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a range of acid concentrations from 3 to 5.5 N . is a product of one of the reactions in the dissolution The effect of bar length and acid mass velocity were process acts as an autocatalyst for the reaction. For a number of runs at temperatures of 30, 40 studied. In general, an increase in temperature or an increase in acid concentration increased the rate and 50" and a t acid normalities of 4.3 and 5.5, the of dissolution. At 30°, the rate of dissolution was effects of variations in mass velocity and bar length influenced by the length of the copper specimen. were shown to be negligible. It was then possible This was particularly true a t the lower acid mass to calculate energies of activation which were a t velocities. Contrariwise, a t 50" the rate of dis- 4.3 N acid 9.73 f 0.30 kcal./gram atom of copper solution was affected to only a minor degree with and 7.83 i 0.14 at 5.5 N acid. In general the changes in bar length. This was true for all the energy of activation was a function of concentration acid mass velocities studied. and temperature. Examination of the copper bars before and after Acknowledgments.-James H. Weber wishes to tests showed that, in the cases where bar length was acknowledge the financial aid in the form of reimportant, more copper was removed from the search grants for this project from the Research upper end of the bar than the lower end. This can Council of the University of Nebraska and the Kabe explained by the fact that the nitrous acid which tional Science Foundation.
SALT EFFECTS I N THE REACTIONS IZETWEES IODATE AND IODIDE BY ANTOXIOINDELLI Chemistry Department of the University o j Ferrura, Ferrara, Italy Received June 19, 1960
The effect of twelve different salts on the rate of the reaction betwecn iodate and iodide has been investigated at 25' in a range of ionic strength from 0.00096 to 0.04. Univalent and divalent cations and anions show no specific effects and an extended form of the Bronsted-Debye equation is obeyed through the whole range with an accuracy of about 2.5%. Lanthanum ion gives a retardation slightly greater and thoriumfion far greater than uni- or divalent cations at the same ionic strength. The causes of the differences in the salt effects between this reaction and the reaction of bromate with iodide are discussed in terms of the charge of the activated complex. h mixture of thorium nitrate with potassium iodate has an ultraviolet absorption spectrum shifted toward shorter wave lengths with respect t o the spectrum calculated for the isolated salts. This could be correlated with the negative salt cffect. The results for Na3P309 suggest that the value reported in the literature for the dissociation constant of HP30g2- is perhaps too low.
the salt effects in a number In previous of reactions have been invedgated, and discrepancies from t,he Brfinsted-Debye equation have been consistently found even at ionic strengths as low as 0.003.3 As a rule the salt effects were found to be more dependent upon the concentration of some particular ion than upon the ionic strength, in agreement with the findings of Olson and Simonson,' specific effects being more important in reactions between ions of the same sign. The charge of t.he act,ivated complex seems, therefore, to be r e l e ~ a n t . ~I n the present paper the reaction bet,ween iodate and iodide has been chosen because the main activated complex has a charge of - l,8,g and because it involves only univalent ions, although the expected salt effect is as large as that expected for a bimolecular reaction between a uni- and a divalent ion. It has also the additional advantage t,hat, because of its very high velocity, it can be studied at very low ionic strengt.hs. This reaction had been studied previously mainly in buffer solut,ions, and little att'ention was given to the salt effect~.~ (1) A. Indelli. Ann. Chim. (Rome), 46, 367 (1956). (2) A. Indelli, ibid.. 47, 586 (1957). (3) A. Indelli and J. E. Prue, J. Chem.Soc., 107 (193(1). (4) A. Indelli and E. S. Amis, J . A m . Chem. Soc., 82, 332 (1900). ( 8 ) A. Indelli, 0 . Nolan and E. S. Amis, ibid., 82, 3233 (l!M9). (0) A. Indelli, C:. Nolan and E. S. Aniis, ibid., 82, 3237 (lncjo). ( 7 ) A. R. Olson and T. R. Simonson. J . Chcm. P h y s . , 17, llG7 (1949). ( 8 ) S. Dnshman, J . P h y s . Chem..8 , 453 (1904). (9) E. 4bel and F. Stadler, Z. p h y s i k . Chem., 122, 49 (1920).
Experimental Materials.-Most of the salts used were "Carlo Erba" R P products, and were recrystallized from conductivity water. Sodium trimetaphosphate was prepared and purified as described elsewhere.'O Potassium iodate was used as a primary standard to standardize the Na2S203and the KI solutions. Disodium ethylenediaminetetraacetate was used as a primary standard to standardize the Ca(NO&, Ba(NO&, La(X03)~and Th(N03)4 solutions. The conductivity water was prepared by passing ordinary distilled water through an ion-exchange column, and had a conductivity less than 10-6 ohm-lcm.-'. Kinetic Measurements.-The experimental procedure was similar to that used to study the bromate iodide reaction: and consisted in adding small amounts of NanSpOa solution to the reacting mixture, and detecting the reappearance of the iodine by the depolarization of a platinum electrode. However, the reaction between iodate and thiosulfate is very fast," and therefore the greatest part of the thiosulfate was added when a substantial amount of iodine was present, and only a very small amount was present in excess before measuring the time. Care was taken t o keep the amount of iodine present almost constant throughout the run, as measured by the potential of the electrode, except in the end of each addition when all the iodine was consumed. The volume of the reacting mixture was 400 ~ m . and ~ , five points were taken in each run. Due to the very low concentration of the reactants and to the commratively substantial amounts of h'a&O:, required in each a,ddition, it was necessary to let the reaction proceed for 11% in most cases, so that the measured r n k could not h r ronsidered the initial rate. It was found, howciw, t,hat, when the concentrations of iodate and nitric acid were stoichiometrical, a second-order plot of the dimensionless quantit >a / ( ~- 2 ) against time gave a nearly straight line. This (10) A . Tndelli, Ann. Chim. (Rome), 43, 84R (19.53). (11) It. Itieder. J . P h w . Chem., 34, 2111 (1930).
Feb., 1961
SALT
EFFECTS IN REACTIONS BETWEEN
should be coincidental, as the order with respect to the hydrogen ion is about 2,8~9and on the other hand a continuous enrichment in K I takes place. The product of the slope of the second-order plot times the initial concentration "a" ha8 been .t,aken equal to the initial rate of the reaction, and the rate Constants have been calculated by dividing the rates by the reactant concentrations raised to the proper orders. Duplicate runs have shown an agreement within about, 3'370. A few runs were made to measure the rate of the reaction between iodate and thiosulfate, which, according to Rieder" is fifth order too, i.e., second order with respect to the hydrogen and the thiosulfate, and first order with respect to the iodate ion. A solution of Na2S20awas mixed w.th a solution of BIOS and " 0 8 , the concentrations of the three reactants being in stoichiometrical ratios, calculated on the basis of NazS406as a final product.1lJa At suit,able intervals samples were withdrawn and titrated with a solution of 1 2 3 x lO-4M, using the polarized electrode as an indicator. When the cooncentration of the reactants was 2.5 X equiv. l.-l, a rate constant of 3.4 X lo1*1.4 equiv.-4 sec.-1 was found. Some runs were also made at a concentration of 5.0 X equiv. 1.+. The half-life was then too short (about 35 seconds). However, values ranging from 2.5 to 4.5 x 10'2 1.4 e q ~ i v . -sec.-I ~ were consistcntly found. All the kinetic runs were made in a thermostat a t 25 i 0.02O.
Spectrophotometric Measurements.-The ultraviolet spectra were taken with a Beckman Spectrophotometer Model DU, equipped with a photomultiplier, the cell length being 10 cm.
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I O D A T E AND I O D I D E
2.0 1.5 1.0 -
0 2.40 2.50 2.60 2.70 2.80 2.90 Fig. 1.-Ultraviolet absorption spectrum of KIO, 0.0003 M , Th(N03), 0.0003 M and their mixture.
since the deviations from the over-all fifth order are comparatively small. TABLE I FIFTH-ORDER RATE CONSTANTS, lo-' k (L. EQUIV.-' S E C . - ~ ) FOR THE REACTION OF IOs- WITH I- A m H + AT 25", AT VARIOUS REACTANT CONCENTRATIONS K[ x io', H N O ~x 104. K I O ~x 104,
N mole 1. - 1 mole 1. -1 mole 1.-1 10-0 k Results 0.833 6.4 1 3.75 5.00 The reaction between iodate and iodide is re.833 6.0 2 1.875 5.00 ported9 to be first order with respect to the iodate, 10.00 ,833 6.8 3 1.875 and second order n?th respect to both iodide and .833 5.1 4 7.50 2.50 hydrogen ions, so that the total order should be 5.6 5.00 .417 5 3.75 five. The order with respect to t.he iodide at very low concentrations of the latt,er is reported to beThe rate constants in the presence of different come l , I 3 as it is for the reaction between bromate salts are reported in Table 11. Apparently the and iodide.6 Some mechanisms have been suggested to account for these order^.'^-'^ More re- differences among salts of the same valence type cently, on the basis of measurements on the rate of are not significant, the only exception being potasthe exchange reaction between iodate and iodine, sium chloride, where probably there is a side rethe order with respect to the hydrogen ion was action of iodate with chloride ion.6 Sulfates greatly found to bo 3, and the reaction rate was found to be decrease the reaction rate, due to the formation of given by the sum of two terms, one of which is sec- the HS04- ion, and a similar effect should be exond order wit,h rcspect to the iodate i011.l~ A third- pected in the presence of Na3P309. Table I11 reorder dependencc upon the hydrogen ion concentra- ports the rate constants corrected for the formation had been found earlier for the iodate-bromide tion of HS04- and HP309,2-Sassuming that the reaction. l8 Throughout t'his paper the rate con- order with respect to the hydrogen ion is 2. It can stants have been calculated on the assumption be seen that the corrected rate constants for the that the orders with respect to the hydrogen, the sulfates are very similar to those for calcium and iodide and the iodate ions are, respectively, 2, 2 barium nitrates, whereas those for Na8P3O9,are and 1. The results for the runs with no added salt higher. are reported in 'Table I, and it can be seen that TABLEI1 all the orders appear t,o be slightly higher than those FIFTH-ORDER RATE CONSTANTS, 10-9 IC ( L . ~ EQUIV.-' referred to above. The order wit'h respect to the OF IO3- WITH I- AND H + AT 25' nitric acid appeass t,o be higher than that relative SEC.-') FOR THE REACTION [KI] = 3.75 X [HNOS] to the potJassiu:m iodide. The causes of these [KIOa] = 0.833 X 0.000958 mole 1.-l, k = 6.4 1.4 deviations have not been further investigated. = 5.00 X lo-*, /I =eauiv. -4 sec. -1. The reaction rate law seems to be given by two Equiv. Added 1. -1 0.002 -0.0044.014 . 0 2 t,erms of different total orders, but the fifth order salt k a k a k a k a one must b e much more important than the other, NaNOa 5.7 0.003 5 . 4 0.005 4 . 7 1 0.011 4 . 2 1 0.021 -t
(12) E. Carriere and L. Faysse, Compt. Rend. Acad. Sci.,ZOl, 1036 (1935). (13) E. Abel and K. Hilierdings, 2. p h ~ s i k Chem.. . 136, 188 (1928). (14) W. Bra,y. J . Am. Chem. SOC.,62, 3380 (1930). (15) E. Abel, H e l o . (:him. Acto, 33, 785 (1950). (16) I
