FREDERICK R. DUKEAND PAULR. QUINNEY
3900
iodide ions, reaction D has been investigated with different concentrations of free iodine and constant concentrations of potassium iodide and iodoacetic acid. In D1 the quantity of free iodine was equivalent to '/3 of the quantity of iodide present, in D2 this ratio was '/zand in D3 i t was 2/3. The concentration of sulfuric acid has been 0.005 molar in all experiments.
tion
tion
A
Concn. K I (moles I. -1) Concn. Iz (moles I. -1) 0 Concn. CHzICOOH (moles I. -1) ILxp. react. const. (moles 1 . - ' sec. -1) 0.0087 Calc. concn. IX 0.0087 (moles total concn. I ,0087 I.-' sec.-l) (calcd. with equil. const. = 0.0014) Calc. concn. I X 0.0087 (moles total concn. I],-I sec,-I) (calcd, assuming 'OoS7 no free 1 2 in 901 )
i
D1 0.0046 .00147 ,0093
Reaction D2
Reaction D3
0.0046
0.0046 ,0030 .0093
,0023 ,0093
y
+ + = total concn. I + 1;-
+
+
a t time t
z = real concn. I*- a t time t k, = exptl. bimolecular reaction constant calcd. with the kt =
aswmption that I - and I; react with equal eaSe theoretical bimolecular reaction constant calcd. with , reacts much slower than the aswmption that 1 1-
The formation velocity of active iodoacetic acid is dx/dt = k.[y(n -
X)
- x ( b - r)]
ke(p
- xb)
(1)
kt(za
- XC)
(2)
,0049
,0041
,0031
or
,0067
.0057
,0049
The exchange between the free iodine and the iodine ions is very fast,s so
'oo59
'Oo4'
'oo'o
d ~ / ' d t= k t [ - [ a -
X)
- X(C
- z)]
z/c = y / b
The results are given in Table 11. Apart from the concentrations used and the exper'mental values of the reaction constant k, the values of the reaction constant found in the monoiodide experiments multiplied with the ratio of the real monoiodide concentration and the concentration of added potassium iodide are given too. This real monoiodide concentration has been calculated with the aid of the equilibrium: I- f 1 2 +r 13, using an equilibrium constant of about 0 . 00146 and also assuming that no free iodine remains in the solution. The formation of higher polyiodides has been neglected. If the exchange of radioactive iodine between the organic molecules and the IsIf?) J . S. Carter, J. Ckcin Soc , 2227 (1928)
[CONTRIBUTION No. 323 FROM
+
a = concn. CHzICOOH C€LI*COOH b = total concn. II*I; 1-; c = real concn. I I*x = concn. CH*I*COOH at time t
OB~ERVED AND CALCULATED REACTION CONSTANTS FOR EsCIIANGE AGAINST RADIOACTIVE MONOIODIDE AND TRIIODIDE Reac-
7G
ions is very slow compared with the exchange with normal iodide ions, we can calculate the experimental velocity constant in the following manner'
TABLE I1
Reac-
VOI.
dxldt
(3)
(4)
(5)
From Table I1 it is seen that the exchange reaction between 1; and CII2ICOOH is much slower than the reaction between I- and CHJCOOH; no conclusions can be drawn regarding the velocity of the former process. Acknowledgment.-This investigation represents part of the research program of the Foundation for Fundamental Research of Matter (F.O.N.). It was performed with the financial aid of the Netherlands Organization for Pure Research (Z.W.O.). ( 7 ) R. D. Heyding a n d C. A. Winkler, Can. J. Chem., 29, 790 (1951). (8) D. E. Hull, C. H. Shitlet a n d S. C . Lind, THIS J O U R N A L , 68, 535 (1936). h\ISTFRT)AM,
THE INSTITCTE FOR
k 5 (yo - rb) tb k, = c/bkt
=
HOTLAYD
ATOMICRESEARCH AND DEPARTMENT O F CHEMISTRY, I O W A COLLEGE1]
STATE
The Kinetics of the Reduction of Perchlorate Ion by Ti(II1) in Dilute Solution BY FREDERICK R. DUKEAND PAULR. OUINNEY RECEIVED FEBRUARY 6, 1954 The reduction of perchlorate ion by Ti(II1) in dilute solutions appears to proceed by several paths. The reaction has [H+] dependent as well as [H+] independent paths. The addition of chloride retards the reaction, apparently due t o the formation of TiC1+2ion. The reaction is linear in C101- at concentrations below about 1.0 M . Above this concentration the reaction rate increases, this probably being due to the increased activity coefficient of the perchlorate ion at higher concentrations. There was insufficient complex ion formation between Ti[ 111) and ClOa- to be kinetically detectable. The activation energies for the various reactions were calculated and a mechanism proposed.
Introduction The reduction of perchlorate ion by Ti(II1) a t room temperatures and slightly above proceeds to completion a t a convenient rate with the production of C1- and Ti(1V) STi(II1)
+ Clod- + 8 H + --f 8Ti(IV) + C1- + 4H20 work
On
the reduction
Of
perchlorate
(1) Work was performed in t h e Ames Laboratory of the Atomic Energy Commission
ion by Ti(II1) in dilute solutions was done by Rredig and ?\.lichel.2 They studied the kinetics of the oxidation of trivalent titanium compounds with perchloric acid in aqueous sulfuric acid and hydrochloric acid solutions, as well as the rate of reaction of HClO4 with Cr(II) and Mo(II) compounds. This present work examines more in detail the Ti(II1)-C104- reaction in solutions containing clllo(2) G. Bredig and J. hlichel, Z. fihysik. Chcm., 100, 124 (1922).
July 20, 1954
REDUCTION OF PERCHLORATE IONBY Tr(II1) : KINETICS
ride. A kinetic mechanism to explain the observed data has been presented.
Experimental All chemicals used were reagent grade. The TiCh solutions were prepared by dissolving titanium metal powder from Metal Hydrides, Inc., Beverly, Massachusetts, in approximately 6 N HCI. All solutions of Tic13 were prepared and stored under an atmosphere of nitrogen. Total chloride was determined by the nitrobenzene modification of the Volhard method for soluble chlorides. The Ti(II1) was first oxidized with excess Ce(1V) to prevent reduction of the Ag+. The Ti(II1) was standardized by “A.O.A.C.” Method I.* The H + was taken to be the difference between the Ti(II1) and total chloride concentration. Naclo4 was prepared by neutralizing 70-72% He104 with NaOH pellets. This was then standardized by passing through an acid exchange resin and titrating the acid produced with standard base. NaCl solutions were made by dissolving the required amount of the salt in water. HCl and HClO, solutions were used and standardized with standard NaOH. TABLE I Experiments run: All, starting Ti(II1) concentration, 0.15 M : TXIII) determined as a function of time: all concentrations in molarity. T ,‘ C . [H+I K1-1 30,40,50 0.23-1.0 2.0 (5 values) 40,50 0.50 0.22 0.24-1 . 4 (7 values) 0.50 40,50 0.50 0.24-1.4 (7 values) 0.50 0.75 0.26-1.6 40,50 (8 values) 0.50 40,50 1.0 0.51-1.4 (6 values) 0.02-2.0 40 0.22 2.0 (15 values) 0.5 2.0 0.1-2.0 40 (15 values) 0.1-2.0 1.0 2.0 40 (15 values) 0.25-6.0 20 0.23 0.24 (13 values) Each run was made in the following manner: The NaC1, NaC104, HClO, and HCl necessary t o make up a solution of the desired concentration were added to a 100-ml. volumetric flask with an extra calibration mark a t 96 ml. The solution was diluted to the 96-ml. mark and placed in a waterbath for about one half hour t o allow it t o come to the desired temperature. The determinations were made in a distilled water-bath using a mercury thermoregulator to control the temperature to =k0.05’. Four ml. of the Ticla solution was added, the solution shaken thoroughly, and the timer started. At known intervals of time 5-ml. samples were removed and quenched in an excess of quenching solution. This consisted of a solution of 0.01 N ceric ammonium sulfate approximately 9 N in H?SO(. The time of quenching was taken as the time a t which the meniscus passed a mark on the pipet, usually the halfway mark. The excess Ce(1V) was then titrated with ferrous ammonium sulfate using ferroin as indicator. First-order plots ~ l time. The experiments run were made from In [ T ~ i +8s. are summarized in Table I.
3801
Ti(II1). The rate expression for the disappearance of Ti(II1) may be written
-
dt
= k [ T ~ i + a ] ([C104-]),([Cl-]),([H+]) F~ (1)
Here k represents the specific rate constant for the reaction and F1is some function of the ion concentrations as yet to be determined. For any one run, the Clod- concentration is large compared to the Ti(III), and the H + and C1- are constant so that this expression becomes
-w = k‘[rTi+3J dt
(2)
where k‘ is the pseudo first-order rate constant. k’ is the slope of a plot of l n [ T ~ i t aus. ] time. By keeping all concentrations of the ions constant except the one whose dependence is being determined it was possible to determine the effects of H+, C1and clod-, on the rate of reaction. The ionic strength was allowed to vary, since two concentrations would have to have been varied simultaneously to keep the ionic strength constant. The reaction was found to be linear in C104- a t concentrations less than about 1.0 M . Above this concentration the rate increases as shown in Fig. 1 4.4
0
4.2
2.21 2.01 o
0.2 I
0.4 I
0.6 I
0.8 I
ID I 1.I2
1.4 I 1.6 I
1.8 I
2.0 I
Fig. 1.-Hydrogen ion dependence 40’; first-order rate constants versus [H+][Ti+J] = 0.015 AI‘, [ClO,-] = 0.5 M , [cl-l = 2.0 111.
Straight lines were obtained when In [ TTii a ] was plotted against time. This confirmed the results of Bredig and Michel, who also found the reaction between Ti(II1) and clod- first order in
The possibility of the formation of any Ti(II1)clod- complexes was investigated. Runs were made a t clod- concentrations up to 6 M and no lowering of the order in ClOd-was observed. The increased rate a t higher Clod- concentration is probably due to the increased activity of c104- and the correspondingly decreased activity of the water. This behavior was shown by Robinson and Baker.4 The hydrogen ion dependence was determined a t C1- = 2 M , clod- = 0.05 M , and original Ti(II1) = 0.015 M , and at 40’. The data are shown
(3) Association of Official Agricultural Chemists, “Methods of Analysis,” 6th Ed., 1045, p. 290.
(4) R . A. Robinson and 0. J. Baker, Trans. Proc. Roy. SOC.New Zealand, 76, 250 (1946).
Results and Discussion
Vol. 76
FREDERICK R. DUKEAND PAULR. QUINNEY
3802
4’21 \ 40
c
11.0 34-
\
I ao
80
-
N ’ 8.0-
0
.X
7.0-
1
5.0-
6.0
2.4
1
41)
-
0
I
I
0.2
0.4
I 0.6
I 0.8
I 1.0
I
I 1.4
I 1.6
I 1.6
I
2.0
Fig. 3.-Perchlorate ion dependence 40’; first-order rate constants versus [Clot] [Tifa] = 0.015 M , [Cl-] = 2.0 M, [H+] = 0 . 5 X .
The equilibrium constant K1 was determined from a study of the chloride dependence. The rate constants were separated and determined by considering the rate expression first a t constant H + and then at constant C1-. At constant H+, the rate equation becomes
where ki’
and ki’
+ k3[H+l = k? + k r [ H + ] = ki
The pseudo first-order rate constant may then be written k’ ~= kl’ + k2’K1[C1-1 1
[C104-1
and
[Tiel++] K 1 = [Tic3][Cl-]
Vi-riting the rate expression, equation 4, in terms of total titanium _ -d[T~i+3I = dt
+ Kl[Cl-I
The constants k,’ and kz’ were determined a t various H+, the C1- and C104- concentrations being the same and constant for each determination. Plots of kl’ and k,’ DS. H + yielded straight lines, the slopes of which were ks and ka, and the intercepts of which were k l and k l , respectively. Considering the rate expression at constant chloride, the pseudo constant becomes
+
k‘(1 ____ K1’)= k,” [c104-1
+
k2”[H+]
(10)
where K,’ ki”
=
+ k~Ki[Ci-]
ks
+ kaKi[Cl-]
and kJ“
=
K1[C1-]
ki
A plot of k ’ 11s. H+was made for different values This is the rate expressian l’ound from experit~~e~lt.of CI -. This yielded a series of straight lilies which
PERIODATE OXIDATION OF ETHYLENE GLYCOL : KINETICS
July 20, 1954
have slopes of kz" as a function of C1-, and intercepts with values of kl" a t various C1-. A plot is then made of kl" vs. C1- and k2" vs. C1-. These plots yield straight lines which have slopes of k1 and k2 and intercepts of kzKl and k4K1, respectively. The rate constants were calculated and the average values are shown in Table 11. TABLE I1 RATEAND EQUILIBRIUM CONSTANTS 40" [Ti+a] = 0.015 M ,[Clod-] = 0.5 M X 109 k, X 103 8.89 f 0.9 3.59 f0.4 ki
ki X 10'
kr X
IO2
3.4040.3 1.36f0.2
The rate constants were determined a t both 40 TABLE I11 ACTIVITY ENERGIES AND ENTROPIES 40' ki X 101 4FZ3~3,kcal. 4H:, kcal. 4S:, e.u.
k, X 101 4 F & kcal. AHi, kcal.
as:,
e.".
k3 >' 10' AF:sis. kcal.
4 H f , kcal. e.u.
AS:.
kr X 102 1.36 f 0.2 4Fzai3, kcal. 21.0 f 0.1 4 H i , kcal. 10 f 50% A s J , e.u. -36 f 50%
3.5 i 0.4 20.4 f 0 . 1 f 14% 25 I50% 17
KI
2.18
AH, cal.
AS, e.u.
* 0.2
17.8 i 0 . 1 2500 -50
+ Clod-
Ti+3 f ClOa-
+H
[complex] +TiO++
+ C103 + clos
+ Z [complex] +Ti(OH)+s
+ ClOa[complex] + TiO++ + c1os f C1- [or TiOCl+] TiCl++ + Clod- 4-H + J _ [complex] Ti(OH)+3 + c l o s + C1- [or TiOHCl+a]
TiCl++
--f
8.89 f 0.9 19.9 +e 0.1 26 f 15% 20 f 50%
A F m , kcal.
and 50'. The energies of activation and the entropies of activation for the various steps in the reaction were then calculated. The results are shown in Table 111. The reaction is thought to proceed through the formation of an activated complex, the disproportionation of which is the rate-determining step. The following four equations have been proposed to explain the kinetics of the reaction Ti+a
Kl 2.18 f 0 . 2
3803
i 50% i 100%
3.40 i 0.3 20.4 f 0.1 6 f 50% -46 i 50%
From the standpoint of the free energies involved and the probability of reaction most likely to occur, the first and third equations above would seem t o be the most important. The radical C103 is rapidly reduced further to the final product C1-. It may well be true that the affinity of the Ti(1V) for OH- and O= accounts for the low activation energy of this reaction as compared with reductions of perchlorate with Fe(II), Sn(II), Cr(I1) and others. AMES,IOWA
[CONTRIBUTION No. 321 FROM
THE
INSTITUTE FOR ATOMICRESEARCH AND DEPARTMENT OF CHEMISTRY, IOWASTATE COLLEGE^ ]
The Kinetics of the Periodate Oxidation of Ethylene Glycol and a Series of Methylated Ethylene Glycols BY FREDERICK R. DUKEAND VERNONC. BULGRIN RECEIVED FEBRUARY 8, 1954 The specific periodate oxidation of the homologous series of glycols, ethylene glycol through pinacol, shows a discontinuity in mechanism between trimethylethylene glycol and pinacol. The latter is very slowly oxidized in an acid-catalyzed bimolecular step, while the former, as with the other glycols studied, proceeds rapidly in a non-catalyzed unimolecular step. The effect of stepwise methyl group substitution in ethylene glycol is determined, both in the free energy of complex formation with periodate and on the unimolecular rate constant.
The mechanism of the periodic acid cleavage of ethylene glycol may involve an intermediate coordination compound of the oxidant and reductant, the rate-determining step being the disproportionation of the intermediate.2 The rate of the reaction has been demonstrated to be nearly independent of hydrogen ion concentration in the pH range three to even.^^^ Waters4 attributed the constancy of the rate in this region to the fact that hydrogen ion catalysis is balanced by decreasing ionization of HJOs to the reactive HJOe-. Crouthamel, Martin and co-workers, in a spectrophotometric study of periodate species in solution,jS6showed that the
concentration of negative ion, predominantly IO4-, is nearly constant in the PH range three to seven. Price3found the rate of oxidation of pinacol to be a peculiar function of hydrogen ion concentration, having a sharp maximum a t p H 2 and a minimum for the range four to six. A kinetic study of the oxidation of the series ethylene glycol through pinacol was undertaken, in an effort t o observe the effect of methyl substitution on the rate and equilibrium constants and to determine a t which point in the series the change in reactivity exhibited by pinacol begins to show up. Experimental
(1) Work was performed in the Ames Laboratory of t h e Atomic Ethylene and propylene glycols obtained from the MatheEnergy Commission. son Co. were twice distilled before using. meso-2,3-Butane(2) F.R. Duke, THIS JOURNAL, 69, 3054 (1947). diol, which had been prepared by the fermentation of corn (3) C.C.Price and M. Knell, i b i d . , 64, 552 (1942). using Aerobacter aerogenes,? was purified by recrystallization (4) W. A. Waters in H. Gilman, "Organic Chemistry," Vol. IV, from dry isopropyl ether.* levo-2,3-Butanediol was preJohn Wiley and Sons, Inc., New York,N. Y., 1953, Chap. 12 (7) E. R.Kooi, E. I. Fulmer and L. A. UnderkoBer, I n d . Eng. Chon., (5) C.Crouthamel, H.V. Meek, D. S. Martin and C. V. Banks, THIS 40, 1440 (1948). JOURNAL, 71, 3031 (1949). (6) C.Crouthamel, A. Hayes and D. S . Martin, ibid.. 73, 82 (19.51). ( 8 ) C . E. Wilson and K. J Lucas. T i i r s J O U R N A L , 58, 2396 (1936).
