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
c. HAYES*AND M. H. LIETZKE
Contribalion f r o m the Chemistry Divisinn, Oak Ridoe National Lnborntoru. Oak Ridge, Tennessee
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
+ 2HC1 + 2Ag r'Hydroquinone + 2'4gC1
(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
~I~CIYHCI
(3)
If one assumes that
RPceived October 17, 1969
(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. A n n . 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, A n n . 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 .9275 ,9274 ,005 ,9284 ,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 Bascd on this in- with those obtained by other investigators. The Sons, Inc., New York, N. Y., 1954. 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,I2 as 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
9.87 9.92 9.85 9.90 9.91 9.90
4.245 6.822 7 . 770 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)
=
snlt(nq)
(1)
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 .
A??(kcnl./mole)
0.5741 1.100 1.359 1 . ti23 1 ,746 I .784 2 . o"2
-0.82 - .79 - .82 - .79 - .75 - .i7 - .73
( 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
