Dec., 1960
STANDARD POTENTL4L
O F THE
SILVER-SILVER CHLORIDE ELECTRODE
1861
ELECTROMOTIVE FORCE STUDIES IK AQUEOUS SOLUTIONS AT ELEVATED TEMPERATURES. 111. THE STANDARD POTENTIAL OF THE SILVER-SILVER BROMIDE ELECTRODE AND THE MEAN IONIC ACTIVITY COEFFICIEXT OF HYDROBROMIC ACID' BY MICHAELB. TOWNS,^ RICHARD S. GREELEY AND &I. H. LIETZKE Chemistry Division, Oak Ridge Nataonal Laboratory, O a f Ridge, Tennessee Receoved April 9 , 1960
The e.m.f. of the cell Pt-Hz(p) jHBr(m)l AgBr-Ag was measured from 25 to 200" using hydrogen pressures of about 1atm. and HBr concentrations from 0.005 to 0.5 m. Additional measurements were made on 1.0 m HBr from 25 to 150 . The standard potential of the cell was determined and found to fit the following equation with a standard error of fit of 1.1 mv. Eo = 0.08289 - 4.0647 X 10-4t - 2.3986 X IO-etZ volts. The mean ionic activity coefficients of HBr were calculated for several concentrations from an extended Debye-Huckel equation, the linear ( B ) parameter of which was obtained from a least squares treatment of the e.m.f. data.
Introduction The standard potential of the silver-silver bromide electrode has been determined from 0 to 60' by Harned, Keston and Donelson3 from measurements of the cell Pt-H&)
jHBr(m)j AgBr-Ag
(A)
but no measurements a t higher temperatures have been reported. I n connection with a general program on the properties of aqueous solutions a t elevated temperatures and, in particular, following a study of the chloride cell analogous to (A) a t temperatures to 275°,4Jit was of interest to extend the measurements of cell (A) to as high a temperature as possible. Therefore cell (A) was investigated over the concentration range 0.005 to 0.5 m HBr, the temperature range 25 to 200', and at hydrogen pressures of about one atmosphere. Additional measurements were made on 1.0 m HBr to 150'. Experimental The Apparatus.-The experimental apparatus has been described previously4 and was used without modification. The electrodes.-The hydrogen electrode was a length of platinum wire platinized prior to each run according to the procedure given by Bates.6 The silver bromide electrodes were of the thermal type made by decomposing a 7: 1 mixture by weight of silver oxide and silver bromate on a platinum wire a t 650" in a manner similar to the method of Keston.7 Enough silver bromide was deposited on the wire to provide an excess over the amount estimated to be dissolved by the solution a t 200". I n addition, excess silver bromide was added to the cell to ensure saturation. The platinum base wire of each type of electrode was long enough to pass through the bomb head, obviating the necessity of crimping the electrodes to a lead wire. The same pair of silver-silver bromide electrodes was used for the entire set of experiments. They were checked at '25' prior to each run by comparing the cell e.m.f. with that given by Harned, Keston and Donelson3 for the particular solution in use. Agreement was always within f0.3mv. (1) This paper is based upon work performed for the United States Atomic Energy Commission at the Oak Ridge National Laboratory operated by Union Carbide Corporation. (3) (a) Oak Ridge Institute of Nuclear Studies summer participant. (b) Department of Chemistry, Tennesaee A. and I. College, Nashville, Tennessee. (3) E.8. Earned, -4. S. Keston and J. G. Donelson, J . Am. Chem. Soc., I S , 989 (1936). (4) R. S. Greeley, e t aZ., THISJOURNAL,64, 652 (1960). (5) R. 9. Oreeley. e l al.. to be published. (6) R. G. Bates, "Electrometrir. p H Determinations," John Wiley and 6011s. 100.. New York, N. Y., 1954, p. 107. (7) A. 8. Keston, J . A n . CImm. Soc., 67, 1571 (1935).
The Solutions and Materials.-All solutions were made from conductivity water, all apparatus was given final rinsings with the same, and all chemicals were recrystallized or washed in conductivity water as deemed necessary. The hydrobromic acid solutions were made up by weight dilutions from twice-distilled constant boiling hydrobromic acid which had been analyzed gravimetrically and found to agree with literature values.* The silver oxide was that made previously for the chloride work.4 Silver bromate "as made according to the procedure given by Keston.? Electrolytic hydrogen obtained in commercial cylinders was passed over Ascarite, Drierite, and a platinized catalyst bed before passing through two bubble towers containing solutions identical to that under test and then into the autoclave. General Procedure.-The procedure for each test was essentially the same as that for the chloride work.4
Results and Calculations E.m.f. measurements were made on 0.005, 0.01, 0.02, 0.05, 0.075, 0.1 and 0.5 m HBr a t 25, 60, 90, 125, 150, 175 and 200' and on 1.0 m HBr a t the above temperatures to 150'. Duplicate tests were made with each solution and the e.m.f. values taken a t the same temperature were reproducible to within about f 0.3 mv. at 25', f 0.5 mv. a t 60 and 90' and f 1 mv. above 90'. Measurements were made only during ascending temperatures except for a final measurement a t room temperature. Agreement between initial and final values at 25' averaged 1 mv. No drift of e.m.f. with time as observed in the chloride work4 mas encountered. Limiting the temperature to 200' (150' for 1.0 m HBr) and limiting the time a t each temperature to the minimum necessary to reach thermal equilibrium resulted in negligible errors due to corrosion of the autoclave head or reaction between hydrogen and silver bromide. No silver mas detected on the hydrogen electrode after any run. In the calculations, the solubility of AgBr was neglected and the mean molality and ionic strength was taken to be equal to the HBr molality. It was assumed that the ratio of solubility of AgBr to AgCl at 100 to 200" was roughly the same as a t 25 to 100" ( v i z . , about l:lO), which would make a contribution to the ionic strength of less than 1% for all molalities of HBr studied to 200". No correction was made for the loss of water or of HBr to the vapor space. The hydrogen pressure was obtained as previously* by subtracting the corrected steam pressure from the observed total pressure. Each e.m.f. value was then corrected to 1.00 atm. H1 by subtracting ( R T / 2 F ) In f ~ where * the fugacity f ~was ? taken to be equal to the hydrogen pressure. (8) "International Critical Tables," First Ed., Vol. 111, McGrawHill Book Co., Ino., Kew York, N. T.,p. 323.
M. B. TOWNS, R. S. GREELEY AND M. R. LIETZKE
1862
The e.m.f. values, taken at temperatures slightly different from the exact values listed above, were corrected to those nominal values by fitting the values at each molality to a cubic equation in the centigrade temperature by the method of least, squares and solving the equations a t the temperatures desired. The standard errors of fit of the cubic equations representing the data were all about 1 mv. or lese. The further treatment of the data involved the computation of EO" for each e.m.f. value
Vol. 64
TABLE I1
MEANIONIC ACTIVITYCOEFFICIENTS OF HYDROBROMIC ACID 60° 150' 175" 200' pn 25' 90" 125' 0.001 0.965 0.963 0.960 0,956 0.953 0.949 0.945 .005 ,929 ,925 .919 .913 .906 .900 .894 .01 .905 .901 .896 .886 .878 .870 .863 .02 .875 ,871 .864 .854 .844 ,834 ,826 .821 .808 .792 .779 .769 .05 .830 .826 ,075 .809 ,804 ,798 .787 .768 ,752 .739 .1 .794 .789 .781 .773 .750 ,731 .715 .618 .580 .5 .808 .819 ,751 ,708 .662 1 .O .873 .852 .773 ,717 ,653 . . . ...
where
E = e.m.f. a t 1.00 atm. H2 and nominal temp. S = Debye-Huckel limiting slope A = denominator coefficient = 50.29 (DZ')-l/zdp~'/*
D
= dieletric constant of water9 d = ion-size parameter PO = density of water10 and m,R , 5 , and T have their usual meaning The values of A were calculated using values of d taken from the previous work on HC1 a t each temperature' for the reason discussed below. Then EO" was fitted by the method of least squares for 12 cells over the range 0.005 to 0.1 m HBr to
where Eo = the standard potential of the silver-silver bromide electrode. Values of EO, B , their standard errors, and the standard errors of fit of equation 2 to the data are listed in Table I. The values of Eo were fitted to a quadratic function of the centigrade temperature by the method of least squares to yield EO = 0.08289 - 4.0647 X IO-*t 2.3986 X 10-et2 volts (3) wit.h a standard error of fit of 1.1 mv. Values of EO calculated from equation 3 are list,ed in the final column of Table I.
-
VALUESOF Eo, B Temp., OC. 25 60
90 125 150 175 200
Eo, V.
TABLE I STANDARD ERROR OF FIT OF EQUATION 2
A N D THE
B,
CEO,
V.
m-1
CB,
m-1
Eo."
mit, V.
V.
4-0.0716 0.00067 +0.112 0.096 0.00079 $0.0712 ,0501 .00084 f .0323 .I60 .0012 ,0499 f ,0251 .0012 - ,012 .I80 .0012 .0269 ,0048 .00077 f .060 .089 .0013 ,0054 - ,0312 .00048 - .Os6 .053 .00081 - .0320 - ,0612 ,00067 .146 ,067 .0010 .0617 - .0951 ,0011 .277 ,106 .0017 ,0943
+ -
+ +
-
-
These values are given in Table 111. Also listed in Table I11 are the average experimental values of the e.m.f. for the 0.5 and 1.0 m HBr solutions.
Discussion Values of Eo" were calculated using six values for the denominator parameter, A , from 0 to 5 for each e.m.f. measurement and these were plotted us. concentration a t each temperature. It was observed that good straight lines were obtained for extrapolation to infinite dilution when the A values corresponding to those used in the chloride work4 were used. E.m.f. values of sufficient precision were not obtained to allow the use of the statistical procedure of minimizing the standard error of fit of equation 2 to determine A , as in the chloride work. The values of Eoobtained at 25 and 60' agree with those of Harned, Keston and Donelson3 within 0.3 and 0.5 mv., respectively. The values of Eo obtained from 25 to 200' fitted a quadratic function of the temperature within experimental error (see equation 3). However a cubic equation was required to fit the experimental e.m.f. values to a function of the temperature. The decrease with temperature in the value of Eo for the Ag-AgBr electrode is slightly greater than the decrease of Eofor the Ag-AgC1 electrode. Upon subtracting the Eofor the former from the Ea for the latter a t each temperature it is found that the Eofor the metathetical reaction AgCl
+ HBr = AgBr + HCl
(7)
decreases from 0.150 volts a t 25' to 0.130 volts at 200O. " Eo = 0.08289 - 4.0647 X l@-4t- 2.3986 X 10-etZ The mean ionic activity coefficients of HBr devolts. crease with temperature a t each molality in a very similar way to y f for HCl. However the activity The mean ionic activity coefficients of hydrobromic acid coefficients for HBr are generally slightly higher from 0.001 to 0.1 m were c,alculated from the Debye-Hiickel than for HC1, becoming one to t x o per cent, higher extended equation ._ at 200'. Agreement with the values listed by Harned, Keston and Donelson3 a t 25 and 60' was logy& = - ' d m +Bm (4) 1+Ay/& within one per cent. As with the chloride work, the tising the above values of A and B . The activity coefficients Debye-Huckel extended equation with a linear, for 0.5, and 1.0 m HBr were calculated from Bm, term was found to be entirely applicable. The extended terms of Gronwell, LaMer and Sandved" log-+ = ___ ( E - EO) - log m (5) were neglected, as was the correction from the 4.606RT where values of E were available. All of these values are rational to the practical activity coefficient scale, listed in Table 11. since these terms were essentially within the experiFinally smooth values of the e.m.f. of cell (A) a t round mental error and since the empirically determined molalities frcm 0.001 to 0.1 m were calculated for several parameters included such corrections. temperatures from The experimental error in this work was sub( 9 ) G . C. Akerlof and H. I. Oshry, J . Am. Chem. SOC.,72, 2844 stantially larger than in the chloride work, viz., * 1 (1860). (10) N. A. Lange, Ed., "Handbook of Chemistry," Seventh Ed. Handbook Publishem, Inc., Sandusky, Ohio, 1949.
(11) T. H.Gronwell, V. K. LaMer and K. Sandved, P h w i k . Z., 39, 358 (1928).
EFFECT OF STRUCTURE OF N,N-DISUBSTITUTED AMIDESON EXTRACTION
Dee., 1960
1863
TABLE I11
SMOOTHEDVALUESOF TJXEE.M.F. FOR
THE
CELL
Pt-H2(p) IHBr(m)( AgBr-Ag Temp., OC.
0.005
0.01
0.02
0.05
25 GO 90 125 150 175 200
0.3476 ,3588 .3620 ,3651) .3623 .3561 .3462
0.3133 ,3205 ,3202 ,3195 .3141 .3052 .2923
0.2795 ,2826 ,2791 .2745 .2664 .2549 ,2396
0.2351 ,2331 ,2249 ,2154 ,2043 ,1894 .1707
and i0.4 inv., respectively. For this reason the change of y i with temperature and the relative partial molal heat content and heat capacity were not calculated. Acknowledgment.-We wish to thank Professor W. T. Smith, Jr., and Dr. R. W. Stoughton for
mHBr
0.075
0.1
0.5
0.2156 ,2113 .2013 .1894 .1769 ,1608 .1408
0.2018 .1959 .1847 ,1709 ,1577 ,1408 ,1201
0.1180 ,1012 ,0862 .0662 .0492 .0293 .0057
1.o 0.0784 ,0591 .0178 .0004 - ,0020
-
......
......
many helpful discussions during the course of the work. We also wish to thank Mrs. Marjorie Lietake for carrying out the computer calculations, Mrs. Laura Meers for assistance with the hand calculations, and Mr. Raymond Jensen for preparing the HBr solutions.
EFFECTS OF STRUCTURE OF N,N-DISUBSTITUTED AMIDES ON THEIR EXTRACTION OF ACTINIDE AXD ZIRCONIUM NITRATES AND OF NITRIC ACID' BY T. H. SIDDALL, I11 Savannah River Laboratory, E. I . du Pont de Nentours & Go., Aiken, South Carolina Received April 18, 1960
The effects of changes in the structure of 21 N,Y-disubstituted amides on the extraction of actinide and zirconium nitrates and of nitric acid are given. Successive alkyl substitution on the cu-carbon atom greatly decreases the extraction of quadrivalent actinides and of zirconium, but only moderately decreases the extraction of hexavalent actinides and nitric acid. The extraction mechanisms for the extraction of uranyl nitrate and nitric acid are very similar to the mechanisms of extraction by trialkyl phosphates and dialkyl alkylphosphonates. However, the extraction mechanisms for the quadrivalent species by amides appear to be quite different from the mechanisms with the phosphorus compounds. A connection between the stretching frequency of the carbonyl bond in the amides and their power as extractants is substantiated in a general and qualitative way, but an exception is noted.
Introduction
extended, with emphasis on the effects of altering
It is to be expected that the disubstituted amides the hydrocarbon substituents in the amide mole-
should be strong extractants. The carbonyl cule on the extraction of the nitrates of quadrivastretching frequency of these amides is considerably lent and hexavalent actinides, zirconium nitrate lower than that for ketones. The lower frequency and nitric acid. has been ascribed t o resonance2 Experimental (I) R-
1
--N
/"
k
6 (11) R-
L=N
R
+'
R'
The contribution of I1 should not only weaken the carbonyl bond, but should also increase the availability of the electrons of the oxygen atom for bond formation. As a consequence amides should be somewhat stronger extractants than ketones, ketones being only moderately strong extractants. Federa coniirmed this prediction by showing that N,N-dibutylacetamide is roughly comparable to tributyl phosphate as an extractant for uranyl nitrate. In this work Feder's investigations were (1) The information contained in this article was developed during the course of work under contract AT(07-21-1 with the U. 9. Atomic Energy Commission, (2) A . Weisaberger, "Technique of Organic Chemistry," Vol. IX, Intericience Publishers. New York, N. Y., 1949, p. 523. (8) E. M. Feder. Argonne National Laboratory, ANL 4675, pp. 66-60, July 50 1951 (Classified Report).
Most of the disubstituted amides were prepared by combining equimolar quantities of the desired acid anhydride and disubstituted amine. The reaction was completed by letting the solution stand overnight, or by warming the solution to 50-60" for a few hours. When the acid anhydride was not available it waa prepared by exchange of the carboxylic acid with acetic anhydride at 140-170". The acid anhydride and residual acetic anhydride were then separated by fractional distillation. The formamides were prepared by dehydration of the amine salts of formic acid at 190-200" for 4-6 hours. As a matter of interest, only very small yields of amide were obtained when the dehydration procedure rather than the anhydride procedure was attempted with pivalic acid. N,N-Dibutylacetamide and N,N-diethyldecanamide were purchased from Distillations Products Industries, Eastman Kodak co. In all cases the crude reaction mixtures of the amides were given preliminary purification by washing with aqueous caustic, hydrochloric acid and water. Most of the amides were given final purification by crystallizing the solid adducts formed with uranyl nitrate. After crystallization the uranyl nitrate was removed by washing with water and aqueous sodium carbonate. Solvent and water were removed under vacuum. Solid adducts were formed with uranyl nitrate more or less readily by all the amides investigated. In order to get
