WARRENL. REYNOLDS
1830 [CONTRIBUTION FROM
THE SCHOOL OF
Vol. 80
CHEMISTRY, UNIVERSITY O F MINNESOTA]
The Reaction between Potassium Ferrocyanide and Iodine in Aqueous Solutions BY WARRENL. REYNOLDS RECEIVED JULY 22, 1957 The reaction between potassium ferrocyanide and molecular iodine in a n aqueous solution has a n initial rate which is first order with respect t o the concentrations of both reagents. At pH values of 7.1 and 9.2 the rate constant for this reaction IS- = 2FeiCy 31was (1.3 f 0.3) X l o 3 liter mole-' set.-'. The equilibrium constant for t h e reaction 2FeoCy calculated from values of rate constants does not agree with the equilibrium constant determined potentiometrically.
+
The reaction between potassium ferrocyanide and iodine appears to have been first reported by Preusz' in 1839. Since that time a number CJf studies have been made on the rate of the reverce reaction between potassium ferricyanide and iodide ion,2 on the equilibrium attained between ferrocyanide, iodine (or triiodide), ferricyanide and iodide3 and on the rate of the forward react i ~ n . Despite ~ this work the mechanisms involved in these reactions are incompletely understood. The experiments reported below were undertaken to determine the kinetic behavior of the forward reaction in particular. Experimental Reagents.-The chemicals employed in the preparation of stock solutions of potassium ferrocyanide, potassium ferricyanide, iodine and potassium iodide and in the preparation of the various buffer solutions and reaction media were all reagent grade chemicals and were used without further purification. The potassium ferrocyanide stock solution was deaerated and kept under a nitrogen atmosphere when stored for short periods of time; the results obtained with these solutions of ferrocyanide were identical to results obtained with freshly prepared air-free solutions of ferrocyanide. Hypoiodous acid was prepared by precipitating the iodide ion from a slightly basic solution of iodine with the theoretical quantity of silver nitrate. The precipitate first formed was somewhat brown in color but upon making the solution slightly acid the color of the precipitate became pale yellow, the color of silver iodide. A portion of the hypoiodous acid solution thus prepared did not give a n iodine color when acidified strongly showing t h a t iodide ion was virtually absent. When iodide ion was added to a portion of the hypoiodous acid solution and the latter acidified, the brown color of 1 3 - appeared immediately. Conductivity water was used in the preparation of all solutions and was prepared by distillation from alkaline permanganate and then from dilute sulfuric acid. Apparatus.-The reaction vessel was a 100-cc. Pyrex bottle fitted with a perforated metal screw cap and self-sealing gasket. Above the middle of the bottle on one side was fitted a n absorption cell of square Pyrex tubing. The reaction vessel could be removed from the constant temperature bath, tipped to fill the absorption cell and the absorbance of the reaction mixture determined without removing a sample. A Beckman DU spectrophotometer was employed for all the optical density m e a s ~ r e m e n t s . ~ A Beckman model H pH meter was used for the cell potential measurements, with a saturated calomel electrode as reference electrode. (1) J. Preusz, Ann., 29, 323 (1839). (2) (a) F. G. Donnan and R. LeRossignol, J. Chcm. Soc.. 703 (1903); b) G. Just, 2 . p h y s i k . Chem., 63, 513 (1908); ( c ) C. Wagner, i b i d . , 113, 2G1 (1924); A. von Kiss, Rec. trau. chim., 62, 289 (1933); H. B. Friedman and B. E. Anderson, THISJOURNAL, 61, 118 (1939). (3) (a) V. K. LaMer and K. Sandved, i b i d . , 60, 2656 (1928); (b) V. K. LaMer and H. B. Friedman, ibid., 62, 876 (1930); ( c ) R. G . Dickinson and S. F. Ravitz, ibid.. 62, 4770 (1930). (4) R . N. J. Saal, Rec. t v a v . chim.,41, 385 (1928). ( 5 ) T h e anthor wishes t o thank Professor R u f u s I.uniry for his kincl permission to u s e this instritnient.
+
Procedures.-The reaction was followed spectrophotometrically by measuring, at various times, the absorbance of the triiodide ion, I3-, at 350 mp when excess iodide was present or by measuring, at various times, the absorbance at 170 mp when no added iodide ion was present. The molar extinction coefficients of ferrocyanide, ferricyanide, triiodide and iodine were determined a t 350 and 470 mp so t h a t absorption resulting from the presence of each species could be accounted for when necessary. Since the presence of air seemed t o decrease the rate of reaction somewhat all solutions, with the exception of the iodine stock solutions, were deaerated by flushing the solutions with purified nitrogen. Because approximately 1 ml. of iodine stock solution was used in a total volume of 25 ml., the amount of oxygen introduced in this way was very small and did not appear to cause a n y irreproducibility. Electrode potentials of solutions containing 13- and I- and of solutions containing ferro- and ferricyanide were measured with a Beckman model H pH meter employing a saturated calomel electrode as reference electrode and a gold wire electrode as indicator electrode.
Results and Discussion The Reactive Iodine Species.-It is known that when I p is added to an aqueous solution the following equilibria are rapidly established. The values given for the equilibrium constants are for 25'. I t T2
+ H?O
I-
+ 1301 + H+; KI = 5.4 X
10-13
(mole/l.)2
*
(I)
IC2
-+ 1- r'I$-; RP= 770 (niole/l.)-l ' K3 I101 J_ H + + 01-; Kj = 2 X (mole/].)* I2
IC,
IT201
'-Y -
H
+
(2)
(3)
-t1101; K,
> 5.4
X
(mole/l.)6
(4)
is also known that IO- will undergo a self oxidation-reduction reaction to yield 1 0 3 - and I-.9 Therefore, in order to postulate a mechanism for the reaction between iodine and ferrocyanide, it was necessary to know which of the iodine species, 13-, Iz, HOI, HIOI+, 01- and IO3- reacted with ferrocyanide. After standing 24 hr. a t room tcmperature a solution containing 1.0 LIT NaOH, 0.01 11s IO3- and 0.005 LU ferrocyanide (FeoCy) remained clear and colorless indicating that no oxidation of FeoCy had occurred. A solution containing 0.001 J f FeoCy and 0.0005 M added iodine in 0.5 M NaOH remained clear and colorless for several hours but became very slightly yellow overnight a t room temperature showing that a very small amount of FeoCy was oxidized by IO- before the latter was converted to 1 0 3 - and I-. It (6) T. 1,. Allen and R. >.I.Keefer, Tnrs JOURNAL, 1 7 , 2957 (1955). (7) >.I.Davies and E. Gwynne, i b i d . , 74, 2748 (1952). (8) RI. L. Josien and G. Sourissean, Bull. SOC. chim. France, 228 (1050). (01 C
!I. T,i ~ v r lC . F. White. THTSJOURNAL, 66, 335 (1043).
April 20, 1958
POTASSIUM FERROCYANIDE AND IODINE IN AQUEOUS SOLUTIONS
is readily calculated from equations 1 and 3 that in 0.5 ill NaOH the hydrolysis of 1 2 to IO- and Iis virtually complete and hence the added iodine was present chiefly as IO- and I-. Ferricyanide, FeiCy, was stable for a t least 24 hr. in the presence of IO- and I- in 0.5 M NaOH. Hence i t may be concluded that the oxidation of FeoCy to FeiCy by IO- is a slow reaction. From equations 3 and 4 i t may be calculated that when IO- solution (containing no I-) is added to an acetate-acetic acid buffer of p H 4.78, the IOion is converted mainly to HOI with negligible amounts of IO- and H2OI+ present. A solution containing 0.001 M FeoCy and 0.0005 M HOI in the above acetate-acetic acid buffer showed no apparent oxidation of FeoCy to FeiCy overnight a t room temperature. However, when the IO- solution is added to a sulfate-bisulfate buffer of fiH 1.45, appreciable quantities of H20I+ should be formed. A solution containing 0.001 M FeoCy and 0.0005 M total hypoiodite in the above sulfate-bisulfate buffer showed an autocatalytic oxidation of FeoCy t o FeiCy. This observation could be explained as the result of a slow reaction between H201+ and FeoCy to form I- and FeiCy followed by a rapid reaction between H201+ (or HOI) and I- to form 1 2 which in turn reacts rapidly with FeoCy to produce more I- and FeiCy. I n most of the experiments on the rate of reaction between FeoCy and iodine, an excess of I- was used so that the iodine was present largely as 13-. It was observed that increasing the excess of Idecreased the rate of reaction approximately inversely as the I - concentration. Thus it appeared that Is- did not react directly with FeoCy. On the other hand, the reaction between 1 2 and FeoCy in the absence of added I- was too rapid to follow a t most of the fiH values employed even a t very low concentrations of the reactants. From the experiments described above i t may be concluded that in aqueous iodine solutions the hydrated iodine molecule is the main oxidizing agent for FeoCy. Products of the Reaction.-The products formed in the reaction between iodine and ferrocyanide are difficult t o define. The equilibrium which is said to obtain3 in the system ferrocyanide, ferricyanide, iodine and iodide is not a stable equilibrium since, in acid solution, ferro- and ferricyanide hydrolyze to form HCN and slightly soluble Prussian Blue whereas in neutral and alkaline solutions iodine hydrolyzes to give hypoiodite ion which in turn undergoes further reaction to form the iodate ion. The exchange of free CN- with bound CN- in ferrocyanide ions in aqueous solutions has been shown to be very slow.1o Thus the rapid disappearance of 1 2 (or 13-) in the presence of FeoCy could not have resulted from the dissociation of ferrocyanide ion to a pentacyanoferrate (11) complex and free CN- followed by reaction of 1 2 with either of the dissociation products. Hypoiodites or higher oxidation states of iodine were not formed as products during the first several hours of a reaction because acidification of a reac(10) A. W. Adamson, J. P. Welker and M. Volpe, TEISJOURNAL, 72, 4030 (1950).
1S31
tion mixture containing a large excess of I- did not produce the characteristic 13- color a t the end of this time. It was assumed, therefore, that I 2 was reduced to I- and that FeoCy was oxidized to FeiCly and that the over-all reaction could be written as
+
2FeoCy
or as
+
2FeoCy
= 2FeiCy
12
+ 21-
= 2FeiCy
18-
+ 31-
Mechanism of Reaction.-Wagner2 postulated reactions 6, 7 and 5 of the mechanism written below as a result of his study of the reaction between FeiCy and I - in the presence of thiosulfate. I n view of the facts that the system FeoCy, FeiCy, I? and I- has been reported3 as coming to equilibrium and that only the reaction between FeoCy and molecular iodine of the possible iodine containing species need be considered, equation 5 was added to the mechanism of Wagner. Assuming a steadyki FeoCy f I Z d FeiCy f 1 2 kz 12- ----f FeoCy Iz FeiCy ks FeoCy 1 2 - --+ FeiCy 21k4 FeiCy 21- --+ FeoCy 12-
+ +
(5)
+ + +
+
(6)
(7) (8)
state concentration for the intermediate 12- and assuming that the equilibrium expressed in (2) was always maintained, one may readily derive the following expression for the rate of disappearance of FeoCy. When I - was used in large excess [I-] was -d[FeoCyl dt
-
2 KzD-I
constant throughout a given rate experiment and the reaction was followed by means of the variation of the triiodide absorbance with time. The absorbances of FeoCy, FeiCy and 1 2 a t 350 mp were small and so the observed absorbance could be corrected readily to give the absorbance E of 13-. Thus, when an excess of I - was employed, one may write ---d[FeoCy] - -2d[Ia-l dt
dt
el
x 37 -dE
=-2
el
xs
(10)
where e is the molar extinction coefficient of Io- a t 350 mp, 1 is the length of the cell used for the absorbance measurements and S is the slope of the absorbance-time plot a t time t . Equation 9 was tested by utilizing equation 10 and rearranging to give kix
- by
- cz =
R2[I-1 el
(11)
where
The quantities x , y and z were known because [13-] was determined from the measured absorbance a t 350 mp, [FeoCy], and [FeiCy] were calculated from the stoichiometry of the reaction, the
1832
WARRENL. REYNOLDS
amount of 13- consumed, and the initial concentrations of FeoCy and FeiCy, and because S a t time t was measured with a mechanical slope-finding device. The data obtained a t various times throughout the reaction were divided into three sets. The method of averages was applied to the data in each set to give the three equations which were solved simultaneously for k l , b and c. Typical experimentally determined absorbancetime curves are shown in Fig. 1. The data presented in Tables I and I1 are typical of the data obtained from the absorbance-time curves.
VOl. 80
TABLE I1 [FeOCy], = 6.24 X If; [I3-la = 3.15 X M; [I-] = 0.0970 M ; [ F e i C y ] ~= 0; e = 2.39 X lo' AT-' C I I I . - ~ ; 1 = 0.78~111.; 24.0 & 0.1'; pH 7.13; /A = 0.20 Time --s x 104, (miri.)
R
set.-'
n X 101
Y
e
3.00 6.00 9.00 12.0 15.0 18.0
11.500 ,446 A00 ,378 .356 ,336
3.59 2.46 1.85 1.43 1.14 0.942
3.95 4.60 5.11 5.70 6.30 6.81
4.65 X 1.99 X 4 . 6 1 X lo-* 8.80 X 0.144 0.214
0.177 ,322 ,444 ,560 . GOO ,752
initial iron concentration, [FeoCy 10 f [FeiCy IO. This result can be obtained from equation 9 by multiplication by -1 to convert the equation to one for the rate of disappearance of FeiCy and by setting [I3--] = 0. Thus
\
2k4 [I-1 * [FeiCyIz if [FeiCyl,
1 [FeoCy] +
[
k3
= kl
Since FeoCy and FeiCy are converted into one another their total concentration remains constant. Wagner obtained a value of 4 for & / & a t l5O.2
From b a value of kd can be obtained because [I-], K2 and k2/k3 are known. Thus kq was cal0 0 culated to be 1.27 X liter2 mole-2 sec.-l. 0 50 100 Vnfortunately, the value of k4 obtained here canTime in minutes not be compared directly with other values which Fig. 1.-Absorbance-time curves for ferrocq-aiiide-iodiiie have been obtainedZbecause the other values have reactions at pH 7.13 and ,u = 0 20. Curve 1, [FeoCy], = been derived a t temperatures other than 24' and M , [I-]o= 0.0938 .If. activation energies have not been determined and 12 5 X 1O-'M, [I,-], = 3.55 X Curve 2, [FeoCy], = 6.24 X 10-6 A l , [I3-I0= 3 15 X 10--6 because the ionic strengths employed differ considerably. However, the order of magnitude of kd X , [I-] = 0.0970 M . compares favorably with 1.22 X liter2 mole-2 When FeoCy was used in excess (Table I), the set.-', the value obtained by Wagner2a t 15' but a t average values of k l , b and c were found to be 8.40X much higher ionic strengths than used here. Since lo4 liter mole-' min.-' (1.40 X l o 3 liter mole-' increase in ionic strength increases the rate of remin.-' and 7.10 X lop3mole: sec.-l), 8.82 X action between similarly charged ions, the effect of liter, respectively. The value of c when combined decrease of temperature on k4 is a t least partly with the values of the other quantities in c gave a conipensated for by the increased ionic strength. value o f 1.8 for the ratio k 2 ' k 3 . This value of k z / k l When approximately equivalent quantities of FeoCy and iodine were used (Table 11) the values TABLE I /FeOCy]a = 12.5 X 10-5 M; [I3-I0 = 3.55 x 10-5 31, calculated for k l , k2,'kJand k d were 1.20 X lo1 liter mole-' set.-', 1.1, and 6.1 X liter2 mole-Z [I-] = 0.0938 M ; [FeiCy], = 0; e = 2.39 X 10' J1-I C I I I . - ~ ; I = 0.78 cm.; 24.0 i 0.1'; p H 7.13; p = 0.20. set.-', respectively. However, the data of Table Time - S , Y 102, I1 cover only the first 30% of the reaction. ThereE min.-I x ioa Y x 103 z fmin.) after the derived equation 11 does not fit the data : on 0.214 0,458 4.33 5.82 0.109 well and, although essentially the same values of kl G.00 2 41 ,340 ,366 0.450 7.60 ,436 and k , 'k? are obtained throughout the reaction, the 1.01 1.03 9 , on .308 8.93 513 value of k4 became negative. Similar observations 1.72 12.0 ,266 1.27 9.32 were made when FeiCy was added initially. Ad15.0 .232 ,582 1.03 2.60 9.58 18.0 ,205 0.800 10.5 ,645 dition of FeiCy retarded the reaction. Use of 3.95 equation 11 as outlined above gave values of K 1 21 0 ,184 ,648 11.4 ,889 5.42 and k 2 / k i which were in agreement with values re,737 24.0 ,167 ,552 11.7 7.07 ported above, but a negative value was again ob1 3 . 1 844 ,370 ,129 12.7 33.0 tained for k 4 . Other reasons given below lead one is in agreement with the work of Donnan and Le to conclude that the postulated mechanism is inRossignol and of Wagner2 who studied the reaction adequate, especially equation S. Initial Rates of Reaction in Presence of Excess between FeiCy and I- a t constant I- concentration by reducing I2 with thiosulfate as rapidly as it Iodide.--When FeiCy is absent initially the initial was formed and concluded that the observed rate rate of reaction should be controlled b y reactions constant was inversely proportional to the total 5 2nd 7 a n d thc rate of disappearance of FeoCy
April 20, 1958
POTASSIUM FERROCYANIDE AND IODINE IN AQUEOUS SOLUTIONS
given by
1S33
Specific effects by K + and by Na+ have been reported on the rate of the FeiCy I- reaction,2 but these have not been systematically investigated. One experiment with 0.100 M NaI in place of 0.100 M K I showed a tendency for k , to decrease but since the effect was hardly outside the experimental error, specific effects of K f and NIL+were not investigated further. Attempts to study the reaction in the presence of excess I - in acetate-acetic acid and in sulfatebisulfate buffers were unsuccessful. At these pH values solutions became turbid shortly after addition of the ferrocyanide to the buffer containing K I even when 0.01 M K I was used. The turbidity appeared to be a result of the presence of I- since i t did not appear when salts such as KNOs or KCl were used in place of KI. Reaction in Absence of Added Iodide.-One rate measurement was carried out in sulfate-bisulfate buffer of pH 1.45 in the absence of added iodide. The reaction was followed by the change in absorbance a t 470 mp in 10 cm. cells. The amount of triiodide a t any time was too small to measure accurately a t 350 mu. At 470 mp Is-, 1 2 and FeiCy absorb and the total absorbance E is given by
+
when I-- is present in excess. blhen integrated, equation 12 becomes 1 = -1 2kit _ E: &+ell(a[I-]
(l3R)
or
depending on whether equivalent or non-equivalent concentrations of ferrocyanide and iodine were used. In (13b), A = el( [FeoCy]~ - 2[13-lo). The subscript zero indicates the value of the quantity a t zero time. Equation 13 was applied for the first 20-25% of the reaction a t different concentrations of Is-, Iand FeoCy and a t two pH values. The results are given in Table 111. The average value of k , was found to be 1.3 X lo3 liter mole-' set:-' with a standard deviation of hO.3 X lo3. The error was large because of inaccuracies in determining the initial part of the absorbance-time curve. This was especially true a t the lower iodide concentrations because the reaction was so rapid there. However, kl does not show any tendency to drift with decreasing iodide concentration. The rate constant k , did not show any dependency on ionic strength p , when p was decreased from 0.2 to 0.11. However, unless k1 was very sensitive to changes in p the effect would be masked by the error in k1. Therefore in the remaining experiments where iodide concentration was varied p was not maintained constant.
E=
eFe!Cy
l[FeiCyl
+ era-Z[T3-1 + ex, Z[I21
+
dt
a
dt
dt \
Expressions for the rate terms inside the parentheses may be found from the postulated reaction mechanism. The material balance and stoichiometry equation may be employed along with
+
[FeoCy] [FeiCy] = [FeoCyJo TABLE I11 1 [FeiCy] [I210 - [I4 THERATECONSTANT FOR THE REACTION BETWEEN IODINE [Izl0- [I21 - [I3-] = 2 A N D FERROCYANIDE BY THE INITIAL RATEMETHOD [I-] [I-] -I- [Ig-] = [FeiCy] 24.0 f 0.1' equations 2 and 14 to give the concentration 9.2
0.20 6.24 3.23 0.097 1.1 .20 6.24 3.12 ,097 1.2 .20 12.5 3.18 .097 1.1 .20 6.24 6.24 ,097 1.5 .20 6.24 6.24 ,097 1.2 7.13 .20 6.24 3.15 .097 1.0 1.8 .20 6.00 3.13 .097 .20 12.5 3.19 .097 1.3 .20 12.5 3.55 ,097 1.2 .20a 12.8 3.14 .097 0.92 .20 6.24 6.38 ,097 0.84 1.0 .20 6.24 6.40 ,097 .20b 6.08 3.11 ,012 1.0 .20b 6.08 3.11 .012 1.1 .I1 6.08 3.11 ,012 1.3 .13 5.20 3.22 ,025 1.8 .13 6.24 3.26 ,025 1.6 .13 6.24 3.26 ,025 1.3 1.5 .13 5.02 3.13 ,031 .15 6.08 3.11 ,050 1.4 0.100 M NaI in place of 0.100 M KI. NaC104 was added to bring ionic strength up to this value.
(14)
where the e's are the molar extinction coefficients of the indicated species and E is the measured absorbance corrected for the absorbance of the solvent. Thus the equation for the rate of change of absorbance may be written dt P I - - - -d[I -1 + e12 d41 (15)
(16a)
(16b) (16~)
of each species in terms of E . Thus an equation giving dE/dt in terms of E can be derived. The numerical values of eFeiCy, er,- and el2 a t 470 m p were determined to be 18, 623 and 530 1. mole--' cm.-', re spectively. For [I2I0= 1.92 X lop5 111 the equa tion was obtained
-dE =
dt
5.00 ks(E - 0.007) k2(0.102
+
{ kZkf(E:)
where f ( ~ = ) -5.16 3.09 x 10-8113
- E)
x
+ k l k d E )i
1 0 - * ~ 5+ 8.9 x 10-8124 + 4.46 x - 2.92 x 1 0 - 9 ~ 2
+ 7.30 X 10-'2 g ( E ) = 6.35 X 10-6E4- 8.50 X 10-5E3 + 1.76 >< 10-6E2- 1.21 X 10-'E + 2.79 X 10-11 10-'OE
and
The function f ( E ) was approximately equal to zero throughout the reaction. Hence the equation for (dE/dt) may be written
!!E= dt
5,00kik3g(E ) k j ( E - 0.007) kz(0.102 - E )
+
WARRENL. REYNOLDS
1534
Two measured values of the slope will thus give k1 and k2/k3. Results are presented in Table IV. The value of k1 remained constant over most of the TABLE IV REACTIONBETWEEN FERROCYANIDE AND IODINE IN ABSENCE OF ADDEDIODIDE [ F e o C y ] ~= 3.85 X M ; [I210 = 1.92 X M; 24.0 f 0.1”; pH 1.48; fi = 0.10. Time
(sec.)
E
0 10 20 30 60 (10 min.)
0.118 .I10 ,104 .099 .086 0.66
g(E) X 108
(dE/df) X lo:, sec.-
h
k,/ks
...
..
..
-8.5 -7.1 -6.0 -3.8
-7.0 -5.8 -5.0 -3.6
1.0 0.7 4.8
156 157 156
x 10-3 x 10-4 x 10-4 .........
1.00 5.00 2.50
.........
eH+ + F ~ ( C N ) O - K.~ ;
TABLE V E.M.F.’s I N HzPOd--HP04-2 BUFFER pH 7.13; 25’; fi = 0.10
OF
[PeOCyl, X 103, IPeiCyl, X 103,
[Is-], Ad
.........
reaction, and the ratio k2/k3 was of the order of unity as found when I- was used in excess. The value of k1 was much less at pH 1.43 than a t 7.13 and 9.2, however. Had this not been so, the direct reaction between FeoCy and I2 could not have been measured in this way a t PH 1.48. At pH 7.13 and 9.2 the ionization HFe(CN)s-a
MEASUREMENT
.
sec.-l
so
As a result of the disagreement between the kinetic and predicted values of K , a t p H 7.13, it was decided to measure the e.m.f.’s of the 13-/1- and FeoCy/FeiCy couples in the reaction medium employed, namely, H2P04--HP04-2 buffer of p =
mole -1
...
...
VOl.
[ I - ] , ‘If
M
M
0.100 .o500
.. ..
..
.om , , ,
..
..,
....
..
..
..
1.00 1.00 0.50
1.00 0.50 1.00
E , v. DS. S.C.E.
E 01
0.300 .318 ,334 ,140 .120 ,158
0.300 ,300 ,298 ,140 ,138 .140
0.100. The results are given in Table V. e.m.f. Eo‘ is defined by the equations
The
for the iodine couple and by
=
5.8 X 10-6 mole/litcrli
where K , is the acid dissociation constant in terms of activities,was complete and k1 was independent of pH. At pH 1.48 the ionization was riot complete. If the assumption that only F ~ ( C N ) B -reacts ~ with I2 was made, the correct [H+] dependence was not obtained. Apparently species other than Fe(CN)6-4 react with 1 2 . The pH effect was not studied further. The Equilibrium Constant.-From equation 9 it is seen that the rate of reaction is zero whenever k&~[FeoCp] 2[13-] = k2k4K2[FeiCy12[1-1
The concentration equilibrium constant is therefore given by where K , is the activity equilibrium constant and Kf is the activity coefficient part of K,. LaMer and Sandved3reported values of K , in the range 0.3 X
