1929
THERMODYNAMICS OF AQUEOUS SOLUTIONS OF HYDRIODIC ACID
diffusion periods required. Chronopotentiometric measurements in Ca(XO& K N 0 3 solvents seem a more promising approach in this case. Finally, we note the interesting prospect of using the glass-forming associated halides such as ZnCh and the Be halides as a bridge for exploring the rela-
+
tionship between simple molten salts and glass-forming oxide melts.
Acknowledgment. The author is indebted to Mr. G. M. Willis for many valuable discussions, and to Professor H . W. Worner for his interest in the work.
Thermodynamilcs of Aqueous Solutions of Hydriodic Acid from Electromotive Force Measurements of IIydrogen-Silver Iodide Cells
by Hannah B. Hetzer, R. A. Robinson, and Roger G. Bates National Bureau of Standards, Washington, D. C.
(Received February T& 1964)
Electromotive force measurements of the cell Pt; Hz(g), HI(m), AgI; Ag have been made a t eleven temperatures from 0 to 50'. The standard e.m.f. (EO) is given within 0.05 mv. by the equation E o = -0.15242 - 3.19 X 10-4(t - 25) - 2.84 X 1Op6(t - 25)2, where 1 is the temperature in degrees Celsius. These values are in excellent agreement with those obtained by Owen from studies of borax-buffered K I solutions at 5, 10, 30, 35, and 40' but differ by 0.14 to 0.17 niv. at 15, 20, and 25'. The activity coefficient (yk) of H I a t molalities (m)from 0.005 to 0.9 has been derived. The relative partial molal enthalpy (E,) of H I at 25' was calculated and compared with that for HC1 and HBr. At 25' and m = 0.1, y* is 0.811, and E, is 130 cal. mole-'.
Introduction From measurements of the e.m.f. of the cell Hz; HI, AgI; Ag containing 0.0980 and 0.0300 m hydriodic, acid, Noyes and Freed' deduced a value of - 0.14778 v. for the standard potential of the silver-silver iodide electrode at 25". The value given in the International Critical Tables2is -0.151 v. which, according to Owen,3 is based on the data of Pearce and Fortsch.* Owen himself compared the e.m.f. of the cell 11,; borax buffer and KX, AgX; Ag, with X = C1 or 1; he made measurements a t intervals of 5" from 5 t o 40" and found E" a t 25" to be -0.15230 v. Hass and Jellinek5 measured the e.m.f. of the cell with liquid junction Ag; AgI, MX, satd. KC1, 0.1 N KCI; calomel, where M = Na or K and X = C1 or I. From their data we celculate E o = -0.1502 v. at 25".
Finally, Cann and Taylor,e using measurements of a cell with a flowing liquid junction, arrived a t a value of E o = -0.1510 v. a t 25", although their value of E o == -0.1507 v. obtained by another method of cxtrapolation might be preferred. The considerable variation among these several values of E" a t 25" led us to investigate the cell Hs; Hl(m), Agl; Ag, where m is molality, in some detail from 0 to 50". We have found E" = -0.15244 v. a t (1) A. A. Noyes and E. 9. Freed, J . Am. Chem. Soc., 42, 476 (1920). (2) "International Critical Tables," Vol. V I , McGraw-Hill Book Co., Ino., New York, N. Y., 1929, p . 332. (3) B. B. Owen, J . Am. Chem. SOC.,5 7 , 1526 (1935). (4) J. N. Pearce and A. R. Fortsch, ibid., 45, 2852 (1923). (5) K. Hass and K. Jellinek, 2 . Phvsik. Chem., 162, 153 (1932). (6) J. Y. Carin and A. C. Taylor, J . Am. Chem. Soc., 59, 1484 (1937).
Volume 68, .\'umber
7
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1930
25". The mean activity coefficient and relative partial molal enthalpy of hydriodic acid have also been determined. Experimental The experimental procedure was essentially that used earlier in this laboratory for determining the standard potentials of the silver-silver chloride' and silver-silver bromide8 electrodes. However, because of the ease of oxidation of hydriodic acid with liberation of free iodine, considerable precaution had to be taken to exclude air, and all solutions were prepared and handled under an atmosphere of nitrogen. I n order to avoid contact of air with the solutions, the apparatus described by Bates and Acreeg was used for the weight dilution of the concentrated hydriodic acid stock solutions. A solution of hydriodic acid, "special reagent" grade, nearly colorless, free from preservatives such as hypophosphorous acid, and having the composition of theazeotrope (57 wt. % HI), was obtained commercially. This material was refluxed over a small amount of red phosphorus and then distilled in an atmosphere of nitrogen. A distillation of this type was made immediately before each run, and the middle fraction of the product was used. Stock solutions were prepared by diluting the distillate with oxygen-free water and their concentrations determined by gravimetric analysis as silver iodide. These stock solutions were then diluted further with oxygen-free water to make the cell solutions. The silver-silver iodide electrodes were of the thermal type, formed on platinum helices by heating a paste of 10% silver iodide and 90% silver oxide a t 450" for 15 min.l0 The two solids were ground together and the paste formed by adding small quantities of water. Two coats of paste were applied to the electrodes, with heating to 450" after each application. The finished electrodes were light gray in color and approximately 4 mm. in diameter, The individual electrodes of each set agreed among themselves to better than 0.08 mv. after conditioning for 6 to 8 hr. in an oxygen-free 0.02 m solution of potassium iodide. No greater differences were observed among electrodes prepared from freshly precipitated silver iodide and those prepared from silver iodide that had been stored in the dark for several years. Two hydrogen electrodes and two silver-silver iodide electrodes were mounted in each cell and a stream of nitrogen saturated with water vapor was passed through the cell for 1 hr. Cell solution was then forced in under hydrogen pressure and the cell rinsed four times before it was finally filled for the e.m.f. measurements. The Journal of P h w i c a l Chemistry
H. B. HETZER, R. A. ROBINSON, AND R. G. BATES
A complete series of measurements with the more dilute solutions required about 70 hr. The e.m.f. at 25" was usually measured three times, namely a t the beginning, in the middle, and a t the end of each series. Of the 18 cells for which three values a t 25" were obtained, the mean difference between the average of the three and the first, second, and third readings was 0.04, 0.05, and 0.05 mv., respectively. The average of these three values is recorded in Table I. Pairs of electrodes in any one cell agreed on the average within 0.08 mv. a t 0", 0.04 mv. a t 25", and 0.02 mv. a t 50". The results were discarded if electrode pairs showed poor agreement or if the e.m.f. at 25" increased with time.'l Yellowing of the solution usually accompanied the increase of e.m.f. and was also a cause for rejection of the results. The life of cells containing the more concentrated acid solutions (>0.2 m ) was much shorter than that of cells containing dilute acid solutions, probably because of the increased solubility of silver iodide in concentrated iodide solutions.10~12For example, the e.m.f. of cells containing acid of molality 0.5 or greater decreased markedly after about 4 hr., and agreement between electrode pairs was poor. Nevertheless, these cells came to equilibrium quickly and satisfactory e.m.f. readings could be obtained in 3 hr. at any one temperature. Measurements at one temperature were therefore followed by a check at 25" and the cell was discarded. Standard Potentials Table I contains the results of e.m.f. measurements of 18 cells a t intervals of 5" from 0 to 50" and, in addition, of seven cells which were measured at 25" only. Data for 10 solutions of molality greater than 0.2 are also included. All values of the e.m.f. have been corrected to a hydrogen partial pressure of 1 atm. The standard e.m.f. of the cell was found by extrapolation of the function
to m = 0. In eq. 1, k is written for (RT In 10)lF. For purposes of extrapolation, E " , computed from all (7) R. G. Bates and V. E. Bower, J . Res. Natl. Bur. Std., 53, 283
(1954). (8) H. B. Hetzer, R. A. Robinson, and R. G. Bates, J . Phys. Chem., 6 6 , 1423 (1962).
(9) R. G. Bates and S. F. Acree, J . Res. Natl. Bur. Std., 30, 129 (1943). (10) R. G. Bates, J . Am. Chem. Soc., 6 0 , 2983 (1938). (11) See J. K. Taylor and E. R. Smith, J . Res. Natl. Bur. Std., 22, 307 (1939). (12) W. Erber, 2. anorg. allgem. Chem., 248, 32 (1941).
THERMODYNAMICS OF AQUEOVS SOLUTIONS OF HYDRIODIC ACID
1931
Volume 68, Zumber 7
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H. B. HETZER, R. A. ROBINSON, AND R. G. BATES
1932
the data at m < 0.16, was expressed as a linear function of m by the method of least sqyares. Values of a = 4.3 and 4.4 A. were used in the extrapolation of data for the corresponding hydrochloric acid cells (0 to 30")' and hydrobromic acid cells (0 to 50°),8respectively. Since the crystal radius of the iodide ion is 0.35 8.higher than that of the chloride ion, it seemed reasonable to use a value of 4.65 8. for a in the extrapolation of the data for hydriodic acid. This value gave a satisfactory linear extrapolation even a t 50". Table I1 gives the standard e.m.f. E", the standard deviation n1 of this quantity, and the values o f F o b tained by OwenJ3 corrected to absolute volts. The average difference between the present results and those of Owen is 0.08 mv. ; the agreement is very good a t the extreme temperatures, 5, 35, and 40", and somewhat less satisfactory at t,he intermediate temperatures where one might have expected better agreement.
Table 11: Standard Electromotive Force ( E " ) of the Cell HP; HI(m), AgI; Ag from 0 to SOo, in v. ,Si,
E O ;
OC.
EO. this work
mv.
Owen'
0 5 10 15 20 25 30 35 40 45 50
-0,14637 -0.14719 -0.14822 - 0.14942 -0,15081 -0.15244 -0 15405 -0.15590 -0.15788 -0.15998 -0.16219
0.04 0.04 0.03 0.03 0.03 0 03 0.03 0.03 0.03 0.04 0.04
.. -0.14717 -0 14810 -0.14925 -0.15067 -0.15230 -0,15401 -0,15591 -0,15792
t,
,
I
.
, . .
The E" values given in Table I1 can be represented by the equation
E"
=
-0.15242 - 3.19 X lOP4(t - 2 5 ) 2.84 X lO-'(t
- 25)a
Activity Coefficients The mean molal activity coefficient y& of hydriodic acid a t 10,26, and 40" was calculated by the equation =
E o - 2lc log m - 2lc log yi
(2)
For concentrations less than 0.1 m, it was found convenient to utilize values of E"' from eq. 1, interpolated a t round values of the molality. Equations 1 and 2 The Journal of Physical Chemietrg
+
For solutions more concentrated than 0.1 m, E"' was again fitted by the method of least squares to an equation linear in m. The slope was, however, slightly different from that characteristic of the dilute range. Values of E"' a t round molalities were then interpolated and the activity coefficients between 0.1 m and 0.9 m were calculated by eq. 3. The results are given in Table 111. Harned and R o b i n ~ o n 'used ~ isopiestic vapor pressure measurements to obtain the activity coefficient of hydriodic acid a t 25". Their values were referred, somewhat arbitrarily, to yh = 0.818 a t 0.1 m. We have found that a choice of yi = 0.798 at 0.1 m brings their results in good agreement with our own over most of the concentration range where the comparison can be made. These revised values are given in Table 111. It may be seen that the agreement is good a t concentrations from 0.3 to 0.9 m, fair a t 0.2 m, and poor a t 0.1 m. As Harned and Robinson had to use platinum dishes in their isopiestic work (silver dishes being corroded by hydriodic acid) and as they also experienced difficulties with the instability of their solutions, the disagreement a t 0.1 m is not surprising. The good agreement a t higher concentrations is gratifying, and it is therefore recommended that the activity coefficients given by Harned and Robinson for solutions from 1 to 3 m also be multiplied by the factor 0.976. Using our values of E " , we have calculated activity coefficients for seven solutions of hydriodic acid from the e.m.f. data given by Pearce and Fortsch4 a t 25". We have also obtained values a t these same concentrations by interpolation in our own data, using the method described at the beginning of this section. The results are listed in Table IV. The agreement at the four highest concentrations is reasonably good, but there are considerable differences for the three most dilute solutions.
Relative Partial Molal Enthalpy
with an average deviation of 0.05 mv.
E
can be combined to give the following expression for 1% Yi Eo' - E o (3) -log Y& = 1 Bad; 2k
The relative partial molal enthalpy, Et, of hydriodic acid is also given in Table 111 and a comparison made with values of already published for hydrochloric acid14and hydrobromic acid.ls The values are given in thermochemical calories per mole; 1 cal. = 4.1840 joules.
zz
(13) H . S. Harned and R. A. Robinson, Trans. Faraday Soc., 37, 302 (1941). (14) H. S. Harned and R. W. Ehlers, J . A n . Chen. SOC.,5 5 , 2179 (1933).
1933
THERMODYNAMICS OF AQUEOUS SOLUTIONS OF HYDRIODIC ACID
of H I , HBr, and HCl a t 25"
Table 111: Activity Coefficients of H I a t 10, 25, and 40";
0.005 0.01 0.02 0.05 0.1 0.2 0.3 0.4 0.5 0.6 0.7
0.8 0.9 1.0 5 This investigation. and Ehlers.I4
6
0,931 0,909 0.882 0,843 0.815 0.799 0.803 0.814 0.831 0.851 0.874 0,900 0.927
0.929 0.907 0.897 0,839 0.811 0.793 0.795 0.805 0.820 0.839 0.861 0.885 0.911
, . .
, . .
0 0 0 0 0 0 0 0 0 0
798 788 792 803 819 839 862 886 913 940
"/**
P and F
this work
0 0 0 0 0 0 0
0 0 0 0 0 0 0
Y*
0 0 0 0 0 0 0
00505 01045 01981 05049 07914 12972 24608
41 60 77 106 130 170 202 ,226 244 263 28 1 297 315 , . .
I
908 884 851 838 823 802 802
929 905
880 838 819 802 793
A simple means of calculating such enthalpy data when activity coefficients are available a t three equally spaced temperatures (TI, T,, and T 3 )is as follows. Let
45 63 85 130 175 242 ...
47 64
85 124 163 222
...
...
t . .
391
338 , , .
,
...
.
I
, . .
...
, . .
...
...
488
61 1
Revised data of Barned and R ~ b i n s o n . ' ~ Data of Harned, Keston, and Donelson.'6
Table IV : Activity Coeflicient of Hydriodic Acid a t 25"
m
0.928 0.904 0.876 0.835 0,806 0.788 0.789 0.798 0.813 0.832 0.853 0.877 0.903
d
Data of Harned
X be defined by X -= T log yi, where T is the temperature in degrees Kelvin. Assuming that X = A BT CTZ,then a t temperature T zis given by
+
+
z2
for a 1:1 electrolyte. The variation of with acid molality is most pronounced with hydrochloric acid and least pronounced with hydriodic acid. Hydrobromic acid is intermediate between the two, with values somewhat greater than the mean of the values for the other two acids. (15) H. S. Harned, A. ~ o c . 5, 8 , 989 (1936).
S. Keston, and J . G. Donelson, J . A m . Chem.
Volume 68, Number 7
J u l y , 1964
