Kinetics of the Periodate Reaction with Some Vicinal Glycols Studied with a Perchlorate Ion Selective Electrode as a Periodate Ion Sensor C. E. Efstathiou and T. P. Had]iioannou* Laboratory of Analytical Chemistry, University of Athens, Athens, Greece
E. McNelis Department of Chemistry, New York University, New York, N. Y. 10003
The perchlorate Ion selective electrode has been used as a periodate sensor. The electrode gave near-Nernstian response M perlodate, wlth a slope of 62-58 at pH 4.0 from 10-1 to depending on the electrode age. iodate, the usual product in periodate organic reactions does not Interfere with periodate determination. Response data prove that the perchlorate electrode can be used for continuous monitoring of periodate reactions. Several perlodate-vicinal glycol reactions have been studied and equilibrlum and rate constants have been determined. Kinetic data obtained wlth the electrode compare favorably wlth those obtalned by tedious sampling techniques.
The kinetic study of reactions of periodate with organic molecules is of great importance because of the selectivity which characterizes periodate oxidations. Information concerning the structure of the oxidized organic molecule can often be deduced from such kinetic data ( I ) . Hitherto, most kinetic studies of periodate reaction have been carried out using the tedious “sampling” technique in which aliquots from the reacting mixture are taken at various time intervals, quenched, and analyzed for the unreacted periodate. This technique is not applicable to fast reactions. Therefore, the relatively fast reactions of periodate with certain glycols must be studied a t low temperatures (around 0 “C). Continuous monitoring of periodate in its reactions can be done spectrophotometrically since periodate absorbs strongly a t 222.5 nm, but the technique is not always applicable because high concentration levels of organic compounds contribute strongly to the absorbance and tedious corrections have to be made. Therefore, a simple device for monitoring periodate reactions would be a useful tool for the chemist. Recently we reported that a perchlorate ion selective electrode responds quantitatively to periodate and described the application of the electrode to the potentiometric determination of vicinal glycols (2).In the work now described we have made a more systematic study of the properties of this electrode in order to use i t as a continuous monitor for reactions involving periodate ion and employed it as a means of determining the equilibrium and rate constants of the reactions of periodate with several vicinal glycols. For models we have used some of the reactions for which Duke et al. ( 3 , 4 )have reported kinetic data and have attained comparable results. In addition, the study has been extended to include glycols not previously reported.
GENERAL CONSIDERATIONS Electrode Storage, Conditioning, and Deterioration Symptoms. The ion exchanger of the perchlorate selective electrode consists of a solution of the salt [Fe(bathophenan410
ANALYTICAL CHEMISTRY, VOL. 49, NO. 3, MARCH 1977
throline)~](C104)2 in 2-nitro-p-cymene. When‘not in use, the electrode is kept in dilute NaC104 solution (-0.01M) and not in NaI04 solution because the electrode performance deteriorates when stored in such a solution. This may be attributed to the slow oxidation of the ferrous complex by periodate and also to the oxidation of the electrode porous membrane. For this reason the electrode should not be used for prolonged periods of time in periodate solutions more concentrated than -0.005 M, especially at elevated temperatures. Before measurements are started, the electrode is immersed for about 10 min in a stirred NaI04 solution (-O.OOlM), in order to take the periodate form and thus avoid small potential drifts. The electrode should be refilled and the membrane replaced if one of the following symptoms is observed: 1)sluggish response; 2) sudden potential jumps of several millivolts; 3) extensive liquid ion-exchanger leakage (in this case the electrode “0”ring should also be checked). Under normal electrode use (working in dilute periodate solutions in the p H range 4 to 8 at or below room temperature) the electrode life is about 40 days. pH Dependence. The electrode potential is practically pH independent in the p H range between 3.8 and 6.5 (Figure 1) and is given by Equation 1:
.RT
E = E’ - -In
ax- = E’
- S log a,-
ZXF
(1)
where E is the measured total potential of the system, E’ is the portion of the total potential due to choice of reference electrodes and internal solution, R and F are the ideal gas and Faraday constants, respectively, T is the absolute temperature, ax- is the activity of the active periodate species x -, z, is the charge of the ion x-, and S is the slope of the linear response function E vs. log a,-. The prelogarithmic factor S of Equation 1 must be accurately known when kinetic studies and determinations of equilibrium and rate constants are conducted with selective ion electrodes. Unfortunately, in ion selective electrodes, S is frequently not exactly Nernstian so that a Slope Normalization Factor, a , which is experimentally determined is introduced to compensate for any deviation from Nernstian response ( a = experimental slope/theoretical slope). Thus, taking into account that 2,- = 1and a x - = [IO4-] in the p H rangb3.8-6.5 Equation 1 becomes
E
=
RT E’ - a -In [IO4-] F
Figure 1shows the perchlorate electrode potential behavior vs. solution pH for periodate solutions in the range 10-4-10-2 M. This peculiar electrode behavior is explained by recalling periodate species equilibria as they have been described by Crouthamel et al. ( 5 ) . According to these equilibria, 1 0 4 -
(metaperiodate) species concentration is p H dependent and the actual concentration of IO4- is given by Equation 3
where K 1 and K2 are the true paraperiodic acid ionization constants and K D is the equilibrium constant for the dehydration of H4IOs- species t o 1 0 4 - ( K 1 = 5.1 X lo-*, Kz = 2.0 X and KD= 40). [PI represents the sum of all periodate species existing in solution. Substituting the 1 0 4 - expression from Equation 3 in Equation 1 we obtain potential vs. p H profiles which are in good agreement with experimental data (Figure 1).Negative deviations at high p H between experimental data and theoretical predictions may be attributed to the effect of p H on the slope normalization factor and to the hydroxide interference which becomes more serious at lower P values. The electrode potential becomes unstable for periodate solutions a t p H lower than 3.5. The electrode may be used a t p H up to 8 provided the solution is properly buffered. The perchlorate electrode cannot function as a periodate monitor a t p H higher than 8 because in such solutions the range of linear response decreases and large deviations from the theoretical slope occur. Also, at high p H values the electrode life is shortened. Vicinal Glycols-Periodate Reaction Kinetics. The oxidation of vicinal glycols by periodate is known to proceed via a rapid reversible formation of a singly-negatively charged intermediate between the glycol and periodate, followed by a slow decomposition of the intermediate to the final products (3).The general reaction scheme is; G
K
+ 1 0 4 - +G
k
IO4-
103-
+ carbonyl compounds
.........theorDtical
t
150
Figure 1. Effect of pH on the perchlorate electrode potential behavior on pure
-d[PI - - W G I [ P l dt
1
+ K[G]
(5)
where [GI is the concentration of the uncoordinated glycol. From Equation 5 we obtain (2) dln[IOs-] = - kK[Gl dt 1 K[G]
+
(6)
Differentiating Equation 2 with respect to time and combining with Equation 6, dE _ - a . - .RT dt
F
kK[G] 1 K[G]
+
(7)
If equilibrium has been established and glycol is always kept in large excess over periodate so that the term [GI remains practically constant, and if the periodate concentration remains within the linear response range of the perchlorate electrode, then the potential should vary linearly with time. Therefore, under such conditions the term dE/dt can be substituted by the term AEElAt, which is the actual measured quantity, and Equation 7 becomes
A E-- a . - .RT _ At
Rearranging Equation 8,
F
hK[G] 1 K[G]
+
Na104 solutions
-.-.-= a RT
At
hKIG1
l+K[G]
=Q
(9)
where Q is an experimentally determinable quantity. Equation 9 can be written as 1 1 _1 ---+-.-
1
Q
h k K [GI Thus, if we plot 1/Qvs, 1/[G], the constants h and K can be determined from the intercept and the slope of the curve. These constants have been determined for four vicinal glycols, ethanediol, propane-l,2-diol, 3-chloropropane-1,2-diol and hexane-1,2,6-triol.
(4) where G is a vicinal glycol, G.104- is the intermediate, and K and k are the equilibrium and rate constants, respectively. Duke et al. ( 3 , 4 )have shown that if glycol is present in large excess over periodate and the p H remains constant in the range 4-7, the kinetics of the reaction are expressed by Equation 5:
response
,,
-experimental
EXPERIMENTAL Instrumentation. Electrodes. An Orion perchlorate ion selective electrode, Model 92-81, was used as an indicator electrode, in conjunction with a double-junction silver-silver chloride electrode, Orion Model 90-02-00, as the reference electrode, as previously reported (2).
Reaction Cell. A thermostated double-walled 50-ml beaker is used as reaction cell. An Ultra-Kryomat, Model TK-SOD, was used for low temperatures. The reaction mixtures were stirred with a Tefloncoated magnetic bar. Recording System. A pH/pIon electrometer (Heath-Schlumberger Model EU-200-30) is used as a follower amplifier for impedance matching. A potentiometric recorder (Heath-Schlumberger Model EU-205 B) is used to mon'itor the output from the electrometer and record the reaction curve. The recorder span is always kept a t 50 mV. The chart speed is adjusted according to the reaction rate. In the present study the following conditions were used: Ethanediol and 3-chloro-propane-l,2,-diol: chart speed = y5 cm/s; concentration: 0.050,0.100,0.50, 0.200, and 0.250 M. Propane-1,2-diol: chart speed = y3 cm/s; concentration: 0.030, 0.060, 0.090, 0.120, and 0.150 M. Hexane-1,2,6-triol: chart speed = y3 cm/s; concentration: 0.020,0.040, 0.060,0.080, and 0.100 M. Reagents. All solutions were prepared with deionized single-distilled water from reagent-grade substances, except where indicated otherwise. S o d i u m Sulfate 0.20 M was used. Acetate Bufffer, p H 4.0 was prepared by adding 5 M NaOH solution to a 0.2 M CH3COOH solution to pH 4.0. S o d i u m Metaperiodate. (a) Stock solution I, 0.0500 M, 10.70 g of NaI04 (G. F. Smith Co., Columbus, Ohio) are dissolved in water and diluted to 11. The solution is kept in amber bottles, (b) Composite (working) solution 11: 5.00 ml of solution I, 184 ml of the sodium sulfate solution, and 25 ml of the buffer solution are diluted with water to 1 1. This solution should be prepared fresh when needed and kept in an amber bottle. Glycols Stock Solutions, 0.5 M / The glycols used were: Ethanediol (Merck, p.a.1, propane-1,2-diol (Fluka, puriss), 3-chloropropane1,2-diol (Fluka, purum), and hexane-1,2,6-triol (Merck-Schuchardt, prosynthesi). The last two diols were tested for impurities producing ANALYTICAL CHEMISTRY, VOL. 49,
NO. 3,
MARCH 1977
411
105
20-
7:
> E
100
.1.5r
e a
20 -
w 0
1
100
200 1/[Gl
. M-'
300
I
Figure 5. Plot of 1/Q vs. l/[G]for propane-1,2-diol at various temperatures
tFigure 3. Recorded curves of cell voltage vs. time for propane-1,2diol-periodate reaction at various temperatures, G: a = 0.030 M, b = 0.060 M, c = 0.090 M, d = 0.120 M, e = 0.150 M = injection of
(4
glycol)
formic acid upon periodate oxidation (e.g., glycerol, polyols, etc.) and found to contain less than 1%on a molar basis. The exact titer was estimated according to the Fleury-Lange method ( I ) , except for 3chloropropane-1,2-diolfor which the method fails. In this case the titer was estimated from the weighed amount of 3-chloropropane-1,2-diol and it was accurate to within fl% (such accuracy is sufficient for the present study). Working solutions of glycols were prepared by appropriate dilution. Procedure. Pipet 20.00 ml of the periodate composite solution into the reaction cell, immerse the electrodes and a thermometer carefully into the solution and sart stirring at the maximum allowable speed so that air bubbles will not be formed. When the temperature is stabilized at the desired value start the recorder, pipet 2.00 ml of the glycol solution quickly into the cell, and record a large portion of the linear part of the reaction curve or until the recording becomes noisy.
RESULTS AND DISCUSSION The perchlorate ion selective electrode exhibits Nernstian to 10-1 M. response for periodate a t concentrations of However, the electrode should not be used for periodate soM, because of possible lutions more concentrated than membrane destruction at higher periodate concentrations. Since iodate is produced during the periodate-vicinal glycol reaction, the selectivity constant K104-,103-was determined using the mixed solution technique (6),and it was found equal to 8 X Thus, iodate interference is negligible and periodate may be determined by direct potentiometry in the presence of thousand-fold excess of iodate. 412
ANALYTICAL CHEMISTRY, VQL. 49, NO. 3, MARCH 1977
The slope normalization factor a was found to depend on storage conditions, frequency of electrode utilization, chemical environment during measurements, pH, and electrode age. The effect of p H and electrode age on a i's shown in Figure 2. Values of a were determined from the slopes of working curves obtained a t various pH values with an electrode of known age. The effect of the aforementioned factors on a necessitates the determination of a in kinetic studies at the particular chemical environment before, during, or soon after a series of experiments. It was found that a is temperature independent in the range 5 to 35 "C. To check the dynamic response of the electrode, its potential was recorded during dilution of periodate solutions by rapidly injecting water under vigorous stirring. It was found that the E M F comes to within 0.2 mV of the steady-state in less than 2 s (including the time for mixing) when decreasing the periodate concentration by 20% in the range 10-4-10-3 M, a t pH 4 to 7, in the temperature range 5 to 25 "C, for solution having initial ionic strength in the range 0.01 to 1.00. At l o p 2M periodate, the response of the electrode is slightly slower, whereas in lov5 M periodate, the response is noisy. From the above data we conclude that the perchlorate electrode can be used for continuous monitoring of reactions involving periodate. Recorded curves for the propane-1,2-diol-periodatereaction a t various temperatures are shown in Figure 3. Similar curves were obtained with the other three glycols tested. It can be seen that the initial fast potential increase corresponding to the rapid reversible formation of a singly-negatively charged intermediate between the glycol and periodate, G.I04-, is more pronounced a t lower temperatures. This is due to the fact that the equilibrium constant K decreases with increasing temperature.
Table I. Equilibrium and Rate Constants and Thermodynamic Parameters for the Reactions:
G Vicinal glycol, G Ethanediol
+ 104-
K
'k
Temperature,
+ 103-
(at pH 4.1-4.3,
K fSK, (M-l)
k fsh,
14 17 20 23
47.7 f 1.4 34.5 f 0.3 25.0 f 0.3 19.2 f 1.0
0.034 f 0.001 0.056 f 0.004 0.088 f 0.001 0.130 f 0.007
14 20 23
81.2 f 1.7 53.6 f 3.9 40.9 f 1.2 30.5 f 1.0
0.092 f 0.002 0.163 f 0.011 0.225 f 0.006 0.334 f 0.010
14 17 20 23
25.6 f 1.7 21.9 & 0.7 15.9 f 0.3 13.1 f 0.7
0.030 f 0.002 0.043 f 0.001 0.069 f 0.001 0.099 f 0.005
14
75 f 6
17
55 f 2 32 f 5
0.17 f 0.02 0.24 f 0.02 0.45 f 0.07 0.65 f 0.05
fO.l O
C
17
3-Chloropropane-1,2-diol
-
3 G 1 0 4 - -aldehydes
20 23
fS A H ,
(s-l)
23 f 3
= 0.1) E a c t f SE,,~,
(Kcal/mol)
(KcaVmol)
-17.2 f 0.5
25.2 f 0.7
-18.1 f 1.2
24.0 f 1.0
-13.1 f 1.2
22.6 f 0.9
-23.0 f 2.0
25.8 f 2.1
Table 11. Comparative Values of K and k Reported values (ref. 4 ) (pH 4-7, F = 0.2) Glycol
Calculated values (pH 4.1-4.2,~= 0.1)
Temperature, " C
K
12
K
k
Ethanediol
0.00 5.25
193 f 10 124 f 10
0.0046 f 0.0001 0.0102 f 0.0002
221 122
0.0036 0.0087
Propane-1,2-diol
0.00 5.25
350 f 20 215 f 15
0.0137 f 0.0005 0.0308 f 0.0020
398
0.0110
212
0.0253
In Figure 4 the dependence of reaction rate on glycol concentration is shown for the four glycols tested. I t can be seen that the reaction rate is larger for propane-1,2-diol than for ethanediol, in agreement with data in the literature ( 4 , 7). Negative deviations from linearity are due to the fact that a t high glycol concentrations the term k [GI cannot be omitted in the denominator of Equation 8. The mathematical manipulations leading to Equation 10 simplify the handling of [GI. This is illustrated in Figure 5 by a plot of l/Qvs, 1/[G] for propane-1,2-diol and linearity of the experimental curves confirms the validity of the reaction scheme. Similar curves were also obtained for the other glycols. Calculated values for the equilibrium constant K and the rate constant k are given in Table I. The slopes and intercepts as well as the standard deviations for the constants were obtained using the least squares method (8). From these data heat of formathe thermodynamic parameters AH and EaCt, tion and activation energy for the intermediate respectively, were obtained using standard plotting techniques. Also, from the same data the K and k values for ethanediol and propane-1,2-diol a t 0 OC are obtained by extrapolation and are compared with those reported in the literature ( 4 ) in Table 11. There is satisfactory agreement between results obtained by the tedious sampling technique ( 4 ) and those obtained using the perchlorate electrode as periodate sensor. From the described evaluation of the perchlorate ion selective electrode, it is concluded that this electrode is a reliable sensor for monitoring and studying periodate reactions and obtaining kinetic data. The logarithmic response of this transducer is advantageous because it greatly simplifies cal-
culations. Other periodate reactions can also be studied using the perchlorate electrode. For example, the electrode has been successfully used (9) in trace analysis by kinetic methods and by catalytic titration since many periodate reactions are catalyzed or promoted by various ions and such catalyzed reactions are inhibited by aminopolycarboxylic acids. ACKNOWLEDGMENT The authors are grateful to J. D. R. Thomas, UWIST (U.K.) and H. Freiser, University of Arizona, Tucson, Ariz., for helpful discussions made possible by NATO Research Grant No 1000, and to H. Pardue, Purdue University, Lafayette, Ind., for valuable suggestions. LITERATURE CITED G. Dryhurst, "Periodate Oxidation of Diol and Other Functional Groups",
Pergamon Press, London, 1970. C. Efstathiou and T. P. Hadjiioannou, Anal. Chem., 47, 864 (1975). F. R. Duke, J. Am. Chem. Soc., 69, 3054 (1947). F. R Duke and V. C. Bulgrin, J. Am. Chem. Soc., 76, 3803 (1954). C. E.Crouthamel,A. M. Hayes, and D. S.Martin, J. Am. Chem. Soc., 73, 82 (1951). G. J. Moody and J. D. R. Thomas, "Selective Ion Sensitive Electrodes", Merrow Publishing Co., Watford, Herts, England, 1971, p 14. G . J. Buist, C. A. Bunton, and J. H. Miles, J. Chem. Soc., 4567 (1957). W. J. Youden, "Statistical Methods for Chemists", John Wiiey & Sons, Inc., New York, 5th printing, 1961, p 40. T. P. Hadjiioannou, M. A. Koupparis, and C. E. Efstathiou, Anal. Chim. Acta, in press.
RECEIVEDfor review August 10, 1976. Accepted November 17, 1976. This research was supported in part by a research grant from the Greek National Institute of Research. ANALYTICAL CHEMISTRY, VOL. 49, NO. 3, MARCH 1977
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