KINETICSTUDY OF
THE
REACTION OF PERIODATE WITH IODIDE IONS
63 1
A Kinetic Study of the Reaction of Periodate with Iodide Ions
by Antonio Indelli, Francesco Ferranti, and Ferdinand0 Secco Department of Chemistry, University of Ferrara, Ferrara, Italy, and Department of Chemistry, University of Camerino, Camerino, Italy (Received July 10, 1966)
~
~~
The rate of the reaction of periodate with iodide in acidic media was studied. The rate k2[1-1 [IO,-] [H+]. equation for acid concentrations of up to 0.003 M is v = kl[I-] [IOd-] At higher acid concentrations the expression also contains a term in [H+Iz. This is confirmed by comparison with the data reported by Abel and Siebenschein. The reaction exhibits salt effects of low specificity, even in the presence of alkaline earth cations. Lanthanum and thorium ions have a retarding effect. The mechanism of reaction is discussed.
+
Introduction In earlier investigations,’ the rates of oxidation of the iodide ions by various oxidizing agents were studied under various conditions, using a microtitration method in which the polarized platinum electrode was used as the indicator. The same method has now been applied to the determination of the rate of the periodat-iodide reaction. This reaction is known to take place in both acidic and neutral media,2 and this is the basis of the analytical distinction between the periodates and the iodate^.^ The reaction is very fast under normal conditions, and no direct measurements appear to have been carried out in acidic media. Only Abel and Siebenschein4 have estimated the rate froin experiments on periodate-iodate-iodide mixtures by measuring the quantities of periodate and iodate remaining after the iodide has been completely consumed. The results obtained by these authors indicate that the over-all process is the result of two simultaneous reactions, one of which is zero order and the other second order with respect to the hydrogen ion. The rate of the reaction in neutral media was studied by Abel and Fiirth5 using a method based, like ours, on the time of reappearance of the iodine. Their results were corroborated by Peschanscki.6 Experimental Section The potassium periodate was a BDH Analar product. LiNOs was prepared from LizS04and Ba(NO& and recrystallized, while the other chemicals were the same as those used in the earlier work.’ The experimental technique was also, in general, the same. A
solution of NazS203was added to the reaction mixture in slight excess with respect to the iodine already present by means of an Agla microsyringe, and the time of reappearance of the iodine was read on a stopwatch. This procedure was repeated many times during each experiment. A small quantity of EDTA (7.5 X M ) was, as usual, present in all runs in order to avoid any possible catalysis by traces of heavy metal ions. Since the concentrations of the reagents, and in particular the equivalent concentration of the periodate, were much lower than in the earlier work, it was generally impossible to measure the initial rate directly, but first-order graphs had to be used. A reproducibility of about 2% was obtained in every case. Results Orders of the Reaction. The first product of the reduction of periodate is the iodate ion, as5 104-
+ 21- + 2Hf +10%-+ Iz + H20
(1)
At sufficiently low iodide ion and hydrogen ion concentrations, the subsequent reaction ~~
~~~
(1) (a) A. Indelli, G. Nolan, and E. S. Amis, J . Am. Chem. SOC.,82, 3233 (1960); (b) A. Indelli, J . Phys. Chem., 65, 240 (1961); (c) A. Indelli and J. E. Prue, J . Chem. Soc., 107 (1969); (d) A. Indelli, J. Phy8. Chem., 68, 3027 (1964). (2) P . Pascal, “Trait6 de Chimie Minkrale,” To1 I, Masson et C . , Paris, 1931, p 760. (3) E. MUller and S . Friedberger, Chem. Ber., 35, 2655 (1920); L. Szekeres, 2. Anal. Chem., 172, 256 (1960). (4) E. Abel and R . Siebenschein, 2.Physik. Chem., 130, 631 (1927). (5) E. Abel and A. Fdrth, ibid., 107, 313 (1924). (6) D. Peschanscki, J . Chim. Phys., 48, 489 (1951).
Volume 70, hiumber 3 March 1966
632
A. INDELLI, F. FERRANTI, AND F. SECCO
IOa--
+ 51- + 6H+
--f
312
+ 3Hz0
(2)
is so slow that it may be disregarded in practice. Moreover, the elementary iodine formed in the reaction was continuously reduced, so that if the concentration of hydrogen ions is much greater than the concentration of periodate, or if it has a secondary effect on the rate, it may be assumed that the apparent order of the reaction for the disappearence of periodate is the true order with respect to the periodate ion. Figure 1 shows graph3 of log ( a - 2) as a function of time for a number of typical examples, where a is the quantity of thiosulfate corresponding to the initial periodate in accordance with reaction 1, and x is the quantity of thiosulfate added at the time in question. It can be seen that some of these graphs are linear for 90% of the reaction, indicating that the reaction is approximately first order with respect to the periodate. At M, hydrogen ion concentrations above 2.5 X however, the graphs are no longer linear; instead, the points lie on a curve with downward concavity. It is nevertheless still possible to determine the initial rate from the initial slope of the curve, by multiplying this by the initial periodate concentration. Table I shows the values of' the initial velocity for various concentrations of the reactants. The rate is approximately
Table I : Dependence of the Rate of Reaction, v, on the Concentrations of HC104, K I , and KIOd a t 25" (concentration:?,mole 1.-1; v, mole 1.-' sec-I)
~~~[ 1080 104[KI] 10%
10'[KIO4] 10%
104[KI] = 2 . 5 ; 104[HC104] 2 . 5 0.1562 0.312 0.625 1.25 1.53 3.87 8.39 17.44
2.9
2.8 2.7 A -
N
2.6
I
52.5 0
4-
2.4
2.3
2.2 2.1 2.0
1000
10.0 36.1 2.5 36.6
2000
Time,
3000
8ec.
Figure 1. First-order diagram for a number of typical experiments.
at values of the latter between 0.5 and 10 X 10-4 mole 1.-'. In this range the rate is a linear function of the hydrogen ion concentration, the relationship being
v
lO4[KIO4] = 0.625; 104[KI] = 2 . 5 H 0c. 5~ 1o. 0J 1 . 5 2 . 0 2 . 5 3 . 0 3 . 5 5 . 0 10.0 7 . 5 9 8.00 8 . 2 6 8 . 3 9 8 . 3 9 8 . 6 2 8 . 8 3 9 . 7 5 10.27 104[KI04] = 0.625; 104[HC104]= 2 . 5 0.625 1.25 2.5 5.0 1.87 3.77 8.39 18.03
1.0
=
VO
+~H[H+]
(3)
Since the orders with respect to iodide and periodat'e are both practically equal to unity, eq 3 may also be written v
=
kl[I-] [IO,-]
+ kz[I-] [IO*-][ H + ]
(4) The following values of IC, and ICz are found from Figure 3 kl = 4.78 1. mole-' sec-'; kz
=
2.73 X lo3L 2 mole+ sec-'
ICl and kz should, however, be expected to depend on proportional to the concentration of iodide and periodate ions, and Figure 2 shows a graph of the logarithms of the rates against the logarithms of these concentrations, the other concentrations being kept constant. The slopes of the straight lines calculated by the method of least squares are 1.1 and 1.0, respectively, for the periodate and iodide ions. The concentrations vary in each case by a factor of 16. It is also obvious from Table .Ithat the rate depends slightly on the hydrogen ion concentration. Figure 3 shows a graph of the rate as a function of the hydrogen ion concentration The Jozirnal of Physical Chemistru
the ionic strength.' In fact, kl relates to a reaction between two anions, and ICz to a reaction between two anions and a cation. If Guntelberg's formula* is used to calculate the activity coefficients of the reagents and of the activated complex, it can be easily shown that log IC1 = log IC10
+ 2AI1'"(1 + 1'12)
(3
(7) A. A. Frost and R. G. Pearson, "Kinetics and Mechanism," John Wiley and Sons, Inc., New York, N. P.. 1953, p 138. (8) E. GUntelberg, 2. Physik. Chem., 123, 199 (1926); E. A. Guggenheim and T. D. Schindler, J. Phys. Chem., 38, 533 (1934).
KINETIC STUDY OF THE REACTION OF PERIODATE WITH IODIDE IONS
log kz = log k*O - 2A11’7(1
+P)
(6)
where I is the ionic strength and A is the Debye-Hiickel constant, which is equal to 0.5085 for water a t 25”. The ionic strength was not kept constant in the experiments reported in Figure 3, but this should not lead to serious errors, since the concentrations were very low in every case. The values of kIo and kzo can be
633
calculated to a good approximation from the values of kl and kz reported above, using eq 5 and 6 and the mean value of the ionic strength for the experiments of Figure 3. This gives the following values for k? and kzo
kI0 = 4.53 1. mole-’ sec-’; kzo = 2.89 X lo31.2mole-2 sec-l At higher perchloric acid concentrations, the determination of the initial rate of disappearance of periodate becomes rather difficult, since the graphs of the type shown in Figure 1 are no longer linear after a very short distance. This is because the iodate formed in
7.5
i 7.0
4
ea
s
8.5
-
3
I
I
n
2
S.0
6.5
4.0 Log concentration.
3.0
z.5
1
Figure 2. Order of the reaction with respect to the I - and the 1 0 4 - ions: circles, log [KIO4]; triangles, log [KI]. 0 0
50
100
250
150 200 Time, aec.
300
Figure 4. Quantities of iodine liberated as a function of time. The curves are calculated on the bmis of eq 7. 104[HC104]: curve A = 10; curve B = 20; curve C = 30; curve D = 40; curve E = 60 mole I.-’.
reaction 1 reacts further in accordance with reaction 2,‘ so that the thiosulfate is consumed by the iodine formed in both reactions. Since the iodide and hydrogen ion concentrations in any single experiment are practically constant, reactions 1 and 2 may be regarded as two consecutive first-order reaction^,^ and the thiosulfate consumed in both can be calculated as a function of time. The following equation is obtained
[s203*-]= 2[10~-1[4(1 3 ~
1.0
2.0 3.0 10~[H+].
4.0
5.0
Figure 3. Dependence of the rate of reaction on the hydrogen ion concentration.
- e-kIot)
hoc
ho, - h o ,
+
(e-kIOal
- e-kIot)]
(7)
(9) K.J. Lsidler, “Chemical Kinetics,” McGraw-Hill Book Co., Inc., Now.Y&, N. Y., IQM),p 22.
Volume 70,Number 3 March 1966
A. INDELLI, F. FERRANTI,AND F. SECCO
634
where [Sz032-]is the total thiosulfate consumed a t time t, [IO,-] is the initial concentration of periodate, kIO, is the pseudo-first-order constant of reaction 1, and kIoa is the pseudo-first-order constant of reaction 2. kIoI is given by
where kl and k2 are corrected by means of eq 5 and 6. k I O a can be obtained from the published data on the reaction of iodate with iodide,'b allowance being made for the fact that reaction 2 is second order with respect to both the iodide and the hydrogen ion and that the dependence of on the ionic strength is therefore given by log kIo3 = log
JC01oa
- 4A11"/(1
+ 1"')
(')
Table I1 gives the values of k ~ o and , kloa obtained in this way for the various conditions studied, and Figure 4 shows the points obtained in individual ex-
, Table I1 : Pseudo-First Order Constants, k10, and k ~ oUsed in Eq 7 (lO4[I
