Hydrolysis of iodine: equilibria at high temperatures Page: 2 of 7
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I20H~ * 1202- + H+
H20 * 0H + H+
I2(vapor) * I2(aqueous)
It should be noted that in our initial model3,
which was based solely on the avialable litera-
ture data, species such as H103, H20I+ and
I202- were found to be insignificant (i.e.,
<0.01% of the total iodine present) under the
conditions pertinent to nuclear accident sce-
narios, viz. initial iodine and iodide con-
centrations of 10-3 to 10-7 gatom/L, pH values
from 5 to 10, and temperatures from 25 to
150*C. However, the equilibria pertaining to
the species are retained in the model for the
sake of completeness. As new experimental
results for equations (1), (3) and (5) will be
presented in this paper, a brief account of the
origin and form of the expressions used to
describe the temperature dependence of the
remaining equilibrium constants is needed.
The combined data4-8 for equation (2),
spanning a temperature range of 0 to 56*C, were
fitted by the equation,
ln K2 = -1392.9/T - 44.764 + 0.070692T
where T represents the temperature in degrees
Kelvin and the subscript refers to the reaction
to which the equilibrium constant is assigned.
Equation (4) was modeled after the tem-
perature dependence for the acid dissociation
of HOBr9, but adjusted to give the measured
value of K4 for HOI at 25*C.10 The equilibrium
expression for this equation is as follows:
log K4 = 29688/T + 81.840 In T - 0.089649T
- 2.046B x 106/T2 - 526.75
Data are available8,11 at only two tem-
peratures for equation (6) and hence a linear
expression must be used. However, because the
charge is the same on both sides of this
equation, the heat capacity change should be
minimal1. This point will be elaborated on
in the ensuing discussion.
log K6 = -2.92/T + 8.27
The heats of reaction for equations (7) and
(8) are unknown and therefore the equilibrium
constants were fixed at their 250C values,
1.4 x 10-10 and 0.045, respectively.10
The dissociation constant for water, Kg, was
taken from Marshall and Franck13, i.e.,
log Kg = -4.098 - 3245.5/T + x.2362 x 105/T2
- 3.984 x 107/T3 + (13.957 -1262.3/T
+ 8.5641 x 105/T2) log Dw
where Dw represents the density of water and is
calculated4 as follows:
Dw = 1.0017 - 2.36582 x 10-5t - 4.77122
x 10-6t2 + 8.27411 x 10-9t3
where t is the temperature in "C.
The distribution coefficient for iodine bet-
ween the aqueous and vapor phases has been
measured by numerous experimentalists.7.15 The
results of these two studies15,16 can be best
fit by the equations:
log K10 = 4220.5/T - 19.991 + 0.02583T
(t < 112*C)
log K10 = 5615.4/T - 25.1798 + 0.02990T
(t > 112*C)
As HOl has not been detected in the vapor phase
as yet, only lower limits can be suggested for
its distribution coefficient (e.g., 104 at
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Palmer, D.A.; Ramette, R.W. & Mesmer, R.E. Hydrolysis of iodine: equilibria at high temperatures, article, January 1, 1984; Tennessee. (https://digital.library.unt.edu/ark:/67531/metadc1207811/m1/2/: accessed April 25, 2019), University of North Texas Libraries, Digital Library, https://digital.library.unt.edu; crediting UNT Libraries Government Documents Department.