1
k
X
4
-2
1 tc
1.79
1.80
1.81
1.82
1.83
IOOO/'K
Figure 7. Arrhenius plot of HMX vaporphase decomposition data Kinetics constants are as follows: E 52.9 kcal/mole, 2 = 1.51 = lozosec-l
=
kcal/mole, 2 = 3.14 x 1013sec-l; HMX, E = 52.9 kcal/mole, 2 = 1.51 X 1020sec-l.
to measure rate constants. Their assumption was correct'; the rate constant for the reaction in the vapor phase at 200 "C sec-l, but in the liquid phase the rate should be 5.5 X sec-' according to Robertson's constant should be 3.6 X constants ( E = 47.5 kcal/mole, 2 = 3.16 X lo1*sec-l) (6) or 3.0 x sec-l according to the liquid-phase values determined in this work ( E = 43.1 kcal/mole, Z = 2.44 x 1016 sec-1). Cosgrove and Owen presented data showing that approximately 0.02 gram of R D X decomposed at 195 "C in a 150-ml flask in one hour, the result being roughly independent of the sample mass. From our rate constants and vapor pressure, ignoring the contribution of decomposition in the solid or liquid phases, we would calculate that 0.007 gram should decompose over solid R D X in one hour or 0.015 gram should decompose over liquid R D X in one hour. Cosgrove and Owen noted that the material liquefied as the reaction proceeded. Therefore, the calculated decompositions should be low, because there would be some unknown contribution from decomposition in the liquid phase. However, the close agreement between our calculated values, which represent minimum estimates, and the measured values of Cosgrove and Owen show that decomposition in the vapor phase can be of primary importance in systems involving RDX.
Cosgrove and Owen ( 5 ) have shown that R D X decomposes much more rapidly in the vapor phase than in the solid phase. They also state that "The rate of reaction in the gas phase is a t least equal to that in the liquid phase," but they were unable
RECEIVED for review August 23, 1972. Accepted November 6, 1972. This work was performed under the auspices of the United States Atomic Energy Commission.
( 5 ) J. D. Cosgrove and A . J. Owen, Chern. Commun., 1968,286.
( 6 ) A. J. B. Robertson, Trans. Faraday Soc., 45, 85 (1949).
Thermodynamic Data for Aqueous Iodine Solutions at Various Temperatures An Exercise in Analytical Chemistry Joanne D. Burger Dow Chemical Company, Freeport, Texas 77541
Herman A. Liebhafsky' Department of Chemistry, Texas A&M University, College Station, Texas 77843
THEAVAILABILITY of the iodide selective electrode ( I ) suggests a new approach, interesting as regards analytical chemistry, to the important problem of determining equilibrium constants for three reactions in aqueous iodine solutions. These equilibria are
Iz + HzOe H I 0
+ Hi- + I-
(hydrolysis of Iz; ZG)
+ I- (dissociation of 13-; K2) 1%+ HzO % H201++ I- (hydrolytic ionization of 1 2 ; 13-
Iz
(1) (2)
K3)
(3) To whom correspondence should be sent. (1) J. H. (1969).
Woodson and H. A. Liebhafsky, ANAL.CHEM., 41, 1894
600
ANALYTICAL CHEMISTRY, VOL. 45, NO. 3, MARCH 1973
The new approach consists in measuring the activity of I(iodide electrode), of H+ (glass electrode), and the concentrations of Iz and Is- (spectrophotometer) before and after shifting the equilibria by dilution or by the addition of IB in known amount. At the low concentrations existing, activities and concentrations differ negligibly for each species. Only concentrations, represented by [ ] as moles,'liter, will be used henceforth. The calculation of K2 is straightforward as each of the three concentrations involved is measured directly. For the calculation of ZG, [HIO] must be known. This concentration is obtained by using the stoichiometric relations attending displacements of the equilibria. Dilution shifts Equilibrium 1 and Equilibrium 2 to the right. Increase of [I2], as by the dissolution of 12(s),shifts Equilibrium 1 to the right and Equilibrium 2 to the left.
Acceptable results are by no means certain. Fortunately, K2 is well known from many prior investigations (2-4) so that good agreement of new results and old shows the determinations of [I2], [I3-], and [I-] t o be reliable. Note, however, that all values of Kl and K, a t [I-] below 5(10W) were rejected; see Ref. ( I ) , Figure 4. The difficulties in the case of Kl are the obtaining of reliable [HIO] values and the possibility that the reverse Dushman reaction-namely 3 IP
+ 3 H20 = 6 H + + 5 I- +
103-
(4)
will interfere. Information about possible interference under our conditions is not available; see, however, Ref. (9,p p 2739 and 2740. The best that can be done is to see whether the Kl values show acceptable standard deviations and agree with the values in the literature, which are less certain and cover a smaller temperature range than those of K2. These tests were met. A less successful attempt to estimate K 3 a t 50 "C will be described later. EXPERIMENTAL
To determine Kl and K2, four solutions were prepared in sequence. Measurements of [I-] and [H'l were made o n all four; [I2] and [I3-] were determined o n the last three. The amounts of I? added were chosen to give satisfactory [HIO] values. To minimize the rate of Reaction 4, only unsaturated iodine solutions were used, and the experiments were completed in a matter of hours. The simple manipulations were carried out so as to minimize loss of iodine by volatilization, and t o avoid the presence of unwanted I~(s). Iodine and HCIOl were of reagent quality, not further purified. The strength of the 0.1M standard iodide solution was found by potentiometric titration to be 0.100 + 0.001. The distilled, deionized water contained no detectable halides. The four solutions were: SOLUTION1. [I-] and [H+] near 10-5, prepared from reagents named above. SOLUTION 2 . Solution 1 with appropriate amounts of 12(s) added and dissolved by shaking at 10-min intervals at thermostat temperature over about 3 hours. SOLUTION 3. 10 ml of Solution 2 diluted a t thermostat temperature to 100 ml in a volumetric flask. Exact value of dilution factor (DF) used in calculations. SOLUTION4. Duplicate of Solution 3 except that a n unsaturated iodine solution of known [I?] was used for dilution. Shaken a t 15-min intervals during 3 or 4 hours in thermostat. The following equations, derived from assumed stoichiometric changes, are available for the calculation of the [HIO] in the last three solutions (solutions identified by subscripts):
+
[H-], = DFLH-12 [HIO], (11) Kl values from Equations 5 and 6 were from 100 to 1000 times comparable values in the literature. Here [HIO] was (2) E. N. Rengevich and E. A. Shilov, Ukr. Khirn. Zh., 28, 1080 (1962); Chem. Absrr.. 59, 4590f (1963). (3) L. I. Katzin and E. Gebert, J . Amer. Chem. Soc., 77, 5814 (1955). (4) M. Davies and E. Gwynne, ibid., 74, 2748 (1952). (5) R . P. Bell and E. Gelles, J. Clrem. SOC.(Lorido/i),1951, 2734.
small relative t o the other concentrations and highly uncertain. Whether this circumstance accounted completely for the high results, or whether the occurrence of Reaction 4 contributed, is not known. At any rate, Equations 5 and 6 gave no Kl values worth reporting. In Equation 9, the last term represents iodine vapor. The term was relatively large and uncertain; it made Equation 9 useless. We are thus left with four values of [HIO] from each complete experiment that are useful for the calculation of Kl. Those from Equations 10 and 11 were usually the most concordant. The primary temperature standard was a pyrometer that could be read to 1 0 . 0 5 O C between 0 and 100 O C . All measurements were made a t thermostat temperature. Spectrophotometric determinations of [In] and [I3-] were done in cuvettes kept within 1 "C of thermostat temperature by means of circulated thermostat water. Electrode measurements of [H+] and [I-] were made in a small, magnetically stirred, water bath on a hot plate; the temperature uncertainty was again + l O C . Ice baths were used at 0 "C. Determinations were made in the following sequence: spectrophotometric, electrode, spectrophotometric. Agreement between successive spectrophotometric determinations was usually satisfactory. Spectrophotometric Measurements. These were done o n a Beckman D U spectrophotometer a t 461, 353, and 286 nm with a hydrogen lamp as light source. From the work of Awtrey and Connick (6), it is clear that
+
A = Ul,xbCi as,xbc2 (12) where A is the combined absorbance (at wavelength X with path length b) due to I? (molar extinction coefficient ul,x and concentration cl) and to 13- (molar extinction coefficient u2,x and concentration cy). For dilute solutions, b was 10 cm; for concentrated solutions, b was 1 cm. The u's of Ref. ( 6 ) were used. Estimated precision, 0.001 absorbance unit at low absorbances, decreasing to 1 0 . 0 1 above absorbances of 0.4. To establish whether a variation with temperature of the u's had to be considered, absorbance measurements were made a t 0, 25, 38, and 50 O C for the evaluation of c1 and c2 according to Equation 12, c1 being many times c?. The sum c1 c2 was compared with total iodine in these solutions as established by thiosulfate titration. For 7 of the 8 solutions measured, the results agreed to within less than 2.5%. For the most concentrated solution, in which [I,] spectrophotoc1 c2 by absorbance metrically determined was 14.3 x by titration, 15.7 X Any temperawas 15.0 X ture variation of the u's is negligible for present purposes. Determination of [H.I' An Orion Research Ionanalyzer Model 801 digital p H meter, a Corning Triple Purpose glass electrode, and a Corning calomel reference electrode were used. At the start of each experiment, the meter was standardized by use of solutions known to have p H values of 4.00 i 0.01 and of 7.01 i 0.01. Estimated precision of determinations, 1 0.03 p H unit. The glass electrode satisfactory for all measurements above 0 O C responded so sluggishly a t this temperature that a new one was substituted. In the first two experiments, the new electrode responded more slowly than normal, requiring 5 t o 10 min to reach a steady reading. It became so sluggish in a third experiment that further work a t 0 " C was discontinued. Determination of [I-]. The digital voltmeter and the reference electrode mentioned above were used in conjunction with a n Orion Specific Ion Electrode, Iodide Model 94-53. The information given below, obtained in part during the investigation of the oscillatory decomposition of HeO,, supplements that in Ref. ( I ) , which pretty well describes how the electrode was used here. In the iodide selective electrode, a mixed crystalline AgIAg2S phase is interposed as a n impermeable membrane be-
*
+
+
~(6) A. D. Awtrey and R . E. Connick, J . Amrr. Cl7em. S O ~ 73, . , 1842 (1951). ANALYTICAL CHEMISTRY, VOL. 45, NO. 3, MARCH 1973
601
Table I. Summary of Results for Reaction 2' kt103) ~~(103) A S' AG ' (entropy T, (New mean (Literature (kcal). units)& "C values) values) 0 0.791 0.704 ( 2 ) 3.86 -0.79 25 1.38 1.34(3) 3.88 -0.77 38 1.81 1.69( 4 ) 3.89 -0.77 3.90 -0.77 50 2.25 2.04 ( 4 ) ... 3.91 -0.79 56 2.47 Calculated from data in second column by use of A G " = A H o - TAS" and the value AH" = 3.65 kcal, new result based on leastsquares line in 1/T plot. Literature value, AH" = 3.6 kcal (3). 5
Table 11. Summary of Results for Reaction 1 I1
T, "C 0 25 38 50 56
Kl (1013) 0.59
5.44 18.9
41.2 68
(number of values) 4 30 22 27 8
AS
(entropy (kcal)c units)c -5.03 16.5 16.7 -5.25 -4.91 16.6 16.8 -5.17 16.8 -5.17 AGc
sa
srnb
0.35 1,773 6.8 12.0 27
0.17 0.32 1.5 2.3 9
All difficulties considered, agreement of the new results with those of Allen and Keefer (7), and with the older results, seems satisfactory. Hydrolytic Ionization. Bell and Gelles (5) proved the existence of Equilibrium 3 and obtained KS = 1.2 x 10-11 at 25 "C from electromotive force measurements. In order to measure KB,it is important to repress Equilibrium 1 by working in at least moderately acid solutions. Unfortunately, measurements have to be made at [I-] so low that the selective electrode cannot be trusted. Consequently, it is necessary to rely upon spectrophotometric methods for [I3-] and for [Iz], and to use these concentrations for calculation of [I-] from K?; and this [I-], and [Iz],for the calculation of [HIO] from Kl. This approach was followed with [HClOJ = 0.01 as solvent. Solution 5 (analogous to Solution 1) contained iodine participating in the three equilibria written above. For this solution,
Tenfold dilution of Solution 5 and addition of Iz gave Solution 6, analogous to Solution 4. For Solution 6, LI-16 = DF[I-Is
n-1
dividual determjnation (xJ of Kl. S~ = s / d n , the measured standard deviation for the mean, f = Ki. A H c = 15.1 kcal, new result based on least-squares line in = 0.41 at 0" 1/T plot. For comparison from Ref. (7): Kl(1013) and 5.4 at 25"; AH" = 16.7 kcal. tween the unknown iodide solution and an Ag/AgCl reference electrode in a suitable electrolyte. Conductivity results almost entirely from the migration of Ag+. When such an electrode is subjected to abrupt temperature changes of about 25 "C, peripheral cracks develop where the circular membrane meets the plastic housing. As a precaution, the electrode was brought slowly to the desired temperature and kept there for at least an hour before measurements were made. The response of the electrode changed with temperature and with time. A 35-mV increase in response at unchanged [I-] was observed as the temperature was changed from 0 to 50 OC. For [I-] = 2.5 X response decreased from -202 to -183 mV in two months. RESULTS AND REACTIONS 1 AND 2
Tri-iodide Dissociation. The new values of I?*, the mean of K z ; data from the recent literature; and values of thermodynamic functions calculated in the usual way for Reaction 2 are given in Table I. Reaction 2 was written as an equilibrium above. The agreement between new and prior results is satisfactory. Iodine Hydrolysis. Table I1 contains for Reaction 1 the kind of information given above for Reaction 2. Because the uncertainties in Kl are greater than those in K?, standard deviations for K1 are included. Comparison is with the spectrophotometric results of Allen and Keefer (7), more recent and probably more reliable than older data. (7) T. L. Allen and R. M. Keefer, J. Amer. Cizem. Soc., 77, 2957 ( 195 5).
602
+ [H201+16 + [HI016 + DF[I3-]6 -
, the measured standard deviation for an in-
ANALYTICAL CHEMISTRY, VOL. 45, NO. 3, MARCH 1973
lI3-16
(14)
from which [HtOIf16 (and hence K3)could be calculated. The result is K3 = (3 + 2), X at 50 "C, a tentative value that requires confirmation. The work was discontinued because [I3-], which lay in the range 10-7-10-8, could not be determined with satisfactory precision.
coNCLUSIO N The results in Tables I and 11 are more comprehensive than those of any prior investigation. The new approach seems worthwhile, and the thermodynamic functions for Reactions 1 and 2 should prove useful. The rejection of Kiand Kz valappears justified. In all ues based on [I-] below 5 X work done here (that on K S included), one must assume that Reaction 4 has not distorted the results. It is an advantage to work in solutions so dilute that the differences between activities and concentrations may be ignored. As the selective iodide electrode cannot be trusted in the determination of K 3 ,the approach becomes less advantageous in this case. The results reported here do not show anomalies in the temperature coefficients of equilibrium constants such as exist for other aqueous halogen equilibria (8). ACKNOWLEDGMENT We thank J. H. Brown, Jr., and W. L. Howard of the Dow Chemical Company for helpful advice, and the Company itself for the use of laboratory facilities. RECEIVED for review August 23, 1972. Accepted October 3, 1972. Financial support was extended to Joanne D. Burger by the Dow Chemical Company. We also thank the Robert A. Welch Foundation for support under Grant A-254. (8) H. A. Liebhafsky, Cizem. Rec., 17,89 (1935).
