Table III. Effect of Reversibility upon the Detection Limits Conditions: p = 0.60f = 200 Hz Eatep= -0.10 mV (0.02 V sec-l) AE = 50 mV A = 6.70 X 10-*cm2 T = 25°C Fe(II1) Ti(1V) Cr(II1) 1M tartaric acid 1M potassium chloride Supporting electrolyte: 0.5M potassium oxalate 1 X 1W6cm sec-l >1 cm sec-1 9 X 10-8 cm sec-1 k a l h (16) 1.5 X 1W6M 5.4 x 10-6M Detection limit conc: 1.8 x 1 0 - 7 ~
tectable concentration is about ten times greater than for a reversible reduction; total irreversibility further diminishes the advantages of SWV. Anodic Stripping Analysis. The utility of ac (17) and square wave ( l a ) polarography in anodic stripping analysis has already been well established. The application of SWV to anodic stripping analysis has been attempted. Well resolved peaks for 5 X 10-'OM Cd(1I) and Pb(I1) in 0.5Mpotassium nitrate were obtained with a 5-minute plating time, utilizing equipment readily available in any laboratory. Reasonable linearity of i,/Ct (where i, is current, t is the plating time, and C is solution depolarizer concentration) was obto 5 X 10-*M. tained in the concentration range of 5 X The authors did not attempt to extend the lower concentration limit any further as they felt that impurities at this concentration (5 x lO-'OM) are present in molar solutions of most analytical reagents that are usually employed either as supporting electrolytes or in the preparation of the sample for analysis. In theory, the plating time employed (5 minutes) could be extended to achieve the detection level (6 X 10-llM)
of first derivative linear sweep anodic stripping analysis (19). However, for stripping times greater than 20 minutes, Kaplan and Sorokovskaya (20),using square wave polarography, have noticed pronounced deviation of i,/Ct to lower values under conditions where linear sweep stripping methods maintain a constant value for this factor. This was probably caused by electrode surface contamination which will affect ac methods to a greater extent than dc methods. Despite its limitations, square wave stripping analysis did achieve reasonable detection levels at shorter analysis times than other methods. In conclusion, it can be seen that square wave voltammetry offers all the advantages of square wave polarography but is accomplished at shorter analysis times, and except for the problem of electrode contamination, which is common to all stationary electrode techniques, it is a rapid, sensitive, simple analytical technique. ACKNOWLEDGMENT
The authors acknowledge the technical assistance received from the Burr-Brown Research Corp., Tucson, Ariz. RECEIVED for review May 1,1969. Accepted June 17,1969.
(16) J. E. B. Randels and K. W. Somerton, Trans. Faraday SOC., 48,938 (1952). (17) W. L. Underkofler and I. Shain, ANAL.CHEM., 37, 218 (1965). (18) F. von Sturm and M. Ressel, 2.Anal. Chem., 186, 63 (1962).
(19) S. P. Perone and J. R. Birk, ANAL.CHEM., 37,9 (1965). (20) B. Ya. Kaplan and I. A. Sorokovskaya, Zavod. Lab., 30, 1177(1964).
Kinetics of the Iodide-CatalyzedReaction between Cerium(lV) and Arsenic(ll1) in Sulfuric Acid Medium Pedro A. Rodriguez' and Harry L. Pardue Department of Chemistry, Purdue University, Lafayette, Ind. 47907 The oxidation-reduction reaction between Ce(lV) and As(ll1) as catalyzed by iodide has been studied in 0.5~4 sulfuric acid medium. At large ratios of [As(lll)]/ [Ce(lV)] the reaction rate is proportional to iodide concentration. For values of this ratio of unity or below, the reaction rate is a nonlinear function of iodide concentration. There is a slight dependence on As(V) but no Ce(ll1) dependence. A reaction pathway is presented which yields a rate equation which satisfies the experimental observations over a wide range of conditions. Effects of HzS04, HC104, Ag(l), Hg(tl), CI-, and temperature on the reaction are discussed.
THECATALYTIC EFFECT of iodine on the oxidation-reduction reaction between Ce(1V) and As(II1) has been used for many years for the determination of ultra-trace quantities of iodine and iodide. The net reaction is: 2Ce(IV)
+ As(II1) = 2Ce(III) + As(V)
(1)
Present address, Prater and Gamble Company, Cincinnati, Ohio.
This reaction was first reported and studied in some detail by Sandell and Kolthoff (I, 2). These authors established empirical conditions for the determination of iodine and iodide at concentrations down to 20 parts per billion (ppb) and also presented an extensive discussion of species which interfere with the determination. Following this initial work Reaction 1 has been applied to a number of analytical problems including the determination of iodide in common salt (3),total iodine (including iodate) in sea water (4), and protein bound iodine (PBI) (5,6). Many of the application studies have included investigations (1) E. B. Sandell and I. M. Kolthoff, J . Amer. Chem. SOC.,56, 1426 (1934). (2) E. B. Sandell and I. M. Kolthoff, Mikrochim. Acta, 1937,9. (3) M. Dubravcic, Analyst (London), 80, 146 (1955). (4) R. A. Barkley and T. G. Thompson, ANAL.CHEM.,32, 154 (1960). (5) J. J. Moran, ibid., 24, 378 (1952). (6) R. D. Strickland and C. M. Maloney, ibid., 29, 1870 (1957). VOL. 41, NO. 11, SEPTEMBER 1969
0
1369
of some of the kinetic aspects of Reaction 1 (3-7). Several general conclusions can be drawn from these studies. It is generally assumed that Reaction 1 is first order in Ce(1V) and shows a complex dependence on As(II1) concentration. It is assumed that for iodide concentrations above about 20 ppb (-1.5 x lO-7M) the reaction is first order in iodine concentration. It has been reported (7) that for iodine concentration below about 4 ppb, kinetic data are erratic and the first order dependence on iodine concentration is lost. Chloride ion is reported to effect the dependence of the reaction on iodide. The data of Sandell and Kolthoff (2) indicate that sodium chloride (and other salts) present in concentrations above about 0.100 gram per 5.00 ml solution inhibits the catalytic reaction. Other authors (7, 8) have reported that lower concentrations of chloride enhance the reaction rate. Despite the rather large volume of literature which exists for this reaction, many questions concerning the kinetics of Reaction 1 remain unanswered. Most of the studies have been directed at solving a specific analytical problem; consequently, attention has been focused upon information pertinent to these problems rather than a complete understanding of the reaction. Virtually all of the kinetic data reported to date have been based upon observations of the integral change in reaction over time periods of several minutes rather than initial rates. Consequently, it is difficult to relate the kinetic information to any specific set of reaction conditions. While this type information may have been perfectly adequate for the types of analytical procedures developed, it is not very useful in obtaining an accurate representation of the kinetic behavior of the reaction. To date there has been no quantitative kinetic expression reported for Reaction 1 which is valid for more than a narrow range of conditions. Detailed kinetic information can be valuable in extending the analytical utility of this type reaction, both from the equilibrium (9) and kinetic points of view. The purpose of the present study is to describe in detail the quantitative dependence of Reaction 1 upon the catalyst (iodide) and all reactants and products over a wide range of conditions. In addition, a limited amount of information on the temperature and acid dependence and the effect of chloride on the reaction is presented. The availability of a photometer system with ultra-low drift characteristics (IO) has made it possible to obtain rate measurements very near zero reaction time where reaction conditions can be defined very accurately. Some of the more pertinent observations are summarized here. It is observed that the dependence on Ce(1V) concentration is first order over a rather wide range of Ce(1V) concentrations. However, when the ratio [Ce(IV)]/[As(III)]becomes large, the Ce(1V) dependence approaches zero order. The exact ratio at which this occurs is a function of the absolute concentrations of these species. The As(II1) dependence is first order at low ratios of [As(III)]/[Ce(IV)I and approaches zero order as this ratio increases. As in the case of Ce(IV), the concentration ratio at which the zero order dependence occurs is a function of the absolute concentrations of Ce(1V) and As(II1). The iodide dependence is observed to be first order at low iodide concentrations (and very low ratios of [I-]/[As(III)] but to approach zero order as the iodide concentration is increased at constant As(II1) concentration. It is observed that even at very low iodide concentrations the slope of the (7) A. Lein and N. Schwartz, ANAL.CHEM., 23, 1507 (1951). (8) J. Deman, Mikrochirn. Acta, 1964, 67. 40, 139 (1968). (9) A. J. Zielen, ANAL.CHEM., (10) H. L. Pardue and P. A. Rodriguez, ibid., 39,901 (1967). 1370
ANALYTICAL CHEMISTRY
rate us. iodide concentration plots is unaffected by chloride concentrations up to 0.125M. The observation (4) that iodate is inactive as a catalyst is confirmed in this work. It is also confirmed that if iodide is subjected to prolonged treatment with Ce(1V) it becomes catalytically inactive and further that subsequent treatment with an excess of As(II1) restores the catalytic activity of the iodide much the same way as As(II1) reduces iodate to a catalytically active form. It is observed that Reaction 1 shows a slight but real dependence on A@), but no detectable dependence on Ce(II1). A mechanism is proposed which satisfies the experimental observations. The proposed mechanism is utilized to derive a kinetic expression which accurately describes the quantitative behavior of the reaction throughout the range of conditions studied. EXPERIMENTAL
All solutions to be added to the reaction cell (except catalyst) were handled using tuberculine type hypodermic syringes equipped with polyethylene needles and Chaney adaptors (Hamilton Co., Whittier, Calif.). All tuberculine syringes used were calibrated by weighing the mass of water dispensed and were found to be reproducible to *0.04%, The Ce(1V) and As(II1) solutions were added using 1.00 ml syringes and the catalyst solution was added using a 100-111 syringe. In most cases a small volume of catalyst solution was added rapidly to a stirred solution of Ce(1V) and As(II1) in the cell. In this way complete mixing (per cent transmittance to within 0.02% of the final value) was achieved within about five seconds. Thus, in most cases, measured reaction rates correspond closely to initial rates. The results reported below are based upon the continuous photometric detection of Ce(1V) in the reaction mixture. Because the kinetics of the formation of a stable Ce(1V) absorbing species in sulfuric acid are complex, extreme care must be exercised to ensure that all experimental observations are accurately translated into concentration changes. When Ce(NO& (NH4)2is dissolved in 0.5M H2S04 the light transmission properties change with time for several hours, presumably due to the slow kinetics of the formation of Ce(1V)-sulfate complexes. Therefore, all Ce(1V) solutions were aged at least 24 hours prior to use. Also, the transmission properties are temperature dependent. For example, for a solution with 25 % T at 407 nm, the temperature coefficient at 25 "C is about -0.08% T per "C. A solution which is removed from the equilibrium temperature in the photometric cell by a few degrees for a short period of time will require several seconds to return to a stable (+0.01 T) reading. Consequently, most of this work was carried out at 25.00 i 0.02 "C with room temperature controlled at 25 rf 1 OC. All solutions and transfer syringes were equilibrated to the working temperature prior to transfer of reagents to the cell. All transfers were performed with minimal delay. The spectral response of CeOV) in sulfuric acid is affected by both As(II1) and As(V) with As(II1) producing a bathochromatic shift and As(V) producing a hyposchromic shift. For example, the molar absorptivity of 7.5 X 10-4M Ce(IV) at 407 nm changes from 650 in 0.5 H2S04to 657 nm in a solution containing 8 x lO-4M As(II1). The molar absorptivity of the solution containing 7.5 X 10-4M Ce(1V) and 8 X 10-4MA~(III)changes from 657 to 616 nm in a solution containing 2.5 X 10-2M As(V). The concentration error introduced by ignoring the effects of As(II1) and As(V) are relatively small for Ce(1V) concentrations below 5 x 10-4M but can result in errors as large as 10% for Ce(IV) concentrations in the 10-ZM concentration range. The errors at high Ce(IV) concentrations are enhanced by the fact that many of the measurements are made at 500 nm where the effects of As(II1) are more pronounced. Because of these complications the photometer response to Ce(1V) was cali-
I
.-
Co
/
/+
12
X
;10 u1
L
.-2
-
.8
-I
- .6P
u
0 .4 +
a
0:6 0.9 1.2 Ratio CCe a3 I C A S m3
I
100 200 300 400 500 C e a - I . reaction time ( s e c ) prior to
Asm addition
Figure 1. Time dependence on deactivation of catalytic activity of iodide by Ce(1V) [Ce(IV)] = 7.96 X 10+M [As(III)] = 1.62 X 10-*M [Itlo = 1.35 X 10-?M
1.5
1.8
Figure 2. Rate dependency on Ce(1V) at small ratios of [Ce(IV)] to [AdIII)] - Predicted rates for [I-] = 8.15 X 10-8M = 2.72 X 1 0 - 7 ~
- - - Predicted rates for [I-] IAs(II1)I
e 8.09 x 10-3 + 8.09 X lo-’ o 8.09 x 10-5 Q 2.43 X lo-‘
11-1
8.15 x 8.15 x 8.15 x 2.72 X
10-8 10-8
10-8 lo-?
Note changes in ordinate magnitude brated for each set of conditions prior to the injection of the catalyst into the reaction mixture. After the molar absorptivity of the Ce(1V)-As(II1) solution is determined, the catalyst is added to the stirred solution in the cell. The volume of catalyst added is small compared to the total volume in the cell. In this way, dilution effects and temperature changes (and, therefore, apparent deviations from Beer’s law) which could otherwise be caused by the addition of catalyst solution, are minimized. The initial T changes by 1 % T or less as a result of catalyst addition. The change in initial T caused by dilution can be accurately predicted using Beer’s law. Agreement with experimental results is within 0.02 T . The photometer utilized in this work (10) yielded a conT us. time. It is easily shown that for a tinuous record of first order reaction detected by following a system which obeys Beer’s law, a plot of % T us. time approaches linearity in the range of 20-60z T . The slope of the resulting curve is proportional to the rate constant. For most of this work T values between 20 and wavelengths were selected to yield 6 0 z T for the Ce(1V) concentrations used. Consequently, straight line traces were obtained and initial reaction rates were obtained by extrapolating these traces back to zero time. Zero time is easily identified by the sharp break in the % T us. time plot at the point where catalyst is injected. Because the photometer is calibrated for each new solution and because the total change in % T observed is small (usually less than 3z T ) , this procedure leads to very accurate and reproducible measurements of reaction rate.
z
z
RESULTS unless Results reported below were obtained in 0.5M stated otherwise. All quoted reaction rates correspond to initial rates unless stated otherwise. Also, iodine is added as potassium iodide in distilled water. Solutions of potassium iodide in distilled water are stable over periods of several weeks. A decrease in catalytic activity is observed when the solutions are prepared in 0.5M sulfuric acid. The solid line curves in Figures 2-6 represent computer fits of the experimental data based on the proposed mechanism and rate equation.
Order of Reagent Addition. The effect of varying the order of addition of reagents was studied prior to determining concentration dependencies. The reaction rate remains unchanged at its maximum value when either iodide or Ce(1V) is added to a stirred solution of As(II1) and the other reagent. However, when As(II1) is added to a stirred solution of Ce(1V) and iodide, there is an appreciable decrease in reaction rate from the observed maximum for the particular set of concentrations in the cell. The extent of loss of catalytic activity depends upon the time iodide has been in contact with Ce(1V) prior to the addition of As(II1). Figure 1 presents representative data for the time dependence of the loss of catalytic activity as a function of time for 1.36 X lO-7M iodide in contact with 7.96 X 10-4MCe(IV) in 0.5M H2SOa for the indicated time periods. The ordinate is plotted as the fraction of the maximum rate which is determined by initiating the reduction by adding either iodide or Ce(1V) initially. If iodide and Ce(1V) are permitted to react in the absence of As(II1) for a period of a few hours, the catalytic activity of iodide is completely destroyed. If an excess of As(II1) is added to a solution which has been treated in this way, and the solution is permitted to react for sufficient time, then the full catalytic activity is restored. The catalytic behavior of the product of the Ce(1V)-Ireaction is the same as iodate. Iodate shows no catalytic activity immediately after addition to a solution of Ce(1V) and As(II1). However, in such a solution, there is a gradual increase in activity with time. Also, if iodate is permitted to react for several minutes at elevated temperature with As(III), then catalytic activity equal to an equivalent amount of iodine is observed. Those observations are discussed quantitatively elsewhere (11). The remaining results were obtained by adding iodide to a stirred solution of Ce(1V) and As(II1). (11) P. A. Rodriguez and H. L. Pardue, ANAL.CHEM., 41, 1376 (1969). VOL. 41, NO. 11, SEPTEMBER 1969
1371
______--.----_____---
Y
a 8
3
6
9
12
Ratio tCe !Kl/CAS !a1
15
18
Figure 3. Rate dependency on Ce(1V) at large ratios of [Ce (IV)] to [As(III)] - Predicted rates for [I-] = 8.15 X 10-8M = 2.72 X lO-7M
- - - Predicted rates for [I-] [As(III)I
11-1 8.15 X 10-8 + 8.09 X lo-‘ 8.15 X 10-8 o 8.09 x 10-5 8.15 X 10-8 o 2.43 x 10--4 2.72 x 10-7 Note changes in ordinate magnitude
e 8.09 x 10-3
0:5
n 5 6 7 6
Rate Dependence on Reactants, Products, and Catalyst. The effects of Ce(IV), Ce(III), As(III), As(V), and I- on the initial reaction rate are presented in this section. CERIUM(IV). The Ce(1V) dependency is represented in Figures 2 and 3, in which reaction rate is plotted as a function of [Ce(IV)]/[As(III)]. Attention should be focused initially on the data represented by the solid lines on the figures. It is observed that at the higher absolute concentrations of As(II1) the reaction is first order in Ce(1V) up to [Ce(IV)]/ [As(III)] ratios as large as 5. On the other hand, at lower As(II1) concentrations the Ce(1V) dependency curves fall off rapidly to approach zero order at [Ce(IV)I/[As(III)] ratios’ as small as 5 . The dashed curves on these figures represent the predicted and observed rates when the As(II1) concentration is low and the iodide concentration is increased. The observed negative deviations presumably result from the conversion of a portion of the iodide to a catalytically inactive species, a phenomenon not accounted for in the reported rate expression. C E R ~ U(111). M The Ce(II1) dependency was evaluated at [Ce(IV)] = 7.5 X 10-4M, [As(III)] = 8.1 X 10-4M, and iodide at several concentrations up to 1.36 X 10-7M. No change in rate (within a 2x experimental error) was observed for [Ce(III)], [Ce(IV)] ratios up to 15. It is concluded there is no Ce(II1) dependency. ARSENIC(111). The As(II1) dependency is represented in Figures 4 and 5 in which initial rate is plotted as a function of the [As(TII)]/[Ce(IV)]ratio for three different Ce(1V) concentrations and a single iodide concentration. It is observed that at low [As(III)I~[Ce(IV)]ratios, the curves show first order dependence on As(II1). However, as the concentration ratio increases, the As(II1) dependence falls off rapidly to zero. The [As(III)I/[Ce(IV)I ratio at which the As(II1) dependence approaches zero order is a function of the absolute As(II1) concentration. ARSENIC(VI. Rate dependency on recrystallized As(V) from several sources was evaluated. In all cases, a decrease in rate was observed which was a linear function of the amount of As(V) added. Working curves on rate rs. iodide concentra1372
ANALYTICAL CHEMISTRY
1:o
1:5
Ratio C A S m3 I CCe m3
Figure 4. Rate dependency on As(I1I) at small ratios of [As(III)I to [Ce(IV)] [I-] = 8.15 X 10-8 [Ce(IV)l e 7.50 x 10-3 + 7.50 x 10-4 Q 7.50 x 10-5 Note changes in ordinate magnitude
n 6 7 8
tion prepared at different As(V) concentrations deviated from linearity. Intercepts along the iodide concentration axis were observed, the magnitude of the intercept was a function of As(V) concentration. On the basis of these experiments it was impossible to determine with certainty whether the decreabe in rate resulted from the As(V) added or some inhibitor (Ag+ or Hg2+)added with the As(V). consequently, an alternate experiment was performed. Stoichiometric amounts of Ce(1V) and As(II1) which were known to be free of impurities (based upon rate measurements) were permitted to react to completion in a sealed ampule generating Ce(II1)
I
0
5
10 Ratio C A s m 3 I CCe m 3
20
15
Figure 5. Rate dependency on As(II1) at large ratios of [As (III)] to [Ce(IV)] [I-] = 8.15 X lo-* tCe(W1 e 7.50 x 10-3 + 7.50 x 10-4 Q 7.50 x 10-5 Note changes in ordinate magnitude
n
6 7 8
a Y
4 8 12 16 20 24 28 Iodide concentration (Moles/ Liter) X IO8
Figure 6. Rate dependency on iodide as a function of [As(III)] at constant [Ce(IV)] - Predicted rates Measured rates [Ce(IV)] = 7.5 X 10-4M [As(III)] = A- 1.62 X 10-3M, B- 8.09 X 10-4M, C- 2.43 X 10-4~
---
'1.08 2.k 3.24 I o d i d e concentration ( M o l e s / L ) x l d
Figure 7.
and As(V). As Ce(II1) had already been shown to have no effect on the reaction (for [Ce(III)]/[Ce(IV)] 15), this solution was used to evaluate the As(V) effect. Conditions were adjusted so that the [Ce(III)]/[Ce(IV)] ratio was always less than 15. A slight As(V) dependence was observed. For example, for [As(III)I = 8.09 X 10-4M, [Ce(IV)I = 7.5 X 10-4M, [I-] = 1.36 X lO-'M, and an [As(V)]/[As(III)] ratio of 10.3, the reaction rate was decreased by about 10.4x from the value it would have if no As(V) were present. It is clear that this effect can be ignored when no As(V) is present initially and initial rates are measured. The decrease is well below that observed when recrystallized As(V) preparations were used. Therefore, it was concluded that these preparations contained inhibiting contaminants which were not removed by repeated recrystallization. IODIDE. The iodide dependency at a single Ce(1V) concentration and several different As(II1) concentrations is represented in Figure 6. It is observed that at the highest As(II1) concentration, the reaction rate is a linear function of iodide concentration throughout the range examined. It should be noted further that at several higher As(II1) concentrations (up to 1.6 X 10-2M) results obtained are identical to those at [As(III)] = 1.62 X 10-3M. However, as the As (111) concentration is decreased, the iodide dependence falls from first to fractional order approaching zero order as the iodide concentration increases above 2.8 X 10-7M for [As (111)] = 2.43 X 10-4M. The arrow on the upper curve represents the iodide concentration at which the data in Figures 4 and 5 were obtained. The rate at this point (7.4 X 10-7 moles/liter-second) is a good agreement with the maximum rates (7.3 X 10' moles/liter-second) observed from Figure 5 for [Ce(IV)] = 7.5 X lO-'M. This upper curve represents the maximum rate for each iodide concentration for the stated value of [Ce(IV)I. Rate Dependence on Diverse Species and Temperature. Effects of Chloride, HC104, HzS04, and temperature are presented in this section. CHLORIDE.Results obtained for the effect of chloride on the reaction rate are presented in Figure 7. These data represent the reaction rate as a function of iodide concentration in the absence of chloride and in the presence of a 4 X 105-
Effect of chloride on iodide dependency [Ce(IV)] = 8.45 X 10-4M [As(III)] = 1.30 X lO-3M [H*SOa] = 0.18M [HCI] = 0.125M [Cl-1 = O.OOM
