March, 1960
NOTES
367
TABLEI EXF~ERIMENTAL RESULTS FOR SODIUM AND CALCIUM EXCHANGED FORMS OF ZEOLITETYPEA Material and activated unit cell formula
Run no.
Sodium A Natt [(A10~),~(Si02)lz]
R C in A.
2.5''
3
15.3 18 2 18.7
4
..
5 1
.. 19.7
2.79 .. 2.6-3.5" 2.86
1 2
$2.1 63.9
4 . I!) 5.15"
1 2
Calcium A, 38% exchanged Nar.rCat.3[(AlOMSiOZ)d Calcium A, 76% exchanged Nar.pCac.r[(AIO,hdSiO)l, I Moisture pickup during the experiment.
-.\rtivntrd---,-p in A . 1
the Si02/A1203 ratio and consequently the number of cations within the unit cell. As in the case of zeolite type A, it is most likely that the scattering observed arises from the cations in the structure being in disorder. As zeolite Y has fewer cations per unit cell than zeolite X, the average distance between cations must be larger in Y than X and hence the slope of the scattering curve should be increased relative to X as is shown by Fig. 5. Upon hydration the c:ttions would be expected to assume positions farther from each ot8hrras they "float" into the zeolite channels thus increasing the slope of the scattering curve. This is observed, Fig. 5,
3.i5
-Hydrntr'dp in A . 2
Ro in
A.
72.4 41.3 .. 39.4 41.4 47.2
5.49 4.14
80.4 78.0
5.78 5.7
.. 4.05 4.15 4.4
for zeolite type Y. The loss of intensity upon hydration comes about either because the cations take up more ordered positions or because the scattering from the cations and water molecules interferes destructively. Acknowledgments.-The author is indebted to Mr. L. G . Dowel1 who supervised the modification of the spectrometer ai!d recorded many of the data; to Mr. L. J. Dobmier who aided in the experimental work; to Dr. D. W. Breck who suggested that zeolites might scatter X-rays at low angles; and to Dr. V. Schomaker for his assistance in clarifying the interpretation of the experimental data.
NOTES THE: ULTRAVIOLET ABSORPTIOK SPECTRUM OF CHLORANIL' BY KARL13. HAWSER AND It. S. MULLIKEN Laboratory of Molrculor Structure and Spectra, Department of Physics, Uniuersrty of Chtcago, Chtcago 57,Illtnors Receroed September 9% 1969
Chloranil is known to be a strong electron acceptor and to form complexes with various electron donating Its absorption spectrum in chloroform down to about 2300 A. has been recorded by Pummerer and c o - ~ o r k e r s . ~More recently, Smith4 has investigated the interaction of chloranil with various electron donors. In the course of this work it was noted that chloranil in n-heptane shows a weak absorption band a t 27,500 cm.-l, in addition to the much stronger absorption (The Snme bands with peak a t 34,600 cm.-'. were recorded by Pummerer, et al., in chloroform; their curve indicates peaks at about 26,600 and 35,000 cm. -l with log t 2.48 and 4.02, respectively.) Because chloranil is a strong electron acceptor, it was thought thnt the weak band at longer wave (1) Thia work was assisted b y the Office of Ordnance Resenrch under Project TB2-0001(505) of Contract DA-11-022-ORD-1002 with the University of Chrcago. (2) E. Welt%.2.Elektroefiem.. 34. 538 (1928); Anyew. Chem.. 66, G5S (1954). (3) Pummerer. Schrnidute and Seifert, Chem. Rer.. 86, 645 (1952). (4) N. 9. Smith with W. G . Brown. to be piihlirrhed.
lengths might be analogous to the strong ultraviolet absorption observed at longer wave lengths than that of 1 2 itself in iodine solutions 111 nheptane and other paraffin hydrocarbons, and thus be identified as a contact charge-transfer spectrum.6 In order to distinguish between this and the possibility that the band in question is chamcteristic of chloranil itself, we have measured the absorption spectrum of chloranil in ethanol, n,heptane, perfluoroheptane, and in the vapor phase (see Fig. 1). The absorption band at 27,500 cm.-1 appears in all three solvents at nearly the same frequency and with the Same intensity within the experimental error. Since perfluoroheptane is known to be an extremely inert solvent, having no detectable interaction wit,h dissolved molecules and giving rise for the Iat.ter to spectra very similar to their vapor spectra,G this indicates that the band a t 27,500 cm.-' is an intrinsic transition of the molecule itself. We have not been able to find in the region above 42,000 em.-' any absorption attributable to an interaction lxtwcwn chloranil and n-heptane or ethanol. Since thr long wive Iciigth :J)sorptioii h i d is only a rnthrr wc:ik h:md (cxtiiiction coefficient E ( 5 ) I). F. I.hnns, J . Clrern. Phys.. 93, 1422 (1955); J . Chem. Soc., 4229 (1957); L. E. Orgcl nnd R. S . Mulliken, J. Am. Chem. SOC.,79, 48 (19.57). (ti) Cf.I). 1'. Evans. J. Chem. I'hua., 93. 1424, 1426. 1429 (1955).
VOl. 64
NOTES
I
5:
4t
'-' ,
I
2
_ _ _L---i _ _L---j
25000
35000 4OOOO 45000 5oooO Y*, tin.-' Fig. 1.-Absorption spectrum of chlorsnil, molar rxtinction coefficient c as function of wavc number, V -, v:tpor 1'FIl-, pcrfluoroheptnne solution; , n-hept:tne solution; Et-.-.-, ethanol solution. In the cahe of the vapor only relative e valucs werc determiued; the curve plotted in Fig. I is based on the nssiimp tion that el,,.x for the peak at 36,500 is the wme as for the highest peak in perflriorohcptane solution.
16
tance as an abrasive and also as a refractory material. In addition to this, it recently has found application as a high-temperature semi-conductor. This should warrant the determination of the thermodynamic properties of silicon carbide. Unfortunately, presumably on account of the refractory nature of this compound, the literature data on the heat of formation of silicon carbide do not agree. Mixterl finds 1.3 kcal., Ruff and Grieger2 26.8 kcal. and Humphrey, et u Z . , ~ 12.3 kcal. In this note the original literature values have been adjusted for a graphite with heat of combustion 94.0 kcal./fwt. The three different free energy of forinlttion functions arising from these heats of forniation are plotted in Fig. 1. These free eiiergy func-
30000
Temperature,
O C
I
+I5
+IO
about 300), arid since the solubility of ch1or:uiil in perfluorohept:uic is very poor (about mol./ l.), it was necessary to use absorption cells of 10 +5 cm. light p:Lth length. With regard to the vapor, uiifortunately a t the highest temperature obtainable with our equipment (about 120') the vapor 0 pressure of chloranil is so small that with the same 10 cm, light path length the maximum extinction rJ of the strong h i d ne:tr 35,000 cm. was about 0.3; U li: coiisequently the extinction expected for the weak - -5 band would be of the order of 0.004 :uid therewith tL below the limit of detectability. Q The center of gravity of the strong :Lbsorption -10 band near 35,000 cm.-' is ne:irly the same in nheptane and ethanol. However, whereas in nheptane one can still cle:vly recognize the doublet structure of the b:md, it is broadened to one single -1 5 somewhnt asymmetric b:md by the stronger interaction with the p o h solvent ethniiol. As oiie would expect, the vapor spectrum is shifted towc-nrtl shorter wave lengths (by nearly two thousand W L V C numbers) :LS comp:tred with the n-1iept:iiic solution spectrum, atid shows :L resolved fine structure. In perfluorohcptane, the center of gravity of the absorption band is closer to 0 0 0 0 0 0 0 0 0 0 0 0 the vapor spectrum than to the n-heptme solution 0 In 0 In 0 9 ( u C Y m m Tt spectrum. Further, it has almost the same fine structure as the vapor spectrum, only slightly Temperature, less resolved. This shows again that in perfluoriFig. 1.-The st:tnd:rrd free cncrgy of formation of silicon nated solvents the iiiteraction between the solvent c:trbide (from graphite and liquid silicon) :icrording to d:it:i and the dissolved molecule, and consequently the of various invcstig:itors. The short dnshtd linc :tcross the iiifluerice of the solvent on the absorption spectrum, AP" = 0 linc indiwtvs the tempcr:tturc, 2830", a t which silicon carbide decomposes into gr:iphitc and a solution of is much smalJer than in hydrocarboiis. O K
c:irt)ori in liquid silicon, a t 35 atmosphrrcs (Scncc :tnd
S1:ick").
tions h:ive hecii calculatcd hy means of t,he cquet,ion ( 1 ) W. G . hlixter. A m . . I . S e i . . I 4 1 24, 130 (lWl7). (2) 0.Iliiff sntl P. G r i w p r , %. nizorn. nllgenr. C h ~ m .111, , 145 (I9:j:j). (1%) C:. L. I I I I I ~ I ~ I IS. I ~S. I ~'l'ii’ /piill = e ( j - i ) A F Q / R T (2) of silicon in graphite. We have assumed it also Drowart, De I\laria and Inghram9 have deter- to be negligibly small. mined by a mass spectroscope technique the partial In Fig. 2 (log I ) ~ ~ ’ log z)ijt’)/(j - i) = pressures of several molecular Specks in vapor cf- 0.4: 0.30 (deeper penetration also corresponds to very small changes in external bath concentration, if the bath ratio is not too small). In applying the “integrated boundary condition” with the linear er curve in method 11, however, the error because of the curvature is compensated to an appreciable extent, as the contribution of the volume element at distance r to the total sorbed quantity is (c s) 2 ?r r dr, thus decreasing in proportion to r . In Table I1 we have made calculations in analogy to those for the plane sheet and the results give some idea about the decreasing accuracy as ro/a decreases. For the sphere the same sort of discussion is applicable. There is still larger increase of curvature with decreasing r , cj. (17),l and in spite of the compensation because of the rapidly decreasing volume elements 4+dr the error of method I1 increases
’
+
March, 19GO
NOTES
more rapidly with decreasing ro/a than for the cylinder as demonstrated in Table I1 (now A in eq. 3 is subst#ituted with V / 2 a ) . Also from the formal point of view method I is less troublesome than method 11. We conclude that the approximation method suggested2 is very useful for a plane sheet to get a great accuracy. It has formal advantages to the steady-state approximation method for a cylinder and also a great,er accuracy in the range indicated here. But it lhas no formal advantage to the steady-state aplproximation method for a sphere and a greater accuracy in a rather narrow range.
Subsequent examination of the data failed to reveal any significant effectof over- or under-cycling or of unequal sweep rates, or of the degree of resolution upon the magnitude or direction of the effect. It should be emphasized that optimal resolution wtts not attained, nor waa bad resolution tolerated. “Ringing” waa never observed for any of the image lines, or for the Cla satellites of (CH,),Si, but was attained a t the ClS lines of the other compounds in about half of the measurements. The widths of the image signals a t half-maximum were 0.70 to 0.85 c./sec. in nearly all cases, while the CI3 satellites (when not ringing) were, if anything, narrower than the corresponding C12 images, except for (CH& Si, for which they averaged ca. 0.12 c./sec. wider. The effect of sweep direction is fairly large, the C13 isotope effect appearing to be larger by 0.0022p.p.m. when the sweep is toward decreasing field, or smaller by the sanie amount when the opposite sweep is used exclusively. Although such error was avoided, it would not have concealed the isotope effect.
PROTON NlJCLEAR SPIN RESONANCE SPECTROSCOPY. XI. A CARBON-13 ISOTOPE EFFECT BY GEORGE VANDYKETIERS Contrtbulion Yo. 166 from the Central Research D e p l . , A4~nncsota Mtning .E. Mfg. Co , Sl. Paul 19, Mrnnesotu Received October 16, 1969
Recently n nuclear spin resonance (n.s.r.) “isotope effect” of CI3upon attached fluorine atoms has been discovered.’ The rather unexpectedly large shifts were always found to be in the direction corresponding to greater shielding by CL3than by C12. Though no shift was found for protons, the much smaller effect anticipated by analogy with the deuterium shifts2would not have been detected. As both the sign and the relative magnitude of such a C13 effect might prove theoretically interpretable, a morc elaborate experimental procedure has been used in the present work. Experimental The compounds studied were examined neat in the customary 5 mm. 0.d. aample tubes, from whirh air had been swept by means of a brisk stream of bubbles of prepurified nitrogen; however, air-saturated CHClj and ( CH3),Si were found to give the same results. The C13 isotopic isomers were present at their natural abnndances. The n.8.r. spectrometer and measurement techniques have been d e ~ c r i b e d . ~For the present study separate reference compounds were not employed, the exceedingly strong sharp signal from the normal (CI2) compound being used instead; “image” lines (akio called “side-bands”) are readily produced from it by audio-oscillator modulation of the magnetic field. When Cl3 ia present, the proton signal is split by it into a doublet, the. coupling constant bring ca. 100 to 250 c./sec. The positions of these two weak CL3satellite lines are measured, sepaxately, rclative to the strong C1* central peak by use of the “image” lines.3 The small but reproducible difference found in earh case results from the isotopic shift, aa otherwise the high- and low-field C13 romponents would be equally spared from the C12central line. Errors random in nature were counteracted by multiple repetition of measurements; all data have been used and weighted e ually. Measurements were made over a threeweek perio8. A t each session six sweeps were run on each of the two C13 lineal for each of the compounds studied. In addition to the routine alternation of sweep direction,a which virtually eliminaki errors due to differential saturation or to “ringing” of the signals, in most rases rare waa taken in the magnet cyrling to obtain a “flat” field and hcnre very symmetrical peaks for both dirertions of sweep; ciweep rates also were controlled to be equal (within 10’70) in both directions. (1) P. C . Lauterbur, private communication: I am indeed grateful for this advance information, without which it is unlikely that the preaent work would have been begun. (2) G . V. D. Tiers, J . Am. Chcm. Soc., 79, 5585 (1957). J . Chcm. Phys.. 49, 963 (1958). (3) G. V. D. Tiers, ’ P H I 8 JOURNAL, 60, 1151 (1958).
373
TABLE I THE EXCESSN.S.R. SHIELDING,A7, PRODUCED BY 0 1 , RELATIVE TO C1*, IN SEVERAL COMPOUNDS Compound
No. of meas.
A T , p.p.m.0 (Cia-Cl*)
J(C”H), c/sb
Shielding value, I C
(CH3)rSi 18 f0.0042 118.20 1O.OOO CHjI 18 .0012 151.17 7.843 CHZClz 12 ,0042 178.24 4.720 CHCla 18 ,0059 209.17 2.755 a Standard deviat,ion of the averaged value was ~k0.0012 p.p.m. in each case. Standard deviation of the averaged value was f0.09 r./sec. in each case. Measured in dilute solution in cc14, as described in ref. 3.
+ + +
Results and Discussion The results presented in Table I demonstrate a small but statistically significant excess shielding by CI3 in three molecules, namely, CHC13, CHBC12 and (CH3)$i. I n the case of CH31the shift is not large enough to be considered as established. The center of the doublet corresponding to protons attached directly to C13 is found a t shielding values higher by ca. 0.004 p.p.m. than the line due to the normal (CI2)compound. This effect is of the same sign but only about 1/44 as large as the corresponding effect upon fluorine.’ The same ratio, 1/40, has been observed for the deuterium isotope effect upon proton and fluorine shieldings.2 I t appears unrelated to the ratios found for the coupling constants in the same compounds, ca. 1/2 for J(C1a-H)/J(C13-F) and ca. 1/4 for J (D-H) /J (D-F) . The data in Table I seem to indicate a significant variability in the magnitude of the effect. The variation observed is not simply related either to the coupling constants, to the relative shieldings, or even to the number of substituents; however, the experimental uncertainty is such as not to warrant more detailed studies at this time. The conclusions reached above are of course entirely dependent upon the successful elimination of directed error in the measurements. Potential errors and the precautions taken have already been discussed in the Experimental section. There may well be further, unrecognized, sources of error; for example, treatment of the C13HC13 spectrum as an “AX” case rather than as an “AB” case‘ in fact must result in an apparent excess shielding by C13, even if there were no isotope effect. By (4) J. A . Pople. W. G. Schneider and H. J . Bernstein. “High-Resolution Nuclear Magnetic Resonance,” McCraw-Hill, Book C o . , Inc., N e w York, N . Y., 1969. pp. 118-123.
374
Vol. 64
NOTES
algebraic manipulation of the relations appropriate to the two cases, the error is shown to be only O.oooO1 p.p.ni. for CHC13, and should be equally negligible for the other molecules studied. The most serious effect of bad resolution should be an increase in random errors, with the result that the CL3effect would be rendered highly uncertain. If resolution varied substantially with sweep direction, :L directed error might be produced; however, no evidence for such could be found by careful analysis of the data. The apparently poor resolution of the C13satellites for the (CH3)& (as judged by the excess line width, ca. 0.12 c./sec., observed for them) is actually due to the exceedingly weak spin-spin coupling between protons on C13 and those on C12. This splitting becomes observable as a result of the "effective chemical shift," produced by the magnetic moment of individual C13 atoms upon their attached pro* tons. The coupling constant, J(C12H3SiC13H3), must be approx. 0.12 c./sec., the multiplet being assumed to be 10-fold, with binomial intensity distribution; J = (0.902- 0.782)i/2/3.6. Acknowledgment.l-I am indebted to Donald Hotchkiss for the excellent careful operation of the 1i.s.r. equipment so nccessary in this work. T H E S't'ANDARD ELECTRODE POTENTIAL OF THE QUINHYDRONE ELECTRODE FROM 225 to 55" BY JoriN
Harned and Wrightll determined the normal electrode potentid of the quinhydrone electrode from 0 to 40' by combining their data with activity coefficients determined by Harned and Ehlers12 by a different method. In the present investigation the standard electrode potentials of the quinhydrone electrode were measured from 25 to 55 a t 5" intervals, using a AgAgCl reference electrode. The establishment of standard electrode potentials a t each interval allows the direct calculation of activity coefficients a t these temperatures. In the calculations, it has been assumcd that the cell reaction is Quinone
(1)
The electromotive force for this reaction is given by the Nernst equation in the forrr.. E = EO,,II
+ RT - In a m F
(2)
Since the quinhydrone used is an equimolar compound of quinone and hydroquinone, and since these substances are non-electrolytes of low solubility in contact with the solid, their activities should be constant. At higher temperatures, the solubility of quinhydrone becomes appreciable (7 g. per 100 g. Hz0),2but as long as the ratio of the activities of quinone and hydroquinone is nearly constant, these activities may be ncglected. Therefore, the Nernst equation may now be written
+
2RT E = Eoc0i~ -F- In
c. HAYES*AND M. H. LIETZKE
Contribalion f r o m the Chemistry Divisinn, Oak Ridoe National Lnborntoru. Oak Ridge, Tennessee RPceived October 17, 1969
+ 2HC1 + 2Ag r'Hydroquinone + 2'4gC1
~I~CIYHCI
(3)
If one assumes that (4)
Since the discovery3 and development4 of the quinhydrone ckctrode, it has found frequent use as a substitute for the hydrogen electrode for pH where A and B are parameters and I = ionic measurements. The electrode is convenient to use strength, S = 1.17202 (23375.2/DT)'/2,and D = dielectric constant of water (at temperature T ) , and it gives results which are easily reprodu~ible.~-~ The earlier research established the fact that meas- determined by the Akerlof equation13then equation urements with the electrode in certain solutions 3 may be rearranged to give contained a "salt error,5 which was due* to a change in the ratio of the activity of hydroquinone and quinone caused by t,he presence of other dis(5) solved substances in solution. While other invesIf the quantities on the right are equated to EO", tigations werc concerned with the "salt-error," they also established the normal potential of the then quinhydrone e l e ~ t r o d eas, ~well as the normal potentials of t.hc hydroquinhydrone and the quinoquinhydrone electrodes. The Eoat any one temperature may be determined ( 1 ) 1'111s paper is hasrd upon work performed for the Atomic Energy by extrapolating the Eo" values to zero ionic Coinrnission at tlie Oak IWae National Laboratory operated by Union strength. In this work, however, the parameters Car'Gic. Coi jmration. A and B and the values of Eo in equation 5 were (2) Reseal-sh I'a,rtiriiiant for the Sumnier, 1969, from IIainline Univi nlity. S t . I'aul. Alinnrsota. determined by a non-linear least squares method (3) 1;. 1laht.r and R. Rims, Z . phvsik. Chrm.. 47, 257 (1904). on a high speed computer (the ORACLE). (4) S ( ; r a n g e - &nd .i. AI. Nelson, J . A m . Chem. Sor., 49, 1401 t5
12.
( IR'L I J .
(A) 15. 1riilcn:in. Ann. C h ? m . ,[9] 16. 109 (1921). ( 6 ) .1. I,. It. Morgan, 0. h l . I,ainmert and hI. A . Campbell. J . A m . Cham. Soc.. 63. 454 (1931). (7) E. Biilnran and .4. L. Jensen. Bull. soc. chrm.. 41, 151 (1027). ( 8 ) S. P. I,. Sijren,scn, AI. Sorensen and K. Iinderstrom-Lang, Ann. Cham., 16, 283 (1!321). (9) F. Ilovorka and W. C. nearing, J . A m . Chcm. Soc.. 67, 440 (1935). ( 1 0 ) 11. 1. Stonchill. T m n s Faradav Snc.. 39, F7 (1943).
Materials and Apparatus Quinhydrone.-The quinhytironr n s r d in this projwt waa Enstman KO.217, rcrrystnllincd from wttcr hcntcd to (11) I[. S. Ilarncd anti D. D. Wright, J . A m . Chem. Soc., 66, ,4849
(1933). (12) H. 8. Harncd and R . W. Ehlcrs. J . A m . Chem. Soc., 66, 2179 (1933~. (13) G. C. Akcrlof and TI. 1. Oshry, J . A m . Chcm. Soc.. 79, 2844 (1950).
NOTE s
March, 1960 65". Thc recrystallized product was dried overnight in :I vaciiiim desiccator. The qiiinhydrone gave a melting point :it 168 i 1'. Nitrogen. -All operations involving quinhydrone were done in an at,mosphere of nitrogen. The biibbling of nitrogen into the cell during e.m.f. measurements not only gave an incrt :itmosphere but also provided stirring of thc solution. Hydrochloric Acid.-Stock solutions of hydrochloric acid were prepared from Fisher Certified Reagent Hydrochlorir Acid. Concentrations were determined by titration of the solutions with standard alkali. Densities were determined by weighing a certain volume of solution. The concentrations then were expressed in terms of molalities to eliminate caorrections for thc change in volume a t different tempemtures. Ag-AgC1 Electrodes.-The silver-silver chloride elert rodes were prepared according to the method described by Greeley.I4 Cell Measurements.-The cell potentials were measured with a vibrating reed elect,rometer (Model 30, Applied Physics Corp.) used as a null instrument. Potentials wpre read from a potentiometer (type B, The Riibicon Co.) using a Brown Recorder t o indicate equilibrium conditions. The cell was held in a constant temperature bath (controlled to +0.02") by meanti of a brass tube holder which had holes drilled in it to facilitate thermal eqiiilibriun. The tube holder :dso shielded tbr. cell from any environmental capacitance. Experimental The cells,, without the quinhydrone, were allowed to equilibrate for a t least one hour in an atmosphere of NI. Measurements of rlotential were started from 5 to 10 minutcs after the quinhydrone had been added. Eqiiilibriiim volt:tges were considered constant when they did not change more thnn 0.2 mv. in a half-hour period. Equilibrium voltnges oft,cn were olxerved within 15 minutes and rc,maincd r,ssent,i:dly constant for as long as 2 or 3 hoiirs.
Discussion of Results The EO valuers for the Ag -AgCl, quinhydrone cell from 25 to 55'' are listed in Table I. The best value of the paramtter A in equation 4 was found to be 1.5. 1 THEAg-AgC1-&UINHYDRONE CELL T.4RLE
EO
(YB)" a
400
300 0 4702
350 0 4762
0.4765
0.0091i 0 0148
0.0120
0.0052 0.0062 0.0118
250 0.4771
450 550 0.4771 0.4770
This term is dcfincd i n cqiiation 5.
TABLE I1 STANDARD ELECTRODE T'OTENTIALS OF ELECTRODE
375 TABLE I11 ACTIVITYCOEFFICIENTS OF HCl 1"
m
250 212
3'6
4'4
5 I6
017
0.001 0.9654 0 . !I656 0.9650 0.9653 0,9650 ,002 .!)524 ,9521 ,9520 ,9525 ,9519 .9283 0.9284 ,9287 ,9280 ,005 ,9287 .!I285 .Ol ,9048 ,9048 ,9045 ,9044 ,8757 .(3040 .02 ,8754 .8755 ,8753 ,8301 ,8747 .05 ,8295 ,8304 .BO8 ,8310 ,8310 .8296 .1 ,7923 ,7964 ,7967 ,7972 ,7938 .?!'I58 30'
m
0.001 0 !3652 0.9650 0.9648 ,002 ,9521 ,9515 ,9518 ,9274 ,005 ,9284 .9275 ,9034 ,9034 .O1 ,9047 .02 ,8759 ,8741 ,8741 .05 ,8319 ,8285 ,8291 .1 ,7981 ,7940 ,7946 m
*
350
0,001 0.9648 0.9647 ,002 ,9515 ,9513 ,005 ,9274 ,9268 .01 .903S ,9025 .02 ,8737 ,8731 .05 .a279 ,8265 .1 .7,9lG ,7918 rn
40"
0.001 0 . !I642 0.0643 0.9642 .002 ,9508 ,9505 .9507 005 ,9261 ,9265 ,9268 .O1 ,9011 ,9016 ,9026 .02 ,8701 ,8715 ,8735 .05 ,8208 ,8246, ,8283 ,1 .77X ,7891 ,7927 m
450
0.001 0.9F39 0.9644 ,002 ,9503 ,9504 ,005 .!I253 ,9261 .01 ,9001 .!loo8 02 ,8090 ,8704 .05 .8194 ,8232 1 ,5780 .7872 m
550
QUINHTDRONE0,001 0.0GS1 0.9636 ,003 .94!)3 ,9497 250 300 350 400 450 550 005 .!I240 ,9240 0.6995" 0 6953 0.6919 0.0886 0.0854 0.6776 01 8987 ,8990 .G995* ,6953 .GO18 ,6885 ,6854 ,8776 02 ,8076 .8680 .05 .8192 ,8195 .699416 6952 ,8919 ,6885 ,6885 ,0776 699711 6!160 ,6923 ,0886 1 ,7801 ,7829 . G9948 " Ilcwilts of this investigation. a Based on this investigation plus Eovaliies for Ag-AgC1 The activity coefficients of hydrochloric acid are from Harned and Ehlers; see also h t c s "Electrometric pH Determinations, Theory and Practice," John Wiley and listed in Table 111. These values agree closely Sons, Inc., New York, N. Y., 1954. Bascd on this in- with those obtained by other investigators. The vcstigation plus ED values for Ag-AgC1 from Greelcy." THE
The standard electrode potentials of the quinhydrone electrode have based based on the EO values of the Ag-AgC1 electrode from the values of Harned and Eh.lers,I2as well as those of GreeleyI* (which were measured in this Laboratory). As shown in Table 11, these values agree usually to within 0.1 mv. (14) 1959.
R. S. Greeley, Ph.D. Thesis, University of Tennessee, May 22,
largest deviations occur a t 40 and 4 5 O , yet they are still less than 1%. Agreement is unusually good at the other temperatures. Hence it appears that the quinhydrone electrode may be used in activit,y coefficient measurements in a manner similar to the hydrogen electrode. (15) R. G Bates and V. E. Bower, J . liesearch N d l . Bur. Standnrds. 63,283 (1951). (16) T . Shedlovsky, .I. A m . Chem. Soc., 79, 3680 (1950). (17) G . J. IIllla and D. J. G. Ives, J . Chem. Soc., 318 (1951).
376
Vol. 64
h70TES
Acknowledgment.-The
authors wish to thank
Dr. R. W. Stoughton for intcrcsting discussions in connection with this work.
nroirs
TABLE I11 HEATSOF SOIJJTION O F IiClO,,(e) IKX)~ x 103/
950
THERMOCHEMISTRY OF POTASSIUM PERMYANGANATE, POTASSIUM MOLYBDATE, POTASSIUM C,HLORATE, SODIlJM CHLORATE, SODIUM CHROMATE AND SODIUM DICHROMATE BYTHOMAS NELSON, CALVIN Moss AND LORENG. HEPLER~ Contribution from Cobb Chemirnl Lnborntorg. UniversifQ o f Virgiiiia Charluttuuille. Va. Received October 66,1969
1111.
A?l(kcnl./molr)
IT20
4.245 6.822 7 . 770
9.87 9.92 9.85 9.90 9.91 9.90
11.01
12.93 16.16
TABLE IV HEATSOF Sor.r~i,ro~ OF KaC103(c) Molrs NaCIO. X 102,' 9.50 nd. 1lzO
Al€(Rcnl./mole~
0.9260 1 ,032 1 .OF2 1,178 I .263 2.270
5.21 Heats of ,solut?ion of KMn04(c), KCIOa(c), 5.21 NaC103(c) and NnnCr04(c) and the heat of re5.23 action of Na26r207(c) with excess OH-(aq) have 5.23 been determimd as part of an undergraduate re5.24 search program. These heats have been used for 5.24 calculation of standard heats of formation of K2M004(c),Na2CrO4(c)and NazCrzOi(c)and standTABLE V ard partial rrtolal entropies of Mn04-(aq) and HEATSOF SOLUTION O F Nn2Cr04(c) IN l o w 3.If OH-(aq) ClOa-(aq).
Experimental The solution calorimeter used for these determinations has been described.*-* All calorimetric determinations were carried out at 25.0 f 0.2' with 050 ml. of water or solution in the calorimeter. All of the saltis were prepared for use in the calorimeter by recrystallization of the appropriate reagents and all except NaC1O8 and KCIO, were analyzed by common volumetric methods. Some samples were dried in vacuum desiccators and some in an air oven.
Rlolrs NaCrOd x lo'/ 950 nil. soln.
AH(konl./mole)
0,5632 ,6645 .979 1 1.016 1,137 1.408
-4.42 -4.44 -4.42 -4.3!) -4.43 -4.36
Heats of reaction of Nx2Crz07(c)with a small Results excess of aqueous NaOII as in the equation Heats of solution of KMn04(c), K2Mo04(c), Na2Cr207(c)+ 2OH-(aq) = 2Nn+(aq) + 2Cr04'(nq) + KC103(c), NaClOl(c) and Na2Cr04(c)are given in N20(1) ( 2 ) Tables I-V. I n each table, AH refcrs to the reare given in Table 1'1. action salt(c)
=
(1)
snlt(nq)
TABLE I HEATS OF SOLUTION O F
hloles IZMnOl X lo*/ 950 nil. HrO
4.810 5.G87 6. 345 6.5i0 7.187 9 !I I58
KMn04(c) A€l(kcal./molr)
10.42 10.30 10.43 10.44 10.42 10.44
TABLE I1 ~ I E A TOF~ SOLUTION OF K2MoOr(c) IN 10-8 A4 OH-(aq) Mole8 KrR'loOi X loz/ 950 nil. s d n .
0.5741
A??(kcnl./mole)
1.100
-0.82 - .79
1.359 1 . ti23 1 ,746 I .784 2 . o"2
- .79 - .75 - .i7 - .73
-
.82
( 1 ) Alfred P. Sloan Foundation Research Fellow. ( 2 ) R. L. Gruhain and L. G . Ilepler, J . Ani. Chem. Sor., 7 8 . 48.16
(1956). ('0 N. RIiilrIrox, .Jr., and L. Q. Ilrpler, ihid. 79, 4045 (195;). ( 4 ) 3 f . R I R i r k y and L. C . llepier, to hr ~ ~ i ~ h l i s h e d .
(..
T A B L E VI HEATSO F SOLUTION O F K:&r207(r)
RTolrs NnrCrKh X 1 0 3 / 950 nll. s d n . AI(il?s 011- X IO2
2.087 2.646 4.302 4.640 5.106 6.549 6.567 8.566 10.482
1 .0 1.0 1.5 1.5 2.0 1.5 1.8 2 .5 2.5
IN
OH-(nq)
Alf(kcnl.,'~nole)
-21.36 -21 .:14 -21.33 -21.49 -21.36
-21.23 -21.32 -21.34 -21 26
Heats of dilution of I
