447
KINETICS OF THE CHLORATE-SULFITE REACTION
April, 1957
data are available to prove this. When plots are made of the B-coefficient of the anion or the cation versus delta crossover temperature, the agreement is best using the anion values although it is not as good as when the B-coefficient for the salt is used. As the B-coefficient has been related to the temperature coefficient of electrical mobility, the lyotropic number, the ionic entropy and the activity coefficient of the ions,44the delta crossover temperature is likewise related t o these properties. The B-coefficient is determined in very dilute solutions whereas the delta crossover temperature is determined in almost saturated solutions. For the other properties previously mentioned such as critical temperature and the dielectric constant, not more than two or three values could be found for comparison. The above correlations demonstrate that the lowering of the crossover temperature and the separation factor are functions of the order-disorder in the solution and that the ions do cause changes in the structure of the water as well as form structures about themselves. Some additional factor must be taken into consideration for the multivalent ions as they did not follow the same relationships as the univalent ions. Small ions, such as the lithium ion and the fluoride ion, do not tear up the water structure rn much as do the large ions, such as the iodide ion and the phosphate ion. Further, anions do affect the sep-
100 1
1
. I
-I
- -*
-n-.--inn 0.000 0.100 0.200 0.300 0.400 Fig. 3.-The lowering of the crossover temperature versus the viscosity B-coefficient for various salts.
aration, and thus the hydrogen bond, more than the cations. It is evident that potassium orthophosphate changes the structure of water to the greatest extent while sodium sulfate changes the structure the least. It is probable that the phosphate ion hydrates the light water preferentially and thus forms a structure about itself in addition to tearing up the water structure. The solubility of sodium sulfate was very slight (as indicated by plugging of the liquid sampling tube), and this could explain the results obtained. This research demonstrates clearly that it is the structure of the liquid phase which causes the crossover temperature as well as a separation factor above and below the crossover temperature. Acknowledgment.-The authors are indebted to the United States Atomic Energy Commission for support of this work.
KINETICS OF THE CHLORATE-SULFITE REACTION BY E. H. GLEASON,* G. MINOAND W. M. THOMAS Contribution from Stamford Laboratories, Research Division, American Cyanantid Company, Stamford, Conn. Receioed Octuber 11, 1966'
Kinetics of the chlorate-sulfite reaction in water were investigated a t 0 and 20" a t pH between 1 and 3.5. The rate of disappearance of sulfite was followed by iodometric titration and the rate of formation of chloride ions was determined amperometrically with silver nitrate. The data show that the reaction is second order and that the rate is pH dependent. A mechanism is proposed involving the reaction between H 8 0 s molecules and chlorate ions as the rate-determining step. I n units of liters, moles and seconds, the reaction rates were formulated a5 -d[ClOa-]/dt = 3.3 X lo7 e-110p0'RT[C103-!. [H.J303]. The reaction is not much influenced by traces of metals, even though these metals strongly accelerate polymenzations initiated by this system. The exact nature of the free-radical intermediates 1s thus stlll in doubt, but I t seems hkely that sulfoxy radical-ions and hydroxyl radicals are involved. Hydrogen ion concentration determlnes the HzSOs level and thus the reaction rate.
I. Introduction Although the iodate-sulfite reaction has been studied extensively, the Only previous kinetic investigation of the chlorate-sulfite reaction appears to be that Of and in The reaction was assumed to be ClOa-
+ 3HzSOa --+ C1- + 4504- + 6H+
ClO3-
+ 3HS03- +Cl- + 3SOa' + 3H'
or
Iodometric
lneasured the loss Of sdfit'e and, from these data, a secold-order constant K of 2.4 liters mole-' min.-' a t 0" was calculated for the equation
*
New York State College of Forestry, Syracuse, New York. (1) A. C . Nixon and K . B. Krauskopf, J. Am. Chen. Soc., 64, 4GOG (1932).
-d[ClOs-]/dt
= K[HzSOa][ClOa-I =
[Hi] [HSOa-][C103-]
KO Fromthe fact that the chlorate-sulfite systeln can be used to initiate polymerization,2 it might have been inferred that one or more intermediates are free radicals. Isotope exchange experiments3 have shown that in the sequence C103C10C1- 4 most of the oxygen is transferred directly from chlorine to sulfur. 11. Experimental
-
Solutions of reagent grade NazSOa, NaC103 and H2SOl in deaerated water were placed in a thermostated flask, acid being added last. The flask was arranged so that 25-ml. (2) A. Ciesswell, U. S. Patent 3,751,374 (195G),t o American Cyanamid Co.; A. Hill, U. 8. Patent 2,1573,192(1954), to Diamond Alkali
Co. (3) J. Halperin and H. Tabbe, J. Am. Chem. Soo., 73, 3310 (1950).
448
E. H. GLEASON,G. MINOAND W. M. THOMAS
Vol. 61
ali uots could be forced into a pipet by COZ pressure. $or sulfite determination, the aliquot was released into used to compute p, was t,aken as 3.10 X loT2at at excess 0.1 N iodine solution and back-titrated with 0.025 N 0" and 1.91 X At pH levels above 2, release of H + during the thiosulfate. For chloride determination, another aliquot was released into sufficient 2 N NaOH to neutralize the solu- reaction is sufficient to shift the value of p. The tion and stop the reaction. A solution of 10% Ba(NOa)8 rate equation was then taken as waB added to precipitate SOa^ and SO4-. After 20 mnutes, the precipitate was filtered and washed free of C1-. The AZsoz = 3K[C1Oa-][H8Os] filtrate was acidified with *NOl to pH 1 and titrated amAt perometrically with 0.05 N AgN03, using a rotating platinum electrode and a calomel electrode. and K was calculated from tangents on plots of For measurement of pH, Leeds and Northrup glass elec- 2302 (total sulfite) DS. t. trode 1199-44 was used a t 0' and glass electrode 1199-30 a t Results of these measurements are summarized higher temperatures.
-
in Table 11. For the two temperatures K is given 111. Results by Rate of formation of the chloride ion was deterK = 2.0 X 109 e-11OOOIRT 1. mole-1 min.-l mined at 1' at a constant pH of 1.20. The rate of K = 3.3 X IO7 e-11o00/RT 1. mole-' 6ec. -1 disappearance of sulfite was determined conTable I1 also summarizes the effect on rate of comitantly on a duplicate solution. Results reported in Table I show that the rate of formation of several additives. Agents expected to influence chloride is one-third the rate of disappearance of radical reactions had little effect, but NaCl and, sulfite. Since analysis for total sulfite is more con- particularly, NazSOaretarded the reaction.
IV. Discussion From the dissociation constants of the sulfoxy acids, one can calculate readily that [SO3-] is negligible below pH 5. Below pH 4,both [HSOt-] and [H,SOs] are important, with HzS03 comprising about 35% of the total sulfite (ZS02) at pH 2 and about 85% at pH 1. A small change near pH 2 drastically alters the ratio of [NzSO3] to [HSO,-]. Since the reaction is faster as [H+] increases, one is led to identify H2SOa and C103as the active species. It was shown that the rate of formation of chloride ions is one-third the rate of disappearance of the total sulfite
1.4 -
2
1.2 1.0 -
-2 Q 0.8 -
T
50.6
-
0.4
-
2
8 10 12 14 16 18 20 Minutes. Fig. 1.-Second-order plots of the chlorate-sulfite rerun 17 a t 20"; (0) run 1 a t 0 '; ( X ) action a t p H 1: ).( data of reference 1. 4
6
-d[C103-] d[C1-] = e
dt
d[ZSOzl
dt
dt
If the bimolecular reaction takes place between a
chlorate ion and a molecule of H2S03, the reaction product, whatever it may be, must react very venient than chloride determination, the kinetic rapidly with two other molecules of HzS03to give runs discussed below relied entirely on total sulfite chloride ions. Only a mechanism of this type is consistent with the above rate expressions. The analysis a t controlled pH. stepwise oxidation-reduction probably involves TABLE I the reduction of the chlorate ion to chlorite ion STOICHIOMETRY OF THE CHLORATE-SULFITE REACTION (CIOz-), followed by subsequent rapid reduction to hypochlorite ion ((310-) and chloride ion (Cl-). AZSOp Min0.0524 utes ZSOn - zsoz c1AC1 The ratio of the total sulfite to H2S03 can be cal5 0.0288 0.0236 0.00790 2.99 culated, a t low pH, as
-
'
10 15 20
.0208 ,0163 .0135
,0310 .03Gl .0389
,0103 .011G .0123
3.02 3.11 3.16
Reaction rate could be followed roughly by change in p H , but the method lacked sensitivity, especially at low pH. In the Discussion, equation 8, we derive the relationship 1 Bo = Kt Bo In A 3PAo A OB where A = [ClOa-1, B = [H2SOx]).p = [KzSOJ/ ZSO, and subscripts zero refer to time zero. Figure 1 shows typical plots of log ABo/AoB D S . t (see Table I1 for conditions). Data of Nixon and Krauskopfl have been replotted according to our scheme. The first ionization constant of H$03,
[H,SOa]
+ [HSOs-] = ZSOz
[HSOa-]
[& [HzSOsl
where K' is the first dissociation constant of sulfurous acid. Since HzS03 and HS03- are in equilibrium, p remains constant throughout the reaction, at constant pH, and H2SOa is produced continuously by the ionic equilibrium H+ HSOa- % H2S03. If A == [C103-], B = [H2S03], C [ClOz-1, D = [CIO-1, E = [Cl-] and X = S O 2 , the reaction
+
(4) Landolt-Bdmstein, "Physikalisch-Chemiacho Tabellen," Vol. 111, Julius Springer, Berlin, 1936, p. 2105; H.V. Tartar and H. 11. Garretson, J . Am. Chem ~ o c . 63, , 808 (1941).
. I
KINETICSOF
April, 1957
THE
449
CHLORATE-SULFITE REACTION
TABLE I1 SUMMARY OF RATEMEASUREMENTS Run
.. 1 2
3 4 5 6 7 8 9 10 11 12 13 14 15
16 17 18 19 20 21
Initial PI-I
..
LNaClOalo
.1 .1 .1 .01 .Ol .01 .Ol .01 .0068 .1
0.01310 ,01200 .01310 .0262 .01922 ,01200 .01314 .01314 .01314 ,01326 .01244 .01388 ,01318 ,01335 .01250 ,01970
Experiments a t 0' 0.0160 ,0160 .0160 ,0080 .0200 ,0160 ,0160 ,0160 ,0160 .0160 .0160 ,0160 ,0160 ,0160 .0160 .0150
0.1 .1 .1 .Ol .01 .0011
0.01314 .01330 .0262 .01325 .01244 .00135
Experiments a t 20" 0.01604 ,01604 .00800 .01604 .01064 ,000567
0.1 .1 .1 .1 .1
1.05 1.05 1.05 1.05 0.85 1.05 1.05 1.05 2.4 2.4 2.4 ea. 2 . 3 2.46 3.37 1.05
....
1.05 1.05 1.05 2.4 2.4 3.0
mechanism can be .written as follows (omitting transient radical species)
+
H + HSOaB (1) A B +C SOa' 2H+ (2) C B +D so4' 2Hf (3) D+ B e E SO4"+ 2H+ (4) At constant pH, B = pX and dB/dt = p(dX/dt)
+
+
+ + +
+ +
where dB/dt is the total rate of change of B. On the other hand, the rate of consumption of A is
-
dt = KIAB
=i
KlpAX
and the total rates of change of D and B are
(5) ,
K
(I. mole-lmin.-
[NaiSOa 10
[HsSOrlo
3.97 4.20 3.77 3.58 3.71 3.46 3.94 3.74 3.80 4.3 4.0,4.5 4.5 1.7 3.2 3.4 3.98 15.2 15.2 15.6 16.2 16.6 19.5
Notes
1)
Recalcd. from ref. 1
0.1 M HC1 replacing &So4 Soln. 1 M in acrylonitrile 10-8 M Cu++ 10-3 M Fe++ Soln. 0.543 M in acrylonitrile 0.017570 benzoquinone 0.1 M Na2S04 0.3 M NaCl
Soln. 0.543 M in acrylonitrile
tial equation which can be readily integrated
-
dt
= KlAB = KlA[Bo
Plots of 2.3 log ABo/AoBversus time are shown in Fig. 1. The slopes of these lines give the value K I (3p AO - BO),from which the second-order rate constant K1 can be calculated. Solving equation 8 for A gives L A = 5eKiLt
dt = KlAB
- K2BC
(6)
dt = K2BC
- KaBD
(7)
+ 3pA - 3pA0l
A0
- 3p
(9)
where L = BO- 3 ~ A o . At constant pH, L is constant and is a measure of the excess of one reactant over the other. (Note if reactions 3and 4 are much faster than 6, C and D that L/p = Bo = 3A0; L / p = ZSOz - 3[C103-]0.) will be used up as formed. Under these conditions, Three situations should be considered. the concentration of C and D will remain very 1. L > 0, Bo > 3pAo; an excess of HzS03 is small and a stationary state will be reached where present. In this case eK1Lt will increase exponentially with time, while the function A(t) will tend dC - = O , - = Od D to zero very rapidly if L is large. This means that, dt dt given several recipes, containing sulfite and chloand KIAB = KzCB = KIDB. The rate of dis- rate in various ratios Bo/Ao,the concentration of appearance of X is the chlorate will decrease more rapidly the larger the excess of HzS03present. - dt = KlAB + KsBC KaDB = 3K1AB 2. L < 0; the chlorate is present in excess. hence In this case e K l L t will decrease exponentially; consequently the concentration of the chlorate dX dA -dt =3-& will tend to the limit-L/3p more rapidly as L increases. and, upon integration, one obtains 3. L = 0, Bo = 3pAo; the chlorate and the X Xo 3A - 3Ao sulfite are present in the quantities required by B = Bo 3pA - 3pA the stoichiometric equation. In this case equation 5 Substituting this value of B into 5 gives a differen- becomes indeterminant. Since B = 3pA
+
+ +
ALVINW. BAKER
450 dA
ai?= and solving for A
Vol 61
- KiAB = - Ki3pA2 1 2
=
1 + 3pKit Ao
Figure 2 shows how the product [C103-] [HzSOs] decreases with time, in four different recipes containing the same initial product [ClOa-]o[HzS03]0 of 5.73 X lods, but different L. The concentration of [C103-] a t various instants was calculated min.-' from equation 9, taking K1 = 4 1. moles and p = 0.763. It is interesting to note that the product decreases with time more rapidly in systems containing either an excess of chlorate or sulfite than in systems containing stoichiometric quantities of the two reactants. The larger the absolute value of L , the faster the product tends to zero. This fact becomes very important when the rate of polymerization is studied as a function of the initiator concentration,6 because systems having the same product [C103-] [HzSOa], but different concentration of [C103-] and [HzSOa],will contain the same number of free radicals only initially. This investigation has not established with certainty the nature of free radical intermediates. One possibility in line with the data at hand is
- 0
0
IO
20 MINUTES.
30
Fig. 2.-Variation in reaction rate with initial chlorate/sulfite ratio: (0) L = 17.2 X 10-8; ( X ) L = 0; ( 0 ) L = 12.2 x 10-3; (0)L = -17.2 x 10-3. OC10HSOs. HO. HS03.
+ HOS.00 + .OH
+ .OH +2H+ + so4' + monomer +polymer
K6 + 6Fe++ + 3 H + +
C10~-
C1-
+ monomer +polymer
6Fe+++
The influence of metals like Fe++ and Cu++ is of interest. Our data indicate that small concentrations mole-') of ferrous ions do not increase the over-all rate of the chlorite-sulfite reaction, but increase appreciably the rate of polymerization of acrylonitrile.6 The following reactions can take place when c103-,HzS03 and Fe++ react in acid medium c103( 5 ) W.
press.
K4 + 3H1S03 + C1- + 6H+ + 3SO4--
(10)
M. Thomas, E. H. Gleason and G. Mino, Polymer Sci., in
+ 3H2SOs + 3&0
+ 6Fe++++ 3 0 H -
K6
+
6Fe++ On summing, 2c103-
+ 3s04-- + 12H+ (12) 2C1- + 6S04-- + 12H+
4- 6H2SOs +
Each of these reactions proceeds through a free radical mechanism and is capable of initiating polymerization. It is interesting to note that the iron present is continuously oxidized and reduced by reactions 11 and 12. If Ka is much larger than K6, a steady state will be reached ih which the, concentration of the ferrous ion will remain constant.
SOLID STATE ANOMALIES I N INFRARED SPECTROSCOPY BYALVINW. BAKER Research Department, Western Division, The Dow Chemical Company, Pittsburg, California Received October 16,ID66
Variations between mull and pellet spectra of organic compounds are due either to an induced physical isomerization or to the samples' having been rendered amorphous in the alkali halide pellet. Factors which influence these changes are: (1) crystal energy of the organic phase; (2) energy of grinding sample and matrix; (3) lattice energy of matrix; (4) particle size of matrix; ( 5 ) ability of sample to recrystallize in the pellet (related to crystal energy); and (6) relative stability of polymorphic forms. Relative merits of mull and ellet techniques are presented and i t is shown how these can supplement each other. From the frequency of occurrence ofpolymorphism, it is concluded that this is a general rather than a rare problem facing the organic chemist.
