KINETICSOF REACTION OF SODIUM CHLORITE WITH POTASSIUM IODIDE
Gschneidner and Vineyard'* applied second-order elasticity theory to obtain the equation
where p is pressure and B is the bulk modulus. This equation is also valid only for dilute solutions. They suggest the approxim at'ion
where V is the molar volume and cy is the niolar heat) capacity. This equation predicts only positive devia-.
Kinetic Study
OKI the
3027
tions from Vegard's law. However, the magnitude of its effect is comparable to the amount by which the first-order theories overestimate the negative deviations. This suggests that a theory combining both first- and second-order elasticity effects would be a considerable improvement over current theories €or predicting deviations from Vegard's law. Acknowledgments. The authors wish to thank B. Goydish for performing the chemical analyses, H. H. Whitaker for performing the mass spectrographic analyses, and J. G. White for measurements of the temperature dependence of the lattice parameter of Ge and Si.
Reaction of Sodium Chlorite with Potassium Iodide
by Antonio Indelli Chemistry Department, University of Ferrara, Ferrara, Italy
(Received June 1 1964) ~
The rate of the reaction of sodium chlorite with potassiuin iodide mas measured by the polarized platinum electrode method a t 25 O and a t different concentrations of the reactants. The orders with respect to the hydrogen, iodide, and chlorite ions were substantially one in all cases. The reaction shows a negative salt effect in the presence of NaNOs and Ba(S03)2, a stronger retardation in the presence of Th(N03)4, and a remarkable acceleration in the presence of U(32(NOa)zand a trace of FeS04. From the rates a t different temperatures, ranging from 1.5 to SO", an activation energy of 14.4 kcal. can be calculated.
The oxidation of the iodides by the oxyacids of the halogens has been investigated by several authors since 18901; recently, the effects of different salts on the reaction between iodate and iodide has been studied in this department, and earlier a similar research had been carried on the reaction between bromate and iodide.3 No data seeim to exist on the analogous reaction between chlorite and iodide, despite the fact that the chlorites are now widely used as bleaching agents in the textile industry.4 A better knowledge of the oxidizing characteristios of the chlorites seems therefore desirable, and a comparison with the other oxy acids of the halogens could also be interesting.
Experimental Materials. Sodium chlorite was prepared from a technical B.D.H. product, containing about 80 % (1) G. Magnanini, Gam. chim. ital., 20, 390 (1890); (b) S. Dushman, J . Phys. C'hem.. 8,453 (1904); ( c ) A. A. Noyes. 2. phystk. Chem., 19, 599 (1896); E. Abel and F. Stadler, ibid., 122, 49 (1926); (d) K . Weber and R. L'alid, Ber., 72B, 1488 (1939); (e) E Abel, Helv. Chim. Acta, 33, 785 (1950); ( f ) K. J. hlorgan, 11 G. Peard, and C . F. Cullis, J . Chem. SOL, 1865 (1951). (2) A. Indelli, J . P h y s . Chem., 6 5 , 240 (1961) (3) A. Indelli, G. Kolan, Jr., and E. S. Amis, J . Am. Chem. Soc., 8 2 , 3233 (1960). (4) G. P. Vincent, E. G. Fenrich, J. F. Synan, and E. R. Woodward, J . Chem. Educ., 2 2 , 283 (1945).
Volume 68, >Vumber10
October, 2964
ANTONIO INDELLI
3028
NaC102. It was dissolved in hot water, filtered through a sintered glass, and crystallized by cooling. Approximately the first 20y0 was discarded, and a second crop of crystals of NaCIOz.2H20,amounting to about 50% of the original quantity, was filtered, washed with a small amount of ice water, and dehydrated and dried for 2 weeks in a vacuum desiccator containing solid E(OHa5 It was analyzed by adding an excess of potassium iodide and of perchloric acid, and, almost immediately, by neutralizing with sodium hydrogen carbonate; finally, the mixture was titrated with sodium thiosulfate, using starch as an indicator. The salt was found to be 98.9% pure in NaC102. The other chemicals wyere Carlo Erba RP products and were used without further purification. Procedure. The rate of the reaction was measured by a method that was similar to that used to investigate the reaction between iodate and iodidea2 The usual care was taken to add the thiosulfate only when a certain amount of iodine was present, so that an excess of reductant was present only at the end of each time interval. From six to twelve points were taken during each run, and the reaction mas followed for about 5 to 12y0 of its course. The reaction rate, under these conditions, did not vary with the time, and the plots of amount of added thiosulfate against time were substantially linear, as Fig. 1 shows. The initial reaction rate was therefore measured with an accuracy of about 5%, as judged from the results of duplicate runs. Each experimental result reported in the tables or in the figures is the average of a t least two experiments.
Some runs were made to check the influence of the reaction of the chlorite with the thiosulfate. In these runs the additions were made immediately after the appearance of the iodine, so that a certain excess of thiosulfate was always present , except during the brief instant necessary to reveal the presence of the free iodine. The concentration of thiosulfate a t the moment of each addition was known and could be varied; this method had already been used to obtain approximate data on the persulfate-thiosulfate reaction.6 The rate of the reaction of the chlorite with the thiosulfate could be determined in this way with reasonable accuracy.
Results Reaction Orders. The logarithms of the reaction rates, calculated as -d [CIOz-]/dt mole 1.-l sec.-l, are plotted in Fig. 2 against the logarithms of the concentrations of the reactants. The slopes of the three straight lines, corresponding to the orders with respect to the chlorite, the iodide, and the hydrogen ions, are 1.15, 0.91, and 0.97, respectively. The deviations from 1.0 are not large and probably are partly due to the experimental uncertainty. In fact, at the lowest iodide concentrations and at the highest chlorite conlog [KI].
2.0
3.5 I
5.5
3.0 I
I
/
0.3
0.1
5.5
100
500 Time, seo.
1000
Figure 1. Typical rate plots, ml. of 0.1 144 NazSz03 against time, for the reaction between chlorite and iodide a t different temperatures. Total volume = 400 ml.; 104[KI] = 5.0; 1 0 4 [ N a C 1 0 ~=] 1.66; 104[HC104] = 5.0; 106 E D T A = 7.5 moles 1.-'. Temperatures: 1, 15"; 2, 20"; 3, 25'; 4, 30"; 5, 35"; G , 40"; 7 , 45"; 8, 50".
The Journal of Physical Chemistry
Z.0 3.5 log [HC101]or [NaClOz].
3.0
3.5
Figure 2. Logarithms of the rate plotted against the logarithms of the concentrations of the reactants. Upper scale, log [KI] ; lower scale, log [HClOJ and log[NaClOzl ; open circles, HC104; filled circles, KI; crosses, KaClOz.
( 5 ) G. R. Levi. A t t i R. accad. Lincei, [ 5 ] 31, 214 (1922); Gam. chim. ital., 5 2 , 418 (1922). (6) A. Indelli and E. S. Amis, J . Am. Chem. Soc., 82, 332 (1960).
KINETICSO F REACTION O F SODIUM CHLORITE WITH
3029
E’OT’ASSIUM IODIDE
37:
1 8.4
8.3
Table 11: Reaction Rates, v, in the Presence of Different Salts a t 25”“
NaNOs
0 100 250 500 1 7 . 3 14.6 1 3 . 6 1 3 . 0 250 500 50 100 13.4 12.9 14.6 14.2 5 10 20 2 9.8 13.7 12.2 11.7 200 39
109~
Ba(N03)~ 109v Th(NOt)a 1090
cOz(N03)z 109v FeSOd 109v
x NaNOl O B a N O , l2
1000 11.8
40 9.2
100 7.1
0 .O O P 476
Th (N O,I,
8.1
a 104[KI] = 5.0; 1O4[SaC1O2]= 1.66; 104[HC10J lOB[EDTA] = 7.5 moles 1.-’. Added salt concentrations No EDTA present. mole 1. -1.
‘
= =
5.0;
1
0.05
0.15
0.10 I”.
Figure 3. Log
00’
(eq. Z!) against the ionic strength, p .
centrations it is difficult to eliminate the influence of the reaction between the chlorite and the thiosulfate. It can be safely coiicluded, therefore, that the main activated complex contains one CL02-, one I-, and one H + ion, although probably there are also other reaction paths of smaller importance. Activation Energy. Table I reports the average rates measured a t different temperatures, v, and the rates, v,,t, calculated by means of the equation log u i n t == 2.813 -
3.15 __ X lo3 T
The concentrations alf the reactants are those indicated under Fig. 1. From four to eight runs were made a t each temperature. Table I : Reaction Rates, v, - d [ ClOz-] /dit (mole 1. sec. - I ) a t Different Temperatures, and Corresponding Interpolated Values, vint’
Temp., “C. lo%, mole 1.-’ sec. 109vint
15 7.5 7.6
20
25
30
35
40
45
1 2 . 1 1 7 . 3 2 6 . 0 39 1 1 . 7 1 7 . 7 2 6 . 4 39
55 57
82
50
81 120 116
a 104[KI] = 5.0; 104[1NaC10~]= 1.66; 104[HC104] = 5.0; 106[EDTA] = 7.5 moles 1 . 3 .
From the temperature dependence of the reaction rate, an activation einergy of 14.4 kcal. can be calcu-. lated. Salt Effects. Table I1 reports the results obtained
in the presence of different salts. The negative salt effect expected for a reaction between two anions and a cation is observed in the presence of sodium and barium nitrates. The rates obey the equation
where p is the ionic strength, A is the Debye constant (for water a t :!5O, A = 0.5085’), and B is an empirical coefficient, which for NaKOs is 0.5 and for Ba(N03)2 0.8. Th(K03)4 gives a much stronger retardation, which cannot be fitted by eq. 1. This is shown in Fig. 3, where log vo’,defined as
is plotted against p S 8 The points for NaN03 and for Ba(N03)2are fitted by straight lines, whereas those for Th(r\TO3)(are on a curve, whose initial slope is very large and negative. The vertical lines represent the rf 5% experimental error. The uranyl and iron salts have a catalytic action, which clearly has nothing to do with the salt effects. Their action is rather irregular, and the reproducibility is poor. The Reaction between Chlorite and Thiosulfate. The product of the reaction of S203-2with ClO2- is S106-2.9 Persulfatelo and iodate” also oxidize thiosulfate to (7) E. A. Guggenheim and R. H. Stokes, Trans. Faraday Soc., 54, 1646 (1958). (8) E. A. Guggenheim and J. E. Prue, “Physicochemical Calculations,” North Holland Publishing Co., Amsterdam, 1956, p. 466. (9) G. R. Levi and C. Castellani-Bisi, Gam. chim. ital., 87, 336 (1957). (10) C. V. King and 0. F. Steinbach, J . A m . Chem. Soc., 5 2 , 4779 (1930). (11) R. Rieder, J . Phys. Chem., 3 4 , 2111 (1930).
Volume 68, Number 10 October, 1964
ANTONIOINDELLI
3030
tetrathionate. Recently n'a2Sz04has been obtained in good yield by reaction of NaCIOz with HC1 and Na2Sz03.12 Therefore, when both iodide and thiosulfate are oxidized by the chlorite in the same time, the two concurrent reactions can be assumed to be ClOz-
+ 4H+ + 41-
--f
C1-
+ 212 + 2Hs0
and ClOz-
(3) &
+ 4SZO3-*-+ 41- + 2S40~-2very fast
+ 4H+ f 4Sz03-'
0 bo
(3') 2.0
+
C1-
+ 2Hz0 + 2S40e-'
(4)
The rate of reaction 3 has been called u and can be considered constant, under our conditions, during each run. If we assume that reaction 4 is first order with respect to the thiosulfate, its rate can be expressed as IC, [Sz03-2].13 The total rate of disappearance of the thiosulfate will be given by -d[Sz03-']/'dt
=
4 ~= ' 4~
+ 4k1[5203-~]
(5) where v is the sum of the rates of the two concurrent reactions. It can be shownethat 0
u'
-
I
2.5
followed by 212
I
I
I
kl[wo,-'lO)
U IC1 [ S z 0 3 - 2 ] 0
h(1+
0
where [S203-2]o is the concentration of the thiosulfate at the moment when the addition is made. ICl can be calculated by successive approximations, and Table I11 reports a selection of values of ICl, together with the experimental u ' , for different [S203-2]o.In all cases IC, was found substantially constant, although [S203-2]o changes by a factor of six, and this justifies the assumption that the reaction is first order with respect to the thiosulfate. The first-order rate constant IC, varies with the concentrations of perchloric acid, sodium chlorite, and potassium iodide, and Fig. 4 reports plots of log k l against log [HClOi], log [NaCIOz], or log [KI]. The orders with respect to CIOz- and H + are a little less than 0.8 and 0.6, respectively. Yo definite order can be given with respect to the iodide, which, however, has a clear inhibiting action.
Discussion The reactions between the halide and the XOa- ions, where X is a halogen, have the common feature that the order with respect to the hydrogen ion is two,1,2 or perhaps even larger.14 This has been interpreted as due to the intervention of a XOz+ ion as an interrnediate.'5 The fact that in this reaction the order with respect to H + is substantially one suggests that neutral HCIOz acts as an intermediate. The following series The Journal of Physical Chemistry
3.5
LO
2.5
3.0
log (HC1041 or [NaClOz].
Figure 4. Log kl against log [HClOd], log [KaC102],and log [ K I ] for the reaction between Ka&hO, and NaC102. Upper scale, log [ H I ] ; lower scale, log [HClOa] and log [SaCIOz].
Table 111: Reaction Rates, v', of Chlorite in the Presence of Different Initial Concentrations of Thiosulfate, and First-Order Rate Constants, k1, of Reaction 4"
10~[SzOs-z]lo, mole I.-' 1O*v, mole 1. -'
5 9.9
10 15 20 14.7 1 9 . 9 25
25 30 28 5 30
sec.
1O2k1,set.-'
5.6
5.0
5 0
5.1
a Temperature: 25"; reactant concentrations: Table 11; 109v = 17.3 mole 1.-l set.-'.
4.9
4.3
same a8 in
of reactions is consistent with the experimental orders.
+ H + J_ HClOz rapid equilibrium (7) HClOz + I-+ HClO + IO- rate determining (8) H + + HClO + 21- --+ C1- + HzO + l 2fast (9) 2H+ + IO- + I-- + HzO + Iz fast (10) ClOz-
Minor reaction paths probably involve a rapid reaction of the iodide ion with some active intermediate formed by decomposition of the chlorite ion and of the chlorous acid. This is indicated both by the deviations from 1.0 of the different reaction orders and by the fact that sodium chlorite reacts with iodides even at pH 8.l6 ~~
~
(12) A. Indelli and V. Bartocci, unpublished results. (13) The subscript 1 or 2 indicates that kl or kz is, respectively, a
first-order or a second-order constant. (14) 0. E. Myers and J. W. Kennedy, J . Am. Chem. Soc., 7 2 , 897 (1950). (15) K. 5. Morgan, M. G. Peard, and C. F. Cullis, J . Chem. Soc., 1865 (1951).
KIXETICS O F REMTIONO F
Equilibrium 7 has been investigated by different authors, and, although there is no coniplete agreement, there is no doubt that pK for the chlorous acid is of the ~ the conditions of the present exorder of t ~ 0 . lUnder periments, therefore, HClOz is largely dissociated. To make a rough evaluation of the activation paranieters of reaction 8, we have neglected the niinor reac1,ion paths and have assumed that the orders with respect to HCIOz and I- are exactly one; also, we have taken for pK of HCIOzat 20’ the value of 8.97 given by David~0n17ba t zero ionic $strength,and for AH of dissociatilon the value of 3.5 kcal.I7” We have corrected the dissociation constants a t the ionic strength of the reaction mixture by means of a Guntelberg forinulal* and we have calculated the concentrations of HCI02 at each temperature. From the rates reported in Table I we have calculated the bimolecular rate constants k2I3 of reaction 8, defined as
k,
2‘
=
[HCl02]
rI-1
For the rate constant, ICz, a t 25O, the activation energy, EA, the frequency factor, -4, the activation enthalpy, A H * , the activation entropy, AX*, the following values are found: k2 = 5 60 1. mole-l sec.-l; EA = 17.8 lccal.; A = 1013.81. mole-’ set.-'; AH* = 17.2 kcal.; AX* = +2.6 e.u. The values for the frequency factor and for the activation entropy are reasonable for a reaction between a neutral molecule and an ion. l 9 The negative salt effects exerted by NaNO:, and Ba( S O & are in agreement with the series of reactions 7 to 10. In fact, the rate constant, of the rate-determining step 8 will not be appreciably influenced, but the concentration of HC1OZwill be decreased because of the influence of the ionic strength on equilibrium 7. This situation is similar to that of the ammonium cyanateurea conversionlZoand any nonchain mechanism which involves an activated complex containing a ClOz-, a H + , and an I- ion will give a rate which obeys the equation 2,
3031
SODIUM CHLORITE WITH POTASSIUM IODIDE
= Vofi3/fi
Introducing forfi the expression
(12)
eq. 1is obtained. B in eq. 1 is given by
B
=
bclo,-
+ b ~ ’ + b ~ -- b*
(14)
The effects of Th(S03)4, UOz(1\’03)2,and FeS04, clearly, are not consistent with this picture. For Th(N09)J there is probably the formation of ion pairs, or complexes, with the chlorite, of decreased reactivity. Additional informations, from independent sources, on this ionic association would be required for a more thorough discussion, but Fig. 3 suggests that at a thorium ion concentration of 0.01 mole 1.-1, most of the chlorite is in the form of ion pairs. There is in fact a definite leveling of the curve after the initial fall. A possible inechanism of the catalytic action of the ferrous sulfate consists in its oxidation to ferric ion, which is reduced by the iodide ion. A similar mechanism had been proposed for the catalysis of the persulfate-iodide reaction, although later work has shown that a mechanism based on the formation of intermediate ion pairs is more For uranyl nitrate an intermediate oxidation-reduction of the uranium is very unlikelylZ2so that in this case, too, an ion-pair mechanism is suggested. Uranyl nitrate is also a catalyst in the persulfate-thiosulfate reaction and in the bromate-iodide reaction. 336 (16) V. P. Tolstikov, Sb. Statei Po Obshch. Khim., Akad. Nauk SSSR, 2, 1249 (1953); Chem. Abstr., 49, 2921e (1955). (17) (a) J. Bjerrum, G. Schwarzenbach, and L. G. Sillen, “Stability Constants,” Part 11, The Chemical Society, London, 1958, p. 27; (b) G. F. Davidson, J . Chem. Soc., 1649 (1954); (c) K. Tachiki, J . Chem. Sac. Japan, 65, 346 (1944); Chem. Abstr., 41, 33479 (1947); (d) >I. W . Lister, Can. J . Chem., 30, 879 (1952). (18) E. A. Guggenheim and T. D. Schindler, J . Phgs. Chem., 38, 543 (1934). (19) (a) E. A. hloelwyn-Hughes, “The Kinetics of Reactions in Solution,” 2nd Ed., CIarendon Press, Oxford, 1956, p. 71 ; (b) S. Glasstone, K. J. Laidler, and H. Eyring, “The Theory of Rate Processes,” McGraw-Hill Book Co., New York, N. Y . , 1941, p. 433; (c) A. Indelli, Rie. Sei., 28, 1676 (1958). (20) (a) A. A. Frost and R. G. Pearson, “Kinetics and Mechanism,” John Wiley and Eons, Inc., New York, N. Y., 1953, p. 260; (b) I . Weil and J. C. Morris, J . Am. Chem. Soc., 71, 1664 (1949). (21) D. A. House, Chem. Rev., 62, 185 (1962). (22) A. Indelli, Ann. Chim. (Rome), 53, 620 (1963).
Volume 68, Sumber 10 October, 1964
