Hydrolysis of iodine: equilibria at high temperatures Page: 3 of 7
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2.1 EMF Measure its
The cell used t determine the equilibrium
constant for equation (3) can be represented as
P1 I2(mL),KI(mLl),HCI04(mL2) HC104(mR2),
Details of the equipment and procedure are
given elsewhere.17 Essentially the cell con-
sisted of two concentric compartments, each
containing the solutions shown above. Three
experiments were carried out with 12 solutions
of simply HI and HIO3, respectively. The ionic
strengths of these solutions, generally <0.1 m,
were matched as closely as possible to minimize
the liquid junction potential. The actual
liquid junction consisted of a porous glass
plug. Platinum electrodes completed the cir-
cuit with the electrode in the iodate solution
being coated with platinum black. The solu-
tions were stirred magnetically throughout the
experiment. The cell was initially purged with
argon and kept lighttight.
The key to this method was to allow the
iodine to distribute between the two solutions
via the gas phase. This process, as well as
the equilibrtion of the 12/103- couple, re-
quired up t) five days at 250C, but was com-
plete within 24 hrs at 200*C. Upon attaining
equilibrium, as indicated by a constant poten-
tial reading, samples were withdrawn from both
solutions: the iodine concentration, mR, was
determined by spectrophotometry using the
absorption maximum at 460 nm (e = 730.9 M-1
cm-1);18 the total iodine, mL, iodate, mR1,
and iodide, mil, concentrations were measured
by potentiostatic coulometry; and the proton
concentrations, mL2 and mR2, were determined by
titration with standard NaOH solutions.
Twenty-two experiments were carried out over
the temperature range 3.8 to 209.00C.
2.2 Spectrophotometric Measurements
The triiode formation equilibrium, equation
(1), was also studied 'in situ' by spectro-
photometry over the temperature range 15.3 to
44.7*C.18 The wavelengths investigated were
370, 350 and 288 nm. The iodine concentration
was maintained at 5 x 10-5 m, while the iodide
concentration was varied from 10-3 to 0.02 m,
with at least nine measurements within this
range at each temperature.
The acid dissociation constant for iodic
acid, equation (5), was determined from the
difference in solubility of TlIO3 in 0.1 m
NaC104 (10-4 m NC104) and 0.1 m HC104 over the
temperature range 2.2 to 75.0*C. The solutions
were analyzed for thallium and iodate by poten-
tiostatic coulometry, the details of which are
described elsewhere.19 Experiments at 100*C
indicated that some oxidation of TI(I) to
TI(III) by iodate had occurred and hence this
imposed an upper limit on the temperature range
that could be effectively investigated.
3. RESULTS AND DISCUSSION
The operating conditions in the emf experi-
ments imposed a number of constraints on
treating the equilibrium data pertaining to
equation (3). Firstly, the coexistence of
iodine and iodide ions at millimolal levels
within the cell resulted in the formation of
substantial concentrations of triiode ion.
Thus, in order to calculate the molalities of
free iodide and iodine, a knowledge of the
thermodynamics of 13- formation is necessary.
This reaction has been well studied in the
past,20 but only over a narrow temperature
range such that the large uncertainty in the
ACp (reported values vary from -263 to +97 J
K-1 mol-1 at 25*C) makes reliable extrapolation
to higher temperatures impossible. Fortu-
nately, this information is provided directly
from the analyses of the solutions in the cell.
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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/m1/3/: accessed April 21, 2019), University of North Texas Libraries, Digital Library, https://digital.library.unt.edu; crediting UNT Libraries Government Documents Department.