JOURNAL OF THE AMERICAN CHEMICAL SOCIETY (Registered in
U.S. P a t e n t Office)
(0Copyright, 1960, b y the American Chemical Society) JULY 20, 1960
VOLUME82
SUMBER 13
PHYSICAL A N D INORGANIC CHEMISTRY [COSTRIBUTION FROM
THE
CHEMISTRY DEPARTMENT OF
THE
VXIVERSITY
OF
ARKANSAS, FAYETTEVILLE, ARKANSAS]
Salt Effects in the Reaction between Bromate and Iodide Ions BY ANTONIOINDELLI, GEORGE XOLAN, JR.,
AND
EDWARD S. AMIS
RECEIVED AUGUST14, 1959 The rate of the reaction of bromate with iodide ions has been measured in the presence of different salts. The reaction is accurately first order with respect to Br03- and I- and second order with respect to H+. The results for the nitrates of potassium, magnesium and lanthanum can be expressed by I -
with individual deviations of the order of 2 to 470. The sodium and barium nitrates and the potassium perchlorate give similar results. The chlorides give higher rate constants, and this can be interpreted as due to a reaction between bromate and chloride ions. The sulfates give much lower rate constants, due to the formation of the HSOa- ion. When the reduction in the hydrogen ion concentration is taken into account, the results for the sulfates are very similar to those for the nitrates a t equal cation concentration. Uranyl nitrate has a specific accelerating effect, which is not due to the increase in H+concentration produced by the hydrolysis. The rate of the reaction catalyzed by uranyl nitrate seems to be of lower order with respect to the H + ion.
In previous papers, the salt effects on the alkaline hydration of condensed phosphates1z2and on the reaction between persulfate and iodide3 have been shown to be very specific and to involve some non-electrostatic factor.2s4 In these cases the activated complexes have a large negative charge, and this can explain to a certain extent the predominance of the specific effects. We thought it would be of interest to study a reaction in which the activated complex has no charge, and we have chosen the reaction between bromate and iodide. This reaction is known to be first order with respect to both Br03- and I-, and second order with respect to H+.5 It can be assumed that the reaction proceeds through some intermediate steps, such asfc BrOlBrOl+
+ 2H' + I-
BrOl+ f H2O rapid equilibrium BrO+ IO- rate determining
= =
it can be proved that the reaction rate is proportional to the fourth power of the activity coefficient of a monovalent ion. The predicted salt effect therefore has the same magnitude as that predicted for the persulfate-iodide reaction, but it is opposite in sign. Only univalent ions, however, are involved. Deviations from the Debye-Hiickel limiting law should therefore start a t concentrations greater than for systems involving multivalent ions.6 Using an approach similiar to that employed in the case of the other reactions indicated above,1 r 3 , 4 , 7 , 8 we have measured the initial rate, to prevent any change in the ionic environment during the reaction. This has been achieved by performing the rate measurements within the first few per cent. of the reaction.
+
Experimental
followed by rapid reactions of the ions BrO+ and IO- with H + and I-. Independent of any mechanism which can account for the observed orders, (1) A. Indelli, A n n . C h i i n (Rome), 48, 332 (1958). (2) A. Indelli, ibid., 4 6 , 717 (1956). (3) A. Indelli a n d J. E. Prue, J . Chem. SOL.,107 (1959). (4) A. Indelli a n d E. S. Amis, THISJOURNAL, 83, 332 (1980). (5) (a) A. A. Noyes, Z . f i h y s i k . Chem., 19, 599 (1896); (b) G. Magnanini, Gam. chim. ilai., 20, 390 (1890); (c) E. Abel, Hclv. Chim. A d a , 8 3 , 785 (1950).
Baker and Adamson quality, C.P., acids and Mallinckrodt reagent quality grade salts were used, except for the Na3P309. This salt was prepared and purified as described in a previous paper.9 The solutions of M ~ ( N O S ) ~ , MgSOr, La( N03)3, MgClz and Lac13 were standardized against a solution of e,thylenediaminotetraacetic acid using
3233
(6) E. A. Guggenheim, Discussioiis F U Y Q ~ U SOL., Y 24, 53 (1957) (7) A. Indelli, A n n . C h i m . ( R o m e ) , 4 6 , 367 (1956). (8) A. Indelli, ibid., 47, 586 (1957). (9) A. Indelli, ibid., 48, 845 (1953).
3234
ANTONIO INDELLI,GEORGENOLAN,JR.,
AND
EDWARD S. AMIS
VOl. 82
eriochrome black T as an indicator. T h e solutions of MgCll TABLE I11 and Lac13 were also standardized against a sc)lutioi: of Ag10+k (L.3 E p u ~ v . - ~ Nos, using fluorescein as an indicator. T h e ?;a2S.0a solu- FOURTHORDER RATE COXSTANTS, tioii was standardized against the stock solutioii of KBrOa, SEC.-') FOR THE REACTION O F Br0,- WITH 1- A S D H + AT which was used, after properdilution, in tile kinetic rutis. The 25" HCl and € I N 0 3si ilution Tvere stmd.trdized agaiiist :i siitliun: rnrhydroxide solution, which liad been, in turn, standardized 031, [KII, Acid, Added Equiv., 10-2 against potassium acid plith,il.ite. All the other solutiolis mole/l. niole/I, mole/l. x lo-' x 10-3 salt 1. - 1 P k were prepared by weighing the s 4 i J salts. The experi- x lo-' 2.08 0 0107 6 5 mental technique was the same as used in the study of t h e persulfate iodide reaction It consisted of atldirig pe3,08 ,0182 9 0 riodically small, awurately measured, arnounts of IY'a&O.i 2 08 , 0 2 5 7 11 0 in slight excess with respect to the iodine already present i'.08 0407 13 0 and of timing the reappcarance of the iodine by the depolaii3.08 0707 19 0 zation of a platinum electrode. The excess of SanSzO:, added each time had to be very small and approximately x 33: 0065 6 9 constant, and therefore the same precautions as indicated 8 33 ,0265 6 4 in a previous paper were used.' In fact the presence of a ,0465 6 2 larger excess, due to a premature addition of the iVa2SzOs 8 33 0265 8 33 6 0 solution, resulted in an increase in the reaction time. This possibly could be due t i , the reactioi: 0465 5 3 8.33 2 09 ,0077 7 0 s.0,- --f S so3' 9 os 0?77 5 4 catalyzed by the hydrogen ions and followed by 2.08 ,047i 4 8 SO3' I2 HzO -+- SOa2H+ 212.08 0477 4 9 T h e reaction rate was calculated I)>- plutting the ainourit of A-a2Sz03adcletl wrsus time. Due to the very small ionic strength p. In many cases the effect of p change in the concentrations, strnight liiics were ohtairied in a reaction between two ions can be expressed with gc~od approximation. T h e fourth order rate cc,nstants, obtained by dividing the slnpes of the straight bv lines by the average cnnce!itr;itioiis during the r w s , are reported in Tables I , I1 and 111.
+
+ +
+
+
TABLE I FOURTH ORDER RATE CONSTANTS, 1'!-*k SEC.-')
( L . a ISQCIV.- 3 \\'IT11 1- AND Ii * AT
FOR THE REACTION O F
in c,->
[I
