepared from 1,he oxides with HC104. Fe"1 concentration was determined iodometrically; CrIII by titration with standardized FeSO,, after oxidation with persulfate; ThrV by precipitation as oxalate, ignition and weighing a:: oxide. Zn", Nil1 and Mgll were titrated with 0.01 hl EDTA.
.
Results and Discussion
It mas assumed above that the depression of the freezing poini; depends only on the Concentration of the foreign ions and is independent of their nature, in other words the molal depression of the freezing point is constant. This assumption which was shown to be approximately correct for a number of solvents' was found to be true also for our perchloric acid solvent. The molal depression constant of the freezing point Kffor a number of metallic perclhlorates was measured, and the results are given in Table I. The cryoscopic constant has an average value of 4.43' and the different ions investigated have a k'f which lies within 7% of this value. The considerable differences in electric charge and radii have only a moderate effect. In solutions of Lif, Ag+ and Ba++ an elevation of the freezing point was observed, probably resulting from precipitation of solid solutions. The range of concentrations of the various cations investigated iis rather narrow because in concentra(7) R. K. Illohfullan and J. D. Corbett, J . Chem. Educ., SS, 313 (1956).
Vol. 65 TABLE
Solute, perrlilorate
I
Range of concn. (moles/1000 g. solvent)
0.0149 to 0 0360 0198 to 0297 0198 to 0297 0117 to 0334 0106 to 0424 0116 to 0194
Magnesium Zinc Nickel Ferric Hexnquocnromlc Thorium
KI
4 24 4 50 4 12 4 53
4 48 4 72
tions lower than 0.01 molal the relative error in At would be too high, and in concentrations above 0.04 molal the viscosity was usually too great to permit adequate stirring. This makes an extrapolation of K f values to infinite dilution impossible, but the results did not reveal any dependence of Kf on concentration in this range, so that such a dependence, if anp, must be small. In the eutectic solution containing 40.7% HClOd a very small degree of supercooling was observed and sometimes it was missing altogether. This caused an appreciable error in the determination of the freezing temperature. We found that with a solution containing 41.7% HCIOl a supercooling of 0.3 to 0.4' could be attained in all measurements and therefore these proportions were used in all experiments. The perchloric acid hydrate began to precipitate a t 57.5' and this process increased the concentration of the solute prior to the attainment of the eutectic freezing point; the values of the concentrations were corrected accordingly. Dinuclear Chromic Perchlorate.-The dcprcsPion of the freezing point caused by dinuclcar chromic perchlorate6 was measured in two solutions containing 0.0288 and 0.0432 gram atom of chromium per 1000 g. of solvent. The average depression was found to be 2.28' for a solution containing one gram atom of chromium per 1000 g. of solvent. If the structure of the chromic ion is dinuclear, the Kf will be 4.56, in exccllent agreement with the results of Table I. Trinuclear or tetranuclear species would give a Kr value of 6 3 4 or 9.12, respectively. It was shown previously6 that t h k chromic species contains two or more chromium atoms, but that it could not be a mixture of different polynuclear species. The rcsults of this work prove that this ion is indeed dinuclear, and so must have either structure I or I1 suggested prcviou~ly.~ H
[(H20)5Cr-O-
I
Cr(IJ20)6]4
+
(H,0)4C':
1
1
,c), i'r(fj20)4 ,'
'\
:;
11
Work is now going on in this Laboratory to dctermine molecular weights of other cationic and anionic polynuclear species in perchloric acid.
ELECTRON lMPACT SPECTROSCOPY O F NITROGEN DIOXIDE1 BYROBERT W. KISERAND I. C. HISATSUNE Department of Chemistry, Kansas State Unicersity, Manhnttan, Kansaa Receiued February 1, 1061
I n the course of some studies of thc nitrogen
August, 1961
NOTES
oxides, nitrogen dioxide was reinvestigated mass spectrometrically. The two principal ions of concern which are formed in the process of electron impact are NO+ and Not+. The determination of the appearance potentials of these ions from nitrogen dioxide was made and we found that they differed significantly from the recent results of Collin and Lowing2 and other^.^-^ Appearance potential measurements were made using a Bendix model 12-100 time-of-flight (TOF) mass spectrometer with an analog output system consisting of a monitor and a scanner. Wiley’ and others8-” have described the TOF mass spectrometer, and recently HarringtonI2 has described in detail the Bendix instrument. The voltage scale was calibrated by simultaneously introducing the sample of NO2 and a rare gczs with a known spectroscopic ionization potential.13 Krypton and xenon were employed in these determinations. Appearance potentials were determined using the extrapolated difference method described by Warren.14j15 The linear por-
tion of the ionization efficiency (i.e.) curves were forced to be parallel in plotting and then the voltage difference, AE, between the two i.e. curves a t any given current, i, was plotted as a function of i. The resultant curve, upon extrapolation to i = 0, gave a value of AE which was algebraically added to the ionization potential of the calibrating gas t o obtain the appearance potential for that particular ion. The ionization potential of NO2 was also determined using the technique of Lossing, Tickner and Bryce,16 the method of extrapolation to zero ion current of the linear portion of the i.e. curves, and the critical slope method developed by Honig.” The nitrogen dioxide showed no lasting deleterious effects upon the mass spectrometer. It was observed that immediately after admission of a sample to the ion source the trap current decreased markedly, probably due to oxidation of the filament (operating a t about 1900°K.). During these studies the 5 mil diameter tungsten filament burned out and had to be replaced. No other
1445
TABLEI
APPEARANCE POTENTIALS OF NO*+ AND NO+ FROM Not , Ion
NO*+
NO
This work
A. P. (e.v.)
Methodo
11.1 1 0 . 2 0 11.35f . l o 11.39 f .27 11.3 1 1 . 2 7 f .17
E.I.L.E. E.I.C.S. E.I.E.D. E.I.L.T.B. Average
No. of detns.
4 4 3 1
A. P. (av.)
11 11 12.3 & 0 . 2 10.0 9.91 11.7 13.98 f 0.12 11.3 & . 4 9.78 1 .05 10.1 .2
*
Literature Methodo
Ref.b
E.I.
18
S.
E.I. E.I.V.C. P.I. E.I. P.I. P.I. E.I.
6
4 3 20 1
19 5 1
13.25 1 0 . 5 3 E.I.L.E. 4 12.91 1 .53 E.I.C.S. 4 12.48 1 .43 E.I.E.D. 4 1 2 . 4 8 f .43 “Best value” NO+ US. NO:!+ 1.85 f .05 E.I.L.E. 4 1.17 1 .04 E.I.C.S. 4 1.15 f .10 E.I.E.D. 4 1 . 1 6 1 .08 “Best value” S. = spectroscopic; P.I. = hotoionization; E. I. = electron impact; L. E. = linear extrapolation: C. S. = critical slo e; E. D. = extrapolated digerence; L. T. B. = method given in reference 16; V. C. = vanishing current. “These re8rences are numbered to correspond to the footnote references in the text. +
(1) This work was supported in part by the U. S. Atomic Energy Commission under Contract No. AT(ll-1)-751 with Kansas State University. ( 2 ) (a) J. Collin and F. P. Lossing, J. Chem. Phgs., 28, 900 (1958); (b) J. Collin, ibid., 30, 1621 (1959). (3) R. J. Kandel, ibid., 23, 84 (1955). (4) R. J. Kandel, Phys. Rep., 91, 436 (1953). ( 5 ) T. Nakayama, M . Y. Kitamura and K. Watanabe, J . Chem. Phys., 30, 1180 (1959). (0) W. C. Price and D. M. Simpson, Trans. Faraday SOC.,37, 106 (1941). (7) W.c. Wiley, ScZence, 124, 817 (1956). (8) W. C. Wiley and I. H. McLaren, Reo. Sei. Instr.. 26, 1150
(1955).
(9) A. E. Cameron and D. F. Eggers, Jr., ibid., 19, 605 (1948). (10) M. M. Wolff and W. E. Stephens, ibid., 24, 616 (1955). (11) R. 9. Golke, Anal. Chem., 31, 535 (1959). (12) D. B. Harrington, “The Time-of-Flight Mass Spectrometer,”
in “Advances in Mass Spectrometry,” edited by J. D. Waldron, Pergamon Press, London, 1959, pp. 249465. (13) C. E. Moore, “Atomic Energy Levels,” National Bureau of Standards Circular No. 467, Vol. 3, 1958. (14) J. W. Warren, Nature. 165, 811 (1950). (15) C. A. McDowell and J. W. Warren, Disc. Faraday Soc,. 10, 53 (1951).
effects or changes were observed. NO2+.-1t is readily apparent from an examination of Table I that the individual measurements of the ionization potential of NO2 by various methods are not in agreement. It is unusual that photoionization results disagree by such a large amount; even the electron impact results show a much greater variation than one would reasonably expect. The present work shows the ionization potential of NOzto be 11.27 f 0.19 e.v., in fair agreement with the results of Stueckelberg and Smyth18and Weisder, et UZ.,~~ but in disagreement with Watanabe3” (16) F. P. Lossing, 4 . W. Tiokner and W. A. Bryce, J. Chem. Phiis., 19, 1254 (1951).
(17) R. E. Honig, ibid., 16, 105 (1948). (18) E. C. G. Stueckelherg and H. D. Smyth, Phys. Reu., 36, 478
(1930).
(19) G. L. Weiasler, 5. A. R. Samson, M. Ogawa and G. R. Cook, J . Opt. Soc. Am., 49, 338 (1959). ( 2 0 ) R. Watanabe, J . Chem. Phya., 26, 542 (1957).
Vol. 65
1446
with the value of 72 kcal./mole for D(0-KO) tabulated by C ~ t t r e l l . ~ ~ We observe that the results reported by us for IP(NOz+) and ,4P(NO+) from NO2 differ significantly from results reported by a number of inv e s t i g a t o r ~ ~ but , ~ - ~agree closely with those reported by Stueckelberg and SmythI8 and others. 19-20 We suggest that additional detailed experimental determinations of KO2 be undertaken in an effort to fix these potentials.
20.0
c
1
10.0 8.0 L
(21) G. G. Cloutier and H. I. Sohiff, i b i d . , 31, 793 (1959). (22) H. Huraeler, M. G. Inghram and J. D. Morrison, ibid., 28, 70
(1958). (23) T. L. Cottrell, "The Strengths of Chemical Bonds," Second Edition, Butterworths Scientific Publications, London, 1958, pp. 210 and 278.
T H E INFRSRED SPECTRA OF SOME DIMETHYL SULFOXIDE COMPLEXES
O ' I /// 0.4
BY RUSSELLS. DRAGO AND DEVON MEEK
-
0,2F I 0.1 13
I
10
12
14
16
18
20
Electron energy, uncorr. (e.v.). Fig. 1.--Ionization efficiency curves for NOz+ and N O + from NOz,using krypton for the electron energy calibration.
and others.2-6 Since the i.e. curves for KO,+from SO2 should be of the norim1 form, the various methods of determining the electron impact ionization potential should all give essentially identical results. Table I show, that this is the case, and a simple average of these values gives the yalue of 11.27 =t 0.19 e.v. That the i.e. curve for SO2+ is of the normal form is readily seen from Fig. 1. During our studies of SO2 we never observed any peaks due to S03+ or N204+. NO+.- The ion of m/e 30 is NO+. It was found to have an appearance potential of 12.48 i= 0.43 e.v. I n addition, individual determinations of the ionization efficiency curves of S O + and NO2+ resulted in hP(NO+) - IP(KO*+) = 1.16 h 0.08 e.v. This coupled with IP(T\;O2+)= 11.27 =t 0.17 e.v. from above gives A P ( N 0 f ) = 12.43 =t 0.19 e.v., in agreement with the value of 12.48 f 0.43 e.v. determined directly. We attempted to determine the appearance potential of SO+from N O with little success. We consistently obtained values about 1.3 v. too high, for which mc= can offer no explanation. We are presently repeating this determination. It may be noted from Fig. 1 that there is a rather sudden change in slope iii the X O + i. e. curve a t approximately 14 v. The latter was also observed in the deterniinatioiis of S O + from NO. Using the value of 12.48 e.v. for AP(NO+) from NO2 and the literature value of IP(NO+) = 9.25 e.v.,21-22we calculated the nitrogen-oxygen bond energy in nitrogen dioxide. D(0-NO) = 3.2 e.v. nr 74 kcal./mded This is in reasonable agreement
W m . A . S o u e s L a b o i a t o ~ y ,Chemistry Department, Lrnieeraity of Illinois, Urbana, Illinois Received Ja?iuary 81, 1961
The preparation of a series of dimethyl sulfoxide complexes has been carried out in this Laboratory and assignments were made for the S-0 st'retching frequency.'& Subsequent to reading proof of the art'icle on this research, assignments which differed from ours were published.2 The purpose of this article is to provide additional data to support some of our original assignments and to aid in solving a problem which is more complex than the present published information would indicate. Results and Discussion The preparation of the complexes and conditions employed to obtain the spectra have been reported,' There is agreement on the assignment of the S-0 stretching frequency for dimethyl sulfoxide (1045 cm.-l in CH3S02 solution), There is also agreement on a very sharp, much less intense absorption a t 950 cm.-l, and a, weaker, broader peak a t 915 cm,-' assigned2 to -CH3 rock. In the complexes there is a discrepancy in the assignments. The addit,ional information we have to offer involves a detailed description of the above two absorption peaks in the complexes as measured in nitromethane solution. Sharper, more easily interpreted spectra are obtained in solution than on Nujol mulls of the solids. Table I contains data for the peaks in the 1000 and 950 cm.-l region. The 1000 crn.-' peak is the one in the complexes that n-e assigned to S-0 stretch and the latter is the oiie assigned by Cotton, et aL2 Some very significant information regarding t8hese spectra entails a qualitative description of the relative intensity of the peaks measured in solution. As the peak in the 1000 cm.-' region moves from 104.7 cm.-l in free sulfoxide to 1025
-
(1) (a) D. W.Meek. D. K. Straub a n d R. S. Drago, J . A m . Chem. SPC,,82, 6013 (1980). (b) For the preparation of these complexes also see F. A . Cotton and R. Francis, abid., 82, 2986 (1960) and H. I,. Schliifer and W. Schaffernioht, Angew. Chsm., 72, 618 (1960). (2) F. A . CotOon, R. Francis and W. D. Horrocks, Jr., J. Phvs. Chem., 6 4 , 1634 (1900).
